,transcriber's notes3 ,in the chemical equations1 superscripts are indicated with a and subscripts are indicated with a 4 ,the affected item is enclosed in curly brackets 4 ,examples are ,h for hydrogen ion and ,h#b,o for water4 ,since the underscore is already being used in this project1 italics are designated by an exclamation point before and after the italicized word or phrase4 ,,an ,,introductory ,,course ,,of ,,quantitative ,,chemical ,,analysis ,,with ,,explanatory ,,notes ,,by ,,henry ,p4 ,,talbot ,,professor ,,of ,,inorganic ,,chemistry ,,at ,,the ,,massachusetts ,,institute ,,of ,,technology ,,sixth ,,edition1 ,,completely ,,rewritten ,,preface ,this ,introductory ,course of ,quantitative ,analysis has been prepared to meet the needs of students who are just entering upon the subject1 after a course of qualitative analysis4 ,it is primarily intended to enable the student to work successfully and intelligently without the necessity for a larger measure of personal assistance and supervision than can reasonably be given to each member of a large class4 ,to this end the directions are given in such detail that there is very little opportunity for the student to go astray2 but the manual is not1 the author believes1 on this account less adapted for use with small classes1 where the instructor1 by greater personal influence1 can stimulate independent thought on the part of the pupil4 ,the method of presentation of the subject is that suggested by ,professor ,a4,a4 ,noyes' excellent manual of ,qualitative ,analysis4 ,for each analysis the procedure is given in considerable detail1 and this is accompanied by explanatory notes1 which are believed to be sufficiently expanded to enable the student to understand fully the underlying reason for each step prescribed4 ,the use of the book should1 nevertheless1 be supplemented by classroom instruction1 mainly of the character of recitations1 and the student should be taught to consult larger works4 ,the general directions are intended to emphasize those matters upon which the beginner in quantitative analysis must bestow special care1 and to offer helpful suggestions4 ,the student can hardly be expected to appreciate the force of all the statements contained in these directions1 or1 indeed1 to retain them all in the memory after a single reading2 but the instructor1 by frequent reference to special paragraphs1 as suitable occasion presents itself1 can soon render them familiar to the student4 ,the analyses selected for practice are those comprised in the first course of quantitative analysis at the ,massachusetts ,institute of ,technology1 and have been chosen1 after an experience of years1 as affording the best preparation for more advanced work1 and as satisfactory types of gravimetric and volumetric methods4 ,from the latter point of view1 they also seem to furnish the best insight into quantitative analysis for those students who can devote but a limited time to the subject1 and who may never extend their study beyond the field covered by this manual4 ,the author has had opportunity to test the efficiency of the course for use with such students1 and has found the results satisfactory4 ,in place of the usual custom of selecting simple salts as material for preliminary practice1 it has been found advantageous to substitute1 in most instances1 approximately pure samples of appropriate minerals or industrial products4 ,the difficulties are not greatly enhanced1 while the student gains in practical experience4 ,the analytical procedures described in the following pages have been selected chiefly with reference to their usefulness in teaching the subject1 and with the purpose of affording as wide a variety of processes as is practicable within an introductory course of this character4 ,the scope of the manual precludes any extended attempt to indicate alternative procedures1 except through general references to larger works on analytical chemistry4 ,the author is indebted to the standard works for many suggestions for which it is impracticable to make specific acknowledgment2 no considerable credit is claimed by him for originality of procedure4 ,for many years1 as a matter of convenience1 the classes for which this text was originally prepared were divided1 one part beginning with gravimetric processes and the other with volumetric analyses4 ,after a careful review of the experience thus gained the conclusion has been reached that volumetric analysis offers the better approach to the subject4 ,accordingly the arrangement of the present 7the sixth7 edition of this manual has been changed to introduce volumetric procedures first4 ,teachers who are familiar with earlier editions will1 however1 find that the order of presentation of the material under the various divisions is nearly the same as that previously followed1 and those who may still prefer to begin the course of instruction with gravimetric processes will1 it is believed1 be able to follow that order without difficulty4 ,procedures for the determination of sulphur in insoluble sulphates1 for the determination of copper in copper ores by iodometric methods1 for the determination of iron by permanganate in hydrochloric acid solutions1 and for the standardization of potassium permanganate solutions using sodium oxalate as a standard1 and of thiosulphate solutions using copper as a standard1 have been added4 ,the determination of silica in silicates decomposable by acids1 as a separate procedure1 has been omitted4 ,the explanatory notes have been rearranged to bring them into closer association with the procedures to which they relate4 ,the number of problems has been considerably increased4 ,the author wishes to renew his expressions of appreciation of the kindly reception accorded the earlier editions of this manual4 ,he has received helpful suggestions from so many of his colleagues within the ,institute1 and friends elsewhere1 that his sense of obligation must be expressed to them collectively4 ,he is under special obligations to ,professor ,l4,f4 ,hamilton for assistance in the preparation of the present edition4 ,,henry ,p4 ,,talbot 6,massachusetts ,institute of ,technology1 ,september1 #aiba64 ,,contents ,,part ,i4 ,,introduction ,,subdivisions ,,of ,,analytical ,,chemistry ,,general ,,directions ,accuracy and ,economy of ,time2 ,notebooks2 ,reagents2 ,wash-bottles2 ,transfer of ,liquids ,,part ,,ii4 ,,volumetric ,,analysis ,,general ,,discussion ,subdivisions2 ,the ,analytical ,balance2 ,weights2 ,burettes2 ,calibration of ,measuring ,devices ,,general ,,directions ,standard and ,normal ,solutions 6,i4 ,neutralization ,methods6 ,,alkalimetry ,,and ,,acidimetry ,preparation and ,standardization of ,solutions2 ,indicators ,,standardization ,,of ,,hydrochloric ,,acid ,,determination ,,of ,,total ,,alkaline ,,strength ,,of ,,soda ,,ash ,,determination ,,of ,,acid ,,strength ,,of ,,oxalic ,,acid 6,,ii4 ,oxidation ,processes6 ,,general ,,discussion ,,bichromate ,,process ,,for ,,the ,,determination ,,of ,,iron ,,determination ,,of ,,iron ,,in ,,limonite ,,by ,,the ,,bichromate ,,process ,,determination ,,of ,,chromium ,,in ,,chrome ,,iron ,,ore ,,permanganate ,,process ,,for ,,the ,,determination ,,of ,,iron ,,determination ,,of ,,iron ,,in ,,limonite ,,by ,,the ,,permanganate ,,process ,,determination ,,of ,,iron ,,in ,,limonite ,,by ,,the ,,zimmermann-,,reinhardt ,,process ,,determination ,,of ,,the ,,oxidizing ,,power ,,of ,,pyrolusite ,,iodimetry ,,determination ,,of ,,copper ,,in ,,ores ,,determination ,,of ,,antimony ,,in ,,stibnite ,,chlorimetry ,,determination ,,of ,,available ,,chlorine ,,in ,,bleaching ,,powder 6,,iii4 ,precipitation ,methods6 ,,determination ,,of ,,silver ,,by ,,the ,,thiocyanate ,,process ,,part ,,iii4 ,,gravimetric ,,analysis ,,general ,,directions ,precipitation2 ,funnels and ,filters2 ,filtration and ,washing of ,precipitates2 ,desiccators2 ,crucibles and their ,preparation for ,use2 ,ignition of ,precipitates ,,determination ,,of ,,chlorine ,,in ,,sodium ,,chloride ,,determination ,,of ,,iron ,,and ,,of ,,sulphur ,,in ,,ferrous ,,ammonium ,,sulphate ,,determination ,,of ,,sulphur ,,in ,,barium ,,sulphate ,,determination ,,of ,,phosphoric ,,anhydride ,,in ,,apatite ,,analysis ,,of ,,limestone ,determination of ,moisture2 ,insoluble ,matter and ,silica2 ,ferric ,oxide and ,alumina2 ,calcium2 ,magnesium2 ,carbon ,dioxide ,,analysis ,,of ,,brass ,electrolytic ,separations2 ,determination of ,lead1 ,copper1 ,iron and ,zinc4 ,,determination ,,of ,,silica ,,in ,,silicates ,,part ,,iv4 ,,stoichiometry ,,solutions ,,of ,,typical ,,problems ,,problems ,,appendix ,,electrolytic ,,dissociation ,,theory ,,folding ,,of ,a ,,filter ,,paper ,,sample ,,notebook ,,pages ,,strength ,,of ,,reagents ,,densities ,,and ,,volumes ,,of ,,water ,,corrections ,,for ,,change ,,of ,,temperature ,,of ,,standard ,,solutions ,,atomic ,,weights ,,logarithm ,,tables ,,quantitative ,,chemical ,,analysis ,,part ,i ,,introduction ,,subdivisions ,,of ,,analytical ,,chemistry ,a complete chemical analysis of a body of unknown composition involves the recognition of its component parts by the methods of 6qualitative analysis61 and the determination of the proportions in which these components are present by the processes of 6quantitative analysis64 ,a preliminary qualitative examination is generally indispensable1 if intelligent and proper provisions are to be made for the separation of the various constituents under such conditions as will insure accurate quantitative estimations4 ,it is assumed that the operations of qualitative analysis are familiar to the student1 who will find that the reactions made use of in quantitative processes are frequently the same as those employed in qualitative analyses with respect to both precipitation and systematic separation from interfering substances2 but it should be noted that the conditions must now be regulated with greater care1 and in such a manner as to insure the most complete separation possible4 ,for example1 in the qualitative detection of sulphates by precipitation as barium sulphate from acid solution it is not necessary1 in most instances1 to take into account the solubility of the sulphate in hydrochloric acid1 while in the quantitative determination of sulphates by this reaction this solubility becomes an important consideration4 ,the operations of qualitative analysis are1 therefore1 the more accurate the nearer they are made to conform to quantitative conditions4 ,the methods of quantitative analysis are subdivided1 according to their nature1 into those of 6gravimetric analysis1 volumetric analysis61 and 6colorimetric analysis64 ,in 6gravimetric6 processes the constituent to be determined is sometimes isolated in elementary form1 but more commonly in the form of some compound possessing a well-established and definite composition1 which can be readily and completely separated1 and weighed either directly or after ignition4 ,from the weight of this substance and its known composition1 the amount of the constituent in question is determined4 ,in 6volumetric6 analysis1 instead of the final weighing of a definite body1 a well-defined reaction is caused to take place1 wherein the reagent is added from an apparatus so designed that the volume of the solution employed to complete the reaction can be accurately measured4 ,the strength of this solution 7and hence its value for the reaction in question7 is accurately known1 and the volume employed serves1 therefore1 as a measure of the substance acted upon4 ,an example will make clear the distinction between these two types of analysis4 ,the percentage of chlorine in a sample of sodium chloride may be determined by dissolving a weighed amount of the chloride in water and precipitating the chloride ions as silver chloride1 which is then separated by filtration1 ignited1 and weighed 7a 6gravimetric6 process72 or the sodium chloride may be dissolved in water1 and a solution of silver nitrate containing an accurately known amount of the silver salt in each cubic centimeter may be cautiously added from a measuring device called a burette until precipitation is complete1 when the amount of chlorine may be calculated from the number of cubic centimeters of the silver nitrate solution involved in the reaction4 ,this is a 6volumetric6 process1 and is equivalent to weighing without the use of a balance4 ,volumetric methods are generally more rapid1 require less apparatus1 and are frequently capable of greater accuracy than gravimetric methods4 ,they are particularly useful when many determinations of the same sort are required4 ,in 6colorimetric6 analyses the substance to be determined is converted into some compound which imparts to its solutions a distinct color1 the intensity of which must vary in direct proportion to the amount of the compound in the solution4 ,such solutions are compared with respect to depth of color with standard solutions containing known amounts of the colored compound1 or of other similar color-producing substance which has been found acceptable as a color standard4 ,colorimetric methods are1 in general1 restricted to the determinations of very small quantities1 since only in dilute solutions are accurate comparisons of color possible4 ,,general ,,directions ,the following paragraphs should be read carefully and thoughtfully4 ,a prime essential for success as an analyst is attention to details and the avoidance of all conditions which could destroy1 or even lessen1 confidence in the analyses when completed4 ,the suggestions here given are the outcome of much experience1 and their adoption will tend to insure permanently work of a high grade1 while neglect of them will often lead to disappointment and loss of time4 ,,accuracy ,,and ,,economy ,,of ,,time ,the fundamental conception of quantitative analysis implies a necessity for all possible care in guarding against loss of material or the introduction of foreign matter4 ,the laboratory desk1 and all apparatus1 should be scrupulously neat and clean at all times4 ,a sponge should always be ready at hand1 and desk and filter-stands should be kept dry and in good order4 ,funnels should never be allowed to drip upon the base of the stand4 ,glassware should always be wiped with a clean1 lintless towel just before use4 ,all filters and solutions should be covered to protect them from dust1 just as far as is practicable1 and every drop of solution or particle of precipitate must be regarded as invaluable for the success of the analysis4 ,an economical use of laboratory hours is best secured by acquiring a thorough knowledge of the character of the work to be done before undertaking it1 and then by so arranging the work that no time shall be wasted during the evaporation of liquids and like time-consuming operations4 ,to this end the student should read thoughtfully not only the 6entire6 procedure1 but the explanatory notes as well1 before any step is taken in the analysis4 ,the explanatory notes furnish1 in general1 the reasons for particular steps or precautions1 but they also occasionally contain details of manipulation not incorporated1 for various reasons1 in the procedure4 ,these notes follow the procedures at frequent intervals1 and the exact points to which they apply are indicated by references4 ,the student should realize that a 6failure to study the notes will inevitably lead to mistakes1 loss of time1 and an inadequate understanding of the subject64 ,all analyses should be made in duplicate1 and in general a close agreement of results should be expected4 ,it should1 however1 be remembered that a close concordance of results in check analyses is not conclusive evidence of the accuracy of those results1 although the probability of their accuracy is1 of course1 considerably enhanced4 ,the satisfaction in obtaining check results in such analyses must never be allowed to interfere with the critical examination of the procedure employed1 nor must they ever be regarded as in any measure a substitute for absolute truth and accuracy4 ,in this connection it must also be emphasized that only the operator himself can know the whole history of an analysis1 and only he can know whether his work is worthy of full confidence4 ,no work should be continued for a moment after such confidence is lost1 but should be resolutely discarded as soon as a cause for distrust is fully established4 ,the student should1 however1 determine to put forth his best efforts in each analysis2 it is well not to be too ready to condone failures and to begin again1 as much time is lost in these fruitless attempts4 ,nothing less than 6absolute integrity6 is or can be demanded of a quantitative analyst1 and any disregard of this principle1 however slight1 is as fatal to success as lack of chemical knowledge or inaptitude in manipulation can possibly be4 ,,notebooks ,notebooks should contain1 beside the record of observations1 descriptive notes4 ,all records of weights should be placed upon the right-hand page1 while that on the left is reserved for the notes1 calculations of factors1 or the amount of reagents required4 ,the neat and systematic arrangement of the records of analyses is of the first importance1 and is an evidence of careful work and an excellent credential4 ,of two notebooks in which the results may be1 in fact1 of equal value as legal evidence1 that one which is neatly arranged will carry with it greater weight4 ,all records should be dated1 and all observations should be recorded at once in the notebook4 ,the making of records upon loose paper is a practice to be deprecated1 as is also that of copying original entries into a second notebook4 ,the student should accustom himself to orderly entries at the time of observation4 ,several sample pages of systematic records are to be found in the ,appendix4 ,these are based upon experience2 but other arrangements1 if clear and orderly1 may prove equally serviceable4 ,the student is advised to follow the sample pages until he is in a position to plan out a system of his own4 ,,reagents ,the habit of carefully testing reagents1 including distilled water1 cannot be too early acquired or too constantly practiced2 for1 in spite of all reasonable precautionary measures1 inferior chemicals will occasionally find their way into the stock room1 or errors will be made in filling reagent bottles4 ,the student should remember that while there may be others who share the responsibility for the purity of materials in the laboratory of an institution1 the responsibility will later be one which he must individually assume4 ,the stoppers of reagent bottles should never be laid upon the desk1 unless upon a clean watch-glass or paper4 ,the neck and mouth of all such bottles should be kept scrupulously clean1 and care taken that no confusion of stoppers occurs4 ,,wash-,,bottles ,wash-bottles for distilled water should be made from flasks of about #gej cc4 capacity and be provided with gracefully bent tubes1 which should not be too long4 ,the jet should be connected with the tube entering the wash-bottle by a short piece of rubber tubing in such a way as to be flexible1 and should deliver a stream about one millimeter in diameter4 ,the neck of the flask may be wound with cord1 or covered with wash-leather1 for greater comfort when hot water is used4 ,it is well to provide several small wash-bottles for liquids other than distilled water1 which should invariably be clearly labeled4 ,,transfer ,,of ,,liquids ,liquids should never be transferred from one vessel to another1 nor to a filter1 without the aid of a stirring rod held firmly against the side or lip of the vessel4 ,when the vessel is provided with a lip it is not usually necessary to use other means to prevent the loss of liquid by running down the side2 whenever loss seems imminent a 6very thin6 layer of vaseline1 applied with the finger to the edge of the vessel1 will prevent it4 ,the stirring rod down which the liquid runs should never be drawn upward in such a way as to allow the solution to collect on the under side of the rim or lip of a vessel4 ,the number of transfers of liquids from one vessel to another during an analysis should be as small as possible to avoid the risk of slight losses4 ,each vessel must1 of course1 be completely washed to insure the transfer of all material2 but it should be remembered that this can be accomplished better by the use of successive small portions of wash-water 7perhaps #e-aj cc471 if each wash-water is allowed to drain away for a few seconds1 than by the addition of large amounts which unnecessarily increase the volume of the solutions1 causing loss of time in subsequent filtrations or evaporations4 ,all stirring rods employed in quantitative analyses should be rounded at the ends by holding them in the flame of a burner until they begin to soften4 ,if this is not done1 the rods will scratch the inner surface of beakers1 causing them to crack on subsequent heating4 ,,evaporation ,,of ,,liquids ,the greatest care must be taken to prevent loss of solutions during processes of evaporation1 either from too violent ebullition1 from evaporation to dryness and spattering1 or from the evolution of gas during the heating4 ,in general1 evaporation upon the steam bath is to be preferred to other methods on account of the impossibility of loss by spattering4 ,if the steam baths are well protected from dust1 solutions should be left without covers during evaporation2 but solutions which are boiled upon the hot plate1 or from which gases are escaping1 should invariably be covered4 ,in any case a watch-glass may be supported above the vessel by means of a glass triangle1 or other similar device1 and the danger of loss of material or contamination by dust thus be avoided4 ,it is obvious that evaporation is promoted by the use of vessels which admit of the exposure of a broad surface to the air4 ,liquids which contain suspended matter 7precipitates7 should always be cautiously heated1 since the presence of the solid matter is frequently the occasion of violent bumping1 with consequent risk to apparatus and analysis4 ,,part ,,ii ,,volumetric ,,analysis ,the processes of volumetric analysis are1 in general1 simpler than those of gravimetric analysis and accordingly serve best as an introduction to the practice of quantitative analysis4 ,for their execution there are required1 first1 an accurate balance with which to weigh the material for analysis2 second1 graduated instruments in which to measure the volume of the solutions employed2 third1 standard solutions1 that is1 solutions the value of which is accurately known2 and fourth1 indicators1 which will furnish accurate evidence of the point at which the desired reaction is completed4 ,the nature of the indicators employed will be explained in connection with the different analyses4 ,the process whereby a 6standard solution6 is brought into reaction is called 6titration61 and the point at which the reaction is exactly completed is called the 6end-point64 ,the 6indicator6 should show the 6end-point6 of the 6titration64 ,the volume of the standard solution used then furnishes the measure of the substance to be determined as truly as if that substance had been separated and weighed4 ,the processes of volumetric analysis are easily classified1 according to their character1 into3 ,i4 ,,neutralization ,,methods2 such1 for example1 as those of acidimetry and alkalimetry4 ,,ii4 ,,oxidation ,,processes2 as exemplified in the determination of ferrous iron by its oxidation with potassium bichromate4 ,,iii4 ,,precipitation ,,methods2 of which the titration for silver with potassium thiocyanate solution is an illustration4 ,from a somewhat different standpoint the methods in each case may be subdivided into 7a7 ,,direct ,,methods1 in which the substance to be measured is directly determined by titration to an end-point with a standard solution2 and 7b7 ,,indirect ,,methods1 in which the substance itself is not measured1 but a quantity of reagent is added which is known to be an excess with respect to a specific reaction1 and the unused excess determined by titration4 ,examples of the latter class will be pointed out as they occur in the procedures4 ,,measuring ,,instruments ,,the ,,analytical ,,balance ,for a complete discussion of the physical principles underlying the construction and use of balances1 and the various methods of weighing1 the student is referred to larger manuals of ,quantitative ,analysis1 such as those of ,fresenius1 or ,treadwell-,hall1 and particularly to the admirable discussion of this topic in ,morse's 6,exercises in ,quantitative ,chemistry64 ,the statements and rules of procedure which follow are sufficient for the intelligent use of an analytical balance in connection with processes prescribed in this introductory manual4 ,it is1 however1 imperative that the student should make himself familiar with these essential features of the balance1 and its use4 ,he should fully realize that the analytical balance is a delicate instrument which will render excellent service under careful treatment1 but such treatment is an essential condition if its accuracy is to be depended upon4 ,he should also understand that no set of rules1 however complete1 can do away with the necessity for a sense of personal responsibility1 since by carelessness he can render inaccurate not only his own analyses1 but those of all other students using the same balance4 ,before making any weighings the student should seat himself before a balance and observe the following details of construction3 #a4 ,the balance case is mounted on three brass legs1 which should preferably rest in glass cups1 backed with rubber to prevent slipping4 ,the front legs are adjustable as to height and are used to level the balance case2 the rear leg is of permanent length4 #b4 ,the front of the case may be raised to give access to the balance4 ,in some makes doors are provided also at the ends of the balance case4 #c4 ,the balance beam is mounted upon an upright in the center of the case on the top of which is an inlaid agate plate4 ,to the center of the beam there is attached a steel or agate knife-edge on which the beam oscillates when it rests on the agate plate4 #d4 ,the balance beam1 extending to the right and left1 is graduated along its upper edge1 usually on both sides1 and has at its extremities two agate or steel knife-edges from which are suspended stirrups4 ,each of these stirrups has an agate plate which1 when the balance is in action1 rests upon the corresponding knife-edge of the beam4 ,the balance pans are suspended from the stirrups4 #e4 ,a pointer is attached to the center of the beam1 and as the beam oscillates this pointer moves in front of a scale near the base of the post4 #f4 ,at the base of the post1 usually in the rear1 is a spirit-level4 #g4 ,within the upright is a mechanism1 controlled by a knob at the front of the balance case1 which is so arranged as to raise the entire beam slightly above the level at which the knife-edges are in contact with the agate plates4 ,when the balance is not in use the beam must be supported by this device since1 otherwise1 the constant jarring to which a balance is inevitably subjected1 will soon dull the knife-edges1 and lessen the sensitiveness of the balance4 #h4 ,a small weight1 or bob1 is attached to the pointer 7or sometimes to the beam7 by which the center of gravity of the beam and its attachments may be regulated4 ,the center of gravity must lie very slightly below the level of the agate plates to secure the desired sensitiveness of the balance4 ,this is provided for when the balance is set up and very rarely requires alteration4 ,the student should never attempt to change this adjustment4 #i4 ,below the balance pans are two pan-arrests operated by a button from the front of the case4 ,these arrests exert a very slight upward pressure upon the pans and minimize the displacement of the beam when objects or weights are being placed upon the pans4 #aj4 ,a movable rod1 operated from one end of the balance case1 extends over the balance beam and carries a small wire weight1 called a rider4 ,by means of this rod the rider can be placed upon any desired division of the scale on the balance beam4 ,each numbered division on the beam corresponds to one milligram1 and the use of the rider obviates the placing of very small fractional weights on the balance pan4 ,if a new rider is purchased1 or an old one replaced1 care must be taken that its weight corresponds to the graduations on the beam of the balance on which it is to be used4 ,the weight of the rider in milligrams must be equal to the number of large divisions 7#e1 #f1 #aj1 or #ab7 between the central knife-edge and the knife-edge at the end of the beam4 ,it should be noted that on some balances the last division bears no number4 ,each new rider should be tested against a #e or #aj-milligram weight4 ,in some of the most recent forms of the balance a chain device replaces the smaller weights and the use of the rider as just described4 ,before using a balance1 it is always best to test its adjustment4 ,this is absolutely necessary if the balance is used by several workers2 it is always a wise precaution under any conditions4 ,for this purpose1 brush off the balance pans with a soft camel's hair brush4 ,then note 7#a7 whether the balance is level2 7#b7 that the mechanism for raising and lowering the beams works smoothly2 7#c7 that the pan-arrests touch the pans when the beam is lowered2 and 7#d7 that the needle swings equal distances on either side of the zero-point when set in motion without any load on the pans4 ,if the latter condition is not fulfilled1 the balance should be adjusted by means of the adjusting screw at the end of the beam unless the variation is not more than one division on the scale2 it is often better to make a proper allowance for this small zero error than to disturb the balance by an attempt at correction4 ,unless a student thoroughly understands the construction of a balance he should never attempt to make adjustments1 but should apply to the instructor in charge4 ,the object to be weighed should be placed on the left-hand balance pan and the weights upon the right-hand pan4 ,every substance which could attack the metal of the balance pan should be weighed upon a watch-glass1 and all objects must be dry and cold4 ,a warm body gives rise to air currents which vitiate the accuracy of the weighing4 ,the weights should be applied in the order in which they occur in the weight-box 7not at haphazard71 beginning with the largest weight which is apparently required4 ,after a weight has been placed upon the pan the beam should be lowered upon its knife-edges1 and1 if necessary1 the pan-arrests depressed4 ,the movement of the pointer will then indicate whether the weight applied is too great or too small4 ,when the weight has been ascertained1 by the successive addition of small weights1 to the nearest #e or #aj milligrams1 the weighing is completed by the use of the rider4 ,the correct weight is that which causes the pointer to swing an equal number of divisions to the right and left of the zero-point1 when the pointer traverses not less than five divisions on either side4 ,the balance case should always be closed during the final weighing1 while the rider is being used1 to protect the pans from the effect of air currents4 ,before the final determination of an exact weight the beam should always be lifted from the knife-edges and again lowered into place1 as it frequently happens that the scale pans are1 in spite of the pan-arrests1 slightly twisted by the impact of the weights1 the beam being thereby virtually lengthened or shortened4 ,lifting the beam restores the proper alignment4 ,the beam should never be set in motion by lowering it forcibly upon the knife-edges1 nor by touching the pans1 but rather by lifting the rider 7unless the balance be provided with some of the newer devices for the purpose71 and the swing should be arrested only when the needle approaches zero on the scale1 otherwise the knife-edges become dull4 ,for the same reason the beam should never be left upon its knife-edges1 nor should weights be removed from or placed on the pans without supporting the beam1 except in the case of the small fractional weights4 ,when the process of weighing has been completed1 the weight should be recorded in the notebook by first noting the vacant spaces in the weight-box1 and then checking the weight by again noting the weights as they are removed from the pan4 ,this practice will often detect and avoid errors4 ,it is obvious that the weights should always be returned to their proper places in the box1 and be handled only with pincers4 ,it should be borne in mind that if the mechanism of a balance is deranged or if any substance is spilled upon the pans or in the balance case1 the damage should be reported at once4 ,in many instances serious harm can be averted by prompt action when delay might ruin the balance4 ,samples for analysis are commonly weighed in small tubes with cork stoppers4 ,since the stoppers are likely to change in weight from the varying amounts of moisture absorbed from the atmosphere1 it is necessary to confirm the recorded weight of a tube which has been unused for some time before weighing out a new portion of substance from it4 ,,weights ,the sets of weights commonly used in analytical chemistry range from #bj grams to #e milligrams4 ,the weights from #bj grams to #a gram are usually of brass1 lacquered or gold plated4 ,the fractional weights are of ,german silver1 gold1 platinum or aluminium4 ,the rider is of platinum or aluminium wire4 ,the sets of weights purchased from reputable dealers are usually sufficiently accurate for analytical work4 ,it is not necessary that such a set should be strictly exact in comparison with the absolute standard of weight1 provided they are relatively correct among themselves1 and provided the same set of weights is used in all weighings made during a given analysis4 ,the analyst should assure himself that the weights in a set previously unfamiliar to him are relatively correct by a few simple tests4 ,for example1 he should make sure that in his set two weights of the same denomination 7i4e41 two #aj-;gram weights1 or the two #ajj-milligram weights7 are actually equal and interchangeable1 or that the #ejj-milligram weight is equal to the sum of the #bjj1 #ajj1 #ajj1 #ej1 #bj1 #bj and #aj-milligram weights combined1 and so on4 ,if discrepancies of more than a few tenths of a milligram 7depending upon the total weight involved7 are found1 the weights should be returned for correction4 ,the rider should also be compared with a #e or #aj-milligram weight4 ,in an instructional laboratory appreciable errors should be reported to the instructor in charge for his consideration4 ,when the highest accuracy is desired1 the weights may be calibrated and corrections applied4 ,a calibration procedure is described in a paper by ,t4,w4 ,richards1 6,j4 ,am4 ,chem4 ,soc461 #bb1 #add1 and in many large text-books4 ,weights are inevitably subject to corrosion if not properly protected at all times1 and are liable to damage unless handled with great care4 ,it is obvious that anything which alters the weight of a single piece in an analytical set will introduce an error in every weighing made in which that piece is used4 ,this source of error is often extremely obscure and difficult to detect4 ,the only safeguard against such errors is to be found in scrupulous care in handling and protection on the part of the analyst1 and an equal insistence that if several analysts use the same set of weights1 each shall realize his responsibility for the work of others as well as his own4 ,,burettes ,a burette is made from a glass tube which is as uniformly cylindrical as possible1 and of such a bore that the divisions which are etched upon its surface shall correspond closely to actual contents4 ,the tube is contracted at one extremity1 and terminates in either a glass stopcock and delivery-tube1 or in such a manner that a piece of rubber tubing may be firmly attached1 connecting a delivery-tube of glass4 ,the rubber tubing is closed by means of a glass bead4 ,burettes of the latter type will be referred to as plain burettes4 ,the graduations are usually numbered in cubic centimeters1 and the latter are subdivided into tenths4 ,one burette of each type is desirable for the analytical procedures which follow4 ,,preparation ,,of ,a ,,burette ,,for ,,use ,the inner surface of a burette must be thoroughly cleaned in order that the liquid as drawn out may drain away completely1 without leaving drops upon the sides4 ,this is best accomplished by treating the inside of the burette with a warm solution of chromic acid in concentrated sulphuric acid1 applied as follows3 ,if the burette is of the plain type1 first remove the rubber tip and force the lower end of the burette into a medium-sized cork stopper4 ,nearly fill the burette with the chromic acid solution1 close the upper end with a cork stopper and tip the burette backward and forward in such a way as to bring the solution into contact with the entire inner surface4 ,remove the stopper and pour the solution into a stock bottle to be kept for further use1 and rinse out the burette with water several times4 ,unless the water then runs freely from the burette without leaving drops adhering to the sides1 the process must be repeated 7,note #a74 ,if the burette has a glass stopcock1 this should be removed after the cleaning and wiped1 and also the inside of the ground joint4 ,the surface of the stopcock should then be smeared with a thin coating of vaseline and replaced4 ,it should be attached to the burette by means of a wire1 or elastic band1 to lessen the danger of breakage4 ,fill the burettes with distilled water1 and allow the water to run out through the stopcock or rubber tip until convinced that no air bubbles are inclosed 7,note #b74 ,fill the burette to a point above the zero-point and draw off the water until the meniscus is just below that mark4 ,it is then ready for calibration4 ,note #a3 ,the inner surface of the burette must be absolutely clean if the liquid is to run off freely4 ,chromic acid in sulphuric acid is usually found to be the best cleansing agent1 but the mixture must be warm and concentrated4 ,the solution can be prepared by pouring over a few crystals of potassium bichromate a little water and then adding concentrated sulphuric acid4 ,note #b3 ,it is always necessary to insure the absence of air bubbles in the tips or stopcocks4 ,the treatment described above will usually accomplish this1 but1 in the case of plain burettes it is sometimes better to allow a little of the liquid to flow out of the tip while it is bent upwards4 ,any air which may be entrapped then rises with the liquid and escapes4 ,if air bubbles escape during subsequent calibration or titration1 an error is introduced which vitiates the results4 ,,reading ,,of ,a ,,burette ,all liquids when placed in a burette form what is called a meniscus at their upper surfaces4 ,in the case of liquids such as water or aqueous solutions this meniscus is concave1 and when the liquids are transparent accurate readings are best obtained by observing the position on the graduated scales of the lowest point of the meniscus4 ,this can best be done as follows3 ,wrap around the burette a piece of colored paper1 the straight1 smooth edges of which are held evenly together with the colored side next to the burette 7,note #a74 ,hold the paper about two small divisions below the meniscus and raise or lower the level of the eyes until the edge of the paper at the back of the burette is just hidden from the eye by that in front 7,note #b74 ,note the position of the lowest point of the curve of the meniscus1 estimating the tenths of the small divisions1 thus reading its position to hundredths of a cubic centimeter4 ,note #a3 ,the ends of the colored paper used as an aid to accurate readings may be fastened together by means of a gummed label4 ,the paper may then remain on the burette and be ready for immediate use by sliding it up or down1 as required4 ,note #b3 ,to obtain an accurate reading the eye must be very nearly on a level with the meniscus4 ,this is secured by the use of the paper as described4 ,the student should observe by trial how a reading is affected when the meniscus is viewed from above or below4 ,the eye soon becomes accustomed to estimating the tenths of the divisions4 ,if the paper is held as directed1 two divisions below the meniscus1 one whole division is visible to correct the judgment4 ,it is not well to attempt to bring the meniscus exactly to a division mark on the burette4 ,such readings are usually less accurate than those in which the tenths of a division are estimated4 ,,calibration ,,of ,,glass ,,measuring ,,devices ,if accuracy of results is to be attained1 the correctness of all measuring instruments must be tested4 ,none of the apparatus offered for sale can be implicitly relied upon except those more expensive instruments which are accompanied by a certificate from the 6,national ,bureau of ,standards6 at ,washington1 or other equally authentic source4 ,the bore of burettes is subject to accidental variations1 and since the graduations are applied by machine without regard to such variations of bore1 local errors result4 ,the process of testing these instruments is called 6calibration64 ,it is usually accomplished by comparing the actual weight of water contained in the instrument with its apparent volume4 ,there is1 unfortunately1 no uniform standard of volume which has been adopted for general use in all laboratories4 ,it has been variously proposed to consider the volume of #ajjj grams of water at #d1 #ae4e1 #af1 #ag4e1 and even #bj;,c41 as a liter for practical purposes1 and to consider the cubic centimeter to be one one-thousandth of that volume4 ,the true liter is the volume of #ajjj grams of water at #d;,c42 but this is obviously a lower temperature than that commonly found in laboratories1 and involves the constant use of corrections if taken as a laboratory standard4 ,many laboratories use #ae4e;,c4 7#fj ,f47 as the working standard4 ,it is plain that any temperature which is deemed most convenient might be chosen for a particular laboratory1 but it cannot be too strongly emphasized that all measuring instruments1 including burettes1 pipettes1 and flasks1 should be calibrated at that temperature in order that the contents of each burette1 pipette1 etc41 shall be comparable with that of every other instrument1 thus permitting general interchange and substitution4 ,for example1 it is obvious that if it is desired to remove exactly #ej cc4 from a solution which has been diluted to #ejj cc4 in a graduated flask1 the #ej cc4 flask or pipette used to remove the fractional portion must give a correct reading at the same temperature as the #ejj cc4 flask4 ,similarly1 a burette used for the titration of the #ej cc4 of solution removed should be calibrated under the same conditions as the measuring flasks or pipettes employed with it4 ,the student should also keep constantly in mind the fact that all volumetric operations1 to be exact1 should be carried out as nearly at a constant temperature as is practicable4 ,the spot selected for such work should therefore be subject to a minimum of temperature variations1 and should have as nearly the average temperature of the laboratory as is possible4 ,in all work1 whether of calibration1 standardization1 or analysis1 the temperature of the liquids employed must be taken into account1 and if the temperature of these liquids varies more than #c or #d from the standard temperature chosen for the laboratory1 corrections must be applied for errors due to expansion or contraction1 since volumes of a liquid measured at different times are comparable only under like conditions as to temperature4 ,data to be used for this purpose are given in the ,appendix4 ,neglect of this correction is frequently an avoidable source of error and annoyance in otherwise excellent work4 ,the temperature of all solutions at the time of standardization should be recorded to facilitate the application of temperature corrections1 if such are necessary at any later time4 ,,calibration ,,of ,,the ,,burettes ,two burettes1 one at least of which should have a glass stopper1 are required throughout the volumetric work4 ,both burettes should be calibrated by the student to whom they are assigned4 ,,procedure4--,weigh a #ej cc41 flat-bottomed flask 7preferably a light-weight flask71 which must be dry on the outside1 to the nearest centigram4 ,record the weight in the notebook4 7,see ,appendix for suggestions as to records47 ,place the flask under the burette and draw out into it about #aj cc4 of water1 removing any drop on the tip by touching it against the inside of the neck of the flask4 ,do not attempt to stop exactly at the #aj cc4 mark1 but do not vary more than #j4a cc4 from it4 ,note the time1 and at the expiration of three minutes 7or longer7 read the burette accurately1 and record the reading in the notebook 7,note #a74 ,meanwhile weigh the flask and water to centigrams and record its weight 7,note #b74 ,draw off the liquid from #aj cc4 to about #bj cc4 into the same flask without emptying it2 weigh1 and at the expiration of three minutes take the reading1 and so on throughout the length of the burette4 ,when it is completed1 refill the burette and check the first calibration4 ,the differences in readings represent the apparent volumes1 the differences in weights the true volumes4 ,for example1 if an apparent volume of #aj4je cc4 is found to weigh #aj4jc grams1 it may be assumed with sufficient accuracy that the error in that #aj cc4 amounts to -#j4jb cc41 or -#j4jjb for each cubic centimeter 7,note #c74 ,in the calculation of corrections the temperature of the water must be taken into account1 if this varies more than #d;,c4 from the laboratory standard temperature1 consulting the table of densities of water in the ,appendix4 ,from the final data1 plot the corrections to be applied so that they may be easily read for each cubic centimeter throughout the burette4 ,the total correction at each #aj cc4 may also be written on the burette with a diamond1 or etching ink1 for permanence of record4 ,note #a3 ,a small quantity of liquid at first adheres to the side of even a clean burette4 ,this slowly unites with the main body of liquid1 but requires an appreciable time4 ,three minutes is a sufficient interval1 but not too long1 and should be adopted in every instance throughout the whole volumetric practice before final readings are recorded4 ,note #b3 ,a comparatively rough balance1 capable of weighing to centigrams1 is sufficiently accurate for use in calibrations1 for a moment's reflection will show that it would be useless to weigh the water with an accuracy greater than that of the readings taken on the burette4 ,the latter cannot exceed #j4ja cc4 in accuracy1 which corresponds to #j4ja gram4 ,the student should clearly understand that 6all other weighings61 except those for calibration1 should be made accurately to #j4jjja gram1 unless special directions are given to the contrary4 ,corrections for temperature variations of less than #d;,c4 are negligible1 as they amount to less than #j4ja gram for each #aj grams of water withdrawn4 ,note #c3 ,should the error discovered in any interval of #aj cc4 on the burette exceed #j4aj cc41 it is advisable to weigh small portions 7even #a cc47 to locate the position of the variation of bore in the tube rather than to distribute the correction uniformly over the corresponding #aj cc4 ,the latter is the usual course for small corrections1 and it is convenient to calculate the correction corresponding to each cubic centimeter and to record it in the form of a table or calibration card1 or to plot a curve representing the values4 ,burettes may also be calibrated by drawing off the liquid in successive portions through a #e cc4 pipette which has been accurately calibrated1 as a substitute for weighing4 ,if many burettes are to be tested1 this is a more rapid method4 ,,pipettes ,a 6pipette6 may consist of a narrow tube1 in the middle of which is blown a bulb of a capacity a little less than that which it is desired to measure by the pipette2 or it may be a miniature burette1 without the stopcock or rubber tip at the lower extremity4 ,in either case1 the flow of liquid is regulated by the pressure of the finger on the top1 which governs the admission of the air4 ,pipettes are usually already graduated when purchased1 but they require calibration for accurate work4 ,,calibration ,,of ,,pipettes ,,procedure4--,clean the pipette4 ,draw distilled water into it by sucking at the upper end until the water is well above the graduation mark4 ,quickly place the forefinger over the top of the tube1 thus preventing the entrance of air and holding the water in the pipette4 ,cautiously admit a little air by releasing the pressure of the finger1 and allow the level of the water to fall until the lowest point of the meniscus is level with the graduation4 ,hold the water at that point by pressure of the finger and then allow the water to run out from the pipette into a small tared1 or weighed1 beaker or flask4 ,after a definite time interval1 usually two to three minutes1 touch the end of the pipette against the side of the beaker or flask to remove any liquid adhering to it 7,note #a74 ,the increase in weight of the flask in grams represents the volume of the water in cubic centimeters delivered by the pipette4 ,calculate the necessary correction4 ,note #a3 ,a definite interval must be allowed for draining1 and a definite practice adopted with respect to the removal of the liquid which collects at the end of the tube1 if the pipette is designed to deliver a specific volume when emptied4 ,this liquid may be removed at the end of a definite interval either by touching the side of the vessel or by gently blowing out the last drops4 ,either practice1 when adopted1 must be uniformly adhered to4 ,,flasks 6,graduated or measuring flasks6 are similar to the ordinary flat-bottomed flasks1 but are provided with long1 narrow necks in order that slight variations in the position of the meniscus with respect to the graduation shall represent a minimum volume of liquid4 ,the flasks must be of such a capacity that1 when filled with the specified volume1 the liquid rises well into the neck4 ,,graduation ,,of ,,flasks ,it is a general custom to purchase the flasks ungraduated and to graduate them for use under standard conditions selected for the laboratory in question4 ,they may be graduated for contents or delivery4 ,when graduated for contents they contain a specified volume when filled to the graduation at a specified temperature1 and require to be washed out in order to remove all of the solution from the flask4 ,flasks graduated for delivery will deliver the specified volume of a liquid without rinsing4 ,a flask may1 of course1 be graduated for both contents and delivery by placing two graduation marks upon it4 ,,procedure4--,to calibrate a flask for 6contents61 proceed as follows3 ,clean the flask1 using a chromic acid solution1 and dry it carefully outside and inside4 ,tare it accurately2 pour water into the flask until the weight of the latter counterbalances weights on the opposite pan which equal in grams the number of cubic centimeters of water which the flask is to contain4 ,remove any excess of water with the aid of filter paper 7,note #a74 ,take the flask from the balance1 stopper it1 place it in a bath at the desired temperature1 usually #ae4e or #ag4e;,c41 and after an hour mark on the neck with a diamond the location of the lowest point of the meniscus 7,note #b74 ,the mark may be etched upon the flask by hydrofluoric acid1 or by the use of an etching ink now commonly sold on the market4 ,to graduate a flask which is designed to 6deliver6 a specified volume1 proceed as follows3 ,clean the flask as usual and wipe all moisture from the outside4 ,fill it with distilled water4 ,pour out the water and allow the water to drain from the flask for three minutes4 ,counterbalance the flask with weights to the nearest centigram4 ,add weights corresponding in grams to the volume desired1 and add distilled water to counterbalance these weights4 ,an excess of water1 or water adhering to the neck of the flask1 may be removed by means of a strip of clean filter paper4 ,stopper the flask1 place it in a bath at #ae4e;,c4 or #ag4e;,c4 and1 after an hour1 mark the location of the lowest point of the meniscus1 as described above4 ,note #a3 ,the allowable error in counterbalancing the water and weights varies with the volume of the flask4 ,it should not exceed one ten-thousandth of the weight of water4 ,note #b3 ,other methods are employed which involve the use of calibrated apparatus from which the desired volume of water may be run into the dry flask and the position of the meniscus marked directly upon it4 ,for a description of a procedure which is most convenient when many flasks are to be calibrated1 the student is referred to the 6,am4 ,chem ,j461 #af1 #dgi4 ,,general ,,directions ,,for ,,volumetric ,,analyses ,it cannot be too strongly emphasized that for the success of analyses uniformity of practice must prevail throughout all volumetric work with respect to those factors which can influence the accuracy of the measurement of liquids4 ,for example1 whatever conditions are imposed during the calibration of a burette1 pipette1 or flask 7notably the time allowed for draining71 must also prevail whenever the flask or burette is used4 ,the student should also be constantly watchful to insure parallel conditions during both standardization and analyst with respect to the final volume of liquid in which a titration takes place4 ,the value of a standard solution is only accurate under the conditions which prevailed when it was standardized4 ,it is plain that the standard solutions must be scrupulously protected from concentration or dilution1 after their value has been established4 ,accordingly1 great care must be taken to thoroughly rinse out all burettes1 flasks1 etc41 with the solutions which they are to contain1 in order to remove all traces of water or other liquid which could act as a diluent4 ,it is best to wash out a burette at least three times with small portions of a solution1 allowing each to run out through the tip before assuming that the burette is in a condition to be filled and used4 ,it is1 of course1 possible to dry measuring instruments in a hot closet1 but this is tedious and unnecessary4 ,to the same end1 all solutions should be kept stoppered and away from direct sunlight or heat4 ,the bottles should be shaken before use to collect any liquid which may have distilled from the solution and condensed on the sides4 ,the student is again reminded that variations in temperature of volumetric solutions must be carefully noted1 and care should always be taken that no source of heat is sufficiently near the solutions to raise the temperature during use4 ,much time may be saved by estimating the approximate volume of a standard solution which will be required for a titration 7if the data are obtainable7 before beginning the operation4 ,it is then possible to run in rapidly approximately the required amount1 after which it is only necessary to determine the end-point slowly and with accuracy4 ,in such cases1 however1 the knowledge of the approximate amount to be required should never be allowed to influence the judgment regarding the actual end-point4 ,,standard ,,solutions ,the strength or value of a solution for a specific reaction is determined by a procedure called 6,standardization61 in which the solution is brought into reaction with a definite weight of a substance of known purity4 ,for example1 a definite weight of pure sodium carbonate may be dissolved in water1 and the volume of a solution of hydrochloric acid necessary to exactly neutralize the carbonate accurately determined4 ,from these data the strength or value of the acid is known4 ,it is then a 6standard solution64 ,,normal ,,solutions ,standard solutions may be made of a purely empirical strength dictated solely by convenience of manipulation1 or the concentration may be chosen with reference to a system which is applicable to all solutions1 and based upon chemical equivalents4 ,such solutions are called 6,normal ,solutions6 and contain such an amount of the reacting substance per liter as is equivalent in its chemical action to one gram of hydrogen1 or eight grams of oxygen4 ,solutions containing one half1 one tenth1 or one one-hundredth of this quantity per liter are called1 respectively1 half-normal1 tenth-normal1 or hundredth-normal solutions4 ,since normal solutions of various reagents are all referred to a common standard1 they have an advantage not possessed by empirical solutions1 namely1 that they are exactly equivalent to each other4 ,thus1 a liter of a normal solution of an acid will exactly neutralize a liter of a normal alkali solution1 and a liter of a normal oxidizing solution will exactly react with a liter of a normal reducing solution1 and so on4 ,beside the advantage of uniformity1 the use of normal solutions simplifies the calculations of the results of analyses4 ,this is particularly true if1 in connection with the normal solution1 the weight of substance for analysis is chosen with reference to the atomic or molecular weight of the constituent to be determined4 7,see problem #bf47 ,the preparation of an 6exactly6 normal1 half-normal1 or tenth-normal solution requires considerable time and care4 ,it is usually carried out only when a large number of analyses are to be made1 or when the analyst has some other specific purpose in view4 ,it is1 however1 a comparatively easy matter to prepare standard solutions which differ but slightly from the normal or half-normal solution1 and these have the advantage of practical equality2 that is1 two approximately half-normal solutions are more convenient to work with than two which are widely different in strength4 ,it is1 however1 true that some of the advantage which pertains to the use of normal solutions as regards simplicity of calculations is lost when using these approximate solutions4 ,the application of these general statements will be made clear in connection with the use of normal solutions in the various types of volumetric processes which follow4 ,i4 ,,neutralization ,,methods ,,alkalimetry ,,and ,,acidimetry ,,general ,,discussion 6,standard ,acid ,solutions6 may be prepared from either hydrochloric1 sulphuric1 or oxalic acid4 ,hydrochloric acid has the advantage of forming soluble compounds with the alkaline earths1 but its solutions cannot be boiled without danger of loss of strength2 sulphuric acid solutions may be boiled without loss1 but the acid forms insoluble sulphates with three of the alkaline earths2 oxalic acid can be accurately weighed for the preparation of solutions1 and its solutions may be boiled without loss1 but it forms insoluble oxalates with three of the alkaline earths and cannot be used with certain of the indicators4 6,standard ,alkali ,solutions6 may be prepared from sodium or potassium hydroxide1 sodium carbonate1 barium hydroxide1 or ammonia4 ,of sodium and potassium hydroxide1 it may be said that they can be used with all indicators1 and their solutions may be boiled1 but they absorb carbon dioxide readily and attack the glass of bottles1 thereby losing strength2 sodium carbonate may be weighed directly if its purity is assured1 but the presence of carbonic acid from the carbonate is a disadvantage with many indicators2 barium hydroxide solutions may be prepared which are entirely free from carbon dioxide1 and such solutions immediately show by precipitation any contamination from absorption1 but the hydroxide is not freely soluble in water2 ammonia does not absorb carbon dioxide as readily as the caustic alkalies1 but its solutions cannot be boiled nor can they be used with all indicators4 ,the choice of a solution must depend upon the nature of the work in hand4 ,a 6normal acid solution6 should contain in one liter that quantity of the reagent which represents #a gram of hydrogen replaceable by a base4 ,for example1 the normal solution of hydrochloric acid 7,h,cl7 should contain #cf4df grams of gaseous hydrogen chloride1 since that amount furnishes the requisite #a gram of replaceable hydrogen4 ,on the other hand1 the normal solution of sulphuric acid 7,h#b,,so#d7 should contain only #di4jc grams1 i4e41 one half of its molecular weight in grams4 ,a 6normal alkali solution6 should contain sufficient alkali in a liter to replace #a gram of hydrogen in an acid4 ,this quantity is represented by the molecular weight in grams 7#dj4ja7 of sodium hydroxide 7,na,o,h71 while a sodium carbonate solution 7,na#b;,,co#c7 should contain but one half the molecular weight in grams 7i4e41 #ec4j grams7 in a liter of normal solution4 ,half-normal or tenth-normal solutions are employed in most analyses 7except in the case of the less soluble barium hydroxide74 ,solutions of the latter strength yield more accurate results when small percentages of acid or alkali are to be determined4 ,,indicators ,it has already been pointed out that the purpose of an indicator is to mark 7usually by a change of color7 the point at which just enough of the titrating solution has been added to complete the chemical change which it is intended to bring about4 ,in the neutralization processes which are employed in the measurement of alkalies 76alkalimetry67 or acids 76acidimetry67 the end-point of the reaction should1 in principle1 be that of complete neutrality4 ,expressed in terms of ionic reactions1 it should be the point at which the ,h ions from an acid,note #a unite with a corresponding number of ,,oh- ions from a base to form water molecules1 as in the equation ,h1 ,cl- ,na1 ,,oh- -- ,na1 ,cl- 7,h#b,o74 ,it is not usually possible to realize this condition of exact neutrality1 but it is possible to approach it with sufficient exactness for analytical purposes1 since substances are known which1 in solution1 undergo a sharp change of color as soon as even a minute excess of ,h or ,,oh- ions are present4 ,some1 as will be seen1 react sharply in the presence of ,h ions1 and others with ,,oh- ions4 ,these substances employed as indicators are usually organic compounds of complex structure and are closely allied to the dyestuffs in character4 ,note #a3 ,a knowledge on the part of the student of the ionic theory as applied to aqueous solutions of electrolytes is assumed4 ,a brief outline of the more important applications of the theory is given in the ,appendix4 ,,behavior ,,of ,,organic ,,indicators ,the indicators in most common use for acid and alkali titrations are methyl orange1 litmus1 and phenolphthalein4 ,in the following discussion of the principles underlying the behavior of the indicators as a class1 methyl orange and phenolphthalein will be taken as types4 ,it has just been pointed out that indicators are bodies of complicated structure4 ,in the case of the two indicators named1 the changes which they undergo have been carefully studied by ,stieglitz 76,j4 ,am4 ,chem4 ,soc461 #be1 #aaab7 and others1 and it appears that the changes involved are of two sorts3 ,first1 a rearrangement of the atoms within the molecule1 such as often occurs in organic compounds2 and1 second1 ionic changes4 ,the intermolecular changes cannot appropriately be discussed here1 as they involve a somewhat detailed knowledge of the classification and general behavior of organic compounds2 they will1 therefore1 be merely alluded to1 and only the ionic changes followed4 ,methyl orange is a representative of the group of indicators which1 in aqueous solutions1 behave as weak bases4 ,the yellow color which it imparts to solutions is ascribed to the presence of the undissociated base4 ,if an acid1 such as ,h,cl1 is added to such a solution1 the acid reacts with the indicator 7neutralizes it7 and a salt is formed1 as indicated by the equation3 7,m4o471 ,,oh- ,h1 ,cl- -- 7,m4o47 ,cl- 7,h#b,o74 ,this salt ionizes into 7,m4o47 7using this abbreviation for the positive complex7 and ,cl-2 but simultaneously with this ionization there appears to be an internal rearrangement of the atoms which results in the production of a cation which may be designated as 7,m'4o'471 and it is this which imparts a characteristic red color to the solution4 ,as these changes occur in the presence of even a very small excess of acid 7that is1 of ,h ions71 it serves as the desired index of their presence in the solution4 ,if1 now1 an alkali1 such as ,na,o,h1 is added to this reddened solution1 the reverse series of changes takes place4 ,as soon as the free acid present is neutralized1 the slightest excess of sodium hydroxide1 acting as a strong base1 sets free the weak1 little-dissociated base of the indicator1 and at the moment of its formation it reverts1 because of the rearrangement of the atoms1 to the yellow form3 ,,oh- 7,m'4o'47 -- ,m'4o'4,,oh -- ,m4o4,,oh4 ,phenolphthalein1 on the other hand1 is a very weak1 little-dissociated acid1 which is colorless in neutral aqueous solution or in the presence of free ,h ions4 ,when an alkali is added to such a solution1 even in slight excess1 the anion of the salt which has formed from the acid of the indicator undergoes a rearrangement of the atoms1 and a new ion1 7,ph'71 is formed1 which imparts a pink color to the solution3 ,h1 7,ph7- ,na1 ,,oh- -- 7,h#b,o7 ,na1 7,ph7- -- ,na1 7,ph'7- ,the addition of the slightest excess of an acid to this solution1 on the other hand1 occasions first the reversion to the colorless ion and then the setting free of the undissociated acid of the indicator3 ,h1 7,ph'7- -- ,h1 7,ph7- -- 7,h,ph74 ,of the common indicators methyl orange is the most sensitive toward alkalies and phenolphthalein toward acids2 the others occupy intermediate positions4 ,that methyl orange should be most sensitive toward alkalies is evident from the following considerations3 ,methyl orange is a weak base and1 therefore1 but little dissociated4 ,it should1 then1 be formed in the undissociated condition as soon as even a slight excess of ,,oh- ions is present in the solution1 and there should be a prompt change from red to yellow as outlined above4 ,on the other hand1 it should be an unsatisfactory indicator for use with weak acids 7acetic acid1 for example7 because the salts which it forms with such acids are1 like all salts of that type1 hydrolyzed to a considerable extent4 ,this hydrolytic change is illustrated by the equation3 7,m4o47 ,c#b;,h#c,o#b- ,h1 ,,oh- -- ,m4o4,,oh ,h1 ,c#b;,h#c,o#b-4 ,comparison of this equation with that on page #cj will make it plain that hydrolysis is just the reverse of neutralization and must1 accordingly1 interfere with it4 ,salts of methyl orange with weak acids are so far hydrolyzed that the end-point is uncertain1 and methyl orange cannot be used in the titration of such acids1 while with the very weak acids1 such as carbonic acid or hydrogen sulphide 7hydrosulphuric acid71 the salts formed with methyl orange are1 in effect1 completely hydrolyzed 7i4e41 no neutralization occurs71 and methyl orange is accordingly scarcely affected by these acids4 ,this explains its usefulness1 as referred to later1 for the titration of strong acids1 such as hydrochloric acid1 even in the presence of carbonates or sulphides in solution4 ,phenolphthalein1 on the other hand1 should be1 as it is1 the best of the common indicators for use with weak acids4 ,for1 since it is itself a weak acid1 it is very little dissociated1 and its nearly undissociated1 colorless molecules are promptly formed as soon as there is any free acid 7that is1 free ,h ions7 in the solution4 ,this indicator cannot1 however1 be successfully used with weak bases1 even ammonium hydroxide2 for1 since it is weak acid1 the salts which it forms with weak alkalies are easily hydrolyzed1 and as a consequence of this hydrolysis the change of color is not sharp4 ,this indicator can1 however1 be successfully used with strong bases1 because the salts which it forms with such bases are much less hydrolyzed and because the excess of ,,oh- ions from these bases also diminishes the hydrolytic action of water4 ,this indicator is affected by even so weak an acid as carbonic acid1 which must be removed by boiling the solution before titration4 ,it is the indicator most generally employed for the titration of organic acids4 ,in general1 it may be stated that when a strong acid1 such as hydrochloric1 sulphuric or nitric acid1 is titrated against a strong base1 such as sodium hydroxide1 potassium hydroxide1 or barium hydroxide1 any of these indicators may be used1 since very little hydrolysis ensues4 ,it has been noted above that the color change does not occur exactly at theoretical neutrality1 from which it follows that no two indicators will show exactly the same end-point when acids and alkalis are brought together4 ,it is plain1 therefore1 that the same indicator must be employed for both standardization and analysis1 and that1 if this is done1 accurate results are obtainable4 ,the following table 7,note #a7 illustrates the variations in the volume of an alkali solution 7tenth-normal sodium hydroxide7 required to produce an alkaline end-point when run into #aj cc4 of tenth-normal sulphuric acid1 diluted with #ej cc4 of water1 using five drops of each of the different indicator solutions4 ,,indicator ,n_/#aj ,n_/#aj ,,color ,,in ,,acid,,color ,,in ,,alka- ,h#b,,so#d ,na,o,h ,,solution ,,line ,,solution cc4 cc4 cc4 ,methyl orange #aj #i4ij ,red ,yellow ,lacmoid #aj #aj4jj ,red ,blue ,litmus #aj #aj4jj ,red ,blue ,rosalic acid #aj #aj4jg ,yellow ,pink ,phenolphthalein #aj #aj4aj ,colorless ,pink ,it should also be stated that there are occasionally secondary changes1 other than those outlined above1 which depend upon the temperature and concentration of the solutions in which the indicators are used4 ,these changes may influence the sensitiveness of an indicator4 ,it is important1 therefore1 to take pains to use approximately the same volume of solution when standardizing that is likely to be employed in analysis2 and when it is necessary1 as is often the case1 to titrate the solution at boiling temperature1 the standardization should take place under the same conditions4 ,it is also obvious that since some acid or alkali is required to react with the indicator itself1 the amount of indicator used should be uniform and not excessive4 ,usually a few drops of solution will suffice4 ,the foregoing statements with respect to the behavior of indicators present the subject in its simplest terms4 ,many substances other than those named may be employed1 and they have been carefully studied to determine the exact concentration of ,h ions at which the color change of each occurs4 ,it is thus possible to select an indicator for a particular purpose with considerable accuracy4 ,as data of this nature do not belong in an introductory manual1 reference is made to the following papers or books in which a more extended treatment of the subject may be found3 ,washburn1 ,e4,w41 ,principles of ,physical ,chemistry 7,mc,graw-,hill ,book ,co471 7,second ,edition1 #aiba71 pp4 #chj-chg4 ,prideaux1 ,e4,b4,r41 ,the ,theory and ,use of ,indicators 7,constable ,co41 ,ltd471 7#aiag74 ,salm1 ,e41 ,a ,study of ,indicators1 6,z4 physik4 ,chem461 #eg 7#aijf71 #dga-eja4 ,stieglitz1 ,j41 ,theories of ,indicators1 6,j4 ,am4 ,chem4 ,soc461 #be 7#aijc71 #aaab-aabg4 ,noyes1 ,a4,a41 ,quantitative ,applications of the ,theory of ,indicators to ,volumetric ,analysis1 6,j4 ,am4 ,chem4 ,soc461 #cb 7#aiaa71 #hae-hfa4 ,bjerrum1 ,n41 ,general ,discussion1 6,z4 ,anal4 ,chem461 #ff 7#aiag71 #ac-bh and #ha-ie4 ,ostwald1 ,w41 ,colloid ,chemistry of ,indicators1 6,z4 ,chem4 ,ind4 ,kolloide61 #aj 7#aiab71 #acb-adf4 ,note #a3 ,glaser1 6,indikatoren der ,acidimetrie und ,alkalimetrie64 ,wiesbaden1 #aija4 ,,preparation ,,of ,,indicator ,,solutions ,a 6methyl orange solution6 for use as an indicator is commonly made by dissolving #j4je-j4a gram of the compound 7also known as ,orange ,,iii7 in a few cubic centimeters of alcohol and diluting with water to #ajj cc4 ,a good grade of material should be secured4 ,it can be successfully used for the titration of hydrochloric1 nitric1 sulphuric1 phosphoric1 and sulphurous acids1 and is particularly useful in the determination of bases1 such as sodium1 potassium1 barium1 calcium1 and ammonium hydroxides1 and even many of the weak organic bases4 ,it can also be used for the determination1 by titration with a standard solution of a strong acid1 of the salts of very weak acids1 such as carbonates1 sulphides1 arsenites1 borates1 and silicates1 because the weak acids which are liberated do not affect the indicator1 and the reddening of the solution does not take place until an excess of the strong acid is added4 ,it should be used in cold1 not too dilute1 solutions4 ,its sensitiveness is lessened in the presence of considerable quantities of the salts of the alkalies4 ,a 6phenolphthalein solution6 is prepared by dissolving #a gram of the pure compound in #ajj cc4 of #ie per cent alcohol4 ,this indicator is particularly valuable in the determination of weak acids1 especially organic acids4 ,it cannot be used with weak bases1 even ammonia4 ,it is affected by carbonic acid1 which must1 therefore1 be removed by boiling when other acids are to be measured4 ,it can be used in hot solutions4 ,some care is necessary to keep the volume of the solutions to be titrated approximately uniform in standardization and in analysis1 and this volume should not in general exceed #abe-aej cc4 for the best results1 since the compounds formed by the indicator undergo changes in very dilute solution which lessen its sensitiveness4 ,the preparation of a 6solution of litmus6 which is suitable for use as an indicator involves the separation from the commercial litmus of azolithmine1 the true coloring principle4 ,soluble litmus tablets are often obtainable1 but the litmus as commonly supplied to the market is mixed with calcium carbonate or sulphate and compressed into lumps4 ,to prepare a solution1 these are powdered and treated two or three times with alcohol1 which dissolves out certain constituents which cause a troublesome intermediate color if not removed4 ,the alcohol is decanted and drained off1 after which the litmus is extracted with hot water until exhausted4 ,the solution is allowed to settle for some time1 the clear liquid siphoned off1 concentrated to one-third its volume and acetic acid added in slight excess4 ,it is then concentrated to a sirup1 and a large excess of #ie per cent4 alcohol added to it4 ,this precipitates the blue coloring matter1 which is filtered off1 washed with alcohol1 and finally dissolved in a small volume of water and diluted until about three drops of the solution added to #ej cc4 of water just produce a distinct color4 ,this solution must be kept in an unstoppered bottle4 ,it should be protected from dust by a loose plug of absorbent cotton4 ,if kept in a closed bottle it soon undergoes a reduction and loses its color1 which1 however1 is often restored by exposure to the air4 ,litmus can be employed successfully with the strong acids and bases1 and also with ammonium hydroxide1 although the salts of the latter influence the indicator unfavorably if present in considerable concentration4 ,it may be employed with some of the stronger organic acids1 but the use of phenolphthalein is to be preferred4 ,,preparation ,,of ,,standard ,,solutions 6,hydrochloric ,acid and ,sodium ,hydroxide4 ,approximate ,strength61 #j4e ,n ,,procedure4--,measure out #dj cc4 of concentrated1 pure hydrochloric acid into a clean liter bottle1 and dilute with distilled water to an approximate volume of #ajjj cc4 ,shake the solution vigorously for a full minute to insure uniformity4 ,be sure that the bottle is not too full to permit of a thorough mixing1 since lack of care at this point will be the cause of much wasted time 7,note #a74 ,weigh out1 upon a rough balance1 #bc grams of sodium hydroxide 7,note #b74 ,dissolve the hydroxide in water in a beaker4 ,pour the solution into a liter bottle and dilute1 as above1 to approximately #ajjj cc4 ,this bottle should preferably have a rubber stopper1 as the hydroxide solution attacks the glass of the ground joint of a glass stopper1 and may cement the stopper to the bottle4 ,shake the solution as described above4 ,note #a3 ,the original solutions are prepared of a strength greater than #j4e ,n1 as they are more readily diluted than strengthened if later adjustment is desired4 ,too much care cannot be taken to insure perfect uniformity of solutions before standardization1 and thoroughness in this respect will1 as stated1 often avoid much waste of time4 ,a solution once thoroughly mixed remains uniform4 ,note #b3 ,commercial sodium hydroxide is usually impure and always contains more or less carbonate2 an allowance is therefore made for this impurity by placing the weight taken at #bc grams per liter4 ,if the hydroxide is known to be pure1 a lesser amount 7say #ba grams7 will suffice4 ,,comparison ,,of ,,acid ,,and ,,alkali ,,solutions ,,procedure4--,rinse a previously calibrated burette three times with the hydrochloric acid solution1 using #aj cc4 each time1 and allowing the liquid to run out through the tip to displace all water and air from that part of the burette4 ,then fill the burette with the acid solution4 ,carry out the same procedure with a second burette1 using the sodium hydroxide solution4 ,the acid solution may be placed in a plain or in a glass-stoppered burette as may be more convenient1 but the alkaline solution should never be allowed to remain long in a glass-stoppered burette1 as it tends to cement the stopper to the burette1 rendering it useless4 ,it is preferable to use a plain burette for this solution4 ,when the burettes are ready for use and all air bubbles displaced from the tip 7see ,note #b1 page #ag7 note the exact position of the liquid in each1 and record the readings in the notebook4 7,consult page #ahh47 ,run out from the burette into a beaker about #dj cc4 of the acid and add two drops of a solution of methyl orange2 dilute the acid to about #hj cc4 and run out alkali solution from the other burette1 stirring constantly1 until the pink has given place to a yellow4 ,wash down the sides of the beaker with a little distilled water if the solution has spattered upon them1 return the beaker to the acid burette1 and add acid to restore the pink2 continue these alternations until the point is accurately fixed at which a single drop of either solutions served to produce a distinct change of color4 ,select as the final end-point the appearance of the faintest pink tinge which can be recognized1 or the disappearance of this tinge1 leaving a pure yellow2 but always titrate to the same point 7,note #a74 ,if the titration has occupied more than the three minutes required for draining the sides of the burette1 the final reading may be taken immediately and recorded in the notebook4 ,refill the burettes and repeat the titration4 ,from the records of calibration already obtained1 correct the burette readings and make corrections for temperature1 if necessary4 ,obtain the ratio of the sodium hydroxide solution to that of hydrochloric acid by dividing the number of cubic centimeters of acid used by the number of cubic centimeters of alkali required for neutralization4 ,the check results of the two titrations should not vary by more than two parts in one thousand 7,note #b74 ,if the variation in results is greater than this1 refill the burettes and repeat the titration until satisfactory values are obtained4 ,use a new page in the notebook for each titration4 ,inaccurate values should not be erased or discarded4 ,they should be retained and marked correct or incorrect1 as indicated by the final outcome of the titrations4 ,this custom should be rigidly followed in all analytical work4 ,note #a3 ,the end-point should be chosen exactly at the point of change2 any darker tint is unsatisfactory1 since it is impossible to carry shades of color in the memory and to duplicate them from day to day4 ,note #b3 ,while variation of two parts in one thousand in the values obtained by an inexperienced analyst is not excessive1 the idea must be carefully avoided that this is a standard for accurate work to be 6generally applied64 ,in many cases1 after experience is gained1 the allowable error is less than this proportion4 ,in a few cases a larger variation is permissible1 but these are rare and can only be recognized by an experienced analyst4 ,it is essential that the beginner should acquire at least the degree of accuracy indicated if he is to become a successful analyst4 ,,standardization ,,of ,,hydrochloric ,,acid ,,selection ,,and ,,preparation ,,of ,,standard ,the selection of the best substance to be used as a standard for acid solutions has been the subject of much controversy4 ,the work of ,lunge 76,ztschr4 angew4 ,chem46 7#aijd71 #h1 #bca71 ,ferguson 76,j4 ,soc4 ,chem4 ,ind46 7#aije71 #bd1 #ghd71 and others1 seems to indicate that the best standard is sodium carbonate prepared from sodium bicarbonate by heating the latter at temperature between #bgj and #cjj;,c4 ,the bicarbonate is easily prepared in a pure state1 and at the temperatures named the decomposition takes place according to the equation #b;,h,na,c,o#c -- ,na#b;,,co#c ,h#b,o ,,co#b and without loss of any carbon dioxide from the sodium carbonate1 such as may occur at higher temperatures4 ,the process is carried out as described below4 ,,procedure4--,place in a porcelain crucible about #f grams 7roughly weighed7 of the purest sodium bicarbonate obtainable4 ,rest the crucible upon a triangle of iron or copper wire so placed within a large crucible that there is an open air space of about three eighths of an inch between them4 ,the larger crucible may be of iron1 nickel or porcelain1 as may be most convenient4 ,insert the bulb of a thermometer reading to #cej;,c4 in the bicarbonate1 supporting it with a clamp so that the bulb does not rest on the bottom of the crucible4 ,heat the outside crucible1 using a rather small flame1 and raise the temperature of the bicarbonate fairly rapidly to #bgj;,c4 ,then regulate the heat in such a way that the temperature rises 6slowly6 to #cjj;,c4 in the course of a half-hour4 ,the bicarbonate should be frequently stirred with a clean1 dry1 glass rod1 and after stirring1 should be heaped up around the bulb of the thermometer in such a way as to cover it4 ,this will require attention during most of the heating1 as the temperature should not be permitted to rise above #caj;,c4 for any length of time4 ,at the end of the half-hour remove the thermometer and transfer the porcelain crucible1 which now contains sodium carbonate1 to a desiccator4 ,when it is cold1 transfer the carbonate to a stoppered weighing tube or weighing-bottle4 ,,standardization ,,procedure4--,clean carefully the outside of a weighing-tube1 or weighing-bottle1 containing the pure sodium carbonate1 taking care to handle it as little as possible after wiping4 ,weigh the tube accurately to #j4jjja gram1 and record the weight in the notebook4 ,hold the tube over the top of a beaker 7#bjj-cjj cc47 and cautiously remove the stopper1 making sure that no particles fall from it or from the tube elsewhere than in the beaker4 ,pour out from the tube a portion of the carbonate1 replace the stopper and determine approximately how much has been removed4 ,continue this procedure until #a4jj to #a4aj grams has been taken from the tube4 ,then weigh the tube accurately and record the weight under the first weight in the notebook4 ,the difference in the two weights is the weight of the carbonate transferred to the beaker4 ,proceed in the same way to transfer a second portion of the carbonate from the tube to another beaker of about the same size as the first4 ,the beakers should be labeled and plainly marked to correspond with the entries in the notebook4 ,pour over the carbonate in each beaker about #hj cc4 of water1 stir until solution is complete1 and add two drops of methyl orange solution4 ,fill the burettes with the standard acid and alkali solutions1 noting the initial readings of the burettes and temperature of the solutions4 ,run in acid from the burette1 stirring and avoiding loss by effervescence1 until the solution has become pink4 ,wash down the sides of the beaker with a 6little6 water from a wash-bottle1 and then run in alkali from the other burette until the pink is replaced by yellow2 then finish the titration as described on page #cg4 ,note the readings of the burettes after the proper interval1 and record them in the notebook4 ,repeat the procedure1 using the second portion of sodium carbonate4 ,apply the necessary calibration corrections to the volumes of the solutions used1 and correct for temperature if necessary4 ,from the data obtained1 calculate the volume of the hydrochloric acid solution which is equivalent to the volume of sodium hydroxide solution used in this titration4 ,subtract this volume from the volume of hydrochloric acid4 ,the difference represents the volume of acid used to react with the sodium carbonate4 ,divide the weight of sodium carbonate by this volume in cubic centimeters1 thus obtaining the weight of sodium carbonate equivalent to each cubic centimeter of the acid4 ,from this weight it is possible to calculate the corresponding weight of ,h,cl in each cubic centimeter of the acid1 and in turn the relation of the acid to the normal4 ,if1 however1 it is recalled that normal solutions are equivalent to each other1 it will be seen that the same result may be more readily reached by dividing the weight in grams of sodium carbonate per cubic centimeter just found by titration by the weight which would be contained in the same volume of a normal solution of sodium carbonate4 ,a normal solution of sodium carbonate contains #ec4j grams per liter1 or #j4jecj gram per cc4 7see page #bi74 ,the relation of the acid solution to the normal is1 therefore1 calculated by dividing the weight of the carbonate to which each cubic centimeter of the acid is equivalent by #j4jecj4 ,the standardization must be repeated until the values obtained agree within1 at most1 two parts in one thousand4 ,when the standard of the acid solution has been determined1 calculate1 from the known ratio of the two solutions1 the relation of the sodium hydroxide solution to a normal solution 7,notes #a and #b74 ,note #a3 ,in the foregoing procedure the acid solution is standardized and the alkali solution referred to this standard by calculation4 ,it is equally possible1 if preferred1 to standardize the alkali solution4 ,the standards in a common use for this purpose are purified oxalic acid 7,h#b;,c#b,o#d4b;,h#b,o71 potassium acid oxalate 7,,khc#b,o#d4;,h#b,o or ,,khc#b,o#d71 potassium tetroxalate 7,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o71 or potassium acid tartrate 7,,khc#d,o#f71 with the use of a suitable indicator4 ,the oxalic acid and the oxalates should be specially prepared to insure purity1 the main difficulty lying in the preservation of the water of crystallization4 ,it should be noted that the acid oxalate and the acid tartrate each contain one hydrogen atom replaceable by a base1 while the tetroxalate contains three such atoms and the oxalic acid two4 ,each of the two salts first named behave1 therefore1 as monobasic acids1 and the tetroxalate as a tribasic acid4 ,note #b3 ,it is also possible to standardize a hydrochloric acid solution by precipitating the chloride ions as silver chloride and weighing the precipitate1 as prescribed under the analysis of sodium chloride to be described later4 ,sulphuric acid solutions may be standardized by precipitation of the sulphate ions as barium sulphate and weighing the ignited precipitate1 but the results are not above criticism on account of the difficulty in obtaining large precipitates of barium sulphate which are uncontaminated by inclosures or are not reduced on ignition4 ,,determination ,,of ,,the ,,total ,,alkaline ,,strength ,,of ,,soda ,,ash ,soda ash is crude sodium carbonate4 ,if made by the ammonia process it may contain also sodium chloride1 sulphate1 and hydroxide2 when made by the ,le ,blanc process it may contain sodium sulphide1 silicate1 and aluminate1 and other impurities4 ,some of these1 notably the hydroxide1 combine with acids and contribute to the total alkaline strength1 but it is customary to calculate this strength in terms of sodium carbonate2 i4e41 as though no other alkali were present4 ,,procedure4--,in order to secure a sample which shall represent the average value of the ash1 it is well to take at least #e grams4 ,as this is too large a quantity for convenient titration1 an aliquot portion of the solution is measured off1 representing one fifth of the entire quantity4 ,this is accomplished as follows3 ,weigh out on an analytical balance two samples of soda ash of about #e grams each into beakers of about #ejj cc4 capacity4 7,the weighings need be made to centigrams only47 ,dissolve the ash in #ge cc4 of water1 warming gently1 and filter off the insoluble residue2 wash the filter by filling it at least three times with distilled water1 and allowing it to drain1 adding the washings to the main filtrate4 ,cool the filtrate to approximately the standard temperature of the laboratory1 and transfer it to a #bej cc4 measuring flask1 washing out the beaker thoroughly4 ,add distilled water of laboratory temperature until the lowest point of the meniscus is level with the graduation on the neck of the flask and remove any drops of water that may be on the neck above the graduation by means of a strip of filter paper2 make the solution thoroughly uniform by pouring it out into a dry beaker and back into the flask several times4 ,measure off #ej cc4 of the solution in a measuring flask1 or pipette1 either of which before use should1 unless they are dry on the inside1 be rinsed out with at least two small portions of the soda ash solution to displace any water4 ,if a flask is used1 fill it to the graduation with the soda ash solution and remove any liquid from the neck above the graduation with filter paper4 ,empty it into a beaker1 and wash out the small flask1 unless it is graduated for 6delivery61 using small quantities of water1 which are added to the liquid in the beaker4 ,a second #ej cc4 portion from the main solution should be measured off into a second beaker4 ,dilute the solutions in each beaker to #ajj cc41 add two drops of a solution of methyl orange 7,note #a7 and titrate for the alkali with the standard hydrochloric acid solution1 using the alkali solution to complete the titration as already prescribed4 ,from the volumes of acid and alkali employed1 corrected for burette errors and temperature changes1 and the data derived from the standardization1 calculate the percentage of alkali present1 assuming it all to be present as sodium carbonate 7,note #b74 ,note #a3 ,the hydrochloric acid sets free carbonic acid which is unstable and breaks down into water and carbon dioxide1 most of which escapes from the solution4 ,carbonic acid is a weak acid and1 as such1 does not yield a sufficient concentration of ,h ions to cause the indicator to change to a pink 7see page #cb74 ,the chemical changes involved may be summarized as follows3 #b;,h1 #b;,cl- #b,na1 ,,co#c-- -- #b,na1 #b;,cl- ,h#b;,,co#c -- ,h#b,o ,,co#b ,note #b3 ,a determination of the alkali present as hydroxide in soda ash may be determined by precipitating the carbonate by the addition of barium chloride1 removing the barium carbonate by filtration1 and titrating the alkali in the filtrate4 ,the caustic alkali may also be determined by first using phenolphthalein as an indicator1 which will show by its change from pink to colorless the point at which the caustic alkali has been neutralized and the carbonate has been converted to bicarbonate1 and then adding methyl orange and completing the titration4 ,the amount of acid necessary to change the methyl orange to pink is a measure of one half of the carbonate present4 ,the results of the double titration furnish the data necessary for the determination of the caustic alkali and of the carbonate in the sample4 ,,determination ,,of ,,the ,,acid ,,strength ,,of ,,oxalic ,,acid ,,procedure4--,weigh out two portions of the acid of about #a gram each4 ,dissolve these in #ej cc4 of warm water4 ,add two drops of phenolphthalein solution1 and run in alkali from the burette until the solution is pink2 add acid from the other burette until the pink is just destroyed1 and then add #j4c cc4 7not more7 in excess4 ,heat the solution to boiling for three minutes4 ,if the pink returns during the boiling1 discharge it with acid and again add #j4c cc4 in excess and repeat the boiling 7,note #a74 ,if the color does not then reappear1 add alkali until it does1 and a 6drop or two6 of acid in excess and boil again for one minute 7,note #b74 ,if no color reappears during this time1 complete the titration in the hot solution4 ,the end-point should be the faintest visible shade of color 7or its disappearance71 as the same difficulty would exist here as with methyl orange if an attempt were made to match shades of pink4 ,from the corrected volume of alkali required to react with the oxalic acid1 calculate the percentage of the crystallized acid 7,h#b;,c#b,o#d4b;,h#b,o7 in the sample 7,note #c74 ,note #a3 ,all commercial caustic soda such as that from which the standard solution was made contains some sodium carbonate4 ,this reacts with the oxalic acid1 setting free carbonic acid1 which1 in turn1 forms sodium bicarbonate with the remaining carbonate3 ,h#b;,,co#c ,na#b;,,co#c -- #b;,h,na,c,o#c4 ,this compound does not hydrolyze sufficiently to furnish enough ,,oh- ions to cause phenolphthalein to remain pink2 hence1 the color of the indicator is discharged in cold solutions at the point at which bicarbonate is formed4 ,if1 however1 the solution is heated to boiling1 the bicarbonate loses carbon dioxide and water1 and reverts to sodium carbonate1 which causes the indicator to become again pink3 #b;,h,na,c,o#c -- ,h#b,o ,,co#b ,na#b;,,co#c4 ,by adding successive portions of hydrochloric acid and boiling1 the carbonate is ultimately all brought into reaction4 ,the student should make sure that the difference in behavior of the two indicators1 methyl orange and phenolphthalein1 is understood4 ,note #b3 ,hydrochloric acid is volatilized from aqueous solutions1 except such as are very dilute4 ,if the directions in the procedure are strictly followed1 no loss of acid need be feared1 but the amount added in excess should not be greater than #j4c-j4d cc4 ,note #c3 ,attention has already been called to the fact that the color changes in the different indicators occur at varying concentrations of ,h or ,,oh- ions4 ,they do not indicate exact theoretical neutrality1 but a particular indicator always shows its color change at a particular concentration of ,h or ,,oh- ions4 ,the results of titration with a given indicator are1 therefore1 comparable4 ,as a matter of fact1 a small error is involved in the procedure as outlined above4 ,the comparison of the acid and alkali solutions was made1 using methyl orange as an indicator1 while the titration of the oxalic acid is made with the use of phenolphthalein4 ,for our present purposes the small error may be neglected but1 if time permits1 the student is recommended to standardize the alkali solution against one of the substances named in ,note #a1 page #da1 and also to ascertain the comparative value of the acid and alkali solutions1 using phenolphthalein as indicator throughout1 and conducting the titrations as described above4 ,this will insure complete accuracy4 ,,ii4 ,,oxidation ,,processes ,,general ,,discussion ,in the oxidation processes of volumetric analysis standard solutions of oxidizing agents and of reducing agents take the place of the acid and alkali solutions of the neutralization processes already studied4 ,just as an acid solution was the principal reagent in alkalimetry1 and the alkali solution used only to make certain of the end-point1 the solution of the oxidizing agent is the principal reagent for the titration of substances exerting a reducing action4 ,it is1 in general1 true that oxidizable substances are determined by 6direct6 titration1 while oxidizing substances are determined by 6indirect6 titration4 ,the important oxidizing agents employed in volumetric solutions are potassium bichromate1 potassium permangenate1 potassium ferricyanide1 iodine1 ferric chloride1 and sodium hypochlorite4 ,the important reducing agents which are used in the form of standard solutions are ferrous sulphate 7or ferrous ammonium sulphate71 oxalic acid1 sodium thiosulphate1 stannous chloride1 arsenious acid1 and potassium cyanide4 ,other reducing agents1 as sulphurous acid1 sulphureted hydrogen1 and zinc 7nascent hydrogen71 may take part in the processes1 but not as standard solutions4 ,the most important combinations among the foregoing are3 ,potassium bichromate and ferrous salts2 potassium permanganate and ferrous salts2 potassium permanganate and oxalic acid1 or its derivatives2 iodine and sodium thiosulphate2 hypochlorites and arsenious acid4 ,,bichromate ,,process ,,for ,,the ,,determination ,,of ,,iron ,ferrous salts may be promptly and completely oxidized to ferric salts1 even in cold solution1 by the addition of potassium bichromate1 provided sufficient acid is present to hold in solution the ferric and chromic compounds which are formed4 ,the acid may be either hydrochloric or sulphuric1 but the former is usually preferred1 since it is by far the best solvent for iron and its compounds4 ,the reaction in the presence of hydrochloric acid is as follows3 #f;,fe,cl#b ,k#b;,cr#b,o#g #ad;,h,cl -- #f;,fe,cl#c #b;,cr,cl#c #b,k,cl #g;,h#b,o4 ,,normal ,,solutions ,,of ,,oxidizing ,,or ,,reducing ,,agents ,it will be recalled that the system of normal solutions is based upon the equivalence of the reagents which they contain to #h grams of oxygen or #a gram of hydrogen4 ,a normal solution of an oxidizing agent should1 therefore1 contain that amount per liter which is equivalent in oxidizing power to #h grams of oxygen2 a normal reducing solution must be equivalent in reducing power to #a gram of hydrogen4 ,in order to determine what the amount per liter will be it is necessary to know how the reagents enter into reaction4 ,the two solutions to be employed in the process under consideration are those of potassium bichromate and ferrous sulphate4 ,the reaction between them1 in the presence of an excess of sulphuric acid1 may be expressed as follows3 #f;,fe,s,o#d ,k#b;,cr#b,o#g #g;,h#b,,so#d -- #c;,fe#b7,,so#d7c ,k#b,,so#d ,cr#b7,,so#d7c #g;,h#b,o4 ,if the compounds of iron and chromium1 with which alone we are now concerned1 be written in such a way as to show the oxides of these elements in each1 they would appear as follows3 ,on the left-hand side of the equation #f7;,fe,o4,,so#c7 and ,k#b,o4#b;,cr,o#c2 on the right-hand side1 #c7;,fe#b,o#c4c,,so#c7 and ,cr#b,o#c4c,,so#c4 ,a careful inspection shows that there are three less oxygen atoms associated with chromium atoms on the right-hand side of the equation than on the left-hand1 but there are three more oxygen atoms associated with iron atoms on the right than on the left4 ,in other words1 a molecule of potassium bichromate has given up three atoms of oxygen for oxidation purposes2 i4e41 a molecular weight in grams of the bichromate 7#bid4b7 will furnish #c ,x #af or #dh grams of oxygen for oxidation purposes4 ,as this #dh grams is six times #h grams1 the basis of the system1 the normal solution of potassium bichromate should contain per liter one sixth of #bid4b grams or #di4jc grams4 ,a further inspection of the dissected compounds above shows that six molecules of ,fe,o4,,so#c were required to react with the three atoms of oxygen from the bichromate4 ,from the two equations #c;,h#b #c,o -- #c;,h#b,o #f7;,fe,o4,,so#c7 #c,o -- #c7;,fe#b,o#c4c,,so#c7 it is plain that one molecule of ferrous sulphate is equivalent to one atom of hydrogen in reducing power2 therefore one molecular weight in grams of ferrous sulphate 7#aea4i7 is equivalent to #a gram of hydrogen4 ,since the ferrous sulphate crystalline form has the formula ,fe,s,o#d4g;,h#b,o1 a normal reducing solution of this crystalline salt should contain #bgg4i grams per liter4 ,,preparation ,,of ,,solutions 6,approximate ,strength #j4a ,n6 ,it is possible to purify commercial potassium bichromate by recrystallization from hot water4 ,it must then be dried and cautiously heated to fusion to expel the last traces of moisture1 but not sufficiently high to expel any oxygen4 ,the pure salt thus prepared1 may be weighed out directly1 dissolved1 and the solution diluted in a graduated flask to a definite volume4 ,in this case no standardization is made1 as the normal value can be calculated directly4 ,it is1 however1 more generally customary to standardize a solution of the commercial salt by comparison with some substance of definite composition1 as described below4 ,,procedure4--,pulverize about #e grams of potassium bichromate of good quality4 ,dissolve the bichromate in distilled water1 transfer the solution to a liter bottle1 and dilute to approximately #ajjj cc4 ,shake thoroughly until the solution is uniform4 ,to prepare the solution of the reducing agent1 pulverize about #bh grams of ferrous sulphate 7,fe,s,o#d4g;,h#b,o7 or about #dj grams of ferrous ammonium sulphate 7,fe,s,o#d47,,nh#d7b,,so#d4f;,h#b,o7 and dissolve in distilled water containing #e cc4 of concentrated sulphuric acid4 ,transfer the solution to a liter bottle1 add #e cc4 concentrated sulphuric acid1 make up to about #ajjj cc4 and shake vigorously to insure uniformity4 ,,indicator ,,solution ,no indicator is known which1 like methyl orange1 can be used within the solution1 to show when the oxidation process is complete4 ,instead1 an outside indicator solution is employed to which drops of the titrated solution are transferred for testing4 ,the reagent used is potassium ferricyanide1 which produces a blue precipitate 7or color7 with ferrous compounds as long as there are unoxidized ferrous ions in the titrated solution4 ,drops of the indicator solution are placed upon a glazed porcelain tile1 or upon white cardboard which has been coated with paraffin to render it waterproof1 and drops of the titrated solution are transferred to the indicator on the end of a stirring rod4 ,when the oxidation is nearly completed only very small amounts of the ferrous compounds remain unoxidized and the reaction with the indicator is no longer instantaneous4 ,it is necessary to allow a brief time to elapse before determining that no blue color is formed4 ,thirty seconds is a sufficient interval1 and should be adopted throughout the analytical procedure4 ,if left too long1 the combined effect of light and dust from the air will cause a reduction of the ferric compounds already formed and a resultant blue will appear which misleads the observer with respect to the true end-point4 ,the indicator solution must be highly diluted1 otherwise its own color interferes with accurate observation4 ,prepare a fresh solution1 as needed each day1 by dissolving a crystal of potassium ferricyanide about the size of a pin's head in #be cc4 of distilled water4 ,the salt should be carefully tested with ferric chloride for the presence of ferrocyanides1 which give a blue color with ferric salts4 ,in case of need1 the ferricyanide can be purified by adding to its solution a little bromine water and recrystallizing the compound4 ,,comparison ,,of ,,oxidizing ,,and ,,reducing ,,solutions ,,procedure4--,fill one burette with each of the solutions1 observing the general procedure with respect to cleaning and rinsing already prescribed4 ,the bichromate solution is preferably to be placed in a glass-stoppered burette4 ,run out from a burette into a beaker of about #cjj cc4 capacity nearly #dj cc4 of the ferrous solution1 add #ae cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab7 and #aej cc4 of water and run in the bichromate solution from another burette4 ,since both solutions are approximately tenth-normal1 #ce cc4 of the bichromate solution may be added without testing4 ,test at that point by removing a very small drop of the iron solution on the end of a stirring rod1 mixing it with a drop of indicator on the tile 7,note #a74 ,if a blue precipitate appears at once1 #j4e cc4 of the bichromate solution may be added before testing again4 ,the stirring rod which has touched the indicator should be dipped in distilled water before returning it to the iron solution4 ,as soon as the blue appears to be less intense1 add the bichromate solution in small portions1 finally a single drop at a time1 until the point is reached at which no blue color appears after the lapse of thirty seconds from the time of mixing solution and indicator4 ,at the close of the titration a large drop of the iron solution should be taken for the test4 ,to determine the end-point beyond any question1 as soon as the thirty seconds have elapsed remove another drop of the solution of the same size as that last taken and mix it with the indicator1 placing it beside the last previous test4 ,if this last previous test shows a blue tint in comparison with the fresh mixture1 the end-point has not been reached2 if no difference can be noted the reaction is complete4 ,should the end-point be overstepped1 a little more of the ferrous solution may be added and the end-point definitely fixed4 ,from the volumes of the solutions used1 after applying corrections for burette readings1 and1 if need be1 for the temperature of solutions1 calculate the value of the ferrous solution in terms of the oxidizing solution4 ,note #a3 ,the accuracy of the work may be much impaired by the removal of unnecessarily large quantities of solution for the tests4 ,at the beginning of the titration1 while much ferrous iron is still present1 the end of the stirring rod need only be moist with the solution2 but at the close of the titration drops of considerable size may properly be taken for the final tests4 ,the stirring rod should be washed to prevent transfer of indicator to the main solution4 ,this cautious removal of solution does not seriously affect the accuracy of the determination1 as it will be noted that the volume of the titrated solution is about #bjj cc4 and the portions removed are very small4 ,moreover1 if the procedure is followed as prescribed1 the concentration of unoxidized iron decreases very rapidly as the titration is carried out so that when the final tests are made1 though large drops may be taken1 the amount of ferrous iron is not sufficient to produce any appreciable error in results4 ,if the end-point is determined as prescribed1 it can be as accurately fixed as that of other methods2 and if a ferrous solution is at hand1 the titration need consume hardly more time than that of the permanganate process to be described later on4 ,,standardization ,,of ,,potassium ,,bichromate ,,solutions 6,selection of a ,standard6 ,a substance which will serve satisfactorily as a standard for oxidizing solutions must possess certain specific properties3 ,it must be of accurately known composition and definite in its behavior as a reducing agent1 and it must be permanent against oxidation in the air1 at least for considerable periods4 ,such standards may take the form of pure crystalline salts1 such as ferrous ammonium sulphate1 or may be in the form of iron wire or an iron ore of known iron content4 ,it is not necessary that the standard should be of #ajj per cent purity1 provided the content of the active reducing agent is known and no interfering substances are present4 ,the two substances most commonly used as standards for a bichromate solution are ferrous ammonium sulphate and iron wire4 ,a standard wire is to be purchased in the market which answers the purpose well1 and its iron content may be determined for each lot purchased by a number of gravimetric determinations4 ,it may best be preserved in jars containing calcium chloride1 but this must not be allowed to come into contact with the wire4 ,it should1 however1 even then be examined carefully for rust before use4 ,if pure ferrous ammonium sulphate is used as the standard1 clear crystals only should be selected4 ,it is perhaps even better to determine by gravimetric methods once for all the iron content of a large commercial sample which has been ground and well mixed4 ,this salt is permanent over long periods if kept in stoppered containers4 ,,standardization ,,procedure4--,weigh out two portions of iron wire of about #j4bd-j4bf gram each1 examining the wire carefully for rust4 ,it should be handled and wiped with filter paper 7not touched by the fingers71 should be weighed on a watch-glass1 and be bent in such a way as not to interfere with the movement of the balance4 ,place #cj cc4 of hydrochloric acid 7sp4 gr4 #a4ab7 in each of two #cjj cc4 ,erlenmeyer flasks1 cover them with watch-glasses1 and bring the acid just to boiling4 ,remove them from the flame and drop in the portions of wire1 taking great care to avoid loss of liquid during solution4 ,boil for two or three minutes1 keeping the flasks covered 7,note #a71 then wash the sides of the flasks and the watch-glass with a little water and add stannous chloride solution to the hot liquid 6from a dropper6 until the solution is colorless1 but avoid more than a drop or two in excess 7,note #b74 ,dilute with #aej cc4 of water and cool 6completely64 ,when cold1 add rapidly about #cj cc4 of mercuric chloride solution4 ,allow the solutions to stand about three minutes and then titrate without further delay 7,note #c71 add about #ce cc4 of the standard solution at once and finish the titration as prescribed above1 making use of the ferrous solution if the end-point should be passed4 ,from the corrected volumes of the bichromate solution required to oxidize the iron actually know to be present in the wire1 calculate the relation of the standard solution to the normal4 ,repeat the standardization until the results are concordant within at least two parts in one thousand4 ,note #a3 ,the hydrochloric acid is added to the ferrous solution to insure the presence of at least sufficient free acid for the titration1 as required by the equation on page #dh4 ,the solution of the wire in hot acid and the short boiling insure the removal of compounds of hydrogen and carbon which are formed from the small amount of carbon in the iron4 ,these might be acted upon by the bichromate if not expelled4 ,note #b3 ,it is plain that all the iron must be reduced to the ferrous condition before the titration begins1 as some oxidation may have occurred from the oxygen of the air during solution4 ,it is also evident that any excess of the agent used to reduce the iron must be removed2 otherwise it will react with the bichromate added later4 ,the reagents available for the reduction of iron are stannous chloride1 sulphurous acid1 sulphureted hydrogen1 and zinc2 of these stannous chloride acts most readily1 the completion of the reaction is most easily noted1 and the excess of the reagent is most readily removed4 ,the latter object is accomplished by oxidation to stannic chloride by means of mercuric chloride added in excess1 as the mercuric salts have no effect upon ferrous iron or the bichromate4 ,the reactions involved are3 #b;,fe,cl#c ,sn,cl#b -- #b;,fe,cl#b ,sn,cl#d ,sn,cl#b #b;,hg,cl#b -- ,sn,cl#d #b;,hg,cl ,the mercurous chloride is precipitated4 ,it is essential that the solution should be cold and that the stannous chloride should not be present in great excess1 otherwise a secondary reaction takes place1 resulting in the reduction of the mercurous chloride to metallic mercury3 ,sn,cl#b #b;,hg,cl -- ,sn,cl#d #b;,hg4 ,the occurrence of this secondary reaction is indicated by the darkening of the precipitate2 and1 since potassium bichromate oxidizes this mercury slowly1 solutions in which it has been precipitated are worthless as iron determinations4 ,note #c3 ,the solution should be allowed to stand about three minutes after the addition of mercuric chloride to permit the complete deposition of mercurous chloride4 ,it should then be titrated without delay to avoid possible reoxidation of the iron by the oxygen of the air4 ,,determination ,,of ,,iron ,,in ,,limonite ,,procedure4--,grind the mineral 7,note #a7 to a fine powder4 ,weigh out accurately two portions of about #j4e gram 7,note #b7 into porcelain crucibles2 heat these crucibles to dull redness for ten minutes1 allow them to cool1 and place them1 with their contents1 in beakers containing #cj cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab74 ,heat at a temperature just below boiling until the undissolved residue is white or until solvent action has ceased4 ,if the residue is white1 or known to be free from iron1 it may be neglected and need not be removed by filtration4 ,if a dark residue remains1 collect it on a filter1 wash free from hydrochloric acid1 and ignite the filter in a platinum crucible 7,note #c74 ,mix the ash with five times its weight of sodium carbonate and heat to fusion2 cool1 and disintegrate the fused mass with boiling water in the crucible4 ,unite this solution and precipitate 7if any7 with the acid solution1 taking care to avoid loss by effervescence4 ,wash out the crucible1 heat the acid solution to boiling1 add stannous chloride solution until it is colorless1 avoiding a large excess 7,note #d72 cool1 and when 6cold61 add #dj cc4 of mercuric chloride solution1 dilute to #bjj cc41 and proceed with the titration as already described4 ,from the standardization data already obtained1 and the known weight of the sample1 calculate the percentage of iron 7,fe7 in the limonite4 ,note #a3 ,limonite is selected as a representative of iron ores in general4 ,it is a native1 hydrated oxide of iron4 ,it frequently occurs in or near peat beds and contains more or less organic matter which1 if brought into solution1 would be acted upon by the potassium bichromate4 ,this organic matter is destroyed by roasting4 ,since a high temperature tends to lessen the solubility of ferric oxide1 the heat should not be raised above low redness4 ,note #b3 ,it is sometimes advantageous to dissolve a large portion--say #e grams--and to take one tenth of it for titration4 ,the sample will then represent more closely the average value of the ore4 ,note #c3 ,a platinum crucible may be used for the roasting of the limonite and must be used for the fusion of the residue4 ,when used1 it must not be allowed to remain in the acid solution of ferric chloride for any length of time1 since the platinum is attacked and dissolved1 and the platinic chloride is later reduced by the stannous chloride1 and in the reduced condition reacts with the bichromate1 thus introducing an error4 ,it should also be noted that copper and antimony interfere with the determination of iron by the bichromate process4 ,note #d3 ,the quantity of stannous chloride required for the reduction of the iron in the limonite will be much larger than that added to the solution of iron wire1 in which the iron was mainly already in the ferrous condition4 ,it should1 however1 be added from a dropper to avoid an unnecessary excess4 ,,determination ,,of ,,chromium ,,in ,,chrome ,,iron ,,ore ,,procedure4--,grind the chrome iron ore 7,note #a7 in an agate mortar until no grit is perceptible under the pestle4 ,weigh out two portions of #j4e gram each into iron crucibles which have been scoured inside until bright 7,note #b74 ,weigh out on a watch-glass 7,note #c71 using the rough balances1 #e grams of dry sodium peroxide for each portion1 and pour about three quarters of the peroxide upon the ore4 ,mix ore and flux by thorough stirring with a dry glass rod4 ,then cover the mixture with the remainder of the peroxide4 ,place the crucible on a triangle and raise the temperature 6slowly6 to the melting point of the flux1 using a low flame1 and holding the lamp in the hand 7,note #d74 ,maintain the fusion for five minutes1 and stir constantly with a stout iron wire1 but do not raise the temperature above moderate redness 7,notes #e and #f74 ,allow the crucible to cool until it can be comfortably handled 7,note #g7 and then place it in a #cjj cc4 beaker1 and cover it with distilled water 7,note #h74 ,the beaker must be carefully covered to avoid loss during the disintegration of the fused mass4 ,when the evolution of gas ceases1 rinse off and remove the crucible2 then heat the solution 6while still alkaline6 to boiling for fifteen minutes4 ,allow the liquid to cool for a few minutes2 then acidify with dilute sulphuric acid 7#a3e71 adding #aj cc4 in excess of the amount necessary to dissolve the ferric hydroxide 7,note #i74 ,dilute to #bjj cc41 cool1 add from a burette an excess of a standard ferrous solution1 and titrate for the excess with a standard solution of potassium bichromate1 using the outside indicator 7,note #aj74 ,from the corrected volumes of the two standard solutions1 and their relations to normal solutions1 calculate the percentage of chromium in the ore4 ,note #a3 ,chrome iron ore is essentially a ferrous chromite1 or combination of ,fe,o and ,cr#b,o#c4 ,it must be reduced to a state of fine subdivision to ensure a prompt reaction with the flux4 ,note #b3 ,the scouring of the iron crucible is rendered much easier if it is first heated to bright redness and plunged into cold water4 ,in this process oily matter is burned off and adhering scale is caused to chip off when the hot crucible contracts rapidly in the cold water4 ,note #c3 ,sodium peroxide must be kept off of balance pans and should not be weighed out on paper1 as is the usual practice in the rough weighing of chemicals4 ,if paper to which the peroxide is adhering is exposed to moist air it is likely to take fire as a result of the absorption of moisture1 and consequent evolution of heat and liberation of oxygen4 ,note #d3 ,the lamp should never be allowed to remain under the crucible1 as this will raise the temperature to a point at which the crucible itself is rapidly attacked by the flux and burned through4 ,note #e3 ,the sodium peroxide acts as both a flux and an oxidizing agent4 ,the chromic oxide is dissolved by the flux and oxidized to chromic anhydride 7,cr,o#c7 which combines with the alkali to form sodium chromate4 ,the iron is oxidized to ferric oxide4 ,note #f3 ,the sodium peroxide cannot be used in porcelain1 platinum1 or silver crucibles4 ,it attacks iron and nickel as well2 but crucibles made from these metals may be used if care is exercised to keep the temperature as low as possible4 ,preference is here given to iron crucibles1 because the resulting ferric hydroxide is more readily brought into solution than the nickelic oxide from a nickel crucible4 ,the peroxide must be dry1 and must be protected from any admixture of dust1 paper1 or of organic matter of any kind1 otherwise explosions may ensue4 ,note #g3 ,when an iron crucible is employed it is desirable to allow the fusion to become nearly cold before it is placed in water1 otherwise scales of magnetic iron oxide may separate from the crucible1 which by slowly dissolving in acid form ferrous sulphate1 which reduces the chromate4 ,note #h3 ,upon treatment with water the chromate passes into solution1 the ferric hydroxide remains undissolved1 and the excess of peroxide is decomposed with the evolution of oxygen4 ,the subsequent boiling insures the complete decomposition of the peroxide4 ,unless this is complete1 hydrogen peroxide is formed when the solution is acidified1 and this reacts with the bichromate1 reducing it and introducing a serious error4 ,note #i3 ,the addition of the sulphuric acid converts the sodium chromate to bichromate1 which behaves exactly like potassium bichromate in acid solution4 ,note #aj3 ,if a standard solution of a ferrous salt is not at hand1 a weight of iron wire somewhat in excess of the amount which would be required if the chromite were pure ,fe,o4,cr#b,o#c may be weighed out and dissolved in sulphuric acid2 after reduction of all the iron by stannous chloride and the addition of mercuric chloride1 this solution may be poured into the chromate solution and the excess of iron determined by titration with standard bichromate solution4 ,,permanganate ,,process ,,for ,,the ,,determination ,,of ,,iron ,potassium permanganate oxidizes ferrous salts in cold1 acid solution promptly and completely to the ferric condition1 while in hot acid solution it also enters into a definite reaction with oxalic acid1 by which the latter is oxidized to carbon dioxide and water4 ,the reactions involved are these3 #aj;,fe,s,o#d #b,k,mn,o#d #h;,h#b,s#d -- #e;,fe#b7,,so#d7c ,k#b,,so#d #b,mn,s,o#d #h;,h#b,o #e;,c#b;,h#b,o#d7b;,h#b,o7 #b,k,mn,o#d #c;,h#b,,so#d -- ,k#b,,so#d #b,mn,s,o#d #aj;,,co#b #a ,h#b,o4 ,these are the fundamental reactions upon which the extensive use of potassium permanganate depends2 but besides iron and oxalic acid the permanganate enters into reaction with antimony1 tin1 copper1 mercury1 and manganese 7the latter only in neutral solution71 by which these metals are changed from a lower to a higher state of oxidation2 and it also reacts with sulphurous acid1 sulphureted hydrogen1 nitrous acid1 ferrocyanides1 and most soluble organic bodies4 ,it should be noted1 however1 that very few of these organic compounds react quantitatively with the permanganate1 as is the case with oxalic acid and the oxalates4 ,potassium permanganate is acted upon by hydrochloric acid2 the action is rapid in hot or concentrated solution 7particularly in the presence of iron salts1 which appear to act as catalyzers1 increasing the velocity of the reaction71 but slow in cold1 dilute solutions4 ,however1 the greater solubility of iron compounds in hydrochloric acid makes it desirable to use this acid as a solvent1 and experiments made with this end in view have shown that in cold1 dilute hydrochloric acid solution1 to which considerable quantities of manganous sulphate and an excess of phosphoric acid have been added1 it is possible to obtain satisfactory results4 ,it is also possible to replace the hydrochloric acid by evaporating the solutions with an excess of sulphuric acid until the latter fumes4 ,this procedure is somewhat more time-consuming1 but the end-point of the permanganate titration is more permanent4 ,both procedures are described below4 ,potassium permanganate has an intense coloring power1 and since the solution resulting from the oxidation of the iron and the reduction of the permanganate is colorless1 the latter becomes its own indicator4 ,the slightest excess is indicated with great accuracy by the pink color of the solution4 ,,preparation ,,of ,a ,,standard ,,solution 6,approximate ,strength #j4a ,n6 ,a study of the reactions given above which represent the oxidation of ferrous compounds by potassium permanganate1 shows that there are #b molecules of ,k,mn,o#d and #aj molecules of ,fe,s,o#d on the left-hand side1 and #b molecules of ,mn,s,o#d and #e molecules of ,fe#b7,,so#d7e on the right-hand side4 ,considering only these compounds1 and writing the formulas in such a way as to show the oxides of the elements in each1 the equation becomes3 ,k#b,o4,mn#b,o#g #aj7;,fe,o4,,so#c7 -- ,k#b,o4,,so#c #b7,mn,o4,,so#c7 #e7;,fe#b,o#c4c,,so#c74 ,from this it appears that two molecules of ,k,mn,o#d 7or #caf4j grams7 have given up five atoms 7or #hj grams7 of oxygen to oxidize the ferrous compound4 ,since #h grams of oxygen is the basis of normal oxidizing solutions and #hj grams of oxygen are supplied by #caf4j grams of ,k,mn,o#d1 the normal solution of the permanganate should contain1 per liter1 #caf4j_/aj grams1 or #ca4fj grams 7,note #a74 ,the preparation of an approximately tenth-normal solution of the reagent may be carried out as follows3 ,,procedure4--,dissolve about #c4be grams of potassium permanganate crystals in approximately #ajjj cc4 of distilled water in a large beaker1 or casserole4 ,heat slowly and when the crystals have dissolved1 boil the solution for #aj-ae minutes4 ,cover the solution with a watch-glass2 allow it to stand until cool1 or preferably over night4 ,filter the solution through a layer of asbestos4 ,transfer the filtrate to a liter bottle and mix thoroughly 7,note #b74 ,note #a3 ,the reactions given on page #fa are those which take place in the presence of an excess of acid4 ,in neutral solutions the reduction of the permanganate is less complete1 and1 under these conditions1 two gram-molecular weights of ,k,mn,o#d will furnish only #dh grams of oxygen4 ,a normal solution for use under these conditions should1 therefore1 contain #caf4j_/f grams1 or #eb4ff grams4 ,note #b3 ,potassium permanganate solutions are not usually stable for long periods1 and change more rapidly when first prepared than after standing some days4 ,this change is probably caused by interaction with the organic matter contained in all distilled water1 except that redistilled from an alkaline permanganate solution4 ,the solutions should be protected from light and heat as far as possible1 since both induce decomposition with a deposition of manganese dioxide1 and it has been shown that decomposition proceeds with considerable rapidity1 with the evolution of oxygen1 after the dioxide has begun to form4 ,as commercial samples of the permanganate are likely to be contaminated by the dioxide1 it is advisable to boil and filter solutions through asbestos before standardization1 as prescribed above4 ,such solutions are relatively stable4 ,,comparison ,,of ,,permanganate ,,and ,,ferrous ,,solutions ,,procedure4--,fill a glass-stoppered burette with the permanganate solution1 observing the usual precautions1 and fill a second burette with the ferrous sulphate solution prepared for use with the potassium bichromate4 ,the permanganate solution cannot be used in burettes with rubber tips1 as a reduction takes place upon contact with the rubber4 ,the solution has so deep a color that the lower line of the meniscus cannot be detected2 readings must therefore be made from the upper edge4 ,run out into a beaker about #dj cc4 of the ferrous solution1 dilute to about #ajj cc41 add #aj cc4 of dilute sulphuric acid1 and run in the permanganate solution to a slight permanent pink4 ,repeat1 until the ratio of the two solutions is satisfactorily established4 ,,standardization ,,of ,a ,,potassium ,,permanganate ,,solution 6,selection of a ,standard6 ,commercial potassium permanganate is rarely sufficiently pure to admit of its direct weighing as a standard4 ,on this account1 and because of the uncertainties as to the permanence of its solutions1 it is advisable to standardize them against substances of known value4 ,those in most common use are iron wire1 ferrous ammonium sulphate1 sodium oxalate1 oxalic acid1 and some other derivatives of oxalic acid4 ,with the exception of sodium oxalate1 these all contain water of crystallization which may be lost on standing4 ,they should1 therefore1 be freshly prepared1 and with great care4 ,at present1 sodium oxalate is considered to be one of the most satisfactory standards4 6,method ,a6 6,iron ,standards6 ,the standardization processes employed when iron or its compounds are selected as standards differ from those applicable in connection with oxalate standards4 ,the procedure which immediately follows is that in use with iron standards4 ,as in the case of the bichromate process1 it is necessary to reduce the iron completely to the ferrous condition before titration4 ,the reducing agents available are zinc1 sulphurous acid1 or sulphureted hydrogen4 ,stannous chloride may also be used when the titration is made in the presence of hydrochloric acid4 ,since the excess of both the gaseous reducing agents can only be expelled by boiling1 with consequent uncertainty regarding both the removal of the excess and the reoxidation of the iron1 zinc or stannous chlorides are the most satisfactory agents4 ,for prompt and complete reduction it is essential that the iron solution should be brought into ultimate contact with the zinc4 ,this is brought about by the use of a modified ,jones reductor1 as shown in ,figure #a4 ,this reductor is a standard apparatus and is used in other quantitative processes4 ,illustration3 ,fig4 #a ,the tube ,a has an inside diameter of #ah mm4 and is #cjj mm4 long2 the small tube has an inside diameter of #f mm4 and extends #ajj mm4 below the stopcock4 ,at the base of the tube ,a are placed some pieces of broken glass or porcelain1 covered by a plug of glass wool about #h mm4 thick1 and upon this is placed a thin layer of asbestos1 such as is used for ,gooch filters1 #a mm4 thick4 ,the tube is then filled with the amalgamated zinc 7,note #a7 to within #ej mm4 of the top1 and on the zinc is placed a plug of glass wool4 ,if the top of the tube is not already shaped like the mouth of a thistle-tube 7,b71 a #fj mm4 funnel is fitted into the tube with a rubber stopper and the reductor is connected with a suction bottle1 ,f4 ,the bottle ,d is a safety bottle to prevent contamination of the solution by water from the pump4 ,after preparation for use1 or when left standing1 the tube ,a should be filled with water1 to prevent clogging of the zinc4 ,note #a3 ,the use of fine zinc in the reductor is not necessary and tends to clog the tube4 ,particles which will pass a #aj-mesh sieve1 but are retained by one of #bj meshes to the inch1 are most satisfactory4 ,the zinc can be amalgamated by stirring or shaking it in a mixture of #be cc4 of normal mercuric chloride solution1 #be cc4 of hydrochloric acid 7sp4 gr4 #a4ab7 and #bej cc4 of water for two minutes4 ,the solution should then be poured off and the zinc thoroughly washed4 ,it is then ready for bottling and preservation under water4 ,a small quantity of glass wool is placed in the neck of the funnel to hold back foreign material when the reductor is in use4 ,,standardization ,,procedure4--,weigh out into ,erlenmeyer flasks two portions of iron wire of about #j4be gram each4 ,dissolve these in hot dilute sulphuric acid 7#e cc4 of concentrated acid and #ajj cc4 of water71 using a covered flask to avoid loss by spattering4 ,boil the solution for two or three minutes after the iron has dissolved to remove any volatile hydrocarbons4 ,meanwhile prepare the reductor for use as follows3 ,connect the vacuum bottle with the suction pump and pour into the funnel at the top warm1 dilute sulphuric acid1 prepared by adding #e cc4 of concentrated sulphuric acid to #ajj cc4 of distilled water4 ,see that the stopcock 7,c7 is open far enough to allow the acid to run through slowly4 ,continue to pour in acid until #bjj cc4 have passed through1 then close the stopcock 6while a small quantity of liquid is still left in the funnel64 ,discard the filtrate1 and again pass through #ajj cc4 of the warm1 dilute acid4 ,test this with the permanganate solution4 ,a single drop should color it permanently2 if it does not1 repeat the washing1 until assured that the zinc is not contaminated with appreciable quantities of reducing substances4 ,be sure that no air enters the reductor 7,note #a74 ,pour the iron solution while hot 7but not boiling7 through the reductor at a rate not exceeding #ej cc4 per minute 7,notes #b and #c74 ,wash out the beaker with dilute sulphuric acid1 and follow the iron solution without interruption with #age cc4 of the warm acid and finally with #ge cc4 of distilled water1 leaving the funnel partially filled4 ,remove the filter bottle and cool the solution quickly under the water tap 7,note #d71 avoiding unnecessary exposure to the oxygen of the air4 ,add #aj cc4 of dilute sulphuric acid and titrate to a faint pink with the permanganate solution1 adding it directly to the contents of the vacuum flask4 ,should the end-point be overstepped1 the ferrous sulphate solution may be added4 ,from the volume of the solution required to oxidize the iron in the wire1 calculate the relation to the normal of the permanganate solution4 ,the duplicate results should be concordant within two parts in one thousand4 ,note #a3 ,the funnel of the reductor must never be allowed to empty4 ,if it is left partially filled with water the reductor is ready for subsequent use after a very little washing2 but a preliminary test is always necessary to safeguard against error4 ,if more than a small drop of permanganate solution is required to color #ajj cc4 of the dilute acid after the reductor is well washed1 an allowance must be made for the iron in the zinc4 6,great care6 must be used to prevent the access of air to the reductor after it has been washed out ready for use4 ,if air enters1 hydrogen peroxide forms1 which reacts with the permanganate1 and the results are worthless4 ,note #b3 ,the iron is reduced to the ferrous condition by contact with the zinc4 ,the active agent may be considered to be 6nascent6 hydrogen1 and it must be borne in mind that the visible bubbles are produced by molecular hydrogen1 which is without appreciable effect upon ferric iron4 ,the rate at which the iron solution passes through the zinc should not exceed that prescribed1 but the rate may be increased somewhat when the wash-water is added4 ,it is well to allow the iron solution to run nearly1 but not entirely1 out of the funnel before the wash-water is added4 ,if it is necessary to interrupt the process1 the complete emptying of the funnel can always be avoided by closing the stopcock4 ,it is also possible to reduce the iron by treatment with zinc in a flask from which air is excluded4 ,the zinc must be present in excess of the quantity necessary to reduce the iron and is finally completely dissolved4 ,this method is1 however1 less convenient and more tedious than the use of the reductor4 ,note #c3 ,the dilute sulphuric acid for washing must be warmed ready for use before the reduction of the iron begins1 and it is of the first importance that the volume of acid and of wash-water should be measured1 and the volume used should always be the same in the standardizations and all subsequent analyses4 ,note #d3 ,the end-point is more permanent in cold than hot solutions1 possibly because of a slight action of the permanganate upon the manganous sulphate formed during titration4 ,if the solution turns brown1 it is an evidence of insufficient acid1 and more should be immediately added4 ,the results are likely to be less accurate in this case1 however1 as a consequence of secondary reactions between the ferrous iron and the manganese dioxide thrown down4 ,it is wiser to discard such results and repeat the process4 ,note #e3 ,the potassium permanganate may1 of course1 be diluted and brought to an exactly #j4a ,n solution from the data here obtained4 ,the percentage of iron in the iron wire must be taken into account in all calculations4 6,method ,b6 6,oxalate ,standards6 ,,procedure4--,weigh out two portions of pure sodium oxalate of #j4be-j4c gram each into beakers of about #fjj cc4 capacity4 ,add about #djj cc4 of boiling water and #bj cc4 of manganous sulphate solution 7,note #a74 ,when the solution of the oxalate is complete1 heat the liquid1 if necessary1 until near its boiling point 7#gj-ij;,c47 and run in the standard permanganate solution drop by drop from a burette1 stirring constantly until an end-point is reached 7,note #b74 ,make a blank test with #bj cc4 of manganous sulphate solution and a volume of distilled water equal to that of the titrated solution to determine the volume of the permanganate solution required to produce a very slight pink4 ,deduct this volume from the amount of permanganate solution used in the titration4 ,from the data obtained1 calculate the relation of the permanganate solution to the normal4 ,the reaction involved is3 #e,na#b;,c#b,o#d #b,k,mn,o#d #h;,h#b,,so#d -- #e,na#b,,so#d ,k#b,,so#d #b,mn,s,o#d #aj;,,co#b #h;,h#b,o ,note #a3 ,the manganous sulphate titrating solution is made by dissolving #bj grams of ,mn,s,o#d in #bjj cubic centimeters of water and adding #dj cc4 of concentrated sulphuric acid 7sp4 gr4 #a4hd7 and #dj cc4 or phosphoric acid 7#he74 ,note #b3 ,the reaction between oxalates and permanganates takes place quantitatively only in hot acid solutions4 ,the temperatures must not fall below #gj;,c4 ,,determination ,,of ,,iron ,,in ,,limonite 6,method ,a6 ,the procedures1 as here prescribed1 are applicable to iron ores in general1 provided these ores contain no constituents which are reduced by zinc or stannous chloride and reoxidized by permanganates4 ,many iron ores contain titanium1 and this element among others does interfere with the determination of iron by the process described4 ,if1 however1 the solutions of such ores are treated with sulphureted hydrogen or sulphurous acid1 instead of zinc or stannous chloride to reduce the iron1 and the excess reducing agent removed by boiling1 an accurate determination of the iron can be made4 ,,procedure4--,grind the mineral to a fine powder4 ,weigh out two portions of about #j4e gram each into small porcelain crucibles4 ,roast the ore at dull redness for ten minutes 7,note #a71 allow the crucibles to cool1 and place them and their contents in casseroles containing #cj cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab74 ,proceed with the solution of the ore1 and the treatment of the residue1 if necessary1 exactly as described for the bichromate process on page #ef4 ,when solution is complete1 add #f cc4 of concentrated sulphuric acid to each casserole1 and evaporate on the steam bath until the solution is nearly colorless 7,note #b74 ,cover the casseroles and heat over the flame of the burner1 holding the casserole in the hand and rotating it slowly to hasten evaporation and prevent spattering1 until the heavy white fumes of sulphuric anhydride are freely evolved 7,note #c74 ,cool the casseroles1 add #ajj cc4 of water 7measured71 and boil gently until the ferric sulphate is dissolved2 pour the warm solution through the reductor which has been previously washed2 proceed as described under standardization1 taking pains to use the same volume and strength of acid and the same volume of wash-water as there prescribed1 and titrate with the permanganate solution in the reductor flask1 using the ferrous sulphate solution if the end-point should be overstepped4 ,from the corrected volume of permanganate solution used1 calculate the percentage of iron 7,fe7 in the limonite4 ,note #a3 ,the preliminary roasting is usually necessary because1 even though the sulphuric acid would subsequently char the carbonaceous matter1 certain nitrogenous bodies are not thereby rendered insoluble in the acid1 and would be oxidized by the permanganate4 ,note #b3 ,the temperature of the steam bath is not sufficient to volatilize sulphuric acid4 ,solutions may1 therefore1 be left to evaporate overnight without danger of evaporation to dryness4 ,note #c3 ,the hydrochloric acid1 both free and combined1 is displaced by the less volatile sulphuric acid at its boiling point4 ,ferric sulphate separates at this point1 since there is no water to hold it in solution and care is required to prevent bumping4 ,the ferric sulphate usually has a silky appearance and is easily distinguished from the flocculent silica which often remains undissolved4 6,zimmermann-,reinhardt ,procedure6 6,method 7,b76 ,,procedure4--,grind the mineral to a fine powder4 ,weigh out two portions of about #j4e gram each into small porcelain crucibles4 ,proceed with the solution of the ore1 treat the residue1 if necessary1 and reduce the iron by the addition of stannous chloride1 followed by mercuric chloride1 as described for the bichromate process on page #ef4 ,dilute the solution to about #djj cc4 with cold water1 add #aj cc4 of the manganous sulphate titrating solution 7,note #a1 page #fh7 and titrate with the standard potassium permanganate solution to a faint pink 7,note #a74 ,from the standardization data already obtained calculate the percentage of iron 7,fe7 in the limonite4 ,note #a3 ,it has already been noted that hydrochloric acid reacts slowly in cold solutions with potassium permanganate4 ,it is1 however1 possible to obtain a satisfactory1 although somewhat fugitive end-point in the presence of manganous sulphate and phosphoric acid4 ,the explanation of the part played by these reagents is somewhat obscure as yet4 ,it is possible that an intermediate manganic compound is formed which reacts rapidly with the ferrous compounds--thus in effect catalyzing the oxidizing process4 ,while an excess of hydrochloric acid is necessary for the successful reduction of the iron by stannous chloride1 too large an amount should be avoided in order to lessen the chance of reduction of the permanganate by the acid during titration4 ,,determination ,,of ,,the ,,oxidizing ,,power ,,of ,,pyrolusite ,,indirect ,,oxidation ,pyrolusite1 when pure1 consists of manganese dioxide4 ,its value as an oxidizing agent1 and for the production of chlorine1 depends upon the percentage of ,mn,o#b in the sample4 ,this percentage is determined by an indirect method1 in which the manganese dioxide is reduced and dissolved by an excess of ferrous sulphate or oxalic acid in the presence of sulphuric acid1 and the unused excess determined by titration with standard permanganate solution4 ,,procedure4--,grind the mineral in an agate mortar until no grit whatever can be detected under the pestle 7,note #a74 ,transfer it to a stoppered weighing-tube1 and weigh out two portions of about #j4e gram into beakers 7#djj-ejj cc47 ,read ,note #b1 and then calculate in each case the weight of oxalic acid 7,h#b;,c#b,o#d4b;,h#b,o7 required to react with the weights of pyrolusite taken4 ,the reaction involved is ,mn,o#b ,h#b;,c#b,o#d7b;,h#b,o7 ,h#b,,so#d -- ,mn,s,o#d #b;,,co#b #d;,h#b,o4 ,weigh out about #j4b gram in excess of this quantity of 6pure6 oxalic acid into the corresponding beakers1 weighing the acid accurately and recording the weight in the notebook4 ,pour into each beaker #be cc4 of water and #ej cc4 of dilute sulphuric acid 7#a3e71 cover and warm the beaker and its contents gently until the evolution of carbon dioxide ceases 7,note #c74 ,if a residue remains which is sufficiently colored to obscure the end-reaction of the permanganate1 it must be removed by filtration4 ,finally1 dilute the solution to #bjj-cjj cc41 heat the solution to a temperature just below boiling1 add #ae cc4 of a manganese sulphate solution and while hot1 titrate for the excess of the oxalic acid with standard permanganate solution 7,notes #d and #e74 ,from the corrected volume of the solution required1 calculate the amount of oxalic acid undecomposed by the pyrolusite2 subtract this from the total quantity of acid used1 and calculate the weight of manganese dioxide which would react with the balance of the acid1 and from this the percentage in the sample4 ,note #a3 ,the success of the analysis is largely dependent upon the fineness of the powdered mineral4 ,if properly ground1 solution should be complete in fifteen minutes or less4 ,note #b3 ,a moderate excess of oxalic acid above that required to react with the pyrolusite is necessary to promote solution2 otherwise the residual quantity of oxalic acid would be so small that the last particles of the mineral would scarcely dissolve4 ,it is also desirable that a sufficient excess of the acid should be present to react with a considerable volume of the permanganate solution during the titration1 thus increasing the accuracy of the process4 ,on the other hand1 the excess of oxalic acid should not be so large as to react with more of the permanganate solution than is contained in a #ej cc4 burette4 ,if the pyrolusite under examination is known to be of high grade1 say #hj per cent pure1 or above the calculation of the oxalic acid needed may be based upon an assumption that the mineral is all ,mn,o#b4 ,if the quality of the mineral is unknown1 it is better to weigh out three portions instead of two and to add to one of these the amount of oxalic prescribed1 assuming complete purity of the mineral4 ,then run in the permanganate solution from a pipette or burette to determine roughly the amount required4 ,if the volume exceeds the contents of a burette1 the amount of oxalic acid added to the other two portions is reduced accordingly4 ,note #c3 ,care should be taken that the sides of the beaker are not overheated1 as oxalic acid would be decomposed by heat alone if crystallization should occur on the sides of the vessel4 ,strong sulphuric acid also decomposes the oxalic acid4 ,the dilute acid should1 therefore1 be prepared before it is poured into the beaker4 ,note #d3 ,ferrous ammonium sulphate1 ferrous sulphate1 or iron wire may be substituted for the oxalic acid4 ,the reaction is then the following3 #b ,fe,s,o#d ,mn,o#b #b;,h#b,,so#d -- ,fe#b7,,so#d7c #b;,h#b,o ,the excess of ferrous iron may also be determined by titration with potassium bichromate1 if desired4 ,care is required to prevent the oxidation of the iron by the air1 if ferrous salts are employed4 ,note #e3 ,the oxidizing power of pyrolusite may be determined by other volumetric processes1 one of which is outlined in the following reactions3 ,mn,o#b #d;,h,cl -- ,mn,cl#b ,cl#b #b;,h#b,o ,cl#b #b,,ki -- ,i#b #b,k,cl ,i#b #b,na#b,s#b,o#c -- ,na#b,s#d,o#f #b,na,i4 ,the chlorine generated by the pyrolusite is passed into a solution of potassium iodide4 ,the liberated iodine is then determined by titration with sodium thiosulphate1 as described on page #gh4 ,this is a direct process1 although it involves three steps4 ,,iodimetry ,the titration of iodine against sodium thiosulphate1 with starch as an indicator1 may perhaps be regarded as the most accurate of volumetric processes4 ,the thiosulphate solution may be used in both acid and neutral solutions to measure free iodine and the latter may1 in turn1 serve as a measure of any substance capable of liberating iodine from potassium iodide under suitable conditions for titration1 as1 for example1 in the process outlined in ,note #e on page #gd4 ,the fundamental reaction upon which iodometric processes are based is the following3 ,i#b #b ,na#b,s#b,o#c -- #b ,na,i ,na#b,s#d,o#f4 ,this reaction between iodine and sodium thiosulphate1 resulting in the formation of the compound ,na#b,s#d,o#f1 called sodium tetrathionate1 is quantitatively exact1 and differs in that respect from the action of chlorine or bromine1 which oxidize the thiosulphate1 but not quantitatively4 ,,normal ,,solutions ,,of ,,iodine ,,and ,,sodium ,,thiosulphate ,if the formulas of sodium thiosulphate and sodium tetrathionate are written in a manner to show the atoms of oxygen associated with sulphur atoms in each1 thus1 #b7,na#b74,s#b,o#b and ,na#b,o4,s#d,o#e1 it is plain that in the tetrathionate there are five atoms of oxygen associated with sulphur1 instead of the four in the two molecules of the thiosulphate taken together4 ,although1 therefore1 the iodine contains no oxygen1 the two atoms of iodine have1 in effect1 brought about the addition of one oxygen atoms to the sulphur atoms4 ,that is the same thing as saying that #bec4hd grams of iodine 7,i#b7 are equivalent to #af grams of oxygen2 hence1 since #h grams of oxygen is the basis of normal solutions1 #bec4hd_/b or #abf4ig grams of iodine should be contained in one liter of normal iodine solution4 ,by a similar course of reasoning the conclusion is reached that the normal solution of sodium thiosulphate should contain1 per liter1 its molecular weight in grams4 ,as the thiosulphate in crystalline form has the formula ,na#b,s#b,o#c4e;,h#b,o1 this weight is #bdh4ab grams4 ,tenth-normal or hundredth-normal solutions are generally used4 ,,preparation ,,of ,,standard ,,solutions 6,approximate ,strength1 #j4a ,n6 ,,procedure4--,weigh out on the rough balances #ac grams of commercial iodine4 ,place it in a mortar with #ah grams of potassium iodide and triturate with small portions of water until all is dissolved4 ,dilute the solution to #ajjj cc4 and transfer to a liter bottle and mix thoroughly 7,note #a74a ,footnote #a3 ,it will be found more economical to have a considerable quantity of the solution prepared by a laboratory attendant1 and to have all unused solutions returned to the common stock4 ,weigh out #be grams of sodium thiosulphate1 dissolve it in water which has been previously boiled and cooled1 and dilute to #ajjj cc41 also with boiled water4 ,transfer the solution to a liter bottle and mix thoroughly 7,note #b74 ,note #a3 ,iodine solutions react with water to form hydriodic acid under the influence of the sunlight1 and even at low room temperatures the iodine tends to volatilize from solution4 ,they should1 therefore1 be protected from light and heat4 ,iodine solutions are not stable for long periods under the best of conditions4 ,they cannot be used in burettes with rubber tips1 since they attack the rubber4 ,note #b3 ,sodium thiosulphate 7,na#b,s#b,o#c4e;,h#b,o7 is rarely wholly pure as sold commercially1 but may be purified by recrystallization4 ,the carbon dioxide absorbed from the air by distilled water decomposes the salt1 with the separation of sulphur4 ,boiled water which has been cooled out of contact with the air should be used in preparing solutions4 ,,indicator ,,solution ,the starch solution for use as an indicator must be freshly prepared4 ,a soluble starch is obtainable which serves well1 and a solution of #j4e gram of this starch in #be cc4 of boiling water is sufficient4 ,the solution should be filtered while hot and is ready for use when cold4 ,if soluble starch is not at hand1 potato starch may be used4 ,mix about #a gram with #e cc4 of cold water to a smooth paste1 pour #aej cc4 of 6boiling6 water over it1 warm for a moment on the hot plate1 and put it aside to settle4 ,decant the supernatant liquid through a filter and use the clear filtrate2 #e cc4 of this solution are needed for a titration4 ,the solution of potato starch is less stable than the soluble starch4 ,the solid particles of the starch1 if not removed by filtration1 become so colored by the iodine that they are not readily decolorized by the thiosulphate 7,note #a74 ,note #a3 ,the blue color which results when free iodine and starch are brought together is probably not due to the formation of a true chemical compound4 ,it is regarded as a solid solution of iodine in starch4 ,although it is unstable1 and easily destroyed by heat1 it serves as an indicator for the presence of free iodine of remarkable sensitiveness1 and makes the iodometric processes the most satisfactory of any in the field of volumetric analysis4 ,,comparison ,,of ,,iodine ,,and ,,thiosulphate ,,solutions ,,procedure4--,place the solutions in burettes 7the iodine in a glass-stoppered burette71 observing the usual precautions4 ,run out #dj cc4 of the thiosulphate solution into a beaker1 dilute with #aej cc4 of water1 add #a cc4 to #b cc4 of the soluble starch solution1 and titrate with the iodine to the appearance of the blue of the iodo-starch4 ,repeat until the ratio of the two solutions is established1 remembering all necessary corrections for burettes and for temperature changes4 ,,standardization ,,of ,,solutions ,commercial iodine is usually not sufficiently pure to permit of its use as a standard for thiosulphate solutions or the direct preparation of a standard solution of iodine4 ,it is likely to contain1 beside moisture1 some iodine chloride1 if chlorine was used to liberate the iodine when it was prepared4 ,it may be purified by sublimation after mixing it with a little potassium iodide1 which reacts with the iodine chloride1 forming potassium chloride and setting free the iodine4 ,the sublimed iodine is then dried by placing it in a closed container over concentrated sulphuric acid4 ,it may then be weighed in a stoppered weighing-tube and dissolved in a solution of potassium iodide in a stoppered flask to prevent loss of iodine by volatilization4 ,about #ah grams of the iodide and twelve grams of iodine per liter are required for an approximately tenth-normal solution4 ,an iodine solution made from commercial iodine may also be standardized against arsenious oxide 7,as#d,o#f74 ,this substance also usually requires purification by sublimation before use4 ,the substances usually employed for the standardization of a thiosulphate solution are potassium bromate and metallic copper4 ,the former is obtainable in pure condition or may be easily purified by re-crystallization4 ,copper wire of high grade is sufficiently pure to serve as a standard4 ,both potassium bromate and cupric salts in solution will liberate iodine from an iodide1 which is then titrated with the thiosulphate solution4 ,the reactions involved are the following3 7a7 ,k,br,o#c #f,,ki #c;,h#b,,so#d -- ,k,br #c;,i#b #c,k#b,,so#d #c;,h#b,o1 7b7 #c;,cu #h;,,hno#c -- #c;,cu7,,no#c7b #b,,no #d;,h#b,o1 #b;,cu7,,no#c7b #d,,ki -- #b;,cu,i #d,,kno#c ,i#b4 ,two methods for the direct standardization of the sodium thiosulphate solution are here described1 and one for the direct standardization of the iodine solution4 6,method ,a6 ,,procedure4--,weigh out into #ejj cc4 beakers two portions of about #j4aej-j4age gram of potassium bromate4 ,dissolve each of these in #ej cc4 of water1 and add #aj cc4 of a potassium iodide solution containing #c grams of the salt in that volume 7,note #a74 ,add to the mixture #aj cc4 of dilute sulphuric acid 7#a volume of sulphuric acid with #e volumes of water71 allow the solution to stand for three minutes1 and dilute to #aej cc4 7,note #b74 ,run in thiosulphate solution from a burette until the color of the liberated iodine is nearly destroyed1 and then add #a cc4 or #b cc4 of starch solution1 titrate to the disappearance of the iodo-starch blue1 and finally add iodine solution until the color is just restored4 ,make a blank test for the amount of thiosulphate solution required to react with the iodine liberated by the iodate which is generally present in the potassium iodide solution1 and deduct this from the total volume used in the titration4 ,from the data obtained1 calculate the relation of the thiosulphate solution to a normal solution1 and subsequently calculate the similar value for the iodine solution4 ,note #a3--,potassium iodide usually contains small amounts of potassium iodate as impurity which1 when the iodide is brought into an acid solution1 liberates iodine1 just as does the potassium bromate used as a standard4 ,it is necessary to determine the amount of thiosulphate which reacts with the iodine thus liberated by making a blank test with the iodide and acid alone4 ,as the iodate is not always uniformly distributed throughout the iodide1 it is better to make up a sufficient volume of a solution of the iodide for the purposes of the work in hand1 and to make the blank test by using the same volume of the iodide solution as is added in the standardizing process4 ,the iodide solution should contain about #c grams of the salt in #aj cc4 ,note #b3 ,the color of the iodo-starch is somewhat less satisfactory in concentrated solutions of the alkali salts1 notably the iodides4 ,the dilution prescribed obviates this difficulty4 6,method ,b6 ,,procedure4--,weigh out two portions of #j4be-j4bg gram of clean copper wire into #bej cc4 ,erlenmeyer flasks 7,note #a74 ,add to each #e cc4 of concentrated nitric acid 7sp4 gr4 #a4db7 and #be cc4 of water1 cover1 and warm until solution is complete4 ,add #e cc4 of bromine water and boil until the excess of bromine is expelled4 ,cool1 and add strong ammonia 7sp4 gr4 #j4ij7 drop by drop until a deep blue color indicates the presence of an excess4 ,boil the solution until the deep blue is replaced by a light bluish green1 or a brown stain appears on the sides of the flask 7,note #b74 ,add #aj cc4 of strong acetic acid 7sp4 gr4 #a4jd71 cool under the water tap1 and add a solution of potassium iodide 7,note #c7 containing about #c grams of the salt1 and titrate with thiosulphate solution until the color of the liberated iodine is nearly destroyed4 ,then add #a-b cc4 of freshly prepared starch solution1 and add thiosulphate solution1 drop by drop1 until the blue color is discharged4 ,from the data obtained1 including the blank test of the iodide1 calculate the relation of the thiosulphate solution to the normal4 ,note #a3 ,while copper wire of commerce is not absolutely pure1 the requirements for its use as a conductor of electricity are such that the impurities constitute only a few hundredths of one per cent and are negligible for analytical purposes4 ,note #b3 ,ammonia neutralizes the free nitric acid4 ,it should be added in slight excess only1 since the excess must be removed by boiling1 which is tedious4 ,if too much ammonia is present when acetic acid is added1 the resulting ammonium acetate is hydrolyzed1 and the ammonium hydroxide reacts with the iodine set free4 ,note #c3 ,a considerable excess of potassium iodide is necessary for the prompt liberation of iodine4 ,while a large excess will do no harm1 the cost of this reagent is so great that waste should be avoided4 6,method ,c6 ,,procedure4--,weigh out into #ejj cc4 beakers two portions of #j4age-j4bjj gram each of pure arsenious oxide4 ,dissolve each of these in #aj cc4 of sodium hydroxide solution1 with stirring4 ,dilute the solutions to #aej cc4 and add dilute hydrochloric acid until the solutions contain a few drops in excess1 and finally add to each a concentrated solution of #e grams of pure sodium bicarbonate 7,na,h,c,o#c7 in water4 ,cover the beakers before adding the bicarbonate1 to avoid loss4 ,add the starch solution and titrate with the iodine to the appearance of the blue of the iodo-starch1 taking care not to pass the end-point by more than a few drops 7,note #a74 ,from the corrected volume of the iodine solution used to oxidize the arsenious oxide1 calculate its relation to the normal4 ,from the ratio between the solutions1 calculate the similar value for the thiosulphate solution4 ,note #a3 ,arsenious oxide dissolves more readily in caustic alkali than in a bicarbonate solution1 but the presence of caustic alkali during the titration is not admissible4 ,it is therefore destroyed by the addition of acid1 and the solution is then made neutral with the solution of bicarbonate1 part of which reacts with the acid1 the excess remaining in solution4 ,the reaction during titration is the following3 ,na#c;,as,o#c ,i#b #b,na,h,c,o#c -- ,na#c;,as,o#d #b,na,i #b;,,co#b ,h#b,o ,as the reaction between sodium thiosulphate and iodine is not always free from secondary reactions in the presence of even the weakly alkaline bicarbonate1 it is best to avoid the addition of any considerable excess of iodine4 ,should the end-point be passed by a few drops1 the thiosulphate may be used to correct it4 ,,determination ,,of ,,copper ,,in ,,ores ,copper ores vary widely in composition from the nearly pure copper minerals1 such as malachite and copper sulphide1 to very low grade materials which contain such impurities as silica1 lead1 iron1 silver1 sulphur1 arsenic1 and antimony4 ,in nearly all varieties there will be found a siliceous residue insoluble in acids4 ,the method here given1 which is a modification of that described by ,a4,h4 ,low 76,j4 ,am4 ,chem4 ,soc46 7#aijb71 #bd1 #ajhb71 provides for the extraction of the copper from commonly occurring ores1 and for the presence of their common impurities4 ,for practice analyses it is advisable to select an ore of a fair degree of purity4 ,,procedure4-- ,weigh out two portions of about #j4e gram each of the ore 7which should be ground until no grit is detected7 into #bej cc4 ,erlenmeyer flasks or small beakers4 ,add #aj cc4 of concentrated nitric acid 7sp4 gr4 #a4db7 and heat very gently until the ore is decomposed and the acid evaporated nearly to dryness 7,note #a74 ,add #e cc4 of concentrated hydrochloric acid 7sp4 gr4 #a4b7 and warm gently4 ,then add about #g cc4 of concentrated sulphuric acid 7sp4 gr4 #a4hd7 and evaporate over a free flame until the sulphuric acid fumes freely 7,note #b74 ,it has then displaced nitric and hydrochloric acid from their compounds4 ,cool the flask or beaker1 add #be cc4 of water1 heat the solution to boiling1 and boil for two minutes4 ,filter to remove insoluble sulphates1 silica and any silver that may have been precipitated as silver chloride1 and receive the filtrate in a small beaker1 washing the precipitate and filter paper with warm water until the filtrate and washings amount to #ge cc4 ,bend a strip of aluminium foil 7#e cm4 x #ab cm47 into triangular form and place it on edge in the beaker4 ,cover the beaker and boil the solution 7being careful to avoid loss of liquid by spattering7 for ten minutes1 but do not evaporate to small volume4 ,wash the cover glass and sides of the beaker4 ,the copper should now be in the form of a precipitate at the bottom of the beaker or adhering loosely to the aluminium sheet4 ,remove the sheet1 wash it carefully with hydrogen sulphide water and place it in a small beaker4 ,decant the solution through a filter1 wash the precipitated copper twice by decantation with hydrogen sulphide water1 and finally transfer the copper to the filter paper1 where it is again washed thoroughly1 being careful at all times to keep the precipitated copper covered with the wash water4 ,remove and discard the filtrate and place an ,erlenmeyer flask under the funnel4 ,pour #ae cc4 of dilute nitric acid 7sp4 gr4 #a4bj7 over the aluminium foil in the beaker1 thus dissolving any adhering copper4 ,wash the foil with hot water and remove it4 ,warm this nitric acid solution and pour it slowly through the filter paper1 thereby dissolving the copper on the paper1 receiving the acid solution in the ,erlenmeyer flask4 ,before washing the paper1 pour #e cc4 of saturated bromine water 7,note #c7 through it and finally wash the paper carefully with hot water and transfer any particles of copper which may be left on it to the ,erlenmeyer flask4 ,boil to expel the bromine4 ,add concentrated ammonia drop by drop until the appearance of a deep blue coloration indicates an excess4 ,boil until the deep blue is displaced by a light bluish green coloration1 or until brown stains form on the sides of the flask4 ,add #aj cc4 of strong acetic acid 7,note #d7 and cool under the water tap4 ,add a solution containing about #c grams of potassium iodide1 as in the standardization1 and titrate with thiosulphate solution until the yellow of the liberated iodine is nearly discharged4 ,add #a-b cc4 of freshly prepared starch solution and titrate to the disappearance of the blue color4 ,from the data obtained1 calculate the percentage of copper 7,cu7 in the ore4 ,note #a3 ,nitric acid1 because of its oxidizing power1 is used as a solvent for the sulphide ores4 ,as a strong acid it will also dissolve the copper from carbonate ores4 ,the hydrochloric acid is added to dissolve oxides of iron and to precipitate silver and lead4 ,the sulphuric acid displaces the other acids1 leaving a solution containing sulphates only4 ,it also1 by its dehydrating action1 renders silica from silicates insoluble4 ,note #b3 ,unless proper precautions are taken to insure the correct concentrations of acid the copper will not precipitate quantitatively on the aluminium foil2 hence care must be taken to follow directions carefully at this point4 ,lead and silver have been almost completely removed as sulphate and chloride respectively1 or they too would be precipitated on the aluminium4 ,bismuth1 though precipitated on aluminium1 has no effect on the analysis4 ,arsenic and antimony precipitate on aluminium and would interfere with the titration if allowed to remain in the lower state of oxidation4 ,note #c3 ,bromine is added to oxidize arsenious and antimonious compounds from the original sample1 and to oxidize nitrous acid formed by the action of nitric acid on copper and copper sulphide4 ,note #d3 ,this reaction can be carried out in the presence of sulphuric and hydrochloric acids as well as acetic acid1 but in the presence of these strong acids arsenic and antimonic acids may react with the hydriodic acid produced with the liberation of free iodine1 thereby reversing the process and introducing an error4 ,,determination ,,of ,,antimony ,,in ,,stibnite ,stibnite is native antimony sulphide4 ,nearly pure samples of this mineral are easily obtainable and should be used for practice1 since many impurities1 notably iron1 seriously interfere with the accurate determination of the antimony by iodometric methods4 ,it is1 moreover1 essential that the directions with respect to amounts of reagents employed and concentration of solutions should be followed closely4 ,,procedure4--,grind the mineral with great care1 and weigh out two portions of #j4ce-j4dj gram into small1 dry beakers 7#ajj cc474 ,cover the beakers and pour over the stibnite #e cc4 of concentrated hydrochloric acid 7sp4 gr4 #a4bj7 and warm gently on the water bath 7,note #a74 ,when the residue is white1 add to each beaker #b grams of powdered tartaric acid 7,note #b74 ,warm the solution on the water bath for ten minutes longer1 dilute the solution very cautiously by adding water in portions of #e cc41 stopping if the solution turns red4 ,it is possible that no coloration will appear1 in which case cautiously continue the dilution to #abe cc4 ,if a red precipitate or coloration does appear1 warm the solution until it is colorless1 and again dilute cautiously to a total volume of #abe cc4 and boil for a minute 7,note #c74 ,if a white precipitate of the oxychloride separates during dilution 7which should not occur if the directions are followed71 it is best to discard the determination and to start anew4 ,carefully neutralize most of the acid with ammonium hydroxide solution 7sp4 gr4 #j4if71 but leave it distinctly acid 7,note #d74 ,dissolve #c grams of sodium bicarbonate in #bjj cc4 of water in a #ejj cc4 beaker1 and pour the cold solution of the antimony chloride into this1 avoiding loss by effervescence4 ,make sure that the solution contains an excess of the bicarbonate1 and then add #a cc4 or #b cc4 of starch solution and titrate with iodine solution to the appearance of the blue1 avoiding excess 7,notes #e and #f74 ,from the corrected volume of the iodine solution required to oxidize the antimony1 calculate the percentage of antimony 7,sb7 in the stibnite4 ,note #a3 ,antimony chloride is volatile with steam from its concentrated solutions2 hence these solutions must not be boiled until they have been diluted4 ,note #b3 ,antimony salts1 such as the chloride1 are readily hydrolyzed1 and compounds such as ,sb,o,cl are formed which are often relatively insoluble2 but in the presence of tartaric acid compounds with complex ions are formed1 and these are soluble4 ,an excess of hydrochloric acid also prevents precipitation of the oxychloride because the ,h ions from the acid lessen the dissociation of the water and thus prevent any considerable hydrolysis4 ,note #c3 ,the action of hydrochloric acid upon the sulphide sets free sulphureted hydrogen1 a part of which is held in solution by the acid4 ,this is usually expelled by the heating upon the water bath2 but if it is not wholly driven out1 a point is reached during dilution at which the antimony sulphide1 being no longer held in solution by the acid1 separates4 ,if the dilution is immediately stopped and the solution warmed1 this sulphide is again brought into solution and at the same time more of the sulphureted hydrogen is expelled4 ,this procedure must be continued until the sulphureted hydrogen is all removed1 since it reacts with iodine4 ,if no precipitation of the sulphide occurs1 it is an indication that the sulphureted hydrogen was all expelled on solution of the stibnite4 ,note #d3 ,ammonium hydroxide is added to neutralize most of the acid1 thus lessening the amount of sodium bicarbonate to be added4 ,the ammonia should not neutralize all of the acid4 ,note #e3 ,the reaction which takes place during titration may be expressed thus3 ,na#c,sb,o#c #b,na,h,c,o#c ,i#b -- ,na#c,sb,o#d #b,na,i ,h#b,o #b;,,co#b4 ,note #f3 ,if the end-point is not permanent1 that is1 if the blue of the iodo-starch is discharged after standing a few moments1 the cause may be an insufficient quantity of sodium bicarbonate1 leaving the solution slightly acid1 or a very slight precipitation of an antimony compound which is slowly acted upon by the iodine when the latter is momentarily present in excess4 ,in either case it is better to discard the analysis and to repeat the process1 using greater care in the amounts of reagents employed4 ,,chlorimetry ,the processes included under the term 6chlorimetry6 comprise those employed to determine chlorine1 hypochlorites1 bromine1 and hypobromites4 ,the reagent employed is sodium arsenite in the presence of sodium bicarbonate4 ,the reaction in the case of the hypochlorites is ,na,cl,o ,na#c;,as,o#c -- ,na#c;,as,o#d ,na,cl4 ,the sodium arsenite may be prepared from pure arsenious oxide1 as described below1 and is stable for considerable periods2 but commercial oxide requires resublimation to remove arsenic sulphide1 which may be present in small quantity4 ,to prepare the solution1 dissolve about #e grams of the powdered oxide1 accurately weighed1 in #aj cc4 of a concentrated sodium hydroxide solution1 dilute the solution to #cjj cc41 and make it faintly acid with dilute hydrochloric acid4 ,add #cj grams of sodium bicarbonate dissolved in a little water1 and dilute the solution to exactly #ajjj cc4 in a measuring flask4 ,transfer the solution to a dry liter bottle and mix thoroughly4 ,it is possible to dissolve the arsenious oxide directly in a solution of sodium bicarbonate1 with gentle warming1 but solution in sodium hydroxide takes place much more rapidly1 and the excess of the hydroxide is readily neutralized by hydrochloric acid1 with subsequent addition of the bicarbonate to maintain neutrality during the titration4 ,the indicator required for this process is made by dipping strips of filter paper in a starch solution prepared as described on page #gf1 to which #a gram of potassium iodide has been added4 ,these strips are allowed to drain and spread upon a watch-glass until dry4 ,when touched by a drop of the solution the paper turns blue until the hypochlorite has all been reduced and an excess of the arsenite has been added4 ,,determination ,,of ,,the ,,available ,,chlorine ,,in ,,bleaching ,,powder ,bleaching powder consists mainly of a calcium compound which is a derivative of both hydrochloric and hypochlorous acids4 ,its formula is ,ca,cl,o,cl4 ,its use as a bleaching or disinfecting agent1 or as a source of chlorine1 depends upon the amount of hypochlorous acid which it yields when treated with a stronger acid4 ,it is customary to express the value of bleaching powder in terms of available chlorine1 by which is meant the chlorine present as hypochlorite1 but not the chlorine present as chloride4 ,,procedure4--,weigh out from a stoppered test tube into a porcelain mortar about #c4e grams of bleaching powder 7,note #a74 ,triturate the powder in the mortar with successive portions of water until it is well ground and wash the contents into a #ejj cc4 measuring flask 7,note #b74 ,fill the flask to the mark with water and shake thoroughly4 ,measure off #be cc4 of this semi-solution in a measuring flask1 or pipette1 observing the precaution that the liquid removed shall contain approximately its proportion of suspended matter4 ,empty the flask or pipette into a beaker and wash it out4 ,run in the arsenite solution from a burette until no further reaction takes place on the starch-iodide paper when touched by a drop of the solution of bleaching powder4 ,repeat the titration1 using a second #be cc4 portion4 ,from the volume of solution required to react with the bleaching powder1 calculate the percentage of available chlorine in the latter1 assuming the titration reaction to be that between chlorine and arsenious oxide3 ,as#d,o#f #d;,cl#b #d;,h#b,o -- #b;,as#b,o#e #h;,h,cl ,note that only one twentieth of the original weight of bleaching powder enters into the reaction4 ,note #a3 ,the powder must be triturated until it is fine1 otherwise the lumps will inclose calcium hypochlorite1 which will fail to react with the arsenious acid4 ,the clear supernatant liquid gives percentages which are below1 and the sediment percentages which are above1 the average4 ,the liquid measured off should1 therefore1 carry with it its proper proportion of the sediment1 so far as that can be brought about by shaking the solution just before removal of the aliquot part for titration4 ,note #b3 ,bleaching powder is easily acted upon by the carbonic acid in the air1 which liberates the weak hypochlorous acid4 ,this1 of course1 results in a loss of available chlorine4 ,the original material for analysis should be kept in a closed container and protected form the air as far as possible4 ,it is difficult to obtain analytical samples which are accurately representative of a large quantity of the bleaching powder4 ,the procedure1 as outlined1 will yield results which are sufficiently exact for technical purposes4 ,,iii4 ,,precipitation ,,methods ,,determination ,,of ,,silver ,,by ,,the ,,thiocyanate ,,process ,the addition of a solution of potassium or ammonium thiocyanate to one of silver in nitric acid causes a deposition of silver thiocyanate as a white1 curdy precipitate4 ,if ferric nitrate is also present1 the slightest excess of the thiocyanate over that required to combine with the silver is indicated by the deep red which is characteristic of the thiocyanate test for iron4 ,the reactions involved are3 ,ag,n,o#c ,,kscn -- ,ag,s,c,n ,,kno#c1 #c,,kscn ,fe7,,no#c7c -- ,fe7,,scn7#c #c,,kno#c4 ,the ferric thiocyanate differs from the great majority of salts in that it is but very little dissociated in aqueous solutions1 and the characteristic color appears to be occasioned by the formation of the un-ionized ferric salt4 ,the normal solution of potassium thiocyanate should contain an amount of the salt per liter of solution which would yield sufficient 7,,cns7- to combine with one gram of hydrogen to form ,,hcns1 i4e41 a gram-molecular weight of the salt or #ig4ag grams4 ,if the ammonium thiocyanate is used1 the amount is #gf4jh grams4 ,to prepare the solution for this determination1 which should be approximately #j4je ,n1 dissolve about #e grams of potassium thiocyanate1 or #d grams of ammonium thiocyanate1 in a small amount of water2 dilute this solution to #ajjj cc4 in a liter bottle and mix as usual4 ,prepare #bj cc4 of a saturated solution of ferric alum and add #e cc4 of dilute nitric acid 7sp4 gr4 #a4bj74 ,about #e cc4 of this solution should be used as an indicator4 ,,standardization ,,procedure4--,crush a small quantity of silver nitrate crystals in a mortar 7,note #a74 ,transfer them to a watch-glass and dry them for an hour at #aaj;,c41 protecting them from dust or other organic matter 7,note #b74 ,weigh out two portions of about #j4e gram each and dissolve them in #ej cc4 of water4 ,add #aj cc4 of dilute nitric acid which has been recently boiled to expel the lower oxides of nitrogen1 if any1 and then add #e cc4 of the indicator solution4 ,run in the thiocyanate solution from a burette1 with constant stirring1 allowing the precipitate to settle occasionally to obtain an exact recognition of the end-point1 until a faint red tinge can be detected in the solution4 ,from the data obtained1 calculate the relation of the thiocyanate solution to the normal4 ,note #a3 ,the thiocyanate cannot be accurately weighed2 its solutions must1 therefore1 be standardized against silver nitrate 7or pure silver71 either in the form of a standard solution or in small1 weighed portions4 ,note #b3 ,the crystals of silver nitrate sometimes inclose water which is expelled on drying4 ,if the nitrate has come into contact with organic bodies it suffers a reduction and blackens during the heating4 ,it is plain that a standard solution of silver nitrate 7made by weighing out the crystals7 is convenient or necessary if many titrations of this nature are to be made4 ,in the absence of such a solution the liability of passing the end-point is lessened by setting aside a small fraction of the silver solution1 to be added near the close of the titration4 ,,determination ,,of ,,silver ,,in ,,coin ,,procedure4-- ,weigh out two portions of the coin of about #j4e gram each4 ,dissolve them in #ae cc4 of dilute nitric acid 7sp4 gr4 #a4b7 and boil until all the nitrous compounds are expelled 7,note #a74 ,cool the solution1 dilute to #ej cc41 and add #e cc4 of the indicator solution1 and titrate with the thiocyanate to the appearance of the faint red coloration 7,note #b74 ,from the corrected volume of the thiocyanate solution required1 calculate the percentage of silver in the coin4 ,note #a3 ,the reaction with silver may be carried out in nitric acid solutions and in the presence of copper1 if the latter does not exceed #gj per cent4 ,above that percentage it is necessary to add silver in known quantity to the solution4 ,the liquid must be cold at the time of titration and entirely free from nitrous compounds1 as these sometimes cause a reddening of the indicator solution4 ,all utensils1 distilled water1 the nitric acid and the beakers must be free from chlorides1 as the presence of these will cause precipitation of silver chloride1 thereby introducing an error4 ,note #b3 ,the solution containing the silver precipitate1 as well as those from the standardization1 should be placed in the receptacle for silver residues as a matter of economy4 ,,part ,,iii ,,gravimetric ,,analysis ,,general ,,directions ,gravimetric analyses involve the following principal steps3 first1 the weighing of the sample2 second1 the solution of the sample2 third1 the separation of some substance from solution containing1 or bearing a definite relation to1 the constituent to be measured1 under conditions which render this separation as complete as possible2 and finally1 the segregation of that substance1 commonly by filtration1 and the determination of its weight1 or that of some stable product formed from it on ignition4 ,for example1 the gravimetric determination of aluminium is accomplished by solution of the sample1 by precipitation in the form of hydroxide1 collection of the hydroxide upon a filter1 complete removal by washing of all foreign soluble matter1 and the burning of the filter and ignition of the precipitate to aluminium oxide1 in which condition it is weighed4 ,among the operations which are common to nearly all gravimetric analyses are precipitation1 washing of precipitates1 ignition of precipitates1 and the use of desiccators4 ,in order to avoid burdensome repetitions in the descriptions of the various gravimetric procedures which follow1 certain general instructions are introduced at this point4 ,these instructions must1 therefore1 be considered to be as much a part of all subsequent procedures as the description of apparatus1 reagents1 or manipulations4 ,the analytical balance1 the fundamentally important instrument in gravimetric analysis1 has already been described on pages #aa to #ae4 ,,precipitation ,for successful quantitative precipitations those substances are selected which are least soluble under conditions which can be easily established1 and which separate from solution in such a state that they can be filtered readily and washed free from admixed material4 ,in general1 the substances selected are the same as those already familiar to the student of ,qualitative ,analysis4 ,when possible1 substances are selected which separate in crystalline form1 since such substances are less likely to clog the pores of filter paper and can be most quickly washed4 ,in order to increase the size of the crystals1 which further promotes filtration and washing1 it is often desirable to allow a precipitate to remain for some time in contact with the solution from which it has separated4 ,the solution is often kept warm during this period of digestion4 ,the small crystals gradually disappear and the larger crystals increase in size1 probably as the result of the force known as surface tension1 which tends to reduce the surface of a given mass of material to a minimum1 combined with a very slightly greater solubility of small crystals as compared with the larger ones4 ,amorphous substances1 such as ferric hydroxide1 aluminium hydroxide1 or silicic acid1 separate in a gelatinous form and are relatively difficult to filter and wash4 ,substances of this class also exhibit a tendency to form1 with pure water1 what are known as colloidal solutions4 ,to prevent this as far as possible1 they are washed with solutions of volatile salts1 as will be described in some of the following procedures4 ,in all precipitations the reagent should be added slowly1 with constant stirring1 and should be hot when circumstances permit4 ,the slow addition is less likely to occasion contamination of the precipitate by the inclosure of other substances which may be in the solution1 or of the reagent itself4 ,,funnels ,,and ,,filters ,filtration in analytical processes is most commonly effected through paper filters4 ,in special cases these may be advantageously replaced by an asbestos filter in a perforated porcelain or platinum crucible1 commonly known1 from its originator1 as a ,gooch filter4 ,the operation and use of a filter of this type is described on page #ajc4 ,porous crucibles of a material known as alundum may also be employed to advantage in special cases4 ,the glass funnels selected for use with paper filters should have an angle as near #fj as possible1 and a narrow stem about six inches in length4 ,the filters employed should be washed filters1 i4e41 those which have been treated with hydrochloric and hydrofluoric acids1 and which on incineration leave a very small and definitely known weight of ash1 generally about 4#jjjjc gram4 ,such filters are readily obtainable on the market4 ,the filter should be carefully folded to fit the funnel according to either of the two well-established methods described in the ,appendix4 ,it should always be placed so that the upper edge of the paper is about one fourth inch below the top of the funnel4 ,under no circumstances should the filter extend above the edge of the funnel1 as it is then utterly impossible to effect complete washing4 ,to test the efficiency of the filter1 fill it with distilled water4 ,this water should soon fill the stem completely1 forming a continuous column of liquid which1 by its hydrostatic pressure1 produces a gentle suction1 thus materially promoting the rapidity of filtration4 ,unless the filter allows free passage of water under these conditions1 it is likely to give much trouble when a precipitate is placed upon it4 ,the use of a suction pump to promote filtration is rarely altogether advantageous in quantitative analysis1 if paper filters are employed4 ,the tendency of the filter to break1 unless the point of the filter paper is supported by a perforated porcelain cone or a small hardened filter of parchment1 and the tendency of the precipitates to pass through the pores of the filter1 more than compensate for the possible gain in time4 ,on the other hand1 filtration by suction may be useful in the case of precipitates which do not require ignition before weighing1 or in the case of precipitates which are to be discarded without weighing4 ,this is best accomplished with the aid of the special apparatus called a ,gooch filter referred to above4 ,,filtration ,,and ,,washing ,,of ,,precipitates ,solutions should be filtered while hot1 as far as possible1 since the passage of a liquid through the pores of a filter is retarded by friction1 and this1 for water at #ajj;,c41 is less than one sixth of the resistance at #j;,c4 ,when the filtrate is received in a beaker1 the stem of the funnel should touch the side of the receiving vessel to avoid loss by spattering4 ,neglect of this precaution is a frequent source of error4 ,the vessels which contain the initial filtrate should 6always6 be replaced by clean ones1 properly labeled1 before the washing of a precipitate begins4 ,in many instances a finely divided precipitate which shows no tendency to pass through the filter at first1 while the solution is relatively dense1 appears at once in the washings4 ,under such conditions the advantages accruing from the removal of the first filtrate are obvious1 both as regards the diminished volume requiring refiltration1 and also the smaller number of washings subsequently required4 ,much time may often be saved by washing precipitates by decantation1 i4e41 by pouring over them1 while still in the original vessel1 considerable volumes of wash-water and allowing them to settle4 ,the supernatant1 clear wash-water is then decanted through the filter1 so far as practicable without disturbing the precipitate1 and a new portion of wash-water is added4 ,this procedure can be employed to special advantage with gelatinous precipitates1 which fill up the pores of the filter paper4 ,as the medium from which the precipitate is to settle becomes less dense it subsides less readily1 and it ultimately becomes necessary to transfer it to the filter and complete the washing there4 ,a precipitate should never completely fill a filter4 ,the wash-water should be applied at the top of the filter1 above the precipitate4 ,it may be shown mathematically that the washing is most 6rapidly6 accomplished by filling the filter well to the top with wash-water each time1 and allowing it to drain completely after each addition2 but that when a precipitate is to be washed with the 6least possible volume6 of liquid the latter should be applied in repeated 6small6 quantities4 ,gelatinous precipitates should not be allowed to dry before complete removal of foreign matter is effected4 ,they are likely to shrink and crack1 and subsequent additions of wash-water pass through these channels only4 ,all filtrates and wash-waters without exception must be properly tested4 6,this lies at the foundation of accurate work61 and the student should clearly understand that it is only by the invariable application of this rule that assurance of ultimate reliability can be secured4 ,every original filtrate must be tested to prove complete precipitation of the compound to be separated1 and the wash-waters must also be tested to assure complete removal of foreign material4 ,in testing the latter1 the amount first taken should be but a few drops if the filtrate contains material which is to be subsequently determined4 ,when1 however1 the washing of the filter and precipitate is nearly completed the amount should be increased1 and for the final test not less than #c cc4 should be used4 ,it is impossible to trust to one's judgment with regard to the washing of precipitates2 the washings from 6each precipitate6 of a series simultaneously treated must be tested1 since the rate of washing will often differ materially under apparently similar conditions1 6,no exception can ever be made to this rule64 ,the habit of placing a clean common filter paper under the receiving beaker during filtration is one to be commended4 ,on this paper a record of the number of washings can very well be made as the portions of wash-water are added4 ,it is an excellent practice1 when possible1 to retain filtrates and precipitates until the completion of an analysis1 in order that1 in case of question1 they may be examined to discover sources of error4 ,for the complete removal of precipitates from containing vessels1 it is often necessary to rub the sides of these vessels to loosen the adhering particles4 ,this can best be done by slipping over the end of a stirring rod a soft rubber device sometimes called a policeman4 ,,desiccators ,desiccators should be filled with fused1 anhydrous calcium chloride1 over which is placed a clay triangle1 or an iron triangle covered with silica tubes1 to support the crucible or other utensils4 ,the cover of the desiccator should be made air-tight by the use of a thin coating of vaseline4 ,pumice moistened with concentrated sulphuric acid may be used in place of the calcium chloride1 and is essential in special cases2 but for most purposes the calcium chloride1 if renewed occasionally and not allowed to cake together1 is practically efficient and does not slop about when the desiccator is moved4 ,desiccators should never remain uncovered for any length of time4 ,the dehydrating agents rapidly lose their efficiency on exposure to the air4 ,,crucibles ,it is often necessary in quantitative analysis to employ fluxes to bring into solution substances which are not dissolved by acids4 ,the fluxes in most common use are sodium carbonate and sodium or potassium acid sulphate4 ,in gravimetric analysis it is usually necessary to ignite the separated substance after filtration and washing1 in order to remove moisture1 or to convert it through physical or chemical changes into some definite and stable form for weighing4 ,crucibles to be used in fusion processes must be made of materials which will withstand the action of the fluxes employed1 and crucibles to be used for ignitions must be made of material which will not undergo any permanent change during the ignition1 since the initial weight of the crucible must be deducted from the final weight of the crucible and product to obtain the weight of the ignited substance4 ,the three materials which satisfy these conditions1 in general1 are platinum1 porcelain1 and silica4 ,platinum crucibles have the advantage that they can be employed at high temperatures1 but1 on the other hand1 these crucibles can never be used when there is a possibility of the reduction to the metallic state of metals like lead1 copper1 silver1 or gold1 which would alloy with and ruin the crucible4 ,when platinum crucibles are used with compounds of arsenic or phosphorus1 special precautions are necessary to prevent damage4 ,this statement applies to both fusions and ignitions4 ,fusions with sodium carbonate can be made only in platinum1 since porcelain or silica crucibles are attacked by this reagent4 ,acid sulphate fusions1 which require comparatively low temperatures1 can sometimes be made in platinum1 although platinum is slightly attacked by the flux4 ,porcelain or silica crucibles may be used with acid fluxes4 ,silica crucibles are less likely to crack on heating than porcelain crucibles on account of their smaller coefficient of expansion4 ,ignition of substances not requiring too high a temperature may be made in porcelain or silica crucibles4 ,iron1 nickel or silver crucibles are used in special cases4 ,in general1 platinum crucibles should be used whenever such use is practicable1 and this is the custom in private1 research or commercial laboratories4 ,platinum has1 however1 become so valuable that it is liable to theft unless constantly under the protection of the user4 ,as constant protection is often difficult in instructional laboratories1 it is advisable1 in order to avoid serious monetary losses1 to use porcelain or silica crucibles whenever these will give satisfactory service4 ,when platinum utensils are used the danger of theft should always be kept in mind4 ,,preparation ,,of ,,crucibles ,,for ,,use ,all crucibles1 of whatever material1 must always be cleaned1 ignited and allowed to cool in a desiccator before weighing1 since all bodies exposed to the air condense on their surfaces a layer of moisture which increases their weight4 ,the amount and weight of this moisture varies with the humidity of the atmosphere1 and the latter may change from hour to hour4 ,the air in the desiccator 7see above7 is kept at a constant and low humidity by the drying agent which it contains4 ,bodies which remain in a desiccator for a sufficient time 7usually #bj-cj minutes7 retain1 therefore1 on their surfaces a constant weight of moisture which is the same day after day1 thus insuring constant conditions4 ,hot objects1 such as ignited crucibles1 should be allowed to cool in the air until1 when held near the skin1 but little heat is noticeable4 ,if this precaution is not taken1 the air within the desiccator is strongly heated and expands before the desiccator is covered4 ,as the temperature falls1 the air contracts1 causing a reduction of air pressure within the covered vessel4 ,when the cover is removed 7which is often rendered difficult7 the inrush of air from the outside may sweep light particles out of a crucible1 thus ruining an entire analysis4 ,constant heating of platinum causes a slight crystallization of the surface which1 if not removed1 penetrates into the crucible4 ,gentle polishing of the surface destroys the crystalline structure and prevents further damage4 ,if sea sand is used for this purpose1 great care is necessary to keep it from the desk1 since beakers are easily scratched by it1 and subsequently crack on heating4 ,platinum crucibles stained in use may often be cleaned by the fusion in them of potassium or sodium acid sulphate1 or by heating with ammonium chloride4 ,if the former is used1 care should be taken not to heat so strongly as to expel all of the sulphuric acid1 since the normal sulphates sometimes expand so rapidly on cooling as to split the crucible4 ,the fused material should be poured out1 while hot1 on to a 6dry6 tile or iron surface4 ,,ignition ,,of ,,precipitates ,most precipitates may1 if proper precautions are taken1 be ignited without previous drying4 ,if1 however1 such precipitates can be dried without loss of time to the analyst 7as1 for example1 over night71 it is well to submit them to this process4 ,it should1 nevertheless1 be remembered that a partially dried precipitate often requires more care during ignition than a thoroughly moist one4 ,the details of the ignition of precipitates vary so much with the character of the precipitate1 its moisture content1 and temperature to which it is to be heated1 that these details will be given under the various procedures which follow4 ,,determination ,,of ,,chlorine ,,in ,,sodium ,,chloride 6,method ,a4 ,with the ,use of a ,gooch ,filter6 ,,procedure4--,carefully clean a weighing-tube containing the sodium chloride1 handling it as little as possible with the moist fingers1 and weigh it accurately to #j4jjja gram1 recording the weight at once in the notebook 7see ,appendix74 ,hold the tube over the top of a beaker 7#bjj-cjj cc471 and cautiously remove the stopper1 noting carefully that no particles fall from it1 or from the tube1 elsewhere than into the beaker4 ,pour out a small portion of the chloride1 replace the stopper1 and determine by approximate weighing how much has been removed4 ,continue this procedure until #j4be-j4cj gram has been taken from the tube1 then weigh accurately and record the weight beneath the first in the notebook4 ,the difference of the two weights represents the weight of the chloride taken for analysis4 ,again weigh a second portion of #j4be-j4cj gram into a second beaker of the same size as the first4 ,the beakers should be plainly marked to correspond with the entries in the notebook4 ,dissolve each portion of the chloride in #aej cc4 of distilled water and add about ten drops of dilute nitric acid 7sp4 gr4 #a4bj7 7,note #b74 ,calculate the volume of silver nitrate solution required to effect complete precipitation in each case1 and add slowly about #e cc4 in excess of that amount1 with constant stirring4 ,heat the solutions cautiously to boiling1 stirring occasionally1 and continue the heating and stirring until the precipitates settle promptly1 leaving a nearly clear supernatant liquid 7,note #c74 ,this heating should not take place in direct sunlight 7,note #d74 ,the beaker should be covered with a watch-glass1 and both boiling and stirring so regulated as to preclude any possibility of loss of material4 ,add to the clear liquid one or two drops of silver nitrate solution1 to make sure that an excess of the reagent is present4 ,if a precipitate1 or cloudiness1 appears as the drops fall into the solution1 heat again1 and stir until the whole precipitate has coagulated4 ,the solution is then ready for filtration4 ,prepare a ,gooch filter as follows3 ,fold over the top of a ,gooch funnel 7,fig4 #b7 a piece of rubber-band tubing1 such as is known as bill-tie tubing1 and fit into the mouth of the funnel a perforated porcelain crucible 7,gooch crucible71 making sure that when the crucible is gently forced into the mouth of the funnel an airtight joint results4 7,a small #a or #a-a_/d-;inch glass funnel may be used1 in which case the rubber tubing is stretched over the top of the funnel and then drawn up over the side of the crucible until an air-tight joint is secured47 ,,illustration3 ,,fig4 #b ,fit the funnel into the stopper of a filter bottle1 and connect the filter bottle with the suction pump4 ,suspend some finely divided asbestos1 which has been washed with acid1 in #bj to #cj cc4 of water 7,note #a72 allow this to settle1 pour off the very fine particles1 and then pour some of the mixture cautiously into the crucible until an even felt of asbestos1 not over #a_/cb inch in thickness1 is formed4 ,a gentle suction must be applied while preparing this felt4 ,wash the felt thoroughly by passing through it distilled water until all fine or loose particles are removed1 increasing the suction at the last until no more water can be drawn out of it2 place on top of the felt the small1 perforated porcelain disc and hold it in place by pouring a very thin layer of asbestos over it1 washing the whole carefully2 then place the crucible in a small beaker1 and place both in a drying closet at #ajj-aaj;,c4 for thirty to forty minutes4 ,cool the crucible in a desiccator1 and weigh4 ,heat again for twenty to thirty minutes1 cool1 and again weigh1 repeating this until the weight is constant within #j4jjjc gram4 ,the filter is then ready for use4 ,place the crucible in the funnel1 and apply a gentle suction1 6after which6 the solution to be filtered may be poured in without disturbing the asbestos felt4 ,when pouring liquid onto a ,gooch filter hold the stirring-rod at first well down in the crucible1 so that the liquid does not fall with any force upon the asbestos1 and afterward keep the crucible will filled with the solution4 ,pour the liquid above the silver chloride slowly onto the filter1 leaving the precipitate in the beaker as far as possible4 ,wash the precipitate twice by decantation with warm water2 then transfer it to the filter with the aid of a stirring-rod with a rubber tip and a stream from the wash-bottle4 ,examine the first portions of the filtrate which pass through the filter with great care for asbestos fibers1 which are most likely to be lost at this point4 ,refilter the liquid if any fibers are visible4 ,finally1 wash the precipitate thoroughly with warm water until free from soluble silver salts4 ,to test the washings1 disconnect the suction at the flask and remove the funnel or filter tube from the suction flask4 ,hold the end of the tube over the mouth of a small test tube and add from a wash-bottle #b-c cc4 of water4 ,allow the water to drip through into the test tube and add a drop of dilute hydrochloric acid4 ,no precipitate or cloud should form in the wash-water 7,note #af74 ,dry the filter and contents at #ajj-aaj;,c4 until the weight is constant within #j4jjjc gram1 as described for the preparation of the filter4 ,deduct the weight of the dry crucible from the final weight1 and from the weight of silver chloride thus obtained calculate the percentage of chlorine in the sample of sodium chloride4 ,note #a3 ,the washed asbestos for this type of filter is prepared by digesting in concentrated hydrochloric acid1 long-fibered asbestos which has been cut in pieces of about #j4e cm4 in length4 ,after digestion1 the asbestos is filtered off on a filter plate and washed with hot1 distilled water until free from chlorides4 ,a small portion of the asbestos is shaken with water1 forming a thin suspension1 which is bottled and kept for use4 ,note #b3 ,the nitric acid is added before precipitation to lessen the tendency of the silver chloride to carry down with it other substances which might be precipitated from a neutral solution4 ,a large excess of the acid would exert a slight solvent action upon the chloride4 ,note #c3 ,the solution should not be boiled after the addition of the nitric acid before the presence of an excess of silver nitrate is assured1 since a slight interaction between the nitric acid and the sodium chloride is possible1 by which a loss of chlorine1 either as such or as hydrochloric acid1 might ensue4 ,the presence of an excess of the precipitant can usually be recognized at the time of its addition1 by the increased readiness with which the precipitate coagulates and settles4 ,note #d3 ,the precipitate should not be exposed to strong sunlight1 since under those conditions a reduction of the silver chloride ensues which is accompanied by a loss of chlorine4 ,the superficial alteration which the chloride undergoes in diffused daylight is not sufficient to materially affect the accuracy of the determination4 ,it should be noted1 however1 that a slight error does result from the effect of light upon the silver chloride precipitate and in cases in which the greatest obtainable accuracy is required1 the procedure described under ,method ,b should be followed1 in which this slight reduction of the silver chloride is corrected by subsequent treatment with nitric and hydrochloric acids4 ,note #e3 ,the asbestos used in the ,gooch filter should be of the finest quality and capable of division into minute fibrous particles4 ,a coarse felt is not satisfactory4 ,note #f3 ,the precipitate must be washed with warm water until it is absolutely free from silver and sodium nitrates4 ,it may be assumed that the sodium salt is completely removed when the wash-water shows no evidence of silver4 ,it must be borne in mind that silver chloride is somewhat soluble in hydrochloric acid1 and only a single drop should be added4 ,the washing should be continued until no cloudiness whatever can be detected in #c cc4 of the washings4 ,silver chloride is but slightly soluble in water4 ,the solubility varies with its physical condition within small limits1 and is about #j4jjah gram per liter at #ah;,c4 for the curdy variety usually precipitated4 ,the chloride is also somewhat soluble in solutions of many chlorides1 in solutions of silver nitrate1 and in concentrated nitric acid4 ,as a matter of economy1 the filtrate1 which contains whatever silver nitrate was added in excess1 may be set aside4 ,the silver can be precipitated as chloride and later converted into silver nitrate4 ,note #g3 ,the use of the ,gooch filter commends itself strongly when a considerable number of halogen determinations are to be made1 since successive portions of the silver halides may be filtered on the same filter1 without the removal of the preceding portions1 until the crucible is about two thirds filled4 ,if the felt is properly prepared1 filtration and washing are rapidly accomplished on this filter1 and this1 combined with the possibility of collecting several precipitates on the same filter1 is a strong argument in favor of its use with any but gelatinous precipitates4 6,method ,b4 ,with the ,use of a ,paper ,filter6 ,,procedure4--,weigh out two portions of sodium chloride of about #j4be-j4c gram each and proceed with the precipitation of the silver chloride as described under ,method ,a above4 ,when the chloride is ready for filtration prepare two #i cm4 washed paper filters 7see ,appendix74 ,pour the liquid above the precipitates through the filters1 wash twice by decantation and transfer the precipitates to the filters1 finally washing them until free from silver solution as described4 ,the funnel should then be covered with a moistened filter paper by stretching it over the top and edges1 to which it will adhere on drying4 ,it should be properly labeled with the student's name and desk number1 and then placed in a drying closet1 at a temperature of about #ajj-aaj;,c41 until completely dry4 ,the perfectly dry filter is then opened over a circular piece of clean1 smooth1 glazed paper about six inches in diameter1 placed upon a larger piece about twelve inches in diameter4 ,the precipitate is removed from the filter as completely as possible by rubbing the sides gently together1 or by scraping them cautiously with a feather which has been cut close to the quill and is slightly stiff 7,note #a74 ,in either case1 care must be taken not to rub off any considerable quantity of the paper1 nor to lose silver chloride in the form of dust4 ,cover the precipitate on the glazed paper with a watch-glass to prevent loss of fine particles and to protect it from dust from the air4 ,fold the filter paper carefully1 roll it into a small cone1 and wind loosely around 6the top6 a piece of small platinum wire 7,note #b74 ,hold the filter by the wire over a small porcelain crucible 7which has been cleaned1 ignited1 cooled in a desiccator1 and weighed71 ignite it1 and allow the ash to fall into the crucible4 ,place the crucible upon a clean clay triangle1 on its side1 and ignite1 with a low flame well at its base1 until all the carbon of the filter has been consumed4 ,allow the crucible to cool1 add two drops of concentrated nitric acid and one drop of concentrated hydrochloric acid1 and heat 6very cautiously61 to avoid spattering1 until the acids have been expelled2 then transfer the main portion of the precipitate from the glazed paper to the cooled crucible1 placing the latter on the larger piece of glazed paper and brushing the precipitate from the smaller piece into it1 sweeping off all particles belonging to the determination4 ,moisten the precipitate with two drops of concentrated nitric acid and one drop of concentrated hydrochloric acid1 and again heat with great caution until the acids are expelled and the precipitate is white1 when the temperature is slowly raised until the silver chloride just begins to fuse at the edges 7,note #c74 ,the crucible is then cooled in a desiccator and weighed1 after which the heating 7without the addition of acids7 is repeated1 and it is again weighed4 ,this must be continued until the weight is constant within #j4jjjc gram in two consecutive weighings4 ,deduct the weight of the crucible1 and calculate the percentage of chlorine in the sample of sodium chloride taken for analysis4 ,note #a3 ,the separation of the silver chloride from the filter is essential1 since the burning carbon of the paper would reduce a considerable quantity of the precipitate to metallic silver1 and its complete reconversion to the chloride within the crucible1 by means of acids1 would be accompanied by some difficulty4 ,the small amount of silver reduced from the chloride adhering to the filter paper after separating the bulk of the precipitate1 and igniting the paper as prescribed1 can be dissolved in nitric acid1 and completely reconverted to chloride by hydrochloric acid4 ,the subsequent addition of the two acids to the main portion of the precipitate restores the chlorine to any chloride which may have been partially reduced by the sunlight4 ,the excess of the acids is volatilized by heating4 ,note #b3 ,the platinum wire is wrapped around the top of the filter during its incineration to avoid contact with any reduced silver from the reduction of the precipitate4 ,if the wire were placed nearer the apex1 such contact could hardly be avoided4 ,note #c3 ,silver chloride should not be heated to complete fusion1 since a slight loss by volatilization is possible at high temperatures4 ,the temperature of fusion is not always sufficient to destroy filter shreds2 hence these should not be allowed to contaminate the precipitate4 ,,determination ,,of ,,iron ,,and ,,of ,,sulphur ,,in ,,ferrous ,,ammonium ,,sulphate1 ,,feso#d47,,nh#d7b,,so#d4f;,h#b,o ,,determination ,,of ,,iron ,,procedure4--,weigh out into beakers 7#bjj-bej cc47 two portions of the sample 7,note #a7 of about #a gram each and dissolve these in #ej cc4 of water1 to which #a cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab7 has been added 7,note #b74 ,heat the solution to boiling1 and while at the boiling point add concentrated nitric acid 7sp4 gr4 #a4db71 6drop by drop6 7noting the volume used71 until the brown coloration1 which appears after the addition of a part of the nitric acid1 gives place to a yellow or red 7,note #c74 ,avoid a large excess of nitric acid1 but be sure that the action is complete4 ,pour this solution cautiously into about #bjj cc4 of water1 containing a slight excess of ammonia4 ,calculate for this purpose the amount of aqueous ammonia required to neutralize the hydrochloric and nitric acids added 7see ,appendix for data71 and also to precipitate the iron as ferric hydroxide from the weight of the ferrous ammonium sulphate taken for analysis1 assuming it to be pure 7,note #d74 ,the volume thus calculated will be in excess of that actually required for precipitation1 since the acids are in part consumed in the oxidation process1 or are volatilized4 ,heat the solution to boiling1 and allow the precipitated ferric hydroxide to settle4 ,decant the clear liquid through a washed filter 7#i cm471 keeping as much of the precipitate in the beaker as possible4 ,wash twice by decantation with #ajj cc4 of hot water4 ,reserve the filtrate4 ,dissolve the iron from the filter with hot1 dilute hydrochloric acid 7sp4 gr4 #a4ab71 adding it in small portions1 using as little as possible and noting the volume used4 ,collect the solution in the beaker in which precipitation took place4 ,add #a cc4 of nitric acid 7sp4 gr4 #a4db71 boil for a few moments1 and again pour into a calculated excess of ammonia4 ,wash the precipitate twice by decantation1 and finally transfer it to the original filter4 ,wash continuously with hot water until finally #c cc4 of the washings1 acidified with nitric acid 7,note #e71 show no evidences of the presence of chlorides when tested with silver nitrate4 ,the filtrate and washings are combined with those from the first precipitation and treated for the determination of sulphur1 as prescribed on page #aab4 ,note #a3 ,if a selection of pure material for analysis is to be made1 crystals which are cloudy are to be avoided on account of loss of water of crystallization2 and also those which are red1 indicating the presence of ferric iron4 ,if1 on the other hand1 the value of an average sample of material is desired1 it is preferable to grind the whole together1 mix thoroughly1 and take a sample from the mixture for analysis4 ,note #b3 ,when aqueous solutions of ferrous compounds are heated in the air1 oxidation of the ,fe ions to ,fe ions readily occurs in the absence of free acid4 ,the ,h and ,,oh- ions from water are involved in the oxidation process and the result is1 in effect1 the formation of some ferric hydroxide which tends to separate4 ,moreover1 at the boiling temperature1 the ferric sulphate produced by the oxidation hydrolyzes in part with the formation of a basic ferric sulphate1 which also tends to separate from solution4 ,the addition of the hydrochloric acid prevents the formation of ferric hydroxide1 and so far reduces the ionization of the water that the hydrolysis of the ferric sulphate is also prevented1 and no precipitation occurs on heating4 ,note #c3 ,the nitric acid1 after attaining a moderate strength1 oxidizes the ,fe ions to ,fe ions with the formation of an intermediate nitroso-compound similar in character to that formed in the ring-test for nitrates4 ,the nitric oxide is driven out by heat1 and the solution then shows by its color the presence of ferric compounds4 ,a drop of the oxidized solution should be tested on a watch-glass with potassium ferricyanide1 to insure a complete oxidation4 ,this oxidation of the iron is necessary1 since ,fe ions are not completely precipitated by ammonia4 ,the ionic changes which are involved in this oxidation are perhaps most simply expressed by the equation #c;,fe ,,no#c- #d;,h -- #c;,fe #b;,h#b,o ,,no1 the ,h ions coming from the acid in the solution1 in this case either the nitric or the hydrochloric acid4 ,the full equation on which this is based may be written thus3 #f;,fe,s,o#d #b;,,hno#c #f;,h,cl -- #b;,fe#b7,,so#d7c #b;,fe,cl#c #b,,no #d;,h#b,o1 assuming that only enough nitric acid is added to complete the oxidation4 ,note #d3 ,the ferric hydroxide precipitate tends to carry down some sulphuric acid in the form of basic ferric sulphate4 ,this tendency is lessened if the solution of the iron is added to an excess of ,,oh- ions from the ammonium hydroxide1 since under these conditions immediate and complete precipitation of the ferric hydroxide ensues4 ,a gradual neutralization with ammonia would result in the local formation of a neutral solution within the liquid1 and subsequent deposition of a basic sulphate as a consequence of a local deficiency of ,,oh- ions from the ,,nh#d,,oh and a partial hydrolysis of the ferric salt4 ,even with this precaution the entire absence of sulphates from the first iron precipitate is not assured4 ,it is1 therefore1 redissolved and again thrown down by ammonia4 ,the organic matter of the filter paper may occasion a partial reduction of the iron during solution1 with consequent possibility of incomplete subsequent precipitation with ammonia4 ,the nitric acid is added to reoxidize this iron4 ,to avoid errors arising from the solvent action of ammoniacal liquids upon glass1 the iron precipitate should be filtered without unnecessary delay4 ,note #e3 ,the washings from the ferric hydroxide are acidified with nitric acid1 before testing with silver nitrate1 to destroy the ammonia which is a solvent of silver chloride4 ,the use of suction to promote filtration and washing is permissible1 though not prescribed4 ,the precipitate should not be allowed to dry during the washing4 6,ignition of the ,iron ,precipitate6 ,heat a platinum or porcelain crucible1 cool it in a desiccator and weigh1 repeating until a constant weight is obtained4 ,fold the top of the filter paper over the moist precipitate of ferric hydroxide and transfer it cautiously to the crucible4 ,wipe the inside of the funnel with a small fragment of washed filter paper1 if necessary1 and place the paper in the crucible4 ,incline the crucible on its side1 on a triangle supported on a ring-stand1 and stand the cover on edge at the mouth of the crucible4 ,place a burner below the front edge of the crucible1 using a low flame and protecting it from drafts of air by means of a chimney4 ,the heat from the burner is thus reflected into the crucible and dries the precipitate without danger of loss as the result of a sudden generation of steam within the mass of ferric hydroxide4 ,as the drying progresses the burner may be gradually moved toward the base of the crucible and the flame increased until the paper of the filter begins to char and finally to smoke1 as the volatile matter is expelled4 ,this is known as smoking off a filter1 and the temperature should not be raised sufficiently high during this process to cause the paper to ignite1 as the air currents produced by the flame of the blazing paper may carry away particles of the precipitate4 ,when the paper is fully charred1 move the burner to the base of the crucible and raise the temperature to the full heat of the burner for fifteen minutes1 with the crucible still inclined on its side1 but without the cover 7,note #a74 ,finally set the crucible upright in the triangle1 cover it1 and heat at the full temperature of a blast lamp or other high temperature burner4 ,cool and weigh in the usual manner 7,note #b74 ,repeat the strong heating until the weight is constant within #j4jjjc gram4 ,from the weight of ferric oxide 7,fe#b,o#c7 calculate the percentage of iron 7,fe7 in the sample 7,note #c74 ,note #a3 ,these directions for the ignition of the precipitate must be closely followed4 ,a ready access of atmospheric oxygen is of special importance to insure the reoxidation to ferric oxide of any iron which may be reduced to magnetic oxide 7,fe#c,o#d7 during the combustion of the filter4 ,the final heating over the blast lamp is essential for the complete expulsion of the last traces of water from the hydroxide4 ,note #b3 ,ignited ferric oxide is somewhat hygroscopic4 ,on this account the weighings must be promptly completed after removal from the desiccator4 ,in all weighings after the first it is well to place the weights upon the balance-pan before removing the crucible from the desiccator4 ,it is then only necessary to move the rider to obtain the weight4 ,note #c3 ,the gravimetric determination of aluminium or chromium is comparable with that of iron just described1 with the additional precaution that the solution must be boiled until it contains but a very slight excess of ammonia1 since the hydroxides of aluminium and chromium are more soluble than ferric hydroxide4 ,the most important properties of these hydroxides1 from a quantitative standpoint1 other than those mentioned1 are the following3 ,all are precipitable by the hydroxides of sodium and potassium1 but always inclose some of the precipitant1 and should be reprecipitated with ammonium hydroxide before ignition to oxides4 ,chromium and aluminium hydroxides dissolve in an excess of the caustic alkalies and form anions1 probably of the formula ,al,o#b- and ,cr,o#b-4 ,chromium hydroxide is reprecipitated from this solution on boiling4 ,when first precipitated the hydroxides are all readily soluble in acids1 but aluminium hydroxide dissolves with considerable difficulty after standing or boiling for some time4 ,the precipitation of the hydroxides is promoted by the presence of ammonium chloride1 but is partially or entirely prevented by the presence of tartaric or citric acids1 glycerine1 sugars1 and some other forms of soluble organic matter4 ,the hydroxides yield on ignition an oxide suitable for weighing 7,al#b,o#c1 ,cr#b,o#c1 ,fe#b,o#c74 ,,determination ,,of ,,sulphur ,,procedure4--,add to the combined filtrates from the ferric hydroxide about #j4f gram of anhydrous sodium carbonate2 cover the beaker1 and then add dilute hydrochloric acid 7sp4 gr4 #a4ab7 in moderate excess and evaporate to dryness on the water bath4 ,add #aj cc4 of concentrated hydrochloric acid 7sp4 gr4 #a4bj7 to the residue1 and again evaporate to dryness on the bath4 ,dissolve the residue in water1 filter if not clear1 transfer to a #gjj cc4 beaker1 dilute to about #djj cc41 and cautiously add hydrochloric acid until the solution shows a distinctly acid reaction 7,note #a74 ,heat the solution to boiling1 and add 6very slowly6 and with constant stirring1 #bj cc4 in excess of the calculated amount of a hot barium chloride solution1 containing about #bj grams ,ba,cl#b4b;,h#b,o per liter 7,notes #b and #c74 ,continue the boiling for about two minutes1 allow the precipitate to settle1 and decant the liquid at the end of half an hour 7,note #d74 ,replace the beaker containing the original filtrate by a clean beaker1 wash the precipitated sulphate by decantation with hot water1 and subsequently upon the filter until it is freed from chlorides1 testing the washings as described in the determination of iron4 ,the filter is then transferred to a platinum or porcelain crucible and ignited1 as described above1 until the weight is constant 7,note #e74 ,from the weight of barium sulphate 7,ba,s,o#d7 obtained1 calculate the percentage of sulphur 7,s7 in the sample4 ,note #a3 ,barium sulphate is slightly soluble in hydrochloric acid1 even dilute1 probably as a result of the reduction in the degree of dissociation of sulphuric acid in the presence of the ,h ions of the hydrochloric acid1 and possibly because of the formation of a complex anion made up of barium and chlorine2 hence only the smallest excess should be added over the amount required to acidify the solution4 ,note #b3 ,the ionic changes involved in the precipitation of barium sulphate are very simple3 ,ba ,,so#d-- -- ,ba,s,o#d ,this case affords one of the best illustrations of the effect of an excess of a precipitant in decreasing the solubility of a precipitate4 ,if the conditions are considered which exist at the moment when just enough of the ,ba ions have been added to correspond to the ,,so#d-- ions in the solution1 it will be seen that nearly all of the barium sulphate has been precipitated1 and that the small amount which then remains in the solution which is in contact with the precipitate must represent a saturated solution for the existing temperature1 and that this solution is comparable with a solution of sugar to which more sugar has been added than will dissolve4 ,it should be borne in mind that the quantity of barium sulphate in this 6saturated solution is a constant quantity6 for the existing conditions4 ,the dissolved barium sulphate1 like any electrolyte1 is dissociated1 and the equilibrium conditions may be expressed thus3 76,conc'n ,ba x ,conc'n ,,so#d--7_/7;,conc'n ,ba,s,o#d7 ,const461 and since 6,conc'n ,ba,s,o#d6 for the saturated solution has a constant value 7which is very small71 it may be eliminated1 when the expression becomes 6,conc'n ,ba x ,conc'n ,,so#d-- ,const461 which is the solubility product of ,ba,s,o#d4 ,if1 now1 an excess of the precipitant1 a soluble barium salt1 is added in the form of a relatively concentrated solution 7the slight change of volume of a few cubic centimeters may be disregarded for the present discussion7 the concentration of the ,ba ions is much increased1 and as a consequence the 6,conc'n ,,so#d6 must decrease in proportion if the value of the expression is to remain constant1 which is a requisite condition if the law of mass action upon which our argument depends holds true4 ,in other words1 ,,so#d-- ions must combine with some of the added ,ba ions to form ,ba,s,o#d2 but it will be recalled that the solution is already saturated with ,ba,s,o#d1 and this freshly formed quantity must1 therefore1 separate and add itself to the precipitate4 ,this is exactly what is desired in order to insure more complete precipitation and greater accuracy1 and leads to the conclusion that the larger the excess of the precipitant added the more successful the analysis2 but a practical limit is placed upon the quantity of the precipitant which may be properly added by other conditions1 as stated in the following note4 ,note #c3 ,barium sulphate1 in a larger measure than most compounds1 tends to carry down other substances which are present in the solution from which it separates1 even when these other substances are relatively soluble1 and including the barium chloride used as the precipitant4 ,this is also notably true in the case of nitrates and chlorates of the alkalies1 and of ferric compounds2 and1 since in this analysis ammonium nitrate has resulted from the neutralization of the excess of the nitric acid added to oxidize the iron1 it is essential that this should be destroyed by repeated evaporation with a relatively large quantity of hydrochloric acid4 ,during evaporation a mutual decomposition of the two acids takes place1 and the nitric acid is finally decomposed and expelled by the excess of hydrochloric acid4 ,iron is usually found in the precipitate of barium sulphate when thrown down from hot solutions in the presence of ferric salts4 ,this1 according to ,kuster and ,thiel 76,zeit4 anorg4 ,chem461 #bb1 #dbd71 is due to the formation of a complex ion 7,fe7,,so#d7b7 which precipitates with the ,ba ion1 while ,richards 76,zeit4 anorg4 ,chem461 #bc1 #chc7 ascribes it to hydrolytic action1 which causes the formation of a basic ferric complex which is occluded in the barium precipitate4 ,whatever the character of the compound may be1 it has been shown that it loses sulphuric anhydride upon ignition1 causing low results1 even though the precipitate contains iron4 ,the contamination of the barium sulphate by iron is much less in the presence of ferrous than ferric salts4 ,if1 therefore1 the sulphur alone were to be determined in the ferrous ammonium sulphate1 the precipitation by barium might be made directly from an aqueous solution of the salt1 which had been made slightly acid with hydrochloric acid4 ,note #d3 ,the precipitation of the barium sulphate is probably complete at the end of a half-hour1 and the solution may safely be filtered at the expiration of that time if it is desired to hasten the analysis4 ,as already noted1 many precipitates of the general character of this sulphate tend to grow more coarsely granular if digested for some time with the liquid from which they have separated4 ,it is therefore well to allow the precipitate to stand in a warm place for several hours1 if practicable1 to promote ease of filtration4 ,the filtrate and washings should always be carefully examined for minute quantities of the sulphate which may pass through the pores of the filter4 ,this is best accomplished by imparting to the filtrate a gentle rotary motion1 when the sulphate1 if present1 will collect at the center of the bottom of the beaker4 ,note #e3 ,a reduction of barium sulphate to the sulphide may very readily be caused by the reducing action of the burning carbon of the filter1 and much care should be taken to prevent any considerable reduction from this cause4 ,subsequent ignition1 with ready access of air1 reconverts the sulphide to sulphate unless a considerable reduction has occurred4 ,in the latter case it is expedient to add one or two drops of sulphuric acid and to heat cautiously until the excess of acid is expelled4 ,note #f3 ,barium sulphate requires about #djj1jjj parts of water for its solution4 ,it is not decomposed at a red heat but suffers loss1 probably of sulphur trioxide1 at a temperature above #ijj;,c4 ,,determination ,,of ,,sulphur ,,in ,,barium ,,sulphate ,,procedure4--,weigh out1 into platinum crucibles1 two portions of about #j4e gram of the sulphate4 ,mix each in the crucible with five to six times its weight of anhydrous sodium carbonate4 ,this can best be done by placing the crucible on a piece of glazed paper and stirring the mixture with a clean1 dry stirring-rod1 which may finally be wiped off with a small fragment of filter paper1 the latter being placed in the crucible4 ,cover the crucible and heat until a quiet1 liquid fusion ensues4 ,remove the burner1 and tip the crucible until the fused mass flows nearly to its mouth4 ,hold it in that position until the mass has solidified4 ,when cold1 the material may usually be detached in a lump by tapping the crucible or gently pressing it near its upper edge4 ,if it still adheres1 a cubic centimeter or so of water may be placed in the cold crucible and cautiously brought to boiling1 when the cake will become loosened and may be removed and placed in about #bej cc4 of hot1 distilled water to dissolve4 ,clean the crucible completely1 rubbing the sides with a rubber-covered stirring-rod1 if need be4 ,when the fused mass has completely disintegrated and nothing further will dissolve1 decant the solution from the residue of barium carbonate 7,note #a74 ,pour over the residue #bj cc4 of a solution of sodium carbonate and #aj cc4 of water and heat to gentle boiling for about three minutes 7,note #b74 ,filter off the carbonate and wash it with hot water1 testing the slightly acidified washings for sulphate and preserving any precipitates which appear in these tests4 ,acidify the filtrate with hydrochloric acid until just acid1 bring to boiling1 and slowly add hot barium chloride solution1 as in the preceding determination4 ,add also any tests from the washings in which precipitates have appeared4 ,filter1 wash1 ignite1 and weigh4 ,from the weight of barium sulphate1 calculate the percentage of sulphur 7,s7 in the sample4 ,note #a3 ,this alkaline fusion is much employed to disintegrate substances ordinarily insoluble in acids into two components1 one of which is water soluble and the other acid soluble4 ,the reaction involved is3 ,ba,s,o#d ,na#b;,,co#c1 -- ,ba,c,o#c1 ,na#b,,so#d4 ,as the sodium sulphate is soluble in water1 and the barium carbonate insoluble1 a separation between them is possible and the sulphur can be determined in the water-soluble portion4 ,it should be noted that this method can be applied to the purification of a precipitate of barium sulphate if contaminated by most of the substances mentioned in ,note #c on page #aad4 ,the impurities pass into the water solution together with the sodium sulphate1 but1 being present in such minute amounts1 do not again precipitate with the barium sulphate4 ,note #b3 ,the barium carbonate is boiled with sodium carbonate solution before filtration because the reaction above is reversible2 and it is only by keeping the sodium carbonate present in excess until nearly all of the sodium sulphate solution has been removed by filtration that the reversion of some of the barium carbonate to barium sulphate is prevented4 ,this is an application of the principle of mass action1 in which the concentration of the reagent 7the carbonate ion7 is kept as high as practicable and that of the sulphate ion as low as possible1 in order to force the reaction in the desired direction 7see ,appendix74 ,,determination ,,of ,,phosphoric ,,anhydride ,,in ,,apatite ,the mineral apatite is composed of calcium phosphate1 associated with calcium chloride1 or fluoride4 ,specimens are easily obtainable which are nearly pure and leave on treatment with acid only a slight siliceous residue4 ,for the purpose of gravimetric determination1 phosphoric acid is usually precipitated from ammoniacal solutions in the form of magnesium ammonium phosphate which1 on ignition1 is converted into magnesium pyrophosphate4 ,since the calcium phosphate of the apatite is also insoluble in ammoniacal solutions1 this procedure cannot be applied directly4 ,the separation of the phosphoric acid from the calcium must first be accomplished by precipitation in the form of ammonium phosphomolybdate in nitric acid solution1 using ammonium molybdate as the precipitant4 ,the yellow precipitate1 as it is often called1 is not always of a definite composition1 and therefore not suitable for direct weighing1 but may be dissolved in ammonia1 and the phosphoric acid thrown out as magnesium ammonium phosphate from the solution4 ,of the substances likely to occur in apatite1 silicic acid alone interferes with the precipitation of the phosphoric acid in nitric acid solution4 ,,precipitation ,,of ,,ammonium ,,phosphomolybdate ,,procedure4--,grind the mineral in an agate mortar until no grit is perceptible4 ,transfer the substance to a weighing-tube1 and weigh out two portions1 not exceeding #j4bj gram each 7,note #a7 into two beakers of about #bjj cc4 capacity4 ,pour over them #bj cc4 of dilute nitric acid 7sp4 gr4 #a4b7 and warm gently until solvent action has apparently ceased4 ,evaporate the solution cautiously to dryness1 heat the residue for about an hour at #ajj-aaj;,c41 and treat it again with nitric acid as described above2 separate the residue of silica by filtration on a small filter 7#g cm47 and wash with warm water1 using as little as possible 7,note #b74 ,receive the filtrate in a beaker 7#bjj-ejj cc474 ,test the washings with ammonia for calcium phosphate1 but add all such tests in which a precipitate appears to the original nitrate 7,note #c74 ,the filtrate and washings must be kept as small as possible and should not exceed #ajj cc4 in volume4 ,add aqueous ammonia 7sp4 gr4 #j4if7 until the precipitate of calcium phosphate first produced just fails to redissolve1 and then add a few drops of nitric acid until this is again brought into solution 7,note #d74 ,warm the solution until it cannot be comfortably held in the hand 7about #fj;,c47 and1 after removal of the burner1 add #ge cc4 of ammonium molybdate solution which has been 6gently6 warmed1 but which must be perfectly clear4 ,allow the mixture to stand at a temperature of about #ej or #fj;,c4 for twelve hours 7,notes #e and #f74 ,filter off the yellow precipitate on a #i cm4 filter1 and wash by decantation with a solution of ammonium nitrate made acid with nitric acid4#a ,allow the precipitate to remain in the beaker as far as possible4 ,test the washings for calcium with ammonia and ammonium oxalate 7,note #c74 ,footnote #a3 ,this solution is prepared as follows3 ,mix #ajj cc4 of ammonia solution 7sp4 gr4 #j4if7 with #cbe cc4 of nitric acid 7sp4 gr4 #a4b7 and dilute with #ajj cc4 of water4 ,add #aj cc4 of molybdate solution to the nitrate1 and leave it for a few hours4 ,it should then be carefully examined for a 6yellow6 precipitate2 a white precipitate may be neglected4 ,note #a3 ,magnesium ammonium phosphate1 as noted below1 is slightly soluble under the conditions of operation4 ,consequently the unavoidable errors of analysis are greater in this determination than in those which have preceded it1 and some divergence may be expected in duplicate analyses4 ,it is obvious that the larger the amount of substance taken for analysis the less will be the relative loss or gain due to unavoidable experimental errors2 but1 in this instance1 a check is placed upon the amount of material which may be taken both by the bulk of the resulting precipitate of ammonium phosphomolybdate and by the excessive amount of ammonium molybdate required to effect complete separation of the phosphoric acid1 since a liberal excess above the theoretical quantity is demanded4 ,molybdic acid is one of the more expensive reagents4 ,note #b3 ,soluble silicic acid would1 if present1 partially separate with the phosphomolybdate1 although not in combination with molybdenum4 ,its previous removal by dehydration is therefore necessary4 ,note #c3 ,when washing the siliceous residue the filtrate may be tested for calcium by adding ammonia1 since that reagent neutralizes the acid which holds the calcium phosphate in solution and causes precipitation2 but after the removal of the phosphoric acid in combination with the molybdenum1 the addition of an oxalate is required to show the presence of calcium4 ,note #d3 ,an excess of nitric acid exerts a slight solvent action1 while ammonium nitrate lessens the solubility2 hence the neutralization of the former by ammonia4 ,note #e3 ,the precipitation of the phosphomolybdate takes place more promptly in warm than in cold solutions1 but the temperature should not exceed #fj;,c4 during precipitation2 a higher temperature tends to separate molybdic acid from the solution4 ,this acid is nearly white1 and its deposition in the filtrate on long standing should not be mistaken for a second precipitation of the yellow precipitate4 ,the addition of #ge cc4 of ammonium molybdate solution insures the presence of a liberal excess of the reagent1 but the filtrate should be tested as in all quantitative procedures4 ,the precipitation is probably complete in many cases in less than twelve hours2 but it is better1 when practicable1 to allow the solution to stand for this length of time4 ,vigorous shaking or stirring promotes the separation of the precipitate4 ,note #f3 ,the composition of the yellow precipitate undoubtedly varies slightly with varying conditions at the time of its formation4 ,its composition may probably fairly be represented by the formula1 7,,nh#d7c,,po#d4ab,mo,o#c4;,h#b,o1 when precipitated under the conditions prescribed in the procedure4 ,whatever other variations may occur in its composition1 the ratio of #ab ,mo,o#c3a ,p seems to hold1 and this fact is utilized in volumetric processes for the determination of phosphorus1 in which the molybdenum is reduced to a lower oxide and reoxidized by a standard solution of potassium permanganate4 ,in principle1 the procedure is comparable with that described for the determination of iron by permanganate4 ,,precipitation ,,of ,,magnesium ,,ammonium ,,phosphate ,,procedure4--,dissolve the precipitate of phosphomolybdate upon the filter by pouring through it dilute aqueous ammonia 7one volume of dilute ammonia 7sp4 gr4 #j4if7 and three volumes of water1 which should be carefully measured71 and receive the solution in the beaker containing the bulk of the precipitate4 ,the total volume of nitrate and washings should not much exceed #ajj cc4 ,acidify the solution with dilute hydrochloric acid1 and heat it nearly to boiling4 ,calculate the volume of magnesium ammonium chloride solution 7magnesia mixture7 required to precipitate the phosphoric acid1 assuming #dj per cent ,p#b,o#e in the apatite4 ,measure out about #e cc4 in excess of this amount1 and pour it into the acid solution4 ,then add slowly dilute ammonium hydroxide 7#a volume of strong ammonia 7sp4 gr4 #j4ij7 and #i volumes of water71 stirring constantly until a precipitate forms4 ,then add a volume of filtered1 concentrated ammonia 7sp4 gr4 #j4ij7 equal to one third of the volume of liquid in the beaker 7,note #a74 ,allow the whole to cool4 ,the precipitated magnesium ammonium phosphate should then be definitely crystalline in appearance 7,note #b74 7,if it is desired to hasten the precipitation1 the solution may be cooled1 first in cold and then in ice-water1 and stirred 6constantly6 for half an hour1 when precipitation will usually be complete47 ,decant the clear liquid through a filter1 and transfer the precipitate to the filter1 using as wash-water a mixture of one volume of concentrated ammonia and three volumes of water4 ,it is not necessary to clean the beaker completely or to wash the precipitate thoroughly at this point1 as it is necessary to purify it by reprecipitation4 ,note #a3 ,magnesium ammonium phosphate is not a wholly insoluble substance1 even under the most favorable analytical conditions4 ,it is least soluble in a liquid containing one fourth of its volume of concentrated aqueous ammonia 7sp4 gr4 #j4ij7 and this proportion should be carefully maintained as prescribed in the procedure4 ,on account of this slight solubility the volume of solutions should be kept as small as possible and the amount of wash-water limited to that absolutely required4 ,a large excess of the magnesium solution tends both to throw out magnesium hydroxide 7shown by a persistently flocculent precipitate7 and to cause the phosphate to carry down molybdic acid4 ,the tendency of the magnesium precipitate to carry down molybdic acid is also increased if the solution is too concentrated4 ,the volume should not be less than #ij cc41 nor more than #abe cc41 at the time of the first precipitation with the magnesia mixture4 ,note #b3 ,the magnesium ammonium phosphate should be perfectly crystalline1 and will be so if the directions are followed4 ,the slow addition of the reagent is essential1 and the stirring not less so4 ,stirring promotes the separation of the precipitate and the formation of larger crystals1 and may therefore be substituted for digestion in the cold4 ,the stirring-rod must not be allowed to scratch the glass1 as the crystals adhere to such scratches and are removed with difficulty4 ,,reprecipitation ,,and ,,ignition ,,of ,,magnesium ,,ammonium ,,phosphate ,a single precipitation of the magnesium compound in the presence of molybdenum compounds rarely yields a pure product4 ,the molybdenum can be removed by solution of the precipitate in acid and precipitation of the molybdenum by sulphureted hydrogen1 after which the magnesium precipitate may be again thrown down4 ,it is usually more satisfactory to dissolve the magnesium precipitate and reprecipitate the phosphate as magnesium ammonium phosphate as described below4 ,,procedure4--,dissolve the precipitate from the filter in a little dilute hydrochloric acid 7sp4 gr4 #a4ab71 allowing the acid solution to run into the beaker in which the original precipitation was made 7,note #a74 ,wash the filter with water until the wash-water shows no test for chlorides1 but avoid an unnecessary amount of wash-water4 ,add to the solution #b cc4 7not more7 of magnesia mixture1 and then dilute ammonium hydroxide solution 7sp4 gr4 #j4if71 drop by drop1 with constant stirring1 until the liquid smells distinctly of ammonia4 ,stir for a few moments and then add a volume of strong ammonia 7sp4 gr4 #j4ij71 equal to one third of the volume of the solution4 ,allow the solution to stand for some hours1 and then filter off the magnesium ammonium phosphate1 which should be distinctly crystalline in character4 ,wash the precipitate with dilute ammonia water1 as prescribed above1 until1 finally1 #c cc4 of the washings1 after acidifying with nitric acid1 show no evidence of chlorides4 ,test both filtrates for complete precipitation by adding a few cubic centimeters of magnesia mixture and allowing them to stand for some time4 ,transfer the moist precipitate to a weighed porcelain or platinum crucible and ignite1 using great care to raise the temperature slowly while drying the filter in the crucible1 and to insure the ready access of oxygen during the combustion of the filter paper1 thus guarding against a possible reduction of the phosphate1 which would result in disastrous consequences both to the crucible1 if of platinum1 and the analysis4 ,do not raise the temperature above moderate redness until the precipitate is white4 7,keep this precaution well in mind47 ,ignite finally at the highest temperature of the ,tirrill burner1 and repeat the heating until the weight is constant4 ,if the ignited precipitate is persistently discolored by particles of unburned carbon1 moisten the mass with a drop or two of concentrated nitric acid and heat cautiously1 finally igniting strongly4 ,the acid will dissolve magnesium pyrophosphate from the surface of the particles of carbon1 which will then burn away4 ,nitric acid also aids as an oxidizing agent in supplying oxygen for the combustion of the carbon4 ,from the weight of magnesium pyrophosphate 7,mg#b,p#b,o#g7 obtained1 calculate the phosphoric anhydride 7,p#b,o#e7 in the sample of apatite4 ,note #a3 ,the ionic change involved in the precipitation of the magnesium compound is ,,po#d--- ,,nh#d ,mg -- ,mg,n,h#d,,po#d4 ,the magnesium ammonium phosphate is readily dissolved by acids1 even those which are no stronger than acetic acid4 ,this is accounted for by the fact that two of the ions into which phosphoric acid may dissociate1 the ,,hpo#d-- or ,h#b,,po#d- ions1 exhibit the characteristics of very weak acids1 in that they show almost no tendency to dissociate further into ,h and ,,po#d-- ions4 ,consequently the ionic changes which occur when the magnesium ammonium phosphate is brought into contact with an acid may be typified by the reaction3 ,h ,mg ,,nh#d ,,po#d--- -- ,mg ,,nh#d ,,hpo#d--2 that is1 the ,,po#d-- ions and the ,h ions lose their identity in the formation of the new ion1 ,,hpo#d--1 and this continues until the magnesium ammonium phosphate is entirely dissolved4 ,note #b3 ,during ignition the magnesium ammonium phosphate loses ammonia and water and is converted into magnesium pyrophosphate3 #b,mg,n,h#d,,po#d -- ,mg#b,p#b,o#g #b,,nh#c ,h#b,o4 ,the precautions mentioned on pages #aaa and #abc must be observed with great care during the ignition of this precipitate4 ,the danger here lies in a possible reduction of the phosphate by the carbon of the filter paper1 or by the ammonia evolved1 which may act as a reducing agent4 ,the phosphorus then attacks and injures a platinum crucible1 and the determination is valueless4 ,,analysis ,,of ,,limestone ,limestones vary widely in composition from a nearly pure marble through the dolomitic limestones1 containing varying amounts of magnesium1 to the impure varieties1 which contain also ferrous and manganous carbonates and siliceous compounds in variable proportions4 ,many other minerals may be inclosed in limestones in small quantities1 and an exact qualitative analysis will often show the presence of sulphides or sulphates1 phosphates1 and titanates1 and the alkali or even the heavy metals4 ,no attempt is made in the following procedures to provide a complete quantitative scheme which would take into account all of these constituents4 ,such a scheme for a complete analysis of a limestone may be found in ,bulletin ,no4 #gjj of the ,united ,states ,geological ,survey4 ,it is assumed that1 for these practice determinations1 a limestone is selected which contains only the more common constituents first enumerated above4 ,,determination ,,of ,,moisture ,the determination of the amount of moisture in minerals or ores is often of great importance4 ,ores which have been exposed to the weather during shipment may have absorbed enough moisture to appreciably affect the results of analysis4 ,since it is essential that the seller and buyer should make their analyses upon comparable material1 it is customary for each analyst to determine the moisture in the sample examined1 and then to calculate the percentages of the various constituents with reference to a sample dried in the air1 or at a temperature a little above #ajj;,c41 which1 unless the ore has undergone chemical change because of the wetting1 should be the same before and after shipment4 ,,procedure4--,spread #be grams of the powdered sample on a weighed watch-glass2 weigh to the nearest #aj milligrams only and heat at #aje;,c42 weigh at intervals of an hour1 after cooling in a desiccator1 until the loss of weight after an hour's heating does not exceed #aj milligrams4 ,it should be noted that a variation in weight of #aj milligrams in a total weight of #be grams is no greater relatively than a variation of #j4a milligram when the sample taken weighs #j4be gram ,,determination ,,of ,,the ,,insoluble ,,matter ,,and ,,silica ,,procedure4--,weigh out two portions of the original powdered sample 7not the dried sample71 of about #e grams each1 into #bej cc4 casseroles1 and cover each with a watch-glass 7,note #a74 ,pour over the powder #be cc4 of water1 and then add #ej cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab7 in small portions1 warming gently1 until nothing further appears to dissolve 7,note #b74 ,evaporate to dryness on the water bath4 ,pour over the residue a mixture of #e cc4 of water and #e cc4 of concentrated hydrochloric acid 7sp4 gr4 #a4b7 and again evaporate to dryness1 and finally heat for at least an hour at a temperature of #aaj;,c4 ,pour over this residue #ej cc4 of dilute hydrochloric acid 7one volume acid 7sp4 gr4 #a4ab7 to five volumes water71 and boil for about five minutes2 then filter and wash twice with the dilute hydrochloric acid1 and then with hot water until free from chlorides4 ,transfer the filter and contents to a porcelain crucible1 dry carefully over a low flame1 and ignite to constant weight4 ,the residue represents the insoluble matter and the silica from any soluble silicates 7,note #c74 ,calculate the combined percentage of these in the limestone4 ,note #a3 ,the relatively large weight 7#e grams7 taken for analysis insures greater accuracy in the determination of the ingredients which are present in small proportions1 and is also more likely to be a representative sample of the material analyzed4 ,note #b3 ,it is plain that the amount of the insoluble residue and also its character will often depend upon the strength of acid used for solution of the limestone4 ,it cannot1 therefore1 be regarded as representing any well-defined constituent1 and its determination is essentially empirical4 ,note #c3 ,it is probable that some of the silicates present are wholly or partly decomposed by the acid1 and the soluble silicic acid must be converted by evaporation to dryness1 and heating1 into white1 insoluble silica4 ,this change is not complete after one evaporation4 ,the heating at a temperature somewhat higher than that of the water bath for a short time tends to leave the silica in the form of a powder1 which promotes subsequent filtration4 ,the siliceous residue is washed first with dilute acid to prevent hydrolytic changes1 which would result in the formation of appreciable quantities of insoluble basic iron or aluminium salts on the filter when washing with hot water4 ,if it is desired to determine the percentage of silica separately1 the ignited residue should be mixed in a platinum crucible with about six times its weight of anhydrous sodium carbonate1 and the procedure given on page #aea should be followed4 ,the filtrate from the silica is then added to the main filtrate from the insoluble residue4 ,,determination ,,of ,,ferric ,,oxide ,,and ,,aluminium ,,oxide 7,,with ,,manganese7 ,,procedure4--,to the filtrate from the insoluble residue add ammonium hydroxide until the solution just smells distinctly of ammonia1 but do not add an excess4 ,then add #e cc4 of saturated bromine water 7,note #a71 and boil for five minutes4 ,if the smell of ammonia has disappeared1 again add ammonium hydroxide in slight excess1 and #c cc4 of bromine water1 and heat again for a few minutes4 ,finally add #aj cc4 of ammonium chloride solution and keep the solution warm until it barely smells of ammonia2 then filter promptly 7,note #b74 ,wash the filter twice with hot water1 then 7after replacing the receiving beaker7 pour through it #be cc4 of hot1 dilute hydrochloric acid 7one volume dilute ,h,cl sp4 gr4 #a4ab to five volumes water74 ,a brown residue insoluble in the acid may be allowed to remain on the filter4 ,wash the filter five times with hot water1 add to the filtrate ammonium hydroxide and bromine water as described above1 and repeat the precipitation4 ,collect the precipitate on the filter already used1 wash it free from chlorides with hot water1 and ignite and weigh as described for ferric hydroxide on page #aaj4 ,the residue after ignition consists of ferric oxide1 alumina1 and mangano-manganic oxide 7,mn#c,o#d71 if manganese is present4 ,these are commonly determined together 7,note #c74 ,calculate the percentage of the combined oxides in the limestone4 ,note #a3 ,the addition of bromine water to the ammoniacal solutions serves to oxidize any ferrous hydroxide to ferric hydroxide and to precipitate manganese as ,mn,o7,,oh7#b4 ,the solution must contain not more than a bare excess of hydroxyl ions 7ammonium hydroxide7 when it is filtered1 on account of the tendency of the aluminium hydroxide to redissolve4 ,the solution should not be strongly ammoniacal when the bromine is added1 as strong ammonia reacts with the bromine1 with the evolution of nitrogen4 ,note #b3 ,the precipitate produced by ammonium hydroxide and bromine should be filtered off promptly1 since the alkaline solution absorbs carbon dioxide from the air1 with consequent partial precipitation of the calcium as carbonate4 ,this is possible even under the most favorable conditions1 and for this reason the iron precipitate is redissolved and again precipitated to free it from calcium4 ,when the precipitate is small1 this reprecipitation may be omitted4 ,note #c3 ,in the absence of significant amounts of manganese the iron and aluminium may be separately determined by fusion of the mixed ignited precipitate1 after weighing1 with about ten times its weight of acid potassium sulphate1 solution of the cold fused mass in water1 and volumetric determination of the iron1 as described on page #ff4 ,the aluminium is then determined by difference1 after subtracting the weight of ferric oxide corresponding to the amount of iron found4 ,if a separate determination of the iron1 aluminium1 and manganese is desired1 the mixed precipitate may be dissolved in acid before ignition1 and the separation effected by special methods 7see1 for example1 ,fay1 6,quantitative ,analyses61 ,first ,edition1 pp4 #ae-ai and #bc-bg74 ,,determination ,,of ,,calcium ,,procedure4--,to the combined filtrates from the double precipitation of the hydroxides just described1 add #e cc4 of dilute ammonium hydroxide 7sp4 gr4 #j4if71 and transfer the liquid to a #ejj cc4 graduated flask1 washing out the beaker carefully4 ,cool to laboratory temperature1 and fill the flask with distilled water until the lowest point of the meniscus is exactly level with the mark on the neck of the flask4 ,carefully remove any drops of water which are on the inside of the neck of the flask above the graduation by means of a strip of filter paper1 make the solution uniform by pouring it out into a dry beaker and back into the flask several times4 ,measure off one fifth of this solution as follows 7,note #a73 ,pour into a #ajj cc4 graduated flask about #aj cc4 of the solution1 shake the liquid thoroughly over the inner surface of the small flask1 and pour it out4 ,repeat the same operation4 ,fill the #ajj cc4 flask until the lowest point of the meniscus is exactly level with the mark on its neck1 remove any drops of solution from the upper part of the neck with filter paper1 and pour the solution into a beaker 7#djj-ejj cc474 ,wash out the flask with small quantities of water until it is clean1 adding these to the #ajj cc4 of solution4 ,when the duplicate portion of #ajj cc4 is measured out from the solution1 remember that the flask must be rinsed out twice with that solution1 as prescribed above1 before the measurement is made4 7,a #ajj cc4 pipette may be used to measure out the aliquot portions1 if preferred47 ,dilute each of the measured portions to #bej cc4 with distilled water1 heat the whole to boiling1 and add ammonium oxalate solution slowly in moderate excess1 stirring well4 ,boil for two minutes2 allow the precipitated calcium oxalate to settle for a half-hour1 and decant through a filter4 ,test the filtrate for complete precipitation by adding a few cubic centimeters of the precipitant1 allowing it to stand for fifteen minutes4 ,if no precipitate forms1 make the solution slightly acid with hydrochloric acid 7,note #b72 see that it is properly labeled and reserve it to be combined with the filtrate from the second calcium oxalate precipitation 7,notes #c and #d74 ,redissolve the calcium oxalate in the beaker with warm hydrochloric acid1 pouring the acid through the filter4 ,wash the filter five times with water1 and finally pour through it aqueous ammonia4 ,dilute the solution to #bej cc41 bring to boiling1 and add #a cc4 ammonium oxalate solution 7,note #e7 and ammonia in slight excess2 boil for two minutes1 and set aside for a half-hour4 ,filter off the calcium oxalate upon the filter first used1 and wash free from chlorides4 ,the filtrate should be made barely acid with hydrochloric acid and combined with the filtrate from the first precipitation4 ,begin at once the evaporation of the solutions for the determination of magnesium as described below4 ,the precipitate of calcium oxalate may be converted into calcium oxide by ignition without previous drying4 ,after burning the filter1 it may be ignited for three quarters of an hour in a platinum crucible at the highest heat of the ,bunsen or ,tirrill burner1 and finally for ten minutes at the blast lamp 7,note #f74 ,repeat the heating over the blast lamp until the weight is constant4 ,as the calcium oxide absorbs moisture from the air1 it must 7after cooling7 be weighed as rapidly as possible4 ,the precipitate may1 if preferred1 be placed in a weighted porcelain crucible4 ,after burning off the filter and heating for ten minutes the calcium precipitate may be converted into calcium sulphate by placing #b cc4 of dilute sulphuric acid in the crucible 7cold71 heating the covered crucible very cautiously over a low flame to drive off the excess of acid1 and finally at redness to constant weight 7,note #g74 ,from the weight of the oxide or sulphate1 calculate the percentage of the calcium 7,ca7 in the limestone1 remembering that only one fifth of the total solution is used for this determination4 ,note #a3 ,if the calcium were precipitated from the entire solution1 the quantity of the precipitate would be greater than could be properly treated4 ,the solution is1 therefore1 diluted to a definite volume 7#ejj cc471 and exactly one fifth 7#ajj cc47 is measured off in a graduated flask or by means of a pipette4 ,note #b3 ,the filtrate from the calcium oxalate should be made slightly acid immediately after filtration1 in order to avoid the solvent action of the alkaline liquid upon the glass4 ,note #c3 ,the accurate quantitative separation of calcium and magnesium as oxalates requires considerable care4 ,the calcium precipitate usually carries down with it some magnesium1 and this can best be removed by redissolving the precipitate after filtration1 and reprecipitation in the presence of only the small amount of magnesium which was included in the first precipitate4 ,when1 however1 the proportion of magnesium is not very large1 the second precipitation of the calcium can usually be avoided by precipitating it from a rather dilute solution 7#hjj cc4 or so7 and in the presence of a considerable excess of the precipitant1 that is1 rather more than enough to convert both the magnesium and calcium into oxalates4 ,note #d3 ,the ionic changes involved in the precipitation of calcium as oxalate are exceedingly simple1 and the principles discussed in connection with the barium sulphate precipitation on page #aac also apply here4 ,the reaction is ,c#b,o#d-- ,ca -- ,ca,c#b,o#d4 ,calcium oxalate is nearly insoluble in water1 and only very slightly soluble in acetic acid1 but is readily dissolved by the strong mineral acids4 ,this behavior with acids is explained by the fact that oxalic acid is a stronger acid than acetic acid2 when1 therefore1 the oxalate is brought into contact with the latter there is almost no tendency to diminish the concentration of ,c#b,o#d-- ions by the formation of an acid less dissociated than the acetic acid itself1 and practically no solvent action ensues4 ,when a strong mineral acid is present1 however1 the ionization of the oxalic acid is much reduced by the high concentration of the ,h ions from the strong acid1 the formation of the undissociated acid lessens the concentration of the ,c#b,o#d-- ions in solution1 more of the oxalate passes into solution to re-establish equilibrium1 and this process repeats itself until all is dissolved4 ,the oxalate is immediately reprecipitated from such a solution on the addition of ,,oh- ions1 which1 by uniting with the ,h ions of the acids 7both the mineral acid and the oxalic acid7 to form water1 leave the ,ca and ,c#b,o#d-- ions in the solution to recombine to form ,ca,c#b,o#d1 which is precipitated in the absence of the ,h ions4 ,it is well at this point to add a small excess of ,c#b,o#d-- ions in the form of ammonium oxalate to decrease the solubility of the precipitate4 ,the oxalate precipitate consists mainly of ,ca,c#b,o#d4;,h#b,o when thrown down4 ,note #e3 ,the small quantity of ammonium oxalate solution is added before the second precipitation of the calcium oxalate to insure the presence of a slight excess of the reagent1 which promotes the separation of the calcium compound4 ,note #f3 ,on ignition the calcium oxalate loses carbon dioxide and carbon monoxide1 leaving calcium oxide3 ,ca,c#b,o#d4;,h#b,o -- ,ca,o ,,co#b ,,co ,h#b,o4 ,for small weights of the oxalate 7#j4f gram or less7 this reaction may be brought about in a platinum crucible at the highest temperature of a ,tirrill burner1 but it is well to ignite larger quantities than this over the blast lamp until the weight is constant4 ,note #g3 ,the heat required to burn the filter1 and that subsequently applied as described1 will convert most of the calcium oxalate to calcium carbonate1 which is changed to sulphate by the sulphuric acid4 ,the reactions involved are ,ca,c#b,o#d -- ,ca,c,o#c ,,co1 ,ca,c,o#c ,h#b,,so#d -- ,ca,s,o#d ,h#b,o ,,co#b4 ,if a porcelain crucible is employed for ignition1 this conversion to sulphate is to be preferred1 as a complete conversion to oxide is difficult to accomplish4 ,note #h3 ,the determination of the calcium may be completed volumetrically by washing the calcium oxalate precipitate from the filter into dilute sulphuric acid1 warming1 and titrating the liberated oxalic acid with a standard solution of potassium permanganate as described on page #gb4 ,when a considerable number of analyses are to be made1 this procedure will save much of the time otherwise required for ignition and weighing4 ,,determination ,,of ,,magnesium ,,procedure4--,evaporate the acidified filtrates from the calcium precipitates until the salts begin to crystallize1 but do 6not6 evaporate to dryness 7,note #a74 ,dilute the solution cautiously until the salts are brought into solution1 adding a little acid if the solution has evaporated to very small volume4 ,the solution should be carefully examined at this point and must be filtered if a precipitate has appeared4 ,heat the clear solution to boiling2 remove the burner and add #be cc4 of a solution of disodium phosphate4 ,then add slowly dilute ammonia 7#a volume strong ammonia 7sp4 gr4 #j4ij7 and #i volumes water7 as long as a precipitate continues to form4 ,finally1 add a volume of concentrated ammonia 7sp4 gr4 #j4ij7 equal to one third of the volume of the solution1 and allow the whole to stand for about twelve hours4 ,decant the solution through a filter1 wash it with dilute ammonia water1 proceeding as prescribed for the determination of phosphoric anhydride on page #abb1 including2 the reprecipitation 7,note #b71 except that #c cc4 of disodium phosphate solution are added before the reprecipitation of the magnesium ammonium phosphate instead of the magnesia mixture there prescribed4 ,from the weight of the pyrophosphate1 calculate the percentage of magnesium oxide 7,mg,o7 in the sample of limestone4 ,remember that the pyrophosphate finally obtained is from one fifth of the original sample4 ,note #a3 ,the precipitation of the magnesium should be made in as small volume as possible1 and the ratio of ammonia to the total volume of solution should be carefully provided for1 on account of the relative solubility of the magnesium ammonium phosphate4 ,this matter has been fully discussed in connection with the phosphoric anhydride determination4 ,note #b3 ,the first magnesium ammonium phosphate precipitate is rarely wholly crystalline1 as it should be1 and is not always of the proper composition when precipitated in the presence of such large amounts of ammonium salts4 ,the difficulty can best be remedied by filtering the precipitate and 7without washing it7 redissolving in a small quantity of hydrochloric acid1 from which it may be again thrown down by ammonia after adding a little disodium phosphate solution4 ,if the flocculent character was occasioned by the presence of magnesium hydroxide1 the second precipitation1 in a smaller volume containing fewer salts1 will often result more favorably4 ,the removal of iron or alumina from a contaminated precipitate is a matter involving a long procedure1 and a redetermination of the magnesium from a new sample1 with additional precautions1 is usually to be preferred4 ,,determination ,,of ,,carbon ,,dioxide 6,absorption ,apparatus6 ,illustration3 ,fig4 #c ,the apparatus required for the determination of the carbon dioxide should be arranged as shown in the cut 7,fig4 #c74 ,the flask 7,a7 is an ordinary wash bottle1 which should be nearly filled with dilute hydrochloric acid 7#ajj cc4 acid 7sp4 gr4 #a4ab7 and #bjj cc4 of water74 ,the flask is connected by rubber tubing 7a7 with the glass tube 7b7 leading nearly to the bottom of the evolution flask 7,b7 and having its lower end bent upward and drawn out to small bore1 so that the carbon dioxide evolved from the limestone cannot bubble back into 7b74 ,the evolution flask should preferably be a wide-mouthed ,soxhlet extraction flask of about #aej cc4 capacity because of the ease with which tubes and stoppers may be fitted into the neck of a flask of this type4 ,the flask should be fitted with a two-hole rubber stopper4 ,the condenser 7,c7 may consist of a tube with two or three large bulbs blown in it1 for use as an air-cooled condenser1 or it may be a small water-jacketed condenser4 ,the latter is to be preferred if a number of determinations are to be made in succession4 ,a glass delivery tube 7c7 leads from the condenser to the small ,u-tube 7,d7 containing some glass beads or small pieces of glass rod and #c cc4 of a saturated solution of silver sulphate1 with #c cc4 of concentrated sulphuric acid 7sp4 gr4 #a4hd74 ,the short rubber tubing 7d7 connects the first ,u-tube to a second ,u-tube 7,e7 which is filled with small dust-free lumps of dry calcium chloride1 with a small1 loose plug of cotton at the top of each arm4 ,both tubes should be closed by cork stoppers1 the tops of which are cut off level with1 or preferably forced a little below1 the top of the ,u-tube1 and then neatly sealed with sealing wax4 ,the carbon dioxide may be absorbed in a tube containing soda lime 7,f7 or in a ,geissler bulb 7,f'7 containing a concentrated solution of potassium hydroxide 7,note #b74 ,the tube 7,f7 is a glass-stoppered side-arm ,u-tube in which the side toward the evolution flask and one half of the other side are filled with small1 dust-free lumps of soda lime of good quality 7,note #c74 ,since soda lime contains considerable moisture1 the other half of the right side of the tube is filled with small lumps of dry1 dust-free calcium chloride to retain the moisture from the soda lime4 ,loose plugs of cotton are placed at the top of each arm and between the soda lime and the calcium chloride4 ,the ,geissler bulb 7,f'71 if used1 should be filled with potassium hydroxide solution 7#a part of solid potassium hydroxide dissolved in two parts of water7 until each small bulb is about two thirds full 7,note #d74 ,a small tube containing calcium chloride is connected with the ,geissler bulb proper by a ground joint and should be wired to the bulb for safety4 ,this is designed to retain any moisture from the hydroxide solution4 ,a piece of clean1 fine copper wire is so attached to the bulb that it can be hung from the hook above a balance pan1 or other support4 ,the small bottle 7,g7 with concentrated sulphuric acid 7sp4 gr4 #a4hd7 is so arranged that the tube 7f7 barely dips below the surface4 ,this will prevent the absorption of water vapor by 7,f7 or 7,f'7 and serves as an aid in regulating the flow of air through the apparatus4 7,h7 is an aspirator bottle of about four liters capacity1 filled with water2 7k7 is a safety tube and a means of refilling 7,h72 7h7 is a screw clamp1 and 7,k7 a ,u-tube filled with soda lime4 ,note #a3 ,the air current1 which is subsequently drawn through the apparatus1 to sweep all of the carbon dioxide into the absorption apparatus1 is likely to carry with it some hydrochloric acid from the evolution flask4 ,this acid is retained by the silver sulphate solution4 ,the addition of concentrated sulphuric acid to this solution reduces its vapor pressure so far that very little water is carried on by the air current1 and this slight amount is absorbed by the calcium chloride in 7,e74 ,as the calcium chloride frequently contains a small amount of a basic material which would absorb carbon dioxide1 it is necessary to pass carbon dioxide through 7,e7 for a short time and then drive all the gas out with a dry air current for thirty minutes before use4 ,note #b3 ,soda-lime absorption tubes are to be preferred if a satisfactory quality of soda lime is available and the number of determinations to be made successively is small4 ,the potash bulbs will usually permit of a larger number of successive determinations without refilling1 but they require greater care in handling and in the analytical procedure4 ,note #c3 ,soda lime is a mixture of sodium and calcium hydroxides4 ,both combine with carbon dioxide to form carbonates1 with the evolution of water4 ,considerable heat is generated by the reaction1 and the temperature of the tube during absorption serves as a rough index of the progress of the reaction through the mass of soda lime4 ,it is essential that soda lime of good quality for analytical purposes should be used4 ,the tube should not contain dust1 as this is likely to be swept away4 ,note #d3 ,the solution of the hydroxide for use in the ,geissler bulb must be highly concentrated to insure complete absorption of the carbon dioxide and also to reduce the vapor pressure of the solution1 thus lessening the danger of loss of water with the air which passes through the bulbs4 ,the small quantity of moisture which is then carried out of the bulbs is held by the calcium chloride in the prolong tube4 ,the best form of absorption bulb is that to which the prolong tube is attached by a ground glass joint4 ,after the potassium hydroxide is approximately half consumed in the first bulb of the absorption apparatus1 potassium bicarbonate is formed1 and as it is much less soluble than the carbonate1 it often precipitates4 ,its formation is a warning that the absorbing power of the hydroxide is much diminished4 6,the ,analysis6 ,,procedure4-- ,weigh out into the flask 7,b7 about #a gram of limestone4 ,cover it with #ae cc4 of water4 ,weigh the absorption apparatus 7,f7 or 7,f'7 accurately after allowing it to stand for #cj minutes in the balance case1 and wiping it carefully with a lintless cloth1 taking care to handle it as little as possible after wiping 7,note #a74 ,connect the absorption apparatus with 7e7 and 7f74 ,if a soda-lime tube is used1 be sure that the arm containing the soda lime is next the tube 7,e7 and that the glass stopcocks are open4 ,to be sure that the whole apparatus is airtight1 disconnect the rubber tube from the flask 7,a71 making sure that the tubes 7a7 and 7b7 do not contain any hydrochloric acid1 close the pinchcocks 7a7 and 7k7 and open 7h74 ,no bubbles should pass through 7,d7 or 7,g7 after a few seconds4 ,when assured that the fittings are tight1 close 7h7 and open 7a7 cautiously to admit air to restore atmospheric pressure4 ,this precaution is essential1 as a sudden inrush of air will project liquid from 7,d7 or 7,f'74 ,reconnect the rubber tube with the flask 7,a74 ,open the pinchcocks 7a7 and 7k7 and blow over about #aj cc4 of the hydrochloric acid from 7,a7 into 7,b74 ,when the action of the acid slackens1 blow over 7slowly7 another #aj cc4 ,the rate of gas evolution should not exceed for more than a few seconds that at which about two bubbles per second pass through 7,g7 7,note #b74 ,repeat the addition of acid in small portions until the action upon the limestone seems to be at an end1 taking care to close 7a7 after each addition of acid 7,note #c74 ,disconnect 7,a7 and connect the rubber tubing with the soda-lime tube 7,k7 and open 7a74 ,then close 7k7 and open 7h71 regulating the flow of water from 7,h7 in such a way that about two bubbles per second pass through 7,g74 ,place a small flame under 7,b7 and 6slowly6 raise the contents to boiling and boil for three minutes4 ,then remove the burner from under 7,b7 and continue to draw air through the apparatus for #bj-cj minutes1 or until 7,h7 is emptied 7,note #d74 ,remove the absorption apparatus1 closing the stopcocks on 7,f7 or stoppering the open ends of 7,f'71 leave the apparatus in the balance case for at least thirty minutes1 wipe it carefully and weigh1 after opening the stopcocks 7or removing plugs74 ,the increase in weight is due to absorption of ,,co#b1 from which its percentage in the sample may be calculated4 ,after cleaning 7,b7 and refilling 7,h71 the apparatus is ready for the duplicate analysis4 ,note #a3 ,the absorption tubes or bulbs have large surfaces on which moisture may collect4 ,by allowing them to remain in the balance case for some time before weighing1 the amount of moisture absorbed on the surface is as nearly constant as practicable during two weighings1 and a uniform temperature is also assured4 ,the stopcocks of the ,u-tube should be opened1 or the plugs used to close the openings of the ,geissler bulb should be removed before weighing in order that the air contents shall always be at atmospheric pressure4 ,note #b3 ,if the gas passes too rapidly into the absorption apparatus1 some carbon dioxide may be carried through1 not being completely retained by the absorbents4 ,note #c3 ,the essential ionic changes involved in this procedure are the following3 ,it is assumed that the limestone1 which is typified by calcium carbonate1 is very slightly soluble in water1 and the ions resulting are ,ca and ,,co#c--4 ,in the presence of ,h ions of the mineral acid1 the ,,co#c-- ions form ,h#b;,,co#c4 ,this is not only a weak acid which1 by its formation1 diminishes the concentration of the ,,co#c-- ions1 thus causing more of the carbonate to dissolve to re-establish equilibrium1 but it is also an unstable compound and breaks down into carbon dioxide and water4 ,note #d3 ,carbon dioxide is dissolved by cold water1 but the gas is expelled by boiling1 and1 together with that which is distributed through the apparatus1 is swept out into the absorption bulb by the current of air4 ,this air is purified by drawing it through the tube 7,k7 containing soda lime1 which removes any carbon dioxide which may be in it4 ,,determination ,,of ,,lead1 ,,copper1 ,,iron1 ,,and ,,zinc ,,in ,,brass ,,electrolytic ,,separations 6,general ,discussion6 ,when a direct current of electricity passes from one electrode to another through solutions of electrolytes1 the individual ions present in these solutions tend to move toward the electrode of opposite electrical charge to that which each ion bears1 and to be discharged by that electrode4 ,whether or not such discharge actually occurs in the case of any particular ion depends upon the potential 7voltage7 of the current which is passing through the solution1 since for each ion there is1 under definite conditions1 a minimum potential below which the discharge of the ion cannot be effected4 ,by taking advantage of differences in discharge-potentials1 it is possible to effect separations of a number of the metallic ions by electrolysis1 and at the same time to deposit the metals in forms which admit of direct weighing4 ,in this way the slower procedures of precipitation and filtration may frequently be avoided4 ,the following paragraphs present a brief statement of the fundamental principles and conditions underlying electro-analysis4 ,the total energy of an electric current as it passes through a solution is distributed among three factors1 first1 its potential1 which is measured in volts1 and corresponds to what is called head in a stream of water2 second1 current strength1 which is measured in amperes1 and corresponds to the volume of water passing a cross-section of a stream in a given time interval2 and third1 the resistance of the conducting medium1 which is measured in ohms4 ,the relation between these three factors is expressed by ,ohm's law1 namely1 that 6,i ,e_/,r61 when ,i is current strength1 ,e potential1 and ,r resistance4 ,it is plain that1 for a constant resistance1 the strength of the current and its potential are mutually and directly interdependent4 ,as already stated1 the applied electrical potential determines whether or not deposition of a metal upon an electrode actually occurs4 ,the current strength determines the rate of deposition and the physical characteristics of the deposit4 ,the resistance of the solution is generally so small as to fall out of practical consideration4 ,approximate deposition-potentials have been determined for a number of the metallic elements1 and also for hydrogen and some of the acid-forming radicals4 ,the values given below are those required for deposition from normal solutions at ordinary temperatures with reference to a hydrogen electrode4 ,they must be regarded as approximate1 since several disturbing factors and some secondary reactions render difficult their exact application under the conditions of analysis4 ,they are3 ,zn ,cd ,fe ,ni ,pb ,h ,cu ,sb ,hg ,ag ,,so#d #j4gg #j4db #j4cd #j4cc #j4ac #j -#j4cd -#j4fg -#j4gf -#j4gi #a4ij ,from these data it is evident that in order to deposit copper from a normal solution of copper sulphate a minimum potential equal to the algebraic sum of the deposition-potentials of copper ions and sulphate ions must be applied1 that is1 #a4ef volts4 ,the deposition of zinc from a solution of zinc sulphate would require #b4fg volts1 but1 since the deposition of hydrogen from sulphuric acid solution requires only #a4ij volts1 the quantitative deposition of zinc by electrolysis from a sulphuric acid solution of a zinc salt is not practicable4 ,on the other hand1 silver1 if present in a solution of copper sulphate1 would deposit with the copper4 ,the foregoing examples suffice to illustrate the application of the principle of deposition potentials1 but it must further be noted that the values stated apply to normal solutions of the compounds in question1 that is1 to solutions of considerable concentrations4 ,as the concentration of the ions diminishes1 and hence fewer ions approach the electrodes1 somewhat higher voltages are required to attract and discharge them4 ,from this it follows that the concentrations should be kept as high as possible to effect complete deposition in the least practicable time1 or else the potentials applied must be progressively increased as deposition proceeds4 ,in practice1 the desired result is obtained by starting with small volumes of solution1 using as large an electrode surface as possible1 and by stirring the solution to bring the ions in contact with the electrodes4 ,this is1 in general1 a more convenient procedure than that of increasing the potential of the current during electrolysis1 although that method is also used4 ,as already stated1 those ions in a solution of electrolytes will first be discharged which have the lowest deposition potentials1 and so long as these ions are present around the electrode in considerable concentration they1 almost alone1 are discharged1 but1 as their concentration diminishes1 other ions whose deposition potentials are higher but still within that of the current applied1 will also begin to separate4 ,for example1 from a nitric acid solution of copper nitrate1 the copper ions will first be discharged at the cathode1 but as they diminish in concentration hydrogen ions from the acid 7or water7 will be also discharged4 ,since the hydrogen thus liberated is a reducing agent1 the nitric acid in the solution is slowly reduced to ammonia1 and it may happen that if the current is passed through for a long time1 such a solution will become alkaline4 ,oxygen is liberated at the anode1 but1 since there is no oxidizable substance present around that electrode1 it escapes as oxygen gas4 ,it should be noted that1 in general1 the changes occurring at the cathode are reductions1 while those at the anode are oxidations4 ,for analytical purposes1 solutions of nitrates or sulphates of the metals are preferable to those of the chlorides1 since liberated chlorine attacks the electrodes4 ,in some cases1 as for example1 that of silver1 solution of salts forming complex ions1 like that of the double cyanide of silver and potassium1 yield better metallic deposits4 ,most metals are deposited as such upon the cathode2 a few1 notably lead and manganese1 separate in the form of dioxides upon the anode4 ,it is evidently important that the deposited material should be so firmly adherent that it can be washed1 dried1 and weighed without loss in handling4 ,to secure these conditions it is essential that the current density 7that is1 the amount of current per unit of area of the electrodes7 shall not be too high4 ,in prescribing analytical conditions it is customary to state the current strength in normal densities expressed in amperes per #ajj sq4 cm4 of electrode surface1 as1 for example1 ,n4,d#ajj #b amps4 ,if deposition occurs too rapidly1 the deposit is likely to be spongy or loosely adherent and falls off on subsequent treatment4 ,this places a practical limit to the current density to be employed1 for a given electrode surface4 ,the cause of the unsatisfactory character of the deposit is apparently sometimes to be found in the coincident liberation of considerable hydrogen and sometimes in the failure of the rapidly deposited material to form a continuous adherent surface4 ,the effect of rotating electrodes upon the character of the deposit is referred to below4 ,the negative ions of an electrolyte are attracted to the anode and are discharged on contact with it4 ,anions such as the chloride ion yield chlorine atoms1 from which gaseous chlorine molecules are formed and escape4 ,the radicals which compose such ions as ,,no#c- or ,,so#d-- are not capable of independent existence after discharge1 and break down into oxygen and ,n#b,o#e and ,,so#c respectively4 ,the oxygen escapes and the anhydrides1 reacting with water1 re-form nitric and sulphuric acids4 ,the law of ,faraday expresses the relation between current strength and the quantities of the decomposition products which1 under constant conditions1 appear at the electrodes1 namely1 that a given quantity of electricity1 acting for a given time1 causes the separation of chemically equivalent quantities of the various elements or radicals4 ,for example1 since #ajg4id grams of silver is equivalent to #a4jjh grams of hydrogen1 and that in turn to #h grams of oxygen1 or #ca4gh grams of copper1 the quantity of electricity which will cause the deposit of #ajg4id grams of silver in a given time will also separate the weights just indicated of the other substances4 ,experiments show that a current of one ampere passing for one second1 i4e41 a coulomb of electricity1 causes the deposition of #j4jjaaah gram of silver from a normal solution of a silver salt4 ,the number of coulombs required to deposit #ajg4id grams is #ajg4id_/j4jjaaah or #if1eej and the same number of coulombs will also cause the separation of #a4jjh grams of hydrogen1 #h grams of oxygen or #ca4gh grams of copper4 ,while it might at first appear that ,faraday's law could thus be used as a basis for the calculation of the time required for the deposition of a given quantity of an electrolyte from solution1 it must be remembered that the law expresses what occurs when the concentration of the ions in the solution is kept constant1 as1 for example1 when the anode in a silver salt solution is a plate of metallic silver4 ,under the conditions of electro-analysis the concentration of the ions is constantly diminishing as deposition proceeds and the time actually required for complete deposition of a given weight of material by a current of constant strength is1 therefore1 greater than that calculated on the basis of the law as stated above4 ,the electrodes employed in electro-analysis are almost exclusively of platinum1 since that metal alone satisfactorily resists chemical action of the electrolytes1 and can be dried and weighed without change in composition4 ,the platinum electrodes may be used in the form of dishes1 foil or gauze4 ,the last1 on account of the ease of circulation of the electrolyte1 its relatively large surface in proportion to its weight and the readiness with which it can be washed and dried1 is generally preferred4 ,many devices have been described by the use of which the electrode upon which deposition occurs can be mechanically rotated4 ,this has an effect parallel to that of greatly increasing the electrode surface and also provides a most efficient means of stirring the solution4 ,with such an apparatus the amperage may be increased to #e or even #aj amperes and depositions completed with great rapidity and accuracy4 ,it is desirable1 whenever practicable1 to provide a rotating or stirring device1 since1 for example1 the time consumed in the deposition of the amount of copper usually found in analysis may be reduced from the #bj to #bd hours required with stationary electrodes1 and unstirred solutions1 to about one half hour4 ,,determination ,,of ,,copper ,,and ,,lead ,,procedure4--,weigh out two portions of about #j4e gram each 7,note #a7 into tall1 slender lipless beakers of about #ajj cc4 capacity4 ,dissolve the metal in a solution of #e cc4 of dilute nitric acid 7sp4 gr4 #a4bj7 and #e cc4 of water1 heating gently1 and keeping the beaker covered4 ,when the sample has all dissolved 7,note #b71 wash down the sides of the beaker and the bottom of the watch-glass with water and dilute the solution to about #ej cc4 ,carefully heat to boiling and boil for a minute or two to expel nitrous fumes4 ,meanwhile1 four platinum electrodes1 two anodes and two cathodes1 should be cleaned by dipping in dilute nitric acid1 washing with water and finally with #ie per cent alcohol 7,note #c74 ,the alcohol may be ignited and burned off4 ,the electrodes are then cooled in a desiccator and weighed4 ,connect the electrodes with the binding posts 7or other device for connection with the electric circuit7 in such a way that the copper will be deposited upon the electrode with the larger surface1 which is made the cathode4 ,the beaker containing the solution should then be raised into place from below the electrodes until the latter reach nearly to the bottom of the beaker4 ,the support for the beaker must be so arranged that it can be easily raised or lowered4 ,if the electrolytic apparatus is provided with a mechanism for the rotation of the electrode or stirring of the electrolyte1 proceed as follows3 ,arrange the resistance in the circuit to provide a direct current of about one ampere4 ,pass this current through the solution to be electrolyzed1 and start the rotating mechanism4 ,keep the beaker covered as completely as possible1 using a split watch-glass 7or other device7 to avoid loss by spattering4 ,when the solution is colorless1 which is usually the case after about #ce minutes1 rinse off the cover glass1 wash down the sides of the beaker1 add about #j4cj gram of urea and continue the electrolysis for another five minutes 7,notes #d and #e74 ,if stationary electrodes are employed1 the current strength should be about #j4a ampere1 which may1 after #ab to #ae hours1 be increased to #j4b ampere4 ,the time required for complete deposition is usually from #bj to #bd hours4 ,it is advisable to add #e cc4 of nitric acid 7sp4 gr4 #a4b7 if the electrolysis extends over this length of time4 ,no urea is added in this case4 ,when the deposition of the copper appears to be complete1 stop the rotating mechanism and slowly lower the beaker with the left hand1 directing at the same time a stream of water from a wash bottle on both electrodes4 ,remove the beaker1 shut off the current1 and1 if necessary1 complete the washing of the electrodes 7,note #f74 ,rinse the electrodes cautiously with alcohol and heat them in a hot closet until the alcohol has just evaporated1 but no longer1 since the copper is likely to oxidize at the higher temperature4 7,the alcohol may be removed by ignition if care is taken to keep the electrodes in motion in the air so that the copper deposit is not too strongly heated at any one point47 ,test the solution in the beaker for copper as follows1 remembering that it is to be used for subsequent determinations of iron and zinc3 ,remove about #e cc4 and add a slight excess of ammonia4 ,compare the mixture with some distilled water1 holding both above a white surface4 ,the solution should not show any tinge of blue4 ,if the presence of copper is indicated1 add the test portion to the main solution1 evaporate the whole to a volume of about #ajj cc41 and again electrolyze with clean electrodes 7,note #g74 ,after cooling the electrodes in a desiccator1 weigh them and from the weight of copper on the cathode and of lead dioxide 7,pb,o#b7 on the anode1 calculate the percentage of copper 7,cu7 and of lead 7,pb7 in the brass4 ,note #a3 ,it is obvious that the brass taken for analysis should be untarnished1 which can be easily assured1 when wire is used1 by scouring with emery4 ,if chips or borings are used1 they should be well mixed1 and the sample for analysis taken from different parts of the mixture4 ,note #b3 ,if a white residue remains upon treatment of the alloy with nitric acid1 it indicates the presence of tin4 ,the material is not1 therefore1 a true brass4 ,this may be treated as follows3 ,evaporate the solution to dryness1 moisten the residue with #e cc4 of dilute nitric acid 7sp4 gr4 #a4b7 and add #ej cc4 of hot water4 ,filter off the meta-stannic acid1 wash1 ignite in porcelain and weigh as ,sn,o#b4 ,this oxide is never wholly free from copper and must be purified for an exact determination4 ,if it does not exceed #b per cent of the alloy1 the quantity of copper which it contains may usually be neglected4 ,note #c3 ,the electrodes should be freed from all greasy matter before using1 and those portions upon which the metal will deposit should not be touched with the fingers after cleaning4 ,note #d3 ,of the ions in solution1 the ,h1 ,cu1 ,zn1 and ,fe ions tend to move toward the cathode4 ,the ,,no#c- ions and the lead1 probably in the form of ,pb,o#b-- ions1 move toward the anode4 ,at the cathode the ,cu ions are discharged and plate out as metallic copper4 ,this alone occurs while the solution is relatively concentrated4 ,later on1 ,h ions are also discharged4 ,in the presence of considerable quantities of ,h ions1 as in this acid solution1 no ,zn or ,fe ions are discharged because of their greater deposition potentials4 ,at the anode the lead is deposited as ,pb,o#b and oxygen is evolved4 ,for the reasons stated on page #ada care must be taken that the solution does not become alkaline if the electrolysis is long continued4 ,note #e3 ,urea reacts with nitrous acid1 which may be formed in the solution as a result of the reducing action of the liberated hydrogen4 ,its removal promotes the complete precipitation of the copper4 ,the reaction is ,,co7,,nh#b7b #b;,,hno#b -- ,,co#b #b,n#b #c;,h#b,o4 ,note #f3 ,the electrodes must be washed nearly or quite free from the nitric acid solution before the circuit is broken to prevent re-solution of the copper4 ,if several solutions are connected in the same circuit it is obvious that some device must be used to close the circuit as soon as the beaker is removed4 ,note #g3 ,the electrodes upon which the copper has been deposited may be cleaned by immersion in warm nitric acid4 ,to remove the lead dioxide1 add a few crystals of oxalic acid to the nitric acid4 ,,determination ,,of ,,iron ,most brasses contain small percentages of iron 7usually not over #j4a per cent7 which1 unless removed1 is precipitated as phosphate and weighed with the zinc4 ,,procedure4--,to the solution from the precipitation of copper and lead by electrolysis1 add dilute ammonia 7sp4 gr4 #j4if7 until the precipitate of zinc hydroxide which first forms re-dissolves1 leaving only a slight red precipitate of ferric hydroxide4 ,filter off the iron precipitate1 using a washed filter1 and wash five times with hot water4 ,test a portion of the last washing with a dilute solution of ammonium sulphide to assure complete removal of the zinc4 ,the precipitate may then be ignited and weighed as ferric oxide1 as described on page #aaj4 ,calculate the percentage of iron 7,fe7 in the brass4 ,,determination ,,of ,,zinc ,,procedure4--,acidify the filtrate from the iron determination with dilute nitric acid4 ,concentrate it to #aej cc4 ,add to the cold solution dilute ammonia 7sp4 gr4 #j4if7 cautiously until it barely smells of ammonia2 then add 6one drop6 of a dilute solution of litmus 7,note #a71 and drop in1 with the aid of a dropper1 dilute nitric acid until the blue of the litmus just changes to red4 ,it is important that this point should not be overstepped4 ,heat the solution nearly to boiling and pour into it slowly a filtered solution of di-ammonium hydrogen phosphate#a containing a weight of the phosphate about equal to twelve times that of the zinc to be precipitated4 7,for this calculation the approximate percentage of zinc is that found by subtracting the sum of the percentages of the copper1 lead and iron from #ajj per cent47 ,keep the solution just below boiling for fifteen minutes1 stirring frequently 7,note #b74 ,if at the end of this time the amorphous precipitate has become crystalline1 allow the solution to cool for about four hours1 although a longer time does no harm 7,note #c71 and filter upon an asbestos filter in a porcelain ,gooch crucible4 ,the filter is prepared as described on page #ajc1 and should be dried to constant weight at #aje;,c4 ,footnote #a3 ,the ammonium phosphate which is commonly obtainable contains some mono-ammonium salt1 and this is not satisfactory as a precipitant4 ,it is advisable1 therefore1 to weigh out the amount of the salt required1 dissolve it in a small volume of water1 add a drop of phenolphthalein solution1 and finally add dilute ammonium hydroxide solution cautiously until the solution just becomes pink1 but do not add an excess4 ,wash the precipitate until free from sulphates with a warm #a per cent solution of the di-ammonium phosphate1 and then five times with #ej per cent alcohol 7,note #d74 ,dry the crucible and precipitate for an hour at #aje;,c41 and finally to constant weight 7,note #e74 ,the filtrate should be made alkaline with ammonia and tested for zinc with a few drops of ammonium sulphide1 allowing it to stand 7,notes #f1 #g and #h74 ,from the weight of the zinc ammonium phosphate 7,zn,n,h#d,,po#d7 calculate the percentage of the zinc 7,zn7 in the brass4 ,note #a3 ,the zinc ammonium phosphate is soluble both in acids and in ammonia4 ,it is1 therefore1 necessary to precipitate the zinc in a nearly neutral solution1 which is more accurately obtained by adding a drop of a litmus solution to the liquid than by the use of litmus paper4 ,note #b3 ,the precipitate which first forms is amorphous1 and may have a variable composition4 ,on standing it becomes crystalline and then has the composition ,zn,n,h#d,,po#d4 ,the precipitate then settles rapidly and is apt to occasion bumping if the solution is heated to boiling4 ,stirring promotes the crystallization4 ,note #c3 ,in a carefully neutralized solution containing a considerable excess of the precipitant1 and also ammonium salts1 the separation of the zinc is complete after standing four hours4 ,the ionic changes connected with the precipitation of the zinc as zinc ammonium phosphate are similar to those described for magnesium ammonium phosphate1 except that the zinc precipitate is soluble in an excess of ammonium hydroxide1 probably as a result of the formation of complex ions of the general character ,zn7,,nh#c7d4 ,note #d3 ,the precipitate is washed first with a dilute solution of the phosphate to prevent a slight decomposition of the precipitate 7as a result of hydrolysis7 if hot water alone is used4 ,the alcohol is added to the final wash-water to promote the subsequent drying4 ,note #e3 ,the precipitate may be ignited and weighed as ,zn#b,p#b,o#g1 by cautiously heating the porcelain ,gooch crucible within a nickel or iron crucible1 used as a radiator4 ,the heating must be very slow at first1 as the escaping ammonia may reduce the precipitate if it is heated too quickly4 ,note #f3 ,if the ammonium sulphide produced a distinct precipitate1 this should be collected on a small filter1 dissolved in a few cubic centimeters of dilute nitric acid1 and the zinc reprecipitated as phosphate1 filtered off1 dried1 and weighed1 and the weight added to that of the main precipitate4 ,note #g3 ,it has been found that some samples of asbestos are acted upon by the phosphate solution and lose weight4 ,an error from this source may be avoided by determining the weight of the crucible and filter after weighing the precipitate4 ,for this purpose the precipitate may be dissolved in dilute nitric acid1 the asbestos washed thoroughly1 and the crucible reweighed4 ,note #h4 ,the details of this method of precipitation of zinc are fully discussed in an article by ,dakin1 6,ztschr4 ,anal4 ,chem461 #ci 7#aijj71 #bgc4 ,,determination ,,of ,,silica ,,in ,,silicates ,of the natural silicates1 or artificial silicates such as slags and some of the cements1 a comparatively few can be completely decomposed by treatment with acids1 but by far the larger number require fusion with an alkaline flux to effect decomposition and solution for analysis4 ,the procedure given below applies to silicates undecomposable by acids1 of which the mineral feldspar is taken as a typical example4 ,modifications of the procedure1 which are applicable to silicates which are completely or partially decomposable by acids1 are given in the ,notes on page #aee4 ,,preparation ,,of ,,the ,,sample ,grind about #c grams of the mineral in an agate mortar 7,note #a7 until no grittiness is to be detected1 or1 better1 until it will entirely pass through a sieve made of fine silk bolting cloth4 ,the sieve may be made by placing a piece of the bolting cloth over the top of a small beaker in which the ground mineral is placed1 holding the cloth in place by means of a rubber band below the lip of the beaker4 ,by inverting the beaker over clean paper and gently tapping it1 the fine particles pass through the sieve1 leaving the coarser particles within the beaker4 ,these must be returned to the mortar and ground1 and the process of sifting and grinding repeated until the entire sample passes through the sieve4 ,note #a3 ,if the sample of feldspar for analysis is in the massive or crystalline form1 it should be crushed in an iron mortar until the pieces are about half the size of a pea1 and then transferred to a steel mortar1 in which they are reduced to a coarse powder4 ,a wooden mallet should always be used to strike the pestle of the steel mortar1 and the blows should not be sharp4 ,it is plain that final grinding in an agate mortar must be continued until the whole of the portion of the mineral originally taken has been ground so that it will pass the bolting cloth1 otherwise the sifted portion does not represent an average sample1 the softer ingredients1 if foreign matter is present1 being first reduced to powder4 ,for this reason it is best to start with not more than the quantity of the feldspar needed for analysis4 ,the mineral must be thoroughly mixed after the grinding4 ,,fusion ,,and ,,solution ,,procedure4--,weigh into platinum crucibles two portions of the ground feldspar of about #j4h gram each4 ,weigh on rough balances two portions of anhydrous sodium carbonate1 each amounting to about six times the weight of the feldspar taken for analysis 7,note #a74 ,pour about three fourths of the sodium carbonate into the crucible1 place the latter on a piece of clean1 glazed paper1 and thoroughly mix the substance and the flux by carefully stirring for several minutes with a dry glass rod1 the end of which has been recently heated and rounded in a flame and slowly cooled4 ,the rod may be wiped off with a small fragment of filter paper1 which may be placed in the crucible4 ,place the remaining fourth of the carbonate on the top of the mixture4 ,cover the crucible1 heat it to dull redness for five minutes1 and then gradually increase the heat to the full capacity of a ,bunsen or ,tirrill burner for twenty minutes1 or until a quiet1 liquid fusion is obtained 7,note #b74 ,finally1 heat the sides and cover strongly until any material which may have collected upon them is also brought to fusion4 ,allow the crucible to cool1 and remove the fused mass as directed on page #aaf4 ,disintegrate the mass by placing it in a previously prepared mixture of #ajj cc4 of water and #ej cc4 of dilute hydrochloric acid 7sp4 gr4 #a4ab7 in a covered casserole 7,note #c74 ,clean the crucible and lid by means of a little hydrochloric acid1 adding this acid to the main solution 7,notes #d and #e74 ,note #a3 ,quartz1 and minerals containing very high percentages of silica1 may require eight or ten parts by weight of the flux to insure a satisfactory decomposition4 ,note #b3 ,during the fusion the feldspar1 which1 when pure1 is a silicate of aluminium and either sodium or potassium1 but usually contains some iron1 calcium1 and magnesium1 is decomposed by the alkaline flux4 ,the sodium of the latter combines with the silicic acid of the silicate1 with the evolution of carbon dioxide1 while about two thirds of the aluminium forms sodium aluminate and the remainder is converted into basic carbonate1 or the oxide4 ,the calcium and magnesium1 if present1 are changed to carbonates or oxides4 ,the heat is applied gently to prevent a too violent reaction when fusion first takes place4 ,note #c3 ,the solution of a silicate by a strong acid is the result of the combination of the ,h ions of the acid and the silicate ions of the silicate to form a slightly ionized silicic acid4 ,as a consequence1 the concentration of the silicate ions in the solution is reduced nearly to zero1 and more silicate dissolves to re-establish the disturbed equilibrium4 ,this process repeats itself until all of the silicate is brought into solution4 ,whether the resulting solution of the silicate contains ortho-silicic acid 7,h#d,si,o#d7 or whether it is a colloidal solution of some other less hydrated acid1 such as meta-silicic acid 7,h#b,si,o#c71 is a matter that is still debatable4 ,it is certain1 however1 that the gelatinous material which readily separates from such solutions is of the nature of a hydrogel1 that is1 a colloid which is insoluble in water4 ,this substance when heated to #ajj;,c41 or higher1 is completely dehydrated1 leaving only the anhydride1 ,si,o#b4 ,the changes may be represented by the equation3 ,si,o#c-- #b;,h -- ,h#b,si,o#c -- ,h#b,o ,si,o#b4 ,note #d3 ,a portion of the fused mass is usually projected upward by the escaping carbon dioxide during the fusion4 ,the crucible must therefore be kept covered as much as possible and the lid carefully cleaned4 ,note #e3 ,a gritty residue remaining after the disintegration of the fused mass by acid indicates that the substance has been but imperfectly decomposed4 ,such a residue should be filtered1 washed1 dried1 ignited1 and again fused with the alkaline flux2 or1 if the quantity of material at hand will permit1 it is better to reject the analysis1 and to use increased care in grinding the mineral and in mixing it with the flux4 ,,dehydration ,,and ,,filtration ,,procedure4--,evaporate the solution of the fusion to dryness1 stirring frequently until the residue is a dry powder4 ,moisten the residue with #e cc4 of strong hydrochloric acid 7sp4 gr4 #a4bj7 and evaporate again to dryness4 ,heat the residue for at least one hour at a temperature of #aaj;,c4 7,note #a74 ,again moisten the residue with concentrated hydrochloric acid1 warm gently1 making sure that the acid comes into contact with the whole of the residue1 dilute to about #bjj cc4 and bring to boiling4 ,filter off the silica without much delay 7,note #b71 and wash five times with warm dilute hydrochloric acid 7one part dilute acid 7#a4ab sp4 gr47 to three parts of water74 ,allow the filter to drain for a few moments1 then place a clean beaker below the funnel and wash with water until free from chlorides1 discarding these washings4 ,evaporate the original filtrate to dryness1 dehydrate at #aaj;,c4 for one hour 7,note #c71 and proceed as before1 using a second filter to collect the silica after the second dehydration4 ,wash this filter with warm1 dilute hydrochloric acid 7,note #d71 and finally with hot water until free from chlorides4 ,note #a3 ,the silicic acid must be freed from its combination with a base 7sodium1 in this instance7 before it can be dehydrated4 ,the excess of hydrochloric acid accomplishes this liberation4 ,by disintegrating the fused mass with a considerable volume of dilute acid the silicic acid is at first held in solution to a large extent4 ,immediate treatment of the fused mass with strong acid is likely to cause a semi-gelatinous silicic acid to separate at once and to inclose alkali salts or alumina4 ,a flocculent residue will often remain after the decomposition of the fused mass is effected4 ,this is usually partially dehydrated silicic acid and does not require further treatment at this point4 ,the progress of the dehydration is indicated by the behavior of the solution1 which as evaporation proceeds usually gelatinizes4 ,on this account it is necessary to allow the solution to evaporate on a steam bath1 or to stir it vigorously1 to avoid loss by spattering4 ,note #b3 ,to obtain an approximately pure silica1 the residue after evaporation must be thoroughly extracted by warming with hydrochloric acid1 and the solution freely diluted to prevent1 as far as possible1 the inclosure of the residual salts in the particles of silica4 ,the filtration should take place without delay1 as the dehydrated silica slowly dissolves in hydrochloric acid on standing4 ,note #c3 ,it has been shown by ,hillebrand that silicic acid cannot be completely dehydrated by a single evaporation and heating1 nor by several such treatments1 unless an intermediate filtration of the silica occurs4 ,if1 however1 the silica is removed and the filtrates are again evaporated and the residue heated1 the amount of silica remaining in solution is usually negligible1 although several evaporations and filtrations are required with some silicates to insure absolute accuracy4 ,it is probable that temperatures above #ajj;,c4 are not absolutely necessary to dehydrate the silica2 but it is recommended1 as tending to leave the silica in a better condition for filtration than when the lower temperature of the water bath is used4 ,this1 and many other points in the analysis of silicates1 are fully discussed by ,dr4 ,hillebrand in the admirable monograph on ,the ,analysis of ,silicate and ,carbonate ,rocks1 ,bulletin ,no4 #gjj of the ,united ,states ,geological ,survey4 ,the double evaporation and filtration spoken of above are essential because of the relatively large amount of alkali salts 7sodium chloride7 present after evaporation4 ,for the highest accuracy in the determination of silica1 or of iron and alumina1 it is also necessary to examine for silica the precipitate produced in the filtrate by ammonium hydroxide by fusing it with acid potassium sulphate and solution of the fused mass in water4 ,the insoluble silica is filtered1 washed1 and weighed1 and the weight added to the weight of silica previously obtained4 ,note #d3 ,aluminium and iron are likely to be thrown down as basic salts from hot1 very dilute solutions of their chlorides1 as a result of hydrolysis4 ,if the silica were washed only with hot water1 the solution of these chlorides remaining in the filter after the passage of the original filtrate would gradually become so dilute as to throw down basic salts within the pores of the filter1 which would remain with the silica4 ,to avoid this1 an acid wash-water is used until the aluminium and iron are practically removed4 ,the acid is then removed by water4 ,,ignition ,,and ,,testing ,,of ,,silica ,,procedure4--,transfer the two washed filters belonging to each determination to a platinum crucible1 which need not be previously weighed1 and burn off the filter 7,note #a74 ,ignite for thirty minutes over the blast lamp with the cover on the crucible1 and then for periods of ten minutes1 until the weight is constant4 ,when a constant weight has been obtained1 pour into the crucible about #c cc4 of water1 and then #c cc4 of hydrofluoric acid4 6,this must be done in a hood with a good draft and great care must be taken not to come into contact with the acid or to inhale its fumes 7,note #b674 ,if the precipitate has dissolved in this quantity of acid1 add two drops of concentrated sulphuric acid1 and heat very slowly 7always under the hood7 until all the liquid has evaporated1 finally igniting to redness4 ,cool in a desiccator1 and weigh the crucible and residue4 ,deduct this weight from the previous weight of crucible and impure silica1 and from the difference calculate the percentage of silica in the sample 7,note #c74 ,note #a3 ,the silica undergoes no change during the ignition beyond the removal of all traces of water2 but ,hillebrand 76loc4 cit467 has shown that the silica holds moisture so tenaciously that prolonged ignition over the blast lamp is necessary to remove it entirely4 ,this finely divided1 ignited silica tends to absorb moisture1 and should be weighed quickly4 ,note #b3 ,notwithstanding all precautions1 the ignited precipitate of silica is rarely wholly pure4 ,it is tested by volatilisation of the silica as silicon fluoride after solution in hydrofluoric acid1 and1 if the analysis has been properly conducted1 the residue1 after treatment with the acids and ignition1 should not exceed #a mg4 ,the acid produces ulceration if brought into contact with the skin1 and its fumes are excessively harmful if inhaled4 ,note #c3 ,the impurities are probably weighed with the original precipitate in the form of oxides4 ,the addition of the sulphuric acid displaces the hydrofluoric acid1 and it may be assumed that the resulting sulphates 7usually of iron or aluminium7 are converted to oxides by the final ignition4 ,it is obvious that unless the sulphuric and hydrofluoric acids used are known to leave no residue on evaporation1 a quantity equal to that employed in the analysis must be evaporated and a correction applied for any residue found4 ,note #d3 ,if the silicate to be analyzed is shown by a previous qualitative examination to be completely decomposable1 it may be directly treated with hydrochloric acid1 the solution evaporated to dryness1 and the silica dehydrated and further treated as described in the case of the feldspar after fusion4 ,a silicate which gelatinizes on treatment with acids should be mixed first with a little water1 and the strong acid added in small portions with stirring1 otherwise the gelatinous silicic acid incloses particles of the original silicate and prevents decomposition4 ,the water1 by separating the particles and slightly lessening the rapidity of action1 prevents this difficulty4 ,this procedure is one which applies in general to the solution of fine mineral powders in acids4 ,if a small residue remains undecomposed by the treatment of the silicate with acid1 this may be filtered1 washed1 ignited and fused with sodium carbonate and a solution of the fused mass added to the original acid solution4 ,this double procedure has an advantage1 in that it avoids adding so large a quantity of sodium salts as is required for disintegration of the whole of the silicate by the fusion method4 ,,part ,,iv ,,stoichiometry ,the problems with which the analytical chemist has to deal are not1 as a matter of actual fact1 difficult either to solve or to understand4 ,that they appear difficult to many students is due to the fact that1 instead of understanding the principles which underlie each of the small number of types into which these problems may be grouped1 each problem is approached as an individual puzzle1 unrelated to others already solved or explained4 ,this attitude of mind should be carefully avoided4 ,it is obvious that ability to make the calculations necessary for the interpretation of analytical data is no less important than the manipulative skill required to obtain them1 and that a moderate time spent in the careful study of the solutions of the typical problems which follow may save much later embarrassment4 #a4 ,it is often necessary to calculate what is known as a chemical factor1 or its equivalent logarithmic value called a log factor1 for the conversion of the weight of a given chemical substance into an equivalent weight of another substance4 ,this is1 in reality1 a very simple problem in proportion1 making use of the atomic or molecular weights of the substances in question which are chemically equivalent to each other4 ,one of the simplest cases of this sort is the following3 ,what is the factor for the conversion of a given weight of barium sulphate 7,ba,s,o#d7 into an equivalent weight of sulphur 7,s78 ,the molecular weight of ,ba,s,o#d is #bcc4e4 ,there is one atom of ,s in the molecule and the atomic weight of ,s is #cb4a4 ,the chemical factor is1 therefore1 #cb4a_/bcc4e1 or #j4acge and the weight of ,s corresponding to a given weight of ,ba,s,o#d is found by multiplying the weight of ,ba,s,o#d by this factor4 ,if the problem takes the form1 ,what is the factor for the conversion of a given weight of ferric oxide 7,fe#b,o#c7 into ferrous oxide 7,fe,o71 or of a given weight of mangano-manganic oxide 7,mn#c,o#d7 into manganese 7,mn78 the principle involved is the same1 but it must then be noted that1 in the first instance1 each molecule of ,fe#b,o#c will be equivalent to two molecules of ,fe,o1 and in the second instance that each molecule of ,mn#c,o#d is equivalent to three atoms of ,mn4 ,the respective factors then become 7#b;,fe,o_/,fe#b,o#c7 or 7#adc4f_/aei4f7 and 7#c,mn_/,mn#c,o#d7 or 7#afd4g_/bbh4g74 ,it is obvious that the arithmetical processes involved in this type of problem are extremely simple4 ,it is only necessary to observe carefully the chemical equivalents4 ,it is plainly incorrect to express the ratio of ferrous to ferric oxide as 7,fe,o_/,fe#b,o#c71 since each molecule of the ferric oxide will yield two molecules of the ferrous oxide4 ,mistakes of this sort are easily made and constitute one of the most frequent sources of error4 #b4 ,a type of problem which is slightly more complicated in appearance1 but exactly comparable in principle1 is the following3 ,what is the factor for the conversion of a given weight of ferrous sulphate 7,fe,s,o#d71 used as a reducing agent against potassium permanganate1 into the equivalent weight of sodium oxalate 7,na#b;,c#b,o#d78 ,to determine the chemical equivalents in such an instance it is necessary to inspect the chemical reactions involved4 ,these are3 #aj;,fe,s,o#d #b,k,mn,o#d #h;,h#b,,so#d -- #e;,fe#b7,,so#d7c ,k#b,,so#d #b,mn,s,o#d #h;,h#b,o1 #e,na#b;,c#b,o#d #b,k,mn,o#d #h;,h#b,,so#d -- #e,na#b,,so#d #aj;,,co#b ,k#b,,so#d #b,mn,s,o#d #h;,h#b,o4 ,it is evident that #aj;,fe,s,o#d in the one case1 and #e,na#b;,c#b,o#d in the other1 each react with #b,k,mn,o#d4 ,these molecular quantities are therefore equivalent1 and the factor becomes 7#aj;,fe,s,o#d_/e,na#b;,c#b,o#d7 or 7#b;,fe,s,o#d_/,na#b;,c#b,o#d7 or 7#cjc4h_/acd74 ,again1 let it be assumed that it is desired to determine the factor required for the conversion of a given weight of potassium permanganate 7,k,mn,o#d7 into an equivalent weight of potassium bichromate 7,k#b;,cr#b,o#g71 each acting as an oxidizing agent against ferrous sulphate4 ,the reactions involved are3 #aj;,fe,s,o#d #b,k,mn,o#d #h;,h#b,,so#d -- #e;,fe#b7,,so#d7c ,k#b,,so#d #b,mn,s,o#d #h;,h#b,o1 #f;,fe,s,o#d ,k#b;,cr#b,o#g #g;,h#b,,so#d -- #c;,fe#b7,,so#c7c ,k#b,,so#d ,cr#b7,,so#d7c #g;,h#b,o4 ,an inspection of these equations shows that #b,,kmo#d react with #aj;,fe,s,o#d1 while ,k#b;,cr#b,o#g reacts with #f;,fe,s,o#d4 ,these are not equivalent1 but if the first equation is multiplied by #c and the second by #e the number of molecules of ,fe,s,o#d is then the same in both1 and the number of molecules of ,k,mn,o#d and ,k#b;,cr#b,o#g reacting with these #cj molecules become #f and #e respectively4 ,these are obviously chemically equivalent and the desired factor is expressed by the fraction 7#f,k,mn,o#d_/e,k#b;,cr#b,o#g7 or 7#idh4j_/adga4j74 #c4 ,it is sometimes necessary to calculate the value of solutions according to the principles just explained1 when several successive reactions are involved4 ,such problems may be solved by a series of proportions1 but it is usually possible to eliminate the common factors and solve but a single one4 ,for example1 the amount of ,mn,o#b in a sample of the mineral pyrolusite may be determined by dissolving the mineral in hydrochloric acid1 absorbing the evolved chlorine in a solution of potassium iodide1 and measuring the liberated iodine by titration with a standard solution of sodium thiosulphate4 ,the reactions involved are3 ,mn,o#b #d;,h,cl -- ,mn,cl#b #b;,h#b,o ,cl#b ,cl#b #b,,ki -- ,i#b #b,k,cl ,i#b #b,na#b,s#b,o#c -- #b,na,i ,na#b,s#d,o#f ,assuming that the weight of thiosulphate corresponding to the volume of sodium thiosulphate solution used is known1 what is the corresponding weight of manganese dioxide8 ,from the reactions given above1 the following proportions may be stated3 #b,na#b,s#b,o#c3;,i#b #caf4d3bec4i1 ,i#b3;,cl#b #bec4i3ga1 ,cl#b3,mn,o#b #ga3hf4i4 ,after canceling the common factors1 there remains #b,na#b,s#b,o#c3,mn,o#b #caf4d3hf4i1 and the factor for the conversion of thiosulphate into an equivalent of manganese dioxide is #hf4i_/caf4d4 #d4 ,to calculate the volume of a reagent required for a specific operation1 it is necessary to know the exact reaction which is to be brought about1 and1 as with the calculation of factors1 to keep in mind the molecular relations between the reagent and the substance reacted upon4 ,for example1 to estimate the weight of barium chloride necessary to precipitate the sulphur from #j4a gram of pure pyrite 7,fe,s#b71 the proportion should read #dhh4 #abj4j #b7;,ba,cl#b4b;,h#b,o73,fe,s#b x3#j4a1 where 6x6 represents the weight of the chloride required4 ,each of the two atoms of sulphur will form upon oxidation a molecule of sulphuric acid or a sulphate1 which1 in turn1 will require a molecule of the barium chloride for precipitation4 ,to determine the quantity of the barium chloride required1 it is necessary to include in its molecular weight the water of crystallization1 since this is inseparable from the chloride when it is weighed4 ,this applies equally to other similar instances4 ,if the strength of an acid is expressed in percentage by weight1 due regard must be paid to its specific gravity4 ,for example1 hydrochloric acid 7sp4 gr4 #a4ab7 contains #bc4h per cent ,h,cl 6by weight62 that is1 #j4bfff gram ,h,cl in each cubic centimeter4 #e4 ,it is sometimes desirable to avoid the manipulation required for the separation of the constituents of a mixture of substances by making what is called an indirect analysis4 ,for example1 in the analysis of silicate rocks1 the sodium and potassium present may be obtained in the form of their chlorides and weighed together4 ,if the weight of such a mixture is known1 and also the percentage of chlorine present1 it is possible to calculate the amount of each chloride in the mixture4 ,let it be assumed that the weight of the mixed chlorides is #j4ae gram1 and that it contains #ec per cent of chlorine4 ,the simplest solution of such a problem is reached through algebraic methods4 ,the weight of chlorine is evidently #j4ae x #j4ec1 or #j4jgie gram4 ,let x represent the weight of sodium chloride present and y that of potassium chloride4 ,the molecular weight of ,na,cl is #eh4e and that of ,k,cl is #gd4f4 ,the atomic weight of chlorine is #ce4e4 ,then x y #j4ae 7#ce4e_/eh4e7x 7#ce4e_/gd4f7y #j4jjgie ,solving these equations for x shows the weight of ,na,cl to be #j4jfbe gram4 ,the weight of ,k,cl is found by subtracting this from #j4ae4 ,the above is one of the most common types of indirect analyses4 ,others are more complex but they can be reduced to algebraic expressions and solved by their aid4 ,it should1 however1 be noted that the results obtained by these indirect methods cannot be depended upon for high accuracy1 since slight errors in the determination of the common constituent1 as chlorine in the above mixture1 will cause considerable variations in the values found for the components4 ,they should not be employed when direct methods are applicable1 if accuracy is essential4 ,,problems 7,the reactions necessary for the solution of these problems are either stated with the problem or may be found in the earlier text4 ,in the calculations from which the answers are derived1 the atomic weights given on page #aie have been employed1 using1 however1 only the first decimal but increasing this by #a when the second decimal is #e or above4 ,thus1 #ci4a has been taken as the atomic weight of potassium1 #cb4a for sulphur1 etc4 ,this has been done merely to secure uniformity of treatment1 and the student should remember that it is always well to take into account the degree of accuracy desired in a particular instance in determining the number of decimal places to retain4 ,four-place logarithms were employed in the calculations4 ,where four figures are given in the answer1 the last figure may vary by one or 7rarely7 by two units1 according to the method by which the problem is solved47 ,,volumetric ,,analysis #a4 ,how many grams of pure potassium hydroxide are required for exactly #a liter of normal alkali solution8 6,answer63 #ef4a grams4 #b4 ,calculate the equivalent in grams 7a7 of sulphuric acid as an acid2 7b7 of hydrochloric acid as an acid2 7c7 of oxalic acid as an acid2 7d7 of nitric acid as an acid4 6,answers63 7a7 #di4je2 7b7 #cf4e2 7c7 #fc2 7d7 #fc4 #c4 ,calculate the equivalent in grams of 7a7 potassium hydroxide2 7b7 of sodium carbonate2 7c7 of barium hydroxide2 7d7 of sodium bicarbonate when titrated with an acid4 6,answers63 7a7 #ef4a2 7b7 #ec4h2 7c7 #he4g2 7d7 #hd4 #d4 ,what is the equivalent in grams of ,na#b;,,hpo#d 7a7 as a phosphate2 7b7 as a sodium salt8 6,answers63 7a7 #dg4cc2 7b7 #ga4j4 #e4 ,a sample of aqueous hydrochloric acid has a specific gravity of #a4ab and contains #bc4ha per cent hydrochloric acid by weight4 ,calculate the grams and the milliequivalents of hydrochloric acid 7,h,cl7 in each cubic centimeter of the aqueous acid4 6,answers63 #j4bffg gram2 #g4cjg milliequivalents4 #f4 ,how many cubic centimeters of hydrochloric acid 7sp4 gr4 #a4bj containing #ci4hj per cent ,h,cl by weight7 are required to furnish #cf4de grams of the gaseous compound8 6,answer63 #gf4cc cc4 #g4 ,a given solution contains #j4ajfc equivalents of hydrochloric acid in #igf cc4 ,what is its normal value8 6,answer63 #j4ajhi ,n4 #h4 ,in standardizing a hydrochloric acid solution it is found that #dg4bf cc4 of hydrochloric acid are exactly equivalent to #a4baf grams of pure sodium carbonate1 using methyl orange as an indicator4 ,what is the normal value of the hydrochloric acid8 6,answer63 #j4dhee ,n4 #i4 ,convert #db4ge cc4 of #j4eafb normal hydrochloric acid to the equivalent volume of normal hydrochloric acid4 6,answer63 #bb4jg cc4 #aj4 ,a solution containing #be4bg cc4 of #j4ajfe normal hydrochloric acid is added to one containing #ib4ba cc4 of #j4edca normal sulphuric acid and #ej cc4 of exactly normal potassium hydroxide added from a pipette4 ,is the solution acid or alkaline8 ,how many cubic centimeters of #j4a normal acid or alkali must be added to exactly neutralize the solution8 6,answer63 #bg4f cc4 alkali 7solution is acid74 #aa4 ,by experiment the normal value of a sulphuric acid solution is found to be #j4eagb4 ,of this acid #ci4fe cc4 are exactly equivalent to #ba4gd cc4 of a standard alkali solution4 ,what is the normal value of the alkali8 6,answer63 #j4idcb ,n4 #ab4 ,a solution of sulphuric acid is standardized against a sample of calcium carbonate which has been previously accurately analyzed and found to contain #ib4dd ,ca,c,o#c and no other basic material4 ,the sample weighing #j4gdbc gram was titrated by adding an excess of acid 7#db4db cc47 and titrating the excess with sodium hydroxide solution 7#aa4bb cc474 #a cc4 of acid is equivalent to #j4iigf cc4 of sodium hydroxide4 ,calculate the normal value of each4 6,answers63 ,acid #j4dcih ,n2 alkali #j4ddji ,n4 #ac4 ,given five #aj cc4 portions of #j4a normal hydrochloric acid1 7a7 how many grams of silver chloride will be precipitated by a portion when an excess of silver nitrate is added8 7b7 how many grams of pure anhydrous sodium carbonate 7,na#b;,,co#c7 will be neutralized by a portion of it8 7c7 how many grams of silver will there be in the silver chloride formed when an excess of silver nitrate is added to a portion8 7d7 how many grams of iron will be dissolved to ,fe,cl#b by a portion of it8 7e7 how many grams of magnesium chloride will be formed and how many grams of carbon dioxide liberated when an excess of magnesium carbonate is treated with a portion of the acid8 6,answers63 7a7 #j4adcd2 7b7 #j4jec2 7c7 #j4ajgi2 7d7 #j4jbgi2 7e7 #j4jdgfe1 and #j4jbb4 #ad4 ,if #cj4jj grams of potassium tetroxalate 7,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o7 are dissolved and the solution diluted to exactly #a liter1 and #dj cc4 are neutralized with #bj cc4 of a potassium carbonate solution1 what is the normal value of the carbonate solution8 6,answer63 #j4gjhd ,n4 #ae4 ,how many cubic centimeters of #j4c normal sulphuric acid will be required to neutralize 7a7 #cj cc4 of #j4e normal potassium hydroxide2 7b7 to neutralize #cj cc4 of #j4e normal barium hydroxide2 7c7 to neutralize #bj cc4 of a solution containing #aj4jb grams of potassium bicarbonate per #ajj cc42 7d7 to give a precipitate of barium sulphate weighing #j4dcbj gram8 6,answers63 7a7 #ej cc42 7b7 #ej cc42 7c7 #ff4gc cc42 7d7 #ab4cc cc4 #af4 ,it is desired to dilute a solution of sulphuric acid of which #a cc4 is equivalent to #j4ajbg gram of pure sodium carbonate to make it exactly #a4bej normal4 #gjj cc4 of the solution are available4 ,to what volume must it be diluted8 6,answer63 #ajhd cc4 #ag4 ,given the following data3 #a cc4 of ,na,o,h #a4aag cc4 ,h,cl4 ,the ,h,cl is #j4dhgf ,n4 ,how much water must be added to #ajj cc4 of the alkali to make it exactly #j4e ,n48 6,answer63 #i4j cc4 #ah4 ,what is the normal value of a sulphuric acid solution which has a specific gravity of #a4hci and contains #ie ,h#b,,so#d by weight8 6,answer63 #ce4fa ,n4 #ai4 ,a sample of ,rochelle ,salt 7,k,na,c#d;,h#d,o#f4d;,h#b,o71 after ignition in platinum to convert it to the double carbonate1 is titrated with sulphuric acid1 using methyl orange as an indicator4 ,from the following data calculate the percentage purity of the sample3 ,wt4 sample #j4iejj gram ,h#b,,so#d used #dc4fe cc4 ,na,o,h used #a4gb cc4 #a cc4 ,h#b,,so#d #a4jfd cc4 ,na,o,h ,normal value ,na,o,h #j4acba ,n4 6,answer63 #hg4gb cc4 #bj4 ,one gram of a mixture of #ej sodium carbonate and #ej potassium carbonate is dissolved in water1 and #ag4cf cc4 of #a4jge ,n acid is added4 ,is the resulting solution acid or alkaline8 ,how many cubic centimeters of #a4jge ,n acid or alkali will have to be added to make the solution exactly neutral8 6,answers63 ,acid2 #a4hf cc4 alkali4 #ba4 ,in preparing an alkaline solution for use in volumetric work1 an analyst1 because of shortage of chemicals1 mixed exactly #df4cb grams of pure ,,koh and #bg4fd grams of pure ,na,o,h1 and after dissolving in water1 diluted the solution to exactly one liter4 ,how many cubic centimeters of #a4jbb ,n hydrochloric acid are necessary to neutralize #ej cc4 of the basic solution8 6,answer63 #gd4ah cc4 #bb4 ,one gram of crude ammonium salt is treated with strong potassium hydroxide solution4 ,the ammonia liberated is distilled and collected in #ej cc4 of #j4e ,n acid and the excess titrated with #a4ee cc4 of #j4e ,n sodium hydroxide4 ,calculate the percentage of ,,nh#c in the sample4 6,answer63 #da4ag4 #bc4 ,in titrating solutions of alkali carbonates in the presence of phenolphthalein1 the color change takes place when the carbonate has been converted to bicarbonate4 ,in the presence of methyl orange1 the color change takes place only when the carbonate has been completely neutralized4 ,from the following data1 calculate the percentages of ,na#b;,,co#c and ,na,o,h in an impure mixture4 ,weight of sample1 #a4ejj grams2 ,h,cl 7#j4e ,n7 required for phenolphthalein end-point1 #bh4he cc42 ,h,cl 7#j4e ,n7 required to complete the titration after adding methyl orange1 #bc4he cc4 6,answers63 #f4fg ,na,o,h2 #hd4bh ,na#b;,,co#c4 #bd4 ,a sample of sodium carbonate containing sodium hydroxide weighs #a4agi grams4 ,it is titrated with #j4cj ,n hydrochloric acid1 using phenolphthalein in cold solution as an indicator and becomes colorless after the addition of #dh4af cc4 ,methyl orange is added and #bd4jh cc4 are needed for complete neutralization4 ,what is the percentage of ,na,o,h and ,na#b;,,co#c8 6,answers63 #bd4ej ,na,o,h2 #fd4ib ,na#b;,,co#c4 #be4 ,from the following data1 calculate the percentages of ,na#b;,,co#c and ,na,h,c,o#c in an impure mixture4 ,weight of sample #a4jjj gram2 volume of #j4be ,n hydrochloric acid required for phenolphthalein end-point1 #bf4dj cc42 after adding an excess of acid and boiling out the carbon dioxide1 the total volume of #j4be ,n hydrochloric acid required for phenolphthalein end-point1 #fg4aj cc4 6,answer63 #fi4ie ,na#b;,,co#c2 #cj4jb ,na,h,c,o#c4 #bf4 ,in the analysis of a one-gram sample of soda ash1 what must be the normality of the acid in order that the number of cubic centimeters of acid used shall represent the percentage of carbon dioxide present8 6,answer63 #j4dedd gram4 #bg4 ,what weight of pearl ash must be taken for analysis in order that the number of cubic centimeters of #j4e ,n acid used may be equal to one third the percentage of ,k#b;,,co#c8 6,answer63 #a4aeb grams4 #bh4 ,what weight of cream of tartar must have been taken for analysis in order to have obtained #ig4fj ,,khc#d;,h#d,o#f in an analysis involving the following data3 ,na,o,h used #cj4jf cc42 ,h#b,,so#d solution used #j4ej cc42 #a cc4 ,h#b,,so#d sol4 #j4jbee gram ,ca,c,o#c2 #a cc4 ,h#b,,so#d sol4 #a4jb cc4 ,na,o,h sol48 6,answer63 #b4hdf grams4 #bi4 ,calculate the percentage of potassium oxide in an impure sample of potassium carbonate from the following data3 ,weight of sample #a4jj gram2 ,h,cl sol4 used #ee4ij cc42 ,na,o,h sol4 used #j4db cc42 #a cc4 ,na,o,h sol4 #j4jjhdgc gram of ,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o2 #b cc4 ,h,cl sol4 #e cc4 ,na,o,h sol4 6,answer63 #fe4fh4 #cj4 ,calculate the percentage purity of a sample of calcite 7,ca,c,o#c7 from the following data3 7,standardization72 ,weight of ,h#b;,c#b,o#d4b;,h#b,o #j4bdfj gram2 ,na,o,h solution used #da4jc cc42 ,h,cl solution used #j4fc2 #a cc4 ,na,o,h solution #a4aij cc4 ,h,cl solution4 7,analysis72 ,weight of sample #j4abjj gram2 ,h,cl used #cf4ch cc42 ,na,o,h used #f4bj cc4 6,answer63 #ig4ig4 #ca4 ,it is desired to dilute a solution of hydrochloric acid to exactly #j4je ,n4 ,the following data are given3 #dd4ig cc4 of the hydrochloric acid are equivalent to #dc4gf cc4 of the ,na,o,h solution4 ,the ,na,o,h is standardized against a pure potassium tetroxalate 7,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o7 weighing #j4bafb gram and requires #di4ad cc4 ,how many cc4 of water must be added to #ajjj cc4 of the aqueous hydrochloric acid8 6,answer63 #aa cc4 #cb4 ,how many cubic centimeters of #c ,n phosphoric acid must be added to #cjj cc4 of #j4d ,n phosphoric acid in order that the resulting solution may be #j4f ,n8 6,answer63 #be cc4 #cc4 ,to oxidize the iron in #a gram of ,fe,s,o#d7,,nh#d7b,,so#d4f;,h#b,o 7mol4 wgt4 #cib7 requires #c cc4 of a given solution of ,,hno#c4 ,what is the normality of the nitric acid when used as an acid8 #f;,fe,s,o#d #b;,,hno#c #b;,h#b,,so#d #c;,fe#b7,,so#d7c #b,,no #d;,h#b,o4 6,answer63 #j4bhce ,n4 #cd4 ,the same volume of carbon dioxide at the same temperature and the same pressure is liberated from a #a gram sample of dolomite1 by adding an excess of hydrochloric acid1 as can be liberated by the addition of #ce cc4 of #j4e ,n hydrochloric acid to an excess of any pure or impure carbonate4 ,calculate the percentage of ,,co#b in the dolomite4 6,answer63 #ch4e4 #ce4 ,how many cubic centimeters of sulphuric acid 7sp4 gr4 #a4hd1 containing #if ,h#b,,so#d by weight7 will be required to displace the chloride in the calcium chloride formed by the action of #ajj cc4 of #j4ajgb ,n hydrochloric acid on an excess of calcium carbonate1 and how many grams of ,ca,s,o#d will be formed8 6,answers63 #j4bih cc42 #j4gcjj gram4 #cf4 ,potassium hydroxide which has been exposed to the air is found on analysis to contain #g4fb water1 #b4ch ,k#b;,,co#c4 and #ij ,,koh4 ,what weight of residue will be obtained if one gram of this sample is added to #df cc4 of normal hydrochloric acid and the resulting solution1 after exact neutralization with #a4jgj ,n potassium hydroxide solution1 is evaporated to dryness8 6,answer63 #c4dg grams4 #cg4 ,a chemist received four different solutions1 with the statement that they contained either pure ,na,o,h2 pure ,na#b;,,co#c2 pure ,na,h,c,o#c1 or mixtures of these substances4 ,from the following data identify them3 ,sample ,i4 ,on adding phenolphthalein to a solution of the substance1 it gave no color to the solution4 ,sample ,,ii4 ,on titrating with standard acid1 it required #ae4bf cc4 for a change in color1 using phenolphthalein1 and #ag4ij cc4 additional1 using methyl orange as an indicator4 ,sample ,,iii4 ,the sample was titrated with hydrochloric acid until the pink of phenolphthalein disappeared1 and on the addition of methyl orange the solution was colored pink4 ,sample ,,iv4 ,on titrating with hydrochloric acid1 using phenolphthalein1 #ae4jj cc4 were required4 ,a new sample of the same weight required exactly #cj cc4 of the same acid for neutralization1 using methyl orange4 6,answers63 7a7 ,na,h,c,o#c2 7b7 ,na,h,c,o#c,na#b;,,co#c2 7c7,na,o,h2 7d7 ,na#b;,,co#c4 #ch4 ,in the analysis of a sample of ,,khc#d;,h#d,o#f the following data are obtained3 ,weight sample #j4dgcb gram4 ,na,o,h solution used #bd4ig cc4 #c4jj cc4 ,na,o,h #a cc4 of ,h#c,,po#d solution of which #a cc4 will precipitate #j4jabbg gram of magnesium as ,mg,n,h#d,,po#d4 ,calculate the percentage of ,,khc#d;,h#d,o#f4 6,answer63 #hh4fg4 #ci4 ,a one-gram sample of sodium hydroxide which has been exposed to the air for some time1 is dissolved in water and diluted to exactly #ejj cc4 ,one hundred cubic centimeters of the solution1 when titrated with #j4ajfb ,n hydrochloric acid1 using methyl orange as an indicator1 requires #ch4fj cc4 for complete neutralization4 ,barium chloride in excess is added to a second portion of #ajj cc4 of the solution1 which is diluted to exactly #bej cc41 allowed to stand and filtered4 ,two hundred cubic centimeters of this filtrate require #bi4fb cc4 of #j4ajfb ,n hydrochloric acid for neutralization1 using phenolphthalein as an indicator4 ,calculate percentage of ,na,o,h1 ,na#b;,,co#c1 and ,h#b,o4 6,answers63 #gh4fc ,na,o,h2 #d4de ,na#b;,,co#c2 #af4ib ,h#b,o4 #dj4 ,a sodium hydroxide solution 7made from solid ,na,o,h which has been exposed to the air7 was titrated against a standard acid using methyl orange as an indicator1 and was found to be exactly #j4a ,n4 ,this solution was used in the analysis of a material sold at #b cents per pound per cent of an acid constituent ,a1 and always mixed so that it was supposed to contain #ae of ,a1 on the basis of the analyst's report4 ,owing to the carelessness of the analyst's assistant1 the sodium hydroxide solution was used with phenolphthalein as an indicator in cold solution in making the analyses4 ,the concern manufacturing this material sells #fjj tons per year1 and when the mistake was discovered it was estimated that at the end of a year the error in the use of indicators would either cost them or their customers #fjjj4 ,who would lose and why8 ,assuming the impure ,na,o,h used originally in making the titrating solution consisted of ,na,o,h and ,na#b;,,co#c only1 what per cent of each was present8 6,answers63 ,customer lost2 #c4id ,na#b;,,co#c2 #if4jf ,na,o,h4 #da4 ,in the standardization of a ,k#b;,cr#b,o#g solution against iron wire1 #ii4he pure1 #db4db cc4 of the solution were added4 ,the weight of the wire used was #j4bb gram4 #c4bg cc4 of a ferrous sulphate solution having a normal value as a reducing agent of #j4ajaa were added to complete the titration4 ,calculate the normal value of the ,k#b;,cr#b,o#g4 6,answer63 #j4ajjf ,n4 #db4 ,what weight of iron ore containing #ef4b ,fe should be taken to standardize an approximately #j4a ,n oxidizing solution1 if not more than #dg cc4 are to be used8 6,answer63 #j4dffg gram4 #dc4 ,one tenth gram of iron wire1 #ii4gh pure1 is dissolved in hydrochloric acid and the iron oxidized completely with bromine water4 ,how many grams of stannous chloride are there in a liter of solution if it requires #i4dg cc4 to just reduce the iron in the above8 ,what is the normal value of the stannous chloride solution as a reducing agent8 6,answer63 #ag4ib grams2 #j4ahhh ,n4 #dd4 ,one gram of an oxide of iron is fused with potassium acid sulphate and the fusion dissolved in acid4 ,the iron is reduced with stannous chloride1 mercuric chloride is added1 and the iron titrated with a normal ,k#b;,cr#b,o#g solution4 #ab4id cc4 were used4 ,what is the formula of the oxide1 ,fe,o1 ,fe#b,o#c1 or ,fe#c,o#d8 6,answer63 ,fe#c,o#d4 #de4 ,if an element has #ih for its atomic weight1 and after reduction with stannous chloride could be oxidized by bichromate to a state corresponding to an ,,xo#d- anion1 compute the oxide1 or valence1 corresponding to the reduced state from the following data3 #j4cbff gram of the pure element1 after being dissolved1 was reduced with stannous chloride and oxidized by #dj cc4 of ,k#b;,cr#b,o#g1 of which one cc4 #j4aifj gram of ,fe,s,o#d7,,nh#d7b,,so#d4f;,h#b,o4 6,answer63 ,monovalent4 #df4 ,determine the percentage of iron in a sample of limonite from the following data3 ,sample #j4ejjj gram4 ,k,mn,o#d used #ej cc4 #a cc4 ,k,mn,o#d #j4jjecag gram ,fe4 ,fe,s,o#d used #f cc4 #a cc4 ,fe,s,o#d #j4jjibjj gram ,fe,o4 6,answer63 #dd4fj4 #dg4 ,if #a gram of a silicate yields #j4ejjj gram of ,fe#b,o#c and ,al#b,o#c and the iron present requires #be cc4 of #j4b ,n ,k,mn,o#d1 calculate the percentage of ,fe,o and ,al#b,o#c in the sample4 6,answer63 #ce4hi ,fe,o2 #aj4jc ,al#b,o#c4 #dh4 ,a sample of magnesia limestone has the following composition3 ,silica1 #c4jj2 ferric oxide and alumina1 #j4bj2 calcium oxide1 #cc4aj2 magnesium oxide1 #bj4gj2 carbon dioxide1 #dc4jj4 ,in manufacturing lime from the above the carbon dioxide is reduced to #c4jj4 ,how many cubic centimeters of normal ,k,mn,o#d will be required to determine the calcium oxide volumetrically in a #a gram sample of the lime8 6,answer63 #bj4jh cc4 #di4 ,if #ajj cc4 of potassium bichromate solution 7#aj gram ,k#b;,cr#b,o#g per liter71 #e cc4 of #f ,n sulphuric acid1 and #ge cc4 of ferrous sulphate solution 7#hj grams ,fe,s,o#d4g;,h#b,o per liter7 are mixed1 and the resulting solution titrated with #j4baba ,n ,k,mn,o#d1 how many cubic centimeters of the ,k,mn,o#d solution will be required to oxidize the iron8 6,answer63 #e4gj cc4 #ej4 ,if a #j4ejjj gram sample of limonite containing #ei4ej per cent ,fe#b,o#c requires #dj cc4 of ,k,mn,o#d to oxidize the iron1 what is the value of #a cc4 of the permanganate in terms of 7a7 ,fe1 7b7 ,h#b;,c#b,o#d4b;,h#b,o8 6,answers63 7a7 #j4jjeahi gram2 7b7 #j4jjehei gram4 #ea4 ,a sample of pyrolusite weighing #j4fjjj gram is treated with #j4ijjj gram of oxalic acid4 ,the excess oxalic acid requires #bc4ie cc4 of permanganate 7#a cc4 #j4jcjch gram ,fe,s,o#d4g;,h#b,o74 ,what is the percentage of ,mn,o#b1 in the sample8 6,answer63 #hd4dg4 #eb4 ,a solution contains #ej grams of ,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o per liter4 ,what is the normal value of the solution 7a7 as an acid1 and 7b7 as a reducing agent8 6,answers63 7a7 #j4eijc ,n2 7b7 #j4ghgb ,n4 #ec4 ,in the analysis of an iron ore containing #fj ,fe#b,o#c1 a sample weighing #j4ejjj gram is taken and the iron is reduced with sulphurous acid4 ,on account of failure to boil out all the excess ,,so#b1 #ch4fj cubic centimeters of #j4a ,n ,k,mn,o#d were required to titrate the solution4 ,what was the error1 percentage error1 and what weight of sulphur dioxide was in the solution8 6,answers63 7a7 #a4fj2 7b7 #b4fg2 7c7 #j4jjcbb gram4 #ed4 ,from the following data1 calculate the ratio of the nitric acid as an oxidizing agent to the tetroxalate solution as a reducing agent3 #a cc4 ,,hno#c #a4bdf cc4 ,na,o,h solution2 #a cc4 ,na,o,h #a4gdc cc4 ,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o solution2 ,normal value ,na,o,h #j4ab4 6,answer63 #d4hhe4 #ee4 ,given the following data3 #be cc4 of a hydrochloric acid1 when standardized gravimetrically as silver chloride1 yields a precipitate weighing #j4edfe gram4 #bd4ce cc4 of the hydrochloric acid are exactly equivalent to #cj4ag cc4 of ,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o solution4 ,how much water must be added to a liter of the oxalate solution to make it exactly #j4jbe ,n as a reducing agent8 6,answer63 #eefd cc4 #ef4 ,ten grams of a mixture of pure potassium tetroxalate 7,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o7 and pure oxalic acid 7,h#b;,c#b,o#d4b;,h#b,o7 are dissolved in water and diluted to exactly #ajjj cc4 ,the normal value of the oxalate solution when used as an acid is #j4acae4 ,calculate the ratio of tetroxalate to oxalate used in making up the solution and the normal value of the solution as a reducing agent4 6,answers63 #b3a2 #j4aegg ,n4 #eg4 ,a student standardized a solution of ,na,o,h and one of ,k,mn,o#d against pure ,,khc#b,o#d4;,h#b;,c#b,o#d4b;,h#b,o and found the former to be #j4jgejj ,n as an alkali and the latter exactly #j4a ,n as an oxidizing agent4 ,by coincidence1 exactly #dg4bf cc4 were used in each standardization4 ,find the ratio of the oxalate used in the ,na,o,h standardization to the oxalate used in the permanganate standardization4 6,answer63 #a3a4 #eh4 ,a sample of apatite weighing #j4fj gram is analyzed for its phosphoric anhydride content4 ,if the phosphate is precipitated as 7,,nh#d7c,,po#d4ab,mo,o#c1 and the precipitate 7after solution and reduction of the ,mo,o#c to ,mo#bd,o#cg71 requires #ajj cc4 of normal ,k,mn,o#d to oxidize it back to ,mo,o#c1 what is the percentage of ,p#b,o#e8 6,answer63 #cc4ha4 #ei4 ,in the analysis of a sample of steel weighing #a4hha grams the phosphorus was precipitated with ammonium molybdate and the yellow precipitate was dissolved1 reduced and titrated with ,k,mn,o#d4 ,if the sample contained #j4jbe per cent ,p and #f4ja cc4 of ,k,mn,o#d were used1 to what oxide was the molybdenum reduced8 #a cc4 ,k,mn,o#d #j4jjgahh gram ,na#b;,c#b,o#d4 6,answer63 ,mo#d,o#e4 #fj4 ,what is the value of #a cc4 of an iodine solution 7#a cc4 equivalent to #j4jcjj gram ,na#b,s#b,o#c7 in terms of ,as#b,o#c8 6,answer63 #j4jjiche gram4 #fa4 #dh cc4 of a solution of sodium thiosulphate are required to titrate the iodine liberated from an excess of potassium iodide solution by #j4cjjj gram of pure ,,kio#c4 7,,kio#c #e,,ki #c;,h#b,,so#d #c,k#b,,so#d #c;,i#b #c;,h#b,o47 ,what is the normal strength of the sodium thiosulphate and the value of #a cc4 of it in terms of iodine8 6,answers63 #j4agec ,n2 #j4jbbbd gram4 #fb4 ,one thousand cubic centimeters of #j4ajgi ,n sodium thiosulphate solution is allowed to stand4 ,one per cent by weight of the thiosulphate is decomposed by the carbonic acid present in the solution4 ,to what volume must the solution be diluted to make it exactly #j4a ,n as a reducing agent8 7,na#b,s#b,o#c #b;,h#b;,,co#c ,h#b,,so#c #b,na,h,c,o#c ,s47 6,answer63 #ajij cc4 #fc4 ,an analyzed sample of stibnite containing #gj4je ,sb is given for analysis4 ,a student titrates it with a solution of iodine of which #a cc4 is equivalent to #j4jjdiej gram of ,as#b,o#c4 ,due to an error on his part in standardization1 the student's analysis shows the sample to contain #gj4cb ,sb4 ,calculate the true normal value of the iodine solution1 and the percentage error in the analysis4 6,answers63 #j4ajjj ,n2 #j4ci4 #fd4 ,a sample of pyrolusite weighing #j4ejjj gram is treated with an excess of hydrochloric acid1 the liberated chlorine is passed into potassium iodide and the liberated iodine is titrated with sodium thiosulphate solution 7#di4ff grams of pure ,na#b,s#b,o#c4e;,h#b,o per liter74 ,if #ch4gb cc4 are required1 what volume of #j4be normal permanganate solution will be required in an indirect determination in which a similar sample is reduced with #j4ijab gram ,h#b;,c#b,o#d4b;,h#b,o and the excess oxalic acid titrated8 6,answer63 #bf4bb cc4 #fe4 ,in the determination of sulphur in steel by evolving the sulphur as hydrogen sulphide1 precipitating cadmium sulphide by passing the liberated hydrogen sulphide through ammoniacal cadmium chloride solution1 and decomposing the ,cd,s with acid in the presence of a measured amount of standard iodine1 the following data are obtained3 ,sample1 #e4jbg grams2 cc4 ,na#b,s#b,o#c sol4 #ab4fh2 cc4 ,iodine sol4 #ae4ei2 #a cc4 ,iodine sol4 #a4jhf cc4 ,na#b,s#b,o#c sol42 #a cc4 ,na#b,s#b,o#c #j4jjejdd gram ,cu4 ,calculate the percentage of sulphur4 7,h#b,s ,i#b #b;,,hi ,s47 6,answer63 #j4ajg4 #ff4 ,given the following data1 calculate the percentage of iron in a sample of crude ferric chloride weighing #a4jjj gram4 ,the iodine liberated by the reaction #b;,fe,cl#c #b;,,hi #b;,h,cl #b;,fe,cl#b ,i#b is reduced by the addition of #ej cc4 of sodium thiosulphate solution and the excess thiosulphate is titrated with standard iodine and requires #g4he cc4 #de cc4 ,i#b solution #de4ie cc4 ,na#b,s#b,o#c solution2 #de cc4 ,as#b,o#c solution #de4bg cc4 ,i#b solution4 #a cc4 arsenite solution #j4jjeafj gram ,as#b,o#c4 6,answer63 #bc4gg4 #fg4 ,sulphide sulphur was determined in a sample of reduced barium sulphate by the evolution method1 in which the sulphur was evolved as hydrogen sulphide and was passed into ,cd,cl#b solution1 the acidified precipitate being titrated with iodine and thiosulphate4 ,sample1 #e4jgf grams2 cc4 ,i#b #bj4hc2 cc4 ,na#b,s#b,o#c #ab4cg2 #dc4de cc4 ,na#b,s#b,o#c #dc4db cc4 ,i#b2 #h4jf cc4 ,k,mn,o#d #dd4ff cc4 ,na#b,s#b,o#c2 #bh4hg cc4 ,k,mn,o#d #j4bjjd gram ,na#b;,c#b,o#d4 ,calculate the percentage of sulphide sulphur in the sample4 6,answer63 #j4jej4 #fh4 ,what weight of pyrolusite containing #hi4ba ,mn,o#b will oxidize the same amount of oxalic acid as #cg4ab cc4 of a permanganate solution1 of which #a cc4 will liberate #j4jage gram of ,i#b from ,,ki8 6,answer63 #j4bdic gram4 #fi4 ,a sample of pyrolusite weighs #j4bdjj gram and is #ib4ej pure ,mn,o#b4 ,the iodine liberated from ,,ki by the manganese dioxide is sufficient to react with #df4bd cc4 of ,na#b,s#b,o#c sol4 ,what is the normal value of the thiosulphate8 6,answer633 #j4aaje ,n4 #gj4 ,in the volumetric analysis of silver coin 7#ij ,ag71 using a #j4ejjj gram sample1 what is the least normal value that a potassium thiocyanate solution may have and not require more than #ej cc4 of solution in the analysis8 6,answer63 #j4jhcci ,n4 #ga4 ,a mixture of pure lithium chloride and barium bromide weighing #j4f gram is treated with #de4ae cubic centimeters of #j4bjag ,n silver nitrate1 and the excess titrated with #be cc4 of #j4a ,n ,,kscn solution1 using ferric alum as an indicator4 ,calculate the percentage of bromine in the sample4 6,answer63 #dj4aa4 #gb4 ,a mixture of the chlorides of sodium and potassium from #j4ejjj gram of a feldspar weighs #j4aejj gram1 and after solution in water requires #bb4ga cc4 of #j4ajab ,n silver nitrate for the precipitation of the chloride ions4 ,what are the percentages of ,na#b,o and ,k#b,o in the feldspar8 6,answer63 #h4bd ,na#b,o2 #i4ad ,k#b,o4 ,,gravimetric ,,analysis #gc4 ,calculate 7a7 the grams of silver in one gram of silver chloride2 7b7 the grams of carbon dioxide liberated by the addition of an excess of acid to one gram of calcium carbonate2 7c7 the grams of ,mg,cl#b necessary to precipitate #a gram of ,mg,n,h#d,,po#d4 6,answers63 7a7 #j4gebf2 7b7 #j4dcig2 7c7 #j4fidj4 #gd4 ,calculate the chemical factor for 7a7 ,sn in ,sn,o#b2 7b7 ,mg,o in ,mg#b,p#b,o#g2 7c7 ,p#b,o#e in ,mg#b,p#b,o#g2 7d7 ,fe in ,fe#b,o#c2 7e7 ,,so#d in ,ba,s,o#d4 6,answers63 7a7 #j4ghgi2 7b7 #j4cfbj2 7c7 #j4fcgh2 7d7 #j4fiij2 7e7 #j4daae4 #ge4 ,calculate the log factor for 7a7 ,pb in ,pb,cr,o#d2 7b7 ,cr#b,o#c in ,pb,cr,o#d2 7c7 ,pb in ,pb,o#b and 7d7 ,ca,o in ,ca,c#b,o#d4 6,answers63 7a7 #i4hjfi-aj1 7b7 #i4cgac-aj2 7c7 #i4icgf-aj2 7d7 #i4fdae-aj4 #gf4 ,how many grams of ,mn#c,o#d can be obtained from #a gram of ,mn,o#b8 6,answer63 #j4hggd gram4 #gg4 ,if a sample of silver coin weighing #j4bejj gram gives a precipitate of ,ag,cl weighing #j4biia gram1 what weight of ,ag,i could have been obtained from the same weight of sample1 and what is the percentage of silver in the coin8 6,answers63 #j4dhih gr42 #ij4je4 #gh4 ,how many cubic centimeters of hydrochloric acid 7sp4 gr4 #a4ac containing #be4ge ,h,cl by weight7 are required to exactly neutralize #be cc4 of ammonium hydroxide 7sp4 gr4 4#ij containing #bh4cc ,,nh#c by weight78 6,answer63 #dg4jc cc4 #gi4 ,how many cubic centimeters of ammonium hydroxide solution 7sp4 gr4 #j4if containing #i4ia ,,nh#c by weight7 are required to precipitate the aluminium as aluminium hydroxide from a two-gram sample of alum 7,k,al7,,so#d7b4ab;,h#b,o78 ,what will be the weight of the ignited precipitate8 6,answers63 #b4bf cc42 #j4baed gram4 #hj4 ,what volume of nitric acid 7sp4 gr4 #a4je containing #i4j ,,hno#c by weight7 is required to oxidize the iron in one gram of ,fe,s,o#d4g;,h#b,o in the presence of sulphuric acid8 #f;,fe,s,o#d #b;,,hno#c #c;,h#b,,so#d #c;,fe#b7,,so#d7c #b,,no #d;,h#b,o4 6,answer63 #j4hj cc4 #ha4 ,if #j4gecj gram of ferric nitrate 7,fe7,,no#c7c4i;,h#b,o7 is dissolved in water and #a4cg cc4 of ,h,cl 7sp4 gr4 #a4aa containing #ba4ib ,h,cl by weight7 is added1 how many cubic centimeters of ammonia 7sp4 gr4 #j4if containing #i4ia ,,nh#c by weight7 are required to neutralize the acid and precipitate the iron as ferric hydroxide8 6,answer63 #b4fc cc4 #hb4 ,to a suspension of #j4cajj gram of ,al7,,oh7#c in water are added #ac4jj cc4 of aqueous ammonia 7sp4 gr4 #j4ij containing #bh4d ,,nh#c by weight74 ,how many cubic centimeters of sulphuric acid 7sp4 gr4 #a4ah containing #bd4g ,h#b,,so#d by weight7 must be added to the mixture in order to bring the aluminium into solution8 6,answer63 #cd4h cc4 #hc4 ,how many cubic centimeters of sulphurous acid 7sp4 gr4 #a4jd containing #ge grams ,,so#b per liter7 are required to reduce the iron in #a gram of ferric alum 7,k,fe7,,so#d7b4ab;,h#b,o78 ,fe#b7,,so#d7c ,,so#b #b;,h#b,o #b;,fe,s,o#d #b;,h#b,,so#d4 6,answer63 #j4he cc4 #hd4 ,how many cubic centimeters of a solution of potassium bichromate containing #bf4cj grams of ,k#b;,cr#b,o#g per liter must be taken in order to yield #j4fjcc gram of ,cr#b,o#c after reduction and precipitation of the chromium8 ,k#b;,cr#b,o#g #c,,so#b ,h#b,,so#d ,k#b,,so#d ,cr#b7,,so#d7c ,h#b,o4 6,answer63 #dd4ci cc4 #he4 ,how many cubic centimeters of ammonium hydroxide 7sp4 gr4 #j4idf containing #ac4hh ,,nh#c by weight7 are required to precipitate the iron as ,fe7,,oh7#c from a sample of pure ,fe,s,o#d47,,nh#d7b,,so#d4f;,h#b,o1 which requires #j4cd cc4 of nitric acid 7sp4 gr4 #a4cej containing #ee4gi ,,hno#c by weight7 for oxidation of the iron8 7,see problem ,no4 #hj for reaction47 6,answer63 #d4gd cc4 #hf4 ,in the analysis of an iron ore by solution1 oxidation and precipitation of the iron as ,fe7,,oh7#c1 what weight of sample must be taken for analysis so that each one hundredth of a gram of the ignited precipitate of ,fe#b,o#c shall represent one tenth of one per cent of iron8 6,answer63 #f4ii grams4 #hg4 ,what weight in grams of impure ferrous ammonium sulphate should be taken for analysis so that the number of centigrams of ,ba,s,o#d obtained will represent five times the percentage of sulphur in the sample8 6,answer63 #j4fhgj gram4 #hh4 ,what weight of magnetite must be taken for analysis in order that1 after precipitating and igniting all the iron to ,fe#b,o#c1 the percentage of ,fe#b,o#d in the sample may be found by multiplying the weight in grams of the ignited precipitate by #ajj8 6,answer63 #j4iffe gram4 #hi4 ,after oxidizing the arsenic in #j4ejjj gram of pure ,as#b,s#c to arsenic acid1 it is precipitated with magnesia mixture 7,mg,cl#b #b,,nh#d;,cl74 ,if exactly #ab4f cc4 of the mixture are required1 how many grams of ,mg,cl#b per liter does the solution contain8 ,h#c;,as,o#d ,mg,cl#b #c,,nh#d,,oh ,mg,n,h#d;,as,o#d #b,,nh#d;,cl #c;,h#b,o4 6,answer63 #cj4ga grams4 #ij4 ,a sample is prepared for student analysis by mixing pure apatite 7,ca#c7,,po#d7b4;,ca,cl#b7 with an inert material4 ,if #a gram of the sample gives #j4djac gram of ,mg#b,p#b,o#g1 how many cubic centimeters of ammonium oxalate solution 7containing #dj grams of 7,,nh#d7b;,c#b,o#d4;,h#b,o per liter7 would be required to precipitate the calcium from the same weight of sample8 6,answer63 #be4fj cc4 #ia4 ,if #j4fgdb gram of a mixture of pure magnesium carbonate and pure calcium carbonate1 when treated with an excess of hydrochloric acid1 yields #j4caag gram of carbon dioxide1 calculate the percentage of magnesium oxide and of calcium oxide in the sample4 6,answers63 #ac4bb ,mg,o2 #dj4ed ,ca,o4 #ib4 ,the calcium in a sample of dolomite weighing #j4ichj gram is precipitated as calcium oxalate and ignited to calcium oxide4 ,what volume of gas1 measured over water at #bj;,c4 and #gfe mm4 pressure1 is given off during ignition1 if the resulting oxide weighs #j4bfjf gram8 7,g4,m4,v4 #bb4d liters2 ,v4,p4 water at #bj;,c4 #ag4d mm47 6,answer63 #bbg cc4 #ic4 ,a limestone is found to contain #ic4je ,ca,c,o#c1 and #e4af ,mg,c,o#c4 ,calculate the weight of ,ca,o obtainable from #c tons of the limestone1 assuming complete conversion to oxide4 ,what weight of ,mg#b,p#b,o#g could be obtained from a #c-;gram sample of the limestone8 6,answers63 #a4efe tons2 #j4bjdd gram4 #id4 ,a sample of dolomite is analyzed for calcium by precipitating as the oxalate and igniting the precipitate4 ,the ignited product is assumed to be ,ca,o and the analyst reports #bi4ej ,ca in the sample4 ,owing to insufficient ignition1 the product actually contained #h of its weight of ,ca,c,o#c4 ,what is the correct percentage of calcium in the sample1 and what is the percentage error8 6,answers63 #bh4df2 #c4fe error4 #ie4 ,what weight of impure calcite 7,ca,c,o#c7 should be taken for analysis so that the volume in cubic centimeters of ,,co#b obtained by treating with acid1 measured dry at #ah;,c4 and #gfc mm41 shall equal the percentage of ,ca,o in the sample8 6,answer63 #j4bcei gram4 #if4 ,how many cubic centimeters of ,,hno#c 7sp4 gr4 #a4ac containing #ba4j ,,hno#c by weight7 are required to dissolve #e grams of brass1 containing #j4fa ,pb1 #bd4ci ,zn1 and #ge ,cu1 assuming reduction of the nitric acid to ,,no by each constituent8 ,what fraction of this volume of acid is used for oxidation8 6,answers63 #ee4jf cc42 #be4 #ig4 ,what weight of metallic copper will be deposited from a cupric salt solution by a current of #a4e amperes during a period of #de minutes1 assuming #ajj current efficiency8 7#a ,faraday #if1ejj coulombs47 6,answer63 #a4cce grams4 #ih4 ,in the electrolysis of a #j4hjjj gram sample of brass1 there is obtained #j4jjcj gram of ,pb,o#b1 and a deposit of metallic copper exactly equal in weight to the ignited precipitate of ,zn#b,p#b,o#g subsequently obtained from the solution4 ,what is the percentage composition of the brass8 6,answers63 #fi4ge ,cu2 #bi4ib ,zn2 #j4cc ,pb4 #ii4 ,a sample of brass 7#fh4ij ,cu2 #a4aj ,pb and #cj4jj ,zn7 weighing #j4idjj gram is dissolved in nitric acid4 ,the lead is determined by weighing as ,pb,s,o#d1 the copper by electrolysis and the zinc by precipitation with 7,,nh#d7b;,,hpo#d in a neutral solution4 7a7 ,calculate the cubic centimeters of nitric acid 7sp4 gr4 #a4db containing #fi4ij ,,hno#c by weight7 required to just dissolve the brass1 assuming reduction to ,,no4 6,answer63 #b4dh cc4 7b7 ,calculate the cubic centimeters of sulphuric acid 7sp4 gr4 #a4hd containing #id ,h#b,,so#d by weight7 to displace the nitric acid4 6,answer63 #j4hc cc4 7c7 ,calculate the weight of ,pb,s,o#d4 6,answer63 #j4jaeb gram4 7d7 ,the clean electrode weighs #aj4ifdj grams4 ,calculate the weight after the copper has been deposited4 6,answer63 #aa4faaf grams4 7e7 ,calculate the grams of 7,,nh#d7b;,,hpo#d required to precipitate the zinc as ,zn,n,h#d,,po#d4 6,answer63 #j4egje gram4 7f7 ,calculate the weight of ignited ,zn#b,p#b,o#g4 6,answer63 #j4fegc gram4 #ajj4 ,if in the analysis of a brass containing #bh4jj zinc an error is made in weighing a #b4e gram portion by which #j4jja gram too much is weighed out1 what percentage error in the zinc determination would result8 ,what volume of a solution of sodium hydrogen phosphate1 containing #ij grams of ,na#b;,,hpo#d4ab;,h#b,o per liter1 would be required to precipitate the zinc as ,zn,n,h#d,,po#d and what weight of precipitate would be obtained8 6,answers63 7a7 #j4jd error2 7b7 #ci4ig cc42 7c7 #a4iji grams4 #aja4 ,a sample of magnesium carbonate1 contaminated with ,si,o#b as its only impurity1 weighs #j4ejjj gram and loses #j4ajjj gram on ignition4 ,what volume of disodium phosphate solution 7containing #ij grams ,na#b;,,hpo#d4ab;,h#b,o per liter7 will be required to precipitate the magnesium as magnesium ammonium phosphate8 6,answer63 #i4jg cc4 #ajb4 #b4fb cubic centimeters of nitric acid 7sp4 gr4 #a4db containing #fi4hj ,,hno#b by weight7 are required to just dissolve a sample of brass containing #fi4bg ,cu2 #j4je ,pb2 #j4jg ,fe2 and #cj4fa ,zn4 ,assuming the acid used as oxidizing agent was reduced to ,,no in every case1 calculate the weight of the brass and the cubic centimeters of acid used as acid4 6,answer63 #j4iib gram2 #a4ig cc4 #ajc4 ,one gram of a mixture of silver chloride and silver bromide is found to contain #j4ffce gram of silver4 ,what is the percentage of bromine8 6,answer63 #ba4cj4 #ajd4 ,a precipitate of silver chloride and silver bromide weighs #j4hacb gram4 ,on heating in a current of chlorine1 the silver bromide is converted to silver chloride1 and the mixture loses #j4adej gram in weight4 ,calculate the percentage of chlorine in the original precipitate4 6,answer63 #f4ac4 #aje4 ,a sample of feldspar weighing #a4jjj gram is fused and the silica determined4 ,the weight of silica is #j4fdfj gram4 ,this is fused with #d grams of sodium carbonate4 ,how many grams of the carbonate actually combined with the silica in fusion1 and what was the loss in weight due to carbon dioxide during the fusion8 6,answers63 #a4ace grams2 #j4dgae gram4 #ajf4 ,a mixture of barium oxide and calcium oxide weighing #b4babj grams is transformed into mixed sulphates1 weighing #e4jbc grams4 ,calculate the grams of calcium oxide and barium oxide in the mixture4 6,answers63 #a4hbd grams ,ca,o2 #j4chgg gram ,ba,o4 ,,appendix ,,electrolytic ,,dissociation ,,theory ,the following brief statements concerning the ionic theory and a few of its applications are intended for reference in connection with the explanations which are given in the ,notes accompanying the various procedures4 ,the reader who desires a more extended discussion of the fundamental theory and its uses is referred to such books as ,talbot and ,blanchard's 6,electrolytic ,dissociation ,theory6 7,macmillan ,company71 or ,alexander ,smith's 6,introduction to ,general ,inorganic ,chemistry6 7,century ,company74 ,the 6electrolytic dissociation theory61 as propounded by ,arrhenius in #ahhg1 assumes that acids1 bases1 and salts 7that is1 electrolytes7 in aqueous solution are dissociated to a greater or less extent into 6ions64 ,these ions are assumed to be electrically charged atoms or groups of atoms1 as1 for example1 ,h and ,br- from hydrobromic acid1 ,na and ,,oh- from sodium hydroxide1 #b,,nh#d and ,,so#d-- from ammonium sulphate4 ,the unit charge is that which is dissociated with a hydrogen ion4 ,those upon other ions vary in sign and number according to the chemical character and valence of the atoms or radicals of which the ions are composed4 ,in any solution the aggregate of the positive charges upon the positive ions 76cations67 must always balance the aggregate negative charges upon the negative ions 76anions674 ,it is assumed that the ,na ion1 for example1 differs from the sodium atom in behavior because of the very considerable electrical charge which it carries and which1 as just stated1 must1 in an electrically neutral solution1 be balanced by a corresponding negative charge on some other ion4 ,when an electric current is passed through a solution of an electrolyte the ions move with and convey the current1 and when the cations come into contact with the negatively charged cathode they lose their charges1 and the resulting electrically neutral atoms 7or radicals7 are liberated as such1 or else enter at once into chemical reaction with the components of the solution4 ,two ions of identically the same composition but with different electrical charges may exhibit widely different properties4 ,for example1 the ion ,mn,o#d- from permanganates yields a purple-red solution and differs in its chemical behavior from the ion ,mn,o#d-- from manganates1 the solutions of which are green4 ,the chemical changes upon which the procedures of analytical chemistry depend are almost exclusively those in which the reacting substances are electrolytes1 and analytical chemistry is1 therefore1 essentially the chemistry of the ions4 ,the percentage dissociation of the same electrolyte tends to increase with increasing dilution of its solution1 although not in direct proportion4 ,the percentage dissociation of different electrolytes in solutions of equivalent concentrations 7such1 for example1 as normal solutions7 varies widely1 as is indicated in the following tables1 in which approximate figures are given for tenth-normal solutions at a temperature of about #ah;,c4 ,,acids ,,substance ,,percentage ,,dissociation ,,in #j4a ,,equivalent ,,solution ,h,cl1 ,h,br1 ,,hi1 ,,hno#c #ij ,h,cl,o#c1 ,h,cl,o#d1 ,h,mn,o#d #ij ,h#b,,so#d -- ,h ,,hso#d- #ij ,h#b;,c#b,o#d -- ,h ,,hc#b,o#d- #ej ,h#b,,so#c -- ,h ,,hso#c- #bj ,h#c,,po#d -- ,h ,h#b,,po#d- #bg ,h#b,,po#d- -- ,h ,,hpo#d-- #j4b ,h#c;,as,o#d -- ,h ,h#b;,as,o#d- #bj ,,hf #i ,,hc#b;,h#c,o#b #a4d ,h#b;,,co#c -- ,h ,,hco#c- #j4ab ,h#b,s -- ,h ,,hs- #j4je ,,hcn #j4ja ,,bases ,,substance ,,percentage ,,dissociation ,,in #j4a ,,equivalent ,,solution ,,koh1 ,na,o,h #hf ,ba7,,oh7#b #ge ,,nh#d,,oh #a4d ,,salts ,,type ,,of ,,salt ,,percentage ,,dissociation ,,in #j4a ,,equivalent ,,solution ,r,r- #hf ,r7,r-7#b #gb 7,r7#b,r-- #gb ,r,r-- #de ,the percentage dissociation is determined by studying the electrical conductivity of the solutions and by other physico-chemical methods1 and the following general statements summarize the results3 6,salts61 as a class1 are largely dissociated in aqueous solution4 6,acids6 yield ,h ions in water solution1 and the comparative 6strength61 that is1 the activity1 of acids is proportional to the concentration of the ,h ions and is measured by the percentage dissociation in solutions of equivalent concentration4 ,the common mineral acids are largely dissociated and therefore give a relatively high concentration of ,h ions1 and are commonly known as strong acids4 ,the organic acids1 on the other hand1 belong generally to the group of weak acids4 6,bases6 yield ,,oh- ions in water solution1 and the comparative strength of the bases is measured by their relative dissociation in solutions of equivalent concentration4 ,ammonium hydroxide is a weak base1 as shown in the table above1 while the hydroxides of sodium and potassium exhibit strongly basic properties4 ,ionic reactions are all1 to a greater or less degree1 6reversible reactions64 ,a typical example of an easily reversible reaction is that representing the changes in ionization which an electrolyte such as acetic acid undergoes on dilution or concentration of its solutions1 6i4e461 ,,hc#b;,h#c,o#b -- ,h ,c#b;,h#c,o#b-4 ,as was stated above1 the ionization increases with dilution1 the reaction then proceeding from left to right1 while concentration of the solution occasions a partial reassociation of the ions1 and the reaction proceeds from right to left4 ,to understand the principle underlying these changes it is necessary to consider first the conditions which prevail when a solution of acetic acid1 which has been stirred until it is of uniform concentration throughout1 has come to a constant temperature4 ,a careful study of such solutions has shown that there is a definite state of equilibrium between the constituents of the solution2 that is1 there is a definite relation between the undissociated acetic acid and its ions1 which is characteristic for the prevailing conditions4 ,it is not1 however1 assumed that this is a condition of static equilibrium1 but rather that there is continual dissociation and association1 as represented by the opposing reactions1 the apparent condition of rest resulting from the fact that the amount of change in one direction during a given time is exactly equal to that in the opposite direction4 ,a quantitative study of the amount of undissociated acid1 and of ,h ions and ,c#b;,h#c,o#b- ions actually to be found in a large number of solutions of acetic acid of varying dilution 7assuming them to be in a condition of equilibrium at a common temperature71 has shown that there is always a definite relation between these three quantities which may be expressed thus3 76,conc'n ,h x ,conc'n ,c#b;,h#c,o#b-7_/;,conc'n ,,hc#b;,h#c,o#b ,constant64 ,in other words1 there is always a definite and constant ratio between the product of the concentrations of the ions and the concentration of the undissociated acid when conditions of equilibrium prevail4 ,it has been found1 further1 that a similar statement may be made regarding all reversible reactions1 which may be expressed in general terms thus3 ,the rate of chemical change is proportional to the product of the concentrations of the substances taking part in the reaction2 or1 if conditions of equilibrium are considered in which1 as stated1 the rate of change in opposite directions is assumed to be equal1 then the product of the concentrations of the substances entering into the reaction stands in a constant ratio to the product of the concentrations of the resulting substances1 as given in the expression above for the solutions of acetic acid4 ,this principle is called the 6,law of ,mass ,action64 ,it should be borne in mind that the expression above for acetic acid applies to a wide range of dilutions1 provided the temperature remains constant4 ,if the temperature changes the value of the constant changes somewhat1 but is again uniform for different dilutions at that temperature4 ,the following data are given for temperatures of about #ah;,c4#a ,,molal ,,fraction ,,molal ,,concentra- ,,molal ,,concentra- ,,value ,,of ,,concentration ,,ionized ,,tion ,,of ,h ,,and ,,tion ,,of ,,undis- ,,constant ,,constant ,,acetate- ,,ions ,,sociated ,,acid #a4j 4#jjd 4#jjd 4#iif 4#jjjjafa #j4a 4#jac 4#jjac 4#jihg 4#jjjjaga #j4ja 4#jdjg 4#jjjdjg 4#jjieic 4#jjjjagb ,footnote #a3 ,alexander ,smith1 6,general ,inorganic ,chemistry61 p4 #egi4 ,the molal concentrations given in the table refer to fractions of a gram-molecule per liter of the undissociated acid1 and to fractions of the corresponding quantities of ,h and ,c#b;,h#c,o#b- ions per liter which would result from the complete dissociation of a gram-molecule of acetic acid4 ,the values calculated for the constant are subject to some variation on account of experimental errors in determining the percentage ionized in each case1 but the approximate agreement between the values found for molal and centimolal 7one hundredfold dilution7 is significant4 ,the figures given also illustrate the general principle1 that the 6relative6 ionization of an electrolyte increases with the dilution of its solution4 ,if we consider what happens during the 7usually7 brief period of dilution of the solution from molal to #j4a molal1 for example1 it will be seen that on the addition of water the conditions of concentration which led to equality in the rate of change1 and hence to equilibrium in the molal solution1 cease to exist2 and since the dissociating tendency increases with dilution1 as just stated1 it is true at the first instant after the addition of water that the concentration of the undissociated acid is too great to be permanent under the new conditions of dilution1 and the reaction1 ,,hc#b;,h#c,o#b -- ,h ,c#b;,h#c,o#b-1 will proceed from left to right with great rapidity until the respective concentrations adjust themselves to the new conditions4 ,that which is true of this reaction is also true of all reversible reactions1 namely1 that any change of conditions which occasions an increase or a decrease in concentration of one or more of the components causes the reaction to proceed in one direction or the other until a new state of equilibrium is established4 ,this principle is constantly applied throughout the discussion of the applications of the ionic theory in analytical chemistry1 and it should be clearly understood that whenever an existing state of equilibrium is disturbed as a result of changes of dilution or temperature1 or as a consequence of chemical changes which bring into action any of the constituents of the solution1 thus altering their concentrations1 there is always a tendency to re-establish this equilibrium in accordance with the law4 ,thus1 if a base is added to the solution of acetic acid the ,h ions then unite with the ,,oh- ions from the base to form undissociated water4 ,the concentration of the ,h ions is thus diminished1 and more of the acid dissociates in an attempt to restore equilbrium1 until finally practically all the acid is dissociated and neutralized4 ,similar conditions prevail when1 for example1 silver ions react with chloride ions1 or barium ions react with sulphate ions4 ,in the former case the dissociation reaction of the silver nitrate is ,ag,n,o#c -- ,ag ,,no#c-1 and as soon as the ,ag ions unite with the ,cl- ions the concentration of the former is diminished1 more of the ,ag,n,o#c dissociates1 and this process goes on until the ,ag ions are practically all removed from the solution1 if the ,cl- ions are present in sufficient quantity4 ,for the sake of accuracy it should be stated that the mass law cannot be rigidly applied to solutions of those electrolytes which are largely dissociated4 ,while the explanation of the deviation from quantitative exactness in these cases is not known1 the law is still of marked service in developing analytical methods along more logical lines than was formerly practicable4 ,it has not seemed wise to qualify each statement made in the ,notes to indicate this lack of quantitative exactness4 ,the student should recognize its existence1 however1 and will realize its significance better as his knowledge of physical chemistry increases4 ,if we apply the mass law to the case of a substance of small solubility1 such as the compounds usually precipitated in quantitative analysis1 we derive what is known as the 6solubility product61 as follows3 ,taking silver chloride as an example1 and remembering that it is not absolutely insoluble in water1 the equilibrium expression for its solution is3 76,conc'n ,ag x ,conc'n ,cl-7_/,conc'n ,ag,cl ,constant64 ,but such a solution of silver chloride which is in contact with the solid precipitate must be saturated for the existing temperature1 and the quantity of undissociated ,ag,cl in the solution is definite and constant for that temperature4 ,since it is a constant1 it may be eliminated1 and the expression becomes 6,conc'n ,ag x ,conc'n ,cl- ,constant61 and this is known as the solubility product4 ,no precipitation of a specific substance will occur until the product of the concentrations of its ions in a solution exceeds the solubility product for that substance2 whenever that product is exceeded precipitation must follow4 ,it will readily be seen that if a substance which yields an ion in common with the precipitated compound is added to such a solution as has just been described1 the concentration of that ion is increased1 and as a result the concentration of the other ion must proportionately decrease1 which can only occur through the formation of some of the undissociated compound which must separate from the already saturated solution4 ,this explains why the addition of an excess of the precipitant is often advantageous in quantitative procedures4 ,such a case is discussed at length in ,note #b on page #aac4 ,similarly1 the ionization of a specific substance in solution tends to diminish on the addition of another substance with a common ion1 as1 for instance1 the addition of hydrochloric acid to a solution of hydrogen sulphide4 ,hydrogen sulphide is a weak acid1 and the concentration of the hydrogen ions in its aqueous solutions is very small4 ,the equilibrium in such a solution may be represented as3 767,conc'n ,h7#b x ,conc'n ,s--7_/,conc'n ,h#b,s ,constant61 and a marked increase in the concentration of the ,h ions1 such as would result from the addition of even a small amount of the highly ionized hydrochloric acid1 displaces the point of equilibrium and some of the ,s-- ions unite with ,h ions to form undissociated ,h#b,s4 ,this is of much importance in studying the reactions in which hydrogen sulphide is employed1 as in qualitative analysis4 ,by a parallel course of reasoning it will be seen that the addition of a salt of a weak acid or base to solutions of that acid or base make it1 in effect1 still weaker because they decrease its percentage ionization4 ,to understand the changes which occur when solids are dissolved where chemical action is involved1 it should be remembered that no substance is completely insoluble in water1 and that those products of a chemical change which are least dissociated will first form4 ,consider1 for example1 the action of hydrochloric acid upon magnesium hydroxide4 ,the minute quantity of dissolved hydroxide dissociates thus3 ,mg7,,oh7#b -- ,mg #b,,oh-4 ,when the acid is introduced1 the ,h ions of the acid unite with the ,,oh- ions to form undissociated water4 ,the concentration of the ,,oh- ions is thus diminished1 more ,mg7,,oh7#b dissociates1 the solution is no longer saturated with the undissociated compound1 and more of the solid dissolves4 ,this process repeats itself with great rapidity until1 if sufficient acid is present1 the solid passes completely into solution4 ,exactly the same sort of process takes place if calcium oxalate1 for example1 is dissolved in hydrochloric acid4 ,the ,c#b,o#d-- ions unite with the ,h ions to form undissociated oxalic acid1 the acid being less dissociated than normally in the presence of the ,h ions from the hydrochloric acid 7see statements regarding hydrogen sulphide above74 ,as the undissociated oxalic acid forms1 the concentration of the ,c#b,o#d-- ions lessens and more ,ca,c#b,o#d dissolves1 as described for the ,mg7,,oh7#b above4 ,numerous instances of the applications of these principles are given in the ,notes4 ,water itself is slightly dissociated1 and although the resulting ,h and ,,oh- ions are present only in minute concentrations 7#a mol4 of dissociated water in #ajg liters71 yet under some conditions they may give rise to important consequences4 ,the term 6hydrolysis6 is applied to the changes which result from the reaction of these ions4 ,any salt which is derived from a weak base or a weak acid 7or both7 is subject to hydrolytic action4 ,potassium cyanide1 for example1 when dissolved in water gives an alkaline solution because some of the ,h ions from the water unite with ,,cn- ions to form 7,,hcn71 which is a very weak acid1 and is but very slightly dissociated4 ,potassium hydroxide1 which might form from the ,,oh- ions1 is so largely dissociated that the ,,oh- ions remain as such in the solution4 ,the union of the ,h ions with the ,,cn- ions to form the undissociated ,,hcn diminishes the concentration of the ,h ions1 and more water dissociates 7,h#b,o -- ,h ,,oh-7 to restore the equilibrium4 ,it is clear1 however1 that there must be a gradual accumulation of ,,oh- ions in the solution as a result of these changes1 causing the solution to exhibit an alkaline reaction1 and also that ultimately the further dissociation of the water will be checked by the presence of these ions1 just as the dissociation of the ,h#b,s was lessened by the addition of ,h,cl4 ,an exactly opposite result follows the solution of such a salt as ,al#b7,,so#d7c in water4 ,in this case the acid is strong and the base weak1 and the ,,oh- ions form the little dissociated ,al7,,oh7#c1 while the ,h ions remain as such in the solution1 sulphuric acid being extensively dissociated4 ,the solution exhibits an acid reaction4 ,such hydrolytic processes as the above are of great importance in analytical chemistry1 especially in the understanding of the action of indicators in volumetric analysis4 7,see page #cb47 ,the impelling force which causes an element to pass from the atomic to the ionic condition is termed 6electrolytic solution pressure61 or ionization tension4 ,this force may be measured in terms of electrical potential1 and the table below shows the relative values for a number of elements4 ,in general1 an element with a greater solution pressure tends to cause the deposition of an element of less solution pressure when placed in a solution of its salt1 as1 for instance1 when a strip of zinc or iron is placed in a solution of a copper salt1 with the resulting precipitation of metallic copper4 ,hydrogen is included in the table1 and its position should be noted with reference to the other common elements4 ,for a more extended discussion of this topic the student should refer to other treatises4 ,,potential ,,series ,,of ,,the ,,metals ,,potential ,,potential ,,in ,,volts ,,in ,,volts ,sodium ,na #b4dd ,lead ,pb -#j4ac ,calcium ,ca ,hydrogen ,h -#j4bh ,magnesium ,mg ,bismuth ,bi ,aluminum ,a#a #a4jj ,antimony -#j4ge ,manganese ,mn ,arsenic ,zinc ,zn #j4di ,copper ,cu -#j4fa ,cadmium ,cd #j4ad ,mercury ,hg -#a4jc ,iron ,fe #j4jfc ,silver ,ag -#a4je ,cobalt ,co -#j4jde ,platinum ,nickel ,ni -#j4jdi ,gold ,tin ,sn -#j4jhe787 ,,the ,,folding ,,of ,a ,,filter ,,paper ,if a filter paper is folded along its diameter1 and again folded along the radius at right angles to the original fold1 a cone is formed on opening1 the angle of which is #fj4 ,funnels for analytical use are supposed to have the same angle1 but are rarely accurate4 ,it is possible1 however1 with care1 to fit a filter thus folded into a funnel in such a way as to prevent air from passing down between the paper and the funnel to break the column of liquid in the stem1 which aids greatly1 by its gentle suction1 in promoting the rate of filtration4 ,such a filter has1 however1 the disadvantage that there are three thicknesses of paper back of half of its filtering surface1 as a consequence of which one half of a precipitate washes or drains more slowly4 ,much time may be saved in the aggregate by learning to fold a filter in such a way as to improve its effective filtering surface4 ,the directions which follow1 though apparently complicated on first reading1 are easily applied and easily remembered4 ,use a #f-;inch filter for practice4 ,place four dots on the filter1 two each on diameters which are at right angles to each other4 ,then proceed as follows3 7#a7 ,fold the filter evenly across one of the diameters1 creasing it carefully2 7#b7 open the paper1 turn it over1 rotate it #ij to the right1 bring the edges together and crease along the other diameter2 7#c7 open1 and rotate #de to the right1 bring edges together1 and crease evenly2 7#d7 open1 and rotate #ij to the right1 and crease evenly2 7#e7 open1 turn the filter over1 rotate #bb-7a_/b7 to the right1 and crease evenly2 7#f7 open1 rotate #de to the right and crease evenly2 7#g7 open1 rotate #de to the right and crease evenly2 7#h7 open1 rotate #de to the right and crease evenly2 7#i7 open the filter1 and1 starting with one of the dots between thumb and forefinger of the right hand1 fold the second crease to the left over on it1 and do the same with each of the other dots4 ,place it1 thus folded1 in the funnel1 moisten it1 and fit to the side of the funnel4 ,the filter will then have four short segments where there are three thicknesses and four where there is one thickness1 but the latter are evenly distributed around its circumference1 thus greatly aiding the passage of liquids through the paper and hastening both filtration and washing of the whole contents of the filter4 6,,sample ,,pages ,,for ,,laboratory ,,records6 6,page ,a6 ,date ,,calibration ,,of ,,burette ,no4 ,,burette ,,difference ,,observed ,,difference ,,calculated ,,readings ,,weights ,,correction #j4jb #af4bg #aj4ab #aj4aj #bf4ce #aj4jh -4#jb #bj4ji #i4ig #cf4bf #i4ia -4#jf #cj4af #aj4jg #df4cd #aj4jh 4#ja #dj4ai #aj4jc #ef4ca #i4ig -4#jf #ej4jj #i4ha #ff4ag #i4hf 4#je ,these data to be obtained in duplicate for each burette4 6,page ,b6 ,date ,,determination ,,of ,,comparative ,,strength ,h,cl vs4 ,na,o,h ,,determination ,i ,,ii ,corrected ,corrected ,final ,reading ,h,cl #dh4ag #dh4jh #dc4bj #dc4ad ,initial ,reading ,h,cl #j4ab 4#ab 4#ag 4#ag ----- ----- ----- ----- #dg4if #db4ig ,corrected ,corrected ,final ,reading ,h,cl #df4cf #df4bi #dj4ea #dj4cg ,initial ,reading ,h,cl #a4ge #a4ge 4#ej 4#ej ----- ----- ----- ----- #dd4ed #ci4hg log cc4 ,na,o,h #a4fdfh #a4fjjh colog cc4 ,h,cl #h4caib #h4cffh ------ ------ #i4ifhj - #aj #i4ifgf - #aj #a cc4 ,h,cl 4#ibij cc4 ,na,o,h 4#ibhb cc4 ,na,o,h ,mean 4#ibhf ,signed 6,page ,c6 ,date ,,standardization ,,of ,,hydrochloric ,,acid ,weight sample and tube #i4agic #h4agca #h4agca #f4iahg ------ ------ ,weight sample #a4jjfb #a4bedd ,final ,reading ,h,cl #ci4ig #ci4hc #di4ij #di4gg ,initial ,reading ,h,cl 4#jj 4#jj 4#jd 4#jd ----- ----- ----- ----- #ci4hc #di4gc ,final ,reading ,na,o,h 4#bf 4#bf 4#fg 4#fg ,initial ,reading ,na,o,h 4#ab 4#ab 4#cf 4#cf --- --- --- --- 4#ad 4#ca 4#ad 4#ca ,corrected cc4 ,h,cl #ci4hc - ----- #ci4fh #di4gc - ----- #di4dj 4#ibhf 4#ibhf log sample #j4jjbe #j4jihc colog cc #h4djad - #aj #h4cjfc - #aj colog milli equivalent #a4bgeg #a4bgeg ------ ------ #i4fgif - #aj #i4fhjc - #aj ,normal value ,h,cl 4#dghb 4#dghi ,mean 4#dghf ,signed 6,page ,d6 ,date ,,determination ,,of ,,chlorine ,,in ,,chloride1 ,,sample ,no4 ,weight sample and tube #af4agba #ae4iigf #ae4iigf #ae4gaag ------- ------- ,weight sample 4#agde 4#bhei ,weight crucible precipitate #ad4ddif #ae4fiae ,constant weights #ad4ddhg #ae4fiae #ad4ddhe ,weight crucible #ad4bbaf #ae4caif ,constant weight #ad4bbaf #ae4caid ,weight ,ag,cl 4#bbfi 4#cgba log ,cl #a4edif #a4edif log weight ,ag,cl #i4ceeh - #aj #i4egjf - #aj log #ajj #b4jjjj #b4jjjj colog ,ag,cl #g4hdch - #aj #g4gdch - #aj colog sample #j4gehc #j4edch ------- ------- #a4ejge #a4ejgh ,cl in sample ,no4 #cb4ah #cb4bj ,signed ,,strength ,,of ,,reagents ,the concentrations given in this table are those suggested for use in the procedures described in the foregoing pages4 ,it is obvious1 however1 that an exact adherence to these quantities is not essential4 ,approx4 ,approx4 ,grams relation relation per to normal to molal liter4 solution solution ,ammonium oxalate1 7,,nh#d7b;,c#b,o#d4;,h#b,o #dj #j4e,n #j4be ,barium chloride1 ,ba,cl#b4b;,h#b,o #be #j4b,n #j4a ,magnesium ammonium chloride 7of ,mg,cl#b7 #ga #a4e,n #j4ge ,mercuric chloride1 ,hg,cl#b #de #j4cc,n #j4ff ,potassium hydroxide1 ,,koh 7sp4 gr4 #a4bg7 #dhj ,potassium thiocyanate1 ,,kscn #e #j4je,n #j4ee ,silver nitrate1 ,ag,n,o#c #ba #j4abe,n #j4abe ,sodium hydroxide1 ,na,o,h #ajj #b4e,n #b4e ,sodium carbonate4 ,na#b;,,co#c #aei #c,n #a4e ,sodium phosphate1 ,na#b;,,hpo#d4ab;,h#b,o #ij #j4e,n or #j4ge,n #j4be ,stannous chloride1 ,sn,cl#b1 made by saturating hydrochloric acid 7sp4 gr4 #a4b7 with tin1 diluting with an equal volume of water1 and adding a slight excess of acid from time to time4 ,a strip of metallic tin is kept in the bottle4 ,a solution of ammonium molybdate is best prepared as follows3 ,stir #ajj grams of molybdic acid 7,mo,o#c7 into #djj cc4 of cold1 distilled water4 ,add #hj cc4 of concentrated ammonium hydroxide 7sp4 gr4 #j4ij74 ,filter1 and pour the filtrate slowly1 with constant stirring1 into a mixture of #djj cc4 concentrated nitric acid 7sp4 gr4 #a4db7 and #fjj cc4 of water4 ,add to the mixture about #j4je gram of microcosmic salt4 ,filter1 after allowing the whole to stand for #bd hours4 ,the following data regarding the common acids and aqueous ammonia are based upon percentages given in the ,standard ,tables of the ,manufacturing ,chemists' ,association of the ,united ,states 6,j4,s4,c4,i461 #bd 7#aije71 #ghg-gij4 ,all gravities are taken at #ae4e;,c4 and compared with water at the same temperature4 ,aqueous ammonia 7sp4 gr4 #j4if7 contains #i4ia per cent ,,nh#c by weight1 and corresponds to a #e4f ,n and #e4f molal solution4 ,aqueous ammonia 7sp4 gr4 #j4ij7 contains #bh4eb per cent ,,nh#c by weight1 and corresponds to a #e4f ,n and #e4f molal solution4 ,hydrochloric acid 7sp4 gr4 #a4ab7 contains #bc4ha per cent ,h,cl by weight1 and corresponds to a #g4c ,n and #g4c molal solution4 ,hydrochloric acid 7sp4 gr4 #a4bj7 contains #ci4hj per cent ,h,cl by weight1 and corresponds to a #ac4a ,n and #ac4a molal solution4 ,nitric acid 7sp4 gr4 #a4bj7 contains #cb4be per cent ,,hno#c by weight1 and corresponds to a #f4a ,n and #f4a molal solution3 ,nitric acid 7sp4 gr4 #a4db7 contains #fi4if per cent ,,hno#c by weight1 and corresponds to a #ae4h ,n and #ae4h molal solution4 ,sulphuric acid 7sp4 gr4 #a4hced7 contains #ic4ai per cent ,h#b,,so#d by weight1 and corresponds to a #cd4h ,n or #ag4d molal solution4 ,sulphuric acid 7sp4 gr4 #a4ah7 contains #bd4gd per cent ,h#b,,so#d by weight1 and corresponds to a #e4i ,n or #b4ie molal solution4 ,the term 6normal6 7,n71 as used above1 has the same significance as in volumetric analyses4 ,the molal solution is assumed to contain one molecular weight in grams in a liter of solution4 ,,densities ,,and ,,volumes ,,of ,,water ,,at ,,temperatures ,,from #ae-cj;,c4 ,temperature ,density4 ,volume4 ,centigrade4 #d #a4jjjjjj #a4jjjjjj #ae #j4iiiabf #a4jjjhgd #af #j4iihigj #a4jjajca #ag #j4iihhja #a4jjabjj #ah #j4iihfbb #a4jjachj #ai #j4iihdcb #a4jjaega #bj #j4iihbcj #a4jjaggc #ba #j4iihjai #a4jjaihe #bb #j4iiggig #a4jjbbjh #bc #j4iigefe #a4jjbdda #bd #j4iigcbc #a4jjbfhe #be #j4iigjga #a4jjbich #bf #j4iifhaj #a4jjcbja #bg #j4iifeci #a4jjcdgc #bh #j4iifbei #a4jjcgee #bi #j4iieiga #a4jjdjdf #cj #j4iiefgc #a4jjdcdf ,authority3 ,landolt1 ,brnstein1 and ,meyerhoffer's 6,tabellen61 third edition4 ,,corrections ,,for ,,change ,,of ,,temperature ,,of ,,standard ,,solutions ,the values below are average values computed from data relating to a considerable number of solutions4 ,they are sufficiently accurate for use in chemical analyses1 except in the comparatively few cases where the highest attainable accuracy is demanded in chemical investigations4 ,the expansion coefficients should then be carefully determined for the solutions employed4 ,for a compilation of the existing data1 consult ,landolt1 ,brnstein1 and ,meyerhoffer's 6,tabellen61 third edition4 ,corrections for #a cc4 ,concentration4 of solution between #ae and #ce;,c4 ,normal 4#jjjbi #j4e ,normal 4#jjjbe #j4a ,normal or more dilute solutions 4#jjjbj ,the volume of solution used should be multiplied by the values given1 and that product multiplied by the number of degrees which the temperature of the solution varies from the standard temperature selected for the laboratory4 ,the total correction thus found is subtracted from the observed burette reading if the temperature is higher than the standard1 or added1 if it is lower4 ,corrections are not usually necessary for variations of temperature of #b;,c4 or less4 ,,international ,,atomic ,,weights #aibj #aibj ,aluminium ,al #bg4a ,molybdenum ,mo #if4j ,antimony ,sb #abj4b ,neodymium ,nd #add4c ,argon ,a #ci4i ,neon ,ne #bj4b ,arsenic ,as #gd4if ,nickel ,ni #eh4fh ,barium ,ba #acg4cg ,nitrogen ,n #ad4jjh ,bismuth ,bi #bjh4j ,osmium ,os #aij4i ,boron ,b #aa4j ,oxygen ,o #af4jj ,bromine ,br #gi4ib ,palladium ,pd #ajf4g ,cadmium ,cd #aab4dj ,phosphorus ,p #ca4jd ,caesium ,cs #acb4ha ,platinum ,pt #aie4b ,calcium ,ca #dj4jg ,potassium ,k #ci4aj ,carbon ,c #ab4jje ,praseodymium ,pr #adj4i ,cerium ,ce #adj4be ,radium ,ra #bbf4j ,chlorine ,cl #ce4df ,rhodium ,rh #ajb4i ,chromium ,cr #eb4j ,rubidium ,rb #he4de ,cobalt ,co #eh4ig ,ruthenium ,ru #aja4g ,columbium ,cb #ic4a ,samarium ,sm #aej4d ,copper ,cu #fc4eg ,scandium ,sc #dd4a ,dysprosium ,dy #afb4e ,selenium ,se #gi4b ,erbium ,er #afg4g ,silicon ,si #bh4c ,europium ,eu #aeb4j ,silver ,ag #ajg4hh ,fluorine ,fl #ai4j ,sodium ,na #bc4jj ,gadolinium ,gd #aeg4c ,strontium ,sr #hg4fc ,gallium ,ga #fi4i ,sulphur ,s #cb4jf ,germanium ,ge #gb4e ,tantalum ,ta #aha4e ,glucinum ,gl #i4a ,tellurium ,te #abg4e ,gold ,au #aig4b ,terbium ,tb #aei4b ,helium ,he #d4jj ,thallium ,tl #bjd4j ,hydrogen ,h #a4jjh ,thorium ,th #bcb4d ,indium ,in #aad4h ,thulium ,tm #afh4e ,iodine ,i #abf4ib ,tin ,sn #aah4g ,iridium ,ir #aic4a ,titanium ,ti #dh4a ,iron ,fe #ee4hd ,tungsten ,w #ahd4j ,krypton ,kr #hb4ib ,uranium ,u #bch4b ,lanthanum ,la #aci4j ,vanadium ,v #ea4j ,lead ,pb #bjg4b ,xenon ,xe #acj4b ,lithium ,li #f4id ,ytterbium ,yb #agc4e ,lutecium ,lu #age4j ,yttrium ,y #hh4g ,magnesium ,mg #bd4cb ,zinc ,zn #fe4cg ,manganese ,mn #ed4ic ,zirconium ,zr #ij4f ,mercury ,hg #bjj4f ,,index ,acidimetry ,acid solutions1 normal standard ,acids1 definition of ,acids1 weak1 action of other acids on action of salts on ,accuracy demanded ,alkalimetry ,alkali solutions1 normal standard ,alumina1 determination of in stibnite ,ammonium nitrate1 acid ,analytical chemistry1 subdivisions of ,antimony1 determination of1 in stibnite ,apatite1 analysis of ,asbestos filters ,atomic weights1 table of ,balances1 essential features of use and care of ,barium sulphate1 determination of sulphur in ,bases1 definition of ,bichromate process for iron ,bleaching powder1 analysis of ,brass1 analysis of ,burette1 description of calibration of cleaning of reading of ,calcium1 determination of1 in limestone ,calibration1 definition of of burettes of flasks ,carbon dioxide1 determination of1 in limestone ,chlorimetry ,chlorine1 gravimetric determination of ,chrome iron ore1 analysis of ,coin1 determination of silver in ,colloidal solution of precipitates ,colorimetric analyses1 definition of ,copper1 determination of1 in brass determination of in copper ores ,crucibles1 use of ,crystalline precipitates ,densities of water ,deposition potentials ,desiccators ,direct methods ,dissociation1 degree of ,economy of time ,electrolytic dissociation1 theory of ,electrolytic separations1 principles of ,end-point1 definition of ,equilibrium1 chemical ,evaporation of liquids ,faraday's law ,feldspar1 analysis of ,ferrous ammonium sulphate1 analysis of ,filters1 folding of how fitted ,filtrates1 testing of ,filtration ,flasks1 graduation of ,funnels ,fusions1 removal of from crucibles ,general directions for gravimetric analysis volumetric analysis ,gooch filter ,gravimetric analysis1 definition of ,hydrochloric acid1 standardization of ,hydrolysis ,ignition of precipitates ,indicators1 definition of for acidimetry preparation of ,indirect methods ,insoluble matter1 determination of in limestone ,integrity ,iodimetry ,ions1 definition of ,iron1 gravimetric determination of volumetric determination of ,jones reductor ,lead1 determination of in brass ,limestone1 analysis of ,limonite1 determination of iron in ,liquids1 evaporation of transfer of ,litmus ,logarithms ,magnesium1 determination of ,mass action1 law of ,measuring instruments ,methyl orange ,moisture1 determination of in limestone ,neutralization methods ,normal solutions1 acid and alkali oxidizing agents reducing agents ,notebooks1 sample pages of ,oxalic acid1 determination of strength of ,oxidation processes ,oxidizing power of pyrolusite ,permanganate process for iron ,phenolphthalein ,phosphoric anhydride1 determination of ,pipette1 calibration of description of ,platinum crucibles1 care of ,precipitates1 colloidal crystalline ignition of separation from filter washing of ,precipitation ,precipitation methods 7volumetric7 ,problems ,pyrolusite1 oxidizing power of ,quantitative ,analyses1 subdivisions of ,reagents1 strength of ,reducing solution1 normal ,reductor1 ,jones ,reversible reactions ,silica1 determination of1 in limestone determination of1 in silicates purification of ,silicic acid1 dehydration of ,silver1 determination of in coin ,soda ash1 alkaline strength of ,sodium chloride1 determination of chlorine in ,solubility product ,solution pressure ,solutions1 normal standard ,standardization1 definition of ,standard solutions1 acidimetry and alkalimetry chlorimetry iodimetry oxidizing and reducing agents thiocyanate ,starch solutions ,stibnite1 determination of antimony in ,stirring rods ,stoichiometry ,strength of reagents ,suction1 use of ,sulphur1 determination of in ferrous ammonium sulphate in barium sulphate ,temperature1 corrections for ,testing of washings ,theory of electrolytic dissociation ,thiocyanate process for silver ,titration1 definition of ,transfer of liquids ,volumetric analysis1 definition of general directions ,wash-bottles ,washed filters ,washing of precipitates ,washings1 testing of ,water1 ionization of densities of ,weights1 care of ,zimmermann-,reinhardt method for iron ,zinc1 determination of1 in brass