Applications in Analysis.
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 removes[289] 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 precipitate[290] 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.
Chapter VII Footnotes
[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).