PRACTICAL AND ANALYTICAL CHEMISTRY. BEING A COMPLETE COURSE IN CHEMICAL ANALYSIS. BY HENRY TRIMBLE, Pn.G., PROFESSOR OF ANALYTICAL CHEMISTRY IN THE PHILADELPHIA COLLEGE OF PHARMACY. OF TOT! UNIVERSITY SECOND EDITION. REVISED, ENLARGED AND ILLUSTRATED. PHILADELPHIA: P. BLAKISTON, SON & CO., No. 1012 WALNUT STREET. 1886. COPYRIGHT, 1886, by P. BLAKISTON, SON & CO. PREFACE TO FIRST EDITION. The increased amount of time devoted by students of phar- macy and medicine to analytical chemistry has directed more attention to the subject of imparting instruction in this science. The object of the present volume is to place before the student as compact a course as possible, in order to enable him to become familiar with the subject in the necessarily limited time at his disposal. The author's experience has led him to believe that a study of Qualitative Analysis should be preceded by some experi- ence in the preparation of the more important gases and a few of the salts. Such practice requires the student to familiarize himself with the construction of apparatus as well as with the processes of filtration, evaporation, crystallization, ignition, etc. The examples for preparation may be increased at the option of the instructor. In Part II the student should perform the reactions of each group, and then be furnished with a solution containing some, or all, of the bases of the group. This should be followed by a solution in which he should search for all the elements previously studied ; such practice being repeated until he can correctly determine all the bases present. In order to enable the student to see the comparative effect of the group reagents readily, a summary has been introduced at the end of each group. This he should be able to write VII viii PREFACE TO THE FIRST EDITION. out, without the use of the book, before attempting to analyze a group solution. By progressive steps he is thus led rapidly on to successfully examine the more complex solutions for both bases and acids. The grouping of the bases is, to a certain extent, new, but it places together those elements which are very closely related, and, in addition, adapts each group to the time of one lesson ; which may be repeated if desirable. In the Part devoted to gravimetric and volumetric analysis, the examples are limited in number ; but this much is intended to give the student an opportunity to learn the methods with the aid of an instructor, so that he may afterwards pursue the subject alone, with the aid of a book like Fresenius's Quanti- tative Analysis. In conclusion, it is but just to state that many of the works on Qualitative Analysis have been consulted. Those of Attfield and Muter furnished many valuable hints, but Fresenius's admirable work has been used as authority, and, to those who wish to pursue the subject in greater detail than here offered, it is recommended for reference. H. T. Philadelphia, August igtk, 1883. PREFACE TO THE SECOND EDITION. Only two important changes have been thought advisable in preparing the second edition of this book. I. The addition of equations explaining the more important reactions which occur in the action of reagents on the various salts. The typical ones are thus expressed, and the others are intended to be written by the student without aid. It is believed, where there is time, that the construction of equations by the student will be of great service to him, and prevent too hasty performance of the reactions. It is the habit of the author to question the student closely on the tests of each element in a group, before allowing him to proceed to the separation. This has been found especially useful in causing the student to think for himself, and to devise, from his knowledge of the reactions, a method for separation of the group. II. While the detection of acids in practice is comparatively simple, it has been found that it is the source of much diffi- culty to the beginner. This arises from the absence of a systematic grouping of the acids, and the separation of a few before attempting all. A plan similar to that adopted for bases is now proposed, which will undoubtedly be found an aid to the student. While it may require a little more time to study the acids in this manner, it cannot but be of service in facilitating the analysis of salts. ix X PREFACE TO THE SECOND EDITION. One word must be said in explanation of the use of formulas and symbols for the reagents, instead of the full chemical name. While it may be inconvenient at first, the student soon learns them, almost without effort, and thus becomes familiar with what he should know, but should not be compelled to memorize. The author's thanks are extended to those who have kindly furnished suggestions for this edition, and also to the great number who, without exception, favorably criticised the first edition. The exhaustion of the first edition in less than a year is satisfactory evidence that an American book on this subject, adapted especially to the requirements of pharmaceutical and medical students, was needed. H. T. Philadelphia, July jth, 1886. CONTENTS. PACK PART FIRST PRACTICAL CHEMISTRY, . . .17 SECTION I PREPARATION AND PROPERTIES OF GASES, . 17 Hydrogen, Chlorine, Hydrochloric Acid, Oxygen, Nitrogen, Ammo- nia, Nitric Acid, Carbon Dioxide, 17~ 2 5 SECTION II PREPARATION OF SALTS, .... 26 Potassium Chloride, Potassium and Sodium Tartrate, Ammonium Nitrate, Ammonium Oxalate, Calcium Phosphate, Magnesium Sulphate, Magnesium Carbonate, Magnesium Oxide, Alum- inium Hydrate, Ferrous Sulphate, Ferric Sulphate, Ferric Hydrate, Copper Sulphate, Lead Acetate, . . . 26-30 PART SECOND QUALITATIVE ANALYSIS, . 33 SECTION I BASES, 33 GROUP I Reactions of Potassium, Sodium, Lithium, Ammonium, 33 Summary of Group I, . . . , . . . . -35 Analysis of Group I ........ 35 GROUP II Reactions of Barium, Strontium, Calcium, Magne- sium, 35-37 Summary of Group II, ........ 3^ Analysis of Group II, ........ 38 Chart for Analysis of Groups I and II, . . . . -39 Precautions and Observations on Chart, .... 39 GROUP III Reactions of Manganese, Zinc, Cobalt, Nickel, . 40-42 Summary and Analysis of Group III, 43 Chart for Analysis of Groups I to III, inclusive, . . -44 xi XI 1 CONTENTS. PAGE GROUP IV Reactions of Iron, Cerium, Aluminium, Chromium, 45-47 Summary and Analysis of Group IV, ..... 48 Chart for Analysis of Groups I to IV, inclusive, ... 49 Precautions and Observations on Chart, . . . . 50 GROUP V Reactions of Arsenic, Antimony, Tin, Gold, Platinum, 50-54 Summary and Analysis of Group V, 55 Chart for Analysis of Groups I to V, inclusive, ... 56 Precautions and Observations on Chart, . . . . -57 GROUP VI Reactions of Mercury(ic), Bismuth, Copper, Cad- mium, 57-59 Summary and Analysis of Group VI, 60 Chart for Analysis of Groups I to VI, inclusive, . . .61 Precautions and Observations on Chart, 62 GROUP VII Reactions of Silver, Mercury (ous), Lead, . . 62-64 Summary and Analysis of Group VII, 64 Chart for Analysis of Groups I to VII, inclusive, ... -65 Precautions and Observations on Chart, .... 66 SECTION II ACIDS, . . . . . . . .67 GROUP I Reactions of Acids Hydrochloric, Hydrobromic, Hydri- odic, Hydrofluoric, Hydrocyanic, 67-69 Summary of Group I, ........ 69 Analysis of Group I, . . . . . . . . .69 GROUP II Hypochlorous, Chloric, Water, Hydrates, Oxides, Hy- drosulphuric, Sulphurous, Sulphuric, Thiosulphuric, Nitric, Hypophosphorous, Orthophosphoric, Pyrophosphoric, Meta- phosphoric, Boric, Carbonic, Silicic, .... 70-75 Analysis of Group II, 75 GROUP III Acetic, Valerianic, Stearic, Oleic, Lactic, Oxalic, Suc- cinic, Malic, Tartaric, Citric, Carbolic, Benzoic, Salicylic, Gallic, Tannic, 77~8o Analysis of Group III, 81 CONTENTS. Xlll PAGE SECTION III DETECTION OF BASES AND ACIDS, . . 82-86 Chart of Solubilities, 87 SECTION IV REACTIONS AND TESTS OF ORGANIC COMPOUNDS, 88 Alkaloids, . . 9 I- 93 Neutral Principles, 94 PART THIRD QUANTITATIVE ANALYSIS, . . 97 SECTION I GRAVIMETRIC ESTIMATION, .... 97 Preliminary Directions, 97 Examples in Gravimetric Estimation, Barium, Chlorine, Copper, Sulphuric Acid, Potassium, Nitric Acid, Calcium, Carbonic Acid, 98-101 SECTION II VOLUMETRIC ESTIMATION, . . . . 102 Examples, Oxalic Acid, Sodium Hydrate, Potassium Bichromate, Iodine, Sodium Hyposulphite, Silver Nitrate, . 102-106 Table of Elements, Symbols and Atomic Weights, . . . 107 Index, 109 PART FIRST. PRACTICAL fiti PRACTICAL AND ANALYTICAL CHEMISTRY. PART I. PRACTICAL. SECTION I. PREPARATION AND PROPERTIES OF GASES. HYDROGEN, H. Preparation. Place a few fragments of zinc in a flat-bottom flask of about one-fourth liter capacity ; cover the zinc with water and adapt a cork, through which pass two tubes (Fig. i), one just reaching through the cork and bent so the long end may be dipped under water, the other running directly from a short distance above the cork nearly to the bottom of the flask, so as to be below the surface of the liquid. The upper end of this tube should have a small funnel placed in it, or be enlarged by softening in the flame, inserting and revolving a file or similar instrument, previously warmed. Add, slowly, a small quantity of sulphuric acid, through this tube, and notice an immediate effervescence, with the escape of bubbles through the water in which the exit tube dips. Fill a test tube with water, and, keeping the open end under the liquid, bring it over the tube, so as to collect the gas. When full, close with the thumb, and, bringing the mouth of the tube near a flame, quickly remove the thumb and allow the gas to ignite. It will burn quietly if the gas be pure, but with a slight explosion if it be mixed with air. Properties. This gas is Hydrogen, and its physical prop- erties may now be studied by observing that it is insoluble in water, and without odor, color or taste. 2 17 18 PRACTICAL CHEMISTRY. EXPERIMENT I. Collect a tube full, and holding it, covered, in a vertical position, bring a lighted taper a short distance above its mouth and remove the cover ; the gas will ignite, showing its great levity. EXPERIMENT II. Another tube, similarly filled, is held in an inverted position, and the cover removed ; it will be found, even after the lapse of some time, that the hydrogen at the mouth of the tube may be ignited, thus demonstrating that the gas is too light to come down and out the mouth of the tube. These two experiments have also demonstrated the combus FIG. i. tibility of the gas, which, when pure, burns quietly, with a colorless flame. If, however, it be mixed with air and flame applied, a violent explosion ensues. Therefore, the tube from the generator should never be brought near a flame until it is certain all the air has been expelled. This is determined by trying a test tube full; if it burn quietly, the jet may be lighted. This precaution should always be observed. EXPERIMENT III. On bringing a lighted taper to the mouth of a tube full of hydrogen, the gas is ignited, but on pushing the burning taper up into the gas its flame is extinguished, PREPARATION AND PROPERTIES OF GASES. 19 thus showing that while hydrogen is combustible it is not a supporter of combustion. EXPERIMENT IV. Continue the addition of acid to the zinc until the latter is nearly all dissolved; disconnect the apparatus, pour the liquid on a filter, collect the filtrate in a small beaker or evaporating dish, concentrate and set aside for twenty-four hours, to crystallize. These crystals are zinc FIG. 2. sulphate, ZnSO 4 , the result of a combination of the sulphuric acid and the zinc, as follows : Zn -f H 2 SO 4 = ZnSO 4 + H 2 . CHLORINE, Cl. Preparation. In a flask, arranged so that heat may be applied (Fig. 2), place a small quantity of manganese dioxide, 20 PRACTICAL CHEMISTRY. add hydrochloric acid, agitate well, to moisten all the powder on the bottom, and apply heat. A yellowish-green gas is evolved, which, being somewhat soluble in water, may be collected by downward displacement, that is, by running the delivery tube to the bottom of the receptacle, loosely covered, the heavy gas displaces the lighter air. If the evolution be moderately active it may be collected over warm water, as only a small loss occurs. Care should be taken to avoid inhaling the gas, by passing it, when not collecting, into a solution of potassium or sodium hydrate. The following expresses the reaction in symbols : MnO 2 + 4HC1 = MnCl 2 + 2H 2 O + C1 2 . Properties. EXPERIMENT I. Pass the gas into water; it is absorbed; if this be continued until the water is saturated, it will be found to have absorbed about twice its volume of the gas ; the Aqua Chlori of the Pharmacopoeia is the resulting product. EXPERIMENT II. A tube full of the gas held with mouth upward, and a lighted taper applied, fails to ignite. Push the taper into the gas ; it is extinguished, or only burns with a small, dense, smoky flame, the result of a combination of the chlorine with the hydrogen of the wax, liberating the carbon. EXPERIMENT III. Into a tube full of the gas put a piece of brightly dyed calico, previously moistened ; it is rapidly bleached. Writing on paper is similarly decolorized, but printing is not affected, as it contains carbon, in the form of lampblack, which is not acted on by the gas. If the experi- ments be made with chlorine which has been passed through sulphuric acid to dry it, and the materials are not moistened, no decoloration takes place. The process of bleaching by chlorine is one of oxidation ; it combines with the hydrogen of the water, forming hydro- chloric acid, while the liberated oxygen in the nascent state readily attacks the coloring matter, water and a colorless com- pound resulting. HYDROCHLORIC ACID, HC1. Preparation. Collect one test tube full of hydrogen and one of chlorine, bring their mouths together (the hydrogen tube above with mouth down, as it is lighter), turn over once PREPARATION AND PROPERTIES OF GASES. 21 or twice, so as to thoroughly mix, and open their mouths to a flame; a sharp report will occur, with the development of strongly acid fumes, which will be recognized by future tests as hydrochloric acid : H 2 + C1 2 = 2HC1. To prepare a quantity of it, the apparatus used for the preparation of chlorine serves best. Put into the flask some sodium chloride (common salt) add sulphuric acid slowly, and when the evolution of gas ceases apply a gentle heat. Collect by downward displacement or over mercury : NaCl + H 2 SO 4 r= NaHSO 4 -f HC1. The above expresses the reaction when an excess of acid has been used, which is preferable, as the resulting acid sodium sulphate is easily dissolved out of the flask with water. On the large scale the following more economical method is used : (NaCl) 2 + H 2 SQ 4 = Na 2 S0 4 + (HC1) 2 . Properties. The pungent, suffocating odor and freedom from color are noted with its production. EXPERIMENT I. A test tube or jar of the gas placed with the open mouth under water will so rapidly dissolve that the liquid rises in the vessel. This solution of the gas in water, when of the proper strength, is the Acidum Hydrochloricum, U. S. P. EXPERIMENT II. A piece of moistened blue litmus paper held in a tube of the gas is instantly reddened. EXPERIMENT III. A lighted taper applied to the gas fails to ignite it, and is extinguished if lowered into it. EXPERIMENT IV. Bring a rod moistened with ammonia over the mouth of a tube full of the gas : dense white fumes of ammonium chloride are formed. The mixture remaining in the flask dissolved in warm water, treated with sodium carbonate so long as effervescence occurs, to neutralize the excess of sulphuric acid, concen- trated, filtered and set aside to crystallize, yields sodium sul- phate (Glauber salt). OXYGEN, O. Preparation. Place a few crystals of potassium chlorate in a test tube, adapt a delivery tube long enough to reach 22 PRACTICAL CHEMISTRY. under the surface of some water near by. Apply a steady flame ; as soon as the bubbles of gas escape freely and the air has been expelled, bring a test tube filled with water over the escaping gas, and collect. It is Oxygen, produced from the potassium chlorate by heat, according to the following reaction : KC10 3 = KC1 + 8 . Properties. The appearance and insolubility of the gas in water are noted as it is collected. EXPERIMENT I. The gas is not ignited by the application of a lighted taper. The taper, however, will burn with greatly- increased energy if it be plunged into the gas. If the flame be extinguished and the taper again brought into the gas, provided a spark remain, the taper is rekindled. EXPERIMENT II. A piece of charcoal, previously ignited, is lowered into the gas ; a rapid combustion ensues, and the charcoal disappears. Pour some lime water into the tube, agitate well ; a white precipitate of calcium carbonate is pro- duced. If this be tried with oxygen previous to the burning of the charcoal, no precipitate will be formed. A number of other substances, as sulphur, phosphorus, and even iron, when once kindled, will burn in oxygen with great brilliancy, form- ing characteristic oxides. EXPERIMENT III. Take two test tubes, one about twice the capacity of the other. Fill the larger with hydrogen and the smaller with oxygen, bring their mouths together, and, after turning once or twice, to thoroughly mix their contents, open them and apply flame. A sharp report is caused by the com- bination to form water. If more than two volumes of hydrogen to one of oxygen are present, the surplus remains uncombined; if oxygen is present in greater proportion, the excess of it remains. NITROGEN, N. Preparation. The usual method of preparing this gas is to deprive air of its oxygen, leaving the nitrogen pure. This is best accomplished by placing a small fragment of phos- phorus on a cork covered with some fireproof material. Float the cork and phosphorus on water, ignite the latter and bring PREPARATION AND PROPERTIES OF GASES. 23 over it a bell-jar. The phosphorus combines with the oxygen, converting it into phosphorus pentoxide, P 2 O 5 , which dissolves in the water present, thus leaving the nitrogen pure. A better method for obtaining larger quantities is to heat a mixture of potassium nitrite and ammonium chloride dissolved in water. When the reaction begins, the temperature must be carefully watched, in order to prevent the too rapid evolution of the gas: KNO a + NH 4 C1 = KC1 + 2 H 2 O + N 2 . Properties. The physical properties have been observed during its preparation and collection. In regard to chemical properties, it is inert in the free state. Its compounds, how- ever, are very energetic. AMMONIA, MH 3 . Preparation. In a test tube or evaporating dish mix equal quantities of powdered calcium oxide (quicklime) and ammo- nium chloride, with a few drops of water; the odor of ammonia will be immediately developed : CaO + (NH 4 C1) 2 = CaCl 2 -f H 2 O (NH 8 ) 2 . In smaller quantities the gas may be recognized by holding over the mixture a strip of moistened red litmus paper ; it will slowly become blue ; or similarly hold a glass rod moistened with hydrochloric acid ; dense white fumes of ammonium chloride will form. To prepare larger quantities of the gas, heat the ordinary water of ammonia, which, at a comparatively low temperature, gives it off freely. If it be desired to dry the gas, it must be passed over quicklime. Ammonia is col- lected by upward displacement, that is, by passing the delivery- tube upward into a jar or test tube inverted over it; being lighter than air the latter is diplaced. Properties. EXPERIMENT I. Place a vessel filled with ammonia gas, mouth downward, into some water, and agitate slightly; the water will rise in the vessel rapidly, nearly filling it, showing the great solubility of the gas in water. The other physical properties, as color, odor, etc., have been noted during its collection. EXPERIMENT II. On applying a lighted taper to the gas it does not burn; if, however, it be mixed with oxygen it will 24 PRACTICAL CHEMISTRY. ignite readily. On introducing the taper into the gas it is extinguished. NITRIC ACID, HNO 3 . Preparation. Place a small quantity of potassium nitrate in a test tube, and cover it with strong sulphuric acid. Apply a gentle heat ; brown, strongly acid fumes are given off. Dilute with a little water and add indigo solution; it is decolor- ized. This is a characteristic test for nitric acid. To prepare a larger quantity a retort is used, to which is adapted a glass receiver. The nitric acid distills over on the application of a moderate heat, forming a reddish-yellow liquid, which rapidly attacks and destroys organic matter. Two reactions may be employed to represent the produc- tion of nitric acid, depending on the relative quantity of the materials used. In the first case an excess of sulphuric acid gives KNO 3 -f H 2 SO 4 = KHSO 4 + HNO 3 . In the second case just a sufficient amount of sulphuric acid is used to decompose the potassium nitrate (KNO 3 ) 2 + H 2 SO 4 = K 2 SO 4 + (HNO 3 ) 2 . On the large scale sodium nitrate is now employed in place of the potassium salt, on account of its lower price. CARBON DIOXIDE, CO 2 . Preparation. The flask used in the preparation of hydro- gen will serve for making carbon dioxide. A few pieces of marble are placed in the flask, covered with water, and hydro- chloric acid added. A brisk effervescence ensues, and the gas being somewhat soluble in water is collected by downward displacement CaC0 3 -+ (HC1) 2 = CaCl 2 + H 2 O + CO 2 . Sulphuric acid should not be used, as it forms an insoluble calcium sulphate which is very difficult to remove from the flask. Properties. EXPERIMENT I. Pour some clear lime water into a jar of the gas and agitate ; the solution immediately becomes cloudy, owing to formation of insoluble calcium car- bonate. If more gas be passed into the mixture it will become clear again, on account of the solubility of the precipitate in carbonic acid. PREPARATION AND PROPERTIES OF GASES. 25 EXPERIMENT II. Add solution of potassium hydrate to a jar of the gas, close and shake well ; the gas is absorbed by the alkali, as may be shown by placing the mouth of the jar under water and removing the stopper, when the water will rush in, nearly filling it. EXPERIMENT III. A lighted taper lowered into the gas is immediately extinguished. The same result is accomplished by opening the vessel some distance above the flame and allowing the gas to flow down upon it. This latter experi- ment also illustrates the great density of the gas, which is twenty-two times heavier than hydrogen. 26 PRACTICAL CHEMISTRY. SECTION II. PREPARATION OF SALTS. POTASSIUM CHLORIDE, KC1. Preparation. One or two cubic centimeters of hydro- chloric acid, diluted with three or four times its bulk of water, are placed in a small beaker glass, and potassium carbonate added so long as effervescence occurs, and until after boiling (to remove CO 2 ) the solution is neutral to litmus paper, that is, when the blue litmus paper is not changed to red nor the red changed to blue. Evaporate to a small bulk and set aside to crystallize. The cubic crystals which separate after standing twenty-four hours may be collected on filter paper and dried at a moderate temperature : K 2 C0 3 + (HC1) 2 = (KC1) 2 + H 2 + C0 2 . Potassium chloride is rarely prepared in this manner, except for practice, as it occurs largely in nature, and is used for preparing many other potassium salts. POTASSIUM AND SODIUM TARTRATE, KNaC 4 H 4 O 6 4H 2 O. (ROCHELLE SALT.) Preparation. Heat, in a porcelain capsule, a solution of sodium carbonate, and add to it potassium bitartrate until effervescence ceases, and the solution (after the escape of CO 2 ) is neutral to litmus. On filtering and cooling, crystals of Rochelle salt are deposited, rapidly and in small crystals if the solution be concentrated, but slowly and in much larger ones if the solution be dilute : Na 2 C0 3 + (KHC 4 H 4 6 ) 2 i= (KNaC 4 H 4 O 6 ) 2 -f H 2 6 + CO 2 . Note on Calculation. In order to calculate the amount of each salt to use in the above process, we notice the number of molecules of each employed, and multiply this by the sum of the atomic weights. (that is, by the molecular weights). In the above case one molecule of Na 2 CO 3 = 106, and two molecules of KHC 4 H 4 O 6 = 2 X 1 88. 1 = 376.2. Therefore every 106 parts of anhydrous sodium carbonate require 376.2 parts of potassium bitartrate, to form Rochelle salt. If we have 50 grams of anhydrous sodium carbonate and wish to convert it into Rochelle salt, we use the following formula : As 106 : 376.2 : : 50 grams : number of grams of potassium bitartrate required, = 177.4 grams. PREPARATION OF SALTS. 27 AMMONIUM NITRATE, NH 4 NO 3 . Preparation. Add to about 20 c.c. of dilute nitric acid, in a beaker glass, sufficient ammonia water to give it a distinct ammoniacal odor; filter, concentrate, keeping the ammonia in slight excess, and set aside in cool place for crystals to form : NH 4 OH -f HNO 3 = NH 4 NO 3 -f H 2 O. Properties. These crystals contain twelve molecules of water of crystallization, which it is desirable to get rid of be- fore using the salt. By exposure to a temperature of 155 C. the water gradually escapes, and the fused or granulated salt is ready to be converted into nitrogen monoxide (laughing gas), which takes place at about 185 C. according to the following reaction : NH 4 N0 3 =N 2 + (H 2 0) 2 . AMMONIUM OXALATE, (NH 4 ) 2 C 2 O 4 . Preparation. Dilute 20 c.c. of solution of ammonia with twice its bulk of water, add a solution of oxalic acid until neutral, concentrate slightly, filter and set aside to crystallize. The crystals may be collected on a filter, and another crop obtained by concentrating the "mother liquor ": (NH 4 OH) 2 + H 2 C 2 4 == (NH 4 ) 2 C 2 4 + (H 2 O) 2 . Ammonium carbonate is sometimes used for combining with the oxalic acid, but the neutralization is not as easily effected, besides it is not desirable on the ground of economy. CALCIUM PHOSPHATE, Ca 3 (PO 4 ) 2 . Preparation. Finely powdered bone ash is digested for a short time with 'diluted hydrochloric acid. The solution filtered, boiled, filtered again, if necessary. The filtrate is treated with ammonia until it smells strongly of it. Collect the precipitate on a filter, wash by pouring on warm water until the washings are tasteless, and dry at a low temperature. The resulting powder is calcium phosphate, which exists in the bone ash and is dissolved by hydrochloric acid, forming acid calcium phosphate, as follows : Ca 3 (P0 4 ) 2 + (HC1) 4 = CaH 4 (P0 4 ) 2 + (CaCl 2 ) 2 . From this solution it is precipitated by ammonia, as follows : CaH 4 (P0 4 ) 2 + (CaCl 2 ) 2 -f (NH 4 OH) 4 = Ca 3 (PO 4 ) 2 + (NH 4 C1) 4 . OF THTC UNIVERSITY H 28 PRACTICAL CHEMISTRY. In addition to the ordinary apparatus with which a student supplies himself, there is required a wash bottle (Fig. 3), which it is well for every student to construct for himself, as it furnishes him valuable practice in cutting and bending glass tubing. This bottle is used in washing all precipitates, and is convenient as a water supply, which may be kept hot, if desired. MAGNESIUM SULPHATE, MgSO 4 .;H 2 O. Preparation. To about 5 c.c. of sulphuric acid, diluted with five or six times its volume of water, heated in a capsule, FIG. 3. add powdered magnesium carbonate until effervescence ceases, and filter. Concentrate and set aside to crystallize. (MgC0 3 ) 4 Mg(OH) 2 . 5 H 2 + (H 2 S0 4 ) 5 = (MgSO 4 ) 5 + (H.O)^ MAGNESIUM CARBONATE, (MgCO 3 ) 4 Mg(OH) 2 .5H 2 O. Preparation. On mixing solutions of magnesium sulphate and sodium carbonate and boiling, we get magnesium car- bonate precipitated, while carbon dioxide escapes. The pre- cipitate is very variable in its composition, depending on the concentration of the solutions. When the U. S. P. product is obtained the following equation expresses the reaction : (MgS0 4 ) 5 + (Na 2 C0 3 ) 5 + (H 2 0) 6 = (MgC0 3 ) 4 Mg(OH J2 . 5 H 2 + (Na 2 S0 4 ) 5 + CO 2 . PREPARATION OF SALTS. 29 The precipitate washed with hot water and dried, serves for the following example of a compound prepared by ignition. MAGNESIUM OXIDE, MgO. Preparation. Heat some of the magnesium carbonate, prepared in the above reaction, in a porcelain crucible until, on taking out a small portion, placing in a test tube with a little water, heating to remove air bubbles, and adding a drop or two of hydrochloric acid, no effervescence is produced. This will require some time, and great care is necessary to determine when the powder fails to give an effervescence with the acid. (MgC0 3 ) 4 Mg(OH) 25 H 2 = (MgO) 5 + (CO 2 ) 4 + (H 2 O) 6 . Zinc oxide may be prepared in a similar manner, from zinc carbonate. This differs from the magnesium oxide by being yellow while hot, and very pale yellow 'when cold. ALUMINIUM HYDRATE, A1 2 (OH) 6 . Preparation. To a solution of alum add a solution of sodium carbonate and boil. Allow the precipitate to settle, decant the clear supernatant liquid on a filter, add more hot water to the precipitate and again decant. Collect the pre- cipitate on the filter, wash well with hot water and dry ; the resulting white powder is Aluminii Hydras, U. S. P. A1 2 (S0 4 ) 3 , K 2 S0 4 -f (Na 2 C0 3 ) 3 + (H 2 O) 3 = A1 2 (OH) 6 + K 2 S0 4 + (Na 2 S0 4 ) 3 + (CO 2 ) 3 . FERROUS SULPHATE, FeSO 4 .;H 2 O. Preparation. Add enough dilute sulphuric acid to some iron filings, or wire, in a beaker, to cover them. AlldWthe reaction to proceed, assisted by a little heat, until effervescence ceases. Filter from the excess of iron, concentrate, filter and crystallize. (Fe) 2 + (H 2 S0 4 ) 2 = (FeS0 4 ) 2 + (H a ) a . These crystals should be rapidly dried and preserved in well stopped bottles, as they quickly become converted into ferric sulphate on exposure to air. FERRIC SULPHATE, Fe 2 (SO 4 ) 3 . Preparation. To a strong solution of ferrous sulphate add one-fourth its bulk of sulphuric acid, heat to the boiling point and drop in nitric acid as long as effervescence is produced 30 PRACTICAL CHEMISTRY. and until the resulting liquid becomes of a clear reddish- brown color. (FeS0 4 ) 6 + (H 2 S0 4 ) 3 + (HN0 3 ) 2 = (Fe 2 (SO 4 ) 3 ) 3 + N 2 O 2 -f (H 2 O) 4 . This is the Liquor Ferri Tersulphatis of the Pharmacopoeia ; and is the most convenient compound to use in the preparation of some of the other iron salts. FERRIC HYDRATE, Fe 2 (OH) 6 . Preparation. Dilute some of the above ferric sulphate solution with an equal bulk of water, add solution of am- monia until, after stirring, it smells strongly. The resulting precipitate is ferric hydrate, the well-known antidote to arsenic. When needed for this purpose, it is sufficient to pour the mixture on a muslin strainer, wash once or twice until the saline taste nearly disappears from the washings, when the compound is ready for use. Fe 2 (S0 4 ) 3 + (NH 4 OH) 6 = Fe 2 (OH) 6 + ( (NH 4 ) 2 SO 4 ) 3 . 7 This preparation should always be freshly prepared when wanted for use as an antidote, as it loses H 2 O on keeping, becoming a mixture of ferric oxide Fe 2 O 3 and hydrate. This change takes place, although more slowly, when the com- pound is kept under water. COPPER SULPHATE, CuSO 4 .sH 2 O. Preparation. Heat copper turnings for some time, with strong sulphuric acid, in a fume closet, until the reaction ceases. Dilute with water, filter and crystallize. Cu -f (H 2 SO 4 ) 2 = CuSO 4 + SO 2 + (H 2 O) 2 . LEAD ACETATE, Pb(C 2 H 8 O 2 ) 2 .3H 2 O. Preparation. Lead oxide (litharge) is boiled with three or four times its weight of acetic acid, in a capsule, adding water from time to time, with more acid, if necessary, until most of the oxide has disappeared. Filter, concentrate, keeping the solution acid, and set aside to crystallize : PbO -f (HC 2 H 3 O 2 ) 2 = Pb(C 2 H 3 O 2 ) 2 + H 2 O. The solution or crystals should not be exposed to the fumes of the laboratory, for if there be only a small quantity of hydrogen sulphide in the room, they will become black. PART SECOND. QUALITATIVE ANALYSIS. PART II. QUALITATIVE ANALYSIS. SECTION I. BASES. GROUP I. POTASSIUM, SODIUM, LITHIUM, AMMONIUM. REACTIONS OF POTASSIUM (K). Use a solution of potassium chloride (KC1). 1. PtCl 4 causes a yellow crystalline precipitate of K 2 PtCl 6 , soluble in excess of water. PtCl 4 + 2KC1 = K 2 PtCl 6 . The delicacy of this reaction is increased by the addition of alcohol. 2. H 2 C 4 H 4 O 6 , in concentrated solution, produces a white crys- talline precipitate of potassium acid tartrate KHQH 4 O 6 , soluble in excess of water, readily in hot water, acids or potas- sium hydrate. H 2 C 4 H 4 O 6 + KC1 = KHC 4 H 4 O 6 + HC1. The addition of alcohol and violent agitation facilitate the formation of this precipitate. 3. A fragment of potassium salt on the loop of a platinum wire, held in the colorless flame of a Bunsen gas lamp imparts a violet color. This reaction is interfered with by the presence of sodium salts, which color ftie flame yellow ; organic matter also colors the flame violet, and should be removed by ignition before testing for potassium. The yellow rays of sodium may be destroyed by viewing the flame through blue glass. 4. Potassium salts are not volatile at a low red heat ; at a white heat they are slowly volatilized. REACTIONS OF SODIUM (Na). Use a solution of sodium chloride (NaCl). I. Sodium salts color the gas flame yellow; so delicate is this reaction that the merest traces are revealed by it. 3 33 34 ANALYTICAL CHEMISTRY. 2. The salts of sodium are not volatile at a low red heat, but slowly volatilize at a white heat. REACTIONS OF LITHIUM (Li). Use a solution of lithium chloride (LiCl). 1. Na 2 HPO 4 added to a strong solution produces, on boiling, a white precipitate of lithium phosphate Li 3 PO 4 . Na 2 HPO 4 -f 3L1C1 = Li 3 PO 4 + 2NaCl -f HC1. This reaction takes place more readily when the solution is first made alkaline with NH 4 OH. 2. Lithium salts impart an intense crimson color to the gas flame. This is somewhat interfered with by sodium salts, but the yellow color of sodium may be excluded by blue glass, which if not too dark will allow the crimson rays of lithium to pass through. These must not be confused with the violet potassium rays, which will pass through a deep blue glass. 3. Lithium salts do not volatilize at a low red heat, but are slowly volatilized at a white heat. REACTIONS. OF AMMONIUM (NH 4 ). Use a solution of ammonium chloride (NH 4 C1). 1. PtCl 4 produces, in strong solution, a yellow crystalline precipitate of ammonium platino-chloride (NH 4 ) 2 PtCl 6 . 2. NaOH on heating causes the evolution of ammonia NH 3 ; detected by the odor; by holding near a glass rod moistened with HC1, which will produce dense white fumes of NH 4 C1 ; or by holding in the mouth of the tube a strip of moistened red litmus paper, when it will immediately become blue. Care must be taken, in this last test, to prevent any of the alkaline liquid coming in contact with the paper, as it would likewise cause the blue color. 3. H 2 C 4 H 4 O 6 , added to a concentrated solution, produces a white precipitate of ammonium acid tartrate NH 4 HC 4 H 4 O 6 , soluble in slight excess of water. 4. Ammonium salts are volatile at a low red heat. BASES. 35 SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP I. K Na Li NH 4 PtCl 4 Yellow Precipitate No Precipitate No Precipitate Yellow Precipitate H 2 C 4 H 4 6 White Precipitate No Precipitate No Precipitate White Precipitate Flame Violet Yellow Crimson None Volatility Not Volatile Not Volatile Not Volatile Volatile DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP I. To a small portion of the solution add NaOH and heat : NH 3 will be given off if ammonium salts are present, and may be detected by the odor, or by moistened red litmus paper. Evaporate another portion of the solution to dryness, transfer to a porcelain crucible, and heat until the white fumes of ammonium salts cease to be given off. Dissolve the residue in a few drops of H 2 O with a drop or two of HC1, and add PtCl 4 ; K, if present, will be precipitated. A loop of platinum wire dipped in the original solution and held in the colorless gas flame will give evidence of Na and Li. If Na be present in excess, so as to obscure the Li flame, evaporate a portion of the original solution to dryness, dis- solve in the smallest possible amount of H 2 O, add Na 2 HPO 4 and boil, filter off the Li 3 PO 4 , wash with a little hot water containing NH 4 OH, and dissolve in a few drops of HC1. With this solution Li may be detected by the flame test. GROUP II. BARIUM, STRONTIUM, CALCIUM, MAGNESIUM. REACTIONS OF BARIUM (Ba). Use a solution of barium chloride (BaCl 2 ). I. H 2 SO 4 produces an immediate precipitate of barium sulphate BaSO 4 , insoluble in boiling hydrochloric or nitric acid. BaCl 2 + H 2 SO 4 = BaSO 4 -f 2HC1. 36 ANALYTICAL CHEMISTRY. 2. K 2 CrO 4 , even in dilute solutions, causes a yellow precipi- tate of barium chromate BaCrO 4 , soluble in hydrochloric or nitric acid, but insoluble in acetic acid. BaCl 2 + K 2 CrO 4 = BaCrO 4 + 2KC1. 3. (NH 4 ) 2 CO 3 precipitates white barium carbonate BaCO 3 , soluble in acetic acid. BaCl 2 + (NH 4 ) 2 CO 3 = BaCO 3 -f 2NH 4 C1. 4. (NH 4 ) 2 HPO 4 produces a white precipitate of barium phosphate BaHPO 4 , soluble in acetic and in hydrochloric acid. BaCl 2 + (NH 4 ) 2 HP0 4 = BaHPO 4 + 2NH 4 C1. 5. (NH 4 ) 2 C 2 O 4 causes a white precipitate of barium oxalate BaC 2 O 4 , slightly soluble in acetic acid. This precipitation will not take place in very dilute solutions. BaCl 2 + (NH t ) 2 C 2 O 4 r= BaC 2 O 4 + 2NH 4 C1. 6. A loop of platinum wire moistened with the solution colors the gas flame green when held in it. REACTIONS OF STRONTIUM (Sr). Use a solution of strontium nitrate (Sr(NO 3 ) 2 ). 1. H 2 SO 4 forms a white precipitate of strontium sulphate SrSO 4 , immediately, if the solution be strong, but not until after some time, if it be very dilute. 2. K 2 CrO 4 produces no precipitate in the presence of acetic acid, but if the solution be made alkaline with KOH, a yellow precipitate, strontium chromate SrCrO 4 , falls. 3. (NH 4 ) 2 CO 3 produces a white precipitate of strontium carbonate SrCO 3 , soluble in acetic and the stronger acids. Na 2 CO 3 produces the same precipitate. 4. (NH 4 ) 2 HPO 4 forms a white precipitate of strontium phosphate SrHPO 4 , soluble in acids. 5. (NH 4 ) 2 C 2 O 4 causes the precipitation of white strontium oxalate SrC 2 O 4 , sparingly soluble in acetic acid, but readily soluble in HC1. 6. Strontium salts impart an intense red to the colorless gas flame. BASES. 37 REACTIONS OF CALCIUM (Ca). Use a solution of calcium chloride (CaCl 2 ). 1. H 2 SO 4 , in moderately dilute solutions, forms a white precipitate of calcium sulphate CaSO 4 , soluble in excess of water. 2. (NH 4 ) 2 CO 3 or Na 2 CO 3 produces a white precipitate of calcium carbonate CaCO 3 , soluble in acids. This precipi- tation is not complete unless the solution is boiled. 3. (NH 4 ) 2 HPO 4 causes the precipitation of calcium phos- phate CaHPO 4 , soluble in acetic and the stronger acids. 4. (NH 4 ) 2 C 2 O 4 produces a white precipitate of calcium oxa- late CaC 2 O 4 , insoluble in acetic acid, soluble in hydrochloric or nitric acid. 5. The salts of calcium color the flame yellowish-red. REACTIONS OF MAGNESIUM (Mg). Use a solution of magnesium sulphate (MgSO 4 ). 1. (NH 4 ) 2 CO 3 forms a white precipitate of magnesium-am- monium carbonate MgCO 3 (NH 4 ) 2 CO 3 , soluble in NH 4 C1. By preceding the addition of the reagent by that of NH 4 C1, a much smaller quantity will suffice to keep the precipitate in solution than will be required to dissolve it after once formed. MgS0 4 + 2(NH 4 ) 2 C0 3 = MgC0 3 (NH 4 ) 2 C0 3 + (NH 4 ) 2 SO 4 . 2. KOH,NaOH or NH 4 OH produces a white precipitate of magnesium hydrate Mg(OH) 2 , soluble in NH 4 C1. 3. (NH 4 ) 2 HPO 4 with NH 4 C1 and NH 4 OH produces a white crystalline precipitate of ammonium-magnesium phosphate Mg(NH 4 )PO 4 , slightly soluble in water, but almost insoluble in water containing NH 4 OH. Violent agitation or stirring assists in the formation of this precipitate. (NH 4 ) 2 HAsO 4 under similar circumstances precipitates white Mg(NH 4 )AsO 4 . MgSO 4 + NH 4 OH + Na 2 HPO 4 = Mg(NH) 4 PO 4 + Na 2 SO 4 -f H 2 O. The ammonium chloride takes no part in the reaction, except to keep magnesium hydrate from precipitating. 4. Magnesium salts impart no color to the flame. 38 ANALYTICAL CHEMISTRY. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP II. Ba Sr Ca Mg H 2 S0 4 White Precipitate insoluble in acids White Precipitate insoluble in acids White Precipitate soluble in excess of H 2 O No Precipitate K 2 CrO 4 Yellow Precipitate insoluble in acetic acid No Precipitate unless alkaline No Precipitate No Precipitate (NH 4 ) 2 C0 3 White Precipitate White Precipitate White Precipitate White Precipitate soluble in NH 4 C1 (NH 4 ) 2 HP0 4 White Precipitate White Precipitate White Precipitate White Precipitate (NH 4 ) 2 C 2 4 White Precipitate in strong solution White Precipitate in strong solution White Precipitate in dilute solution No Precipitate unless concentrated NH 4 OH- No Precipitate No Precipitate No Precipitate White Precipitate DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP II. Add NH 4 C1, NH 4 OH and (NH 4 ) 2 CO 3 , boil and filter. Ppt Ba, Sr, Ca. Wash, dissolve in HC 2 H 3 O 2 , add K 2 CrO 4 , filter. Filt. Mg. Add (NH 4 ) 2 HPO 4 , agitate. White ppt. if Mg be present. Ppt. Ba. Yellow. Filtrate Sr, Ca. Add very dilute H 2 SO 4 , allow to stand 10 minutes, filter. Ppt. Sr. Confirm by flame test. Filt. Ca. Add NH 4 OH and (NH 4 ) 2 C 2 4 , white ppt. In order to thoroughly acquaint the student with the method of analysis by this and subsequent charts, the following expla- nation is given, in the belief that a careful study of it, until perfectly understood, will enable the student to follow all the charts which come after, and which are simply enlargements of this scheme. To a small quantity of the solution in a test tube add an equal volume of solution of NH 4 C1, close and invert the tube, so as to thoroughly mix the contents, add NH 4 OH until, after mixing, the solution smells distinctly of it, and then add (NH 4 ) 2 CO 3 , boil and filter. We suppose all four of the bases under consideration to be present until their absence is proven ; so by this treatment we divide them into two groups. The precipitate consists of Ba, Sr and Ca, and the filtrate of Mg. A little more (NH 4 ) 2 CO 3 BASES. 39 should be added to this filtrate and the whole again boiled, to make sure that all the insoluble carbonates have been precipitated. If this is found to be the case, (NH 4 \HPO 4 is added to the filtrate, and Mg, if present, will form a white precipitate. The first precipitate, consisting of Ba, Sr and Ca, having, in the meantime, been washed by forcing a jet of water from the wash bottle on it, is dissolved in HC 2 H 3 O 2 , and to the solu- tion K 2 CrO 4 added ; this again divides the solution into a pre- cipitate and a filtrate. The yellow precipitate represents Ba, while the filtrate contains the Sr and Ca. To this filtrate very dilute H 2 SO 4 (made by adding a small quantity of H 2 SO 4 to about ten times its volume of water) is added, and the mix- ture allowed to stand ten minutes, for the SrSO 4 to form and subside, then filtered, which again gives us a precipitate, repre- senting Sr, and a filtrate indicating, after the addition of NH 4 OH and (NH 4 \C 2 O 4 , the presence or absence of Ca. DIRECTIONS FOR THE ANALYSIS OF A SOLUTION CONTAINING SOLUBLE SALTS OF ALL THE PRECEDING ELEMENTS. Add NH 4 C1, NH 4 OH and (NH 4 ) 2 CO 3 , boil and filter. Ppt. Ba, Sr, Ca. Wash, dissolve in HC,H,O 2 , add K 2 CrO 4 , filter. Filtrate Mg, K, Na, Li, NH 4 . Add (NH 4 ) 2 HPO 4 , agitate, filter. SJ: Filtrate K, Na, Li, NH 4 . Evaporate to dryness, ignite, dissolve in a small quantity of H,O, add Na 2 HP0 4 , boil, filter. Ppt. Ba. Yellow. Filtrate Sr, Ca. Add very dilute H 2 SO 4 , allow to stand, filter. Ppt. Sr. Confirm by flame test. Filt. Ca. Add NH 4 OH and (NH 4 ) 2 C 2 4 , white ppt. * Confirm by flame test. Filtrate K, Na, NH 4 . Concentrate, add HC1 and PtCl 4 yellow ppt. K. Test for Na and NH 4 in original solution. PRECAUTIONS TO BE OBSERVED IN THE PRECEDING CHARTS. Ammonium chloride must be added in excess, in order to keep the magnesium salts in solution when the hydrate and carbonate are added. Ammonium hydrate is added until the liquid smells of it. Ammonium carbonate is added as long as a precipitate is produced. In order to determine this to a certainty, a portion of the filtrate is tested with a little more of the reagent, when, if no precipitate occurs, the analysis may be proceeded with. This precaution of applying more of 40 ANALYTICAL CHEMISTRY. the reagent to a portion of the filtrate, to prove the complete precipitation, should be exercised in every case, as it is im- portant to add just sufficient of the reagent to accomplish the object, but always to avoid a large excess. The analysis of the above solutions may be much simplified in many cases by adding a solution of CaSO 4 to the original solution. If a precipitate form immediately, Ba is present, Sr and Ca may be. If a precipitate form after some time, Ba is absent, Sr is present, and Ca may be present. If no pre- cipitate be formed, Ba and Sr are absent, and Ca may be tested for in another portion with (NH 4 ) 2 C 2 O 4 . The above charts suppose all the elements to be present, in which case they afford the simplest means of detecting them. GROUP III. MANGANESE, ZINC, COBALT, NICKEL. REACTIONS OF MANGANESE (Mn). Use a solution ofmanganous sulphate (MnSO 4 ). 1. NH 4 HS, in neutral or alkaline solution, precipitates the flesh-colored manganous sulphide MnS, which, on exposure to air, becomes brown. HC1, HNO 3 and HC 2 H 3 O 2 dissolve this precipitate, but it is insoluble in alkalies. NH 4 C1 facili- tates the separation of the precipitate, while the salts of the organic acids and excess of NH 4 OH prevent it. 2. KOH or NaOH produces a whitish precipitate of man- ganous hydrate Mn(OH) 2 , insoluble in excess. MnSO 4 -j- 2KOH =Mn(OH; 2 + K 2 SO 4 . 3. NH 4 OH likewise precipitates Mn(OH) 2 , first white, but becoming rapidly brown, partly soluble in excess. This pre- cipitation is prevented by the previous addition of NH 4 C1. 4. (NH 4 ) 2 CO 3 produces a white precipitate of manganous carbonate MnCO 3 , insoluble in excess. 5. HNO 3 and red oxide of lead will, on heating and allow- ing the precipitate to subside, impart to the supernatant liquid a red color, due to permanganic acid H 2 Mn 2 O 8 . Hydro- chloric acid and chlorides interfere with this reaction. This is known as Crum's process for detecting manganese. BASES. 41 6. A fragment of a manganese salt fused on platinum foil with K 2 CO 3 and KNO 3 will form a green mass containing potassium manganate K 2 MnO 4 . /. A borax bead (formed by fusing on the loop of a plati- num wire some borax until it becomes a clear glass) with manganese, in the oxidizing blowpipe flame, becomes violet while hot, and a fine amethyst color on cooling. REACTIONS OF ZINC (Zn). Use a solution of zinc sulphate (ZnSO 4 ). 1. NH 4 HS produces a white precipitate of zinc sulphide ZnS, insoluble in acetic acid, readily soluble in dilute hydro- chloric acid. 2. KOH, NaOH and NH 4 OH give white precipitates of zinc hydrate Zn(OH) 2 , readily soluble in excess, forming zincates as Zn(OK) 2 . Zn(OH) 2 is again precipitated on boil- ing. 3. (NH 4 ) 2 CO 3 forms a white precipitate of basic zinc car- bonate (ZnCO 3 ) 2 (Zn(OH) 2 ) 3 , readily soluble in excess. 5(NH 4 ) 2 C0 3 -f 5 ZnS0 4 + 3 H 2 O = 2ZnC0 33 Zn(OH) 2 + 5 (NH 4 ) 2 SO 4 + 3 CO 2 . 4. K 2 CO 3 or Na 2 CO 3 produces a similar precipitate, insoluble in excess. 5. On charcoal, before the blowpipe, metallic zinc volatilizes and burns, forming an incrustation of oxide, which is yellow while hot, becoming white on cooling ; if this coating be moistened with a drop of cobaltous nitrate, and again heated in the outer flame, it becomes green. REACTIONS OF COBALT (Co). Use a solution of cobaltous nitrate (Co (NO 3 ) 2 ). i. NH 4 HS produces a black precipitate of cobaltous sul- phide CoS, insoluble in acetic acid and cold dilute hydro- chloric acid. The precipitation is promoted by the presence of NH 4 C1. 2NH 4 HS + Co(NO 3 ) 2 = CoS + 2NH 4 NO 8 -f H 2 S. 42 ANALYTICAL CHEMISTRY. 2. KOH or NaOH produces a blue precipitate of cobaltous hydrate Co(OH) 2 , insoluble in excess, and becoming pink on boiling or exposure to air. 3. NH 4 OH causes a similar precipitate of Co(OH) 2 , soluble in excess with a red color. Sugar and some other organic compounds prevent the precipitations by the alkalies. The alkaline carbonates behave like their respective hydrates. 4. KCN gives a red-brown precipitate of cobaltous cyanide Co(CN) 2 , soluble in excess and reprecipitated by HC1 ; if, however, the solution be boiled with only a few drops of HC1, the cobaltous cyanide will not be precipitated on the further addition of HC1, on account of the formation of potassium cobalti-cyanide K 6 Co 2 (CN) 12 . This experiment should be per- formed in a fume closet, in order to avoid inhaling the fumes of hydrocyanic acid. 5. Salts of cobalt color the borax bead blue before the blowpipe. REACTIONS OF NICKEL (Ni). Use a solution of nickelous sulphate (NiSO 4 ). 1. NH 4 HS forms a black precipitate of nickelous sulphide NiS, insoluble in acetic acid and cold dilute hydrochloric acid. The precipitation is promoted by the presence of NH 4 C1. 2. KOH or NaOH produces a green precipitate of nickel- ous hydrate Ni(OH) 2 , insoluble in excess. 3. NH 4 OH gives a similar precipitate, soluble in excess, with a blue color. Sugar and some other organic compounds pre- vent the precipitation by the alkalies. The alkaline carbon- ates behave like their respective hydrates. 4. KCN produces a yellowish-green precipitate of nickel- ous cyanide Ni(CN) 2 , soluble in excess, and re-precipitated by HC1 even after boiling, also precipitated, after adding HCl and boiling, by KOH. This is used as a method of dis- tinguishing and separating nickel and cobalt, the latter not precipitating under these circumstances with KOH. This experiment should be performed in a fume closet. BASES. 43 5. The salts of nickel color the borax bead violet while hot, and reddish-brown when cold. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP III. Mn Zn Co Ni NH 4 HS Flesh colored precipitate. White precipitate Black precipitate Black precipitate KOH White precipitate insoluble in excess White precipitate soluble in excess Blue precipitate insoluble in excess Green precipitate insoluble in excess NH 4 OH White precipitate soluble in excess White precipitate soluble in excess Blue precipitate soluble in excess Green precipitate soluble in excess (NH 4 ) 2 C0 3 White precipitate insoluble in excess White precipitate soluble in excess Blue precipitate soluble in excess Green precipitate soluble in excess DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP III. Mn, Zn, Co, Ni. Add NH 4 OH, then acidify slightly with HC 2 H 3 O 2 and add H 2 S until the liquid smells strongly ; falter. Precipitate Zn, Co, Ni. Wash, dissolve in HC1 and HNO 3 , evaporate excess of acid, add KOH in excess, filter. Precip. Co, Ni. Wash, dissolve in HC1, add KCN in excess and boil with HC1 until all odor of HCN disappears, then add KOH in excess, filter. Precip. Ni. Green. Filt. Co. Evaporate to dryness and test with borax bead. Filt. Zn. Add NH 4 HS white i pt. Filtrate Mn. Add NH 4 OH and NH 4 HS, ppt. pink. 44 ANALYTICAL CHEMISTRY. ^ ^^ W d en .H fr 5 S"ScJ =< >- ^ _o ^ u hrJ'Ml CO . vpL, '' -g, 1 ^ 5 2 W V3 S"" IS ^ c Z ^ ^ ^_J Or, ,? ^ H 3'SP " H S us ^ SK"o ' J 1 3"fiS 'o ,-J ^ ' T) c^T & 5? 4J. <* ^ J i >% Xi -jj te ^ s" "o " ; ~ cn -w -^ o H rt" rt 0.1^ ^ S Is S o -^ X rt id PH O rt 5-5 cn M ^ rt ^| 4) c/l fc ^ . nS o g3 c3| > ii a; 1 ' c .gs^. Si < 5? >, SffiK'a, g 3 c > /5 5 ~ Q, :i* g 4 'o rt ffi* fe ^1 & |g ^ (9J25 ~ o cn " cn" P3 ~- a IS W W K 1^ * t .ffi " fi u SB H 'Is g B*3 fa _ u^ 3 TD rt ^ 42 2^^ < ^ c fc " uffi jj c o O HS^ 5 d uf a d FOR THE DETECT 15 Ppt. Zn, Co, Ni. n HCI and HNO 3 , evapc Add KOH in excess, fill 'pt. Co, Ni. i HCI, add KCN in exce !1 until all odor of HCN -ars, add KOH in xcess, filter. C . o A c! sent it may remain in sol cn "v 1*4.3 VJ Q, W w g fc -3 j> > c3 p< "e '3-S.S2 .2 1 1- 'S IjT y 2 d Pi f^ ^Q 42 3 rt rt > * BASES. 45 GROUP IV. IRON, CERIUM, ALUMINIUM, CHROMIUM. REACTIONS OF IRON in ferrous state (Fe) 11 . Use a solution of ferrous sulphate (FeSO 4 ). 1. K 4 Fe(CN) 6 /in a neutral or acid solution, gives a white (rapidly changing to light blue) precipitate of potassium ferrous ferrocyanide K 2 Fe 2 (CN) 6 , also known as Everett's salt. Alkalies decompose this precipitate, forming ferrous hy- drate and a ferrocyanide of the base used. K 4 Fe(CN) 6 + FeS0 4 = K 2 Fe 2 (CN) 6 + K 2 SO 4 . 2. K 6 Fe 2 (CN) 12 , in a neutral or slightly acid solution, forms a dark blue precipitate of ferrous ferricyanide Fe 3 Fe 2 (CN) 12 , known as Turnbull's blue. If the solutions be very dilute there is produced merely a deep blue-green coloration. K 6 Fe 2 (CN) 12 + 3 FeS0 4 = Fe 3 Fe 2 (CN) 12 + 3 K 2 SO 4 . 3. KCNS produces no change. 4. H 2 S in acid solution does not form a precipitate. 5. NH 4 HS, with a neutral or alkaline solution, forms a black precipitate of ferrous sulphide, soluble in HC1 or HNO 3 . NH 4 C1 promotes the formation of this precipitate. 6. NH 4 OH in the absence of NH 4 C1 produces a dirty green precipitate of ferrous hydrate Fe(OH) 2 . This precipitate rapidly becomes reddish-brown, owing to absorption of oxygen. 7. KOH produces a dirty green precipitate of ferrous hy- drate Fe(OH) 2 , similar to that produced by ammonia. Non- volatile organic substances, as sugar and some acids, retard the precipitations by NH 4 OH and KOH. 8. Na 2 CO 3 causes a white precipitate of ferrous carbonate FeCO 3> which rapidly becomes brown, from absorption of oxygen. This rapid oxidation is prevented by the use of dis- tilled water and sugar. REACTIONS OF IRON in ferric state (Fe 2 ) VI . Use a solution of ferric chloride (Fe 2 Cl 6 ). i. K 4 Fe(CN) 6 , in a neutral or acid solution, produces a dark blue precipitate of ferric ferrocyanide (Fe 2 ) 2 (Fe(CN) 6 ) 3 , de- composed by alkalies. 3 K 4 Fe(CN) 6 + 2Fe 2 Cl 6 i= (Fe 2 ) 2 (Fe(CN) 6 ) 3 + 12 KC1. 46 ANALYTICAL CHEMISTRY. 2. K 6 Fe 2 (CN) 12 forms no precipitate, but produces a deep reddish-brown color. The olive-green color sometimes pro- duced in this reaction is due to traces of ferrous salt. 3. KCNS imparts to acid solutions a deep blood-red color, due to the formation of ferric sulphocyanate. This color is immediately destroyed by HgG 2 . Dilute solutions show this reaction best. 4. H 2 S forms in acid solutions a white turbidity due to separation of sulphur ; the ferric salt being at the same time reduced to the ferrous condition. 2Fe 2 Cl 6 + 2H 2 S == 4FeCl 2 + 4HC1 + S 2 . In alkaline solutions this reagent acts as an alkaline sulphide. 5. NH 4 HS causes a black precipitate of ferrous sulphide FeS, sulphur separating at the same time. 6. NH 4 OH precipitates reddish-brown ferric hydrate Fe 2 (OH) 6 ; non-volatile organic acids and sugar prevent this precipitation. 7. KOH and NaOH react like NH 4 OH. 8. Na 2 CO 3 and the other alkaline carbonates precipitate ferric hydrate Fe,(OH) 6 . 9. With borax in the oxidizing blowpipe flame, ferrous and ferric compounds give dark yellow to red colored beads while hot, and yellow when cold. In the reducing flame the beads change to bottle-green. REACTIONS OF CERIUM (Ce). Use a solution of cerous chloride (Ce 2 Cl 6 ). 1. NH 4 HS causes a white precipitate of cerous hydrate Ce 2 (OH) 6 . 2. NH 4 OH produces the same white precipitate of Ce 2 (OH) 6 , as also do KOH and NaOH. 3. (NH 4 ) 2 C 2 O 4 or H 2 C 2 O 4 forms a white precipitate of cerous oxalate Ce 2 (C 2 O 4 ) 3 . Organic matter does not interfere with the formation of this precipitate. 4. With a borax bead before the blowpipe the salts of cerium behave like those of iron. BASES. 47 REACTIONS OF 'ALUMINIUM (Al). Use a solution of alum (K 2 SO4A1 2 (SO 4 )3). 1. NH 4 HS produces a white precipitate of aluminium hy- drate A1 2 (OH) 6 , while H 2 S escapes. 6NH 4 HS + K 2 S0 4 A1 2 (S0 4 ) 3 + 6H 2 O = A1 2 (OH) 6 + 3 (NH 4 ) 2 SO 4 -f K 2 SO 4 + 6H 2 S. 2. NH 4 OH forms a white precipitate of aluminium hydrate A1 2 (OH) 6 , insoluble in excess ; this is an important distinc- tion from zinc. 3. KOH and NaOH produce a similar white precipitate of A1 2 (OH) 6 , soluble in excess, forming aluminates of the base, as A1 2 (OK) 6 . This is not reprecipitated by boiling (distinction from chromium). 4. Na 2 CO 3 and the other alkaline carbonates precipitate white gelatinous aluminium hydrate A1 2 (OH) 6 , insoluble in excess, CO 2 escaping at the same time. 3 Na 2 C0 3 + K 2 S0 4 A1 2 (S0 4 ) 3 -f 3H 2 O = A1 2 (OH) 6 + 3Na 2 SO 4 + K 2 S0 4 -j- 3 C0 2 . The presence of non-volatile organic acids and sugar prevent the complete precipitation in the above reactions. REACTIONS OF CHROMIUM (Cr). Use a solution of chromic chloride (Cr 2 Cl 6 ). 1. NH 4 HS produces a greenish precipitate of chromic hy- drate Cr 2 (OH) 6 , H 2 S escaping. 2. NH 4 OH forms the same greenish precipitate of Cr 2 (OH) 6 , insoluble in excess. 3. KOH and NaOH produce the same precipitate soluble in excess, but reprecipitated on boiling (distinction from alu- minium). 4. Na 2 CO 3 and the other alkaline carbonates, precipitate green basic carbonates. 5. KNO 3 and K 2 CO 3 fused with chromium compounds be- come yellow from formation of potassium chromate K 2 CrO 4 . AgNO 3 and Pb(C 2 H 3 O 2 ) 2 are important tests for this compound ; the former precipitates red silver chromate, the latter yellow lead chromate. 6. With the borax bead in the inner blowpipe flame, chro- mium compounds give a green color. 48 ANALYTICAL CHEMISTRY. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP IV. Fe(ous) Fe(ic) Ce Al Cr K 4 Fe(CN) 6 White ppt. turning blue. Deep blue precipitate. K 6 Fe 2 (CN) 12 Deep blue precipitate. No ppt. brownish-red color. KCNS No change. Blood-red color. NH 4 HS Black ppt. Bk ck ppt. White ppt. White ppt. Greenish ppt. NH 4 OH Dirty green ppt. Reddish- brown ppt. White ppt. White ppt. in- soluble in excess. Greenish ppt. in- soluble in excess. KOH Dirty green ppt. Reddish- brown ppt. White ppt. White ppt. solu- ble in excess, not reprecipitated by boiling. Green ppt. soluble in excess, reprecipitated by boiling. Na 2 C0 3 " White ppt. be- coming dark. Reddish- brown ppt. White ppt. White ppt. Green ppt. DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP IV. Evaporate a portion of the solution to dryness, fuse on platinum foil with Na 2 CO 3 and KNO 3 , boil with water and filter. Residue, Fe, Ce. Dissolve in concentrated H 2 SO 4 and C 2 H 5 OH, divide in two parts. Filtrate, Al, Cr. Yellow if Cr be present, divide in two parts. Fe. Dilute, test with K 4 Fe(CN) 6 . Test original solution with K 4 Fe(CN) 6 and K 6 Fe 2 (CN) 12 for Fe(ous) and Fe(ic). Ce. Dilute, add H 3 C 6 H 6 7 , NH 4 OH and Na 2 HP0 4 , white ppt. Al. Add NH 4 C1 and warm, white gelatinous ppt. Cr. Add HC,H,O 2 and Pb(C 2 H 3 2 ) 2 yellow ppt. BASES. 49 4 ll SB.:? U ^ s ^ 6, ^ ^| a offi^i Filt. M ff excess o HC 2 H 3 2 Filt. K, Na, Li, NH 4 . Evaporate to dryness ignite, dissolve in ff o STJJ5 rt< E fc 1 c SB -a 3 . O &* 50 ANALYTICAL CHEMISTRY. PRECAUTIONS AND OBSERVATIONS ON THE PRECEDING CHART. 1. It is particularly desirable not to filter immediately after adding NH 4 HS in the first precipitation ; otherwise, the Mn will not be thoroughly precipitated. 2. When the first precipitate is dissolved in HC1 and HNO 3 , great care must be exercised to thoroughly oxidize the Fe by boiling with the HNO 3 , otherwise, the precipitation by NH 4 OH will not be complete. 3. The precipitate of aluminium hydrate is difficult to observe, as it floats in the solution instead of falling to the bottom. Warming the solution will usually render it visible. 4. When phosphoric acid is present, the members of Group II precipitate with Group IV, and require an entirely different method of separation. It is desirable, however, for the student first to familiarize himself with these simpler soluble salts, and undertake the more difficult cases of salts insoluble in water after the acids have been considered. See page 83. GROUP V. ARSENIC, ANTIMONY, TIN, GOLD, PLATINUM. REACTIONS OF ARSENIC (As). (a) ARSENIOUS COMPOUNDS. Use a solution of As 2 O 3 , in water. 1. H 2 S passed into the solution produces a yellow color, but no precipitate until HC1 is added, when a yellow precipi- tate of arsenious sulphide As 2 S 3 , falls. This precipitate is insoluble in strong HC1, but soluble in NH 4 HS, NH 4 OH and (NH 4 ) 2 C0 3 . As 2 3 + 3 H 2 S = As 2 S 3 -f 3 H 2 0. 2. NH 4 HS causes the formation of arsenious sulphide, which remains in solution as ammonium sulpharsenite (NH 4 ) 3 AsS 3 . On the addition of HC1 arsenious sulphide is precipitated. 3. AgNO 3 produces no precipitate until a few drops of dilute ammonia solution are added, when a yellow precipitate of silver arsenite Ag 3 AsO 3 , falls, soluble in HNO 3 and in NH 4 OH. BASES. 51 4. CuSO 4 , under similar circumstances, produces a yellowish- green precipitate of cupric arsenite CuHAsO 3 . (&) ARSENIC COMPOUNDS. Use a solution of sodium arsenate (Na 2 HAsO 4 ). 1. H 2 S causes, in acid solution only, a yellow precipitate of arseniou ssulphide As 2 S 3 , mixed with sulphur. This re- action takes place slowly, but is accelerated by heat. 5H 2 S 4- 2Na 2 HAsO 4 + 4HC1 = As 2 S 3 + S 2 -f- 4NaCl 4- 8H 2 O. 2. NH 4 HS produces no precipitate, but forms arsenic sulphide As 2 S 5 , which remains in solution as ammonium sulpharsenate (NH 4 ) 3 AsS 4 . Upon the addition of HC1 As 2 S 5 is precipitated, and not As 2 S 3 and S. 3. AgNO 3 , with a small amount NH 4 OH, produces a choco- late-colored precipitate of silver arsenate Ag 3 AsO 4 , soluble in HNO 3 and NH 4 OH. 4. CuSO 4 , under similar circumstances, forms a bluish-green precipitate of cupric arsenate CuHAsO 4 . The following tests are applicable to both arsenious and arsenic compounds. 1. Mars/is Test. Generate hydrogen in the usual way, allowing it to escape through a glass tube drawn out at the end so as form a small orifice (Fig. 4). In very exact cases, the gas should be dried by passing over calcium chloride. When all the air has been expelled (which should be deter- mined by collecting a small test tube full and holding its mouth to a flame; if the gas burn quietly, without explosion, it is pure), ignite the escaping gas ; it should burn with a colorless or yellow flame ; in the latter case it is due to the sodium in the glass. A piece of cold porcelain a small cru- cible lid is best is pressed down on the flame ; there should be no deposit on it. Add now, through the funnel tube, a solution of arsenic, washing it down with a little water. The flame will become of a pale blue color, due to the formation of hydrogen arsenide H 3 As. On bringing the crucible lid into the flame now, a blackish-brown deposit, with metallic lustre, will form on it. This deposit is readily soluble in a solution of sodium or calcium hypochlorite. 2. ReinscJCs Test. Boil some strips of copper with dilute 52 ANALYTICAL CHEMISTRY. HC1 ; if no discoloration of the copper takes place, the arsenic solution may be added. The copper immediately becomes coated with an iron-gray metallic film. Pour off the liquid, dry the copper by holding it in the lamp with the fingers so it may not become too hot, place in a clean, dry, narrow test tube, and heat gently, when a white ring of As 2 O 3 will form on the tube above the copper, readily distinguished by the characteristic octahedral shape of the crystals. Fleitmarfs Test. Generate hydrogen in a test tube with FIG. 4. zinc and solution of potassium hydrate ; moisten a piece of filter paper with one drop of solution of silver nitrate, place it over the mouth of the tube and heat ; there should be no coloration of the spot on the paper. Now add some com- pound of arsenic ; the silver nitrate will immediately become black, owing to production of metallic silver. H 3 As + (AgN0 3 ) 6 + (H 2 0) 3 = H 3 As0 3 + (HNO 3 ) 6 + (Ag 2 ) 3 . Before the blowpipe, on charcoal, arsenic volatilizes, with the characteristic odor of garlic. BASES. 53 REACTIONS OF ANTIMONY (Sb). Use a solution of tartar emetic (KSbOC 4 H 4 O 6 ), acidified with HC1/ 1. H 2 S forms an orange precipitate of antimonous sul- phide Sb 2 S 3 , soluble in NH 4 HS, and in concentrated HC1, but insoluble in (NH 4 ) 2 CO 3 . 2. NH 4 HS produces an orange precipitate of antimonous sulphide, readily soluble in excess, forming ammonium sulph-antimonite (NH 4 ) 3 SbS 3 , from which HC1 again pre- cipitates Sb 2 S 3 . 3. KOH or NaOH precipitates white, bulky antimonous hydrate Sb(OH) 3 , soluble in excess. 4. NH 4 OH precipitates the same compound insoluble in excess. 5. Marsh's Test gives the same result as with arsenic ; the black spot, however, is insoluble in sodium or calcium hypo- chlorite solution, but soluble in NH 4 HS. 6. Reinsch's Test causes a deposit on copper, as with arsenic, but when heated in a tube there is formed a white amorphous ring, which is readily distinguished from the crystalline one of arsenic. 7. Fleitman's Test gives no result with antimony com- pounds. 8. On charcoal, with Na 2 CO 3 , before the blowpipe, a metallic globule of antimony is produced, while characteristic fumes of the oxide are given off. REACTIONS OF TIN (Sn) v (a) STANNOUS COMPOUNDS. Use a solution of stannous chloride (SnCl 2 ). 1. H 2 S precipitates dark brown stannous sulphide SnS, soluble in concentrated HC1 and in (NH 4 ) 2 S, insoluble in NH 4 HS. 2. KOH or NaOH precipitates white stannous hydrate Sn(OH) 2 , soluble in excess ; on boiling this solution SnO pre- cipitates. 3. NH 4 OH precipitates the same compound, insoluble in excess. 54 ANALYTICAL CHEMISTRY. 4. HgCl 2 causes a white precipitate of Hg 2 Q 2 , converting the SnCl 2 into SnCl 4 . 2HgCl 2 + SnCl 2 = Hg 2 Cl 2 + SnQ 4 . This precipitate blackens on the addition of NH 4 OH. (b) STANNIC COMPOUNDS. Use" a solution of stannic chloride (SnCl 4 ). 1. H 2 S produces a yellow precipitate of stannic sulphide SnS 2 , soluble in NH 4 HS and in concentrated HC1. 2. KOH or NaOH precipitates white stannic acid H 2 SnO 3 , soluble in excess ; on boiling no reprecipitation takes place distinction from stannous salts. 3. NH 4 OH produces the same precipitate, insoluble in excess. 4. Heated on charcoal, before the blowpipe, with Na 2 CO 3 , metallic tin is formed, with the production of a white incrust- ation of the oxide. REACTIONS OF GOLD (Au). Use a solution of auric chloride (AuCl 3 ). 1. H 2 S precipitates black auric sulphide Au 2 S 3 , insoluble in HC1, soluble in (NH 4 ) 2 S. 2. H 2 C 2 O 4 or FeSO 4 precipitates metallic gold as a finely divided brown powder. 3. SnCl 2 mixed with SnCl 4 (prepared by SnCl 2 and chlo- rine water) produces a purple-red precipitate or coloration (Purple of Cassius), consisting of the mixed oxides of gold and tin. 4. Heated on charcoal, before the blowpipe, metallic gold is produced. REACTIONS OF PLATINUM (Pt). Use a solution of platinic chloride (PtCl 4 ). 1. H 2 S causes a brown precipitate of platinic sulphide PtS 2 , insoluble in HC1, soluble in (NH 4 ) 2 S. 2. KC1 produces a yellow crystalline precipitate of potas- sium platinic chloride K 2 PtCl 6 . 3. NH 4 C1 forms a similar precipitate of (NH 4 ) 2 PtCl 6 . The BASES. 55 precipitation in both cases is facilitated by the addition of alcohol. 4. Zn, Fe and some other metals precipitate metallic platinum. 5. Heating on charcoal, before the blowpipe, produces the metal. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP V. As Sb Sn Au Pt H 2 S Yellow ppt. soluble in NH 4 HS. Orange-red ppt. soluble in NH^HS. Brown or Yellow ppt. soluble in (NH 4 ) 2 S. Black ppt. soluble in (NH 4 ) 2 S. Brown ppt. soluble in (NH 4 ) 2 S KOH No change. White ppt. soluble in excess. White ppt. soluble in excess. No change. With excess of HC1 yellow ppt. NH 4 OH No change. White ppt. insoluble in excess. White insoluble in excess. Red ppt. fulminating gold. With excess of HCt yellow ppt. Marsh's Test Black spot with metallic lustre, soluble in Ca(OCl) 2 Black, sooty spot, soluble in (NHJHS, insoluble in Ca(OCl) 2 . Reinsch's Test White sublimate of octahedral crystals. White amorphous sublimate. Fleitman's Test Dark spot of metallic silver. No change. DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP V. Add HC1 and H 2 S, collect the precipitate, transfer to a dish, boil with concentrated HC1 ; filter. Ppt. As, Au, Pt. If yellow, As only is present ; if dark, digest with (NH 4 ) 2 CO 3 , filter. wash, Ppt. Au, Pt. Dissolve in HC1 and HNO 3 . Divide in two parts. Filt. As. Acidify with HC1, yellow ppt. Au. Add SnCl 2 purple. Confirm by testing original solution. Pt. Add KC1, yellow ppt. Confirm by testing original solution. Filt. Sb, Sn. Dilute with H 2 O, boil in a dish with a strip of platinum foil, and a small piece of zinc so the metals touch. The platinum will be coated with black Sb, while the Sn will be deposited as a black sediment. This black sediment is dissolved in strong HC1 and tested with HgCl 2 , white ppt. if Sn be present. Treat the foil with a few drops of HNO 3 , dissolve in solution H 2 C 4 H 4 O 6 , add H 2 S, orange ppt. if Sb be present. 56 ANALYTICAL CHEMISTRY. < Q H w 5? rt O W ^ Q a w w ^ ffi w 2 M H H Cr, Mn, Zn, Co, Ni, , after warming, the , rt rt fcv * Al, Cr, Mn, Zn, Co, OH until after agita Co, Ni 4 OH i N 1 ^ -T3 e-g s JIJ. ! *S * U rt 73 -T3 o rt o 5 cO C ffi C/3 E 2; O ^ .- S ^_faf o^O UK^ N"-^ -$8 C * U ft S -d ojl ^ "2 '" ' 51 o. BASES. 57 PRECAUTIONS AND OBSERVATIONS ON THE PRECEDING CHART. A great deal of time may be saved by carefully noting the color of the precipitate produced by H 2 S. If stannic salts are absent and the precipitate yellow, As only is present; if orange, Sb is present and As may be ; when such an orange precipitate is obtained and Sn(ic) is absent, the readiest method of separation is to wash, and add to the precipitate (NH 4 ) 2 CO 3 . As will be dissolved, and may be detected in the filtrate by adding HC1, while the Sb remains on the filter. When the precipitate is dark, Au and Pt should be sought for in the original solution, as well as separated by the chart Only a small piece of zinc is necessary to effect the separa- tion of Sn and Sb. GROUP VI. MERCURY (1C), BISMUTH, COPPER, CADMIUM. REACTIONS OF MERCURY as mercuric salt (Hg(ic)). Use a solution of mercuric chloride (HgCl 2 ). 1. H 2 S or NH 4 HS produces, when in small proportion, a whitish precipitate of (HgS) 2 HgCl 2 ; a further addition of the reagent, together with the application of heat, causes the formation of a black precipitate of mercuric sulphide HgS, insoluble in either HC1 or HNO 3 , but soluble in a mixture of the two. 2. KOH or NaOH produces a yellow precipitate of mer- curic oxide HgO ; unless the reagent be in excess, a brown precipitate of a basic salt is formed. 2KOH -f- HgCl 2 = HgO + 2KC1 + H 2 O. 3 NH 4 OH precipitates white mercur-ammonium chloride NH 2 HgCl. 2NH 4 OH + HgCl 2 = NH 2 HgCl + NH 4 C1 -f 2H 2 O. This precipitate is readily soluble in HC1 and in HC 2 H 3 O 2 . 4. K 2 CrO 4 produces a red precipitate of mercuric chro- mate HgCrO 4 . 5. KI precipitates mercuric iodide HgI 2 , first yellow, but rapidly becoming scarlet. This precipitate is readily soluble in excess of KI or HgCl 2 . 6. SnCl 2 , in small quantity, in the presence of HC1, pre- 58 ANALYTICAL CHEMISTRY. cipitates mercurous chloride Hg 2 Cl 2 . On the addition of a larger quantity of the reagent, the mercurous chloride is reduced to the metal, which may be collected into a globule. 7. Na 2 CO 3 produces a reddish-brown precipitate of basic carbonate HgCO 3 (HgO) 3 . 8. Before the blowpipe, HgO breaks up into Hg and O. HgS, under similar circumstances, sublimes unchanged. REACTIONS OF BISMUTH (Bi). Use a solution of bismuth nitrate (Bi(NO 3 ) 3 ). 1. H 2 S or NH 4 HS produces a black precipitate of bismuth trisulphide Bi 2 S 3 , insoluble in dilute acids and alkalies, soluble in boiling HNO 3 . 2. KOH, NaOH or NH 4 OH forms a white precipitate of bismuth hydrate Bi(OH) 3 , converted by boiling into the yellow oxide Bi 2 O 3 . 3. K 2 CrO 4 precipitates yellow bismuth chromate Bi 2 (Cr0 4 ) 3 . 4. KI forms a brown precipitate of bismuth iodide BiI 3 , soluble in excess of the reagent. 5. Na 2 CO 3 precipitates white oxycarbonate of bismuth (BiO) 2 C0 3 .H 2 0. 3Na 2 CO 3 + 2 Bi(NO 3 ) 3 -f H 2 O = (BiO) 2 CO 3 H 2 O + 6NaNO 3 + 2CO 2 . 6. H 2 O in excess, when there is not an excess of free acid, precipitates bismuth subnitrate BiONO 3 .H 2 O. Bi(NO 3 ) 3 -f 2H 2 O = BiONO 3 .H 2 O + 2HNO 3 . When the chloride is so diluted the oxychloride BiOCl, separates. 7. Bismuth on charcoal, before the blowpipe, forms a hard bead of metal, with a characteristic incrustation of oxide deep orange-yellow while hot, pale when cold. REACTIONS OF COPPER (Cu). Use a solution of cupric sulphate (CuSO 4 ). i. H 2 S and NH 4 HS precipitate black cupric sulphide CuS, insoluble in dilute acids and alkalies, slightly soluble in NH 4 HS, and entirely dissolved by boiling HNO 3 . This pre- cipitation is prevented by KCN. BASES. 59 2. KOH or NaOH produces a light blue precipitate of cupric hydrate Cu(OH) 2 , insoluble in excess, and converted by boiling into black cupric oxyhydrate (CuO) 2 Cu(OH) 2 . In the presence of non-volatile organic acids this precipitation does not take place, but a blue color results. 2KOH + CuSO 4 = Cu(OH) 2 -f K 2 SO 4 . 3Cu(OH) 2 = (CuO) 2 Cu(OH) 2 + 2H 2 O. 3. NH 4 OH in small quantity forms a greenish-blue precipi- tate, readily soluble in excess, forming tetra-ammonio-cupric sulphate (NH 3 ) 4 CuSO 4 .H 2 O. 4NH 4 OH + CuSO 4 = (NH 3 ) 4 CuSO 4 .H 2 O + sH 2 O. 4. K 4 Fe(CN) 6 precipitates reddish-brown cupric ferro- cyanide Cu 2 Fe(CN) 6 . 5. Metallic Fe or Zn precipitates red metallic Cu. 6. In the outer blowpipe flame copper salts color the borax bead green while hot, blue when cold. In the inner flame, after moistening with SnCl 2 , it becomes red, owing to formation of Cu 2 O. REACTIONS OF CADMIUM (Cd). Use a solution of cadmium sulphate (CdSO 4 ). # 1. H 2 S or NH 4 HS precipitates yellow cadmium sulphide CdS, soluble in hot HNO 3 , but insoluble in NH 4 HS or KCN. 2. KOH or NaOH produces a white precipitate of cadmium hydrate Cd(OH) 2 , insoluble in excess. 3. NH 4 OH causes the same white precipitate of Cd(OH) 2 , soluble in excess. 4. Na 2 CO 3 produces a white precipitate of cadmium car- bonate CdCO 3 , insoluble in excess, but slightly soluble in ammonium salts, entirely soluble in NH 4 OH. 5. On charcoal, before the blowpipe, the salts of cadmium are reduced to metal and volatilize, forming a brownish in- crustation of oxide. 60 ANALYTICAL CHEMISTRY. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP VI. Hg(ic) Bi Cu Cd H 2 S or NH 4 HS Black ppt. insoluble in HNO 3 . Black ppt. Black ppt. soluble in KCN. Yellow ppt. insoluble in KCN. KOH Yellow ppt. White ppt. Blue ppt. White ppt. NH 4 OH White ppt. White ppt. Blue ppt. soluble in excess. White ppt. soluble in excess. Na 2 C0 3 Reddish-brown ppt. White ppt. Blue ppt. White ppt. DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP VI. Add HC1 and H 2 S, collect, boil with HNO 3 ; filter. Ppt. Hg. Black. Fill. Bi, Cu, Cd. Add NH 4 OH in excess ; filter. Ppt. Bi. White. Fill. Cu, Cd. Add H 2 S, collect and wash ; boil with dilute H a SO 4 ; filter. Ppt. Cu. Black. Filt. Cd. Dilute, add H 2 S, yellow ppt. BASES. 61 3 O *s 8 55 co O fc g3 ^ 2 p 5S w w [o ffi o g s? w J ^w ~ f-\ rt s <: e, fig ffi CO o . c u 33 " Ppt. Co Wash, dis KCN in e l with HCI add , bo HC ce , until odor of HCN removed, add KOH excess, filter. .- c/i o S-S S 5 ac'S-Q Filt. Al, Cr. Yellow if Cr be present. Divide i two parts. t! 4> .rt M, SS, 1 : ! J| si Sg . als Filt. Cu, Cd. Add H 2 S. Collect and wash, boil with dilute H 2 SO 4 , filter. 62 ANALYTICAL CHEMISTRY. PRECAUTIONS AND OBSERVATIONS ON THE PRECEDING CHART. 1. Unless excess of H 2 S be used, and the solution warmed, Hg will not be thoroughly precipitated. 2. When Sn, Au and Pt are present, the yellow ammonium sulphide (NH 4 ) 2 S must be used to dissolve Group V. When Cu is present some of it may be dissolved, so that any dark precipitate in that group should be tested for Cu. 3. When the precipitate insoluble in NH 4 HS is boiled with HNO 3 , a black mass of sulphur is sometimes obtained. This is readily distinguished from Hg, because it floats as one mass on the liquid, while the mercury collects at the bottom as a heavy, black precipitate. GROUP VII. SILVER, MERCURY(OUS), LEAD. REACTIONS OF SILVER (Ag). Use a solution of silver nitrate (AgNO 3 ). 1. HC1 or soluble chlorides precipitate white, curdy silver chloride AgCl, insoluble in HNO 3 , soluble in NH 4 OH, form- ing ammonio-silver chloride (AgCl) 2 (NH 3 ) 3 , from which the chloride is again precipitated by acids. HC1 + AgNO 3 = AgCl -f HNO 3 . 2AgCl + 3 NH 4 OH = (AgCl) 2 (NH 3 ) 3 + 3 H 2 O. 2. H 2 S or NH 4 HS produces a black precipitate of silver sulphide Ag 2 S, insoluble in dilute acids and in alkalies, soluble in boiling HNO 3 . 3. KOH or NaOH forms a grayish-brown precipitate of silver oxide Ag 2 O, insoluble in excess, but soluble in NH 4 OH. 2KOH + 2AgNO 3 = Ag 2 O -f 2KNO 3 -f H 2 O. 4. NH 4 OH in small quantity, precipitates silver oxide, soluble in excess. 5. K,CrO 4 produces a red precipitate of silver chromate Ag 2 CrO 4 , soluble in concentrated HNO 3 and in NH 4 OH. 6. KI and KBr produce precipitates of silver iodide Agl, yellow, insoluble in NH 4 OH, and silver bromide AgBr, yellowish-white, slowly soluble in NH 4 OH. 7. KCN precipitates white silver cyanide AgCN, soluble in excess and in concentrated HNO S . BASES. 63 8. Heated with Na 2 CO 3 'on charcoal, before the blowpipe, compounds of silver form a bright, metallic button, soluble in HN0 3 . REACTIONS OF MERCURY as mercurous salt (Hg(ous)). Use a solution of mercurous nitrate (Hg 2 (NO 3 ^ 2 ). 1. HC1 or soluble chlorides precipitate mercurous chloride Hg 2 Cl 2 , converted by strong HNO 3 into a mixture of HgCl 2 and Hg(NO 3 ) 2 , also becoming black on the addition of NH 4 OH, forming NH 2 Hg 2 CL 2. H 2 S or NH 4 HS precipitates a mixture of Hg with HgS. 3. KOH or NaOH produces a black precipitate of mer- curous oxide Hg 2 O, insoluble in excess. 4. NH 4 OH causes a black precipitate of mercurous- ammonium nitrate NH 2 Hg 2 NO 3 . 2NH 4 OH + Hg 2 (N0 3 ) 2 = NH 2 Hg 2 N0 3 + NH 4 NO 3 -f 2 H 2 O. 5. K 2 CrO 4 forms an orange precipitate of mercurous chromate Hg 2 CrO 4 , 6. KI precipitates green mercurous iodide Hg 2 I 2 . 7. Before the blowpipe, mercurous salts volatilize, some being converted into mercuric salt and mercury, both of which sublime. REACTIONS OF LEAD (Pb). Use a solution of lead acetate (Pb(C 2 H 3 O 2 ) 2 ). 1. HC1, or soluble chlorides, produce a white precipitate of lead chloride PbCl 2 , soluble in hot water. 2. H 2 S or NH 4 HS precipitates black lead sulphide PbS, insoluble in HC1, soluble in hot HNO 3 . 3. KOH or NaOH produces a white precipitate of lead hydrate Pb(OH) 2 , soluble in large excess, forming potas- sium or sodium plumbate K 2 PbO 2 or Na 2 PbO 2 . 4. NH 4 OH precipitates white basic lead hydrate. 5. K 2 CrO 4 produces a yellow precipitate of lead chromate -PbCrO 4 , soluble in KOH and in strong HNO 3 . 6. KI forms a yellow precipitate of lead iodide PbI 2 , soluble in boiling water. 7. H 2 SO 4 produces a white precipitate of lead sulphate 64 ANALYTICAL CHEMISTRY. PbSO 4 , insoluble in acids, but soluble in solution of ammo- nium acetate or tartrate. 8. Na 2 CO 3 precipitates white basic lead carbonate (PbC0 3 ) 2 Pb(OH) 2 . 3 Na 2 C0 3 + 3 Pb(C 2 H 3 2 ) 2 + H 2 = (PbCO 3 ) 2 Pb(OH) 2 + 6NaC 2 H 3 O 2 -f CO 2 . 9. Before the blowpipe, on charcoal, lead compounds are converted into a malleable globule of the metal, with the formation of some yellow oxide. SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP VII. Ag Hg(ous) Pb HC1 White ppt. soluble in NH 4 OH. White ppt., turning black with NH 4 OH. White ppt. soluble in hot H 2 O. H 2 S and NH 4 HS Black ppt. Black ppt. Black ppt. KOH Brown ppt. Black ppt. White ppt. NH 4 OH Brown ppt. Black ppt. White ppt. Na 2 CO 3 Brown ppt. Black ppt. White ppt. K 2 CrO 4 Red ppt. Orange ppt. Yellow ppt. KI Yellow ppt. Green ppt. Yellow ppt. DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF GROUP VII. Add HC1, collect, wash and pour on the filter boiling H 2 O. Ppt. Ag, Hg(ous). Pour on the filter NH 4 OH. Filt. Pb. Add H 2 SO 4 , white ppt. Ppt. Hg(ous). Black. Filt. Ag. Add HNO 3 in excess, white ppt. "D 3s S-T &g m-z Uffi X jy.cO - r >M .o 2 a N) 5 s fig cO x .2-= a Sfl! 1 w^ u . - ~ a. 4lgiM^Jli a* .i^tSjr^s.? 13 e "4fj|lf?l&Stf 4^3ff-Jj|a P^HH .-c S| b^ sjljj -"S-^^: rj S -l Na 2 HPO 4 " Mg, j Notprecipitatedina | R> N ^ _ G L group When one of these reagents fails to produce a precipitate, it indicates the absence of that group, and the next group reagent is immediately added. , ACIDS. 67 SECTION II. ACIDS. GROUP I. HYDROCHLORIC ACID, HYDROBRO- MIC ACID, HYDRIODIC ACID, HYDROCY- ANIC ACID, HYDROFLUORIC ACID. REACTIONS OF HYDROCHLORIC ACID (HC1). Use a solution of potassium chloride (KC1). 1. AgNO 3 produces a white, curdy precipitate of silver chloride AgCl, insoluble in HNO 3 , readily soluble in NH 4 OH. This precipitate should be preserved to compare with those of AgBr and Agl. 2. Hg 2 (NO 3 ) 2 causes a white precipitate of mercurous chloride Hg 2 Cl 2 , insoluble in HNO 3 , blackening on the addition of NH 4 OH. 3. Pb(C 2 H 3 O 2 ) 2 forms a white crystalline precipitate of lead chloride PbCl 2 , soluble in 33 parts of boiling water. 4. Warming with H 2 SO 4 and MnO 2 causes the evolution of Chlorine, recognized by its odor and color. 5. On warming with H 2 SO 4 , hydrochloric acid is given off, recognized by its odor and intensely acid reaction; also by the dense white fumes of NH 4 C1 produced by holding a rod moistened with NH 4 OH near the mouth of the tube. REACTIONS OF HYDROBROMIC ACID (HBr). Use a solution of potassium bromide (KBr). 1. AgNO 3 produces a yellowish-white precipitate of silver bromide AgBr, insoluble in HNO 3 , slowly soluble in NH 4 OH. This precipitate should be compared with those of AgCl and Agl. 2. HgCl 2 forms a white precipitate of mercuric bromide HgBr 2 , soluble in a large quantity of water. 3. Hg 2 (NO 3 ) 2 precipitates yellowish-white mercurous bro- mide Hg 2 Br 2 . 68 ANALYTICAL CHEMISTRY. 4. H 2 SO 4 , when concentrated, causes the evolution of red vapors of bromine ; this occurs more readily in the presence of MnO 2 . 5. Chlorine water and starch paste cause a yellow color, due to formation of starch bromide ; in dilute solutions it is necessary to agitate the mixture with ether or chloroform, which will separate, carrying the bromine in solution, with a red or reddish-brown color. REACTIONS OF HYDRIODIC ACID (HI). Use a solution of potassium iodide (KI). 1. AgNO 3 produces a yellowish precipitate of silver iodide Agl, insoluble in HNO 3 and almost insoluble in NH 4 OH. 2. Hg 2 (NO 3 ) 2 precipitates green mercurous iodide Hg 2 I 2 . 3. HgCl 2 causes a red precipitate of mercuric iodide HgI 2 , soluble in excess of either reagent. 4. Pb(C 2 H 3 O 2 ) 2 produces a yellow precipitate of lead iodide PbI 2 , soluble in boiling water. 5. Chlorine water and starch paste form a blue color of starch iodide, which color disappears on heating and returns on cooling, also destroyed by excess of chlorine water. 6. A concentrated solution of one part CuSO 4 and three parts FeSO 4 produces a grayish precipitate of cuprous iodide Cu 2 I 2 . This reaction is useful in separating iodine from chlorine and bromine. HYDROFLUORIC ACID (HF). The evolution of intensely irritating fumes of this acid, on the addition of H 2 SO 4 to calcium fluoride, which etch glass, is sufficiently characteristic. REACTIONS OF HYDROCYANIC ACID (HCN). Use a solution of potassium cyanide (KCN). 1. AgNO 3 produces a white precipitate of AgCN, soluble in KCN, sparingly soluble in NH 4 OH, and insoluble in dilute HN0 3 . 2. NH 4 HS evaporated with KCN to dryness, on a water bath, will give on dissolving in water and adding Fe 2 Cl 6 a deep blood red color, due to formation of sulphocyanate. 3. Fe 2 Cl 6 and FeSO 4 , then NaOH until a precipitate is pro- ACIDS. 69 duced, heat and add HC1, will give a deep blue residue of ferric ferrocyanide (Prussian blue). SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP I. HC1 HBr HI HCN AgN0 3 White ppt. soluble in NH 4 OH Yellowish- white ppt. slowly soluble in NH 4 OH. Yellowish ppt. insoluble in NH 4 OH. White ppt. sparingly soluble - in NH 4 OH. Hg 2 (N0 3 ) 3 White ppt. Yellowish-white ppt. Greenish ppt. White ppt. HgCl a No ppt. White ppt. Red ppt. No ppt. Pb(C 2 H 3 2 ) 2 White ppt. White ppt. Yellow ppt. White ppt. Chlorine water and starch solution. No color. Yellow color. Blue color. No color. DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- TION CONTAINING SOLUBLE SALTS OF GROUP I. Add HNO 3 and boil for some time; the HCN is driven off. This operation should be performed with caution, so as to avoid inhaling these poisonous fumes. After all the HCN has been driven off, add AgNO 3 and shake well ; collect the precipitate of AgCl, AgBr, Agl on a filter, and, after washing, pour on the filter NH 4 OH ; AgCl will be dissolved, and may be detected in the filtrate by acidifying with HNO 3 . To detect HBr and HI take another portion of the original solution ; add a few drops of starch solution, and then, slowly, chlorine water, until, after agitation, the liquid smells of it. If there is a blue color HI is present. Continue the addition of chlorine water until the blue color is destroyed, when, if the solution is yellow, HBr is indicated. This may be still further verified by agitating with chloroform, which will, after separating, assume a red or red brown color if HBr is present. 70 ANALYTICAL CHEMISTRY. GROUP II. HYPOCHLOROUS ACID. CHLORIC " HYDROXYL. HYDROSULPHURIC ACID. SULPHUROUS SULPHURIC THIOSULPHURIC NITRIC HYPOPHOSPHOROUS ACID, ORTHOPHOSPHORIC PYROPHOSPHORIC " METAPHOSPHORIC BORIC CARBONIC SILICIC HYPOCHLOROUS ACID (HC1O). This acid is known in combination with calcium Ca(ClO) 2 , bleaching powder, and with sodium NaCIO, Labarraque's solution. The chlorine-like odor devel- oped on the addition of an acid is sufficient evidence of its presence. CHLORIC ACID (HC1O 3 ). The compounds of potassium and sodium with this acid are best known. 1. H 2 SO 4 , added to one of them, causes the evolution of yellow chlorine tetroxide C1 2 O 4 , having a characteristic odor. When the dry salt is used this reaction takes place with explosive violence. 2. AgNO 3 produces no precipitate. This is an important distinction from HC1. WATER (H 2 O), HYDRATES AND OXIDES. Water is distinguished by having no odor or taste, not changing litmus, and evaporating without residue. The soluble hydrates KOH, NaOH, LiOH, Ca(OH) 2 , Sr(OH) 2 , Ba(OH) 2 , yield solutions with water which change red litmus paper blue. The insoluble hydrates give off steam when heated in a dry tube. The soluble oxides K 2 O, Na 2 O, Li 2 O, CaO, SrO, BaO, are known by forming hydrates with water. The insoluble oxides are recognized by giving negative tests for acids when dis- solved and tested in the usual way. HYDROSULPHURIC ACID (H 2 S). Sufficient evidence of the presence of this acid is afforded by the characteristic odor. In the case of sulphides, first adding H 2 SO 4 and warming if necessary. A trace may be detected by holding over the mouth of the tube a piece of filter paper moistened with lead acetate, which will become black (lead sulphide) in the presence of H 2 S. ACIDS. 71 REACTIONS OF SULPHUROUS ACID (H 2 SO 3 ). 1. In solution, uncombined, this acid is recognized by its odor of burning sulphur, by strong bleaching action, by decolorizing potassium permanganate solution, and by causing the evolution of hydrogen sulphide when added to a mixture of zinc and hydrochloric acid. 2. Sulphites are distinguished by the characteristic odor of SO 2 on the addition of a strong acid. 3. Salts of Ag, Hg or Pb, produce precipitates which blacken on heating, owing to formation of sulphides. 4. BaCI 2 with neutral solutions forms a white precipitate of barium sulphite BaSO 3 , soluble in HC1. REACTIONS OF SULPHURIC ACID (H 2 SO 4 ). Use dilute H 2 SO 4 , or an alkali sulphate. 1. BaCl 2 produces a white precipitate of barium sulphate BaSO 4 , insoluble in boiling concentrated acids. 2. Pb(C 2 H 3 O 2 ) 2 causes a precipitate of white lead sulphate, insoluble in dilute acids, but soluble in hot concentrated acids. Alcohol increases the delicacy of this reaction. 3. A sulphate fused on charcoal with Na 2 CO a , the fused mass placed on a bright silver coin and moistened with a drop of dilute HC1 will cause a black stain, due to formation of silver sulphide. This reaction is especially adapted to the detection of insoluble sulphates. REACTIONS OF THIOSULPHURIC (HYPOSUL- PHUROUS) ACID (H 2 S 2 O 3 ). Use a solution of sodium thitfsulphate (Na^Os). 1. H 2 SO 4 causes the evolution of sulphurous oxide SO 2 , recognized by the. odor. A deposit of sulphur takes place at the same time, which is an important distinction from sulphites. 2. AgNO 3 produces a white precipitate of silver thio- sulphate Ag 2 S 2 O 3 , soluble in excess. After a time (immedi- ately on heating) the precipitate becomes dark and then black, silver sulphide and sulphuric acid being formed. 3. BaCl 2 produces a white precipitate soluble in excess of H 2 O, and decomposed by HC1. 72 ANALYTICAL CHEMISTRY. 4. Added to a mixture of zinc and hydrochloric acid, hydrogen sulphide is evolved. REACTIONS OF NITRIC ACID (HNO 3 ). Use a solution of potassium nitrate (KNO 8 ). 1. H 2 SO 4 , on heating, will cause the nitric acid to volatilize. If copper turnings be added with the sulphuric acid colorless nitrogen dioxide N 2 O 2 will be given off, which in contact with air will form red nitrogen tetroxide N 2 O 4 , readily recognized by the color and odor. If alcohol be added to the mixture the characteristic odor of nitrous ether is developed. 2. FeSO 4 , acidified with H 2 SO 4 , added in a test tube, so as to form a layer on a solution of a nitrate, acidified with H 2 SO 4 , will cause a dark layer to form at the line of contact. 3. Indigo solution, strongly acidified with H 2 SO 4 , is decolor- ized by a nitrate. 4. Zinc and potassium hydrate cause the reduction of the acid radical to ammonia, which may be detected in the usual way. This is a valuable test for distinguishing nitric acid in the presence of chloric acid. 5. Heated on charcoal, deflagration takes place, the char- coal burning at the expense of the oxygen of the nitrate. HYPOPHOSPHOROUS ACID (HH 2 PO 2 ). 1. By ignition the hypophosphites are resolved into spontaneously inflammable hydrogen phosphide PH 3 , and phosphate. 2. AgNO 3 produces at first a white precipitate of silver hypophosphite AgH 2 PO 2 , which soon becomes black, owing to formation of metallic silver. 3. HgQ 2 in excess causes a white precipitate of mercurous chloride Hg 2 Cl 2 > this takes place more rapidly on warming, and in the presence of HC1. REACTIONS OF ORTHOPHOSPHORIC ACID (H 3 PO 4 ). Use a solution of sodium phosphate (Na 2 HPO 4 ). 1. AgNO 3 causes a light yellow precipitate of silver phos- phate Ag 3 PO 4 , soluble in HNO 3 and in NH 4 OH. 2. Fe 2 Cl 6 , in presence of sodium acetate, produces a yel- lowish-white, gelatinous precipitate of ferric phosphate Fe 2 (PO 4 ) 2 . An excess of Fe 2 Cl 6 must be avoided. ACIDS. 73 3. (NH 4 ) 2 MoO 4 , in neutral or acid solution, causes a yellow precipitate to separate slowly, which is ammonium phospho- molybdate (NH 4 ) 3 PO 4 (MoO3) 10 + 2H 2 O, insoluble in HNO 3 , soluble in NH 4 OH. 4. Magnesia mixture (consisting of MgSO 4 , NH 4 C1, NH 4 OH) causes a white precipitate of ammonium magnesium phos- phate MgNH 4 PO 4 + 6H 2 O. Agitation facilitates the for- mation of this precipitate. 5. BaCl 2 , in neutral solution, produces a white precipitate of barium hydrogen phosphate BaHPO 4 . 6. Albumen (white of egg) does not cause a precipitate. REACTIONS OF PYROPHOSPHORIC ACID (H 4 P 2 O 7 ). Use a solution of sodium pyrophosphate (Na 4 P 2 O 7 ). 1. AgNO 3 precipitates white silver pyrophosphate Ag 4 - P 2 O 7 , soluble in HNO 3 and NH 4 OH. 2. MgSO 4 causes a white precipitate of magnesium pyro- phosphate Mg 2 P2O 7 , soluble in excess of the reagent ; from this solution NH 4 OH does not reprecipitate it in the cold. This reaction may be used to distinguish ortho- from pyro- phosphoric acid. 3. Neither (NH 4 ) 2 MoO 4 nor albumen produces a precipitate. 4. BaCl 2 , in neutral solution, precipitates white barium pyrophosphate Ba 2 P 2 O 7 . REACTIONS OF METAPHOSPHORIC ACID (HPO S ). Use a solution of sodium metaphosphate (NaPO 3 ). 1. AgNO 3 produces a white precipitate of silver meta- phosphate AgPO 3 , soluble in HNO 3 and NH 4 OH. 2. Albumen forms a white precipitate with the free acid, and with the salts on the addition of acetic acid. This is an important distinction from the ortho- and pyro-acids. 3. Neither (NH 4 ) 2 MoO 4 nor magnesia mixture produces a precipitate; should one form with the latter reagent, it is readily soluble in NH 4 C1. 4. BaCl 2 forms a white precipitate of barium metaphos- phate Ba(PO 3 ) 2 . 74 ANALYTICAL CHEMISTRY. 5. Solutions of the meta- and pyro-acids in water are con- verted into the ortho-acid by boiling. REACTIONS OF BORIC ACID (H 3 BO 3 ). Use a solution of sodium borate (Na 2 B 4 O 7 ). 1. AgNO 3 produces a white precipitate of silver borate, soluble in HNO 3 . 2. BaCl 2 precipitates white barium borate, soluble in excess of water and in NH 4 C1. 3. H 2 SO 4 or HC1 causes the separation of the acid, H 3 BO 3 , in crystalline form, from strong solutions. 4. Alcohol added to the acid, and ignited, burns with a characteristic green flame. In the case of salts, the addition of alcohol is preceded by that of H 2 SO 4 , in order to liberate the free acid. Salts of copper when present should be re- moved before this test is applied, as they likewise impart a green color to the alcohol flame. REACTIONS OF CARBONIC ACID (H 2 CO 3 ). Use a solution of sodium carbonate (Na 2 CO 3 ). 1. All free acids except HCN and H 2 S decompose carbon- ates with effervescence. The escaping gas passed into a solu- tion of Ba(OH) 2 or Ca(OH) 2 causes a white precipitate. 2. BaCl 2 , in neutral solution, precipitates .white barium carbonate BaCO 3 , soluble in acids. 3. CaCl 2 precipitates white calcium carbonate CaCO 3 , soluble in acids with effervescence. REACTIONS OF SILICIC ACID (H 4 SiO 4 ). 1. Insoluble silicates are determined by fusing on platinum foil some of the fine powder with Na 2 CO 3 , treating the fused mass with H 2 O and HC1, evaporating to dryness and redis- solving in H 2 O, when the silica, SiO 2 , will remain as a fine white precipitate. 2. Before the blowpipe, with a bead of microcosmic salt, silica forms the so-called silica skeleton, which is very charac- teristic. ACIDS. 75 3. Soluble silicates give gelatinous precipitates of silicic acid H 4 SiO 4 , on the addition of H 2 SO 4 or HC1; on evapor- ating this to dryness with a little HC1, and redissolving in H 2 O, silica SiO 2 , remains. 4. NH 4 C1 also precipitates H 4 SiO 4 when added to a soluble silicate. REACTIONS OF HYDROFERROCYANIC AND HYDROFERRICYANIC ACIDS. These acids have been sufficiently characterized under iron. DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- TION CONTAINING SOLUBLE SALTS OF GROUP II. I. Try the solution with litmus paper ; if alkaline, hydrates, carbonates, borates, silicates and, possibly, phosphates, may be present. If the solution is acid, neutralize with NaOH before applying the following tests : II. Evaporate a portion of the solution to dryness, and add concentrated H 2 SO 4 . The following acids will give character- istic reactions, and may be recognized by the further applica- tion of tests previously given : HC1O and HC1O 3 give the odor of chlorine. They are readily distinguished by the chlorate deflagrating with char- coal, by its liberating oxygen on heating, and by the yellowish- green gas, C1 2 O 4 , which is given off. The addition of H 2 SO 4 to a chlorate should be performed with very small quantities and with great care, on account of the tendency of the chlorine tetroxide to decompose with explosive violence. The most characteristic difference of HC1O is the white precipitate of AgCl which it gives with AgNO 3 . H 2 S is recognized by its peculiar odor, and by blackening a piece of filter paper moistened with solution of lead acetate. H 2 SO 3 and H 2 S 2 O 3 give off SO 2 , readily recognized by its odor and bleaching property. They are distinguished from each other by H 2 SO 4 producing with a concentrated solution of thiosulphate a white precipitate of sulphur in addition to the liberation of SO 2 . HNO 3 is detected by the peculiar acid fumes, which become red when metallic copper is added with the H 2 SO 4 . All the 76 ANALYTICAL CHEMISTRY. special tests, previously given, of this acid should be applied before a conclusion is reached concerning the presence or absence of it. HH 2 PO 2 is readily detected by its odor. H 3 BO 3 is easily detected when, in addition to the H 2 SO 4 , some alcohol is added and ignited. In the absence of salts of copper, the green flame is evidence of this acid. H 2 CO 3 is recognized by effervescence with the dilute acid in the cold without odor. When other gases are given off at the same time, the CO 2 may be detected by passing into lime water, which will cause a precipitate of CaCO 3 . H 4 SiO 4 is precipitated as a white, gelatinous mass when the H 2 SO 4 is added to a soluble silicate. Insoluble silicates will be treated of later. H 4 Fe(CN) 6 and H 6 Fe 2 (CN) 12 give the odor of HCN. Their presence is confirmed by the use of FeSO 4 and Fe 2 Cl 6 . III. The acids of this group not detected by H 2 SO 4 are sul- phuric and the three phosphoric acids. H 2 SO 4 is precipitated, on the addition of BaCl 2 , as BaSO 4 , insoluble in HC1 or HNO 3 . H 3 PO 4 is precipitated yellow on the addition of AgNO 3 , the precipitate being soluble in HNO 3 (distinction from HC1, HBr and HI). (NH 4 ) 2 MoO 4 is the most distinctive test. It causes a yellow precipitate when a few drops of the solution are added to some of the reagent in a test tube and warmed gently (dis- tinction from H 4 P 2 O 7 and HPO 3 ). Magnesia mixture is also a characteristic test for this acid. H 4 P 2 O 7 gives a white precipitate with AgNO 3 . It is further distinguished from the ortho-acid by (NH 4 ) 2 MoO 4 , and from the meta-acid by its behavior with MgSO 4 , and with albumen. HPO 3 also gives a white precipitate with AgNO 3 . It is best distinguished from the other varieties by albumen, by negative tests, and by boiling for some time and then applying the tests for the ortho-acid. H 3 AsO 4 and H 2 CrO 4 might be classified with this group, but a little care on the part of the student will detect them among the bases, and a little thought will tell him how to prove whether they are present as acids or bases. ACIDS. 77 GROUP III. ACETIC ACID. VALERIANIC ACID. STEARIC OLEIC LACTIC OXALIC SUCCINIC MALIC TARTARIC ACID. CITRIC CARBOLIC " BENZOIC SALICYLIC GALLIC TANNIC REACTIONS OF ACETIC ACID (HC 2 H 3 O 2 ). 1. In the free state acetic acid is readily recognized by its odor. 2. H 2 SO 4 , added to an acetate and the mixture warmed, gives the characteristic odor. 3. H 2 SO 4 and C 2 H 5 OH in equal volumes added to an acetate form acetic ether, readily recognized by its odor. 4. Fe 2 Cl 6 , with a neutral acetate, forms a deep red color, due to ferric acetate Fe 2 (C 2 H 3 O 2 ) 6 . VALERIANIC ACID (HC 5 H 9 O 2 ). The odor is sufficient evidence of the presence of this acid. This odor is developed by moisture and heat, and in the case of the salts by the addition of H 2 S0 4 . STEARIC ACID Stearic acid is a white, fatty solid, melting at 70.5 C., giving when combined with potassium a soft soap, and with sodium a hard soap ; from both it separates as an oily liquid on the addition of HCl, becoming solid on cooling. The lead salt, lead stearate, Pb(C 18 H 35 O 2 ) 2 is insoluble in ether. OLEIC ACID (HC 18 H33O 2 ). Oleic acid is an oily liquid at ordinary temperatures, but becomes solid at 4 C. and remains so until the temperature rises to 14 C., when it again be- comes liquid. Lead Oleate Pb(C 18 H 33 O 2 ) 2 , prepared by neutralizing sodium oleate with acetic acid and adding lead acetate, is almost entirely soluble in ether. This is an important distinction from stearic acid. LACTIC ACID (HC 3 H 5 O 3 ). Lactic acid is a colorless, syrupy liquid, of a slight unpleasant odor and a very sour taste, soluble in water, alcohol and ether, but insoluble in chloroform. When heated with H 2 SO 4 , CO is evolved. Lactates are all soluble in water, most of them sparingly ; they are insoluble in ether. Hg 2 (NO 3 ) 2 boiled with strong solution of a lactate, deposits crimson mer- curous lactate Hg 2 (C 3 H 5 O 3 j 2 . 78 ANALYTICAL CHEMISTRY. REACTIONS OF OXALIC ACID (H 2 C 2 O 4 2H 2 O). Use a solution of ammonium oxalate ((NH 4 ) 2 C 2 O 4 ). 1. K 2 Mn 2 O 8 , acidulated with H 2 SO 4 , is decolorized. 2. BaCl 2 produces a white precipitate of barium oxalate BaC 2 O 4 , sparingly soluble in acetic or oxalic acid, and freely soluble in HC1, HNO 3 and NH 4 CL 3. AgNO 3 precipitates white silver oxalate Ag 2 C 2 O 4 , readily soluble in hot concentrated HNO 3 , sparingly so in dilute acid, soluble in NH 4 OH. 4. CaCl 2 produces, even in highly diluted solutions, a white precipitate of calcium oxalate CaC 2 O 4 , soluble in HC1 and HNO 3 , but insoluble in acetic acid. 5. FeSO 4 with dilute solutions causes a yellow color; with more concentrated solutions and warming, a yellow precipitate ferrous oxalate FeC 2 O 4 , falls. This precipitate is insoluble in acetic acid, but is dissolved by HC1 and HNO 3 . SUCCINIC ACID (H 2 C 4 H 4 O 4 ). Use a solution of sodium succinate (Na 2 C 4 H 4 O 4 ). 1. Bad 2 produces no precipitate, but on the addition of alcohol a white pre- cipitate of barium succinate BaC 4 H 4 O 4 , falls, soluble in NH 4 C1. 2. Fe 2 Cl 6 causes a brownish-red, bulky precipitate of ferric succinate Fe 2 (C 4 H 4 4 ) 3 . 3. Pb(C 2 H 3 O 2 ) 2 precipitates white amorphous lead succinate PbC 4 H 4 O 4 , soluble in HNO 3 . MALIC ACID (H 2 C 4 H 4 5 ). 1. CaCl 2 fails to give a precipitate with the acid or its salts, until the mixture is boiled, when calcium malate CaC 4 H 4 O 5 H 2 O, separates. The addition of alcohol will also cause this precipitate. 2. Pb(C 2 H 3 O 2 ) 2 precipitates white lead malate PbC 4 H 4 O 5 3H 2 O. REACTIONS OF TARTARIC ACID (H 2 C 4 H 4 O 6 ). Use a solution of sodium tartrate (Na 2 C 4 H 4 O 6 ). 1. BaCl 2 precipitates white barium tartrate BaC 4 H 4 O 6 , soluble in ammonium salts, and in HC1. 2. CaCl 2 produces a white precipitate of calcium tartrate CaC 4 H 4 O 6 4H 2 O, ammonium salts prevent this precipitation. On adding KOH to this precipitate, it dissolves ; boil this ACIDS. 79 solution, calcium tartrate is again precipitated. Acetic acid is also a solvent for this precipitate. 3. AgNO 3 causes a white precipitate of silver tartrate Ag 2 C 4 H 4 O 6 ; on adding a few drops of NH 4 OH to dissolve the precipitate, and boiling, a mirror of metallic silver forms on the test tube. The test tube must be perfectly cleaned for this reaction; to accomplish this, rinse the tube with KOH solu- tion and then thoroughly with water. 4. Ca(OH) 2 in excess causes a precipitate of calcium tar- trate CaC 4 H 4 O 6 . This precipitate is flocculent at first, in which state it is readily soluble in NH 4 C1, but after standing for some time it becomes crystalline, in which case it is insoluble in NH 4 C1. 5. Tartaric acid and tartrates char on heating, and with H 2 SO 4 the odor of burnt sugar is given off. REACTIONS OF CITRIC ACID (H 3 C 6 H 5 O 7 H 2 O). Use a solution of sodium citrate (Na 3 C 6 H 5 O 7 ). 1. BaCl 2 produces a white precipitate of barium citrate Ba 3 (C 6 H 5 O 7 ) 2 , soluble in excess of water, in ammonium salts, and in acids. 2. CaCl 2 precipitates white calcium citrate Ca 3 (C 6 H 5 O 7 ) 2 , more insoluble in hot than in cold water, soluble in cold NH 4 C1, but insoluble in KOH. 3. AgNO 3 produces a white precipitate of silver citrate Ag 3 C 6 H 5 O 7 ; on boiling no metallic mirror is formed. 4. Ca(OH) 2 , in excess, does not cause a precipitate until the mixture is boiled. This is an important distinction from tartaric acid. 5. Citric acid and citrates char on heating, and with H 2 SO 4 give off the odor of burnt sugar. CARBOLIC ACID (C 6 H 5 OH). 1. HNO 3 , with an aqueous solution of the acid, forms a yellow color, due to picric acid C 6 H 2 (NO 2 ) 3 OH. 2. Fe 2 Cl 6 produces a violet blue color. 3. A piece of pine wood dipped in the acid, and then exposed to the fumes of HC1, becomes after a short time colored blue. 4. Bromine water causes a white flocculent precipitate. 5. Albumen is coagulated by the free acid. 80 ANALYTICAL CHEMISTRY. BENZOIC ACID (HC 7 H 5 O 2 ). npitates from neutral solutions, flesh-colored ferric benzoate Fe 2 (C 7 H 5 O 2 ) 6 , soluble in HC1, with separation of benzoic acid. 2. HC1 causes the separation of benzoic acid from cold solutions of the benzoates. 3. Bad 2 and CaCl 2 produce no precipitates with either the free acid or its salts. SALICYLIC ACID (HC 7 H 5 O 3 ). 1. Fe 2 Cl 6 produces a deep violet color, which is very characteristic. 2. Warmed with HjSO^ and methyl alcohol, the fragrant odor of methyl salicjfiate (oil of winter green) is developed. 3. HC1 causes the separation of the free acid from cold solutions of the salicylates. GALLIC ACID (HC 7 H 5 O 5 .H 2 O). 1. FeSO 4 produces no change. 2. Fe 2 Cl 6 produces a bluish-black precipitate, which dis- appears oh heating. 3. KOH, if not in excess, develops slowly a deep green color, which becomes red on the addition of acids. Alka- line carbonates cause the same green color, although more slowly. 4. No precipitate is produced with either gelatin or the alkaloids. With the former, however, a precipitation takes place when gum is present. TANNIC ACID (C 14 H 10 O 9 ). 1. FeSO 4 , when perfectly pure, causes no change; in the presence of oxygen, however, a dark color rapidly develops, which on standing, slowly becomes a precipitate. 2. Fe 2 Cl 6 produces a bluish-black precipitate. 3. Normal solution of iodine mixed with a small quantity of ammia, previously diluted with ten times its volume of water, jm)duces a brilliant fed color. This reaction will take placa wJh only traces of tannin. 4. Gelatin causes a white flocculent precipitate. This re- action is more delicate in the 'presence of small quantities of aluni. 5. Alkaloids produce white precipitates, soluble in acetic acid and alcohol. 6. K(SbO)C 4 H 4 O 6 causes a white, gelatinous precipitate. Most metallic salts cause precipitates with tannin. ACIDS. 81 DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- TION CONTAINING SOLUBLE SALTS OF GROUP III. As many of the acids in this group would indicate their presence by odor or physical appearance, a method of sepa- rating only the more important and closely related ones will be given. The list will^therefore be limited to Acetic acid,, Oxalic acid, Tartaric acid, Citric acid. If the Solution is Acid to Litmus Paper, neutralize with NaOH. I. To a small portion add H 2 SO 4 and warm. Acetic acid will, if present, be detected by its odor. II. To another portion add NH 4 OH until slightly alkaline, and then CaCl 2 ; allow to stand (avoiding heat) for ten rruhutes, and filter. Ppt. H 2 C 2 4 , H 2 C 4 H 4 6 . Wash, pour on the filter HC 2 H 3 O 2 . Filt. H 3 C 6 H B 7 . Boil to remove NH 4 OH, a white ppt. slowly forms on sides of the tube. Ppt. H 2 C 2 4 . Confirm by testing original solution with HC 2 H 3 O 3 andCaCl 2 . Filt. H 2 C 4 H 4 6 . Add NH 4 OH until slightly alkaline, white ppt. Confirm by forming mirror with AgNO 3 in original solution. 82 ANALYTICAL CHEMISTRY. SECTION III. DETECTION OF BASES AND ACIDS. SPECIAL OBSERVATION. The bases must always be determined first in a portion of the solution or powder, according to chart, page 65. If only the alkali metals are present and the reaction is neutral, proceed to search for the acids according to DIVI- SION I. If other than the alkali metals are present and the reaction is acid, proceed to search for the acids according to DIVISION II. When the substance is not entirely dissolved by H 2 O, HC1, or a mixture of HC1 and HNO 3 , proceed for both bases and acids according to DIVISION III. It will be seen from the above that before beginning the analysis of a solution it is necessary to try its action on litmus paper. If the reaction is alkaline it is necessary to neutralize with HC1, bearing in mind the precautions on page 66. When the analysis of a solid is to be performed, the action on litmus paper of the portion soluble in water is also to be noted, and the solution made neutral if necessary. All the physical properties of the solid should be carefully observed. EXCEPTION IN SEARCHING FOR BASES. When phosphates are present in acid solution, determined by adding a few drops of the solution to some HNO 3 and (NH 4 ) 2 MoO 4 and warming gently, the following chart must be used in the separation of Group IV. That is, bring the pre- cipitate containing Fe, Ce, Al, Cr, from chart, page 65, and work according to this. Oxalates also cause an exception to the general chart for separation of bases ; when, therefore, their presence is suspected, heat the substance to redness before examining for bases. n r ?-T' DETECTION OF BASES AND ACIDS. 83 DIRECTIONS FOR THE ANALYSIS OF GROUP IV WHEN PHOSPHATES ARE PRESENT. In addition to Fe, Ce, Al, Cr, there may be present the phosphates of Ca, Sr, Ba. Mn, Mg. Dissolve the precipitate in HC1, add Na 2 HP0 4 in excess, then excess of NH 4 C 2 H 3 O 2 ; boil, filter. Ppt. Fe 2 (P0 4 ) 2 A1 2 (P0 4 ) 2 Ce 2 (P0 4 ) 2 . Filt. Ba, Sr, Ca, Mn, Mg, Cr. (If green Cr is present.) Add K 2 CrO 4) warm, filter. Wash with hot H 2 O, then pour on the filter hot solution KOH. Ppt. Filt. Sr, Ca, Mn, Mg. Ba. Yellow. Add very dilute H 2 SO 4 , allow to stand, filter. An excess of H 2 SO 4 should be avoided. Ppt. Fe, Ce. Wash, dissolve in Filt. A1 2 (P0 4 ) 2 . HC1 and test a portion with K 4 Fe(CNj e Acidify with HC 2 H 3 O a white Confirm Filt. Ca, Mn, Mg. Add NH 4 C 2 H 3 O 2 and (NH 4 ) 2 C 2 O 4 , filter. for Fe. ppt. = Al. fl by name test. Ppt. Ca. White. Filt. Mg, Mn. Evaporate a portion to dryness and test with borax bead for Mn. To another por- tion add H 3 C 6 H 6 O 7 , then NH 4 OH in excess and To the remaining portion add Fe 2 Cl e , filter out the Fe 2 (PO 4 ) 2 , to filtrate add (NH 4 ) 2 C 2 4) NH 4 C1, NH 4 OH and NH 4 HS, white precipitate = Ce. to separate Fa and Mn, filter and test filtrate with Na 2 HPO 4 , white ppt. = Mg. DIVISION I. When the bases are the alkali metals, and the re- action is neutral. RULE I. Evaporate a portion of the solution to dryness, and slowly heat to redness. If the mass chars, one or more of the following organic acids are indicated : ACETATES, TARTRATES, CITRATES, GAL- LATES, TANNATES. NOTE. Oxalates do not char, although if the heating take place slowly a grayish coloration may be noticed, the residue in this case giving off CO 2 on the addition of H 2 SO 4 . Fe 2 Cl 6 immediately detects TANNATES and GALLATES. Gallic acid crystallizes in the cold on acidifying the solution. Tan- nic acid precipitates with solution of gelatin. The other acids are separated and detected according to chart, page 81. RULE II. Add to a second portion of the concentrated solu- tion, or the dry salt, strong H 2 SO 4 , warm gently, and note any of the following effects : Effervescence with dilute acid in the cold, \ no odor ......................................... J Carbonates. Effervescence on heating, no odor ...... / \ Oxalates ' confirm CaC1 and HC 2 H 3 2 . 84 ANALYTICAL CHEMISTRY. Effervescence with dilute acid on heating, ) , c TJ o r Sulphides, odor of H 2 S / Odor of SO 2 Sulphites. Odor of SO 2 with precipitation of S Thiosulphates. c Iodides, confirm by starch and Cl Dark brown color and violet fumes < i , j. , r f Bromides, confirm by starch and Dark red color and reddish fumes 1 I Cl water. Odor of HCN Cyanides. ,. TT .-,,, ., ,,. j . ( Ferro- or Ferri-cyanides, confirm Odor of HCN with crystalline deposit 1 I by Fe 2 Cl 6 and FeSO 4 . Odor of acetic acid Acetates. With dilute acid, odor of chlorine Hypochlorites. With strong acid, odor of chlorine Chlorates. Strongly acid suffocating fumes Chlorides, confirm by AgNO 3 . Strongly acid fumes becoming red when j Nitrates, confirm by FeSO 4 and metallic Cu is added \ H 2 SO 4 , also by indigo solution. c Benzoates, Succinates, Valerianates, Characteristic odors J ( Carbolates, Hypophosphites. RULE III. To a third portion add BaCl 2 . A white precipitate insoluble in HC1 indicates SULPHATES. RULE IV. To a fourth portion add CaCl 2 . A white precipitate soluble in excess of H 2 O indicates SULPHATES ; if insoluble in excess of H 2 O and in acetic acid, OXALATES are indicated. A white precipitate soluble in KOH, reprecipitated on boiling, indicates TARTRATES; confirm by boiling with NH 4 OH and AgNO 3 , forming a mirror of silver on the test tube. A white precipitate soluble in NH 4 C1 and reprecipitated on boiling indicates CITRATES. Citrates and Tartrates are also detected by heating some of the dry salt with H 2 SO 4 , when the odor of burnt sugar is developed. When all four of these acids are suspected to be present, treat the solution with HC1 and BaCl 2 to remove sulphates ; then neutralize the excess of HC1 with NaOH, and proceed according to chart, page 81. RULE V. To a fifth portion, acidified with HNO 3 , add AgNO 3 . A white, curdy precipitate, immediately and completely soluble in NH 4 OH indicates CHLORIDES. If slowly soluble in NH 4 OH, BROMIDES may be present; confirm by chlorine water and starch. DETECTION OF BASES AND ACIDS. 85 A yellowish precipitate insoluble in NH 4 OH = IODIDES. A white precipitate soluble in strong, hot HNO 3 = CYA- NIDES. When bromides and iodides are present they may be separated, by adding to the original solution chlorine water and starch ; continue the addition of chlorine water until the blue color of the starch iodide is discharged, then shake with chloroform ; when a red color is imparted to the separated chloroform, Br is indicated. For further instruction regarding the separation of these acids, consult chart, page 69. RULE VI. Add to a sixth portion of the solution magnesia mixture. A white precipitate indicates PHOSPHATES and ARSENATES. In another portion separate the As by HC1 and H 2 S, and repeat the reaction for H 3 PO 4 in the filtrate by neutralizing and adding magnesia mixture. RULE VII. To a seventh portion add Fe 2 Cl 6 . This reagent readily indicates TANNATES, GALLATES, FERRO- and FERRI-CYANIDES. RULE VIII. To an eighth portion apply the special test with H 2 SO 4 and FeSO 4 / containing 4.0 grams in i liter. As sodium hydrate is never, or very rarely absolutely pure, it is necessary to standardize this solution. For this take something more than the theoretical amount (about 50 grams) and dissolve in a liter of water. Place 100 c.c. of the standard oxalic acid solution in a beaker, and, having added the indi- cator, bring it under a burette containing some of the soda solution, and note the number of c.c. necessary to exactly neutralize the acid solution. Take of the alkaline solution ten times the number of c.c. necessary to neutralize the 100 c.c. of the acid solution, and add sufficient water to bring the measure to 1000 c.c. For instance, if it required 95 c.c. of the alkaline solution to neutralize the 100 c.c. of the oxalic acid solution, then 95 c.c. X 10 = 950 c.c. which amount diluted to 1000 c.c. would make it exactly equal in strength to the oxalic acid solution. One c.c. containing 0.040 of sodium hydrate NaOH, is the equivalent of GRAM. Acetic Acid, absolute, HC 2 H 3 O 2 0.0600 Citric Acid, "crystallized, H 3 C 6 H 5 O 7 H 2 O 0.0700 Hydrobromic Acid, absolute, HBr 0.0808 Hydrochloric Acid, absolute, HC1 0.0364 Hydriodic Acid, absolute, HI 0.1276 Lactic Acid, absolute, HC 3 H 5 O 3 0.0900 Nitric Acid, absolute, HNO 3 0.0630 Oxalic Acid, crystallized, H 2 C 2 O 4 2H 2 O 0.0630 Sulphuric Acid, absolute, H 2 SO 4 0.0490 Tartaric Acid, crystallized, H 2 C 4 H 4 O 6 0.0750 VOLUMETRIC SOLUTION OF POTASSIUM BI- CHROMATE (K 2 Cr 2 O 7 =: 294.8). A viginti-normal solution is used, containing 14.74. grams to i liter. Place the necessary amount of the salt (14.74 grams) in a liter flask, add about 800 c.c of water, agitate until dissolved, and bring the measure with water to 1000 c.c. This solution, when acidified with H 2 SO 4 , is used in the estimation of iron in VOLUMETRIC ESTIMATION. 105 the ferrous condition. The end of the reaction is determined by taking out a drop of the iron solution, and testing on a white porcelain surface with a drop of potassium ferricyanide solution ; when this ceases to give a blue color the reaction is at an end. One c.c. containing 0.01474 gram of potassium bichromate, K 2 Cr 2 O 7 , is the equivalent of GRAM. Iron in ferrous condition, Fe 0.01677 Ferrous Carbonate, FeCO 3 0.03477 Ferrous Sulphate, FeSO 4 7H 2 O . . . 0.08337 Ferrous Sulphate, (dry) FeSO 4 H 2 O 0.05097 VOLUMETRIC SOLUTION OF IODINE (I = 126.6). A deci-normal solution is used, containing 12.66 grams in I liter. Weigh the necessary amount of iodine (12.66 grams) in a stoppered tube, to prevent loss, as well as the corrosive action of the fumes on the balance. Also weigh 18 grams of potas- sium iodide and place in a liter flask with the iodine. Add about 200 c.c. of water. The iodine dissolves more readily in this strength of potassium iodide solution, besides it admits of more thorough agitation. When the solution is complete add water until the liquid measures 1000 c.c. Starch solution is used as an indicator in the determinations with the iodine solution. One c.c. containing 0.01266 iodine is the equivalent of GRAM. Arsenious Oxide, As 2 O 3 0.00494 Potassium Sulphite, crystallized, K 2 SO 3 2H 2 O 0.0097 Sodium Bisulphite, NaHSO 3 0.0052 Sodium Hyposulphite, crystallized, Na 2 S 2 O 3 5H 2 O 0.0248 Sodium Sulphite, crystallized, Na 2 SO 3 7H 2 O 0.0126 Sulphurous Oxide, SO 2 0.0032 VOLUMETRIC SOLUTION OF SODIUM HYPO- SULPHITE (Na 2 S 2 O 3 5H 2 O = 248). A deci-normal solution is used, containing 24.8 grams in I liter. Sodium hyposulphite cannot be prepared sufficiently pure to be relied on, consequently this solution must be standard- ized, therefore more than the theoretical amount is taken. 106 QUANTITATIVE ANALYSIS. The U. S. Pharmacopoeia recommends 32 grams. Dissolve this amount in 1000 c.c. of water ; place 100 c.c. of the standard solution of iodine in a beaker and run in the hyposulphite solution until the color of iodine nearly disappears, then add a small quantity of starch solution and continue until the blue color is discharged. Multiply the number of cubic centi- meters of the hyposulphite solution used by 10, and to that amount add sufficient water to bring the measure to 1000 c.c. The substances estimated by this solution either contain free iodine, or develop it on the addition of potassium iodide, so that starch solution may be used as an indicator. One c.c. containing 0.0248 gram of sodium hyposulphite is the equivalent of GRAM. Bromine, Br 0.00798 Chlorine, Cl 0.00354 Iodine, I 0.01266 VOLUMETRIC SOLUTION OF SILVER NITRATE (AgNO,= 169.7). A deci-normal solution is used, containing 16.97 grams in i liter. As silver nitrate can be obtained or prepared perfectly pure, the necessary amount (16.97 grams) is dissolved in sufficient distilled water to make 1000 c.c. In testing one of the following compounds, ammonium chloride for instance, a weighed amount is taken, dissolved in water, and a few drops of potassium bichromate solution added. The silver nitrate solution is then run in until a red precipitate remains permanently. The silver combines with the chlorine until the latter is all used, when it forms with the chromic acid red silver chromate, so that its formation indicates the end of the reaction. One c.c. containing 0.01697 gram of nitrate silver is the equivalent of GRAM. Ammonium Bromide, NH 4 Br 0.00978 Ammonium Chloride, NH 4 C1 0.00534 Ferrous Bromide, FeBr 2 0.01077 Ferrous Iodide, FeI 2 0.01545 Hydrocyanic Acid, absolute, HCN 0.00270 VOLUMETRIC ESTIMATION. 107 GRAM. Hydriodic Acid, HI 0.01276 Potassium Bromide, KBr 0.01188 Potassium Chloride, KC1 . 0.00744 Potassium Cyanide, KCN 0.01300 Sodium Bromide, NaBr 0.01028 Sodium Chloride, NCI 0.00584 TABLE OF ELEMENTS. ELEMENTS. SYMBOLS. ATOMIC WEIGHTS. Aluminium Al 27 Antimony Sb 120 Arsenic As 74.9 Barium Ba 136.8 Beryllium Be 9 Bismuth Bi 210 Boron B n Bromine Br 79.8 Cadmium Cd in.8 Caesium Cs . 132-6 Calcium Ca 40 Carbon C 12 Cerium Ce 140.2 Chlorine Cl 35.4 Chromium Cr 52.4 Cobalt Co 58.9 Copper Cu 63.2 Didymium Di !4 2 -3 Erbium . E I ^>S-9 Fluorine Fl 19 Gallium G 68.8 Gold Au 196.2 Hydrogen H I Indium : In U3-4 Iodine I 126.6 Iridium Ir l 9 2 -7 Iron . . . . ' Fe 55.9 Lanthanum La 138.2 Lead Pb 206.5 Lithium Li 7 Magnesium Mg . . 24 Manganese Mn 54 Mercury Hg l 99-7 Molybdenum Mo 95.5 Nickel Ni 58 Niobium Nb 94 Nitrogen N , . . . 14 108 QUANTITATIVE ANALYSIS. TABLE OF ELEMENTS. Continued. ELEMENTS. SYMBOLS. ATOMIC WEIGHTS. Osmium Os 198.5 Oxygen O 1 6 Palladium Pd 105.7 Phosphorus P 31 Platinum Pt 194-4 Potassium K 39 Rhodium Rh 104.1 Rubidium Rb 85.3 Ruthenium Ru 104.2 Scandium Sc 44 Selenium Se 78.8 Silicon Si 28 Silver Ag 107.7 Sodium Na 23 Strontium Sr 87.4 Sulphur S 32 Tantalum Ta 182 Tellurium Te 128 Thallium Tl 203.7 Thorium Th 233 Tin Sn II7-7 Titanium Ti 48 Tungsten W 183.6 Uranium U 2 3&>5 Vanadium V $1.3 Ytterbium Yb . 172.7 Yttrium Y 89.8 Zinc Zn 64.9 Zirconium Zr 90 Davyum, Decipium, Germanium, Neptunium and a few others, are hardly sufficiently known to warrant their being placed in the above list. INDEX. PAGE Acetates, detection of. ......................... 81, 83 reactions of. .................................. 77 Acetic acid ............................................ 77 Acid, acetic .......................................... 77 benzoic ......................................... 80 boric ............................................ 74 carbolic ....................................... 79 carbonic ....................................... 74 estimation of. ............................. 101 chloric ......................................... 70 citric ............................................ 79 gallic ............................................ 80 hydriodic ...................................... 68 hydrobromic ................................. 67 hydrochloric ............................. 20, 67 hydrocyanic ................................. 68 hydroferricyanic ........................... 75 hydroferrocyanic ........................... 75 hydrofluoric ................................. 68 hydrosulphuric .............................. 70 hypochlorous ............................... 70 hypophosphorous .......................... 72 hyppsulphurous ............................. 71 lactic ............................................ 77 malic ................ ........................... 78 metaphosphoric ............................. 73 nitric ......................................... 72 estimation of .............................. 100 oleic .............. . .............................. 77 orthophosphoric ........................... 72 oxalic .......................................... 78 phosphoric ................................... 72 pyrophosphoric ..................... ....... 73 salicylic ....................................... 80 silicic ......................................... 74 stearic .......................................... 77 succinic ....................................... 78 sulphuric ..................................... 71 estimation of. ............................. 100 sulphurous .................................... 71 tannic .......................................... 80 tartaric ........................................ 78 thiosulphuric ................................. 71 valerianic ......................... , ........... 77 Acids, analytical reactions of. ................. 67 detection of. ............................... 82-87 Acidum hydrochloricum .......................... 21 Alcohol, reactions and tests of. ................. 89 amylic, detection of. ...................... 89 Alkaloids, reactions and tests of ........... 91 Aluminium, analytical reactions of. ......... 47 hydrate ........................................ 29 Ammonia, preparation of. ....................... 23 properties of ................................. 23 Ammonium, analytical reactions of. ........ 34 nitrate ......................................... 27 Analysis of Acids Group I .......................... . ............. 69 " II ....................................... 75 " HI ....................................... 81 Bases Group 1 ....................................... 35 " II ....................................... 38 VII PACK Analysis of soluble salts 83 insoluble salts 86 Antimony, analytical reactions of 53 Marsh's test for 53 Reinsch's test for 53 Aqua chlori 20 Arsenates, analytical reactions of 51 detection of. 85 Arsenic, analytical reactions of 50 Fleitman's test for 52 Marsh's test for 51 Reinsch's test for 51 Arsenious oxide 50 Atomic weights, list of. 107 Atropine 93 Balance, rules for using 97 Barium, analytical reactions of. 35 chloride, estimation of. 98 Benzoic acid 80 Bismuth, analytical reactions of. 58 Bleaching 20 Boric acid, analytical reactions of 74 Bromides, detection of. 69, 84 Cadmium, analytical reactions of. 59 Caffeine 93 Calcium, analytical reactions of. 37 carbonate, estimation of. 101 phosphate, preparation of. 27 Cane sugar 91 Carbolic acid 79 Carbon dioxide, preparation of. 24 Carbonic acid, analytical reactions of. 74 estimation of. 101 Cerium, analytical reactions of 46 Chloral hydrate, reactions and tests of..... 89 Chloric acid 70 Chlorides, detection of. 69,84 Chlorine, estimation of. 99 preparation of 19 water 20 Chloroform, reactions and tests of. 88 Chromium, analytical reactions of. 47 Cinchonidine, reactions and tests of. 92 Cinchonine, reactions and tests of. 93 Citric acid, analytical reactions of. 79 detection of. 81, 84 Cobalt, analytical reactions of. 41 Copper, analytical reactions of 58 estimation of. 99 sulphate, estimation of. 99 preparation of. 30 Cyanides, detection of. 69, 85 Deci-normal solution 102 Dextrose 90 Drying 98 Elements, table of. 107 Gallic acid, analytical reactions of. 80 detection of. 83 Glauber salt 21 Glucose, reactions and tests of 90 Glycerin, reactions and tests of. 90 Gold, analytical reactions of. 54 Grape sugar go Gravimetric estimation 97 Group I, Acids 67 " II, " 70 " HI, " 77 " I, Bases 33 109 110 INDEX. PAGE Group II, Bases 35 III, " 40 IV, " : 45 V, " 50 VI, " 57 VII, " 62 Group reagents 66 Hydrates, analytical reactions of 70 Hydriodic acid, analytical reactions of 68 Hydrobromic acid, analytical reactions of 67 Hydrochloric acid, analytical reactions of.. 67 preparation of 20 properties of. 21 Hydrocyanic ecid, analytical reactions of. 68 Hydroferricvanic acid 75 Hydroferrocyanic acid 75 Hydrofluoric acid 68 Hydrogen 17 preparation of 17 properties of. 17 arsenide 51 Hydrosulphuric acid 70 Hypochlorous acid 70 Hypophosphorous acid 72 Hype-sulphurous acid 71 Ignition 98 Indicator 102 Iodides, detection of 69,85 lodoform, reactions and tests of 88 Iron, analytical reactions of. 45 Lactic acid 77 Lead, acetate of. 30 analytical reactions of 63 Lithium, analytical reactions of 34 Magnesium, analytical reactions of. 37 carbonate 28 oxide 29 sulphate 28 Malic acid 78 Manganese, analytical reactions of. 40 Marsh's test for arsenic 51 Mercuric compounds 57 analytical reactions of. 57 Mercurous compounds 63 analytical reactions of 63 Metaphosphoric acid 73 Morphine, reactions and tests of 91 Neutralization 102 Nickel, analytical reactions of. 42 Nitrates, detection of 85 Nitric acid, analytical reactions of 72 preparation of 24 properties of. 24 Nitrogen, preparation of. 22 properties of 23 Normal solution 102 Oleic acid 77 Oxidation -. 20 Oxides 70 Oxygen, preparation of 21 properties of 22 Phenolphtalein 102 Phosphates, detection of. 85 Phosphoric acid, analytical reactions of... 72 Platinum, analytical reactions of. 54 Potassium, analytical reactions of. 33 chloride, preparation of. 26 estimation of 100 PAGE Potassium hydrate, estimation of. 103 and sodium tartrate 26 Precautions on Group II, Bases " IV, V, " " VI, " " VII, " Precipitation Purple of Cassius 54 Quantitative analysis, preliminary direc- tions 97 Quinine, reactions and tests of. 92 Reinsch's test for arsenic 51 Rochelle salt 26 Saccharose, reactions and tests of. 91 Salicin 94 Salicylic acid 80 Santonin, reactions and tests of. 94 Seminormal solution 102 Silicic acid ... 74 Silver, analytical reactions of. 62 Sodium, analytical reactions of 33 sulphate 21 Solubilities, chart of. 87 Stannic compounds 54 Stannous compounds 53 Starch 91 Stearic acid 77 Strontium, analytical reactions of 36 Strychnine, reactions and tests of 92 Succinic acid 78 Sulphates, detection of. 76, 84 Sulphuretted hydrogen 70 Sulphuric acid, analytical reactions of. 71 Sulphurous acid, analytical reactions of... 71 Summary Group I, Acids 69 II, III, " . I, Bases. II, III, IV, V, VI, VII, 55 6o 64 Symbols, table of. 107 Tannates, detection of 85 Tannic acid, analytical reactions of. 80 Tartaric acid, analytical reactions of 78 Tartrates, detection of 81,84 Thalleioquin 92 Thiosulphuric acid 71 Tin, analytical reactions of. 53 Valerianic acid 77 Veratrine, reactions and tests of. 93 Volumetric analysis 102 estimation of alkalies 103 solutions 102 iodine 105 oxalic acid 102 potassium bichromate 104 silver nitrate 106 sodium hydrate 104 sodium hyposulphite 105 Wash bottle 28 Water 70 Zinc, analytical reactions of 41 sulphate, preparation of..... 19 14 DAY USE RETURN TO DESK FROM WHICH BORROWED LOAN DEPT. f This book is due on the last date stamped below, or on the date to which renewed. Renewed books are subject to immediate recall. i RZC D LLJ RECTD LD MAY 9 1960 SEP 24 1961 14 Mar'6%^ l8Nla;/6t B1 ?~J5-6i- k't^ APR 1 4 1962 t7 ' / MAY 18 1961 [ 1 JHN'GIRT iEu J LD .j APR 13 tan u/vs D IBbJ 12Dec'62JG -..,_.*... ... - I^MJ IxtC'O LD 30" , NOV 2 9 IQfi-* ' ''JOc. LD 21A-50m-4,'60 (A9562slO)476B General Library University of California Berkeley QV-7ST T7S- UNIVERSITY OF CALIFORNIA LIBRARY