LIBRARY UNIVERSITY OF CALIFORNIA. GIFT OF PP(OF,.W,B, Class . I .: '.;. 1111 iyPii i 16 which, on the further addition of the reagent, becomes yellow, then brown, and finally is converted into black HgS, mercuric sulphide : 3HgCl 2 + 2H 2 S = Hg 3 S 2 Cl 2 + 4HC1 ; Hg 3 S 2 Cl 2 + H 2 S = 3HgS -f 2HC1. HgS is not dissolved by ammonium sulphide ; sometimes, however, on being treated with ammonium sulphide, it is converted from black HgS into red HgS (cinnabar). Na^S, sodium sulphide, and K 2 S, potassium sulphide (par- ticularly in the presence of sodium or potassium hydroxide), dissolve mercuric sulphide, with the formation of HgS 2 Na 2 (3) or HgS 2 K 2 .> Mercuric sulphide is insoluble in boiling hydrochloric acid or in nitric acid, but by the continued action of hot concen- trated nitric acid it is converted into the white insoluble double salt Hg 3 S 2 (NO 3 ) 2 (as in the case of Hg 2 S, mercurous sulphide). Nitro-hydrochloric acid dissolves mercuric sulphide, with the formation of HgCl 2 , mercuric chloride : 3HgS 4- 6HC1 + 2HN0 3 = 3HgCl 2 + 2NO + 4H 2 O + S 3 . 2. NaOH, sodium hydroxide, and KOH, potassium hy- droxide, produce a brownish-red precipitate of a basic salt, which, upon further addition of the reagent, is converted into yellow HgO, mercuric oxide. 3. NH 4 OH, ammonium hydroxide, produces a white pre- cipitate ; thus, in a solution of HgCl 2 , mercuric chloride, white NH 2 HgCl, mercuric ammonium chloride is produced : 17 HgCl 2 + 2NH 4 OH = NH 2 HgCl + NH 4 C1 + 2H 2 O. 4. SnCl 2 , stannous chloride, added in small quantities to mercuric chloride or to mercuric salts containing a very slight quantity of free hydrochloric acid, precipitates white Hg 2 Cl 2 , mercurous chloride, which, on the addition of more stannous chloride, is reduced to gray, finely-divided, metallic mercury (very delicate reaction). 5. KI, potassium iodide, precipitates red HgI 2 , mercuric iodide, soluble in excess of the reagent, with the formation of HgI 2 (KI) 2 , potassium mercuric iodide. 6. A drop or two of a solution of a mercuric salt placed on clean copper foil produces a discoloration, due to the depo- sition of metallic mercury, as in the case of mercurous salts. On gently rubbing the spot it becomes mirror-like in appear- ance, and on heating the foil the spot disappears, due to the volatilization of the mercury. 7. Many of the mercuric salts, when heated in a glass re- duction-tube, sublime undecomposed, as, for example, HgCl 2 , mercuric chloride (corrosive sublimate), while others yield sublimates whic.li, because of an admixture of basic salts, are colored yellow. If the white (or yellow) sublimate be covered with dry sodium carbonate and again heated, red mercuric oxide is produced, which, on being more strongly heated, breaks up into metallic mercury and oxygen. COPPER, Cu (CUPRUM). Atomic weight, 63.18; valence, II. Reddish metal ; specific gravity, 8.94 ; melting-point, 1054 C. CuSOt, cupric sulphate, may be employed in making the tests. 1 . H 2 S, hydrogen sulphide, or (NH 4 ) 2 S, ammonium sulphide, precipitates black CuS, cupric sulphide, insoluble in dilute " b 2* 18 acids, insoluble in N^S, sodium sulphide, and in K 2 S, potas- sium sulphide. Ammonium sulphide (particularly the yellow ammonium sulphide) dissolves traces of the precipitate, with the formation of Cu^NHJ, = (CuS) 2 (NH 4 ) 2 S 5 . Boiling nitric acid dissolves CuS, with the formation of Cu(NO 3 ) 2 , cupric nitrate. It is also soluble in KCN, potassium cyanide : 2CuS + 4KCN = Cu 2 C 4 N 4 K 2 + K 2 S 2 . The precipitate (CuS), when moist and exposed to the air, readily absorbs oxygen, with the formation of CuSO 4 , cupric sulphate. 2. NaOH, sodium hydroxide, or KOH, potassium hy- droxide, produces in cold solution a voluminous flocculent precipitate of bluish-white Cu(OH) 2 , cupric hydroxide, insol- uble in excess of the reagent, but easily soluble in ammonium hydroxide. The precipitate, on being boiled with excess of sodium or potassium hydroxide, loses water and forms black CuO, cupric oxide. On adding sodium or potassium hydrox- ide to copper solutions containing non-volatile organic acids, and particularly containing such organic substances as glucose, (grape sugar,) glycerin, etc., and agitating the liquid, the bluish-white cupric hydroxide at first produced is immediately dissolved, with the production of a deep-blue liquid. (l) 3. NH 4 OH, ammonium hydroxide, added in small quanti- ties, produces a bluish- white precipitate of a basic salt, which is soluble in an excess of the reagent, producing a deep-blue 1 This property is made use of in the preparation of an alkaline copper solution in which cupric sulphate solution is added to a strong sodium hydroxide solution containing KNaC 4 H 4 O 6 , potassium sodium tartrate (Rochelle salt), the whole forming a deep-blue liquid (Fehling's solution), which is employed as a reagent for the detection of glucose (grape sugar). On boiling a solution containing glucose to which Fehling's solution has been added, insoluble red Cu 2 O, cuprous oxide, or yellow Cu 2 (OH) 2 , cuprous hydroxide, separates. 19 solution, due to the formation of Cu(NH 3 ) 4 SO 4 , cupric am- monium sulphate (very delicate reaction). Strong acid solu- tions are not generally precipitated by ammonium hydroxide. 4. Na 2 HPO 4 , sodium hydrogen phosphate, produces a bluish-green, flocculent precipitate of Cu 3 (PO 4 ) 2 , soluble in ammonium hydroxide. 5. K 4 Fe(CN) 6 , potassium ferrocyanide, precipitates brown- ish-red Cu 2 Fe(CN) 6 , cupric ferrocyanide (very delicate reac- tion). 6. KCN, potassium cyanide, added in excess to a neutral or ammoniacal solution of a salt of copper, produces a color- less solution of Cu 2 (CN) 2 (KCN) 2 , potassium cuprous cyanide : Cu(NO 3 ) 2 + 2KCN = Cu(CN) 2 -f 2KNO 3 ; Cu C 2 N 2 K 2Cu(CN) 2 -f2KCN= I + C 2 N 2 . Cu C 2 N 2 K The copper of this potassium salt of hydrocuprocyanic acid (1) is not precipitated by hydrogen sulphide (correspond- ing to the iron in potassium ferro- and ferricyanide, which is not precipitated by the ordinary reagents). 7. A bright piece of iron (knife-blade) free from grease placed in a solution of copper is soon covered with a reddish deposit of metallic copper. 8. Compounds of copper mixed with sodium carbonate and strongly heated on charcoal in the reducing flame yield reddish spangles or globules of metallic copper. 9. Compounds of copper fused in a bead of borax, Na 2 B 4 O 7 , held in a loop of platinum wire in the oxidizing flame of the blowpipe, yield a bluish-green bead. Fused in a bead of sodium ammonium phosphate, NaNH 4 HPO 4 (microcosmic Cu C,N 2 H. 20 salt), they yield, in the oxidizing flame, a bluish-green bead, which, when heated in the reducing flame, becomes reddish brown and opaque, due to the presence of separated metallic copper. The addition to the bead of a little metallic tin facilitates the reduction. BISMUTH, Bi. Atomic weight, 2O7.5; valence, III. Reddish-white metal ; specific gravity, 9.82 ; melting-point, 270 C. or BiCl 3 may be employed in making the tests. 1. H,jS, hydrogen sulphide, or (NH 4 ) 2 S, ammonium sul- phide, precipitates brownish-black BijS 3 , bismuth sulphide, insoluble in dilute acids and in ammonium sulphide. It is dissolved by boiling nitric acid, forming Bi(NO 3 ) 3 , bismuth nitrate. 2. NaOH, sodium hydroxide, KOH, potassium hydroxide, or NH 4 OH, ammonium hydroxide, precipitates white, amor- phous BiO-OH, bismuth hydroxide, insoluble in excess of the reagent. 3. K 2 CrO 4 , potassium chromate, precipitates yellow, crys- talline Bi 2 O(CrO 4 ) 2 , basic bismuth chromate, insoluble in sodium hydroxide, soluble in nitric acid. 4. A clear solution of a bismuth salt, when poured into a large quantity of water (provided the bismuth solution does not contain too much free acid), produces a milky turbidity, due to the separation of a white basic salt of bismuth. BiCl 3 , bis- muth chloride, yields BiOCl, bismuth oxychloride. Bi(NO 3 ) 3 , bismuth nitrate, yields first BiONO 3 , bismuth oxynitrate, and afterwards, especially on heating the liquid, Bi 2 O 2 NO 3 OH. (1) OH 21 A few drops of hydrochloric acid or of NH 4 C1, ammonium chloride, added to a bismuth nitrate solution before it is poured into the water, causes the separation of the bismuth as BiOCl, bismuth oxychloride. The reaction with BiCl 3 is the more delicate. Tartaric acid does not interfere with this reaction. 5. Bismuth salts, mixed with sodium carbonate and heated in the reducing flame on charcoal, yield brittle globules of me- tallic bismuth and a yellow incrustation of Bi 2 O 3 , bismuthous oxide. ARSENIC, As (ARSENICUM). Atomic weight, 74.9; valence, III, V. Steel-gray non-metal ; specific gravity, 5.72 at 14 C. Arsenic forms two compounds with oxygen, As 2 O 3 , arsen- ious oxide or anhydride, and As 2 O 5 , arsenic oxide or anhy- dride. BEHAVIOR OF ARSENIC IN THE ARSENIOUS CONDITION, AS ARSENIOUS ACID. As 2 3 , arsenious oxide, which, when dissolved in water, forms H 3 As0 3 , arsenious acid, may be employed in making the tests. 1. H 2 S, hydrogen sulphide, precipitates, from w^arm solu- tions acidulated with hydrochloric acid, yellow As 2 S 3 , arseni- ous sulphide, which is soluble in ammonium sulphide and in (NH 4 ) 2 CO 3 , ammonium carbonate, but is insoluble in hydro- chloric acid. Dissolved in ordinary colorless ammonium sulphide it forms (NH 4 ) 3 AsS 3 , ammonium sulpharsenite, and from this solution it may be reprecipitated by acids as As 2 S 3 , arsenious sulphide : 2(NH 4 ) 3 AsS 3 + 6HC1 = As 2 S 3 + 6NH 4 C1 + 3H 2 S. Dissolved in yellow ammonium sulphide it forms (NH 4 ) 3 AsS 4 , ammonium sulpharseniate. From this solution it is precip- itated by acids as As 2 S 5 , arsenic sulphide : 4 22 3(NH 4 ) 2 S + S 2 = 2(NH 4 ) 3 AsS 2(NH 4 ) 3 AsS 4 + 6HC1 = AS& + 6NH 4 C1 + 3H 2 S. Ammonium carbonate dissolves As 2 S 3 with the formation of ammonium sulpharsenite and ammonium arsenite : As& + 3(NH 4 ) 2 C0 3 = (NH 4 ) 3 AsS 3 + (NH 4 ) 3 A S O 3 + 3CO 2 . Acids reprecipitate it from this solution as As 2 S 3 : (NH 4 ) 3 AsS 3 + (NH 4 ) 3 As0 3 + 6HC1 = As 2 S 3 -f 6NH 4 C1 + 3H 2 0. As 2 S 5 , arsenic sulphide, dissolved in ammonium carbonate forms ammonium sulpharseniate and ammonium arseniate : 4AS& -f 12(NH 4 ) 2 C0 3 = 5(NH 4 ) 3 AsS 4 + 3(NH 4 ) 3 AsO 4 + 12CO 2 . From this solution acids reprecipitate it as As 2 S 5 : 5(NH 4 ) 3 AsS 4 + 3(NH 4 ) 3 AsO 4 + 24HC1 = 4As 2 S 5 -f- 24NH 4 C1 4- 12H 2 0. 2. AgNO 3 , argentic nitrate, added to an aqueous solution of arsenious acid and ammonium hydroxide added drop by drop produces a yellow, curdy precipitate of Ag 3 AsO 3 , ar- gentic arsenite, soluble in nitric acid and in ammonium hydroxide. 3. CuSO 4 , cupric sulphate, added to an aqueous solution of arsenious acid, and ammonium hydroxide subsequently added drop by drop, produces a greenish, flocculent precipi- tate of CuHAsO 3 , cupric arsenite (Scheele's green), soluble in excess of ammonium hydroxide and in acids. 4. Marsh's Test. When a few drops of a solution of ar- senious acid or a soluble arsenite are placed in an apparatus in which hydrogen is being evolved, the nascent hydrogen reduces the arsenical compound, and gaseous AsH 3 , hydrogen arsenide (arsenu retted hydrogen), is evolved with the free hydrogen : H 3 AsO 3 + 3Zn + 3H 2 SO 4 = AsH 3 -f 3ZnSO 4 -f 3H 2 O. When this mixture of hydrogen and hydrogen arsenide is 23 slowly passed through a glass tube heated to incipient red- ness, the hydrogen arsenide is decomposed, and the arsenic is deposited in the metallic state on the inner surface of the tube just beyond the heated part, as a lustrous brown, gray, or black coating. For this purpose the apparatus of Marsh, Fig. 1, is best adapted. FIG. 1. The apparatus consists of a small generating flask, A, a drying-tube, B, (l) containing small pieces of calcium chlo- ride, and a reduction-tube, (7, of hard glass, contracted at intervals. The metallic zinc and concentrated sulphuric acid (diluted with about four volumes of water) used in the operation must be free from arsenic, and therefore the following con- trol test should always be made to determine their purity. Zinc is placed in the flask A, and the drying-tube .5, to- gether with the reduction-tube (7, is connected with the flask. Dilute sulphuric acid (1-4) is introduced through the funnel-tube until the zinc is covered. When, after some minutes, (2) the evolved hydrogen has expelled the air from the 1 The drying-tube is sometimes dispensed with, and the reduction-tube connected directly with the delivery-tube of the flask. 2 If the action is slow, as is usually the case when pure zinc is em- ployed, it may be accelerated by the addition of a few drops of platinic chloride. 24 entire apparatus/ 1 * the flame of a Bunsen burner is applied to that part (at D) of the reduction-tube between the contracted portion and the drying-tube, and the tube then heated to incipient redness. After the flame has been applied for sev- eral minutes, the contracted part of the reduction-tube is ex- amined for the presence of a brownish, gray, or black lus- trous deposit. Should such a deposit be found, it indicates that either the zinc or the sulphuric acid, or both, are con- taminated with arsenic and therefore unfit for use in the test. If no deposit is produced by the above test, the application of the heat is continued, and the solution containing arsenic is introduced into the flask, through the funnel-tube. After the lapse of some minutes the contracted part of the tube immediately beyond the flame is examined for the presence of a brownish, gray, or black deposit. (2) A deposit having formed, the reduction-tube is detached (leaving it open at both ends, to permit the free access of air), inclined over a small flame, and gently heated at the part containing the deposit. The arsenic volatilizes, combines with oxygen, and deposits beyond the part heated, as As 2 O 3 , arsenious oxide, in minute octahedral crystals/ 3 ^ If the gas is ignited as it escapes from the contracted end of the tube, and the temperature of the flame is reduced by holding a piece of cold porcelain in it, incomplete combustion 1 Unless the air is expelled, an explosion, which may cause personal injury, is likely to occur when the flame is applied to the reduction- tube. 2 Antimony yields a deposit much resembling in color that produced by arsenic. The arsenical deposit is soluble in fresh NaOCl, sodium hypochlorite, whereas the antimony deposit is insoluble in that reagent. 3 The antimony deposit volatilizes and yields a white sublimate, which is generalty amorphous, or consists of minute granules and opaque gran- ular masses ; but it may contain well-defined octahedral crystals of Sb 2 O 3 , antimonious oxide. 25 occurs, and the arsenic is deposited on the porcelain in the metallic state in lustrous brown, gray, or black spots : 2AsH 3 + O 3 = As 2 + 3H 2 O. The arsenical deposit is soluble in fresh NaCIO, sodium hy- pochlorite. (Distinction from antimony.) As 2 H- SNaOCl + 3H 2 O = 2H 3 AsO 3 + 3NaCl. As hydrogen arsenide is exceedingly poisonous, it should not be allowed to escape in the room, but should be decom- posed by igniting it as it escapes from the tube, or conducting it into a solution of argentic nitrate, whereby reduction of the silver salt occurs with the separation of metallic silver, the arsenic remaining in solution : AsH 3 + 6 AgNO s + 3H 2 O 6 Ag + H 3 AsO 3 + 6HNO 8 . 5. Reinsch's Test. Metallic copper reduces arsenious oxide in acid solution to metallic arsenic, which is deposited on the copper as Cu 5 As 2 , cupric arsenide. The arsenical solution is acidulated with about one-seventh of its volume of hydro- chloric acid, a clean piece of copper foil placed in the solution, and the whole heated and kept almost at the boiling-point for several minutes. In this hot solution the arsenic is deposited on the foil as a grayish or black coating. The foil is taken from the liquid and im- mersed several times in water to wash off the hydrochloric acid, then pressed (without rubbing) between filter paper to free it from adherent moisture, and finally completely dried by being heated in a porcelain dish on a water- bath. It is then placed in a reduction-tube near the contracted part, the tube inclined, and the part containing the foil gently heated over a small flame (Fig. 2). 26 Volatilization of the arsenic and combination with oxygen take place, and octahedral crystals of As 2 O 3 , arsenious oxide, are deposited in the cooler part of the tube. 6. Arsenious oxide heated in a reduction-tube sublimes unchanged, and is deposited in the cooler portion of the tube in octahedral crystals. Heated in a dry reduction-tube with charcoal, a grayish or black mirror-like deposit of metallic arsenic is formed in the cooler part of the tube : As 2 O 3 + C 3 = As 2 + SCO. 7. Arsenious oxide or compounds of arsenic heated on charcoal in the reducing flame produce a garlic-like odor. (The arsenious oxide is first reduced to metallic arsenic, which volatilizes and combines with oxygen to form As 2 O 3 , which sometimes collects as an incrustation on the charcoal.) 8. Arsenious oxide or an arsenite, mixed with six times its bulk of a dry mixture consisting of equal parts of sodium carbonate and potassium cyanide and heated in a reduction- tube, is reduced, with the formation of a black glistening sublimate of metallic arsenic in the cool part of the tube. ARSENIC ACID. As 2 5 , arsenic oxide, which, dissolved in water, forms H 3 As0 4 , arsenic acid; or Na 3 As0 4 , sodium arseniate, may be employed in making the tests. 1. H 2 S, hydrogen sulphide, does not at first produce a precipitate, but reduces the arsenic acid to arsenious acid. Heating the solution facilitates the reduction : As. 2 O 5 -f 2H 2 S = As 2 O 3 + 2H 2 O + S 2 . Continuing the addition of hydrogen sulphide, As 2 S 3 , arsen- ious sulphide, is precipitated : 3H 2 S = As 2 S 3 + 3H 2 O. 27 The final precipitate is therefore a mixture of arsenious sul- phide and sulphur (As 2 S 3 -j- S 2 ). 2. AgNO 3 , argentic nitrate, added to a solution of arsenic acid which has been exactly neutralized with ammonium hydroxide, or to an arseniate, precipitates reddish-brown Ag 3 AsO 4 , argentic arseniate, soluble in nitric acid and in ammonium hydroxide : H 3 AsO 4 + 3AgNO 3 -f 3NH 4 OH t* Ag 3 AsO 4 + 3HN 4 NO 3 + 3H 2 0. 3. CuSO 4 , cupric sulphate, added to a solution of arsenic acid, followed by the addition of ammonium hydroxide drop by drop, or to an arseniate, produces a bluish-green precipi- tate of CuHAsO 4 , cupric arseniate, soluble in an excess of ammonium hydroxide and in acids. 4. The behavior of arsenic acid in Marsh's test or Reinsch's test, in the reduction-tube, mixed with charcoal, and on char- coal itself is identical with arsenious acid. 5. MgSO 4 , magnesium sulphate, added to a solution of ar- senic acid or an arseniate, followed by the addition of NH 4 C1, ammonium chloride, (1) and ammonium hydroxide (magnesia mixture), precipitates white, crystalline MgNH 4 AsO 4 -f- 6H 2 O, ammonium magnesium arseniate : H 3 As0 4 + MgS0 4 -f 3NH 4 OH = MgNH 4 AsO 4 -f- (NH 4 ) 2 SO 4 -f 3H 2 0. In concentrated solutions the precipitate forms immediately, and in dilute solutions gradually ; but is always perceptibly crystalline. It is soluble in 15,293 parts of cold water and less soluble in water containing ammonium hydroxide ; easily soluble in dilute acids, from which solutions it is reprecipitated by the addition of ammonium hydroxide. 1 The addition of ammonium chloride is for the purpose of preventing the precipitation of magnesium hydroxide. 28 6. NH 4 HMoO 4 , ammonium molybdate, added to a solution of arsenic acid rather strongly acidulated with nitric acid, and the whole gently warmed, produces a yellow precipitate of, possibly, (NH 4 ) 3 AsO 4 (MoO 3 ) 10 , ammonium molybdoarseni- ate, soluble in ammonium hydroxide and reprecipitated from this solution by nitric acid. The presence of hydrochloric acid or of chlorides interferes with the delicacy of the reaction. 7. To detect arsenic acid in the presence of arsenious acid (providing the compounds are soluble in water) their behavior with magnesia mixture (compare above, 5) is made use of; arsenious acid produces no precipitate with magnesia mixture. In case the compound is insoluble in water, it is dissolved in hydrochloric acid, and the arsenious acid is precipitated in cold solution with hydrogen sulphide. The resulting arseni- ous sulphide is removed by filtration, the filtrate is warmed, and hydrogen sulphide again conducted into the liquid. The production of a precipitate indicates the presence of arsenic acid. ANTIMONY, Sb (STIBIUM). Atomic weight, 119.6; valence, III, V. Silvery-white metal ; specific gravity 6.7 ; melting-point, 425 C. Antimony forms two typical compounds with oxygen, Sb 2 O 3 , antimonious oxide, and Sb 2 O 5 , antimonic oxide. BEHAVIOR OF ANTIMONY IN THE ANTIMONIOUS CONDITION. SbCl 3 , antimonious chloride, may be employed in making the I . H 2 S, hydrogen sulphide, produces in solutions of anti- monious salts which are not too strongly acidulated an orange- red precipitate of Sb 2 S 3 , antimonious sulphide, insoluble in dilute acids, soluble in concentrated hydrochloric acid (without 29 the separation of sulphur) and also in ammonium sulphide and in sodium or potassium sulphide ; insoluble in ammonium carbonate (distinction from arsenic). When dissolved in colorless ammonium sulphide it forms (NH 4 ) 3 SbS 3 , ammonium sulphantimonite : Sb 2 S 3 + 3(NH 4 ) 2 S - 2(NH 4 ) 3 SbS 3 ; and when dissolved in yellow ammonium sulphide it forms (NH 4 ) 3 SbS 4 , ammonium sulphantimonate : Sb 2 S 3 + 3(1S T H 4 ) 2 S + S 2 = 2(NH 4 ) 3 SbS 4 . Hydrochloric acid precipitates from the sulphantimonite so- lution Sb 2 S 3 , and from the sulphantimonate solution Sb 2 S 5 : 2(NH 4 ) 3 SbS 3 + 6HC1 = Sb 2 S 3 + 6NH 4 C1 + 3H 2 S ; 2(NH 4 ) 3 SbS 4 + 6HC1 = Sb 2 S 5 + 6NH 4 C1 + 3H 2 S. 2. NaOH, sodium hydroxide, as well as KOH, potassium hydroxide, produces a white voluminous precipitate of Sb(OH) 3 , antimonious hydroxide, readily soluble in an excess of the reagent, forming SbO(ONa), sodium antimonite, or SbO(OK), potassium antimonite. On being boiled in the alkaline liquid the precipitate of Sb(OH) 3 is converted into Sb 2 O 3 , antimonious oxide. 3. NH 4 OH, ammonium hydroxide, precipitates white Sb(OH) 3 , antimonious hydroxide, insoluble in an excess of the reagent. Tartaric acid prevents the precipitation. 4. On pouring a solution of an antimonious salt, as, for example, SbCl 3 , antimonious chloride, into a large quantity of water, a white precipitate of a mixture of SbOCl, anti- monious oxychloride, and Sb 4 O 5 Cl 2 is produced : SbCl 3 + H 2 O = SbOCl + 2HC1 ; 4SbCl 3 + 5H 2 O = Sb 4 O 5 Cl 2 -f 10HC1. A milkiness is produced in water by even the slightest quan- tity of antimonious chloride. Tartaric acid prevents the pre- cipitation by dissolving the precipitate : SbOCl + H 2 C 4 H 4 6 = (SbO)HC 4 H 4 6 + HC1. 3* 30 5. Soluble salts of antimony, placed in a flask in which hydrogen is being generated from zinc and dilute sulphuric acid (14), are decomposed, with the formation of gaseous SbH 3 , antimonious hydride (antimouuretted hydrogen) : 2SbCl s + Zn 6 + 3H 2 SO 4 = 2SbH 3 -f 3ZnCl 2 + 3ZnSO 4 . The apparatus of Marsh is best adapted for this purpose, and the same precautions as given under arsenic should be ob- served. (See page 23.) On heating the reduction-tube of Marsh's apparatus to dull redness and slowly passing antimonious hydride through the tube, the compound is reduced, and a lustrous brown or black deposit of metallic antimony is formed in the part of the tube before the flame, or on both sides of the flame. If the gas is ignited as it escapes from the contracted end of the tube and the temperature of the flame reduced by hold- ing a piece of cold porcelain in it, incomplete combustion occurs, and the antimony is deposited on the porcelain in dull brownish or black spots : 2SbH 3 -f O 3 = Sb 2 -f 3H 2 O. The deposit of metallic antimony is insoluble in fresh sodium hypochlorite (distinction from arsenic). The deposit of metallic antimony in the tube, on being gently heated over a small flame with free access of air, (1) vol- atilizes, combines with oxygen, and condenses in the cooler part of the tube as white Sb 2 O 3 , antimonious oxide. The sublimate is usually entirely amorphous, but occasionally may contain octahedral crystals of antimonious oxide. 6. Compounds of antimony in acid solution are reduced on being heated with a piece of bright copper foil, with the deposition of the antimony as a grayish or black coating upon the copper. On washing the foil with water, drying, 1 As in Marsh's test for arsenic. (See page 24.) 31 and gently heating it in a small reduction-tube over a flame, the antimony volatilizes, combines with oxygen, and deposits in the cooler part of the tube as amorphous Sb 2 O 3 , antimoni- ous oxide, which may sometimes contain octahedral crystals of antimonious oxide. (See Reinsch's Test for Arsenic, page 25.) 7. Metallic zinc reduces antimonious solutions, the anti- mony separating as a black powder. If a drop of the anti- monious solution is placed on a piece of platinum foil and a small fragment of zinc is placed in the solution, the antimony is deposited on the foil as a brown or black adherent coating, insoluble in hydrochloric acid : 2SbCl 3 + Zn s = Sb 2 -f- 3ZnCl 2 . 8. Compounds of antimony, when heated in the reducing flame with sodium carbonate on charcoal, yield a white, brittle globule of metallic antimony, usually coated with a white incrustation of Sb 2 O 3 , antimonious oxide. BEHAVIOR OF ANTIMONY IN THE ANTIMONIC CONDITION. SbCl 5 , antimonic chloride, may be employed in making the tests. 1. H 2 S, hydrogen sulphide, precipitates from acid solu- tions orange-red Sb 2 S 5 , antimonic sulphide, insoluble in dilute acids and in ammonium carbonate, soluble in concentrated hydrochloric acid, forming SbCl 3; antimonious chloride (with the separation of sulphur) : Sb 2 S 5 + 6HC1 == 2SbCl 3 + 3H 2 S + S 2 , and soluble in ammonium sulphide and in sodium or potas- sium sulphide, with the formation of sulphantimonates : Sb 2 S 5 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 SbS 4 . 2. The behavior of antimonic compounds is similar to 32 that of antimonious compounds in respect to the tests with zinc and dilute sulphuric acid, copper foil and hydrochloric acid, and platinum foil and zinc. 3. To detect antimonious compounds in the presence of antimonic compounds, the behavior of an alkaline solution of antimonious oxide with a silver solution is taken advan- tage of. On adding argentic nitrate to the alkaline solution and gently heating it, a precipitate composed of Ag 2 O, argentic oxide, and metallic silver is formed. Ammonium hydroxide has the property of dissolving only the argentic oxide, leaving the metallic silver undissolved : SbO(OK) + Ag 2 O = KSbO 3 + Ag 2 . After washing the precipitate and then treating it with ammonium hydroxide, metallic silver will remain undis- solved in case an antimonious compound was originally present. To detect antimonic compounds in the presence of anti- monious compounds, the alkaline solution is acidulated with hydrochloric acid, KI, potassium iodide, added, and then boiled ; in the presence of an antimonic compound iodine is separated : SbCl 5 + 2HI = SbCl 3 + 2HC1 + I 2 . TIN, Sn (STANNUM). Atomic weight, 117.35; valence, II, IV. Bluish-white metal ; specific gravity, 7.29 ; melting-point, 235 C. Tin forms two series of compounds, named respectively stannous and stannic compounds. SnO, stannous oxide, may be taken as the type of the stannous, and SnO 2 as the type of the stannic compounds. 33 BEHAVIOR OF TIN IN THE STANNOUS CONDITION. SnCl 2) stannous chloride, may be employed in making the tests. 1. H 2 S, hydrogen sulphide (also ammonium sulphide), precipitates dark-brown SnS, stannous sulphide, insoluble in colorless ammonium sulphide, but easily soluble in yellow ammonium sulphide, witli the formation of ammonium sulphostannatc, (NH 4 ) 2 SnS 3 : SnS + (NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 3 . From this solution acids precipitate yellow SnS 2 , stannic sul- phide : (NH 4 ) 2 SnS 3 -f 2HC1 = SnS 2 + 2NH 4 C1 -f H 2 S. 2. NaOH, sodium hydroxide, as well as KOH, potassium hydroxide, precipitates white Sn(OH) 2 , stannous hydroxide, soluble in excess of the reagent. On boiling a solution of a stannous salt to which an insufficient quantity of sodium or potassium hydroxide has been added, the Sn(OH) 2 is con- verted into black SnO, stannous oxide. 3. NH 4 OH, ammonium hydroxide, precipitates white Sn(OH) 2 , stannous hydroxide, insoluble in excess of the reagent. 4. HgCl 2 , mercuric chloride, added to an excess of SnCl 2 , stannous chloride, produces a grayish precipitate of finely- divided metallic mercury : If, on the other hand, an excess of mercuric chloride is added to a stannous chloride solution, a white precipitate of Hg 2 Cl 2 , mercurous chloride, is formed : 2HgCl 2 + SnCl 2 = Hg 2 Cl 2 -f- SnCl 4 (a very delicate test and a means of distinction between stan- nous and stannic salts). 34 5. A fragment of metallic zinc placed in a solution of stannous chloride precipitates grayish metallic tin : SnCl 2 + Zn = ZnCl 2 -f Sn. If performed on platinum foil (see 7 under Antimony, page 31) the tin which separates does not adhere to the platinum foil as a black coating (distinction from antimony). 6. Both stannous sails and stannic salts, when fused with sodium carbonate, or with a mixture of sodium carbonate and potassium cyanide, in the reducing flame on charcoal, yield white, ductile globules of metallic tin together with a slight incrustation of SnO 2 , stannic oxide. Stannic oxide moistened with Co(NO 3 ) 2 , cobaltous nitrate, and heated in the blowpipe- flame becomes bluish green in color. BEHAVIOR OF TIN IN THE STANNIC CONDITION. SnCl 4 , stannic chloride, may be employed in making the tests. 1 . H 2 S, hydrogen sulphide, precipitates yellow SnS 2 , stannic sulphide, insoluble in ammonium carbonate, but soluble in colorless and also in yellow ammonium sulphide, with the formation of (NH 4 ) 2 SnS 3 , ammonium sulphostannate. From this solution SnS 2 is reprecipitated on the addition of acids. SnS 2 is soluble in concentrated hydrochloric acid. 2. NaOH, sodium hydroxide, KOH, potassium hydroxide, or NH 4 OH, ammonium hydroxide, produces in solutions of stannic salts white precipitates. The precipitate produced in hydrochloric acid solutions of ordinary SnO 2 , stannic oxide, is Sn(OH) 4 , stannic hydroxide, and is easily soluble in dilute sodium or potassium hydroxide ; that produced in solutions of metastannic acid is metastannic hydroxide, only slightly soluble in excess of the reagent. 3. Na 2 SO 4 , sodium sulphate, or ]S T H 4 NO 3 , ammonium ni- 35 trate, in saturated solution, added in excess to a hydrochloric acid solution of stannic oxide, precipitates the tin, particularly on the application of heat, as white Sn(OH) 4 , stannic hydrox- ide, or as (Sn(OH) 4 ) M , metastannic hydroxide : SnCl 4 + 4Na 2 S0 4 + 4H 2 O = Sn(OH) 4 -f- 4NaCl -j- 4NaHSO 4 ; SnCl 4 + 4NH 4 NO 3 + 4H 2 O=Sn(OH) 4 + 4NH 4 C1 + 4HNO 3 . (Distinction from stannous salts.) 4. Metallic zinc reduces stannic salts in solution to metallic tin in the same manner as it reduces stannous salts. (See 5, page 34.) CADMIUM, Cd. Atomic weight, 111.7 ; valence, II. Bluish-white metal ; specific gravity, 8.54 ; melting-point, 315 C. CdS0 4) cadmium sulphate, may be employed in making the tests. 1. H 2 S, hydrogen sulphide, or ammonium sulphide pro- duces a yellow precipitate of CdS, cadmium sulphide, in- soluble in dilute acids, in ammonium and sodium sulphides, and in potassium cyanide, soluble in boiling nitric acid, with the formation of Cd(NO 3 ) 2 , cadmium nitrate. 2. NaOH, sodium hydroxide, as well as KOH, potassium hydroxide, precipitates white Cd(OH) 2 , cadmium hydroxide, insoluble in excess of the reagent. 3. NH 4 OH, ammonium hydroxide, precipitates white Cd(OH) 2 , cadmium hydroxide, soluble in excess of the re- agent, probably with the formation of a double salt of cad- mium and ammonium, Cd(ONH 4 ) 2 . 4. KCN, potassium cyanide, added to a neutral or ammo- niacal solution of a cadmium salt, precipitates white Cd(CN) 2 , 36 cadmium cyanide, which is soluble in an excess of potassium cyanide, forming a colorless solution of Cd(CN) 2 (KCN) 2 : CdS0 4 + 2KCN = Cd(CN) 2 + K 2 SO 4 ; Cd(CN) 2 + 2KCN = Cd(CN) 2 (KCN) 2 . Hydrogen sulphide precipitates from this solution yellow CdS, cadmium sulphide. 5. Cadmium compounds, mixed with sodium carbonate and fused in the reducing flame on charcoal, yield yellow to brown incrustations of CdO, cadmium oxide. GOLD, Au (AURUM). Atomic weight, 196.2; valence, III. Yellow metal ; specific gravity, 19.26 ; melting-point, 1035 C. AuCls, auric chloride, may be employed in making the tests. 1. H 2 S, hydrogen sulphide, produces in a cold solution of auric chloride a black precipitate of AiigSj, auric sulphide, soluble in ammonium sulphide. From hot auric chloride solutions hydrogen sulphide pre- cipitates brownish metallic gold : 8 AuCl s + 3H 2 S + 1 2H 2 O = Au 8 + 24HC1 + 3H 2 SO 4 . 2. NaOH, sodium hydroxide, and also potassium hydrox- ide precipitate reddish-yellow, amorphous Au(OH) 3 , auric hydroxide, soluble in excess of the reagent. 3. NH 4 OH, ammonium hydroxide, produces a reddish- yellow precipitate of (NH 3 ) 2 Au 2 O 3 , ammonium aurate (fulmi- nating gold) : 2AuCl 3 + 8NH 4 OH = (NH s ) 2 Au 2 O s + 6NH 4 C1 + 5H 2 O. 4. FeSO 4 , ferrous sulphate, precipitates in the presence of a free mineral acid, even in the cold, but especially on heating, 37 metallic gold, brownish in color because of its finely-divided condition : 2AuCl 3 + 6FeSO 4 = Au 2 + Fe 2 Cl 6 + 2Fe 2 (SO 4 ) 3 . 5. H 2 C 2 O 4 , oxalic acid, also precipitates metallic gold from auric chloride solutions : 2AuCl 3 + 3H 2 C 2 O 4 = Au 2 + 6HC1 + 6CO 2 . The precipitation proceeds slowly, but is complete. Warm- ing the solution facilitates the reduction. The presence of considerable free mineral acid interferes with the precipitation. 6. SnCl 2 , stannous chloride, especially in very dilute solu- tion, added to auric chloride solutions, produces a purplish- red coloration or a purplish-red precipitate (purple of Cassius), consisting probably of a mixture of finely-divided gold and stannic oxide. 7. Compounds of gold fused with sodium carbonate or with borax on charcoal yield yellow, glistening, ductile spangles of metallic gold. PLATINUM, Pt. Atomic weight, 194.3; valence, IV. Tin- white metal ; specific gravity, 21.46 ; melting-point, 1775 C. platinic chloride, may be employed in making the 1. H 2 S, hydrogen sulphide, produces in cold platinic chlo- ride solutions a brownish coloration, but after some time has elapsed a brownish-black precipitate of PtS 2 , platinic sulphide, separates. The precipitate appears at once on heating the solution. The precipitate is insoluble in hydrochloric acid and also in nitric acid, but soluble in nitro-hydrochloric acid (aqua regia) and also in ammonium sulphide. 2. KNO 3 , potassium nitrate, to which a drop of hydro- 4 38 chloric acid has been added, or potassium chloride, added to a concentrated solution of platinic chloride, produces a yellow, crystalline precipitate of (KCl) 2 PtCl 4 , potassium platinic chlo- ride, slightly soluble in water, insoluble in alcohol. The test is best made in a watch-glass, the liquid being stirred with a glass rod. Alcohol facilitates precipitation. 3. NH 4 C1, ammonium chloride, produces a yellow, crystal- line precipitate of (NH 4 Cl) 2 PtCl 4 , ammonium platinic chlo- ride, slightly soluble in water, insoluble in alcohol. This test is best made in a watch-glass as in 2, above. 4. Compounds of platinum heated in the reducing flame are reduced to spongy metallic platinum. THIRD GROUP. Metals precipitated as hydroxides by NH 4 OH, ammonium hydroxide : Iron, Aluminium, and Chromium. IRON, Fe (FERRUM). Atomic weight, 55.88; valence, II, IV. Silver-white metal ; specific gravity, 7.84. Iron forms two typical series of compounds, named re- spectively ferrous and ferric compounds. FeO, ferrous oxide, may be taken as the type of the ferrous compounds, and Fe 2 O 3 as the type of the ferric compounds. BEHAVIOR OF IRON IN THE FERROUS CONDITION. FeSOu ferrous sulphate, may be employed in making the tests. 1. NH 4 OH, NaOH, or KOH precipitates, in solutions of ferrous salts which are free from dissolved air, white 39 Fe(OH) 2 , ferrous hydroxide, which, by the absorption of oxygen, quickly changes in color to green, black, and finally to reddish brown. The presence of ammonium chloride or sulphate retards the precipitation ; nevertheless, in these alka- line solutions, in consequence of the absorption of oxygen, black ferrous hydroxide and reddish-brown ferric hydroxide gradually separate. 2. (NH 4 ) 2 S precipitates black FeS, ferrous sulphide, insol- uble in excess of the reagent, easily soluble in hydrochloric acid and in nitric acid. Very dilute ferrous solutions are colored green by ammonium sulphide. Moist ferrous sul- phide is oxidized on exposure to the air and changes to reddish-brown Fe 2 O(SO 4 ) 2 , basic ferric sulphate. 3. K 4 Fe(CN) 6 , potassium ferrocyanide, produces in ferrous solutions free from ferric salts a white precipitate, which quickly changes to bluish-white K 2 Fe(Fe(CN) 6 ), potassium ferrous ferrocyanide (Everett's salt), insoluble in acids. On exposure to air the bluish-white precipitate gradually absorbs oxygen and changes to blue Fe 4 (Fe(CN) 6 ) 3 , ferric ferrocyanide (Prussian blue) : 4K 2 Fe(Fe(CN) 6 ) + O 2 + 4HC1 = Fe 4 (Fe(CN) 6 ) 3 + K 4 Fe(CN) 6 + 4KC1 + 2H 2 O. 4. K 3 Fe(CN) 6 , potassium ferricyanide, precipitates dark- blue Fe 3 (Fe(CN) 6 ) 2 , ferrous ferricyanide (Turnbull's blue), insoluble in acids. 5. KCNS, potassium sulphocyanide, does not produce a claret-red coloration in solutions of ferrous salts free from ferric salts. (Distinction from ferric salts.) 6. Ferrous compounds and also ferric compounds when ignited with sodium carbonate on charcoal yield a black magnetic oxide. 7. All compounds of iron when fused in the oxidizing flame in a bead of borax yield while hot a yellow or reddish- 40 brown bead, which on cooling becomes lighter in color or colorless. Fused in the reducing flame the bead becomes bottle-green in color. BEHAVIOR OF IRON IN THE FERRIC CONDITION. Fe^Clgj ferric chloride, may be employed in making the tests. 1. NH 4 OH, NaOH, or KOH produces in solutions of ferric salts a voluminous, reddish-brown precipitate of Fe.j(OH) 6 , ferric hydroxide, insoluble in excess of the reagent and in ammonium salts. 2. (NHJgS precipitates black FeS together with free sul- phur : Fe 2 Cl 6 + 3(NH 4 ) 2 S 2FeS + S + 6NH 4 C1. 3. K 4 Fe(CN) 6 , potassium ferrocyanide, precipitates, even in exceedingly dilute solutions of ferric salts, blue Fe 4 (Fe(CN) 6 ) 3 , ferric ferrocyanide (Prussian blue), insoluble in acids, but decomposed by alkalies. 4. K 3 Fe(CN) 6 , potassium ferricyanide, does not produce a precipitate in ferric solutions, but imparts a green or brown coloration to the solution. (See 3, page 80.) 5. KCNS, potassium sulphocyanide, produces an intense claret-red coloration in ferric solutions, due to the formation of soluble Fe 2 (CNS) 6 , ferric sulphocyanide. In exceedingly dilute solutions the color is pale red. HgCl 2 , mercuric chlo- ride, destroys the coloration, soluble Hg(CNS) 2 being formed. 6. NaC 2 H 3 O 2 , sodium acetate, added to a ferric salt colors the solution dark red, due to the formation of Fe 2 (C 2 H 3 O 2 ) 6 , ferric acetate, which, on boiling the sufficiently diluted solu- tion, separates with part of the acetic acid as a brownish- red, flocculent precipitate of Fe 2 (OH) 4 (C 2 H 3 O 2 ) 2 , basic ferric acetate : A)e + 4H 2 = Fc 2 (OH)<(C 2 H 3 2 \ + 4HC 2 H 3 O 8 . 41 7. H 2 S, hydrogen sulphide, reduces ferric salts in solution to ferrous salts, with the separation of sulphur : Fe 2 Cl 6 + H 2 S = 2FeCl 2 + 2HC1 + S. 8. For the behavior of ferric salts on charcoal and in the borax bead see under Ferrous Salts, 7, page 39. ALUMINIUM, Al. Atomic weight, 27, 04; valence, IV. Tin-white metal ; specific gravity, 2.56 (spec. grav. of the hammered metal 2.67) ; melting-point, about 700 C. Al 2 (80 4 \j aluminium sulphate, or NH^Al(80^ ammonium aluminium sulphate (ammonia alum), may be employed in making the tests. 1 . NH 4 OH precipitates white, gelatinous A1 2 (OH) 6 , alumin- ium hydroxide, slightly soluble in an excess of the reagent. The precipitation is complete only when the excess of am- monia has been driven off by boiling the solution. 2. NaOH or KOH precipitates gelatinous A1 2 (OH) 6 , aluminium hydroxide, soluble in an excess of either reagent, with the formation of Al 2 O 2 (ONa) 2 , sodium aluminate, or A1 2 O 2 (OK) 2 potassium aluminate : A1 2 (OH) 6 -f 2NaOH = Al 2 O 2 (ONa) 2 + 4H 2 O. As aluminium hydroxide is insoluble in ammonium hydrox- ide (providing the latter is not present in great excess), the aluminium hydroxide may be reprecipitated from its solution as aluminate by the addition of ammonium chloride : Al 2 2 (ONa) 2 + 2NH 4 C1 + 2H 2 O= A1 2 (OH) 6 + 2NH 8 +2NaCl. Boiling does not decompose the aluminates. The solutions of aluminates have an alkaline reaction. 3. (NH 4 ) 2 S completely precipitates aluminium from its solu- tion as A1 2 (OH) 6 , aluminium hydroxide, with the evolution of hydrogen sulphide. 42 4. N^HPO^ sodium hydrogen phosphate, precipitates in neutral solutions white gelatinous A1 2 (PO 4 ) 2 , aluminium phos- phate, insoluble in acetic acid and in ammonium hydrox- ide, soluble in mineral acids and in sodium or potassium hydroxide, with the formation of alumi nates : A1 2 (PO 4 ) 2 + SNaOH = Al 2 O 2 (ONa) 2 + 2Na 3 PO 4 + 4H 2 O. Ammonium chloride reprecipitates the aluminium phosphate from its solution in sodium or potassium hydroxide : Al 2 O 2 (ONa) 2 -f- 2Xa 3 PO 4 + 8NH 4 C1 == A1 2 (PO 4 ) 2 + SNaCl + 8NH 3 + 4H 2 O. 5. Na^Og, sodium hyposulphite, added to a neutral solu- tion of a salt of aluminium precipitates white, gelatinous A1 2 (OH) 6 , with the separation of free sulphur and libera- tion of SO 2 , sulphurous anhydride. Complete precipitation takes place only when the solution of the aluminium salt is dilute and is boiled, after the addition of the hyposulphite, until the odor of sulphurous anhydride can no longer be detected : A1 2 C1 6 + SNaAOs + 3H 2 O = A1 2 (OH) 6 + 6NaCl + 3SO 2 6. Compounds of aluminium, mixed with sodium carbon- ate and ignited on charcoal, yield white, infusible aluminium oxide ; on moistening the mass with cobaltous nitrate and again igniting, an infusible blue residue is obtained, due to the combination of CoO, cobaltous oxide, with A1 2 O 3 , alumin- ium oxide (Thenard's blue). CHROMIUM, Cr. Atomic weight, 52.45; valence, II, IV. Light-gray, crystalline powder; specific gravity, 6.81. O 2 C7 6 , chromic chloride, may be employed in making the tests. 1. NH 4 OH precipitates bluish-gray, gelatinous Cr 2 (OH) 6 , 43 chromic hydroxide ; if the precipitation has taken place in a cold solution, a small quantity of chromic oxide will remain dissolved in the ammonium hydroxide ; on boiling this pink- ish solution all of the chromium is precipitated as chromic hydroxide. 2. XaOH or KOH precipitates from solutions of both the green and the violet salts of chromium greenish, flocculent Cr 2 (OH) 6 , chromic hydroxide, soluble in an excess of the reagent, forming Cr 2 O 2 (ONa) 2 , sodium chromite, and impart- ing a greenish color to the solution : O 2 (OH) 6 + 2NaOH = Cr 2 O 2 (ONa) 2 + 4H 2 O. From this solution chromic hydroxide is reprecipitated by the addition of ammonium chloride or by long-continued boiling : O 2 O a (ONa) a 4. 4H 2 O == Cr 2 (OH) 6 + 2NaOH. The precipitated chromic hydroxide obtained by boiling its alkaline solution appears to be insoluble in sodium or potas- sium hydroxide. It is probably a chromic hydroxide some- what deficient in water of hydration. The solubility of chromic hydroxide in sodium hydroxide is very much re- tarded by the presence of ferric oxide. 3. (NH 4 ) 2 S precipitates Cr 2 (OH) 6 , chromic hydroxide, with the evolution of hydrogen sulphide : Cr 2 Cl 6 + 3(NH 4 ) 2 S + 6H 2 O - Cr 2 (OH) 6 + 6NH 4 C1 + 3H 2 S. 4. A salt of chromium, fused on platinum foil with a mixture of sodium carbonate and potassium nitrate or potas- sium chlorate, yields a mass containing a salt of chromic acid, i.e., a chromate ; on exhausting the mass with water a yellow solution of Na 2 CrO 4 , sodium chromate, or K 2 CrO 4 , potassium chromate, is obtained, which when treated with plumbic acetate yields a yellow precipitate of PbCrO 4 , plumbic chromate : O 2 O 3 + 2Na 2 CO 3 + O 3 = 2Na 2 CrO 4 + 2CO 2 ; Pb(C 2 H 3 O 2 ) 2 = PbCrO 4 + 2NaC 2 H 3 O 2 . 44 5. Compounds of chromium when ignited with sodium car- bonate on charcoal yield a green fused mass containing oxides of chromium. 6. Fused in a bead of borax or of microcosmic salt, in either the oxidizing or the reducing flame, chromium com- pounds yield a yellowish-green bead, which becomes emerald- green on cooling. FOURTH GROUP. Metals precipitated as sulphides from neutral solutions by (NH 4 ) 2 S, ammonium sulphide : Manganese, Zinc, Cobalt, and Nickel. MANGANESE, Mn. Atomic weight, 54.8; valence, II, IV. Grayish- white metal ; specific gravity, about 8. MnSO, manganous sulphate, may be employed in making the tests. 1. (NH 4 ) 2 S precipitates pale-salmon-colored MnS, manga- nous sulphide, containing water, easily soluble in acetic acid and in hydrochloric acid. (Occasionally, especially after standing some time, the pale-salmon-colored precipitate con- taining water is converted into green MnS, manganous sul- phide, which is free from water.) Manganous sulphide readily oxidizes on exposure to the air and becomes dark brown, due to the formation of MnO(OH) 2 , hydrated peroxide of manganese. 2. NaOH or KOH precipitates white Mn(OH) 2 , manganous hydroxide, insoluble in excess of the reagent. On exposure to the air the precipitate rapidly becomes brown, due to the 45 formation of Mn 2 (OH) 6 , manganic hydroxide. Manganous hydroxide is soluble in ammonium chloride, owing to the production of a double salt, whereas manganic hydroxide is insoluble in that reagent ; on this account ammonium chloride solutions of manganous hydroxide containing free ammonia become brown on exposure to the air, due to the separation of manganic hydroxide : Mn(OH) 2 -f 4NH 4 C1 = MnCl 2 (NH 4 Cl) 2 -j- 2H 2 O + 2NH 3 ; 2MnCl 2 (NH 4 Cl) 2 + 4NH 3 + 5H 2 O + O = Mn 2 (OH) 6 + 8NH 4 C1. 3. NH 4 OH precipitates in neutral solutions, and also in solutions free from salts of ammonium, white Mn(OH) 2 , manganous hydroxide; in the presence of salts of ammonium or of free acids, excess of ammonium hydroxide fails to pro- duce a precipitate, because of the formation of a soluble double salt of manganous hydroxide with the ammonium salts. The action of the oxygen of the air converts the solu- ble manganous salt into Mn 2 (OH) 6 , manganic hydroxide, which separates as a brown precipitate. 4. Compounds of manganese, fused on platinum foil with sodium carbonate and potassium nitrate, yield a bluish-green mass containing manganates of sodium and potassium : 3Mn(OH) 2 -f- Na 2 CO 3 + 4KNO 3 = Na 2 MnO 4 + 2K 2 MnO 4 + 4NO + CO 2 + 3H 2 O. The test is an exceedingly delicate one, and only a minute quantity of a salt of manganese need be used. On exhausting the mass with water, soluble KMnO 4 , potassium permanganate, and insoluble brown MnO(OH) 9 , hydrated peroxide of manganese, are formed : 3K 2 MnO 4 + 3H 2 O = 2KMnO 4 -f MnO(OH) 2 + 4KOH. The potassium permanganate dissolves, imparting a purplish- red color to the water. 5. Compounds of manganese, fused in the oxidizing flame 46 in a bead of borax or of microcosmic salt, yield an amethyst- colored bead ; fused in the reducing flame, the bead becomes colorless. ZINC, Zn. Atomic weight, 64.88; valence, II. Bluish-white metal ; specific gravity, 6.9 ; melting-point, 433 C. ZnS0 4 , zinc sulphate, may be employed in making the tests. 1. (NH 4 ) 2 S precipitates white ZnS, zinc sulphide, easily soluble in hydrochloric acid, insoluble in acetic acid. 2. NaOH or KOH precipitates white, gelatinous Zn(OH) 2 , zinc hydroxide, soluble in excess of the reagent, with the formation of Zn(ONa) 2 , sodium zincate, or Zn(OK, 2 , potas- sium zincate : Zn(OH) 2 -f 2NaOH = Zn(ONa) 2 + 2H 2 O. These solutions, which have an alkaline reaction, yield a precipitate of ZnS, zinc sulphide, on the addition of hydrogen sulphide. 3. NH 4 OH precipitates in neutral solutions white, floccu- lent Zn(OH) 2 , zinc hydroxide, soluble in excess of the reagent. 4. K 4 Fe(CN) 6 , potassium ferrocyanide, produces a white, flocculent precipitate of Zn 2 Fe(CN) 6 , zinc ferrocyanide, in- soluble in acids and in ammonium hydroxide. The precipi- tate while in suspension often has a pale-yellowish appear- ance, due to the color imparted to the liquid by the presence of an excess of potassium ferrocyanide. 5. Compounds of zinc, ignited with sodium carbonate in the reducing flame on charcoal, yield a coating of ZnO, zinc oxide, which is yellow when hot and white when cold. On moistening the deposit with cobaltous nitrate and again ignit- ing, the deposit becomes green in color. 47 COBALT, Co. Atomic weight, 58.6; valence, II, IV. Steel-gray metal ; specific gravity, 8.6. Co(N0 3 ) 2 , cobaltous nitrate, or CoCl 2 , cobaltous chloride, may be employed in making the tests. 1. (NH 4 ) 2 S precipitates black CoS, cobalt sulphide, insolu- ble in excess of colorless ammonium sulphide and in dilute hydrochloric acid; soluble in nitro-hydrochloric acid, with the formation of CoCl 2 , cobaltous chloride : 3CoS + 6HC1 -j- 2HNO 3 = 3CoCl 2 + S 3 + 2NO + 4H 2 O. 2. NaOH or KOH precipitates in cold cobaltous solu- tions a bluish basic salt, and in boiling solutions rose-red Co(OH) 2 , cobaltous hydroxide. Both precipitates become oxidized on exposure to the air and turn olive-green in color. They are insoluble in excess of the reagent. 3. NH 4 OH precipitates in cold cobaltous solutions a bluish basic salt, and in boiling solutions rose-red Co(OH) 2 , cobalt- ous hydroxide. Both of these precipitates are soluble in an excess of the reagent, imparting a reddish color to the liquid, which, on exposure to the oxidizing action of the air, soon changes to brown. 4. KCN, potassium cyanide, precipitates brownish-white Co(CN) 2 , cobaltous cyanide, soluble in excess of the reagent, with the formation of (KCN) 2 Co(CN) 2 , potassium cobaltous cyanide, from which solution cobaltous cyanide is reprecipi- tated by hydrochloric acid. 5. KNO 2 , potassium nitrite, added in excess to a somewhat concentrated solution of a salt of cobalt, to which sufficient acetic acid has previously been added, produces a yellow', crystalline precipitate of K 6 Co 2 (]S"O 2 ) 12 , potassium cobaltic nitrite = (KNO 2 ) 6 Co 2 (NO 2 ) 6 : 48 2CoCl 2 + 14KN0 2 + 4HC 2 H 3 2 - K 6 Co 2 (NO 2 ) 12 + 4KC 2 H 3 O 2 + 4KC1 + 2NO + 2H 2 O. In concentrated cobalt solutions the precipitate appears im- mediately, while in dilute solutions it requires some time for it to form. The presence of free acetic acid is necessary to liberate the nitrous acid (required in the oxidation) from the potassium nitrite. Free hydrochloric acid must not be present ; in case of its presence in the solution, it must be neutralized by the addition of NaC 2 H 3 O 2 , sodium acetate, previous to the addi- tion of the acetic acid. To insure complete precipitation of the cobalt, particularly in the case of dilute solutions, the so- lution should be allowed to stand in a warm place for about twenty-four hours. 6. Compounds of cobalt, ignited -with sodium carbonate in the reducing flame on charcoal, yield dark, metallic, magnetic spangles. 7. Compounds of cobalt, fused in a bead of borax or of microcosmic salt in either the reducing or the oxidizing flame, impart to the bead a beautiful sapphire-blue color. NICKEL, Ni. Atomic weight, 58.6; valence, II, IV. Silver- white metal ; specific gravity, 8.9. NiSOt, nickelous sulphate, or NiCl 2 , nickelous chloride, may be employed in making the tests. 1. (NH 4 ) 2 S precipitates black NiS, nickelous sulphide, sol- uble in excess of ammonium sulphide (particularly in the presence of ammonia), imparting a brownish color to the solution. On boiling the ammonium sulphide solution of nickelous sulphide, it undergoes decomposition (particularly 49 after the addition of acetic acid), with the separation of the dissolved nickclous sulphide. The precipitate is insoluble in dilute hydrochloric acid, but soluble in nitre-hydrochloric acid. 2. NaOH or KOH precipitates amorphous, apple-green Ni(OH)s, nickelous hydroxide, insoluble in an excess of the reagent. 3. NH 4 OH, added in small quantity, precipitates, in neu- tral solutions of nickelous salts free from ammonium salts, apple-green Ni(OH) 2 , nickelous hydroxide, soluble in excess of ammonium hydroxide, imparting a bluish color to the solution. 4. KCN, potassium cyanide, precipitates light-green Ni(ON)j, nickelous cyanide, soluble in excess of the re- agent, with the formation of (KCN) 2 Ni(CN) 2 , potassium nickelous cyanide. Hydrochloric acid reprecipitates from this solution Ni(CN) 2 , nickelous cyanide. (Distinction from co- balt.) 5. KNO 2 , potassium nitrite (under the conditions given for cobalt, 5, page 47), fails to produce a precipitate in solu- tions of nickelous compounds. 6. Compounds of nickel, ignited with sodium carbonate on charcoal, yield dark, magnetic, metallic spangles. 7. Compounds of nickel, fused in the oxidizing flame in a bead of borax, yield a bead which is purplish red while hot and pale brownish yellow when cold. In the reducing flame the bead becomes gray and opaque, due to the separation of metallic nickel. Fused in a bead of microcosmic salt in the oxidizing or the reducing flame, salts of nickel yield a reddish-brown bead which becomes yellow or yellowish red on cooling. 50 FIFTH GROUP. Metals precipitated as carbonates from neutral solutions by (NH 4 ) 2 CO 3 , ammonium carbonate : Barium, Strontium, and Calcium. Complete precipitation does not take place in solutions which were originally acid, or when ordinary commercial am- monium carbonate is employed, unless the solution is boiled after the addition of the ammonium carbonate. Commer- cial ammonium carbonate consists of equal molecules of NH 4 HCO 3 , acid ammonium carbonate, and NH 4 NH 2 COO, ammonium carbamate. (1) Dissolving the commercial carbonate in water converts the ammonium carbamate into neutral ammonium carbonate : NH 4 HC0 3 + NH 4 NH 2 COO + H 2 O = NH 4 HCO 3 + (NH 4 ) 2 C0 3 . In precipitating with ammonium carbonate containing acid ammonium carbonate, part of the precipitate will consist of acid salts, for example, J>a(HCO 3 ) 2 , which are converted into neutral salts on boiling : Ba(HCO 3 ) 2 = BaCO 3 + CO 2 -f H 2 O. BARIUM, Ba. Atomic weight, 136.86; valence, II. Silver-white metal ; specific gravity, about 4.0. BaCl 2j barium chloride, may be employed in making the tests. 1. (NH 4 ) 2 CO 3 , ammonium carbonate, precipitates white, 1 According to other views, commercial ammonium carbonate consists of one molecule of neutral and two molecules of acid ammonium carbon- ate, thus : (NH 4 ) 2 CO 3 + 2NH 4 HCO 3 51 flocculent BaCO 3 , barium carbonate. The precipitate is easily soluble in dilute hydrochloric acid, in nitric acid, and in acetic acid, insoluble in pure water, slightly soluble in ammonium chloride, and, like all the carbonates of the alka- line earths, soluble in water containing carbonic acid. 2. H 2 SO 4 , sulphuric acid, and soluble sulphates, including solutions of calcium and strontium sulphates, precipitate white, finely-pulverulent BaSO 4 , barium sulphate, insoluble in acids. If the precipitation occur in a cold solution, the particles of the precipitate are so minute that they readily pass through a filter ; whereas, if the precipitation take place in a hot solution, the precipitate that is formed is crystalline and readily retained by a filter. 3. (NH 4 ) 2 C 2 O 4 , ammonium oxalate, precipitates white, pul- verulent BaC 2 O 4 , barium oxalate, which when freshly precip- itated is soluble in acetic acid and in H 2 C 2 O 4 , oxalic acid. 4. Na 2 HPO 4 , sodium hydrogen phosphate, precipitates white, flocculent BaHPO 4 , di-basic barium phosphate, soluble in hydrochloric, nitric, and acetic acids. 5. K 2 CrO 4 , potassium chromate, produces in neutral or acetic acid solutions of salts of barium yellow BaCrO 4 , barium chromate, soluble in hydrochloric acid and in nitric acid. 6. Compounds of barium, held in the flame of a Bunsen burner by means of a platinum wire, impart a yellowish- green color to the flame. STRONTIUM, Sr. Atomic weight, 87.3; valence, II. Yellowish metal ; specific gravity, 2.5. $r(JV0 3 ) 2 , strontium nitrate, may be employed in making the 1. (NH 4 ) 2 CO 3 , ammonium carbonate, precipitates white 52 SrCO 3 , strontium carbonate, easily soluble in dilute hydro- chloric acid, in nitric acid, and in acetic acid. 2. H 2 SO 4 , sulphuric acid, and soluble sulphates, including calcium sulphate, precipitate white, usually crystalline SrSO 4 , strontium sulphate, insoluble in alcohol. In dilute solutions, and also on using calcium sulphate as the precipitating re- agent, the precipitation takes place gradually. 3. (NH 4 ) 2 C 2 O 4 , ammonium oxalate, precipitates white, pulverulent SrC 2 O 4 , strontium oxalate, soluble Avith difficulty in acetic acid and in oxalic acid. 4. Na^HPO^ sodium phosphate, precipitates white SrHPO 4 , di-basic strontium phosphate, soluble in hydro- chloric, nitric, and acetic acids. 5. K 2 CrO 4 , potassium chromate, does not produce a pre- cipitate with salts of strontium. 6. Compounds of strontium impart a crimson color to the flame. CALCIUM, Ca. Atomic weight, 39.9; valence, II. Yellowish metal ; specific gravity, 1.57. CdCl 2j calcium chloride, may be employed in making the tests. 1. (NH 4 ) 2 CO 3 , ammonium carbonate, precipitates white CaCO 3 , calcium carbonate, easily soluble in dilute hydro- chloric acid, in nitric acid, and in acetic acid. 2. H 2 SO 4 and soluble sulphates precipitate immediately, in concentrated solutions of salts of calcium, white, crystalline CaSO 4 , calcium sulphate, insoluble in alcohol, but soluble in boiling hydrochloric acid. Precipitation takes place in dilute solutions either gradually or not at all. 3. (NH 4 ) 2 C 2 O 4 , ammonium oxalate, precipitates white, pul- 53 verulent CaC 2 O 4 , calcium oxalate, easily soluble in hydro- chloric or in nitric acid, insoluble in acetic and oxalic acids. 4. Na 2 HPO 4 , sodium hydrogen phosphate, precipitates white CaHPO 4 , di-basic calcium phosphate, soluble in hydro- chloric, nitric, and acetic acids. 5. K 2 CrO 4 , potassium chromate,. does not produce a pre- cipitate in solutions of calcium salts. 6. Compounds of calcium impart a yellowish-red color to the flame. SIXTH GROUP. Bases not precipitated by any particular group reagent: Magnesium, Potassium, Sodium, Ammonium, and Lithium. MAGNESIUM, Mg. Atomic weight, 23.94; valence- II. Silver- white metal ; specific gravity, 1.75. MgSOu magnesium sulphate, may be employed in making the tests. 1. NH 4 OH precipitates, in neutral solutions of salts of magnesium, part of the magnesium as flocculent Mg(OH) 2 , magnesium hydroxide, leaving the other part in solution as a double salt of magnesium and ammonium : 2MgSO 4 + 2NH 4 OH = Mg(OH) 2 + MgSO 4 (NH 4 ) 2 SO 4 . This double salt is not decomposed by a slight excess of ammonium hydroxide. Compounds of magnesium are not precipitated by ammonium hydroxide in the presence of an excess of ammonium chloride, the latter reagent having the property of dissolving magnesium hydroxide : Mg(OH) 2 + 4NH 4 C1 = MgCl 2 (NH 4 Cl) 2 + 2NH 4 OH. 54 . 2. NaOH or KOH precipitates, particularly on boiling, white Mg(OH) 2 , magnesium hydroxide. 3. NagCOg, sodium carbonate, or K 2 CO 3 , potassium car- bonate, precipitates Mg 4 (CO 3 ) 3 (OH) 2 , basic magnesium car- bonate. (The carbonic acid liberated in the reaction retains part of the magnesium in solution as an acid carbonate ; this is precipitated by boiling the solution.) The precipitate is soluble in ammonium chloride. 4. (NH 4 ) 2 CO 3 produces no precipitate immediately, but after standing some time a crystalline precipitate of MgCO 3 (NH 4 ) 2 CO 3 appears. In the presence of a sufficient quantity of ammonium chloride the precipitation does not take place. o. Na 2 HPO 4 produces in concentrated solutions a white, flocculent precipitate of MgHPO 4 , di-basic magnesium phos- phate. If ammonium chloride and ammonium hydroxide are added to the solution of the magnesium salt, and after- wards sodium hydrogen phosphate added, a white, crystalline precipitate of MgNH 4 PO 4 , ammonium magnesium phosphate, is produced : MgSO 4 -f Na 2 HPO 4 + NH 4 OH = MgNH 4 PO 4 + Na 2 SO 4 + H 2 0. The ammonium chloride is added to the solution in order to prevent the precipitation of the magnesium salt by the am- monium hydroxide. The precipitate is always crystalline; in dilute solutions it forms gradually, the formation is facili- tated, however, by gently rubbing the inner sides of the vessel with a glass rod. 6. Compounds of magnesium ignited on charcoal are some- what luminous in the flame. On moistening the mass with cobaltous nitrate and again strongly igniting, a pale-pink color, which is more evident on cooling, is imparted to the mass. 55 POTASSIUM, K (KALIUM). Atomic weight, 39.O3 ; valence, I. Silver- white metal ; specific gravity, 0.87 ; melting-point, 62.5 C. KN0 3 , potassium nitrate, or KCl, potassium chloride, may be employed in making the tests. 1. PtCl 4 , platinic chloride, precipitates from neutral or acid solutions yellow, crystalline (KCl) 2 PtCl 4 , potassium platinic chloride, slightly soluble in water, insoluble in alcohol. The test is best made in a watch-glass, and the liquid should be stirred with a glass rod. In dilute solutions the precipitate forms slowly. The addition of a little alcohol and, if the potassium salt is not a chloride, a drop of hydrochloric acid facilitates the precipitation. 2. NaHC 4 H 4 O 6 , acid sodium tartrate, produces, in rather concentrated neutral solutions of salts of potassium, a white, granular, crystalline precipitate of KHC 4 H 4 O 6 , acid potas- sium tartrate. Stirring the liquid with a glass rod or the addition of alcohol promotes the precipitation. If the po- tassium solution has an alkaline reaction, it must be neutral- ized with acetic acid previous to the addition of the acid so- dium tartrate. H 2 C 4 H 4 O 6 , tartaric acid, may be used instead of acid sodium tartrate, but in using it NaC 2 H 3 O 2 , sodium acetate, must be added to the solution : KN0 3 + H 2 C 4 H 4 6 + NaC 2 H 3 2 = KHC 4 H 4 O 6 + NaNO s + HC 2 H 3 O 2 . In dilute solutions of potassium salts precipitation occurs only after standing some time. 3. Compounds of potassium impart to the flame a violet color, 56 SODIUM, Na (NATRIUM). Atomic weight, 22.99; valence, I. Silver-white metal; specific gravity, 0.97; melting-point, 95.6 C. NaCly sodium chloride, may be employed in making the tests. 1. K 2 H 2 Sb 2 O 7 , potassium pyroantimonate, produces, in neutral or slightly alkaline concentrated solutions of salts of sodium, a white, crystalline precipitate of Na 2 H 2 Sb 2 O 7 , sodium pyroantimonate. In dilute solutions the precipitate forms only after the liquid has been standing some time. Gently rubbing the inner sides of the vessel facilitates the formation of the precipitate. If the sodium solution have an acid reaction, it must be neutralized with potassium car- bonate before the addition of the potassium pyroantimonate. Metals other than sodium or potassium must not be present, as they interfere by forming insohible antimonates. 2. PtCl 4 , platinic chloride, as well as H 2 C 4 H 4 O 6 , tartaric acid, fails to produce a precipitate in solutions of salts of sodium. 3. Compounds of sodium impart to the flame an intense yellow color. AMMONIUM, NH 4 . The radical NH 4 in its behavior with acid radicals corre- sponds to potassium and sodium. Another analogy between the radical ammonium and the metals is the existence of an ammonium amalgam. NHCl, ammonium chloride, may be employed in making the tests. 1. The ammonium salts (in combination with volatile 57 acids) are characterized by their complete volatility when heated to high temperatures ; ammonium borate and ammo- nium phosphate, however, on being strongly heated leave a residue respectively of boric acid and of phosphoric acid. 2. PtCl 4 , platinic chloride, precipitates in concentrated so- lutions yellow, crystalline (NH 4 Cl) 2 PtCl 4 , ammonium platinic chloride, slightly soluble in water, insoluble in alcohol. The test is best made in a watch-glass, stirring the liquid with a glass rod. The addition of a little alcohol and, if the am- monium salt is not a chloride, a drop of hydrochloric acid hastens the formation of the precipitate. 3. NaOH or KOH added to a solution of a salt of am- monium liberates ammoniacal gas, on boiling the solution, which may be detected by its odor ; by its producing white clouds of ammonium acetate when a glass rod wet with acetic acid is held above the liquid ; by its action upon turmeric paper moistened with water, which becomes brown when held above the liquid in which the liberation has occurred ; and by its action upon filter paper moistened with mercurous nitrate, which becomes black when held in the evolved gas (due to the formation of black NH 2 Hg 2 NO 3 ). 4. H 2 C 4 H 4 O 6 , tartaric acid, or NaHC 4 H 4 O 6 , acid sodium tartrate, produces in concentrated solutions of ammonium salts white, crystalline NH 4 HC 4 H 4 O 6 , acid ammonium tartrate. (For conditions favoring precipitation see Potassium, 2, page 55.) LITHIUM, Li. Atomic weight, 7.O; valence, I. Silver-white metal ; specific gravity, 0.59 ; melting-point, 180 C. LiCl, lithium chloride, may be employed in making the tests. 58 1. NajCOg precipitates, in cold concentrated solutions of salts of lithium, white Li 2 CO 3 , lithium carbonate. 2. NaJHPC^ added to a solution of a salt of lithium pro- duces a white, crystalline precipitate of Li 3 PO 4 , lithium phos- phate. 3. Compounds of lithium impart to the flame a carmine- red color. II. PROPERTIES OF THE ACIDS. FIRST GROUP. ACIDS which are precipitated by BaCl 2 , barium chloride, from neutral and from acid solutions : Sulphuric Acid, Hydrofluosilicic Acid. SULPHURIC ACID, H 2 SO 4 . (Sulphuric acid combines with bases to form salts called sulphates.) MgSOt, magnesium sulphate, may be employed in making the tests. 1. The neutral sulphates, with the exception of barium, strontium, calcium, and lead sulphates, are easily soluble in water. Basic sulphates of the heavy metals are soluble in hydrochloric acid or in nitric acid. Lead sulphate and the sulphates of the alkaline earths are decomposed and con- verted into carbonates by sodium or potassium carbonate. 2. BaCl 2 , barium chloride, precipitates, from solutions con- taining sulphates or free sulphuric acid, white, pulverulent BaSO 4 , insoluble in acids. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white PbSO 4 , plumbic sulphate, insoluble in dilute acetic acid, somewhat soluble in boiling concentrated acids. Plumbic sulphate is easily soluble in (NH 4 ) 2 C 4 H 4 O 6 , ammonium tartrate ; from this solution potassium chromate precipitates the lead as yellow PbCrO 4 , plumbic chromate. 4. Sulphates, fused with sodium carbonate on charcoal, 59 60 yield a residue containing !N"a 2 S, sodium sulphide. On placing a portion of the mass on a clean silver coin and adding a few drops of water, a brownish or black stain of Ag 2 S is produced : C 2 == NaJ3 -f 2CO 2 ; -f Ag 2 -f 2H 2 O = Ag 2 S -f 2NaOH + H 2 . HYDROFLUOSILICIC ACID, H.,SiF 6 . (Hydrofluosilicic acid combines with bases to form salts called silicofluorides.) iFfi, sodium silicofluoride, may be employed in making the tests. 1 . Most of the silicofluorides are soluble in water ; when gently heated with concentrated sulphuric acid, they evolve gaseous SiF 4 , silicon fluoride, and HF, hydrofluoric acid : K 2 SiF 6 -f H 2 S0 4 = SiF 4 + 2HF + K 2 SO 4 . If a piece of platinum foil containing a drop of water be inverted over the vessel in which the decomposition is effected, the water becomes milky in appearance, due to the formation of insoluble H 2 SiO 3 , silicic acid : 3H 2 O + 3SiF 4 = H 2 SiO 3 -f 2H 2 SiF 6 . 2. BaCl 2 , barium chloride, precipitates, in solutions of hydrofluosilicic acid and of silicofluorides, crystalline BaSiF 6 , barium silicofluoride, insoluble in dilute acids. 3. KNO 3 , potassium nitrate, precipitates, in solutions that are not too dilute, translucent, gelatinous K 2 SiF 6 , potassium silicofluoride, soluble with difficulty in water, insoluble in alcohol. 4. NH 4 OH produces NH 4 F, ammonium fluoride, and H 2 SiO 3 , silicic acid, both of which are precipitated : 6NH 4 OH -f H 2 SiF 6 = H 2 SiO 3 + 6NH 4 F 4- 3H 2 O. 61 SECOND GROUP. Acids which are precipitated by BaCl 2 , barium chloride, in neutral solutions, the barium salts of which are soluble in hydrochloric acid : Sulphurous Acid, Hypo sulphurous Acid, Phosphoric Acid, Boric Acid, Hydrofluoric Acid, Carbonic Acid, Silicic Acid, Chromic Acid, Arsenic Acid, Arsenious Acid. SULPHUROUS ACID, H 2 SO 3 . (Sulphurous acid combines with bases to form salts called sulphites.) Na 2 SO B , sodium sulphite, may be employed in making the tests. 1. Of the neutral sulphites only those of the alkalies are soluble in water ; the others are easily soluble in acids, with the evolution of SO 2 , sulphurous anhydride : BaSO 3 + 2HC1 = BaCl 2 + SO 2 + H 2 O. 2. Dilute acids decompose sulphites, with the evolution of SO 2 , sulphurous anhydride, which may be recognized by its odor (that of burning sulphur). The presence of sulphurous anhydride may be detected in gaseous mixtures by its be- havior with KIO 3 , potassium iodate. A piece of filter paper saturated with a solution of potassium iodate and starch paste, brought while wet in contact with gaseous mixtures contain- ing sulphurous anhydride, becomes blue in color, owing to the reduction of HIO 3 , iodic acid, to iodine, and the action of the latter on the starch. By means of the sulphurous anhydride, in the presence of water, the iodic acid is reduced to HI, hydriodic acid : 3SO 2 + 3H 2 O + HIO 3 = HI + 3H 2 SO 4 . The hydriodic acid, by the action of the remaining iodic acid, is reduced, with the liberation of free iodine : 5HI-f HI0 3 = I 6 + 3H 2 0. 6 62 (As the free iodine is reconverted by an excess of sulphurous anhydride into hydriodic acid : I 2 -f SO 2 + 2H 2 O = 2HI -f- H 2 SO 4 , an excess of the sulphurous anhydride causes a disappearance of the color.) 3. BaCl 2 precipitates white BaSO 3 , barium sulphite, soluble in acids. 4. Pb(C 2 H 3 O 2 ) 2 precipitates white PbSO 3 , plumbic sulphite, soluble in nitric acid. 5. AgNO 3 , argentic nitrate, precipitates white Ag 2 SO 3 , argentic sulphite, soluble in nitric acid. On boiling the precipitate with w^ater, it is decomposed into metallic silver and sulphuric acid, the liquid becoming gray in color, due to the separated metallic silver : Ag 2 S0 3 + H 2 = Ag 2 -f- H 2 S0 4 . 6. ZnSO 4 , zinc sulphate solution, containing a little Na 2 NOFe(CN) 5 , sodium nitroprusside, added to a solution of a sulphite which, if not neutral, has been neutralized with acetic acid, produces a red coloration ; or a flocculent, purplish-red precipitate, if the solution contain a considerable quantity of the sulphite. When operating with dilute solu- tions of a sulphite, the test may be .made more delicate by the addition of a few drops of potassium ferrocyanide solu- tion. (Distinction from hyposulphites.) 7. H 2 S conducted into a solution of sulphurous acid de- composes the latter, with the separation of sulphur and the probable formation of H 2 S 5 O 6 , pentathionic acid : 5S0 2 + 5H 2 S = H 2 S 5 6 -f S 5 + 4H 2 O. 8. Sulphites ignited with sodium carbonate on charcoal yield a yellowish residue containing sodium sulphide, as in the case of sulphates. A portion of the residue placed on a clean silver coin and moistened with a few drops of water pro- duces a brown or black stain of argentic sulphide on the coin. 63 HYPOSULPHUROUS ACID, H 2 S 2 O 3 (THIOSULPHURIC ACID). (Hyposulphurous or thiosulphuric acid combines with bases to form salts called hyposulphites or thiosulphites.) Na 2 S 2 O s , sodium hyposulphite, may be employed in making the tests. 1. Most of the hyposulphites (thiosulphites) are soluble in water. 2. HC1 or H 2 SO 4 added to a solution of a hyposulphite liberates H 2 S 2 O 3 , hyposulphurous acid, which quickly breaks up into sulphur, sulphurous anhydride, and water : Na 2 S 2 O 3 + 2HC1 =-- H 2 S 2 O 3 + 2NaCl ; Thus hyposulphites, on the addition of either of the above acids, are decomposed and yield sulphurous anhydride, which may be recognized by its odor, and free sulphur. 3. BaCl 2 , barium chloride, produces in concentrated solu- tions of hyposulphites a white precipitate of BaS 2 O 3 , barium hyposulphite, soluble in a large quantity of water. It is also soluble in hydrochloric acid, with the evolution of sulphurous anhydride and the separation of sulphur. 4. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white PbS 2 O 3 , plumbic hyposulphite, soluble in nitric acid. 5. AgNO 3 , argentic nitrate, precipitates white Ag 2 S 2 O 3 , argentic hyposulphite, soluble in excess of sodium hyposul- phite solution : Ag 2 S 2 3 + Na 2 S 2 3 = 2NaAgS 2 3 . The precipitate becomes almost immediately yellow, then brown, and finally black, due to the formation of argentic sulphide : Ag 2 S 2 O 3 + H 2 O = Ag 2 S + H 2 SO 4 . 64 6. Fe 2 Cl 6 , ferric chloride, immediately colors hyposulphite solutions reddish violet. (Distinction from sulphites.) 7. Hyposulphites ignited with sodium carbonate on char- coal yield a residue containing sodium sulphide, as in the case of sulphates and of sulphites. A portion of the residue placed on a clean silver coin and moistened with water pro- duces a brown or black stain of argentic sulphide. PHOSPHORIC ACID, H 3 PO 4 . (Phosphoric acid combines with bases to form salts called phosphates.) Na 2 HPO^ sodium hydrogen phosphate, may be employed in making the tests. 1. The phosphates of the alkalies are soluble in water, the others are soluble in acids. 2. BaCl 2 , barium chloride, precipitates in solutions of the neutral phosphates white BaHPO 4 or Ba 3 (PO 4 ) 2 , soluble in hydrochloric acid or in nitric acid. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white Pb 3 (PO 4 ) 2 , soluble in nitric acid. 4. AgNO 3 , argentic nitrate, produces in solutions of the phosphates a light-yellow precipitate of Ag 3 (PO 4 ), soluble in nitric acid and in ammonium hydroxide. 5. NH 4 C1, ammonium chloride, NH 4 OH, ammonium hy- droxide, and MgSO 4 , magnesium sulphate, (l) added in turn to a solution of a phosphate, produce a white, crystalline pre- cipitate of MgNH 4 PO 4 , ammonium magnesium phosphate : Na 2 HPO 4 + NH 4 OH + MgSO 4 = MgNH 4 PO 4 + (The ammonium chloride is added to prevent the precipitation of the magnesium as magnesium hydroxide by the ammo- 1 The three reagents composing the so-called "magnesia mixture." 65 nium hydroxide.) The precipitate is sparingly soluble in pure water, and very slightly soluble in water containing ammonium hydroxide. In precipitating very dilute solu- tions the precipitate forms more rapidly when the inner sides of the vessel are gently rubbed with a glass rod. 6. NH 4 HMoO 4 , ammonium molybdate, added in excess, with a considerable quantity of nitric acid, to a solution of phosphoric acid or a phosphate, produces a yellow precipitate, probably of (NH 4 ) 3 PO 4 (MoO 3 ) 10 , ammonium phosphomolyb- date : 10NH 4 HMoO 4 + H 3 PO 4 + 7HNO 3 == (NH 4 ) 3 PO 4 (MoO 3 ) 10 + 7NH 4 N0 3 + 10H 2 0. The precipitate is insoluble in dilute nitric acid, but easily soluble in ammonium hydroxide ; it is reprecipitated from the ammoniacal solution by the addition of excess of nitric acid. The precipitate is also soluble in excess of a phosphate, and thus is explained the non-appearance of a precipitate when only a little ammonium molybdate is added to a solu- tion containing much phosphoric acid. In dilute solutions the precipitate forms slowly. The precipitation is hastened by warming the solution to a temperature of 40 C. A higher temperature should be avoided. (Pyrophosphates yield with argentic nitrate white precip- itates of Ag 4 P 2 O 7 , argentic pyrophosphate. Metaphosphates likewise yield white precipitates of AgPO 3 . Only the meta- phosphates coagulate albumen.) BORIC ACID, H 3 B0 3 . (Boric acid combines with bases to form salts called borates.) JVa 2 I> 4 7 , sodium biborate (borax), may be employed in making the tests. 66 1. Of the borates those of the alkalies are easily soluble in water. 2. BaCl 2 , barium chloride, produces in concentrated solu- tions of borates white Ba(BO 2 ) 2 , barium metaborate, easily soluble in excess of barium chloride and in ammonium chlo- ride : 2BaCl 2 + H 2 O = 2Ba(BO 2 ) 2 + 2NaCl + 2HC1. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates in concen- trated solutions white Pb(BO 2 ) 2 , 4japmm. metaborate, soluble in excess of the reagent. 4. AgNO 3 , argentic nitrate, precipitates in concentrated solutions of neutral borates white AgBO 2 , argentic metaborate, which is occasionally tinged with yellow, due to the presence of argentic oxide. In solutions of acid borates the precipitate is Ag 6 B 8 O 15 . Both precipitates are easily soluble in nitric acid. 5. Boric acid, in the powdered condition, placed in a porce- lain dish and covered with alcohol gives a greenish flame on igniting the alcohol. Borates, in the powdered condition, impart the same greenish color to the flame, but must be moistened with a few drops of concentrated sulphuric acid before the addition of the alcohol. The greenish color im- parted to the flame is due to (C 2 H 5 ) 3 BO 3 , the ethyl ester of boric acid, formed in the reaction. Boric acid (without the addition of alcohol) when strongly heated on platinum wire imparts a greenish color to the flame. (Compounds of barium and of copper and compounds containing chlorine interfere with the test.) 6. Turmeric paper dipped in an aqueous solution of boric acid, or in a solution of a borate acidified with hydrochloric acid, and warmed until dry, becomes reddish brown in color. On bringing dilute sodium or potassium hydroxide solution in contact with the reddish-brown paper, the color becomes blue and then greenish black. 67 HYDROFLUORIC ACID, HF. (Hydrofluoric acid combines with bases to form salts called fluorides.) KFj potassium fluoride, or NaF, sodium fluoride, may be employed in making the tests. 1 . Of the fluorides those of the alkalies are easily soluble in water, the others are soluble only with great difficulty. 2. BaCl 2 precipitates from solutions of fluorides white BaF 2 , soluble in hydrochloric acid ; Pb(C 2 H 3 O 2 ) 2 precipitates white PbF 2 , easily soluble in nitric acid ; and AgNO 3 pre- cipitates white AgF, also easily soluble in water. 3. CaCl 2 , calcium chloride, precipitates white, gelatinous CaF 2 , calcium fluoride, almost insoluble in water, soluble with difficulty in mineral acids. 4. Hydrofluoric acid has the property of etching glass, forming with the silicic oxide of the glass volatile SiF 4 , silicon fluoride : , - SiO 2 + 4HF = SiF 4 + 2H 2 O. The test is performed in a platinum crucible covered with a watch-glass. The convex side of the watch-glass is covered with melted wax, and, after the wax has cooled, a design or figure is, by means of a sharpened piece of wood or the point of a knife-blade, graven of sufficient depth in the wax to expose an uncoated surface of glass. The pulverized fluoride is placed in the crucible, moistened with concen- trated sulphuric acid, quickly covered with the watch-glass (waxed side down), and the whole placed on a moderately warm iron plate or porcelain dish. (l) After some time the watch-glass is taken from the crucible, and when the w r ax 1 The watch-glass should be filled with cold water, to prevent the melt- ing of the wax. 68 is removed the graven design will appear etched in the glass. 5. In decomposing fluorides containing considerable silicic acid with concentrated sulphuric acid, gaseous SiF 4 is evolved, which, when conducted through a glass tube moistened with Avater, undergoes decomposition, rendering the water turbid, with the deposition of silicic acid : 3SiF 4 + 3H 2 = H 2 Si0 3 + 2H 2 SiF 6 . The result of the reaction is especially observable on drying the tube. CARBONIC ACID, H 2 CO 3 - (Carbonic acid combines with bases to form salts called car- bonates.) Na 2 CO B , sodium carbonate, may be employed in making the 1. The carbonates of the alkalies are soluble in water, the other carbonates are insoluble in water. Many of the latter are, however, soluble in water containing carbon dioxide, forming acid carbonates : CaCO 3 + C0 2 + H,0 = Ca(HC0 3 ) 2 . Carbonates in general readily dissolve in dilute acids, with effervescence (due to the liberation of carbon dioxide). The metal of the carbonate forms a salt with the acid used as a solvent : BaC0 3 + 2HC1 = BaCl 2 -f- CO 2 + H 2 O. 2. HC1 or any dilute acid (except hydrocyanic acid), added to a carbonate either in solution or in the solid condition, (l) produces effervescence, due to the evolution of carbon dioxide. 1 The minerals magnesite (MgCO 3 ), dolomite (CaCO 3 ,MgCO 3 ), and siderite (FeCO 3 ) produce effervescence with a dilute acid only after being warmed. 69 The latter may be detected by inclining the test-tube in which the effervescence has taken place so as to pour only the gas- eous CO 2 into another test-tube containing clear calcium hy- droxide solution. The CO 2 , being specifically heavier than air, displaces the air in the tube containing calcium hy- droxide, and, on closing the latter tube with the thumb and agitating the liquid, a turbidity is produced, due to the for- mation of CaCO 3 , calcium carbonate : C0 2 + Ca(OH) 2 = CaC0 3 + H 2 O. 3. BaCl 2 , barium chloride, precipitates white BaCO 3 , barium carbonate ; Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white PbCO 3 ; both are soluble with effervescence in dilute acids. 4. AgNO 3 , argentic nitrate, precipitates white Ag 2 CO 3 , argentic carbonate, which in a little time becomes yellowish, and on being boiled with an excess of sodium carbonate changes to brownish-gray Ag 2 O, argentic oxide. The argen- tic oxide is soluble in ammonium hydroxide and in ammo- nium carbonate. SILICIC ACID, H 2 SiO 3 . (Silicic acid combines with bases to form salts called sili- cates.) Na 2 Si0 3 , sodium silicate, may be employed in making the tests. 1. Of the silicates only those of the alkalies are solu- ble in water, the others are partially soluble in concentrated acids. The addition of an acid (as hydrochloric acid) to a solution of a silicate of an alkali causes the separation of silicic acid, which, if the solution is of sufficient concentration, appears as a gelatinous precipitate ; ammonium chloride also separates silicic acid from solutions of silicates of the alkalies : 70 Na 2 Si0 3 + 2HC1 = H 2 SiO 3 + 2NaCl ; Na 2 Si0 3 -f 2JS T H 4 C1 + 2H 2 O = H 2 SiO 3 + 2NaCl + 2NH 4 OH. The silicic acid separated in this manner is somewhat soluble in dilute acids. On evaporating the solution containing silicic acid i.e., the solution with the precipitate in suspension to the dryness of dust on a water-bath, the silicic acid loses water and amorphous silicic acids are produced, i.e., poly- silicic acids, H 2 Si 4 O 9 , for example, which are entirely insolu- ble in water : 4H 2 Si0 3 = H 2 Si 4 9 +3H 2 0. On extracting the residue with water containing a little hy- drochloric acid, the metal which had originally been in com- bination with the silicic acid is dissolved as a chloride, while the silicic acid remains undissolved. Evaporation over a free flame is not advised, as thereby (because of the stability of silicic acid when heated) a part of the salt might be recon- verted into silicates, as, for example : H 2 Si 4 O 9 + 2NaCl = Na^A + 2HC1. For the methods employed in dissolving and disintegrating silicates insoluble in water, see Silicates, page 112. 2. BaCl 2 , barium chloride, precipitates in solutions of silicates of the alkalies white BaSiO 3 , barium silicate; Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white PbSiO 3 , plumbic silicate ; AgNO 3 , argentic nitrate, precipitates yellow- ish Ag 2 SiO 3 , argentic silicate ; all soluble in acids, the argen- tic silicate being also soluble in ammonium hydroxide. 3. NH 4 HMoO 4 , ammonium molybdate, together with an excess of nitric acid, added to a solution of a silicate pro- duces a yellowish coloration, and, in the presence of con- siderable ammonium chloride, a lemon-yellow precipitate. Warming facilitates the reaction. 4. On fusing a silicate with microcosmic salt in the loop of a platinum wire, the sodium metaphosphate which is produced 71 dissolves the base, while the silicic acid remains undissolved and swims in small opaque particles in the otherwise trans- parent bead while the latter is in a state of fusion (" skeleton of silica") ; for example : CaSi0 3 + NaP0 3 = CaNaPO 4 + SiO 2 . Uncombined silicic acid produces the same result. The reac- tion is made more evident by coloring the bead with a com- pound of copper or of iron. ARSENIOUS ACID, H 3 AsO 3 . (See page 21.) ARSENIC ACID, H 3 AsO 4 . (See page 26.) CHROMIC ACID, H 2 CrO 4 . (Chromic acid combines with bases to form salts called chromates.) K 2 CrO^ potassium chromatej may be employed in making the tests. 1. Most of the chromates are insoluble in water. The chromates of the alkalies (the neutral salts) are easily solu- ble ; the bichromates (the so-called acid salts) are soluble, with the production of a reddish-yellow color. 2. BaCl 2 , barium chloride, precipitates from solutions of chromates yellow BaCrO 4 , barium chromate, soluble with great difficulty in water, soluble in hydrochloric and in nitric acids. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates yellow, crys- talline PbCrO 4 , plumbic chromate (chrome-yellow), insoluble in water and in acetic acid, soluble in nitric acid and in sodium hydroxide, in the latter with the formation of Na 2 CrO 4 , sodium chromate, and Pb(ONa) 2 , sodium plum- bite ; acetic acid reprecipitates lead chromate from the sodium hydroxide solution. 72 4. AgNO 3 , argentic nitrate, precipitates in solutions of chromates purplish-red Ag 2 CrO 4 , argentic chromate, and in solutions of bichromates purplish-red Ag 2 O 2 O 7 , argentic bi- chromate, both soluble in nitric acid and in ammonium hydroxide. 5. H 2 S, hydrogen sulphide, conducted into a solution of a chromate containing considerable free hydrochloric acid or sulphuric acid, reduces the chromate, with the formation of a soluble chromic salt and the separation of sulphur, the solution at the same time becoming green in color : 2K 2 CrO 4 + 3H 2 S + 10HC1 = O 2 C1 6 -f S 3 + 4KC1 + 8H 2 O. In case the acid is present in small quantity, greenish Cr 2 (OH) 6 , chromic hydroxide, or (especially on warming the solution) brown chromium chromate is precipitated : 2K 2 CiO 4 +3H 2 S + 4HC1 = O 2 (OH) 6 + S 3 -f 4KC1 + 2H 2 0; 3K 2 CrO 4 + 3H 2 S + 6HC1 = (CrO) 2 CrO 4 + S 3 + 6KC1 + 6H 2 0. The action of ammonium sulphide in neutral or alkaline solutions of chromates is similar to that of hydrogen sulphide. 6. On adding C 2 H 5 OH, alcohol, to a solution of a chromate or bichromate containing free hydrochloric or sulphuric acid, and warming the liquid, the chromate is reduced to a chromic salt, while the alcohol is oxidized to C 2 H 4 O, aldehyde; in consequence, the liquid becomes green in color and the odor of aldehyde becomes evident : K 2 Cr 2 7 -j- 4H 2 S0 4 -f 3C 2 H 5 OH = Cr 2 (SO 4 ) 3 + 3C 2 H 4 O + K 2 SO 4 + 7H 2 O. 7. Chromates fused in a bead of borax or of microcosmic salt impart a yellowish-green color to the bead while hot, which becomes emerald-green on cooling. 73 THIRD GROUP. Acids which are not precipitated by BaCl 2 , barium chloride, but are precipitated by AgNO 3 , argentic nitrate : Hydrochloric Acid, Hydrobromic Acid, Hydriodic Acid, Hydrocyanic Acid, Hydroferro cyanic Acid, Hydroferricyanic Acid, Sulphydric Acid (Hydrogen Sulphide), Nitrous Acid, Hypochlorous Acid. HYDROCHLORIC ACID, HC1. (Hydrochloric acid combines with bases to form salts called chlorides.) NaCl, sodium chloride, may be employed in making the tests. 1. The chlorides are soluble in water, with the exception of argentic chloride, mercurous chloride, and plumbic chlo- ride; the latter, however, being sparingly soluble in cold water. (For dissolving insoluble chlorides, see Dissolving Oxides and Salts, page 104.) t 2. AgNO 3 , argentic nitrate, precipitates white, curdy AgCl, argentic chloride, insoluble in dilute nitric acid, easily soluble in ammonium hydroxide. From its solution in ammonium hydroxide the argentic chloride is reprecipitated by nitric acid. The precipitate is also soluble in KCN, potassium cyanide, and in Na 2 S 2 O 3 , sodium hyposulphite. When ex- posed to sunlight the precipitate changes in color to violet and then to black. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates, in hydro- chloric acid and in solutions of chlorides, white, sometimes crystalline PbCl 2 , plumbic chloride, sparingly soluble in cold water, easily soluble in hot water, from which, when in con- centrated solution, it crystallizes, on cooling, in glistening rhombic needles. Precipitation does not occur in very dilute solutions of chlorides. 74 4. On placing a dry mixture of a chloride and potassium bichromate in a small retort or tubulated fractionating flask, adding concentrated sulphuric acid, and carefully distilling the contents of the retort, CrO 2 Cl 2 , chlorochromic anhydride, as a brownish-red gas, (1) is produced, which, when conducted into a receiving flask, condenses into a brownish-red liquid : 4KC1 + K 2 Cr 2 7 + 6H 2 SO 4 = 2CrO 2 Cl 2 + 6KHSO 4 + 3H 2 O. Sodium hydroxide added to the brownish-red distillate pro- duces a yellowish solution of N^CrO^, sodium chromate (2> (together with sodium chloride) : CrOaCLj + 4NaOH == NajCrO* + 2NaCl + 2H 2 O. If the yellowish solution is acidified with acetic acid and plumbic acetate added, the production of a yellow precipitate of plumbic chromate gives indirect but conclusive evidence of chlorine. HYDROBROMIC ACID, HBr. (Hydrobromic acid combines with bases to form salts called bromides.) KBr, potassium bromide, may be employed in making the 1. The bromides in general are soluble in water. Argentic bromide and mercurous bromide are insoluble ; plumbic bro- mide is sparingly soluble in water. 2. AgNO 3 , argentic nitrate, precipitates yellowish-white, curdy AgBr, argentic bromide, insoluble in dilute nitric acid, sparingly soluble in dilute and more easily soluble in concen- trated ammonium hydroxide. The precipitate is easily soluble in potassium cyanide and in sodium hyposulphite. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates in hydro- 1 Distinction from iodides, which furnish violet-colored vapors of free iodine. 2 Distinction from bromides, which do not impart a color to the liquid. 75 bromic acid and in solutions of bromides white, crystalline PbBr 2 , plumbic bromide, sparingly soluble in cold water, more easily soluble in hot water. 4. Dry bromides, on being distilled in a retort with potas- sium bichromate and concentrated sulphuric acid (see under Chlorides, 4, page 74), yield brown vapors of bromine, which condense in the receiver as a brown distillate of bromine, free from chromium : 6KBr + K 2 Cr 2 O 7 + 7H 2 SO 4 = Cr 2 (SO 4 ) 3 + 4K 2 SO 4 -f 7H 2 + Br 6 . Sodium hydroxide added to the distillate decolorizes it, form- ing sodium bromide and NaBrO, sodium hypobromite : Br 2 + 2NaOH = NaBr -f NaBrO -f H 2 O. 5. Chlorine-water added in small quantity to a solution of a bromide liberates bromine, which remains dissolved in the water. On adding a small quantity of chloroform or of carbon disulphide (both of which are insoluble in water and sink to the bottom of the test-tube), closing the mouth of the tube with the thumb, and thoroughly shaking it, the chloro- form or carbon disulphide extracts the bromine and collects at the bottom of the tube as a yellowish or brownish liquid. The depth of coloration depends upon the quantity of bromine present, If an excess of chlorine- water is used, decolorization of the liquid occurs, due to the formation of HBrO 3 , bromic acid : BrCl 5 + 3H 2 O = HBrO 3 -f 5HC1. HYDRIODIC ACID, HI. (Hydriodic acid combines with bases to form salts called iodides.) y potassium iodide, may be employed in making the tests. 76 1. Most of the iodides are soluble in water; tl^e others are soluble in acids, with the exception of argentic iodide. Plumbic iodide is sparingly soluble in cold water. 2. AgNO 3 , argentic nitrate, precipitates yellowish, amor- phous Agl, argentic iodide, insoluble in nitric acid and in ammonium hydroxide, soluble in potassium cyanide and in sodium hyposulphite. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates in solutions of hydriodic acid and of iodides yellow, crystalline PbI 2 , plumbic iodide, soluble in hot water, from which, on cooling, it separates in glistening yellow, six-sided plates. 4. Dry iodides, distilled in a retort with potassium bichro- mate and concentrated sulphuric acid (see under Chlorides, 4, page 74), yield violet vapors of iodine : (1) 6KI + K 2 Cr 2 7 -f 7H 2 S0 4 = Cr 2 (SO 4 ) 3 + 4K 2 SO 4 + 7H 2 O + 1 6 . The iodine contained in the distillate is soluble in sodium hydroxide, forming Nal, sodium iodide, and NaIO 3 , sodium iodate, the distillate at the same time becoming colorless : I 6 + 6NaOH = 5NaI + NaIO s + 3H 2 O. 5. Chlorine-water added in small quantity to a solution of an iodide liberates iodine, which imparts a yellowish or brownish-yellow color to the solution. On adding a small quantity of chloroform or of carbon disulphide to the liquid, closing the tube with the thumb, and thoroughly shaking it, the chloroform or carbon disulphide will settle at the bottom of the tube, and be found to possess a blue color, due to the free iodine extracted from the aqueous solution. If, instead of chloroform or carbon disulphide, a drop of dilute starch paste is added, the solution becomes blue, due to the action of the free iodine upon the starch. The test is exceedingly delicate, and when considerable iodine is present 1 Distinction from chlorine and bromine. 77 the liquid becomes black upon the addition of the starch ; therefore strong solutions of iodides should be diluted before making this test. The addition of an excess of chlorine-water causes the oxidation of the iodine to iodic acid, with a consequent decolorization of the liquid : IC1 5 + 3H 2 = HI0 3 + 5HC1. HYDROCYANIC ACID, HCN. (Hydrocyanic acid combines with bases to form salts called cyanides.) KCN,potassium cyanide, may be employed in making the tests. 1. Of the cyanides those of the alkalies and of the alkaline earths are soluble in water (also mercuric cyanide) ; the cya- nides of the heavy metals are insoluble in water, although many of them are soluble in potassium cyanide, with the formation of double salts ; for example : AgCN + KCN == AgCN KCN. By the addition of an acid to these solutions, the cyanide of the heavy metal is usually but not invariably reprecipitated, with the evolution of hydrocyanic acid : AgCN KCN + HNO 3 = AgCN + HCN -f-KNO 3 . For methods of dissolving and fusing cyanides, see 4, page 111. 2. AgNO 3 , argentic nitrate, precipitates in solutions of hydrocyanic acid and of cyanides white, curdy AgCN, in- soluble in nitric acid, easily soluble in ammonium hydroxide. From this solution it is reprecipitated by nitric acid : AgCN + NH 3 = NH 3 AgCN ; NH 3 AgCN + HNO 3 = AgCN + NH 4 NO 3 . Argentic cyanide is soluble in potassium cyanide ; therefore 78 a precipitate appears only after an excess of argentic nitrate has been added. It is also soluble in sodium hyposulphite. On igniting argentic cyanide it breaks up into metallic silver and cyanogen gas (together with some argentic paracyanide) : ' 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, produces in solutions of cyanides a white precipitate of Pb(CN) 2 , plumbic cyanide, soluble in nitric acid. 4. If NaOH, sodium hydroxide, FeSO 4 , ferrous sulphate, and Fe 2 Cl 6 , ferric chloride, are added in small quantities to a solution of hydrocyanic acid or to a cyanide, the mixture warmed, and finally acidulated with hydrochloric acid, a blue precipitate of Fe 4 (Fe(CN) 6 ) 3 , ferric ferrocyanide (Prussian blue), is formed ; while the ferrous hydroxide first produced is dissolved by the acid. The ferrous sulphate with the sodium hydroxide produces Fe(OH) 2 , ferrous hydroxide : FeSO 4 + 2NaOH = Fe(OH) 2 + Na 2 SO 4 , which, on being warmed with the cyanide solution, yields a ferrocyanide ; for example, with potassium cyanide it yields K 4 Fe(CN) 6 , potassium ferrocyanide : Fe(OH) 2 + 6KCN = K 4 Fe(CN) 6 + 2KOH, which combines with the iron of the ferric chloride to form blue ferric ferrocyanide (Prussian blue). 5. To detect hydrocyanic acid which is being evolved from a liquid, a drop of yellow ammonium sulphide and of ammo- nium hydroxide is placed on the concave side of a watch- glass, the watch-glass inverted and placed as a cover over the vessel in which the hydrocyanic acid is being evolved, so that the vapors of the acid coming in contact with the ammoniacal liquid can be absorbed. After some time the watch-glass is removed, placed on a water-bath, and warmed, whereby NH 4 CNS, ammonium sulphocyanide, is produced : HCN + (NH 4 \S 2 +NH 4 OH = NH 4 CNS + (NH 4 ) 2 S + H 2 O, 79 which remains as a dry residue on the complete evaporation of the liquid. This residue is dissolved in a little water ; a few drops of hydrochloric acid (to decompose any (NHJ 2 S remaining) and a drop of ferric chloride are added, whereby a claret-red coloration is produced, due to the formation of Fe 2 (CNS) 6 , ferric sulphocyanide. 6. When heated in a reduction- tube the cyanides of the heavy metals are decomposed ; the cyanides of the noble metals break up into metal and cyanogen gas ; other cyanides break up into metal, carbon, and nitrogen. Argentic and mercuric cyanides, in which the cyanogen cannot be detected by the ordinary reagents, can be recognized in this manner. Mercuric cyanide in aqueous solutions, when treated with hydrogen sulphide, decomposes and forms mercuric sulphide and hydrocyanic acid. HCN and CN are virulent poisons. HYDROFERROCYANIC ACID, H 4 Fe(CN) 6 . (Hydroferrocyanic acid combines with bases to form salts called ferrocyanides.) KFe(CN)v potassium ferrocyanide, may be employed in making the tests. 1. The ferrocyanides, with the exception of those of the alkalies and of the alkaline earths, are mostly insoluble in water. Regarding their solution and fusion, see page 112. 2. AgNO 3 , argentic nitrate, precipitates white Ag 4 Fe(CN) 6 , argentic ferrocyanide, insoluble in nitric acid and in ammo- nium hydroxide, soluble in potassium cyanide. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, in solutions of ferro- cyanides precipitates white Pb 2 Fe(CN) 6 , plumbic ferrocyanide, insoluble in dilute nitric acid. 4. Ferrous salts (FeSO 4 , ferrous sulphate) produce in solu- tions of ferrocyanides (when the ferrocyanide is in excess) a 80 white precipitate, which, on exposure to the air, rapidly changes to bluish-white K 2 Fe(Fe(CN) 6 ), potassium ferrous ferrocyanide (Everett's salt). When the ferrous salt is in excess, Fe 2 Fe(CN) 6 , ferrous ferrocyanide, is produced. 5. Ferric salts (Fe 2 Cl 6 , ferric chloride) precipitate dark-blue Fe 4 (Fe(CN) 6 ) 3 , ferric ferrocyanide (Prussian blue), insoluble in acids. 6. CuSO 4 , cupric sulphate, precipitates brownish-red Cu 2 Fe(CN) 6 , cupric ferrocyanide. HYDROFERRICYANIC ACID, H 3 Fe(CN) 6 . (Hydroferricyanic acid combines with bases to form salts called ferricyanides.) K 3 Fe(CN) 6 , potassium ferricyanide, may be employed in making the tests. 1. Of the ferricyanides those of the alkalies and of the alkaline earths are soluble in water, while those of the heavy metals are mostly insoluble in water. Regarding their solu- tion and fusion, see page 112. 2. AgNO 3 , argentic nitrate, precipitates from solutions of ferricyanides reddish-brown Ag 3 Fe(CN) 6 , argentic ferri- cyanide, insoluble in nitric acid, soluble in ammonium hy- droxide and in potassium cyanide. 3. Ferrous salts (FeSO 4 , ferrous sulphate) precipitate Fe 3 (Fe(CN) 6 ) 2 , ferrous ferricyanide (TurnbulPs blue), insolu- ble in acids. Ferric salts fail to produce a precipitate, but cause a dark coloration ; possibly soluble Fe 2 (Fe(CN) 6 ) 2 , ferric ferricyanide, is produced : Fe 2 Cl 6 + 2K 3 Fe(CN) 6 = Fe^F^CN)^ + 6KC1. 4. CuSO 4 , cupric sulphate, precipitates greenish-yellow Cu 3 (Fe(CN) 6 ) 2 , cupric ferricyanide. 81 SULPHYDRIC ACID, H 2 S (HYDROGEN SULPHIDE). (Sulphydric acid combines with bases to form salts called sulphides.) Na^S, sodium sulphide, may be employed in making the tests. 1. The sulphides, with the exception of those of the alkalies and of the alkaline earths, are insoluble in water. Most of them are soluble in hydrochloric and in nitric acids ; some are soluble only in nitro-hydrochloric acid. (See Sulphides of the Heavy Metals, page 111.) They may be recognized by their giving off hydrogen sulphide when dis- solved in hydrochloric acid, or by the separation of sulphur when dissolved in nitric acid or in nitro-hydrochloric acid. 2. AgNO 3 , argentic nitrate, precipitates black Ag 2 S, argen- tic sulphide, soluble in nitric acid when warmed. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates in solutions of sulphides or of hydrogen sulphide black PbS, plumbic sulphide, soluble in nitric acid when warmed. 4. To detect hydrogen sulphide gas a strip of filter-paper is moistened with plumbic acetate and held in the atmosphere containing the gas. In the presence of hydrogen sulphide the paper becomes brown or black, due to the formation of PbS, plumbic sulphide. 5. A few drops of an alkaline solution of plumbic oxide (Pb(OK) 2 , potassium plumbite), added to a solution contain- ing hydrogen sulphide or a sulphide of a metal, produces a perceptible brownish coloration, even if only the slightest trace of the sulphide be present. 6. ]N"a 2 NOFe(CN) 5 , sodium nitro-prusside, solutions are colored violet by sulphides, but not by solutions of free hydrogen sulphide. 7. Many of the sulphides of the metals, when heated in a reduction-tube, yield a sublimate of sulphur. Sulphides, 82 heated in a glass tube open at both ends and held obliquely in the flame, are oxidized, with the formation of SO 2 , sul- phurous anhydride. Ignited with sodium carbonate in the reducing flame on charcoal, they yield sodium sulphide, which, when placed on a clean silver coin and moistened with water, produces a black discoloration of argentic sulphide. NITROUS ACID, HNO 2 . (Nitrous acid combines with bases to form salts called nitrites.) KN0 2 , potassium nitrite, may be employed in making the tests. 1. Most of the nitrites are soluble in water. Treated with hydrochloric or sulphuric acid they evolve brownish-red fumes of NO 2 , nitrogen dioxide. 2. AgNO 3 , argentic nitrate, precipitates white AgNO 2 , argentic nitrite, soluble with difficulty in water. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, colors solutions of nitrous acid yellow. 4. H 2 S, hydrogen sulphide, is decomposed by nitrous acid with the separation of sulphur : H 2 S + 2HNO 2 .= 2NO + S + 2H 2 O. 5. FeSO 4 , ferrous sulphate, added to a solution of a nitrite containing a few drops of sulphuric acid, (1) produces a brown or black coloration, due to the formation of NO, nitrogen monoxide, which enters into combination with the ferrous sulphate : (a+3)FeSO 4 + H 2 SO 4 + 2HNO, = *FeSO 4 (NO) 2 + Fe 2 (S0 4 ) 3 + 2H 2 0. Heating the liquid causes the coloration to disappear. 1 The nitrites of commerce usually contain free nitrous acid, and there- fore respond to the test without the addition of sulphuric acid. 83 6. KI, potassium iodide, (or CdI 2 , cadmium iodide,) starch paste, and dilute sulphuric acid, added to a solution of a nitrite, immediately produce a blue coloration in the liquid. The nitrous acid liberates iodine from the hydriodic acid : 2HNO 2 + 2HI = 2NO + 2H 2 O + I 2 . The free iodine combining with the starch forms the blue compound. (In this test cadmium iodide or potassium iodide free from iodic acid should be used, as hydriodic and iodic acid undergo decomposition when together, with the liberation of iodine : HYPOCHLOROUS ACID, HC1O. (Hypochlorous acid combines with bases to form salts called hypochlorites.) NaCIO, sodium hypochlorite, may be employed in making the 1 . The hypochlorites, as a rule, contain chlorides, produced, during the preparation of the hypochlorite, by the action of the chlorine upon hydroxides : 2NaOH + C1 2 = NaCIO + NaCl + H 2 O. On the addition of acids they are decomposed, with the evolution of chlorine : NaCIO + 2HC1 = NaCl + C1 2 + H 2 O ; NaCIO + NaCl + H 2 SO 4 = Na 2 SO 4 + C1 2 -f H 2 O. 2. AgNO 3 , argentic nitrate, added to a solution of a hypo- chlorite produces soluble AgCIO, argentic hypochlorite, which immediately breaks up into white, insoluble AgCl, argentic chloride, and soluble AgClO 3 , argentic chlorate : GAgCIO 2AgClO 3 + 4AgCl. 3. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, produces at first a white precipitate of PbCl 2 , plumbic chloride, which soon becomes 84 yellow and finally brown, due to the formation of PbO 2 , lead dioxide. (In like manner MnSO 4 , manganous sulphate, yields brown MnO(OH) 2 , hydrated peroxide of manganese.) FOURTH GROUP. Acids which are not precipitated by barium chloride or by argentic nitrate : Nitric Acid, Chloric Acid. NITRIC ACID, HNO 3 . (Nitric acid combines with bases to form salts called nitrates.) KNOfr potassium nitrate, may be employed in making the tests. 1. The nitrates, with the exception of a few basic salts, are soluble in water. Some nitrates (for example, Ba(NO 3 ) 2 , barium nitrate) are only sparingly soluble in nitric acid. 2. BaCl 2 , Pb(C 2 H 3 O 2 ) 2 , and AgNO 3 do not produce pre- cipitates in solutions of nitrates. 3. On placing a small crystal of FeSO 4 , ferrous sulphate, in a cooled mixture of concentrated sulphuric acid and a solution of a nitrate, a brownish-black ring is formed around the crystal. In the reduction of the nitric acid NO, nitrogen monoxide, is produced, which combines with, the ferrous sulphate to form an unstable compound : (*+6)FeS0 4 + 3H 2 S0 4 + 2HNO 3 = [>FeSO 4 (NO) 2 ] + 3Fe 2 (SO 4 ) 3 + 4H 2 O. The test is best made in a flat porcelain dish, or in a watch- glass placed on white paper. Heat destroys the black ring. 4. KI, potassium iodide, (or CdI 2 , cadmium iodide,) starch paste, and dilute sulphuric acid, added to a solution of a nitrate, produce no reaction (distinction from nitrites), but, 85 on placing a fragment of zinc in the liquid, nitrous acid is evolved, which, acting upon the potassium iodide, liberates the iodine, which with the starch produces a blue coloration : HN0 3 +Zn + H 2 S0 4 = HNO 2 + ZnSO 4 + H 2 O ; 2HNO 2 + 2HI = 2NO + 2H 2 O + I 2 . 5. Nitrates of the alkalies when heated in a reduction-tube are reduced to nitrites, with the evolution of oxygen : The nitrates of the heavy metals when heated in a reduc- tion-tube evolve reddish-brown fumes of nitrogen dioxide : Pb(N0 3 ) 2 = PbO + 0-1- 2N0 2 . The latter reaction also takes place when a nitrate of an alkali mixed with cupric sulphate is heated in a reduction-tube : 2KNO 8 + CuSO 4 = K 2 SO 4 +CuO + O + 2NO 2 . Nitrates deflagrate when ignited on charcoal. CHLORIC ACID, HC1O 3 . (Chloric acid combines with bases to form salts called chlorates.) KQIO Z , potassium chlorate, may be employed in making the tests. 1. The chlorates are soluble in water. 2. BaCl 2 , barium chloride, does not produce a precipitate in solutions of chlorates. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, and AgNO 3 , argentic nitrate, do not produce precipitates if the solution of the chlorate be free from chlorides. 3. On warming a solution of a chlorate with hydrochloric acid the liquid becomes greenish yellow in color, and greenish- yellow fumes of a mixture of chlorine and C1 2 O 4 , chlorine tetroxide (chlorine peroxide), are evolved : KC1O 3 +6HC1 = KC1 + C1 8 + 3H 2 O ; 2KC1O 3 + 4HC1 = 2KC1 + C1 2 O 4 + C1 2 + 2H 2 O. 86 4. Concentrated sulphuric acid poured over a very small piece of a chlorate in a porcelain dish causes a decomposition of the chlorate, with the production of a perchlorate and chlorine tetroxide (chlorine peroxide) : 3KC10 3 + 2H 2 S0 4 = 2KHS0 4 + KC1O 4 + C1 2 O 4 + H 2 O. Great care should be used in making this test, and only small quantities of chlorate should be employed. Warming should be avoided, as explosions, which may cause personal injury, are likely to occur on the application of heat. 5. Chlorates heated in a reduction-tube undergo decomposi- tion, and are converted into chlorides, with the evolution of oxygen : KC1O 3 = KC1 + O s . (Bromates and iodates undergo a similar decomposition on being heated, forming respectively bromides and iodides, with the evolution of oxygen.) APPENDIX: OEGANIC ACIDS. Acetic Acid, Oxalic Acid, Tartaric Acid. ACETIC ACID, HC 2 H 3 2 . (Acetic acid combines with bases to form salts called acetates.) NaC 2 H 3 2 , sodium acetate, may be employed in making the tests. 1. Most of the acetates are easily soluble in water. 2. BaCl 2 , barium chloride, and Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, do not produce precipitates in solutions of acetates. 3. AgNO 3 argentic nitrate, precipitates, in concentrated acetic acid or in concentrated solutions of acetates, crystalline AgC 2 H 3 O 2 , argentic acetate, soluble in a large quantity of water and in ammonium hydroxide. 87 4. Fe 2 Cl 6 , ferric chloride, added to a neutral acetate, or to acetic acid, which must afterwards be exactly neutralized with ammonium hydroxide, produces a reddish-brown solution of Fe 2 (C 2 H 3 O 2 ) 6 , ferric acetate : 6NaC 2 H 3 2 + Fe 2 Cl 6 = Fe 2 (C 2 H 3 O 2 ) 6 + 6NaCl. On warming this solution a precipitate of brownish-red Fe 2 (OH) 4 (C 2 H 3 O 2 ) 2 , basic ferric acetate, separates, while the supernatant liquid becomes colorless : Fe 2 (C 2 H 3 2 ) 6 + 4H 2 = Fe 2 (OH) 4 (C 2 H 3 O 2 ) 2 + 4HC 2 H 3 O 2 . 5. On adding sulphuric acid to a solution of an acetate and warming the liquid, HC 2 H 3 O 2 , acetic acid, is liberated, which may be recognized by its odor of vinegar. 6. On adding C 2 H 5 OH, alcohol, to a cool solution of an acetate containing sulphuric acid and then warming the liquid, C 2 H 5 C 2 H 3 O 2 , ethyl acetate (acetic ether), is produced, which may be recognized by its characteristic apple-like odor : C 2 H 5 OH + H 2 S0 4 = C 2 H 5 HS0 4 + H 2 O ; C 2 H 5 HS0 4 -f HC 2 H 3 2 = C 2 H 5 C 2 H 3 2 + H 2 SO 4 . The alcohol should not be added while the liquid is hot, as violent ebullition might occur with consequent spurting of the liquid. 7. Acetates on being ignited are decomposed, without the separation of carbon, into volatile products (for example, acetone) and carbonates or oxides of the metals which were in combination as acetates. OXALIC ACID, H 2 C 2 O 4 . (Oxalic acid combines with bases to form salts called oxalates.) ( J /V r J fZ 4 ) 2 C 2 4 , ammonium oxalate, may be employed in making the tests. 1. Of the oxalates those of the alkalies are soluble in water ; most of the others are insoluble in water. 88 2. BaCl 2 , barium chloride, precipitates in solutions of neutral oxalates white BaC 2 O 4 , barium oxalate, easily soluble in hydrochloric and in nitric acid. 3. CaCl 2 , calcium chloride, precipitates from neutral solu- tions of oxalates white CaCgO^ calcium oxalate, soluble in hydrochloric and in nitric acid, insoluble in acetic acid. 4. Pb(C 2 H 3 O 2 ) 2 , plumbic acetate, precipitates white PbC 2 O 4 , plumbic oxalate, soluble in nitric acid. 5. AgNO 3 , argentic nitrate, precipitates white Ag 2 C 2 O 4 , argentic oxalate, soluble in nitric acid and in ammonium hydroxide. 6. Concentrated sulphuric acid, on being warmed with oxalic acid or oxalates, decomposes them into water, CO 2 , carbon dioxide, and CO, carbon monoxide : H 2 C 2 O 4 + H 2 SO 4 == H 2 O + CO 2 + CO + H 2 SO 4 . On pouring the gases into a test-tube containing clear solution of calcium hydroxide (lime-water), closing the tube with the thumb, and shaking it, the production of a milky turbidity, due to the formation of calcium carbonate, indicates the presence of carbon dioxide. 7. Oxalates on ignition are decomposed into carbon mon- oxide and carbonates or oxides of the metals which were in combination as oxalates. Pure oxalates on being ignited do not become black in color : TARTARIC ACID, H 2 C 4 H 4 O 6 . (Tartaric acid combines with bases to form salts called tart rates.) KNaCHiO & , potassium sodium tartrate, may be employed in making the tests. 1. The tartrates of the alkalies and some of the tartrates 89 of the heavy metals are soluble in water, the other tartrates are soluble in acids. 2. BaCl 2 , barium chloride, added in excess precipitates white BaC 4 H 4 O 6 , barium tartrate, soluble in hydrochloric and in nitric acid. 3. CaCl 2 , calcium chloride, added in excess precipitates white, crystalline CaC 4 H 4 O 6 , calcium tartrate, soluble in hydrochloric, nitric, and acetic acids. The precipitate is also soluble in potassium or sodium hydroxides, forming a clear liquid, from which, on boiling, the calcium salt separates in gelatinous masses. Probably a salt, Ca~Na 2 C 4 H 2 O 6 , is pro- duced in which the hydrogen atoms of the alcoholic hydroxyl of the tartaric acid have also been replaced by a metal : CaC 4 H 4 6 + 2NaOH = CaNa 2 C 4 H 2 O 6 + 2H 2 O. This compound, on being boiled with considerable water, is reconverted into the original calcium tartrate : CaNa 2 C 4 H 2 6 + 2H 2 O = CaC 4 H 4 O 6 + 2NaOH. 4. Pb(C 2 H 3 O 2 ) 2; plumbic acetate, precipitates white PbC 4 H 4 O 6 , plumbic tartrate, soluble in nitric acid and in ammonium hydroxide. 5. AgNO 3 , argentic nitrate, precipitates in solutions of neutral tartrates' Ag 2 C 4 H 4 O 6 , argentic tartrate, soluble in nitric acid and in ammonium hydroxide. On boiling the precipitate it is decomposed, with the separation of metallic silver. 6. Tartrates on being ignited are decomposed, with the production of an odor resembling burnt sugar, the separation of carbon and the formation of carbonates. III. PRELIMINARY EXAMINATION. (A) PRELIMINARY TESTS IN THE DRY WAY. THE special tests for bases and acids (testing in the Wet Way) are always preceded by a short preliminary examina- tion (in the Dry Way), in order to obtain general information regarding the nature of the substance to be analyzed. It is hardly possible to determine the best method to be employed in the preparation of the substance for analysis without re- sorting to this preliminary examination. It should therefore never be omitted. When solutions are to be analyzed, a portion is evaporated to dryness at a moderate temperature (without ignition), and the residue used for the preliminary tests. 1. EXAMINATION IN THE REDUCTION-TUBE. To ascertain the behavior of the substance at higher tem- peratures, a small portion of it, or of the residue obtained by evaporation, is placed in a narrow glass tube closed at one end, and heated, at first slightly, afterwards more strongly, and then to redness. The occurrence of any of the following changes should especially be noted : 1. Separation of Carbon: Indicates the presence of organic compounds. Simultaneously a generation of empyreumatic vapors takes place, or, if nitrogen is present, an odor of burnt feathers. 2. Elimination of Water : Indicates the presence of water of crystallization or of adherent moisture ; frequently a 90 91 change of color occurs, as in the transformation of the blue hydrous sulphate of copper (CuSO 4 -f 5H 2 O) into the an- hydrous salt (CuSO 4 ). Intumescence may take place as in the case of borax (Na 2 B 4 O 7 -j- 10H 2 O), or decrepitation as in sodium chloride (in consequence of the violent expulsion of water confined between the lamellae of the crystals). 3. Change in Color: Indicates the presence of combina- tions of heavy metals. The change may be caused by the elimination of water (see 2, page 90), or by the conversion of salts into oxides ; for example, cupric nitrate and cupric carbonate become black in color when heated, due to their conversion into cupric oxide : Cu(NO 3 ) 2 = CuO -f 2NO 2 + O ; CuCO 3 = CuO + CO 2 . Many compounds differ in color when hot and when cold ; for example, oxide of zinc is yellow when hot and white when cold. 4. Formation of a Sublimate: Indicates the presence of volatile compounds. (a) White sublimate : Salts of mercury, ammonium salts, arsenious oxide, antimonious oxide. On heating the sublimate with dry sodium carbonate, salts of mercury become red, due to the formation of mercuric oxide (frequently metallic mer- cury is produced at the same time) ; for example : HgCl 2 + NaaCOs = HgO + CO 2 + 2NaCl. Ammonium salts evolve ammoniacal gas, which may be recognized by the odor, and by its coloring moistened tur- meric-paper brown, and red litmus-paper blue : 2NH 4 C1 -f- ]Sa 2 C0 3 = 2NH 3 + CO 2 -f H 2 O + 2NaCl. Arsenical vapors and antimonious oxide are apparently not changed when heated with sodium carbonate. The arsenic sublimes in octahedral crystals; the antimony forms an amorphous sublimate, which sometimes contains crystals. 92 (b) Yellow sublimate : Mercuric iodide (becomes red when stirred), arsenious sulphide. (c) Yellow to red : Compounds of mercury (formation of basic salts). (d) Yellow to brownish yellow : Sulphur (when hot col- lects in reddish-brown drops). Free sulphur, or sulphides rich in sulphur, for example, Sb 2 S 5 = Sb 2 S 3 -f- S 2 . (e) Gray to black : Mercury (globules) ; mercuric sulphide (black, red when rubbed) ; iodine (violet vapors, characteristic odor of iodine) ; arsenic (mirror). It is to be remembered that, in addition to the sublimates mentioned, quite a number of compounds exist which are more or less volatile, for example, many chlorides. 5. Evolution of Vapors : (a) Colorless vapors should be tested for their reaction with litmus-paper. The acids frequently form clouds when escaping from the tube (in consequence of their changing from the anhydrous to the hydrous state). (b) Reddish-brown vapors : Nitrogen dioxide, bromine. Nitrogen dioxide, resulting from the decomposition of nitrates of the heavy metals, for example, Pb(NO 3 ) 2 = PbO -f 2NO 2 -j- O, does not color starch-paper, and is recognized by its odor. Bromine, also recognizable by its odor, colors starch-paper reddish yellow. (c) Violet vapors : Iodine. Characteristic odor ; frequently simultaneous formation of a black sublimate. Colors starch- paper blue to brownish black. 6. Production of an Odor : (a) Odor of ammonia : Ammonium salts ; compounds of cyanogen or organic compounds containing nitrogen. (b) Odor of sulphurous anhydride : Resulting from the decomposition of sulphates. (c) Odor of cyanogen : Compounds of cyanogen. Cyan- 93 ogen gas burns when ignited, with a flame pinkish lavender in color : (d) Odor of garlic : Compounds of arsenic, resulting from reduction. 7. Evolution of Oxygen (may be recognized by the flaring or re-igniting of a glowing stick held at the mouth of the tube) : Indicates the presence of peroxides, for example, pyrolusite, MnO 2 : 3MnO 2 = Mn 3 O 4 -f O 3 ; of mercuric oxide : HgO = Hg + 0; of salts rich in oxygen, for example : KC1O 3 = KC1 + O 3 . 2. EXAMINATION ON CHARCOAL. To determine the behavior of substances in the reducing flame a small portion of the substance, generally mixed with dry sodium carbonate, is heated in a cavity in the charcoal by means of the reducing flame of the blowpipe. (1) The sodium carbonate is added in order to transform salts and sulphates into carbonates and oxides respectively, for example : CaSO 4 + Na 2 CO 3 = CaCO 3 -f ]S T a 2 SO 4 ; CuCl 2 -f Na 2 CO 3 = CuO + CO a + 2NaCl. The addition of sodium carbonate is not necessary in the case of metals that form metallic globules, oxides, and salts which are easily decomposed, as the alkalies and their salts are ab- sorbed by the charcoal (because of their easy fusibility). The 1 The reducing flame is obtained by holding the blowpipe near the flame and by gentle blowing directing it upon the substance to be heated. The oxidizing flame is obtained by placing the blowpipe in the interior of the flame and blowing with force. 94 oxides of the remaining elements may be recognized by the following characteristics : 1. The oxides of the heavy metals heated in the reducing flame are reduced by the charcoal. The metals themselves are either volatile or noil- volatile, may oxidize or not, and may be fusible or infusible ; therefore fused globules may be obtained, or infusible masses and incrustations, the latter resulting from the presence of metals that volatilize and oxidize. From plumbic oxide, for example, metallic lead is obtained : part of which volatilizes, combines with the oxygen of the air, and is deposited on the cooler part of the charcoal as an incrustation of yellow oxide : The metallic globules differ in their behavior in the oxi- dizing flame: some change into oxides and others remain unchanged. The ductility should also be ascertained ; for this purpose the globule is placed in a mortar and struck with the pestle ; those which are ductile are flattened into plates, while those which are brittle break into pieces and may be pulverized by subsequent rubbing. (a) Fused metallic globules, without incrustation, are produced : Yellow : gold, ductile, not oxidizable. White : silver, ductile, not oxidizable. Ked : copper, (1) ductile, oxidizable. With incrustation, White globule, incrustation yellow : Ductile : lead, oxidizable. Brittle : bismuth, oxidizable. White globule, incrustation white : 1 Generally obtained as metallic spangles. 95 Ductile: tin, (1) oxidizable. Brittle : antimony, oxidizable. (6) Incrustation without metallic globule : White (when hot, yellow when cold) : zinc. Yellowish red to brown : cadmium. (c) Gray, infusible masses : Iron "^ Cobalt Nickel Manganese J Platinum : not oxidizable. (d) Neither globule nor incrustation : Volatile with odor of garlic : arsenic. Volatile without odor of garlic : mercury. In examining metallic globules it is to be remembered that in the presence of different metals alloys may be formed. 2. White infusible masses remain on the charcoal if salts of the alkaline earths, magnesium or aluminium, are present. (By the action of Na 2 CO 3 , carbonates and oxides are formed.) The white masses, moistened with a solution of cobaltous nitrate and strongly heated in the oxidizing flame, yield as follows : Aluminium : blue masses (infusible). Magnesium : pink-colored masses. Barium ^| Strontium J> gray masses. Calcium J 1 Tin and antimony are obtained with difficulty in the form of globules when sodium carbonate is employed. Therefore on the appearance of a white incrustation a second test is made, in which, in addition to sodium carbonate, potassium cyanide is added to the salt, and the whole heated in the reducing flame. (KCN thereby changes into KCNO : for example, SnO 2 -f 2KCN = Sn + 2KCNO.) Compare also its behavior with cobalt solution, see 6, page 34. 96 The cobaltous nitrate on being heated is converted into co- baltous oxide : Co(NO 3 ) 2 = CoO + 2NO 2 -f O, which combines with aluminium and magnesium compounds. With barium, strontium, and calcium, mixtures only of the oxides are obtained. Many silicates and phosphates which are fusible with diffi- culty, and also many borates and arseniates, may form blue masses when ignited with cobaltous nitrate ; frequently these double salts of cobalt are easily fusible. Zinc oxide, when ignited with cobaltous nitrate, becomes yellowish green in color ; antimonious oxide, a dirty green ; stannic oxide, bluish green. (Compounds of CoO are pro- duced with the different oxides.) 3. Green fused masses (consisting of chromic oxide) indi- cate salts of chromium and chromates. 4. Yellow or brown fused masses, consisting of sodium sul- phide, indicate the presence of compounds containing sulphur. A portion is placed on a silver coin and moistened with water to ascertain whether a black discoloration of Ag 2 S is produced. (See 7, page 81.) As the formation of sodium sulphide by the reduction of salts containing acids of sulphur requires time, and, like all alkali compounds, the sulphide impregnates the charcoal on continued heating, these tests must be made just after the reduction has taken place and before the sodium sul- phide has been absorbed by the charcoal. Many of the compounds containing sulphur, for example, the sulphides, when heated in a small glass tube open at both ends and held obliquely in the flame, yield sulphurous anhydride, which is easily recognized by its odor. 97 3. EXAMINATION BY MEANS OF MICROCOSMIC SALT. A portion of sodium ammonium phosphate (microcosmic salt) is heated in a loop of platinum wire until it melts, and forms a bead. A very small portion of the substance to be examined is then attached to the clear bead, which is again heated in the oxidizing flame or in the oxidizing space of a Bunsen flame. The ]S T aNH 4 HPO 4 -f 4H 2 O, sodium ammo- nium phosphate, when fused, first loses its water of crystal- lization and then changes into sodium metaphosphate : NaNH 4 HPO 4 == NaPO 3 -f H 2 O + NH 3 . Th^ sodium metaphosphate dissolves most of the oxides and salts (in the latter a replacement of the acids takes place), and forms beads, generally characteristic in color : CuO -f NaP0 3 = CuNaP0 4 ; CuSO 4 -f JS T aPO 3 = CuNaPO 4 + SO 3 . Some of the beads change color in the reducing flame or in the reducing space of the Burisen flame (in consequence of the reduction of the phosphates) ; for example, the transparent bluish-green copper bead by reduction becomes brownish-red and opaque : CuNaPO, + C = NaP0 3 + Cu + CO ; the violet manganic oxide bead in the reducing flame is con- verted into the colorless manganous bead : Mn 2 (NaPO 4 ) 3 -f C = 2MnNaPO 4 + NaPO 3 + CO. Borax (Na 2 B 4 O 7 -f 10H 2 O) with oxides and salts yields beads similar to microcosmic salt, which are likewise reducible : Na 2 B 4 O 7 -f CuO = 2NaBO 2 + Cu(BO 2 ) 2 ; Na 2 B 4 O 7 + CuSO 4 = 2NaBO 2 + Cu(BO 2 ) 2 + SO 3 ; 2NaBO 2 -f Cu(BO 2 ) 2 + C = Na 2 B 4 O 7 + Cu + CO. The reduction of the oxide in the beads is often facilitated by adding a small piece of tin foil : 2Cu]SaPO 4 + Sn = Sn(NaPO 4 ) 2 + Cu 2 . E g 98 The following elements (1) produce characteristic colorations in the bead of microcosmic salt : Oxidizing Flame. Reducing Flame. Iron : yellow to dark red when Green to colorless, hot, light yellow to colorless when cold. Nickel : same as iron. As in the oxidizing flame. (2) Cobalt : blue. Blue. Manganese : violet. Colorless. Chromium : green. Green. Copper : blue-green. Brownish, opaque. The remaining oxides yield colorless, transparent or .trans- lucent, enamel-like beads. The behavior of silicic acid and of the silicates in the bead of microcosmic salt is characteristic. Silicic acid does not dissolve in the bead, but, while the bead is in a state of fusion, swims in distinctly outlined masses. The silicates are decom- posed in the bead, with the separation of undissolved silicic acid : NaPO 3 + CaSiOj = CaNaPO 4 + SiO 2 . 4. EXAMINATION IN THE FLAME. If the presence of alkalies or alkaline earths is suspected, a small portion of the substance, or of the residue obtained by evaporation, is attached to a loop of dean platinum wire/ 3) moistened with a drop of hydrochloric acid, and held in the flame of a Bunsen burner. The flame is colored by the salts of Potassium : violet. Sodium : intense yellow. *For beads produced by the rare elements, see Appendix. 2 For the behavior of nickel in the borax bead, see 7, page 49. 3 Or the wire may be dipped in the concentrated solution of the sub- stance to be examined. 99 Barium : green. Strontium : crimson. Calcium : yellowish red. It must be remembered that, if two or more of these ele- ments are present, one colored flame may interfere with the other. Salts of copper and also boric acid color the flame green. For colored flames produced by the rare elements, see Ap- pendix. (B) PRELIMINARY TESTS FOR ACIDS. Important conclusions regarding the presence or absence of certain acids may be drawn from the behavior of their salts with dilute and concentrated sulphuric acid, and also with alcohol and sulphuric acid. 1. If a portion of the substance or solution be placed in a test-tube and dilute sulphuric acid poured over it, there may be evolved : Colored Gases: Greenish-yellow chlorine in presence of hypochlorites (see 1, page 83). Moistened potassium iodide starch paper held in the fumes is colored blue. Red vapors of nitrogen dioxide from nitrites (see 1, page 82). Colorless Gases recognized by their Odor : Sulphurous an- hydride, from sulphites or hyposulphites; in the presence of the latter, separation of sulphur also takes place (see 2, page 63). Detection of sulphurous an- hydride by potassium iodate (see 2, page 61). Hydrocyanic acid, from many of the cyanides, recognized by its odor of bitter almonds, and also by the sulpho- cyanide reaction (see 5, page 78). 100 Acetic acid in presence of acetates. Hydrogen sulphide, from many of the sulphides, blackens paper saturated with solution of plumbic acetate (see 4, page 81). Polysulphides evolve hydrogen sul- phide, with the separation of sulphur; sulpho-acids may also separate. (See pages 121, 122.) Colorless and Odorless Gas : Carbon dioxide is liberated with effervescence from carbonates (to be confirmed with calcium hydroxide, see 2, page 68). 2. If a small portion of the substance is treated with three or four times its volume of concentrated sulphuric acid and gently warmed, there may be evolved : Colored Gases : Greenish-yellow chlorine in presence of hypo- chlorites; also when both chlorides and nitrates, or chlorides and peroxides are present. (When chlorides and nitrates are present, hydrochloric acid and nitric acid are simultaneously liberated and react upon each other (see d, page 106). When chlorides and per- oxides are present, the liberated hydrochloric acid acts upon the peroxides (see c, page 106).) Greenish-yellow explosive mixture of chlorine and chlo- rine tetroxide, derived from chlorates (see 4, page 86). Brownish bromine together with hydrobromic acid de- rived from bromides ; the gas colors starch paper red- dish yellow. Brownish-red chromium oxychloride when chlorides and chromates are both present (see 4, page 74). Reddish-brown fumes indicate nitrites (see 1, page 82). Violet vapors of iodine from iodides color moistened starch paper blue. Colorless Gases recognized by their Odor: Hydrochloric acid 101 vapors from chlorides; pungent odor, and render argentic nitrate solution (on glass rod) turbid (see 2, page 73). Hydrobromic acid (see above). Hydrofluoric acid from fluorides ; of a strongly acid odor, etches glass (see 4, page 67). Nitric acid from nitrates, pungent odor. Red vapors arise when ferrous sulphate is added. Sulphurous anhydride from sulphites and hyposulphites (see 2, page 61, and 2, page 63. N.B. May also result from the reduction of the sulphuric acid employed). Hydrogen sulphide from sulphides (see 4, page 81). Acetic acid from acetates, odor of vinegar (see 5, page 87). Colorless and Odorless Gases: Oxygen (recognized by test with glowing wood, see 7, page 93) in presence of per- oxides, chromates, and permanganates ; for example : Mn0 2 + H 2 SO 4 = MnS0 4 + H 2 O + O ; 2K 2 CrO 4 + 5H 2 SO 4 = Cr 2 (SO 4 ) 3 -f- 2K 2 SO 4 + 5H 2 O + 3 ; 2KMnO 4 -f 3H 2 SO 4 = 2MnSO 4 + K 2 SO 4 + 3H 2 O + 0, Chromates become green in color; permanganates are decolorized. Carbon dioxide from carbonates, effervescence (see 2, page 68). Carbon monoxide (burns with bluish flame) from organic substances, usually with blackening of the substance and the evolution of carbon dioxide and sulphurous anhydride, as in the case of tartaric acid. Carbon monoxide together with carbon dioxide is evolved from oxalic acid (without blackening, see 6, page 88). From cyanides, ferrocyanides, etc. (Cyanides, page 110). 9* 102 In presence of the latter a transitory bluish coloration appears. 3. If a portion of the substance is heated with concentrated sulphuric acid and alcohol, there is produced, in the presence of acetates, ethyl acetate, Avhich may be recognized by its apple-like odor (see 6, page 87). If the alcohol is ignited, the flame assumes a green color in presence of boric acid (see 5, page 66). IV. SOLUTION AND FUSION. SOLIDS must necessarily be in solution in order to make the tests in the Wet Way. The method employed in dissolving the solid depends upon the nature of the substance; with this in view, substances may be divided into the following five groups : 1. Oxides and salts (in general). 2. Metals and alloys. 3. Sulphides (of the heavy metals). 4. Cyanides (of the heavy metals). 5. Silicates. A distinction may be made between solution and fusion. Many salts cannot be directly dissolved in water or acids, but must undergo a special treatment to separate the acids from the bases, as in the case of barium sulphate ; the sul- phuric acid is separated from the barium by fusing with sodium carbonate. By fusion, new compounds of the bases and acids are obtained which are soluble in water or acids. In case a substance is not entirely soluble in any one of the solvents, it should be treated by each solvent in turn, and the solutions analyzed separately, as two simple analyses are more quickly made than one complex one. For example, the substance is first boiled with water, the solution obtained is filtered off and set aside for examination, any residue insoluble in hot water is treated with nitric acid and the solution ex- amined separately, any residue remaining after treatment with nitric acid is treated with hydrochloric acid. By this pro- 103 104 cedure a more distinct insight into the nature of the sub- stance to be analyzed is obtained. Hard bodies, minerals, etc., must be pulverized in a por- celain or agate mortar before they are dissolved. Very hard minerals are first crushed in a steel mortar, and the coarse powder thus obtained is afterwards pulverized in an agate mortar. It is advisable to sift the powder through a linen cloth (previously washed and dried), remove the coarser particles and again pulverize them, and repeat the operation. If the substance to be analyzed be an organic compound or contain organic material (as shown by the preliminary ex- amination), the organic substance must be destroyed by igni- tion and the residue then dissolved in water (removing by filtration any separated carbon). 1. DISSOLVING OXIDES AND SALTS. (a) A portion of the substance to be dissolved is heated in a test-tube, with water. In case it enters into solution, a larger portion is dissolved and the liquid employed in testing for bases and acids. If the substance is apparently undis- solved, it is separated by filtration and the filtrate evaporated to dryness, to ascertain whether any of the original substance entered into solution. (6) Substances insoluble in water are further tested as to their solubility in dilute nitric acid. An excess of nitric acid should be avoided, as many nitrates soluble in water are insoluble in excess of strong acids. On dissolving oxides with nitric acid, nitrates are formed, and on dissolving salts, nitrates of the bases are produced with the liberation of the acids which were in combination ; for example : Ca 3 (P0 4 ) 2 + 6HNO S *= 3Ca(NO 3 ) 2 + 2H 3 PO 4 ; CuC0 3 + 2HX0 3 - Cu(X0 3 ) 2 + CO 2 + H 2 O. 105 Thus the presence of volatile acids becomes evident : Carbonic acid : effervesces ; odorless gas ; renders calcium hydroxide solution turbid (see 2, page 68). Hydrocyanic acid : odor of bitter almonds ; forms am- monium sulphocyanide with ammonium sulphide (see 5, page 78). Hydrogen sulphide : recognizable by its odor ; blackens paper saturated with solution of plumbic acetate (see 4, page 81). Sulphurous acid : odor of burning sulphur ; colors potas- sium iodate starch paper blue (see 2, page 61). Under certain conditions the presence of iodine, bromine, or chlorine may become evident (see 2, page 100). In using nitric acid as a solvent, acids which are soluble with difficulty may separate : Boric acid, crystalline, easily soluble in hot water ; silicic acid, gelatinous. Reddish-brown fumes of nitrogen dioxide result from the processes of oxidation ; for example, when mercurous com- pounds are converted into mercuric compounds : Hg 2 + 6HN0 3 = 2Hg(N0 3 ) 2 + 2NO 2 + 3H 2 O. These oxidations may interfere with the results of the analysis, especially when compounds of mercury are present. After the oxidation with nitric acid it is impossible to determine the original condition of oxidation of the salt ; for example, in the case of mercury, after oxidation it cannot be ascertained whether the salt was present originally as a mercurous or a mercuric salt. Salts of mercury which are insoluble in water or in moderately warm dilute nitric acid are decom- posed by boiling in sodium hydroxide (compare page 106, e). Compounds of arsenic should be dissolved, when pos- sible, in hydrochloric acid, in order to prevent the conver- sion of arsenious acid into arsenic acid. Plumboso-plumbic oxide (red lead) when treated with dilute nitric acid is decom- 106 posed into soluble plumbic nitrate and insoluble, brown lead dioxide : Pb 3 O 4 -f 4HNO 3 = 2Pb(NO 3 ) 2 + PbO 2 -f 2H 2 O. The latter is converted into plumbic chloride by concentrated hydrochloric acid. (c) Those substances which are insoluble in dilute nitric acid must be treated with concentrated hydrochloric acid. If in dissolving the substance in hydrochloric acid chlorine gas is evolved, peroxides and similar compounds, such as manga- nese dioxide, chromic acid, or permanganic acid, are present : 2CrO 3 + 12HC1 =* O 2 C1 6 + C1 6 +6H 2 O ; Mn 2 O 7 -f 14HC1 = 2MnCl 2 + C1 10 + 7H 2 O. Lead dioxide is converted into plumbic chloride, which crys- tallizes as the solution cools ; it is best decomposed with sodium carbonate (page 106, e). (d) Many compounds insoluble in nitric acid or in hydro- chloric acid are soluble in nitro-hydrochloric acid (aqua regia). In dissolving with nitro-hydrochloric acid chlorine (l; is liberated, which is the active agent in effecting solution : 3HC1 + HNO 3 = C1 3 + NO + 2H 2 O. Nitro-hydrochloric acid is prepared by mixing about three volumes of concentrated hydrochloric acid with one volume of concentrated nitric acid ; the reaction takes place upon the application of heat. When nitro-hydrochloric acid is em- ployed as a solvent, oxidation necessarily occurs if the sub- stance is capable of being oxidized, as, for example, with compounds of mercury. (e) Many compounds that are insoluble in water and in acids are decomposed by boiling or fusing with carbonates of the alkalies, that is, they are converted into soluble com- 1 Besides (NOC1) nitrosyl chloride and (NO 2 C1) nitroxyl chloride. 107 pounds. Among them are plumbic sulphate, the sulphates of the alkaline earths, plumbic chloride, plumbic iodide, stannic oxide, etc. Of the sulphates, plumbic sulphate and calcium sulphate are easily decomposed by boiling in a solution of sodium carbonate. Precipitated strontium sulphate is also decom- posed in the same manner, although with more difficulty. Precipitated barium sulphate is only partly decomposed by boiling with sodium carbonate solution. These sulphates (as well as minerals) are readily decomposed by being fused with from four to six parts of sodium potassium carbonate.^ In these decompositions the acid of the substance fused unites with the alkalies, and the base is converted into a carbonate ; for example, with BaSO 4 and NaKCO 3 the compounds NaKSO 4 , soluble in water, and BaCO 3 , soluble in acids, are formed : NaKCO 3 + BaSO 4 = NaKSO 4 + BaCO 3 . The fused mass is completely extracted with hot water and the insoluble residue (after separation by filtering) is dissolved in hydrochloric acid or nitric acid. The aqueous solution is to be examined for the acid, and the acid solution for the base. Plumbic chloride and plumbic iodide, etc., when boiled with a solution of sodium carbonate, are decomposed respec- tively into chloride and iodide of sodium and plumbic carbonate : 1 The double salt NaKCO 3 fuses more easily than the sodium or potas- sium salt alone. The fusion is best made in a platinum crucible, as porcelain is attacked by the alkali carbonates. The following substances should never be fused in a platinum crucible : potassium and sodium hydroxide, nitrates and cyanides of the alkalies, metals and metallic sulphides, or any substance from which a metal may be obtained by reduction or substances from which chlorine may be evolved. 108 PbCl 2 + Na 2 CO 3 = 2NaCl + PbCO 3 . (Plumbic carbonate is slightly soluble in sodium carbonate.) Stannic oxide (cassiterite) when fused with a carbonate of an alkali is converted into a stannate of the alkali, which is soluble in water and in hydrochloric acid : SnO 2 + K 2 CO 3 = SnO(OK) 2 + CO 2 ; SnO(OK) 2 + 6HC1 = SnCl 4 + 2KC1 + 3H 2 O. Fusion is continued until carbon dioxide ceases to be evolved. As stannic oxide is acted upon only with great difficulty by sodium carbonate, it is best fused in a silver crucible with sodium or potassium hydroxide, (1; and the fused mass treated with water and hydrochloric acid, as mentioned above. Many substances are unacted upon by the carbonates of the alkalies, but are readily decomposed on being foiled with sodium or potassium hydroxide ; for example, mercury and silver compounds. An oxide of the metal is formed, while {he acid remains in solution in combination with the alkali. The oxide after being washed is dissolved in nitric acid. Hg 2 Cl 2 -f 2NaOH = Hg 2 O -f 2XaCl + H 2 O. In dissolving compounds of mercury cold dilute nitric acid should be used, in order to avoid the oxidation of mercurous salts to mercuric salts. (Mercuric iodide, which partly redis- solves in a carbonate of an alkali, i.e., in the iodide of the alkali which is formed, should be dissolved in nitro-hydro- chloric acid.) (/) Compounds of fluorine (for example, fluor spar) are decomposed by being gently heated with concentrated sul- phuric acid in a platinum crucible : 1 Or it may be fused in a porcelain crucible with three parts of sodium carbonate and three parts of sulphur to one part of the substance, and the fused mass, after cooling, extracted with water. The yellow solution con- tains the tin as sulphostannate, Na 2 SnS 3 ; the insoluble residue, containing sulphides, is to be examined further according to 3, page 111. 109 CaF 2 + H 2 S0 4 - CaS0 4 + 2HF. The hydrofluoric acid is recognized by its etching glass (see 4, page 67); the residue in the crucible, consisting of sul- phates, is dissolved in hydrochloric acid or, if necessary, fused with sodium carbonate. Silicates containing fluorine, if treated in this manner, yield silicon fluoride, according to the reaction : 2CaF 2 + SiO 2 H- 2H 2 SO 4 = SiF 4 + 2CaSO 4 + 2H 2 O. If the evolved gas be conducted through a glass tube moist- ened with water, silicic acid together with hydrofluosilicic acid is produced : 3SiF 4 + 3H 2 = 2H 2 SiF 6 + H 2 SiO 3 . The silicic acid will appear, either directly or on drying the tube, in the form of a white coating (see 5, page 68). (g) Chromic oxide, chromite, aluminium oxide, and ferric oxide are best fused by mixing them with ten parts of acid potassium sulphate. If the heat applied is not too great, neutral sulphates (together with basic salts) are formed : A1 2 3 + 6KHS0 4 * A1 2 (S0 4 ) 3 + 3K 2 SO 4 + 3H 2 O, which, on cooling, may be dissolved by water or hydrochloric acid. Chromite is best fused with acid potassium sulphate, and the fused mass obtained again fused with potassium chlorate and potassium carbonate, to convert the chromic oxide into chromic acid : Cr 2 (S0 4 ) 3 + 3K 2 C0 3 = Cr 2 O 3 + 3K 2 SO 4 + 3CO 2 ; Cr 2 O 3 + 2K 2 CO 3 + KC1O 3 = 2K 2 CrO 4 + KC1 + 2CO 2 . The fused mass yields potassium chromate when extracted with water ; the residue, consisting of ferric oxide (with some chromic oxide) is dissolved in hydrochloric acid. (h) Carbon (charcoal, graphite) and sulphur are recognized by their appearance and their behavior when heated. 10 110 2. THE DISSOLVING OF METALS AND ALLOYS. Metals and alloys are treated with concentrated nitric acid and heated until red vapors cease to be produced upon the further addition of acid. The excess of nitric acid (which would interfere with the solubility of the nitrates in water) is evaporated on the water-bath, and the residue dissolved with water and a little nitric acid. Most of the metals enter into solution as nitrates, gold, platinum, etc., remain unchanged, and tin and antimony remain as oxides or hydroxides. u) In the 'presence of tin or antimony arsenic may be found in the residue, in the arsenic condition. The residue, after thorough washing, is digested with yellow ammonium sul- phide, whereby tin, antimony, and arsenic enter into solution as sulpho-salts : . Sn(OH) 4 -f 3(NH 4 ) 2 S = (NH 4 ) 2 SnS 3 + 4NH 3 + 4H 2 O ; Sb 2 O 3 + 6(NH 4 ) 2 S + S 2 = 2(NH 4 ) 3 SbS 4 + 6NH 3 + 3H 2 O ; Sb 2 O 4 + 7(NH 4 ) 2 S + S 2(NH 4 ) 3 SbS 4 + 8NH 3 -f- 4H 2 O ; Sb 2 O 5 + 8(NH 4 ) 2 S = 2(NH 4 \SbS 4 + 10NH 3 -f 5H 2 O ; AsA + 8(NH 4 ) 2 S = 2(NH 4 ) s AsS 4 + 10NH 3 + 5H 2 O. If an insoluble residue remain, it is again treated with nitric acid ; if it still fail to dissolve, it is finally treated with nitro-hydrochloric acid, which dissolves gold and platinum as chlorides : Au -f- 3HC1 -|- HNO S = AuCl s -f- NO -f 2H 2 O ; 3Pt + 12HC1 + 4HNO 3 = 3PtCl 4 + 4NO + 8H 2 O. 1 In dissolving metals in nitric acid different oxides of nitrogen are pro- duced, depending upon the concentration of the acid employed. With nitric acid of 1.42 specific gravity NO 2 is produced ; with an acid of 1.35 specific gravity, principally N 2 O 3 ; with an acid of 1.2 specific gravity NO ; and with an acid of 1.1 specific gravity N 2 O. (With an acid of greater dilution ammonia is produced.) Nitric acid of 1.1 or less specific gravity is decomposed only by the more strongly positive metals ; for example, Zn and Fe. in 3. SULPHIDES OF THE HEAVY METALS. The sulphides of the heavy metals generally possess a metallic lustre ; like the metals they are treated with concen- trated nitric acid, whereby most of them are dissolved as nitrates : CuS + 4HNO 3 Cu(NO 3 ) 2 + S + 2NO 2 + 2H 2 O. The procedure is as given under 2, page 110. The sulphur which separates first is oxidized by the nitric acid to sulphuric acid. The insoluble residue, in addition to the oxides of tin, antimony, and arsenic, may contain PbSO 4 , BiONO 3 (formed on treating the nitrate with water), and HgS. This residue is treated with yellow ammonium sulphide, which dissolves tin, antimony, and arsenic. Any residue remaining is filtered off and treated with nitric acid to dissolve the lead and bis- muth (which at this stage may be found again as sulphides), and any residue of mercuric sulphide is collected on a filter and dissolved by nitro-hydrochloric acid : 3HgS + 6HC1 + 2HNO 3 = 3HgCl 2 + 2NO + 4H 2 O -f S 3 . Finally silicious gangue, barite, etc., may remain, which should be examined according to 5 (page 113) and 1 (page 104) respectively. The sulphides are easily recognized by their appearance, and also by their behavior in the preliminary examination. 4. CYANIDES. The simple cyanides, which are insoluble in water, may be decomposed into chlorides and hydrocyanic acid, by boiling with concentrated hydrochloric acid. Argentic cyanide and mercuric cyanide, in which the cyan- ogen cannot be detected by the ordinary methods, may be readily recognized by their behavior when heated, as they separate into metal and cyanogen. If the cyanide is heated 112 in a narrow glass tube, the escaping cyanogen may be ignited, burning with a pinkish-lavender flame. The gas is also rec- ognizable by its odor of bitter almonds. Mercuric cyanide may be decomposed by dissolving in water and passing hy- drogen sulphide through the solution ; mercuric sulphide is precipitated and hydrocyanic acid enters into solution. The insoluble compounds of ferrocyanogen and ferricy- anogen are decomposed by boiling with sodium carbonate or sodium hydroxide; sodium ferrocyanide and ferricyanide respectively are formed, together with an insoluble carbonate or an oxide of the metal : Pb 2 Fe(CN) 6 + 2X0,00, = Na 4 Fe(CN) 6 + 2PbCO 3 ; Cu 2 Fe(CX) 6 + 4XaOH = JSa 4 Fe(CJN T ) 6 + 2CuO + 2H 2 O. The aqueous solution is filtered and the filtrate tested for the acid, while the carbonates or oxides are dissolved in dilute nitric acid. If sodium hydroxide is employed as the decom- posing agent, lead, zinc, and aluminium, and also arsenic, antimony, and tin, may enter into solution. In such cases a portion of the alkaline solution is tested for lead, zinc, and aluminium by saturating the solution with hydrogen sul- phide, thereby precipitating the first two metals as sulphides and the last as hydroxide. The filtrate from any precipitate which may have been produced, or the clear solution if no precipitate was produced by hydrogen sulphide, is acidulated with hydrochloric acid to precipitate arsenic, antimony, and tin as sulphides. If sodium hydroxide is used in decomposing the ferri- cyanide, sodium ferricyanide is formed, providing the metallic oxide produced in the operation is not further oxidizable : Cu s (Fe(CN) 6 > 2 + 6NaOH - 2Na 3 Fe(CN) 6 + 3CuO + 3H 2 O. If, however, the separated oxide is capable of further oxida- tion, this oxidation takes place, accompanied by the reduction of the sodium ferricyanide to sodium ferrocyanide : 113 Fe 3 (Fe(CN) 6 ) 2 + SNaOH = 2Na 3 Fe(CN) 6 + 3Fe(OH) 2 + 2NaOH ; 2Na 3 Fe(CN) 6 + 3Fe(OH) 2 + 2NaOH = 2Na 4 Fe(CN) 6 + Fe 2 (OH) 6 + Fe(OH) 2 . Consequently, in such cases to detect the acid the substance is fused, whenever possible, with sodium carbonate. To detect alkalies in ferrocyanogen and ferricyanogen com- pounds, the latter are decomposed into sulphates, carbon monoxide, and ammonium sulphate, by being heated with concentrated sulphuric acid : CuK 2 Fe(CN) 6 + 6H 2 SO 4 + 6H 2 O = FeSO, + CuSO 4 + K 2 SO 4 + 6CO + 3(N.H 4 ) 2 SO 4 . 5. SILICATES. Before silicates can be analyzed they must be finely pul- verized (page 104). (a) Silicates soluble in water or silicates that may be decom- posed by acids are best decomposed by being boiled with con- centrated hydrochloric acid ; by this procedure silicic acid and chlorides of the respective metals are formed ; for example : K 2 Si0 3 + 2HC1 = H 2 Si0 3 + 2KC1. Boiling is continued until complete decomposition has taken place, and no gritty particles are detected on stirring with a glass rod. The solution is then evaporated to the dryness of dust on a water-bath (see 1, page 69), to convert the solu- ble silicic acid into insoluble amorphous silicic acid. The dry residue is then moistened with a little concentrated hydro- chloric acid, to convert any basic chlorides (of Fe, Al, Mg, etc.) into neutral chlorides, thereby rendering them soluble ; finally the chlorides of the bases are extracted with water and dilute hydrochloric acid. (b) Silicates that are not decomposed by acids must either be fused with a carbonate of an alkali or decomposed by h 10* 114 hydrofluoric acid. To determine which method should be employed, the silicate is tested for the presence of an alkali. For this purpose a small portion of the powdered silicate, moistened with hydrochloric acid, is placed on a platinum wire and held in the non-luminous flame of a Bunsen burner, to observe whether a color is imparted to the flame (sodium, yellow ; potassium, violet). If alkalies are absent the method of decomposition by means of sodium carbonate is to be em- ployed ; whereas if alkalies are present, in order to test for them, the silicate must be decomposed with hydrofluoric acid. (c) In case the fusion is to be made with sodium carbonate (preferably with sodium potassium carbonate), one part of the finely pulverized substance is thoroughly mixed with six parts of sodium potassium carbonate, placed in a platinum cruci- ble, and the mixture fused by means of the blast-lamp. The silicate is decomposed by the carbonate of an alkali, with the production of a silicate of the alkali (or at least silicates that are decomposed by acids) and a carbonate of the metal : CaSi0 3 + NaKCO s - NaKSiO 8 + CaCO 3 ; CaSi 2 O 5 + 2NaKCO 3 = 2NaKSiO 8 + CaCO 3 + CO 2 . On disintegrating the fused mass with hydrochloric acid, ac- cording to a, page 113, silicic acid remains insoluble, and the chlorides of the metals together with sodium chloride and potassium chloride enter into solution. (d) In using hydrofluoric acid as a solvent the finely pulver- ized substance is placed in a platinum crucible, and treated with the pure acid (1) until a thin paste is formed. The mixture is stirred with a platinum wire (not with a glass rod) and digested, at a very gentle heat, until the substance is completely dissolved. By this treatment the silicates are converted into fluosilicates : 1 The hydrofluoric acid must be free from alkalies, and, when possible, freshly distilled in a platinum still. 115 CaSiO 3 + 6HF = CaSiF 6 + 3H 2 O ; CaSi A + 12HF = CaSiF 6 + H 2 SiF 6 + 5H 2 O. When completely dissolved concentrated sulphuric acid is added and heat applied, gently at first, but afterwards more strongly, to drive off the excess of acid. The sulphuric acid converts the fluosilicates into sulphates, while hydrofluosilicic acid is evolved : CaSiF 6 + H 2 SO 4 = H 2 SiF 6 + CaSO 4 . The residue of sulphates is dissolved in water and a little hydrochloric acid. When this method is employed to decompose silicates con- taining barium, strontium, or calcium^ it is necessary espe- cially with barium and strontium, -to afterwards fuse the residue containing the barium, strontium, or calcium sulphate with a carbonate of an alkali (page 106)o In mineral analyses it is often of interest to ascertain whether the minerals contain, in addition to the silicates not decomposable by acids, others that may be decomposed, thus making separation possible= With this in view, after having mechanically separated the gangue and any other impurities from the mineral proper, it is finely pulverized, treated with hydrochloric acid, the solution evaporated to dryness as above described (page 113, a), and the chlorides resulting from the decomposed silicates dissolved in water,, The insoluble residue, which may contain silicic acid and undecomposed silicates, is boiled with sodium carbonate, which dissolves the silicic acid derived from the decomposed silicate. After acidulating the sodium carbonate solution with hydrochloric acid, evaporating to dryness, and extracting with hot water, the silicic acid re- mains as a light, white powder. If a residue remain after boiling a second time with sodium carbonate, it is to be con- sidered an undecomposable silicate, which is to be further tested according to 6, c, and d, page 114. V. DETECTION OF BASES IN THE WET WAY. IF the substance to be analyzed is a solid it is to be dis- solved, as before described (page 104). To test for bases in organic substances the latter should be incinerated and the bases extracted from the ash by water or acids. Organic acids, etc., interfere with a number of the reactions used in the detection of bases. The reaction of solutions to be examined should be tested with litmus and turmeric paper to ascertain whether they are neutral, acid, or alkaline. A number of substances may be present in acid solutions, which, in neutral solutions, may be disregarded ; for example, in the third group, acid solutions must be tested for phosphates and oxalates, whereas if the solu- tions are neutral the tests for these acids need not be made. Regarding combinations that may arise during the exam- ination of alkaline solutions see 6, page 120. 116 PRECIPITATION OF THE DIFFERENT GROUPS. To separate the bases into groups the following group reagents are employed : 1. Hydrochloric acid. 2. Hydrogen sulphide. 3. Ammonium hydroxide. 4. Ammonium sulphide. 5. Ammonium carbonate. By each of these reagents a series of bases called a group is precipitated. Bases that are not precipitated by group reagents are classed as a sixth group. (Table L, pages 118 and 119o) The rare elements are not considered in this plan. 117 118 TABLE I. GROUP PRECIPITATIONS. GROUP I. Metals precipitated by Hy- drochloric Acid. GROUP II. Metals precipitated by Hy- drogen Sulphide. GROUP HI. Metals precipitated by Am- monium Hydroxide. as white, curdy AgCl, argentic chloride. Mercurous salts, as white, pulverulent Hg 2 Cl 2 , mercurous chlo- ride. Lead, as white PbCl 2 , plumbic chloride. Lead, as black PbS, plumbic sulphide. Mercuric salts, as black HgS, mercuric sulphide. Copper, as black CuS, cupric sulphide. Bismuth, as brownish-black Bi.>S 3 , bismuthous sul- " phide. Stannous salts, as brownish-black SnS, stannous sulphide. Cadmium, as yellow CdS, cadmium sulphide. Arsenic, as yellow arsanious sulphide (mixed with sulphur if precipitated from arsenic acid solutions). Stannic salts, as yellow SnS 2 , stannic sulphide. Anlimonious salts, as orange-red SbgSg, antimonious sul- phide. Antimonic salts, as orange-red Sb 2 S 6) antimonic sulphide (together with and sulphur). Gold, as black Au 2 Ss, auric sulphide. Platinum, as brownish-black PtS 2 , platinic sulphide. Iron, as reddish- browu Fe 2 (OH) 6) ferric hydroxide. Chromium, as dirty-green Cr 2 (OH) 6 , chromic hydrox- ide. Aluminium, as white, gelatinous A1 2 (OH) 6 , aluminium hy- droxide. In presence of phosphoric acid iron and alumin- ium respectively are pre- cipitated as phosphates, thus : Iron, as white Fe 2 (PO 4 ) 2) ferric phosphate. Aluminium, as white A1 2 (PO.|)2, aluminium phosphate. In presence of phosphoric acid or oxalic acid cal- cium, strontium, and barium are precipitated as phosphates or oxa- lates, as white Ca 3 (PO 4 ) 2 , SrC 2 O 4 , etc. Magnesium in the pres- ence of phosphoric acid is precipitated in this group as white MgNH 4 PO 4 . ammonium magnesium phosphate. In presence of iron, man- ganese may be precipi- tated as white, changing to brown, Mn(OH) 2 , manganous hy- droxide. 119 TABLE I. GKOUP PKECIPITATIONS. Continued. GROUP IV. Metals precipitated by Am- monium Sulphide. GROUP V. Metals precipitated by Am- monium Carbonate. GROUP VI. For which there is no Spe- cial Group Reagent. Manganese, Barium, Magnesium. as light-salmon-colored MnS, manganous sul- phide. as white BaCO 3 , barium carbonate. Strontium, Potassium. Sodium. Zinc, as white ZnS, zinc sulphide. as white SrCO 3 , strontium carbon- ate. Lithium. Ammonium. Nickel, Calcium, as black as white NiS, nickelous sulphide. CaCO 3 , calcium carbon- ate. Cobalt, as black CoS, cobaltous sulphide. 120 If on the addition of the reagent a precipitate is formed, it is filtered off and carefully washed. The filtrate should be tested to ascertain whether the precipitation was complete, that is, whether no precipitation takes place on further addi- tion of the reagent. The precipitate must be collected only on properly-cut filters which fit closely to the inner surface of the funnel, and the liquid in which the precipitate is sus- pended should be poured into the filter down a glass rod. The precipitates should be thoroughly washed before proceed- ing with the examination. This is not only good preliminary practice for quantitative work, but is absolutely necessary to obtain exact results in qualitative analysis. Concentrated solutions should be diluted with water before the examination is commenced. This dilution may cause turbidity, in consequence of the formation of basic salts or oxychlorides of bismuth, antimony, or mercury. These, how- ever, may be redissolved by the addition of a little nitric or hydrochloric acid. FIRST GROUP. (a) Neutral or acid solutions are treated with a few drops of dilute hydrochloric acid. There will be precipitated : Silver, as white, curdy AgCl, argentic chloride. Mercurous salts, as white, pulverulent Hg 2 Cl 2 , mercu- rous chloride. Lead, as white, crystalline PbCl 2 , plumbic chloride. The latter is incompletely precipitated, as it is slightly soluble in water ; therefore a test for it must also be made in the second group. The solution in which the precipitation takes place must be cold, as plumbic chloride is easily soluble in hot water and might remain in solution ; moreover, small quantities of mercurous salts might be overlooked in the 121 presence of nitric acid, as, when hydrochloric acid and nitric acid are both present and the solution is warm, mercurotis chloride is transformed into soluble mercuric chloride. Furthermore, it should be observed whether the precipitate redissolves on the addition of an excess of the hydrochloric acid. On the addition of dilute hydrochloric acid, dilute solutions of compounds of bismuth yield a white precipitate of BiOCl, bismuth oxychloride, which on the further addition of hydrochloric acid is redissolved as BiCl 3 , bismuthous chlo- ride. Compounds of antimony, especially K(SbO)C 4 H 4 O 6 , potassium antimonious tartrate, with dilute hydrochloric acid form SbOCl, antimonious oxychloride, which is soluble in an excess of the acid as SbCl 3 , antimonious chloride. KHC 4 H 4 O 6 , acid potassium tartrate, if it should have sepa- rated, would be redissolved on the further addition of hydro- chloric acid : KHC 4 H 4 6 + HC1 = H 2 C 4 H 4 6 -fKCl. Furthermore, there may be precipitated in the first group : boric acid (crystalline), organic acids, and sulphur. (Sulphur separates from hyposulphites and polysulphides ; Na 2 S 2 3 + 2HC1 = 2NaCl + S -f- SO 2 + H 2 O ; (NH 4 ) 2 S 3 + 2HC1 = 2NH 4 C1 + S 2 + H 2 S. In the first case sulphurous anhydride, in the latter case hy- drogen sulphide, is evolved with the sulphur. Polysulphides are always alkaline in reaction.) Attention should be paid to any gases evolved on treat- ment with hydrochloric acid (with reference to the manner of distinguishing them see pages 99 to 102 and 105). Sul- phurous anhydride must be driven off by heating ; other- wise, on the addition of hydrogen sulphide, separation of sulphur would occur (together with the formation of pen- tathionic acid) : 5S0 2 + 5H 2 S = H 2 S 5 6 + S 5 + 4H 2 O. P 11 122 Chlorine, nitrogen dioxide, etc., should also be expelled by heating the liquid. (6) Alkaline solutions should be treated with hydrochloric acid until acid in reaction, and any formation of precipitates or evolution of gases observed. From alkaline solutions there may be separated : 1. Sulphur and sulphides of the metals, accompanied by the evolution of hydrogen sulphide. The sulphides are the following sulpho-acids : As 2 S 3 , As 2 S 5 , Sb 2 S 3 , Sb 2 S 5 , SnS 2 : they should be tested according to the directions given in the chapter treating of them under the second group (see B, page 135). Under certain conditions CuS, HgS, and NiS might also be encountered at this stage. The filtrate from the sulphur or the sulphides which have separated may be examined directly for the metals of the fifth and sixth groups. 2. Cyanides of the heavy metals (which were dissolved in cyanides of the alkalies), with the evolution of hydrocyanic acid. Concentrated hydrochloric acid is added to the liquid containing the precipitate and the whole heated. The cyanides are thereby converted into chlorides, which finally dissolve, argentic chloride alone remaining undissolved. The solution is then examined for the presence of metals of the second, third, and subsequent groups ; argentic chloride in the residue is confirmed by testing its solubility in ammonium hydroxide. 3. Silicic acid : gelatinous ; should be confirmed in the bead of microcosmic salt (see 4, page 70). The solution, together with the precipitate, is treated with an excess of hydrochloric acid, and evaporated to dryness on the water- bath to render the silicic acid insoluble. The residue is extracted with water and a little hydrochloric acid (page 113, a), and the filtrate examined for bases. It usually contains nothing but the alkalies. 123 4. Precipitates of plumbic hydroxide, aluminium hydrox- ide, chromium hydroxide, and zinc hydroxide may be formed, but on acidifying with hydrochloric acid will immediately disappear, being converted into soluble chlorides. SECOND GROUP. Hydrogen sulphide is conducted into the acid nitrate ob- tained from the precipitate of the first group, or into the solution in which hydrochloric acid failed to produce a pre- cipitate, until a distinct odor of the gas is observable in the liquid. There will be precipitated : Lead, as black PbS, plumbic sulphide. Mercuric salts, as black HgS, mercuric sulphide. Copper, as black CuS, cupric sulphide. Bismuth, as brownish-black Bi 2 S 3 , bismuthous sulphide. Gold, as black Au 2 S 3 , auric sulphide. Platinum, as brownish-black PtS 2 , platinic sulphide. Cadmium, as yellow CdS, cadmium sulphide. Arsenious compounds, as yellow As 2 S 3 , arsenious sulphide. Arsenic compounds, as yellow As 2 S 3 , arsenious sulphide (with sulphur). Antimonious compounds, as orange-red Sb 2 S 3 , antimoni- ous sulphide. Antimonic compounds, as orange-red Sb 2 S 5 , antimonic sulphide (together with Sb 2 S 3 and S). Stannous compounds, as brownish-black SnS, stannous sulphide. Stannic compounds, as yellow SnS 2 , stannic sulphide. From solutions containing hydrochloric acid, when hydro- gen sulphide is not present in sufficient quantity, lead is pre- cipitated as red Pb 2 SCl 2 , plumbic sulphochloride, which is 124 converted by further addition of hydrogen sulphide into black PbS. In solutions of mercuric salts white precipitates of double salts (for example, Hg 3 S 3 Cl 2 ) are formed, which on continuing the addition of hydrogen sulphide become yellow, then brown, and finally are converted into black HgS. Arsenious acid is precipitated at once, arsenic acid gradually ; the precipitation is accelerated, however, by heating (see 1, page 26). Sulphur may also separate when hydrogen sulphide is introduced into the solution. This separation may be caused by: 1. Chlorine, bromine, iodine, nitrous acid, nitrogen dioxide, etc. (in consequence of their oxidizing action upon hydrogen sulphide) ; for example : N 2 3 -f H 2 S = 2M) + H 2 + S. On passing hydrogen sulphide into solutions containing an excess of nitric acid or nitro-hydrochloric acid, sulphur is separated. The excess of acid should be driven off by evaporation and, after diluting with water, the introduction of hydrogen sulphide should be repeated. 2. Sulphurous acid (page 121). 3. Ferric salts, in consequence of their reduction to ferrous salts : Fe 2 Cl 6 + H 2 S = 2FeCl 2 + 2HC1 -f S. Decolorization of the solution results from the reduction. 4. Chromic acid and chromates, in consequence of their reduction to chromic salts : 2H 2 CrO 4 -f- 3H 2 S + 6HC1 = Cr 2 Cl 6 + 8H 2 O + S 3 . The solution changes in color from yellow to green. By repeated introduction of hydrogen sulphide accompanied by renewed additions of hydrochloric acid, the chromic acid is completely decomposed. If the acid is not added in sufficient 125 quantity, a precipitate is formed consisting either of green chromic hydroxide : 2H 2 CrO 4 + 3H 2 S = Cr 2 (OH) 6 + S 3 + 2H 2 O, or of brown chromium chromate : 3H 2 Cr0 4 + 3H 2 S = (CrO) 2 CrO 4 + S 3 + 6H 2 O. 5. Permanganic acid and permanganates, in consequence of their reduction to manganous compounds : 2HMnO 4 + 5H 2 S + 4HC1 2MnCl 2 + S 5 + 8H 2 O. The purplish-red solution is decolorized. The procedure is the same as in 4, page 124. (If the hydrochloric acid is not added in sufficient quantity, brown precipitates are formed.) THIRD GROUP. From the nitrate of the second group, or from the solution in which hydrogen sulphide failed to produce a precipitate, the hydrogen sulphide is expelled by boiling. A small quantity of nitric acid is added and the solution warmed to oxidize the bases, ammonium chloride and afterwards ammo- nium hydroxide (the latter in not too great excess) are added, and the solution is boiled until the odor of ammoniacal gas can no longer be detected. There will be precipitated : Iron, as reddish-brown Fe 2 (OH) 6 , ferric hydroxide. Chromium, as dirty-green Cr 2 (OH) 6 , chromic hydroxide. Aluminium, as white, gelatinous A1 2 (OH) 6 , aluminium hydroxide. In the presence of phosphoric or oxalic acids, Ferric phosphate, Fe 2 (PO 4 ) 2 (white). Aluminium phosphate, A1 2 (PO 4 ) 2 (white). Phosphates and oxalates of col- 1 Ca 3 (PO 4 ) 2 , etc. (white). eium, strontium, barium, ) CaC 2 O 4 , etc. (white). Ammonium magnesium phosphate, MgNH 4 PO 4 (white). 126 In presence of iron some manganese may be precipitated as Mn(OH) 2 , manganous hydroxide. The hydrogen sulphide must be expelled, so that, on the addition of the ammonium hydroxide, ammonium sulphide may not form and thereby precipitate the fourth group with the third. ' By means of the nitric acid ferrous salts are converted into ferric salts ; in presence of ammonium chloride the fer- rous salts are not precipitated, or are precipitated only incom- pletely. If the oxidation is not complete, a greenish precip- itate is obtained in the presence of ferrous salts, which, when exposed to the air, oxidizes and gradually changes to black and finally to reddish-brown (ferric hydroxide). In solutions containing silicates the ammonium hydrox- ide may precipitate gelatinous H 2 SiO 3 , silicic acid. H 2 SO 4 may possibly be formed (by the oxidation of the H 2 S passed into the solution), and precipitate barium and strontium as sulphates. Ammonium chloride is added to prevent the pre- cipitation of manganese and magnesium (see 3, page 45, and 1, page 53). The ammonium chloride should be added in excess, but not too great excess, as thereby the precipitation of the fifth group is unnecessarily rendered more difficult. After the addition of the ammonium hydroxide it is neces- sary to boil the liquid until the odor of ammonia disappears, in order to completely precipitate aluminium and chromium (see 1, page 41, and 1, page 42). By this procedure the excess of ammoniacal gas is expelled ; but the boiling should not be continued too long, as the solution may become acid (in con- sequence of the decomposition of NH 4 C1 with the liberation of NH 3 ). 127 FOURTH GROUP. To the filtrate from the third group (to which ammonium hydroxide is again added), or to the solution in which am- monium chloride and ammonium hydroxide failed to produce a precipitate, colorless or slightly yellow ammonium sulphide is added. There will be precipitated : Manganese, as light-salmon-colored MnS, manganous sulphide. Zinc, as white ZnS, zinc sulphide. Nickel, as black NiS, nickelous sulphide. Cobalt, as black CoS, cobaltous sulphide. Nickelous sulphide is slightly soluble in an excess of yellow ammonium sulphide, imparting a brown color to the solution. The nickelous sulphide is completely separated on boiling the solution, especially after the addition of acetic acid. Ammo- nium sulphide might also precipitate iron as ferrous sulphide, in case the iron were held in solution by organic substances. FIFTH GROUP. From the filtrate of the fourth group, or from the solution in which ammonium sulphide failed to produce a precipitate, the ammonium sulphide is expelled by boiling, any sulphur which may have separated is filtered off, ammonium hydrox- ide and ammonium carbonate are added to the filtrate, and the whole is boiled as long as carbon dioxide is evolved. There will be precipitated : Barium, as white BaCO 3 , barium carbonate. Strontium, as white SrCO 3 , strontium carbonate. Calcium, as white CaCO 3 , calcium carbonate. On the addition of commercial ammonium carbonate, acid carbonates soluble in water as, for example, Ca(HCO 3 ) 2 128 are produced (page 50), which, on boiling, are converted into neutral, insoluble carbonates, with the liberation of CO 2 and H 2 O: Ca(HC0 3 ) 2 - CaC0 3 + CO 2 + H 2 O. The carbonates are soluble in an excess of ammonium chlo- ride on long-continued boiling : CaC0 3 + 2NH 4 C1 = CaCl 2 + 2NH 3 + CO 2 + H 2 O. SIXTH GROUP. In this group are classed magnesium, potassium, sodium, and lithium. Ammonium is also classed with this group, but the test for it must be made in the original substance presented for analysis. (With reference to their separation see Separation of the Sixth Group, page 152.) With this group may also be found the ferro- and ferri- cyanides of the alkalies, cobalticyanides of the alkalies, etc., from which the iron and cobalt are not precipitated by the ordinary reagents. Furthermore, aluminium may have re- mained in solution, because of the presence of organic sub- stances. These compounds are to be treated with concen- trated sulphuric acid (page 113) and separated by the regular group precipitations. SEPARATION OF THE BASES CONTAINED IN THE GROUP PRECIPITATES. The group precipitates thus obtained are now examined separately. The precipitates of the second and fourth groups must be examined immediately, as they oxidize when exposed to the air. Precipitates of the third group must be quickly 129 filtered, in order to prevent the formation and precipitation of manganic hydroxide ; for example : 2MnCl 2 (NH 4 Cl) 2 + 4NH.OH + H 2 O + O = Mn 2 (OH) 6 + 8NH 4 C1. If no precipitate is formed in the third group, ammonium sulphide should be added rapidly, to prevent the separation of manganic hydroxide. If arsenic or tin be found in the second group, a portion of the filtrate is reserved for the tests for acids and the other portion is used in testing for bases. The filtrate, including the wash-water, from each group precipitation is reserved for treatment with the succeeding group reagent. If no precipitate is produced by a group reagent, it indicates that the metals of that particular group are absent. The solution is then treated with the succeeding group reagent. SEPARATION OF THE FIRST GROUP. The precipitate produced by hydrochloric acid, after being washed with cold water, is treated while on the filter with hot water ; any plumbic chloride present is dissolved by the hot water, and may be tested for in the filtrate by the addi- tion of sulphuric acid, the formation of a white precipitate of PbSO d indicating the presence of lead. Argentic chloride and mercurous chloride would remain on the filter, undis- solved by the hot water. Any residue remaining on the filter is washed with hot water until free from lead (test washings with H 2 SO 4 ), and then treated with ammonium hydroxide : mercurous chloride is converted into black, in- soluble NH 2 Hg 2 Cl, dimercurous ammonium chloride, while argentic chloride is dissolved as NH 3 AgCl, argent-ammonium chloride. The ammoniacal filtrate is treated with nitric acid until acid in reaction. A white precipitate of AgCl indicates the presence of silver. 130 TABLE II.-SEPAKATION OF THE FIEST GKOUP. The precipitate, which may contain AgCl, Hg 2 Cl 2 , PbCl 2 , is treated, while on the filter, with hot water : Filtrate. PbCl 2 . Treat with H 2 SO 4 : white precipitate of PbSO 4 indi- cates presence of lead. Insoluble Residue. AgCl, Hg 2 Cl 2 . Treat with ammonium hydroxide : Filtrate. Ag (as NH 3 AgCl). Treat with HN0 3 : white, curdy pre- cipitate of AgCl indicates pres- ence of silver. Residue. Hg as black NH 2 Hg 2 Cl indi- cates presence of mercurous salts. To detect small quantities of argentic chloride in the pres- ence of mercurous chloride, the dry mixture of the chlorides is heated in a small glass tube : mercurous chloride will vola- tilize, while argentic chloride remains as a horny mass, which may be further tested on charcoal with the blowpipe. SEPARATION OF THE SECOND GROUP. Of the sulphides of the second group some are basic and others acid in character ; therefore some of them are unacted upon by ammonium sulphide, while others are dissolved as sulpho-salts. Soluble. Arsenious sulphide. Antimonious sulphide. Antimonic sulphide. Stannous sulphide. Stannic sulphide. Auric sulphide. Platinic sulphide. Insoluble. Lead sulphide. Mercuric sulphide. Cupric sulphide. Bismuthous sulphide. Cadmium sulphide. 131 Cupric sulphide is slightly soluble in ammonium sulphide, insoluble, however, in sodium sulphide and in potassium sul- phide. Mercuric sulphide is insoluble in ammonium sulphide, but soluble in sodium sulphide and in potassium sulphide containing free alkali. Stannous sulphide is insoluble in colorless ammonium sulphide, but easily soluble in yellow ammonium sulphide. To ascertain whether sulphides of both basic and acid divisions or of only one division are present, the precipitate produced by hydrogen sulphide is examined regarding its be- havior with ammonium sulphide. For this purpose a small portion of the precipitate in a test-tube is treated with ammo- nium hydroxide and then with yellow ammonium sulphide and gently warmed, any residue remaining undissolved is fil- tered off, and the filtrate is acidified with dilute hydrochloric acid, to ascertain whether a sulpho-salt is present in the solu- tion, that is, whether a (colored) precipitate of sulphide is formed. If none of the sulphides have gone into solution, and only a milkiness, due to the separation of sulphur from the yellow ammonium sulphide, is produced on the addition of the hydrochloric acid, basic sulphides only are present, and the remainder of the precipitate should be treated according to directions given under A, page 132. If the precipitate is completely dissolved, acid sulphides only are present, and the remainder of the precipitate should be examined according to B, page 132. If a portion of the precipitate remain undissolved and another portion enter into solution, the entire remainder of the precipitate is treated with ammonium hydroxide and am- monium sulphide, the insoluble part filtered off and examined according to A, while the solution (filtrate) is treated accord- ing to B. (See Table III a, page 132.) 132 1 .i eS G ^5 11 iSpi o w co s 00 -2V 1 s 12-s < GC 3 X 13 gs 02 i'c z'S B.^JBS^'t 05 <^s f|5|||||||||I^| ! Illllll O J >- ^3 a ill . *- * >& f^SP 00:^: fe-o d c o o M2re A ca ~ = o^ .^Bw^-SgB _- ig^H&gli .22^Sos3?i-uiu-iO.G 5 133 .s 3-i SP-SS 'Si -I 3 .2 Qd s ^! ^^3 c5 1 fifl-SSw 3 ^ A lack argentic sulphide. As argentic hyposulphite is soluble in an excess of a hyposulphite of an alkali, precipitation occurs only when an excess of the argentic nitrate is added. To detect sulphuric acid and other acids in the presence of 164 hyposulphurous acid, the latter acid must be decomposed by gently warming with hydrochloric acid, the liquid filtered, and the tests for sulphuric acid and acids other than hydro- chloric acid made in the filtrate. Phosphoric Add. (Phosphates.) Ammonium chloride, ammonium hydroxide, and mag- nesium sulphate (magnesia mixture), added in turn to a solu- tion of a phosphate, produce a white, crystalline precipitate of ammonium magnesium phosphate. Ammonium molybdate, added in excess with a consider- able quantity of nitric acid, produces a yellow precipitate of ammonium phosphomolybdate. (If arsenic acid is present, it must be completely removed by precipitation with hydrogen sulphide before testing for phosphoric acid. With reference to the behavior of silicic acid with ammonium molybdate, see 3, page 70.) Boric Acid. (Borates.) Turmeric paper dipped in an aqueous solution of boric acid, or of a borate acidified with hydrochloric acid, and warmed until dry, becomes reddish brown in color. Boric acid alone, or borates placed in a dish and moistened with a few drops of concentrated sulphuric acid, covered with alcohol, and the latter ignited, impart a greenish color to the flame. (Other substances, as copper, which might also im- part a green color to the flame should be removed before making the test.) Hydrofluoric Acid. (Fluorides.) Etches glass (see 4, page 67). Carbonic Acid. (Carbonates^) On adding an acid to a carbonate, effervescence occurs, due to the evolution of carbon dioxide. The presence of carbon 165 dioxide is confirmed by the production of a white turbidity or precipitate, due to the formation of calcium carbonate with clear calcium hydroxide solution (see 2, page 68). Silicic Acid. (Silicates.) A portion fused in a bead of microcosmic salt yields a bead in which the silica is not dissolved, but swims in the fused bead as small opaque particles (see 4, page 70). Arsenious Acid. (Arsenites.) Arsenic Acid. (Arseniates.) Chromic Acid. ( Chromates.) These will have been detected in the preliminary exam- ination and in the examination for bases. Argentic nitrate added to a solution of arsenious acid, fol- lowed by the addition of ammonium hydroxide, drop by drop, or to a solution of an arsenite, produces a yellow, curdy precipitate of argentic arsenite, soluble in ammonium hydroxide and in nitric acid. Argentic nitrate added to a solution of arsenic acid, fol- lowed by the addition of ammonium hydroxide, drop by drop, or to a solution of an arseniate, produces a reddish- brown precipitate of argentic arseniate, soluble in ammonium hydroxide and in nitric acid. Arsenious acid may be detected in the presence of arsenic acid by its being immediately precipitated as arsenious sul- phide by hydrogen sulphide, whereas arsenic acid is pre- cipitated only after continuing the introduction of hydrogen sulphide for some time. Arsenic acid is detected in the presence of arsenious acid by the formation of a white, crystalline precipitate of am- monium magnesium arseniate on the addition of magnesia mixture (see 7, page 28). Arsenious acid does not produce a precipitate with magnesia mixture. 166 Chromic acid produces a yellow precipitate with plumbic acetate and a purplish-red precipitate with argentic nitrate (see 3, page 71, and 4, page 72). Oxalic Add. (Oxalates.) Tartaric Acid. (Tartrates.) Produce white precipitates with calcium chloride. They may be distinguished when together by the behavior of their calcium salts : calcium oxalate is insoluble and calcium tar- trate soluble in acetic acid. Hydrochloric Acid. (Chlorides.) Hydrobromic Acid. (Bromides.) Hydriodic Acid. (Iodides.) Hydrocyanic Acid. (Cyanides.) All are precipitated by argentic nitrate respectively as chloride, bromide, iodide, and cyanide of silver, and are dis- tinguished by the behavior of their silver salt with ammonium hydroxide. Argentic chloride and argentic cyanide are easily soluble in dilute ammonium hydroxide. Argentic bromide is soluble in concentrated ammonium hydroxide ; argentic iodide is insoluble in ammonium hydroxide. If the precipitate produced on the addition of argentic nitrate is insoluble in nitric acid but soluble in ammonium hydroxide, hydriodic acid is absent, but hydrochloric acid, hydrocyanic acid, and hydrobromic acid may be present. A test for hydrocyanic acid may be made by means of the Prussian-blue reaction (see 4, page 78), and for hydrobromic acid with chlorine- water and chloroform or carbon disulphide (see 5, page 75). If neither hydrobromic acid nor hydrocyanic acid is present, the solubility of the silver precipitate in ammonium hydrox- ide proves the presence of hydrochloric acid. If hydrobromic acid or hydrocyanic acid is present, the distillation test with 167 potassium bichromate and sulphuric acid must be made for hydrochloric acid (see 4, page 74). If the precipitate is insoluble or only partly soluble in ammonium hydroxide, the presence or absence of hydriodic acid must be established by means of chlorine-water and chloroform or carbon disulphide (see 5, page 76). If a violet color is produced, chlorine-water is added drop by drop until either decolor ization occurs (absence of hydrobromic acid) or the yellow color, due to the presence of bromine (which was obscured by the violet color produced by iodine), appears (see 5, page 75). The test for hydrocyanic acid should be made as before described (page 125). The distillation test for the detection of hydrochloric acid is to be employed when, in addition to hydriodic acid, hydrobromic acid or hydrocyanic acid is present. If hydrobromic acid or hydrocyanic acid -is ab- sent, hydrochloric acid may be detected in the presence of hydriodic acid by the solubility of the silver precipitate in ammonium hydroxide. It should be remembered that chloride, bromide, iodide, and cyanide of silver are soluble in hyposulphites of the alkalies ; therefore in the presence of hyposulphites the hypo- sulphurous acid should be removed by gently warming with nitric acid. As argentic nitrate fails to produce a precipitate in solu- tions of mercuric cyanide, the presence of mercuric cyanide must be proved according to 4, page 111. Hydroferrocyanic Acid. (Ferrocyanides.} Ferric chloride produces a dark-blue precipitate of ferric ferrocyanide (Prussian blue), insoluble in acids (see 5, page 80). Cupric sulphate produces a precipitate of brownish-red cupric ferrocyanide. 168 Hydroferricyanic Acid. (Ferricyanides.) Ferrous sulphate precipitates blue ferrous ferricyanide (Turnbull's blue), insoluble in acids. Ferric chloride fails to produce a precipitate, but produces a dark coloration in the liquid, due probably to the produc- tion of ferric ferricyanide (see 3, page 80). Cupric sulphate precipitates greenish-yellow cupric ferri- cyanide. Hydriodic add (iodides), hydrobromic acid (bromides), and hydrochloric acid (chlorides) are detected in the presence of hydroferrocyanic acid and hydroferricyanic acid by means of chlorine-water and chloroform, and by distillation with potas- sium bichromate and sulphuric acid. To detect hydrocyanic acid in the presence of hydroferro- cyanic acid and hydroferricyanic acid, the solution is acidu- lated with hydrochloric acid, and calcium carbonate is imme- diately added until carbon dioxide ceases to be evolved. A test for hydrocyanic acid is then made by the ammonium sulphocyanide reaction (see 5, page 78). The hydrochloric acid liberates hydrocyanic acid as well as hydroferrocyanic acid and hydroferricyanic acid, but only the two latter pos- sess the property of decomposing carbonates to form salts; hydrocyanic acid, therefore, remains in the free state. Hydrogen Sulphide. (Sulphides.) The presence of sulphides is detected when making the preliminary examination. A soluble sulphide placed on a clean silver coin and moistened with a few drops of water produces a brownish or black stain on the coin. Nitric acid or nitro-hydrochloric acid decomposes sulphides, with the separation of sulphur ; hydrochloric acid causes the 169 evolution of hydrogen sulphide if it should have any action at all. Plumbic acetate produces in solutions of sulphides a black precipitate of plumbic sulphide. Sodium nitro-prusside solutions are colored violet by sul- phides, but not by free hydrogen sulphide. Nitrous Acid. (Nitrites.) On the addition of an acid to a nitrite, brownish-red fumes of nitrogen dioxide are evolved. On adding a few drops of sulphuric acid to a solution of a nitrite, cooling the liquid, and adding ferrous sulphate, a brown or black coloration is produced (see 5, page 82). Potassium iodide (or cadmium iodide), starch paste, and a few drops of dilute sulphuric acid added to a solution of a nitrite produce a blue coloration (see 6, page 83). Before testing for nitric acid in the presence of nitrous acid, the nitrous acid must be decomposed by being boiled a suffi- cient length of time with a solution of ammonium chloride : KNO 2 + NH 4 C1 = KC1 + N 2 + 2H 2 O. Hypochlorous Add. (Hypochlorites.) Hydrochloric acid decomposes hypochlorites, with the evo- lution of chlorine. Plumbic acetate produces in solutions of hypochlorites at first a white precipitate of plumbic chloride, which soon becomes yellow and finally brown, due to the formation of lead dioxide (see 3, page 83). Nitric Acid. (Nitrates.} On adding a small quantity of concentrated sulphuric acid to a solution of a nitrate, cooling, and placing a crystal of ferrous sulphate in the liquid, a brownish or black ring is formed around the crystal (see 3, page 84). H 15 170 With potassium iodide (or cadmium iodide), starch paste, and dilute sulphuric acid nitrates do not produce a blue discoloration unless a fragment of metallic zinc is added. (Distinction from nitrites. See 4, page 84.) Before testing for nitric acid in the presence of hydriodic acid or hydrobromic acid, the two latter acids must be re- moved by precipitation with argentic sulphate or with plumbic acetate, the precipitate filtered off, and the tests for nitric acid made in the filtrate. Chloric Acid. (Chlorates.) On warming a solution of a chlorate with hydrochloric acid the liquid becomes greenish yellow in color, and greenish- yellow fumes of a mixture of chlorine and chlorine tetroxide are evolved. Concentrated sulphuric acid poured upon a very small portion of a solid chlorate causes the evolution of chlorine tetroxide (see 4, page 86). Chlorates in the solid state on being strongly heated are converted into chlorides. On dissolving the residue in water and testing for a chloride with argentic nitrate, a white pre- cipitate of argentic chloride will be produced. (A chlorate itself, free from chlorides, does not yield a precipitate with argentic nitrate.) Acetic Acid. (Acetates.') On adding ferric chloride to a solution of an acetate and boiling the liquid, a brownish-red precipitate of basic ferric acetate is formed. Sulphuric acid added to an acetate and the liquid warmed liberates acetic acid, which is recognized by its odor of vinegar. Alcohol added to a cooled solution of an acetate containing free sulphuric acid and then warmed produces acetic ether, which is recognized by its apple-like odor (see 6, page 87). APPENDIX. BEHAVIOE OF THE COMPOUNDS OF THE RAKE ELEMENTS. THE deportment of the rare elements and their compounds when subjected to the usual preliminary examination, as well as the behavior of these elements with the ordinary group reagents, will be treated of in the following pages. It is not intended to give a detailed description of the separation of the rare elements from one another or from the more fre- quently occurring elements, but in the latter part of the appendix a few examples are given of the separation of the rare elements in minerals which may be easily procured. I. BEHAVIOR IN THE PRELIMINARY EXAMINA- TION. (a) On heating the substance in a glass reduction-tube : Titanic acid becomes yellow to brown. Niobic acid becomes yellow. Tantalic o,cid becomes pale yellow. Selenium and selenides yield a reddish-brown sublimate : a portion heated in a tube open at both ends and held obliquely in the flame gives a radish-like odor. 171 172 Tellurium sublimes ; heated in a tube open at both ends, it burns, producing dense white clouds. (6) On heating the substance with the blowpipe flame on charcoal, there are produced : Fused metallic globules : Gold : yellow, ductile, without incrustation. Thallium : white, ductile, yellow incrustation. Indium : white, ductile, white incrustation. Incrustation, without metallic globule : Tellurium. Infusible metallic masses : Tungsten, Molybdenum, Platinum, Palladium, etc. White masses (When heated with cobaltous nitrate solution) : Titanic acid becomes yellowish green. Niobic add becomes dirty green. Tantalie acid becomes flesh color. Beryllia becomes gray. Brownish-red masses : Selenium compounds. Tellurium compounds. When placed on a silver coin and moistened with water, a brown or black stain is produced on the coin. When treated with hydrochloric acid, hydrogen sele- nide and hydrogen telluride are evolved. (c) On fusing the substance in a bead of microcosmic salt the following colored beads result in the 173 Oxidizing Flame. Uranium : yellow when hot, yellowish green when cold. Cerium: reddish yellow when hot, lighter reddish yellow to colorless when cold. Didymium : colorless. Titanium : colorless. Niobium: colorless. Tungsten: colorless. Molybdenum : colorless. Vanadium : colorless. Reducing Flame. Green. Colorless. Amethyst changing to violet. Violet, Blue or violet. Blue. Black. Green. Gold and platinum are not dissolved in the bead of micro- cosmic salt. (d) Examination in the flame. The non-luminous flame is colored by Lithium : carmine-red. Rubidium: violet. Caesium: violet. Indium : bluish violet. Selenium : ultramarine-blue. Tellurium : blue bordered with green. Thallium: intense green. Molybdic acid : yellowish green. Lithium, rubidium, ccesium, indium, thallium, and gallium are best detected by means of the spectroscope. Erbium and didymium also furnish absorption spectra. 15* 174 II. BEHAVIOR WITH THE GROUP REAGENTS. FIRST GROUP. Hydrochloric acid precipitates : Thallium : as white curdy T1C1, thallous chloride, solu- ble with difficulty in water. From alkaline solutions : Molybdic acid : as white H 2 MoO 4 , molybdic acid, soluble in an excess of hydrochloric acid. Tungstic acid : as white H 2 WO 4 , tungstic acid, insoluble in an excess of hydrochloric acid ; becomes yellow on boiling. Tantalic acid : as white HTaO 3 , tantalic acid, soluble in an excess of hydrochloric acid, producing opalescence in the liquid. SECOND GROUP. Hydrogen sulphide precipitates : Palladium : as black PdS, palladious sulphide. 1 | Osmium : as brownish-black OsS, osmic sulphide. g & Rhodium : as brown Rh 2 S 3 , rhodic sulphide. |- ~ Ruthenium : as brown Ru 2 S 3 , ruthenic sulphide. J f 1 (The liquid at first becomes azure-blue in color.) Gold : as black Au 2 S 3 , auric sulphide. ] Platinum: as brownish-black PtS 2 , platinic sulphide. | Indium : as brown Ir 2 S 3 , iridic sulphide. Molybdenum : as brown MoS 3 , molybdic sulphide. (A small quantity of hydrogen sulphide colors the j- s solution blue.) Selenium : as yellow, which on warming changes to ,| reddish-yellow SeS 2 , selenic sulphide. Tellurium : as brown TeS 2 , telluric sulphide. The solution may become blue in color if compounds of tungsten or vanadium are present. 175 THIRD GROUP. Ammonium hydroxide in the presence of ammonium chloride precipitates : Uranium: as yellow (NH 4 ) 2 Ur 2 O 7 (?), ammonium uranate. Indium : as white In(OH) 3 , indie hydroxide, soluble in NaOH. Beryllium: as white Be(OH) 2 , beryllic hydroxide, soluble in NaOH. Zirconium: as white Zr(OH) 4 , zirconic hydroxide, insoluble in NaOH. Thorium : as white Th(OH) 4 , thoric hydroxide, in- soluble in NaOH. Yttrium: as white Y(OH) 2 , yttric hydroxide, in- soluble in NaOH. Cerium : ~\ Lanthanum : > as basic salts. Didymium : . Titanium : as white Ti(OH) 4 , titanic hydroxide. Tantalum : as white TaO 2 (OH), acid tantalic hydroxide, or as an acid ammonium salt. Niobium : as white NbO 2 (OH), acid niobic hydroxide, or as an acid ammonium salt. FOURTH GROUP. Ammonium sulphide precipitates : Thallium : as black T1S, thallous sulphide. If the nitrate from the Fourth Group precipitate is treated with hydrochloric acid, there will be precipitated : Tungsten : as brown WS 3 , tungstic sulphide. 176 Vanadium: as brown vanadium sulphide, containing oxygen and varying in composition. Molybdenum: (if present) as brown MoS 3 , molybdic sulphide. FIFTH GROUP. In this group may be found : Lithium. Ccesium. Rubidium. (To be detected by means of the spectroscope.) III. EXAMPLES FOR PRACTICE, WOLFRAMITE. Wolframite may be recognized by the blue color it imparts to the bead of microcosmic salt in the reducing flame, and by the yellow residue of tungstic acid remaining when the finely- pulverized mineral is boiled with hydrochloric acid. The finely-pulverized wolframite is boiled with concentrated hydrochloric acid, and a few drops of concentrated nitric acid are added from time to time, until the undissolved resi- due is pure yellow in color and does not undergo further change. Tungstic acid remains undissolved as a yellow powder, while the bases enter into solution as chlorides. The liquid containing the insoluble residue is evaporated to dry- ness on a water-bath, the residue extracted with water con- taining a small quantity of hydrochloric acid, filtered, and the filtrate tested for bases. The insoluble residue contains tungstic acid, and frequently silicic acid and niobic acid. The residue is treated with am- monium hydroxide, which dissolves the tungstic acid as an 177 ammonium salt, leaving an undissolved residue consisting of silicic acid and possibly niobic acid. This residue is thor- oughly washed with ammonium hydroxide, to render it free from tungstic acid, and then tested in a bead of microcosmic salt for niobic acid. The ammoniacal solution containing ammonium tungstate should give the following reactions : Hydrochloric acid precipitates white H 2 WO 4 , tungstic acid, which, on boiling, becomes yellow. Metallic zinc and an excess of hydrochloric acid impart to the precipitate of tungstic acid a blue color changing to brown (due to the formation of lower oxides of tungsten). Stannous chloride produces a yellow precipitate ; on adding hydrochloric acid and warming, the yellow color changes to blue. Ammonium sulphide added to the solution of ammonium tungstate produces no precipitate, but forms the soluble sulpho-salt (NH 4 ) 2 WS 4 . On adding hydrochloric acid to this solution, brown WS 2 , tungstic sulphide, is precipitated. The supernatant liquid is generally blue in color. MOLYBDENITE. Molybdenite, when fused in a bead of microcosmic salt, yields a colorless bead in the oxidizing flame and a black bead in the reducing flame. It imparts a yellowish-green color to the non-luminous flame. When heated on charcoal, it yields a reddish-brown mass. It is soluble in nitro-hydro- chloric acid, imparting a green color to the liquid. On evaporating the excess of acid, diluting with water, and con- ducting hydrogen sulphide into the solution, a blue coloration is produced, and gradually brownish-black MoS 3 , molybdic sulphide, is precipitated. Molybdic sulphide is soluble in 178 ammonium sulphide, which dissolves it as a sulpho-salt, (NH 4 ) 2 MoS 4 , from which solution it is reprecipitated by hydrochloric acid as MoS 3 , molybdic sulphide. The nitrate from the precipitate produced by the intro- duction of hydrogen sulphide may still contain molybdenum in solution ; therefore, before testing for the metals of the Fifth Group, ammonium hydroxide is added to the solution, which is gently warmed, filtered if a precipitate be formed, and the solution, which now contains (NH 4 ) 2 MoS 4 , is treated with hydrochloric acid, which precipitates the remaining molybdenum as molybdic sulphide. The molybdic sulphide reprecipitated in the Second Group from the ammonium sulphide solution by hydrochloric acid is collected on a filter, dried, and placed in an uncovered crucible, which is placed obliquely over the flame and gently heated, whereby the molybdic sulphide is oxidized and con- verted into molybdic acid, with the evolution of sulphurous anhydride. When the sulphurous anhydride ceases to be evolved, the yellow residue is dissolved in ammonium hy- droxide and the resulting solution of ammonium molybdate tested as follows : a small portion of the solution is tested for copper with a few drops of ammonium sulphide, and the remainder of the solution is used in making the tests for molybdenum. (The precipitated molybdic sulphide obtained from the Fourth Group is heated in an uncovered crucible and treated in the same manner.) Hydrochloric acid causes the precipitation of white ILjMoO^ molybdic acid, soluble in an excess of hydrochloric acid. Stannous chloride produces in the hydrochloric acid solution of molybdic acid a blue coloration, changing to green and finally to brown metallic zinc produces a similar coloration, but not so promptly. The change in color in the two pre- ceding reactions is due to the reduction of molybdic acid. 179 Ammonium sulphocyanide added to the ammoniacal solu- tion, followed by the addition of hydrochloric acid and zinc, produces a carmine-red coloration (in consequence of reduction with the formation of sulphocyanides of the oxides). Concentrated nitric acid and sodium hydrogen phosphate added to the ammoniacal solution precipitate yellow am- monium phosphomolybdate. WULFENITE. Wulfenite, when heated in the blowpipe flame on charcoal, yields a globule of metallic lead ; when fused in a bead of microcosmic salt, it yields a colorless bead in the oxidizing flame and a black bead in the reducing flame. Wulfenite is soluble in hydrochloric acid with the separa- tion, upon cooling, of crystalline plumbic chloride. The hydrochloric acid solution yields with hydrogen sulphide in the Second Group a precipitate of molybdic sulphide, soluble in ammonium sulphide, thus furnishing a means of separating it from plumbic sulphide, which is insoluble in ammonium sulphide. After precipitating the Fourth Group, the reddish- brown nitrate is treated with hydrochloric acid to precipitate the remainder of molybdic sulphide. The molybdic sulphide is further examined as given under Molybdenite, page 134. URANINITE (PITCHBLENDE). Uraninite, when fused in the bead of microcosmic salt, yields a yellow bead in the oxidizing flame and a green bead in the reducing flame ; treated with nitric acid it dissolves, leaving a residue of silicic acid and insoluble oxides (see page 110). The nitric acid solution yields in the Third Group a precipitate containing uranium, the uranium being precipi- tated as yellow ammonium uranate. In order to separate 180 uranium the precipitate of the Third Group is digested at a moderate heat with a concentrated solution of ammonium carbonate ; uranium enters into solution as uranyl ammonium carbonate, UrO 2 CO 3 ((NH 4 ) 2 CO 3 ) 2 , imparting a yellow color to the solution. The oxides of the other metals remain undissolved, and after being collected on a filter may be examined according to the usual scheme of analysis. To detect uranium a portion of the yellow filtrate contain- ing uranyl ammonium carbonate is acidulated with acetic acid and treated with potassium ferrocyanide : a reddish- brown precipitate of (UrO 2 ) 2 Fe(CN) 6 , uranyl ferrocyanide, indicates the presence of uranium. The remainder of the solution of uranyl ammonium carbonate is carefully concen- trated on a water-bath, and on cooling glistening yellow crystals of uranyl ammonium carbonate separate, which when strongly ignited leave a residue of dark-green uranous-uranic oxide, Ur 3 O 8 . On treating the filtrate containing ammonium sulphide from the precipitate of the Fourth Group with hydrochloric acid, nickelous sulphide together with vanadic sulphide may be precipitated. If a precipitate is obtained by treatment with hydrochloric acid, it is collected on a filter, washed, dried, mixed with potassium nitrate, fused, and the resulting fused mass extracted with water, whereby vanadium as an acid vanadium salt enters into solution. On filtering, neutralizing the filtrate with nitric acid, and adding a concentrated solution of ammonium chloride, vanadic acid in combination with ammonium is gradually precipitated as a white ammonium salt. On collecting the precipitate on a filter, dissolving in water, and adding a small quantity of hydrochloric acid, (1) 1 The solution becomes yellow or red in color. 181 followed by the addition of metallic zinc, the solution becomes blue in color. RUTILE (TITANIFEROUS IRON). Rutile, when fused in a bead of microcosmic salt in the reducing flame, yields a violet- to blood-red-colored bead. Fused in the oxidizing flame (when a sufficient quantity of the mineral is employed) it yields microscopic tabular crystals of anatase (TiO 2 ). The mineral is best decomposed by being fused with acid potassium sulphate at not too high a temperature. The fused mass after cooling is pulverized and dissolved in cold water : the titanium enters into solution as sulphate. The liquid is filtered and a portion of the filtrate tested with metallic zinc or tin for titanic acid : in the presence of titanic acid a pale violet or a blue coloration is imparted to the solution and afterwards a blue precipitate separates which gradually changes to white. On boiling the remainder of the precipitate for some time, meta-titanic acid separates as a white powder. The meta- titanic acid is filtered off and the filtrate diluted with water and again boiled and filtered. The filtrate is then tested for the remaining bases in the usual manner. BERYL. Beryl, when fused in a bead of microcosmic salt, slowly dissolves without the formation of a skeleton of silica. The fragment of beryl remains in the bead of microcosmic salt and gradually diminishes in size, forming, on cooling, an opalescent bead. Varieties of beryl containing chromium (for example, emeralds) impart, a green color to the bead. As beryl is insoluble in acids, it must be decomposed by fusing with sodium potassium carbonate. The fused mass is 16 182 treated with hydrochloric acid and evaporated to dryness on a water-bath, in order to separate silicic acid. The residue is extracted with water containing a small quantity of hydro- chloric acid and filtered ; the filtrate contains BeCl 2 , beryllium chloride. On the addition of ammonium hydroxide to the filtrate, Be(OH) 2 , beryllium hydroxide, separates as a white precipitate, and is collected on a filter and dissolved in an excess of sodium hydroxide. The sodium hydroxide solution contains sodium aluminate and sodium beryllate, and may also contain sodium chromite, in which case the solution would be green in color. If chromium is not present, the solution is treated with a considerable quantity of ammonium chloride, which precipi- tates aluminium hydroxide and beryllium hydroxide. On boiling the liquid for some time, beryllium hydroxide enters into solution as beryllium chloride, with the evolution of ammoniacal gas, while aluminium hydroxide remains un- dissolved, is collected on a filter, and, when heated with cobaltous nitrate on charcoal, yields a blue mass. The ni- trate, which contains beryllium chloride, is treated with ammonium hydroxide to precipitate beryllium hydroxide. Beryllium hydroxide is soluble in an excess of ammonium carbonate, and from this solution it separates on boiling as a basic beryllium carbonate. The beryllium hydroxide, when heated with cobaltous nitrate in the oxidizing flame on char- coal, yields a gray mass. If chromium is present, the solution is diluted with water and boiled for some time : aluminium hydroxide remains in solution, while beryllium hydroxide and chromium hydroxide are precipitated. The precipitate is collected on a filter, dried, and the beryllium hydroxide separated from the chro- mium hydroxide by fusing with a mixture of sodium carbon- ate and potassium nitrate. On extracting the yellow fused 183 mass with water, beryllium oxide remains undissolved, while chromium enters into solution as a chromate of the alkali. CERITE. Cerite, heated in the blowpipe flame on charcoal, is infusi- ble, but becomes dirty yellow in color. Fused in a bead of microcosmic salt, it yields in the oxidizing flame a reddish- yellow bead, and in the reducing flame a colorless bead, together with a skeleton of silica. On heating a portion of the finely-pulverized mineral with hydrochloric acid, diluting with water, filtering, and adding oxalic acid to the filtrate, a white precipitate is produced. On heating the mineral with concentrated hydrochloric acid, evaporating to dryness on a water-bath, and extracting the residue with water and a few drops of hydrochloric acid, the bases enter into solution while silicic acid remains undis- solved. In the Second Group, in addition to other metals, molyb- denum may be precipitated as molybdic sulphide. (See Molybdenite, page 134.) Cerium, lanthanum, and didymium are precipitated by ammonium hydroxide in the Third Group as hydroxides : Ce(OH) 2 , La(OH) 2 , and Di(OH) 2 . Cerium hydroxide and lanthanum hydroxide are white. The former when exposed to the air oxidizes and becomes yellow ; didymium hydroxide is pink in color. On collecting the precipitate of the Third Group on a filter, dissolving in hydrochloric acid, and add- ing oxalic acid, cerium, lanthanum, and didymium are pre- cipitated as white oxalates insoluble in dilute acids; the filtrate from the precipitate of oxalates may be examined for the remaining bases of the Third Group. The oxalates of cerium, lanthanum, and didymium when ignited yield a brown residue consisting of a mixture of 184 oxides (Ce 3 O 4 , LaO, and DiO). On heating a portion of this mixture of oxides with concentrated sulphuric acid, sul- phates are formed which are soluble in water and impart a yellow color to the liquid. In this solution concentrated potassium sulphate produces a lemon-yellow precipitate con- sisting of a mixture of double salts. The remainder of the mixture of oxides is treated with hydrochloric acid and alcohol and then heated : the oxides are dissolved thereby, with the formation of chlorides (CeCl 2 , LaCl 2 , DiCl 2 ). If didymium is present, on passing a ray of light through the solution, the spectrum shows dark absorp- tion bands (four between Frauenhofer's lines D and F and two between F and G). On adding sodium acetate and passing chlorine through the solution, or on the addition of sodium hypochlorite, light- yellow Ce 3 O 7 H 6 , cerium hydroxide, is precipitated, which is soluble when warmed with hydrochloric acid, forming CeCl 2 , with the evolution of chlorine. (For the separation of cerium, lanthanum, and didymium the reader is referred to more extensive works on qualitative analysis.) ZIRCON (HYACINTH). Zircon, heated in the blowpipe flame on charcoal, is infusi- ble, but becomes lighter in color. It is not dissolved in the bead of microcosmic salt. It is decomposed by being fused for some time with sodium potassium carbonate. The fused mass is treated with hydro- chloric acid, evaporated to dryness on a water-bath (to render the silica insoluble), the residue extracted with water and hydrochloric acid and filtered. The filtrate contains zirco- nium as zirconium chloride. From this solution the zirco- nium is precipitated by ammonium hydroxide in the Third Group as Zr(OH) 4 , zirconium hydroxide. Zirconium hy- 185 droxide is insoluble in sodium hydroxide, but soluble in ammonium carbonate. From the solution in ammonium carbonate it is reprecipitated by boiling. On dissolving a portion of this precipitate in sulphuric acid and adding a concentrated solution of potassium sulphate, a white double salt of zirconium is precipitated. The hydrochloric acid solution is not precipitated by oxalic acid, but the neutral solution is precipitated by ammonium oxalate ; the precipitate redissolves in an excess of ammonium oxalate. Turmeric paper, moistened with the hydrochloric acid solution, becomes reddish brown in color when dry. LEPIDOLITE. Lepidolite when fused in a bead of microcosmic salt yields a skeleton of silica. When placed on a platinum wire and held in the non-luminous flame it imparts a carmine-red color to the flame, especially after moistening the lepidolite with a few drops of hydrochloric acid. Treated with concentrated sulphuric acid it responds to the fluorine reactions (see 4, page 67, and 5, page 68). A portion of the mineral is heated in a platinum crucible (without the addition of carbonates of the alkalies as a flux) until melted. The fused mass is pulverized and then decom- posed by being boiled with concentrated hydrochloric acid. The solution is evaporated to dryness on a water-bath to render the silica insoluble; the residue is extracted with water and a few drops of hydrochloric acid, the insoluble silica filtered oif, and the filtrate examined for bases. Lithium belongs to the Sixth Group. The hydrochloric acid filtrate, obtained as described above, is treated with am- monium carbonate to precipitate the metals of the Fifth Group, and with sodium hydrogen phosphate to precipitate magnesium, filtered, the filtrate treated with barium chloride, 16* 186 and the liquid again filtered. The filtrate now contains, in addition to the excess of barium chloride, the alkaline metals as chlorides. The liquid is evaporated to dryness, the residue gently heated over a free flame to expel ammonium chloride, and then placed in a small flask and extracted with a mixture of alcohol and ether. Lithium chloride enters into solution, while the chlorides of the other metals remain undissolved. The alcoholic solution is filtered and the filtrate evaporated to dryness on a water-bath ; the residue remaining consists of lithium chloride, which is recognized by its imparting a carmine-red color to the non-luminous flame and by the examination with the spectroscope. The residue, insoluble in the mixture of alcohol and ether, is dissolved in water, and sulphuric acid added to the solution to precipitate barium as barium sulphate, which is filtered off and the filtrate ex- amined for potassium and sodium. INDEX. PAGE Acetates 86, 1 70 Acetic acid 86, 170 Acid, acetic 86, 170 arsenic 26 arsenious 21 boric 65, 164 carbonic 68, 164 chloric 85, 170 chromic 71, 165 ferricyanic 80, 168 ferrocyanic 79, 168 hydriodic 75, 166 hydrobromic 74, 166 hydrochloric 73, 166 hydrocyanic 77, 166 hydroferricyanic 80, 168 hydroferrocyanic IT- .: . 79, 167 hydrofluoric 67 hydrofluosilicic 60, 162 hypochlorous 83, 169 hyposulphurous 63, 163 molybdic 173 niobic . 172 nitric ^ , 84, 169 nitrous 82, 169 oxalic ' 87 phosphoric 64, 164 silicic 69, 165 sulphuric 59, 162 sulphurous 61, 163 sulphydric 81, 168 tantalic 172 187 188 PAGE Acid, tartaric gg thiosulphuric 53 titanic .172 Acid sulphides, separation of 135 Acids, behavior with group reagents 160 examination for 158 preliminary tests for 99 Aluminium 41 Ammonium 55 carbonate 50 Antimony 28 Marsh's test for 30 Reinsch's test for 30 in antimonic condition 31 in antimonious condition 28 Arsenic 21 Marsh's test for 22 Reinsch's test for 25 in arsenic condition 26 in arsenious condition 21 Arsenic acid 26 Arsenious acid 21 Barium 50 Bases, detection of, in the Wet Way 116 precipitation of groups ' 117 properties of 9 separation of First Group 129 Second Group . . 130 Third Group 139 Fourth Group 146 Fifth Group /v : 148 Sixth Group 152 Basic sulphides, separation of 135 Beryl 181 Beryllium ' 175 Bismuth . . v 20 Borates 65, 164 Boric acid 65, 164 Bromides . . 74, 166 189 PAGE Cadmium 35 Caesium 176 Calcium 52 Carbonates 68, 164 Carbonic acid 68, 164 Cerite 183 Cerium 175 Charcoal, examination on 93 Chlorates 85, 170 Chloric acid 85, 170 Chlorides 73, 166 Chromates * 71, 165 Chromium 42 Cobalt 47 Copper 17 Cyanides 77, 111, 166 Detection of bases in the Wet Way 116 Didymium 175 Dissolving metals and alloys 110 oxides and salts 104 Everett's salt 39,80 Examination by microcosmic salt 97 for acids . , 158 in the flame 98 in the reduction-tube 90 on charcoal 93 Ferricyanides 80, 168 Ferrocyanides 79, 167 Flame, examination in 98 Fluorides 67 Gold 36 Group (acids) : First 59 Second 61 Third 73 Fourth 84 192 PAGE Keinsch's test 25 Rhodium 174 Rubidium 176 Ruthenium 174 Rutile -... 181 Selenides 171 Selenium 174 Separation of acid sulphides 135 of bases : First Group 129 Second Group 130 Third Group 139 Fourth Group 146 Fifth Group 148 Sixth Group , ~~r~v-\ 152 of basic sulphides 134 Silicates 69, 113, 165 Silicic acid 69, 165 Silicofluorides . 60, 162 Silver 9 Sodium 56 Strontium 51 Sulphates 59, 162 Sulphides 81, 168 of heavy metals Ill Sulphites 61, 163 Sulphuric acid 59, 162 Sulphurous acid 61, 163 Sulphydric acid 81, 168 Table of group precipitations 118 Tantalic acid 172 Tantalum 175 Tartaric acid 88 Tartrates 88 Tellurium 174 Tests, preliminary, in Dry "Way 90 Thallium 174 Thenard's blue 42 Thiosulphites 63 193 PAGE Thiosulphuric acid 63 Thorium 175 Tin 32 in stannic condition 34 in stannous condition 33 Titanic acid 172 Titaniferous iron 181 Titanium 175 Turnbull's blue 39, 80 Uraninite 179 Uranium 175 Vanadium 176 Wolframite 176 Wulfenite 179 Yttrium 175 Zinc 46 Zircon . . 184 17