These relations clearly recall the characteristic behavior of ammoniacal and cyanide solutions, in which complex ions are formed, and the interference of the organic compounds with precipitation is of a similar nature—complex ions are formed by metal ions with these organic compounds, and the complexes are, in many instances, sufficiently stable to reduce the concentrations of the metal ions to the point, where only very difficultly soluble salts can be precipitated.
The relations may be illustrated by the discussion of the complex formed by the cupric-ion with tartrates. The structure of sodium tartrate is expressed by NaO2C─CH(OH)─CH(OH)─CO2Na. The underscored hydroxide groups OH are known in organic chemistry as alcohol groups.487 Now, alcohols resemble water in very many properties and, among others, in the capacity to form metal derivatives or alcoholates, in which the hydrogen (ion) of the hydroxide group is replaced by metal ions. The alcoholates correspond, thus, to the metal hydroxides, which are the analogous [p240] derivatives of water. Exactly as there is a vast difference in the readiness with which the various metal hydroxides, or bases, ionize, many of them being only slightly ionizable (the weakest bases), so certain alcoholates are much less readily ionizable than others. The alkali alcoholates are most readily ionized.
When sodium tartrate is mixed with an excess of sodium hydroxide, some of the readily ionizable sodium salt of the alcohol groups of the sodium tartrate is formed and we have:
When cupric sulphate is added to this mixture, a slightly ionizable complex cupri-tartrate-ion is formed by the union of the cupric-ion with the "alcoholate-tartrate-ion":
The complex ion is not perfectly stable and so the action is a reversible one, as indicated in the equation. The greater portion of the copper, however, is present as part of the complex negative ion of cupric-tartaric acid and its salts. This may be demonstrated by subjecting the solution to electrolysis in a U-tube (p. 45). It is readily seen (exp.) that a deep blue ion, obviously containing copper, migrates to the positive pole.488 The concentration of cupric-ion is so small that its hydroxide and its phosphate are not precipitated from the solution by the addition, respectively, of alkali or of a soluble phosphate (exp.). Cupric sulphide, however, is so insoluble that it may be precipitated completely from the solution by the addition of a sulphide (exp.), the concentration of cupric-ion being much smaller in the saturated solution of the sulphide than in the solution of sodium cupri-tartrate.
Citrates, sugars, glycerine, contain alcoholic groups of the same nature as found in the tartrates, and they are capable of forming similar complexes, or little ionizable alcoholates, with metal ions.
Certain organic acids, which contain no alcohol groups,489 are [p241] also capable of forming fairly stable complexes with metal ions: thus, acetates form a complex with lead-ion, that is sufficiently stable to render lead sulphate, which is difficultly soluble in water, readily soluble in ammonium acetate solution490 (exp.). Soluble oxalates readily combine with ferric, ferrous, cupric and other oxalates and interfere, more or less, with the detection of the metal ions, as the result of complex formations.
All of these complexes are decomposed rather readily by the addition of strong acids, whose hydrogen-ion breaks up the complex ions, by suppressing the anions491 of the much weaker organic acids and alcohols. Consequently, these organic compounds do not interfere with tests which may be carried out in strongly acid solution. For instance, the addition of potassium ferrocyanide to a solution of ammonium ferrioxalate (NH4)3Fe(C2O4)3, to which an excess of ammonium oxalate has been added, gives only a slight indication of the presence of ferric-ion (a greenish blue solution is obtained); when hydrochloric acid, in excess, is added to the mixture, Prussian blue, ferric ferrocyanide, is immediately precipitated in quantity (exp.).
[425] The naming of the complex ions, which ammonia forms with metal ions, has not yet been satisfactorily settled. English writers frequently speak of "ammonio-argentic" ion and "ammonio-argentic" nitrate. German writers speak of "Silber-ammoniak" ion (Abegg, Handb. der anorg. Chem., II, 728), which would read "silver-ammonia" ion in English terms. The terminology "silver-ammonium" ion, used in this book, is based on the idea, that all these complex ions are essentially of the same nature as the well-known ammonium-ion, NH4+, the positive charge being, almost certainly, carried by nitrogen in these complex ions, as it is in ammonium-ion. The latter is a complex ion of ammonia with hydrogen-ion. The name "ammonio-argentic" ion does not bring out this close relationship and puts the emphasis on the silver, which is probably little concerned in the reactions of the complex as such. The names "silver-ammonia" ion and "silver-ammonia" nitrate sound badly and do not emphasize the relation to ammonium, potassium, sodium and similar positive ions and their salts. The term "ammonium" is, for the reasons given, used here in a generic sense for all complex ions of ammonia with simple metal ions (such as H+, Ag+, Cu2+, Zn2+ etc.), and the number of ammonia molecules, entering into the composition of a complex ion, is not indicated in the names. A similar nomenclature has long been in vogue, and has worked well, for the complex ions of metal ions with the cyanide-ion (see below). We speak of ferrocyanide, Fe(CN)64−, argenticyanide, Ag(CN)2− and Ag2(CN)32−, etc., without indicating the number of cyanide groups, CN, in the complex, and we use the same generic ending "cyanide" as is used to designate the simple cyanide ion, e.g. to designate the ion formed from potassium cyanide, KCN ⇄ K+ + CN−.
[426] The hydroxide-ion appears with the same coefficient, 1, on both sides of the equilibrium equation and need not be included in the mathematical statement; it would appear as a factor in both terms of the ratio given and would cancel out.
[427] Bodländer and Fittig, Z. phys. Chem., 39, 602 (1903).
[428] Bonsdorff, Ber. d. chem. Ges., 36, 2324 (1903).
[429] It is also frequently called the dissociation constant of the complex ion, indicating the tendency of the complex ion to dissociate into its components.
[430] Two independent experimental methods were used and gave concordant results—one having as its basis the solubility of silver salts (chloride, bromide), the other the electrolytic potentials of silver against ammoniacal silver solutions (see Chap. XV).
[431] Bull. de la Soc. Chim. de Paris, (3), 13, 386 (1895).
[432] We may consider the salt to be ionized to about the same extent as ammonium or potassium nitrate in 0.05 molar solutions, or, approximately, 87%. If we call x the concentration of silver-ion, formed by the decomposition of the silver-ammonium-ion, then 2 x is the concentration of the free ammonia, and (0.05 × 0.87 − x) is the concentration of the complex ion. Since x is a small number in comparison with 0.0435, we may write, with sufficient accuracy for our purposes,
Then, x = [Ag+] = 0.0009.
[433] Kohlrausch and Holborn, loc. cit., p. 202.
[434] The dilution of the silver-ammonium nitrate (10 c.c. to 11 c.c.) and the decrease in ionization due to the added salt reduce the concentration of silver-ion from 0.0009 to 0.00085. [Ag(NH3)2+] = (0.05 × 10 / 11) × 0.8 = 0.0364 and 4 x3 = 6.8E−8 × 0.0364 (see footnote, p. 220). Then x = [Ag+] = 0.00085.
[435] Thiel (cf. Bodländer and Fittig, loc. cit.). The solubility given in the table at the end of the laboratory manual refers to 18°. The constant for the complex ion was determined at 25°.
[436] The combined concentration of the salts is 0.055 and their degree of ionization may be taken as 87%, the same as the degree of ionization of 0.05 to 0.06 molar KNO3. Then [Cl−] = (0.1 × 1 / 11) × 0.87 = 0.008. [Ag(NH3)2+] = (0.05 × 10 / 11) × 0.87 = 0.04 and x = [Ag+] = 0.00089 (see the method of calculation in the footnote, p. 220).
[437] The strong solution of ammonia is used in order to avoid unnecessary dilution, and in the experiment, described below, the dilution of the liquids by the added ammonia is considered negligible.
[438] The following solubilities have been determined at 25°:
| [Ag+] × [Cl−] = 2E−10; | [Ag+] = 1.4E−5. |
| [Ag+] × [I−] = 1E−15; | [Ag+] = 1E−8. |
| [Ag+]2 × [S2−] = 4E−50; | [Ag+] = 4.3E−17. |
[439] 100 c.c. molar ammonia dissolves at 25° only 0.6 milligram of silver iodide (Bodländer, loc. cit., p. 606).
[440] Fresenius, Qualitative Analysis, p. 378.
[441] In regard to Cu(NH3)42+ see Locke and Forssall, Am. Chem. J., 31, 268, 297 (1904), and Dawson, J. Chem. Soc. (London), 89, 1674 (1906).
[442] Euler, Ber. d. chem. Ges., 36, 3403 (1903).
[445] The acid HAg(CN)2, corresponding to the salt, is crystallizable and is a strong acid. It is largely decomposed, by water, into silver cyanide and hydrocyanic acid.
[447] In solutions containing an excess of potassium cyanide greater than 0.05 molar, the salt K2[Ag(CN)3] is formed. The dissociation or instability constant for the complex ion Ag(CN)32− is 1E−22.
[448] Z. anorg. Chem., 39, 222 (1904).
[449] The solubility-product constant for silver chloride at 25° is 2E−10. If the concentration of chloride-ion be made 1.0 by the addition of potassium chloride to a 0.05 molar solution of KAg(CN)2, then the concentration of silver-ion, necessary for the precipitation of the chloride, would be KS.P. / [Cl−] = 2E−10 gram-ion. Neglecting the fact that the complex salt is not completely ionized and putting [Ag(CN)2−] = 0.05, and calling x the concentration of the cyanide-ion just necessary to prevent the precipitation of the chloride, we have:
We find x = 5E−7 mole, or approximately 0.03 milligram potassium cyanide (cyanide-ion) per liter. This minute quantity of free cyanide, if not originally present in the solution used, would be formed by the liberation of potassium cyanide from the complex (according to KAg(CN)2 + KCl → AgCl + 2 KCN) as soon as 2.5E−7 mole, or 0.036 milligram, of silver chloride per liter have been formed, a quantity too small to be perceptible.
When potassium cyanide is added to a silver nitrate solution, the precipitate formed is found to be silver argenticyanide, Ag[Ag(CN)2], the silver salt of the extremely stable complex, rather than the simple salt, silver cyanide, AgCN [cf. Bodländer, Z. anorg. Chem., 39, 223 (1904)]. Ag[Ag(CN)2] is even less soluble than silver chloride, the solubility-product constant for [Ag+] × [Ag(CN)2−] being 2.25E−12. An excess of only 2E−6 mole, or about 0.15 milligram, of potassium cyanide (cyanide-ion) per liter is sufficient to prevent the precipitation of silver cyanide (silver argenticyanide) from a 0.1 molar solution of KAg(CN)2, and, conversely, at least this minute excess of potassium cyanide is used in the preparation of a clear 0.1 molar solution of KAg(CN)2, by the addition of potassium cyanide to silver nitrate, until the silver cyanide, first precipitated, is just redissolved (Bodländer, loc. cit.). This excess, as just explained, is more than sufficient to prevent the precipitation of silver chloride from the cyanide solution, even by a large excess of potassium or sodium chloride. Unless one takes into account, in the manner indicated, this marked influence of a minute excess of cyanide-ion in decidedly reducing the concentration of silver-ion in these solutions, one could be led, wrongly, to infer from the value of the instability constant of the complex ion and that of the solubility-product constant of silver chloride, that silver chloride should still be precipitated by the addition of sodium chloride to a solution of KAg(CN)2.
[450] For this reason potassium cyanide is an excellent cleansing agent for stained silverware (sulphide stains), and, since it is an intense poison, cleaning powders should be examined for it.
[451] For the quantitative relations see Lucas, Z. anorg. Chem., 41, 192 (1904).
[452] See Bodlaender, Z. phys. Chem., 39, 597 (1902); Ber. d. chem. Ges., 36, 3933 (1903).
[453] Potassium cyanide is a powerful reducing agent (see p. 89) and is readily oxidized to potassium cyanate. The action, presumably, takes the following course (see Chapters XIV and XV):
[454] Put [Cu+] = x, and [CN−] = 3 x, and, neglecting the degree of ionization, [Cu(CN)32−] = 0.1, x being so small that it need not be subtracted from 0.1. Then x × (3 x)3 = 0.1 × 0.5E−27, and x = 3.7E−8.
[455] Treadwell and Girsewald, Z. anorg. Chem., 38, 92 (1904).
[456] Euler, Ber. d. chem. Ges., 36, 3404 (1903).
[457] Putting [Cd2+] = y, we have y × (4 y)4 = 0.1E−17, and y = 8E−5. In view of the values of the constants, a small excess of potassium cyanide will have a much smaller suppressing effect on the cadmium-ion than on the cuprous-ion. For the excess [CN−] = 0.01, [Cu+] = 5E−23, [Cd2+] = 10−10 as compared with [Cu+] = 4E−8 in a 0.1 molar solution of the salt K3[Cu(CN)3], and with [Cd2+] = 8E−5 in a 0.1 molar solution of K2Cd(CN)4.
[458] Bromine water is a convenient agent for oxidizing cobaltous to cobaltic ions (see Chapter XV).
[459] The heavy arrows → indicate the main course the reversible actions take, under the influence of the reagents used. Since the oxidation of nickel-ion by bromine is accomplished only after the bromine has oxidized any excess of cyanide used—potassium cyanide is a powerful reducing agent (p. 89)—the addition of cyanide, beyond a very small excess, must be avoided (see laboratory instructions).
[460] E.g. for the precipitation of silver, copper, nickel, cobalt and certain other metals from cyanide solutions; cf. Edgar F. Smith, Electro-Analysis (1907).
[461] Z. phys. Chem., 43, 705 (1903). Vide also Haber, Z. Elektrochem., 11, 847 (1905).
[462] 2 Fe2+ + Hg2+ → 2 Fe3+ + Hg ↓. If the treatment with mercuric oxide is carried to completion the final products of the reaction are ferric hydroxide, mercuric cyanide, mercury and potassium hydroxide (Rose, Z. anal. Chem., 1, 300 (1862)):
[463] Bodlaender, loc. cit.
[464] The solubility-product constant of silver sulphide at 25° is 0.5E−51; for [S2−] = 0.8E−5 (p. 202), we would have in the present case [Ag+]2 × [S2−] = 2E−46, which is greater than the constant. Vide quantitative data by Lucas, loc. cit.
[465] Z. f. Elektrochem., 10, 433 and 773 (1904).
[467] Ostwald, Allgem. Chemie, Vol. II, part 1, p. 881 (1893).
[468] Haber, loc. cit.
[469] Then TDecomposition = 10−4 × 1022 = 1018 seconds, and, since there are 3.15E7 seconds in a year, TDecomposition = 3E10 years.
[471] Since there are still smaller "instability constants" than that of the argenticyanide-ion (e.g. for the gold-cyanide-ion the constant is 1 / 1028), there is a large margin of safety for the plausibility of Haber's argument. For full details, his articles (loc. cit.), and the discussion (by Abegg, Bodlaender, Danneel, ibid.) aroused by them should be consulted.
[472] See Le Blanc and Schick, Z. phys. Chem., 46, 213 (1903), on measurements of the speed of ionic actions. The values obtained agree, in general, with Haber's contention.
[473] The concentrations of silver-ion are large, in comparison with those in cyanide solution, and the action is, most likely, essentially an ionic one; but the argument applies with equal force to cyanide systems.
[474] Loc. cit.
[475] An equilibrium constant, as we have seen, is a ratio of velocity constants of balanced reactions (pp. 94, 233) and involves therefore at least two unknown velocity constants. By determining the actual rate of change with known concentrations of reacting components, i.e. by determining the velocity constants themselves, rather than their ratio, a definite conclusion as to the mechanism or path of a given reaction can often be reached (see p. 80).
[476] Proceedings Amer. Academy, 1892.
[477] In the absence of any added cyanide, it combines with itself. Silver cyanide, according to Bodländer's results, is, in saturated solutions, chiefly (AgCN)2 or Ag[Ag(CN)2], i.e. Ag─[N═C═C═N─Ag].
[479] Werner has developed quite a different theory of the structure of complex ions. (Cf. Nernst, Theoretical Chemistry, p. 374 (1904).)
[480] Sherrill, Z. phys. Chem., 43, 721 (1903).
[481] In Nessler's reagent, Fresenius' Qualitative Analysis, p. 141.
[482] Cf. Remsen, Am. Chem. J., 11, 291 (1899); 14, 81 (1892) (Stud.).
[483] For instance, for arsenious acid we have
and, therefore, [As3+] × [HO−]3 / ([AsO33−] × [HO+]3) = k1. Since [H+] = k′HOH / [HO−] (p. 176), we have further, [As3+] × [HO−]6 / [AsO33−] = k2. And since we may derive the relation [HO−]2 = k3 × [O2−], by considering the primary and the secondary ionization of water (see pp. 246, 278), we have, finally, [As3+] × [O2−]3 / [AsO33−] = K. The constants for the primary and the secondary ionization of water are included in the value of K.
[484] Fitzgerald and Lapworth, J. Chem. Soc. (London), 93, 2163 (1908); Lapworth, ibid., 2187. Vide also Franklin on the characteristics of the NH4+ ion in liquid ammonia, Am. Chem. J., 23, 305 (1900).
[485] See the laboratory instructions, in regard to the precautions used, to avoid errors from this source.
[486] On the other hand, colloidal organic substances, such as casein, glue or albumen, interfere with the precipitation of even the most insoluble sulphides, by producing colloidal suspensions of the latter (see Chap. VII; cf. Müller, Allgemeine Chemie der Kolloide, p. 56 (1907)).
[487] In alcohols the hydroxide group is held by a carbon atom, whose remaining valences are satisfied by hydrogen or carbon atoms, as in ordinary or ethyl alcohol, H3C─CH2(OH).
[488] Küster, Z. Elektrochem., 4, 117 (1897).
[489] The most common organic acids contain the acid group —CO(OH), as in acetic acid, CH3CO(OH). The hydroxide group OH of the alcohols, e.g. in CH3CH2(OH), is still found in these organic acids, but its tendency to form hydrogen-ion is very much increased by the replacement of two hydrogen atoms of the alcohols by the oxygen atom, as found in the acids. To a certain degree, the properties of the alcohol hydroxide are maintained in the properties shown by the acid hydroxide group. Thus, the organic acids, on the whole, are still rather weak acids, and their salts, in many instances, are appreciably less ionizable than the salts of strong inorganic acids. The organic acids, further, may combine, to a certain extent, with water and thus form hydrates (e.g. CH3COOH + H2O ⇄ CH3C(OH)3) containing a number of hydroxide groups: the second and third hydroxide groups must have a very much smaller tendency to form hydrogen ions and ionizable salts, than has the first hydroxide group (p. 102), and the former, thus show, more nearly, the behavior of alcoholic hydroxide groups. Finally, organic acids also show a tendency to combine with themselves, forming complex acids (e.g., (CH3COOH)2 or CH3C(OH)2OOCCH3), from which complex salts may be derived, which may be little ionizable. The power of the organic acids to form complex ions—which they share with many inorganic acids—is most likely intimately connected with the relations described.
[490] Lead acetate, itself, is less ionized than most salts and this property contributes to the solubility of lead sulphate in acetate solutions. (Cf. Noyes and Bray, loc. cit.)
The analytical groups, which we have heretofore discussed, contain elements, whose oxides are preëminently base-forming. The methods of separation of these groups, from each other, involve, primarily, physical492 differences between the groups—in the matter of the relative insolubility of analogous salts. Thus, barium, strontium and calcium carbonates are precipitated, and separated from the alkalies, by means of ammonium carbonate, not because the alkalies do not form carbonates when their salts, in solution, are treated with ammonium carbonate, but wholly because barium, strontium and calcium carbonates are very difficultly, the alkali carbonates easily, soluble in water. The hydroxides of the aluminium group and the sulphides of the zinc group are less soluble than the hydroxides and sulphides of the alkaline earths and alkalies. The sulphides of the copper and the arsenic groups, again, are still less soluble than the sulphides of the zinc group, and thus the former may be precipitated by hydrogen sulphide, even when its precipitating power is reduced by the suppression of its sulphide (and hydrosulphide) ions by the addition of a strong acid.
On the other hand, the separation of the arsenic group (arsenic, antimony, tin, gold and platinum) from the copper group, with which it is precipitated by hydrogen sulphide from acid solutions, depends, essentially, on a chemical difference between the groups. The oxides, especially the higher oxides, of the arsenic group, are preëminently acid-forming; the higher oxides form such acids as arsenic acid, H3AsO4, antimonic acid, H3SbO4, stannic acid, H2SnO3, platinic acid, H2PtO3, and auric acid, HAuO2. These [p243] hydroxides are, however, all more or less weakly basic in character as well. The hydroxides of the lower oxides of the metals are, as one must expect, much more strongly basic, but most of them—arsenious, antimonous and stannous hydroxides—still show sufficient acid character to be distinctly amphoteric in behavior. But, with the exception of arsenious acid, the basic ionization of the hydroxides of the lower oxides is more pronounced than their acid ionization.
The basic ionization of the hydroxides of their lower and higher oxides brings these elements into the plan of analysis for the metal or positive ions in systematic analysis. In the presence of hydrochloric acid they form chlorides, which yield positive ions in sufficient quantity493 to allow their extremely insoluble sulphides to be precipitated by hydrogen sulphide in acid solution, together with the, likewise, very insoluble sulphides of the copper group.
The acid-forming properties of the oxides of the arsenic group are maintained in their sulphides. Again, this is especially evident in the higher sulphides. The element sulphur is substituted for the closely related element oxygen without any profound change in the chemical behavior of the compounds. Advantage is taken of this acid-forming power to separate the sulphides of the arsenic group from the sulphides of the copper group, which either are not acid-forming at all, or exhibit this property only to a very slight degree.494