Systematic Analysis for Acid Ions, Based on the Removal of Metal Ions other than the Alkali Metal Ions.

—On account of this kind of interference, by cations other than those of the alkali group, with a number of group and specific tests for anions, that may be made in neutral or alkaline solutions, provision is made, in most systems of analysis, for the removal of such ions by proper treatment of the substance under examination with sodium carbonate. Interfering metal ions are thereby converted into carbonates or hydroxides which are insoluble in water, while the acids form sodium salts which pass into solution in water. Occasionally, recourse is also taken to hydrogen sulphide to remove ions of the arsenic and copper groups.

The treatment with sodium carbonate, while advantageous in certain cases, is not uniformly successful590: it is also frequently complicated by the presence of amphoteric bases or of organic substances, and frequently demands treatments and tests beyond those made with the solution thus prepared. Furthermore, if a group is found to be present, say a group of acid ions forming barium or calcium salts insoluble in water, but soluble in acids,591 all the acids of the group, which have not been found to be present or absent in the analysis for metal ions,592 must be specifically [p301] tested for, although they may all be absent and the only representative of the group present may be an ion previously found in the analysis for metal ions.593 Then, too, when the group is found to be absent by the group test, more sensitive tests for some of the acid ions of the group must still be made to insure their complete absence.594

While much may be said in favor of the systematic analysis for acid ions, based on the preparation of a solution containing only the alkali salts of the anions, and while one should be familiar with the plan and be able to have recourse to it at will, yet the drawbacks mentioned suggest another basis for the analysis.

Systematic Analysis for Acid Ions in Acid Solution.

—If the systematic analysis for acid ions is carried out entirely in acid solutions, interference of cations with tests for anions is rarely met with and in those rare cases may be easily provided against. Such a method of systematic analysis in acid solution is frequently more direct and more convenient than a method based on the removal of cations other than the alkali metal ions. Almost all of the most characteristic tests for anions, as it is, are carried out in acid solutions, and only a very few good tests, which must be made in neutral or alkaline solution, are sacrificed by placing the emphasis on those carried out in acid solutions. The method, on the whole, has proved a time-saving and convenient one, without loss in the trustworthiness of the results.

Desirability of Experience with Both Methods.

—It is desirable to have experience with both methods, and to learn by such experience, when to have recourse to the one or the other method. Suggestions as to the choice of method are given in the Laboratory Manual, p. 119.

The Groups of Acid Ions.

—The arrangement of the acid ions into groups, for analysis in acid solution, does not differ in any essential respect from the arrangement based on the use of solutions of sodium salts of the acid ions—only one group test, which must be made in neutral solutions, is omitted when the acid solution is used, and the individual members of this group are tested [p302] for, specifically. As that is also very frequently necessary when solutions of sodium salts are used, no notable sacrifice in convenience is made.

Since the grouping for both methods of analysis may be made the same, the following grouping of acid ions has been adopted. Only the group characteristics are given; the members of the groups are described in detail in the Laboratory Manual (Part III).

I. Ions of Amphoteric Acids and of Related Acids. This group includes those acids whose amphoteric character, or whose ready reduction by hydrogen or ammonium sulphide, leads to their being found, or indicated, in the systematic analysis for metal ions.

II. The Carbonate Group. This group includes those acids whose physical properties (insolubility), or the physical properties of decomposition products of which (carbon dioxide is a decomposition product of carbonic acid), usually lead to their discovery in the course of the preparation of solutions or of the analysis for cations.

III. The Sulphate Group. The barium salts of this group of acid ions are insoluble in acid solution. Barium nitrate, added to a solution acidified with nitric acid, is the group reagent.

IV. The Chloride Group. The anions of this group form silver salts, which are insoluble in nitric acid. Silver nitrate, added to a solution acidified with nitric acid, is the group reagent.

V. The Phosphate Group. A test for this group, as a whole, can be made only if cations other than the alkali metal ions are absent: the barium salts of the acid ions of the group are insoluble in neutral, but soluble in strongly acid, solutions. Barium nitrate, used with a neutral solution, is the group reagent in the absence of metal ions other than the alkali metal ions. In the presence of other cations, the three members of the group, which are not found in some other group,595 namely: phosphate, borate and fluoride ions, are tested for specifically, and the group test omitted. Phosphate-ion is tested for in nitric acid solution, in the same solution as is used for the tests for groups III and IV. [p303]

VI. The Nitrate Group. The salts of the acids of this group are readily soluble in water and specific tests for the acid ions are made; there is no group test.

VII. The Group of Organic Acids. This group need only be considered when a test for organic matter reveals its presence.

Applications of Physico-Chemical Principles and Theories.

—The physico-chemical principles and theories, which have been developed in the previous chapters, naturally apply also to the reactions by which acid ions are identified. In many cases such characteristic reactions are identical with reactions studied in connection with the metal ions. For instance, the precipitation of silver chloride, used to identify the silver-ion (reagent, chloride-ion), may be used, with certain precautions, to identify chloride-ion as well (reagent, silver-ion).

In the following, only a few typical and interesting applications of the principles and theories to acid ions will be given; numerous other applications will suggest themselves in connection with the laboratory work on the acids.

Fractional Precipitation of Salts with a Common Ion.

For saturated solutions of silver chloride, bromide and iodide, we have, according to the principle of the solubility-product596:
[Ag+]1 × [Cl−]1 = KAgCl = 1E−10;
[Ag+]2 × [Br−]2 = KAgBr = 4E−13;
[Ag+]3 × [I−]3 = KAgI = 3E−16.

For a solution saturated simultaneously with the three silver salts, the value of the concentration of the silver-ion is the same in the three solubility-products (see p. 164). Consequently, for such a solution, with which the three solid salts are in equilibrium, the ratios of the concentrations of the anions must be: [Cl−] : [Br−] : [I−] = KAgCl : KAgBr : KAgI = 3 × 105 : 1300 : 1. That is, if silver nitrate is added to a mixture of iodides, bromides and chlorides, silver iodide must be precipitated first, until the concentration of bromide-ion, in solution, is 1300 times as great as the concentration of the iodide-ion left in solution. Then bromide and traces of iodide of silver will be precipitated, until the concentration of chloride-ion is 300,000 times as great as the concentration of iodide-ion and some 250 times as great as the concentration of the bromide-ion. In other words, if silver nitrate is added gradually to such a mixture, iodide-ion and bromide-ion will be almost completely removed from solution before a precipitate of silver chloride can be in equilibrium with the solution. This gives us a convenient and rapid method of detecting chlorides, if present, [p304] in more than small quantities, with iodides and bromides. Silver nitrate, a few drops at a time, is added to the solution and the mixture vigorously shaken after each addition. As long as a yellow (AgI) or yellowish (AgBr) silver salt is precipitated on the addition of silver nitrate to the supernatant liquid (the precipitate settles quickly), silver nitrate is added as before; when the color becomes quite pale, the solution is filtered and silver nitrate added a drop at a time; if a pure white precipitate results finally, chloride-ion is present in the mixture (exp.).

Complex Ions.

Instances of the rôle of complex ions in the analysis for acid ions are numerous. One of the most interesting illustrations is the application of the equilibrium conditions for the complex silver-ammonium-ion (p. 224) to the separation of silver chloride, bromide and iodide. A rather more convenient and more sensitive method597 for detecting the three halide ions in the presence of each other than the method just considered, may be discussed from this point of view.

The condition of equilibrium between silver-ion, ammonia and silver-ammonium-ion is expressed in the relation:

[Ag+] × [NH3]2 / [Ag(NH3)2+] = K = 1 / 107.

The concentration of silver-ion, which may exist in an ammoniacal solution, evidently must decrease rapidly with increasing concentrations of the free ammonia. Now, let us imagine only sufficient free ammonia, in solution, added to a mixture of silver chloride, bromide and iodide, to keep the concentration of silver-ion, which can exist in the solution, say at [Ag+] = 6E−9, which is just 1 / 100th of the concentration of silver-ion in a saturated aqueous solution of silver bromide. Such a solution of ammonia, in contact with the three silver salts mentioned, will dissolve silver chloride, if sufficient is present, until [Cl−] = KAgCl / 6E−9 = 0.017 molar. At the same time, silver bromide would be dissolved until [Br−] = KAgBr / 6E−9 = 0.000,06 molar. In other words, silver chloride could be dissolved in some quantity, while silver bromide is dissolved only in traces (the ratio of [Cl−] : [Br−] is again about 250 : 1). When such an ammoniacal extract is acidified with nitric acid, almost pure (white) silver chloride would be precipitated and only traces of bromide would be lost. After the extraction of the chloride, an increased concentration of ammonia would lead, similarly, to a solution in which silver bromide would dissolve readily and only traces of the iodide be lost, and thus a separation of bromide and iodide may be effected.

In Hagar's method, the concentration of ammonia, required to dissolve silver chloride with but traces of bromide, is attained by the use of a solution of ammonium sesqui-carbonate,598 in which free ammonia is present only in small concentration, as a result of the hydrolysis of the salt. After the [p305] extraction of the chloride by this solution, the bromide is extracted with a 5% solution of ammonia.

Complex Ions of Acid Ions with Other Acids.

—In the study of complex ions we found that positive ions (silver, cupric, etc.) may form complex positive ions with ammonia599 or complex negative ions with acid ions (e.g. with cyanide-ion). In the study of the acid ions we also meet instances of complexes formed by the union of two acids to form a new complex acid. Ammonium phosphomolybdate, an important salt that is extremely useful in detecting the presence of phosphate-ion, is the most interesting instance of the salt of such an acid, which is met in elementary qualitative analysis.600

Ammonium phosphomolybdate, (NH4)3PO4, 12 MoO3, is the salt of a complex phosphomolybdic acid, formed from phosphoric acid, O:P(OH)3, and molybdic acid, O2Mo(OH)2, by a loss of water, much as potassium dichromate is formed from potassium acid chromate [KO(CrO2)OH + HO(CrO2)OK ⇄ KO(CrO2)O(CrO2)OK]. The only difference between the two actions lies in the fact that, in the case of the dichromate, anhydride formation occurs between two molecules of a single acid; in the case of the phosphomolybdate, anhydride formation takes place between molecules of different acids, and a much larger number of molecules is involved. If we suppose the combination between the two acids to proceed symmetrically,601 we may consider the following to be the action:

O:P(OH)3 +
3 [HO(MoO2)OH + HO(MoO2)OH + HO(MoO2)OH + HO(MoO2)OH] ⇄
O:P[O(MoO2)O(MoO2)O(MoO2)O(MoO2)OH]3 + 12 H2O.

Intermediate complex acids, containing less molybdic acid, are no doubt formed first (the action is a relatively slow one), and the action proceeds until the formation of an insoluble salt leads to the final precipitation of all of the phosphate in this form. The precipitate shows the characteristic behavior of an acid anhydride—alkalies dissolve it readily and form phosphate and molybdate—e.g. ammonium hydroxide forms [NH4]2HPO4 and (NH4)2MoO4 (exp.). Dichromates, in a similar way, are converted by alkalies into chromates, an action which may readily be followed by the change in color (exp.).

Oxidation and Reduction.

—While fractional precipitation of silver iodide, bromide and chloride, and fractional solution of the silver salts in ammonia are convenient methods for detecting the three halide ions in the presence of one another, the most accurate and most convenient methods for this purpose depend on the different sensitiveness which iodide, bromide and chloride ions exhibit towards oxidizing agents. Of the three halogens, iodine shows the smallest tendency to form its ion (see the table, p. 294), chlorine the greatest. Vice versa, of the three halide ions, iodide-ion is most readily, chloride-ion least readily, oxidized. Treatment with a mild oxidizing agent, such as ferric-ion [p306] (see Chap. XIV and Laboratory Manual under iodide-ion), suffices to oxidize iodide-ion to iodine: 2 Fe3+ + 2 I− ⥂ 2 Fe2+ + I2. Bromide-ion and chloride-ion are left practically unaffected by this agent (see Chap. XIV). A somewhat stronger oxidizing agent, chromic acid (or its ion Cr6+, see Chap. XV), oxidizes bromide-ion and leaves chloride-ion practically unaffected: 2 Cr6+ + 6 Br− ⥂ 2 Cr3+ + 3 Br2. This method of fractional oxidation forms one of the most convenient and sensitive methods for detecting the three halide ions in the presence of one another.602

We shall discuss here only one other oxidation-reduction reaction, taken in connection with the laboratory work—the oxidation of hydroiodic acid by exposure to the air and the resistance to oxidation shown by an iodide, such as potassium iodide, under the same conditions. The following method of proximate analysis of the chief relations involved may also be used to interpret the contrast in the behavior of hydroiodic acid and that of hydrobromic or hydrochloric acid (Laboratory Manual, q. v.). In all of these cases the actual relations are rendered more complex in consequence of secondary reactions, than is indicated in the text that follows: it is intended only to outline the most effective of the factors involved and to illuminate the qualitative results observed.

Oxidation of Hydroiodic Acid by Air.

—The oxidation of hydroiodic acid, or of potassium iodide, by the oxygen of the air may be considered (Chapters XIV and XV) to involve primarily the action
4 I− + O2 + 2 HOH ⇄ 2 I2 + 4 HO−.
(1)

The condition for equilibrium will be

[I−]4 × [O2] / ([I2]2 × [HO−]4) = Kequil.
(2)603

A system in which I− is directly in equilibrium with I2 (for which [I−]12 : [I2]1 = KI−, Iodine = 5.6E29, at room temperature (p. 298)) and in which, at the same time, HO− is directly in equilibrium with O2 (for which at room temperature [HO−]14 : [O2]1 = KHO−, Oxygen = 8.2E49 (p. 298)) would also represent a condition of equilibrium for the four components. We find thus

Kequil. = KI−, Iodine2 / KHO−, Oxygen = (5.6E29)2 / (8.2E49) = 4E9.
(3)

With the aid of this constant and of equation (2) we can obtain, at least, an approximate interpretation of the results of the exposure of hydroiodic acid and of potassium iodide to the influence of atmospheric oxygen.604 We may [p307] calculate, first, what concentration of free iodine would be required to prevent oxidation of hydroiodic acid, in molar solution, by the oxygen of the air, i.e. to establish equilibrium. We will call x that concentration of I2. As hydroiodic acid is a very strong acid, ionized to the extent of about 80% in molar solution, we may, with sufficient accuracy for our purpose, consider it completely ionized and put [I−] = 1 and [H+] = 1. Since at 25° [H+] × [HO−] = 1.2E−14 (p. 104), we may put [HO−] = 10−14. The concentration of oxygen in the air, at room temperature, may be considered to be approximately [O2] = (1/5) × (1 / 23.9). Inserting all these given values in equation (2), we have

[I−]4 × [O2]  =  1 × (1/5) × (1 / 23.9)  = 4 × 109.
[I2]2 × [HO−]4 x2 × (10−14)4
(4)

Solving for x, we find x = 1022 = [I2]. That is, free iodine of this enormous concentration would be required to prevent oxidation of hydroiodic acid in molar solution by the oxygen of the air at room temperatures. It is obvious that hydroiodic acid must be extremely sensitive to oxidation by exposure to air.

One might estimate, in a similar way, the extent to which hydroiodic acid, of a given concentration, would be oxidized by air before equilibrium would be reached. The process would involve simultaneous changes in three factors—iodide-ion is destroyed, iodine is formed and hydroxide-ion increases, as the result of the neutralization of hydrogen-ion by the hydroxide-ion formed in the action (see above). The solution of the equilibrium equation is too involved for the elementary purposes of this discussion: it leads to the same qualitative conclusion as was just reached.

Oxidation of Potassium Iodide by Air.

We may now ask what the relations would be, if we used a molar solution of potassium iodide in place of the free acid. [I−] and [O2] would have the same value as before. The solution being originally neutral, [HO−] would at first have the value √(1.2E−14) = 1.1E−7. But when potassium iodide is exposed to the air, if iodine is liberated, the solution becomes alkaline605 (HO− is formed according to equation (1)) and the concentration of HO− consequently grows continuously greater. We will, therefore, formulate the problem as follows: how much iodine606 must be liberated, by oxidation of iodide-ion, in molar potassium iodide solution in order to establish equilibrium? For every two molecules of iodine liberated, four HO− ions are formed (equation (1)). If we call y the concentration of iodine at the point of equilibrium, then 2 y is the concentration of [HO−] at that point, formed by the oxidation process. Inserting the given values607 in equation (2), we have
[I−]4 × [O2]  =  1 × (1/5) × (1 / 23.9)  = 4E9.
[I2]2 × [HO−]4 y2 × (2 y)4
[p308]

Solving for y, we find y = 0.007. That is, in molar solution, about 1.4% of the iodide608 would be oxidized (carbonic acid and other acids being excluded); in 5 c.c. (see Lab. Manual, p. 73) 9  milligrams of iodine609 would be liberated to reach a condition of equilibrium.610

It is thus clear that the conditions for equilibrium between a solution of an iodide and air would be satisfied, in the case of an alkali iodide, by the liberation of a mere trace of iodine, whereas, as was previously shown, in the case of hydrogen iodide, a very large proportion of iodine must be liberated before equilibrium could obtain. A careful comparison of the two developments shows that the difference in result610 is plainly due to the higher oxidizing power, the higher potential of oxygen (p. 280), in acid solutions, containing only a minute concentration of hydroxide-ion, as compared with its efficiency in neutral or slightly alkaline solution.

Chapter XVI Footnotes

[589] In the few cases when there is interference, it is provided against.

[590] Cf. Fresenius, Qualitative Analysis, p. 520.

[591] The group includes phosphate, borate, fluoride, oxalate, silicate, arsenite, arseniate, chromate and tartrate ions.

[592] The ions of the amphoteric acids, arsenic and arsenious acids, and chromate-ion, which is reduced by hydrogen sulphide to chromium-ion, are found in the systematic analysis for metal ions.

[593] The ions of the amphoteric acids, arsenic and arsenious acids, and chromate-ion, which is reduced by hydrogen sulphide to chromium-ion, are found in the systematic analysis for metal ions.

[594] See Fresenius, loc. cit., p. 511, footnote, and p. 520.

[595] Other acid ions which would show the group test—precipitation of a barium salt in a neutral solution—are determined in other groups, as follows: arsenite, arseniate and chromate ions in the group of amphoteric acids, etc. (I); carbonate and silicate ions in the carbonate group (II); and oxalate and tartrate ions in the group of organic acids (VII).

[596] The constants refer to 18°. The subindices are used to distinguish the (unequal) concentrations of the silver-ion and of the halide ions in the different solutions referred to in the text.

[597] Hagar's method. See Fresenius, loc. cit., pp. 356 and 378.

[598] See Fresenius, loc. cit., pp. 356 and 378, for the preparation of the solution and for details of the method, and see Smith, General Inorganic Chemistry, p. 566, as to the nature of the sesqui-carbonate.

[599] They also form complex ions with substances related to ammonia, such as the organic amines.

[600] Ammonium arsenomolybdate is an analogous salt (see Laboratory Manual, Part III). A similar complex acid, phosphotungstic acid, is used in alkaloidal analysis.

[601] The exact structure of the complex acid is not known.

[602] In the Laboratory Manual a second, similar method is also given.

[603] It is considered that water has a constant concentration in a dilute solution and that for its active components [H+] × [HO−] is a constant (p. 176).

[604] In the calculation which follows, which is meant merely for a rough survey, no account is taken of the formation of complex ions I3−, or of the tendency of hydroiodic acid to decompose spontaneously into iodine and hydrogen: 2 H+ + 2 I− ⇄ H2 + I2, a reaction which could also be studied profitably with the aid of the equilibrium constants for I2 ⇄ 2 I− and for H2 ⇄ 2 H+. The value of the iodide constant is also uncertain (see p. 273).

[605] 4 K+ + 4 I− + O2 + 2 HOH ⇄ 2 I2 + 4 K+ + 4 HO−.

[606] The formation of complex ions I3− and other secondary reactions (formation of hypoiodite, iodate, etc.) are ignored.

[607] y has so small a value that we may consider [I−] practically unchanged.

[608] 0.007 I2 = 0.014 I−.

[609] A mole of I2 = 2 × 127 = 254 grams; 0.007 × 254 × 5 / 1000 = 0.009 gram.

[610] The tendency of iodine to form hypoiodous acid, iodates, etc., is not taken into consideration here and involves another relation.

INDEX

Numbers marked (†) refer to subjects illustrated by experiments, heavy numbers refer to tables.