CHAPTER IX
«SYSTEMATIC ANALYSIS FOR THE COMMON METAL IONS. THE IONS OF THE ALKALIES AND OF THE ALKALINE EARTHS. ORDER OF PRECIPITATION OF DIFFICULTLY SOLUBLE SALTS WITH A COMMON ION»
In systematic analysis it is most convenient to make separate examinations for the metal and for the acid ions. The examination for metal ions usually precedes that for the acid ions, and the scheme of analysis for the former will be considered first.
The analytical grouping of the metallic elements is not a natural one, as far as their chemical behavior is concerned. Such a grouping is found in the Periodic System of Mendeléeff and is used in systematic inorganic chemistry.[326] The groups in analysis are based chiefly, but not exclusively, on the physical property of greater or smaller solubility of certain salts of the metals. According to the salts chosen, different systems vary somewhat in detail. Frequently elements of the same natural family are also found in the same analytical group, relationship in chemical properties being often coincident with relationship in the physical behavior of the salts of the metals.
In the following list, the common metal ions are arranged in groups, which are given in the order in which they are precipitated in the method of systematic analysis adopted. In each case, a group name and the characteristic reagents used in separating a group from those following it, are given.
«The Silver Group.»—Ions whose chlorides are insoluble in dilute acid solutions. The precipitating agent is HCl. [p158]
«Pb»^{2+} (the chloride is somewhat soluble), «Ag»^{+}, «Hg»^{+}.
«The Copper and the Arsenic Groups.»—Ions whose sulphides are insoluble in dilute acids. The precipitating agent for the two groups is H_{2}S, ‹in acid solution›. The ‹sulphides of the arsenic group are soluble in a mixture of› (NH_{4})_{2}S ‹and› (NH_{4})_{2}S_{x} and are separated thereby from the sulphides of the ‹copper group›.
«The Copper Group.»—«Hg»^{2+}, («Pb»^{2+}), «Bi»^{3+}, «Cu»^{2+}, «Cd»^{2+}.
«The Arsenic Group.»—«As»^{3+}, «As»^{5+}, «Sb»^{3+}, «Sb»^{5+}, «Sn»^{2+}, «Sn»^{4+}, «Pt»^{2+}, «Pt»^{4+}, «Au»^{+}, «Au»^{3+}.
«The Aluminium and the Zinc Groups.»—Ions whose sulphides or hydroxides are insoluble in neutral or slightly alkaline (ammonium hydroxide) solutions. The two groups are precipitated together by a mixture of NH_{4}Cl, NH_{4}OH, (NH_{4})_{2}S. The ‹aluminium group› may be ‹separated› from the ‹zinc group› by treatment of solutions of the chlorides or nitrates with ‹barium carbonate›.
«The Aluminium Group.»—«Fe»^{3+}, «Al»^{3+}, «Cr»^{3+}.
«The Zinc Group.»—«Fe»^{2+}, «Ni»^{2+}, «Co»^{2+}, «Mn»^{2+}, «Zn»^{2+}.
«The Alkaline Earth Group.»—Ions whose carbonates and phosphates are insoluble in neutral or alkaline solutions. The precipitating agent is NH_{4}Cl, NH_{4}OH, (NH_{4})_{2}CO_{3} for «Ba»^{2+}, «Sr»^{2+}, «Ca»^{2+}, and (NH_{4})_{2}HPO_{4} for «Mg»^{2+}.[327]
«The Alkalies.»—Ions whose chlorides, sulphides, hydroxides, carbonates and phosphates are soluble. «K»^{+}, «Na»^{+}, «NH_{4}»^{+}.
In considering the analytical reactions and the analysis of these groups of metal ions, we shall take up the groups in the order reversed to that given in the table. We shall begin with the group of alkali metals, follow this group with the alkaline earths, then take the aluminium and zinc groups, the copper and silver groups, and finish with the arsenic group. This order is chosen because the chemistry of the reactions involved is simplest in the groups to be studied first and grows more complicated as we advance to those to be studied later.
It is not intended to discuss in detail all the reactions and methods; our attention will be limited rather to the study of ‹typical general relations›, and the student is expected to acquire the power to apply the general conclusions reached, to any specific case demanding it. [p159]
«The Alkali Group.»—The group includes the ions of «sodium» and «potassium», the most common and most important of the alkali metals, and the «ammonium»-ion. Characteristic of the ions of the group is the fact that all of their common salts are easily soluble in water. They remain in solution in systematic analysis, while other metal ions are removed in the form of various insoluble precipitates.
The ammonium-ion is recognized, and may be removed from a mixture of the salts of the group, on the basis of a fundamental distinction in its chemical behavior, namely its instability and the instability of its compounds. Sodium-ion and potassium-ion, are recognized, and separated from each other, by physical methods. All ammonium compounds, NH_{4}X, decompose more or less readily into ammonia and the free acids,[328] according to the reversible reaction,
NH_{4}X ⇄ NH_{3} + HX.
The stronger the acid combined with ammonia, the more stable is the salt, and the higher is the temperature, at which the salt decomposes readily and rapidly. Ammonium chloride, one of the most stable of the salts, is decomposed rapidly only at about 350°, which is, however, still below red heat; ammonium carbonate, the salt of a very much weaker acid, decomposes appreciably at ordinary temperatures, and exposed to the air, it gradually disappears as ammonia, carbon dioxide and water. If the acid of the salt is volatile at the dissociation temperature of the salt, the whole salt is volatilized, and if the ammonia and volatile acid vapor reach a colder space, recombination to form the original solid salt occurs to a considerable extent (the "smoking off" of ammonium chloride). If the acid is not volatile, the salt, nevertheless, loses its ammonia at temperatures below red heat, while the acid remains. Sodium-ammonium phosphate, for instance, when heated, loses its ammonia, and sodium-dihydrogen phosphate is left as a nonvolatile residue.
Ammonium happens to form salts which closely resemble the corresponding salts of potassium in physical properties, such as [p160] solubility and insolubility, salts which could readily be mistaken for potassium salts. Advantage is taken of the chemical instability of the ammonium salts, just described, to remove ammonium completely from mixtures, by ignition, before tests for potassium are made.
Water is a far weaker acid (see table, p. 104) than carbonic acid and it is not surprising to find the compound formed by water and ammonia, ammonium hydroxide, one of the least stable of the ammonium compounds. Even at ordinary temperature, the hydroxide is more or less decomposed, according to a reversible reaction of the same type as that found for the ammonium salts, NH_{4}OH ⇄ NH_{3} + HOH.
Chemists have always been interested in the problem of the exact degree of stability of ammonium hydroxide, and, more particularly, in the problem whether ammonia gives solutions in water showing very much weaker basic strength[329] than equivalent solutions of potassium and sodium hydroxides, primarily ‹because only a small proportion of the ammonia is combined with water to form the real base, ammonium hydroxide› (which would be present then in much more dilute solution than the alkali metal hydroxides are in the solutions with which the comparison is made), or chiefly ‹because ammonium hydroxide is much less readily ionizable› than potassium or sodium hydroxide.
For the reversible action NH_{3} + HOH ⇄ NH_{4}OH we would have, at a constant temperature, according to the law of chemical equilibrium,
[NH_{3}] × [HOH] / [NH_{4}OH] = ‹k›.
For a dilute solution at a given temperature the concentration of the water may be considered a constant, and therefore
[NH_{3}] / [NH_{4}OH] = ‹k› / [HOH] = ‹k›_{NH_{3}}. (I)
For the ionization of ammonium hydroxide, NH_{4}OH ⇄ NH_{4}^{+} + HO^{−}, we would have in turn,
[NH_{4}^{+}] × [HO^{−}] / [NH_{4}OH] = ‹k›_{base}. (II)
This constant represents ‹the real ionization constant of ammonium hydroxide as a base›, and its approximate value has only recently been determined by Moore[330] and found to be about 5E−5. The ratio [NH_{3}] / [NH_{4}OH] was found to be approximately 2 at 20°. According to this result, ammonium hydroxide is really a much weaker, less readily ionized base than potassium or sodium hydroxide. The ‹efficiency of ammonium hydroxide as a base› depends [p161] on both conditions of equilibrium; the first equation states what proportion of ammonium hydroxide can exist, as such, in solution, if a given amount of ammonia is dissolved in a given amount of water, and the second equation shows the proportion of the hydroxide, which is ionized. The equations may be combined[331] into one expression,
[NH_{4}^{+}] × [HO^{−}] / ([NH_{4}OH] + [NH_{3}]) = K. (III)
That is, ‹the ratio of› [NH_{4}^{+}] × [HO^{−}] ‹to the total concentration of nonionized ammonium hydroxide and ammonia, is a constant›. This constant has the value 0.000,018 at 18° (as given in the table, p. 106) and, as said, comprises in a single expression a statement measuring the ‹efficiency›, as a base, of a solution of ammonium hydroxide and ammonia in aqueous solutions. The concentration of the hydroxide-ion, on which the efficiency as a base depends, can be ascertained directly from the expression, provided we know the total concentration of the ammonia and ammonium hydroxide and the concentration and degree of ionization of any ammonium salt, which may be present with the base (see p. 161). These data are easily obtained by direct measurement.
The instability of ammonium hydroxide is used as a means for ‹detecting ammonium-ion› in its salts. The latter are treated with some strong base, such as sodium or calcium hydroxide, and the ammonia, liberated by the decomposition of its hydroxide, is recognized by its odor or by the more sensitive test of its action on moist litmus. The delicacy of the test is dependent on the conditions expressed in the equilibrium equation (p. 160).
Sodium and potassium resemble each other so profoundly in the chemical behavior of their compounds, that they are recognized, and separated from each other, by ‹physical› methods. A very simple physical test is based on the color of their heated vapors, the color imparted to the nonluminous bunsen flame by the introduction and volatilization of their salts. The sodium flame is so intense that, if sodium is present in any quantity, the color of its flame easily masks the faint color of potassium vapor. The color of the flame is best examined, in such a case, with the aid of a spectroscope, in which the light emitted by the two elements may be readily recognized side by side, or with the aid of cobalt glass, which absorbs the sodium light. [p162]
The ions of the two metals may be separated and identified by means of difficultly soluble salts. There are so few of these in the case of both ions, that recourse must be taken to the salts of comparatively uncommon acids. ‹Potassium chloroplatinate› K_{2}PtCl_{6}, precipitated by the addition of chloroplatinic acid H_{2}PtCl_{6} to concentrated solutions of potassium salts, gives very satisfactory results, both in qualitative and in quantitative work. The acid tartrate, KHC_{4}H_{4}O_{6}, the picrate, KC_{6}H_{2}N_{3}O_{7}, and the cobaltinitrite, K_{3}Co(NO_{2})_{6},[332] are difficultly soluble and are sometimes used to ‹identify potassium-ion›. The corresponding ammonium salts are also difficultly soluble and resemble the potassium salts, and ammonium-ion must therefore be removed, as stated above, if present, before any of these precipitates may be used for the identification of potassium-ion. In the case of sodium-ion, recourse is taken to the salt of a still more uncommon acid; ‹pyroantimonate of sodium›, Na_{2}H_{2}Sb_{2}O_{7}, 6 aq., is sufficiently difficultly soluble and characteristic to be used as a means of ‹identifying sodium-ion›.
«The Alkaline Earth Group.»—This group includes «magnesium», «calcium», «strontium», and «barium.» Chemically, the analogous compounds of the four alkaline earths resemble one another so much, that physical differences alone are used in their separation and identification. For qualitative work, the colors of their heated vapors,[333] especially when examined in the spectroscope, give us sensitive and reliable tests for their presence. The alkaline earth ions, especially the ions of barium, strontium and calcium, form a great number of insoluble salts, which may be used to separate them from each other. Salts used for this purpose and the methods of employing them are considered in detail in the laboratory work (see