Thus, colloidal As2S3, carrying a negative charge, is precipitated by the positive ions of added electrolytes. The addition of a few drops of a molar solution of ammonium nitrate (the precipitating ion is NH4+) to 20 c.c. of the colloidal suspension282 produces a slight precipitate; complete precipitation requires 3.5 to 3.6 c.c. of the ammonium nitrate solution.283 Only 0.06 c.c. of an equivalent solution284 of magnesium nitrate (the precipitating ion is Mg2+) is required, and as little as 0.015 c.c. of an equivalent solution of aluminium nitrate285 (the precipitating ion is Al3+) has the same effect (exp.). An increase in valence of the negative ion, which is not the precipitating ion in this case, does not affect the result appreciably: 3.5 c.c. of a solution286 of ammonium sulphate, (NH4)2SO4, equivalent to the solution of NH4NO3, is also required for the complete precipitation of the colloidal As2S3 (exp.), although the one contains the univalent ion, NO3−, the other the bivalent ion, SO42−.
Conversely, a positively charged colloid, like ferric hydroxide, may be precipitated by much smaller quantities287 of bivalent negative ions than of univalent ions, etc. [p136]
In the second place, if precipitations are attempted either in rather dilute solutions or in solutions of little ionized substances (arsenious acid and hydrogen sulphide), the addition of an electrolyte is frequently required to insure precipitation. Thus, the presence of ammonium chloride, or nitrate, in excess, is helpful in the precipitation of the sulphides of the zinc group; the addition of hydrochloric acid (or other electrolyte) is required to effect the precipitation of arsenious sulphide from a solution of the oxide (p. 126).
In the next place, account must be taken, in analytical work, of the fact that colloids carry down with them the precipitating ion by which they are coagulated, a fact which may lead to the loss of ions which, it is intended, should be kept in solution. To a certain extent, this loss may also be avoided by insuring the presence of electrolytes (acids, ammonium salts) in sufficient concentration to cause the coagulation without the aid of the ions which, it is intended, should not be precipitated. In view of the much weaker precipitating power of univalent ions (of hydrochloric acid, ammonium nitrate and chloride), as compared with that of polyvalent ions, which may be present, the acid and ammonium salts must not be used in too small concentrations. In quantitative analysis, when conditions permit it, ammonium or sodium sulphate is frequently substituted for the ammonium salts of the univalent monobasic acids. The washing of the precipitated colloid with such salt solutions gradually removes289 the ions which are precipitated with the colloid and forms a further safeguard against their loss. But this source of loss is avoided only with great difficulty and is seldom absolutely removed.
Finally, the presence of protective colloids, especially of the [p138] gelatine and albumen type, may interfere so decidedly with the common precipitation tests for ions, that their destruction is imperative, before these tests can be applied with any degree of confidence. Thus, the mixing of solutions (0.1 molar) of silver nitrate and hydrochloric acid, each containing one per cent of gelatine, fails to produce the ordinary, characteristic precipitate290 of silver chloride, the reaction which is used to determine the presence of the silver-ion in systematic analysis.
The mixture is opalescent and, in reflected light, looks opaque-white; on somewhat prolonged standing a white milk is produced, but no precipitate. When the mixture is boiled, the same deep white milk is formed, but no coagulated precipitate, the mixture running unchanged through a filter. Hydrogen sulphide converts the mixture into a similar suspension of the black sulphide.
[227] The ratio is affected somewhat by the fineness of division of liquids and solids as a result of surface tension phenomena, as explained below.
[228] An aqueous solution of iodine and potassium iodide shaken with chloroform gives similar results, and the difference in color between the two layers is an advantage for a lecture experiment. But the iodine is partially combined with the iodide, according to KI + I2 ⇄ KI3, or I− + I2 ⇄ I3−, and the theoretical relations are not so simple as for bromine in aqueous and chloroform solutions.
[229] To accelerate the action, the mixture is shaken vigorously. After the separation of the layers, the bromine may be recognized in the aqueous layer by its color, or by the addition of potassium iodide (liberation of iodine).
[230] [Br]aq. and [Br]ch. are used to express the actual concentrations at any given moment we wish to consider, for instance at the moment equilibrium is reached. The ratio k2 / k1 for any substance S is found to be equal to the ratio of the solubilities of the substance in the two solvents. That it must be so can be proved by applying the law of physical equilibrium to the mixed solvents in contact with an excess of the substance, i.e. to its saturated solutions (see below).
[231] Similar relations hold for the condition of equilibrium between a liquid and its vapor.
[232] Curie, Bull. Soc. Min., 8, 145 (1885); Ostwald, Grundlagen der Anal. Chem., p. 22 (1894).
[233] Hulett, Z. Phys. Chem., 37, 384 (1901). See also Ostwald, ibid., 34, 495 (1900), on the solubilities of finely divided mercuric oxide ("yellow oxide") and of larger crystals ("red oxide").
[234] A 1 / 1000 solution of AuCl3, 2 aq., and a 1.2 / 1000 solution of SnCl2, 2 aq., may be used conveniently. When equal volumes of the solutions are mixed, the desired precipitate is formed.
[235] Stannous chloride is usually sufficiently contaminated with stannic salt. Add a few drops of bromine- or chlorine-water to 100 c.c. of a pure stannous chloride solution.
[236] The action is an exceedingly sensitive qualitative test for gold. By a modification of the test Donau was able to detect as little as 2E−9 gram of gold (Monatshefte f. Chem., 25, 545 (1904)).
[237] The nature of the reduction reaction is discussed in Chapters XIV and XV.
[238] Faraday, Proc. Royal Soc., 8, 356 (1857), Phil. Mag. (4), 13, 401 (1857) (Stud.).
[239] Zsigmondy, Liebig's Annalen, 301, 30 (1898).
[240] Bredig, Z. f. Elektrochem., 4, 514, 547 (1898), and Anorganische Fermente, 1901. Colloidal preparations of platinum, silver and many other metals may be prepared in the same way. In ether, colloidal preparations of the alkali metals may be made (Svedberg, Ber. d. chem. Ges., 38, 3616 (1905), 39, 1708 (1906)).
[241] Argentum Credé may be conveniently used. It contains, with the metallic silver, a small percentage of albumen, which is added for reasons discussed below. Brown solutions are formed at once.
[242] First a thin film of black carbon is produced round the metal, then the latter appears in the form of a filigree of silver.
[243] For details of the preparation, see A. A. Noyes, J. Amer. Chem. Soc., 27, 94 (1905).
[244] This process of separation of substances, which do not pass through membranes, from such as do, is called dialysis. It was first used by Graham, Trans. Royal Soc., London, 151, 183–224 (1861) (Stud.).
[245] Graham made the first extended investigations in this field: Trans. Royal Soc., 151, 183 (1861); J. Chem. Soc. (London), 17, 318 (1864) (Stud.). He found that amorphous, gelatinous bodies like ferric hydroxide, aluminium hydroxide, silicic acid, gelatine, glue, dextrin, caramel, albumen and similar bodies do not pass through membranes and may be obtained by dialysis in the colloidal condition. Such substances were called "colloids" by Graham, the name referring to the Latin for gelatine. Substances which pass through membranes readily were found by Graham to resemble in behavior such bodies as are crystallizable when solid; such compounds were classified by him as "crystalloids." That liquids containing substances in the colloidal condition (e.g. arsenious sulphide, gold, silver and many other substances) may be prepared by methods other than dialysis, was found later by many investigators and, in a few cases, previous to Graham, e.g. by Faraday, loc. cit. A brief history of the chemistry of colloids is found as an introduction to Wo. Ostwald's Kolloidchemie, pp. 1–63 (1909).
[246] Cf. Wo. Ostwald, loc. cit., p. 193.
[247] Before Graham's time, and for the few colloidal liquids then known, this view was held by such men as J. B. Richter, Berzelius and Faraday (loc. cit.) (cf. Wo. Ostwald, loc. cit., p. 19). The first extended experimental investigation in support of it was made by Barus and Schneider, Z. phys. Chem., 8, 278 (1891). Bredig was also an early and consistent champion of this view (vide his Anorganische Fermente, 1901).
[248] Cf. Wo. Ostwald, loc. cit., pp. 102–114. Graham considered "colloidal silicic acid a liquid miscible with water in all proportions." According to modern ideas, no true miscibility exists, but a suspension or emulsion is formed (see Ostwald, p. 237).
[249] Siedentopf and Zsigmondy, Drude's Annalen, 10, 1 (1903). Zsigmondy, Colloids and the Ultramicroscope (1909), Chapter V.
[250] Zsigmondy, Z. für Elektrochem., 8, 684 (1902); Siedentopf and Zsigmondy, loc. cit.
[251] Zsigmondy, loc. cit., p. 161. A µµ is 1E−6 mm. The hydrogen molecule is considered to have a diameter of 0.1 µµ (O. E. Meyer), the alcohol molecule one of 0.5 µµ (Zsigmondy, loc. cit., plate IV, p. 157).
[252] Weimarn, Chem. Zentralblatt, 1907, II, p. 1293.
[253] Paal, Ber. d. chem. Ges., 39, 1436, 2859 (1906).
[254] Other varieties of heterogeneous colloidal mixtures are tabulated by Wo. Ostwald, loc. cit., p. 96.
[255] The "Cassius' purple" test for gold is an instance where the colloidal condition is used in analysis for a positive test. See Wo. Ostwald, loc. cit., p. 68, for other, similar applications for positive tests.
[256] Vide Wo. Ostwald's Kolloidchemie, 1909, and the references given by Noyes, loc. cit., p. 86.
[257] A general discussion of the preparation and properties of colloidal mixtures is given by A. A. Noyes, J. Am. Chem. Soc., 27, 86–104 (1905) (Stud.).
[258] Picton and Linder, J. Chem. Soc. (London), 61, 160 (1892), 67, 63 (1895), 71, 568 (1897), etc., and others. Wo. Ostwald, loc. cit., p. 240, gives a list of references.
[259] The arrangement of the experiment is described by Noyes, loc. cit., p. 98.
[260] Cf. Wo. Ostwald, loc. cit., p. 108. Billitzer has found that gelatine is positive in acid solution, negative in alkaline, Z. phys. Chem., 51, 147 (1905). The charges are, however, relatively small ones.
[261] Billitzer, Z. f. Elektrochem., 8, 638 (1902). This is probably true of all amphoteric colloids (Chapter X); it is also true of many other substances, which are not pronouncedly amphoteric. (Cf. Perrin, Comp. rend., 136, 1388 (1903); Billitzer, Z. phys. Chem., 51, 157 (1905).)
[262] Hardy, J. of Physiology, 24, 288 (1899); Z. phys. Chem., 33, 387 (1900).
[263] Billitzer, loc. cit., p. 159; Müller's Allgemeine Chemie der Kolloide, 1907, p. 79.
[265] In a slightly acid solution colloidal silicic acid is negatively charged; in a strong acid solution, positively—a relation which agrees with its predominantly acid character.
[266] The general class of substances, showing both basic and acid properties, of which albumen is a derivative, is described in a footnote on glycocoll, Chapter X, p. 188.
[267] Hardy, loc. cit.
[268] J. Loeb, University of California Publications, Physiology, 2, 149 (1904).
[269] Very little is known about the nature of contact electricity. It is even doubtful whether it is different, in principle, from ionization.
[270] W. Ostwald, Lehrbuch der Chem., 2, (1) 553 (1903).
[271] Hardy, Proc. Royal Soc., 66, 110 (1899); Z. phys. Chem., 33, 391 (1900).
[272] Picton and Linder, J. Chem. Soc. (London), 67, 63 (1895).
[273] Whitney and Ober, J. Am. Chem. Soc., 23, 852–856 (1901) (Stud.).
[274] Picton and Linder, loc. cit. 71, 572 (1897); Lottermoser, Anorganische Kolloide, p. 76; Biltz, Ber. d. chem. Ges., 37, 1095 (1904).
[275] Precipitation is complete only when the colloids are used in the proportions required to neutralize each other's charges [Billitzer, Z. phys. Chem., 51, 140 (1905)]. The proportions to be used must be determined in each case, most simply by trial (Noyes, loc. cit., p. 101), but quantitative methods for determining the charges, by titration, are also known (cf. Billitzer, loc. cit.).
[276] E. A. Schneider, Z. anorg. Chem., 5, 80 (1894).
[277] See the above discussion on silicic acid. Stannic acid has a greater tendency to form a base than has silicic acid.
[278] Zsigmondy, Liebig's Annalen, 301, 361 (1898).
[279] Cf. Fresenius, p. 334, or Smith's Inorganic Chemistry, p. 468.
[280] Mylius, Ber. d. chem. Ges., 36, 775 (1903); Biltz, ibid., 37, 1116 (1904).
[281] J. für prakt. Chem., 25, 431 (1882).
[282] The suspension used is prepared by saturating, with hydrogen sulphide, an aqueous solution of arsenious oxide. The latter is saturated on a steam bath, cooled to 20°, filtered and diluted with an equal volume of water before it is used.
[283] Eight grams of NH4NO3 per 100 c.c.
[284] 0.6 c.c. of a tenth-normal solution is used, containing 1.28 gram Mg(NO3)2, 6 aq., in 100 c.c. Precipitation was found to be incomplete with 0.5 c.c.
[285] 0.15 c.c. of a tenth-normal solution is used, containing 1.2 gram Al(NO3)3, 4 aq., in 100 c.c. Precipitation was found to be incomplete with 0.1 c.c. of the solution.
[286] 3.3 grams (NH4)2SO4 in 100 c.c.
[287] Freundlich found, for instance, that NaCl, KCl, BaCl2, in equivalent concentrations, had practically the same effect on colloidal ferric hydroxide, but only one-fortieth as much of a sulphate (the precipitating ion is SO42− versus Cl−) was required; Z. phys. Chem., 44, 129 (1903).
[288] For other reasons (e.g. to prevent oxidation of the sulphides), hydrogen sulphide is also used in the solution for washing the arsenic group, and ammonium sulphide in that for the zinc group (see Lab. Manual, pp. 101 and 110).
[289] Picton and Linder, loc. cit., and Whitney and Ober, loc. cit.
[290] A. A. Noyes describes a similar experiment with sodium chloride and silver nitrate, loc. cit.
It frequently happens that we have to deal, simultaneously, with conditions of chemical and of physical equilibrium, obtaining in the same system. For instance, a gas like carbon dioxide, in contact with its saturated solution in water, is in equilibrium with the dissolved carbon dioxide, and this, in turn, is in equilibrium with its hydrate, carbonic acid. A substance may be distributed between two solvents and show a different molecular weight in the two (see p. 18); it may exist, in the one, primarily in polymeric form, and only to a slight extent in the simple form, the two forms being in equilibrium (chemical equilibrium). In the other solvent, it may exist only in its simple molecular form, and this will be in equilibrium with the same simple molecular form in the first solvent (physical equilibrium). In matters dealing with the solubility of electrolytes in water, and, therefore, in questions of their precipitation or solution, such simultaneous conditions of chemical and physical equilibrium are constantly occurring. Since qualitative analysis deals to a very considerable extent with just such precipitates of salts, acids and bases, these cases are of particular importance to us.
When water is added to solid silver acetate, the salt will dissolve. If an excess of the acetate is used, equilibrium will result between the solid salt and its solution, when the solution is saturated at the temperature used. As the salt dissolves, it is more or less ionized, and in the saturated solution we have a [p140] condition of chemical equilibrium between the salt and its ions:
If the law of chemical equilibrium is applied to this reversible action, we have (p. 98)
The nonionized silver acetate is present in two phases, in the solid phase and also in solution:
Applying the law of physical equilibrium to this system, we have further (p. 121)
The concentration of a pure solid, as we have seen, may be considered a constant at a given temperature. Consequently, if we consider the question of the size of the solid particles as a minor factor and negligible, we shall conclude, that, for saturated solutions of silver acetate, the concentration of the solid silver acetate being a constant, the concentration of the nonionized or molecular silver acetate, the first term of our constant ratio II, must also have some definite, constant value at a given temperature. We may call this concentration the "molecular solubility" of silver acetate and may put
for a saturated solution of silver acetate in water at the given temperature. Now, since the concentration of the nonionized silver acetate, [CH3COOAg], in the saturated solution also forms the second term of equation I, representing the condition of chemical equilibrium between the acetate and its ions, we obtain, by combining I and III,
Further, if the assumption is made that the presence of foreign electrolytes, in not too concentrated solutions, does not affect either the molecular solubility, Kmol. sol., of silver acetate or its tendency to ionize—as expressed in KIonization—then, the [p141] relation, which has been developed, would hold for saturated aqueous solutions of silver acetate in the presence of foreign electrolytes, as well as for a saturated, pure, aqueous solution. A single, simple equation would thus express the conditions for simultaneous chemical and physical equilibrium between a difficultly soluble ionogen, of the type of silver acetate, and its saturated solutions, at a given temperature, in the presence or the absence of foreign electrolytes.
It appears, however, that while the soundness of this theoretical development of the relations expressed by the solubility-product must be questioned, nevertheless as a matter of experiment, the product of the ion concentrations of a difficultly soluble salt is found, in dilute solutions, to be a constant, or sufficiently close to a constant to satisfy all but the most rigorous requirements.295
It is, in fact, quite evident, that a decreasing value for the second term of the ratio I—namely, for [CH3COOAg], the molecular solubility of the salt—as the total concentration of the electrolytes present increases, together with an increasing value for the whole ratio I under the same conditions, are not incompatible with a constant value of the first term of the ratio. That is, the product of the ion concentrations, [CH3COO−] × [Ag+], may well remain constant (equation IV), or approximately constant, in dilute salt solutions, even if equations I and III do not hold for salt solutions. [p143]
Whether in the case of all difficultly soluble salts, as the total salt concentration increases, the increasing values of the chemical equilibrium ratio (equation I) will be so nicely balanced by the decreasing values of the molecular solubility, that the first term of the first ratio (the solubility-product) will always be constant, is a question demanding further extended investigation.296 The range of the investigation must be extensive, because it must include several other classes297 of salts (e.g. Me2X, MeY2, etc.), for which the first equation has a different form; for instance, for Me2X,
For the present we must remain content with the result of the past investigations and consider the principle of the constant solubility-product to be essentially an empirical one. It is an extremely convenient condensation, into a very simple mathematical form, of the main factors involved in the precipitation and solution of difficultly soluble salts, acids, and bases. It should be used with due knowledge of its character and limitations.
Washburn298 derives the principle of the constancy of the solubility-product, without involving in his derivation the relation between ions and nonionized molecules—a relation which, as was stated above, deviates from the law of chemical equilibrium. The deviation, it will be recalled (p. 109), is generally supposed to be due to the fact that the fundamental kinetic assumption which must be made to derive the law of chemical equilibrium from the kinetic theory, the assumption that there should be none but negligible forces of attraction and repulsion between the molecules (of a gas or solute) which are in equilibrium, is not fulfilled in the case of solutions of strong electrolytes (p. 109). According to Washburn, if it is assumed that the ions of an electrolyte fulfill this fundamental condition and that only the nonionized molecules do not—the latter causing the deviation from the law of chemical equilibrium—then the principle of the solubility-product constant follows.299 He sees an approximate confirmation of the assumption made, in the fact that the principle is found, empirically, to be true, and that other relations, developed on the basis of the same assumption, agree with the observations made. [p144]
This theoretical derivation of the principle, like the derivations of the law of chemical equilibrium and of all our laws of dilute solutions, assumes300 that the nature of the solvent, and consequently of the solution-process, is not changed by added substances, for instance by an excess of the precipitating ionogen. There can be no question, however, that the nature of the solvent must change, as a continuous function, by the addition of electrolytes to solutions. The changed solubilities of inert gases in salt solutions,301 and a mass of other evidence,302 lead to this conclusion. The addition of a half mole of sodium chloride to a liter of water reduces the dissolving power of the liquid towards oxygen at 25° by 15%, i.e. by 30% per mole of salt. A weak electrolyte, such as acetic acid, has practically no effect at this concentration, and so the effect must be chiefly due to the ions of sodium chloride; since the salt, in half-molar solution, is ionized 73%, the reduction in the dissolving power would be 30 / 0.73 = 41% per mole of fully ionized salt. The principle of the constant solubility-product cannot be considered as established for solutions more concentrated than 0.2 to 0.3 molar; but it is evident that, in any comprehensive theoretical formulation of the principle for the range in which it is found empirically to hold, the change in the nature of the solvent, which in some cases is conspicuous in 0.5 molar solution, must be taken into consideration as a factor even in more dilute solutions (say 0.05 to 0.3 molar). It seems at present, quite possible, perhaps even probable, that the constancy, in all but the most dilute solutions, is the result of the approximate balancing of two (or more) opposing factors.303 When we leave the range of concentrations mentioned, and go to more concentrated solutions, these factors seem to be less well balanced and the validity of the principle ceases.304 For the present it will be safe to consider the principle as an empirical one, holding for solutions of total salt content up to 0.25 or 0.3 molar.305 For quite dilute solutions the effect of the electrolyte on the solvent would be negligible, and only to such solutions would the theoretical derivation brought forward by Washburn be applicable.