CHAPTER XV

«OXIDATION AND REDUCTION. II. OXIDATION BY OXYGEN, PERMANGANATES, ETC.; OXIDATION OF ORGANIC COMPOUNDS»

We will turn now to the consideration of the question, how the principles of the theory of electric oxidation and reduction may be applied to the most important oxidizing agent, oxygen, and to such vigorous and common oxidizing agents as permanganates, dichromates, nitric acid, and similar substances.

«Oxidation of Hydrogen by Oxygen.»—The oxidation of hydrogen by oxygen may be first considered, as representing a typical and, in some respects, the most important case of oxidation by oxygen. Hydrogen, like copper, zinc, and other elements, has a certain tendency to form its ion, H^{+}, and the latter, in turn, has a tendency to be reduced to hydrogen. We may put H_{2} ⇄ 2 H^{+} and the condition for equilibrium, at a given temperature, will be

[H^{+}]^2 / [H_{2}] = K_{H^{+}, Hydrogen}.

In this equation [H_{2}] represents the concentration of the hydrogen in contact with the electrode (see below) and with the solution, and [H^{+}] represents the concentration of hydrogen-ion in the solution bathing the electrode. The ionization of hydrogen, at a given temperature, depends, according to this equation, on ‹two variables›, the concentration, or pressure, of the gas and the concentration, or osmotic pressure, of the hydrogen-ion in a given solution.

If platinum gauze, coated with platinum black, is charged with hydrogen, then, the greater the pressure of the gas, the more soluble the hydrogen will be in the platinum (p. 121). Such a charged gauze may be used as a hydrogen electrode (Fig. 13, p. 281), the concentration of the hydrogen in which is proportional to the concentration, or pressure, of the hydrogen gas surrounding it; the platinum will allow of the ready transmission of electric charges from and to the hydrogen dissolved in it. [p278]

The value of the constant[553] K_{H^{+}, Hydrogen} = [H^{+}]^2 / [H_{2}], at 18°, is found to be 5.55E−9, and ‹hydrogen, at 18°, under atmospheric pressure, is directly in equilibrium with hydrogen-ion of the concentration› [H^{+}] = 1.52E−5.

If such an electrode, in contact with hydrogen of atmospheric pressure, is dipped into the solution of some neutral salt, say sodium chloride, in which the concentration of the hydrogen-ion, formed by the ionization of water, at 18°, is 0.9E−7, which is less than 1.5E−5, the hydrogen in the electrode must tend to ionize more rapidly than it is formed from the ion, and the ‹electrode must receive a negative charge›, exactly as in the case of zinc, placed in a zinc sulphate solution.

For oxygen similar relations may be developed.[554] We have: O_{2} ⇄ 2 O^{2−} and

[O^{2−}]^2 / [O_{2}] = K_{1}. (1)

If we use the relation of the oxide-ion, O^{2−}, to the more stable hydroxide-ion, HO^{−}, we also have:[555] [p279]

[HO^{−}]^4 / [O_{2}] = K_{2} = K_{HO^{−}, Oxygen} (2)

The value of the constant K_{HO^{−}, Oxygen} is 8.2E49 at 18°, and ‹oxygen of atmospheric pressure at this temperature should be in equilibrium with solutions containing hydroxide-ion at a concentration[556] of› 1.36E12.

An electrode of platinum gauze, charged with oxygen under atmospheric pressure, when dipped into the solution of a neutral salt, ‹acquires a very strong positive charge›, the minute concentration of hydroxide-ion, 0.9E−7, being very much smaller than the value required by the constant, and the oxygen ionizing very much more rapidly, in consequence, than it is formed by the discharge of hydroxide ions (see p. 259). [p280]

When we combine the hydrogen and the oxygen electrodes, dipping into a solution of sodium chloride, we find a current is, in fact, established (the apparatus discussed on p. 281 is used), and it flows in the direction anticipated from the above development, the positive current entering the voltmeter from the oxygen electrode.[557]

The potential of the hydrogen electrode, for a constant pressure of hydrogen, is dependent on the concentration of hydrogen-ion in the solution surrounding the electrode, exactly as the potential of a copper plate, against a solution of cupric-ion, depends on the concentration, or osmotic pressure, of cupric-ion in the solution in which the plate is immersed. The concentration of hydrogen-ion, in the present instance, is very small (0.9E−7, at 18°), the solution being practically neutral; but the ‹addition of an alkali› must reduce its concentration far below even this value, since for water the product of the concentrations of hydrogen-ion and hydroxide-ion is a constant (p. 176) and the increase in the concentration of hydroxide-ion, produced by the addition of alkali, must decrease the concentration of the hydrogen-ion proportionally. We would expect, then, that the potential of the hydrogen electrode must ‹increase›, when we add alkali to the solution surrounding it, the hydrogen now ionizing against a much smaller concentration of its ion. Such is in fact the case (‹exp.›), and the increase is found to be subject to a logarithmic function for the relation between potential and the concentration of the ion, similar to that found to hold for copper and its ion.[558] In the same way, the potential of the oxygen electrode must depend on the concentration[559] of the hydroxide-ion in the solution bathing it. The addition of a strong acid, like sulphuric acid, to this solution, by suppressing the hydroxide-ion, small as its concentration is, should increase the potential of the [p281] electrode and the total potential of the cell. This, in fact, is the case (‹exp.›); the cell working under these conditions shows us the largest potential yet observed.[560]

The arrangement of the apparatus and the course of the current are shown in Fig. 14. The glass tube of the hydrogen electrode is connected with a hydrogen generator, the tube of the oxygen electrode with a cylinder or gasometer filled with oxygen. The hydrogen electrode is connected with the negative post of the voltmeter, the oxygen electrode with the positive post. Since the hydrogen ionizes, under the conditions used, more rapidly than it is formed from the small concentration of the hydrogen ions surrounding the hydrogen electrode, ‹hydrogen ions pass from the electrode into solution A›, leaving a negative charge on the electrode; there is a migration of sodium ions through the salt bridge (see p. 254) to solution B, and the hydrogen ions formed combine with hydroxide ions and produce water. In a similar way, ‹oxygen passes into solution B in the form of hydroxide ions› and these combine with hydrogen ions of the sulphuric acid, forming water; SO_{4}^{2−} ions migrate from the solution B through the salt bridge toward solution A and thus prevent polarization (p. 254). While water is an actual product of the action of the cell, working under these conditions, ‹the essential feature of the oxidation of hydrogen is its ionization›—H_{2} → 2 H^{+}; it would be in the same condition of oxidation if the hydrogen ions combined with any negative ions other than HO^{−}, or if they remained ions (as they would, if sodium chloride surrounded the hydrogen electrode). Similarly, the essential feature of the reduction of oxygen is its ionization in the form of HO^{−} ions; in the present [p282] instance, these actually combine with hydrogen ions and form water, but the reduction of oxygen would also be accomplished, if the hydroxide ions remained ionized (as they would, if sodium chloride bathed the oxygen electrode). The formation of water is the result of a union of ions, following the oxidation-reduction reaction, which may be expressed in the following condensed form:

2 H_{2} ⥂ 4 H^{+}

2 HOH + O_{2} ⥂ 4 HO^{−}

4 HO^{−} ⥂ 4 H_{2}O

«Summary.»—We thus find that the oxidation of hydrogen by oxygen may be used to develop an electric current, exactly in the same way as the other oxidation and reduction reactions, which we have discussed, are found available for the same purpose. And at a given temperature, the oxidation is subject to the influence of analogous factors,—the solution-tension ‹constants, the concentrations of the corresponding ions› in the solutions surrounding the electrodes, and the ‹concentrations› of the ‹gases›.

«Interpretation of Oxidation-Reduction Reactions in Terms of the Oxygen-Hydrogen Gas Cell.»—It is possible to interpret all classes of reversible oxidations and reductions, carried out in ‹aqueous› solutions, in terms of this so-called "oxygen-hydrogen gas cell," if the assumption is made that ‹each oxidizing agent›, such as nitric acid, permanganate, dichromate, etc., ‹has a tendency to liberate›, either from its own molecules or by its action on water, ‹oxygen of a definite concentration or pressure›, and that ‹each reducing agent›, in turn, ‹has a tendency to liberate hydrogen›, from water, ‹of a definite pressure or concentration›. The potential of the oxygen-hydrogen cell is dependent on the concentrations of the gases with which the electrodes are in contact:[561] therefore, each oxidation and reducing agent, yielding its own characteristic concentration of oxygen or hydrogen, respectively, would have ‹a characteristic constant›, corresponding to the solution-tension constants of the elements and measuring its oxidizing or reducing power. This interpretation of oxidation-reduction reactions has received extended attention and recognition. The study, just made, of the oxidation of hydrogen by oxygen, sufficiently suggests the treatment of oxidation and reduction from this viewpoint.

«Interpretation of Oxidation-Reduction Reactions in Terms of Direct Transfers of Electric Charges.»—In the study of the [p283] oxidation of zinc by cupric-ion, Zn ↓ + Cu^{2+} ⥂ Cu ↓ + Zn^{2+}, of the oxidation of iodides and sulphides by ferric salts, 2 Fe^{3+} + 2 I^{−} ⥂ 2 Fe^{2+} + I_{2}, and 2 Fe^{3+} + S^{2−} ⥂ 2 Fe^{2+} + S↓, and of similar actions, it has been possible to represent the oxidation-reduction actions as the result of ‹direct transfers of electric charges› between atoms and their ions.[562] In the following discussions, the attempt will be made to interpret, in the same terms, the oxidizing power of the most important remaining oxidizing agents, which include such compounds as nitric, permanganic, chromic, arsenic and similar oxygen acids. The interpretation will avoid the assumption of the liberation of oxygen and hydrogen, under hypothetical,[563] and, sometimes, enormous pressures, as ‹intermediate products› in the actions.

«Arsenic Acid as an Oxidizing Agent.»—Arsenic acid is occasionally used as an oxidizing agent (‹e.g.› in the aniline-dye industry), and, although it is not a very powerful one, its study is of theoretical interest. If to a solution of potassium arseniate some potassium iodide is added, practically no iodine is liberated (‹exp.›). If dilute hydrochloric acid, in excess, is added to this mixture, iodine is slowly liberated (‹exp.›). But the addition of concentrated hydrochloric acid causes iodine to be liberated at once in very large amounts (‹exp.›). We may ask in what way the [p284] addition of the concentrated acid causes such a decided difference in the ease and speed with which arsenic acid oxidizes the iodide-ion to free iodine and is reduced, in turn, to arsenious acid and its derivatives.[564]

We may recall the fact that a solution of potassium arseniate, to which dilute hydrochloric acid has been added, will remain clear for some time when the mixture is saturated with hydrogen sulphide (‹exp.›). If a considerable excess of concentrated hydrochloric acid is added to this mixture, hydrogen sulphide immediately forms a dense precipitate (‹exp.›) of arsenic ‹pentasulphide›—presumably through the union of quinquivalent arsenic-ion with the sulphide-ion: 2 As^{5+} + 5 S^{2−} ⇄ As_{2}S_{5} ↓ (see p. 247). This behavior suggested that arsenic acid, although a moderately strong acid, might nevertheless be ‹somewhat amphoteric›, might have ‹slight› basic properties, as well as its ordinary acid functions. The relation is expressed in the equations:[565]

3 H^{+} + AsO_{4}^{3−} ⇄ (HO)_{3}AsO

(HO)_{3}AsO + HOH ⇄ As(OH)_{5} ⥃ As^{5+} + 5 HO^{−}.

Since oxidations by arsenic acid involve its reduction to arsenious acid, containing ‹trivalent›,[566] in place of ‹quinquivalent arsenic›, one might well suspect, that the ‹oxidizing component› is the ‹quinquivalent arsenic-ion›, As^{5+}, ‹the discharge of two of whose positive charges would cause oxidation› (‹e.g.› of iodide-ion), exactly as the discharge of positive charges at the positive pole of an electric current causes oxidation (p. 252): As^{5+} + 2 I^{−} ⥂ As^{3+} + I_{2}. [p285]

In a solution of potassium arseniate, we would have only the faintest trace of the ion As^{5+}, since the addition of an alkali to the system, expressed in the above equations, would carry the reversible reactions towards the left. The addition of dilute hydrochloric acid to the system must carry the reactions towards the right and ‹increase› the concentration of As^{5+}; the addition of concentrated acid must increase the concentration of As^{5+} very much more. Even if the concentration of As^{5+} remained minute, the oxidizing power would be increased proportionally to the ‹ratio› of the concentrations in the first and the last solutions. A millionfold increase in concentration, even when we are dealing with very small numbers, would imply a millionfold increase in the activity of the solution. If, then, the ‹oxidizing component of arsenic acid› is the ‹quinquivalent ion›, As^{5+}, which would tend to discharge two of its positive (oxidizing) charges, arsenic acid should be a much more powerful oxidizing agent in strong acid solution than in alkaline or neutral solutions.

We thus arrive at the conclusion that the addition of hydrochloric acid to a mixture of arseniate and iodide may be effective, in bringing about the reduction of the arseniate and the oxidation of the iodide, ‹primarily because of its action on arsenic acid›, perhaps by facilitating its ionization as a base, and that it is not effective through any action on the iodide, for instance by producing free hydroiodic acid, as is often assumed. This conclusion may easily be tested with the aid of the chemometer (see p. 253): potassium arseniate against potassium iodide gives only the faintest possible current, barely perceptible with the aid of a very sensitive voltmeter.[567] The addition of hydrochloric acid to the beaker containing the potassium iodide does ‹not increase› the potential (it rather decreases it somewhat), whereas the addition of the concentrated acid to the potassium arseniate solution ‹produces a most decided increase in the potential›[568] (‹exp.›). It is evident, therefore, that the addition of the acid is primarily and directly ‹intended to increase the oxidizing power of the arsenic acid›, rather than to increase the reducing power of the iodide. [p286]

The more common methods of expressing oxidation-reduction reactions of this type are illustrated in the following equations:

Na_{3}AsO_{4} + 2 HI ⇄ Na_{3}AsO_{3} + H_{2}O + I_{2} (1)[569]

and AsO_{4}^{3−} + 2 H^{+} + 2 I^{−} ⇄ AsO_{3}^{3−} + H_{2}O + I_{2}. (2)

Both of these forms of expression give the net results of the action correctly. Neither attempts to interpret the interesting and important fact that the reduction of arsenic acid is facilitated by the presence of acids (of hydrogen-ion). It is, at least, also ‹permissible› to consider As^{5+} ions to be present and to express the oxidation-reduction reaction with the aid of this conception,[570] as has been done in the previous discussion. In the final analysis, this method seems to have the advantage of showing directly the changes of the valences[571] (electric charges) of the atoms involved in the oxidation-reduction, and it also expresses, clearly and definitely, the relation of the hydrogen-ion to the action.[571] The following case furnishes an illustration as to how the new point of view works out from the standpoint of a quantitative study of an oxidation-reduction reaction of this type[572]: uranyl salts, such as the sulphate UO_{2}SO_{4}, are oxidizing reagents, which are readily reduced, particularly in acid solutions, to uranous salts (‹e.g.› to the sulphate, U(SO_{4})_{2}). The potential of a mixture of uranyl and uranous salts is found[573] to depend on the action expressed in the equation UO_{2}^{2+} + 4 H^{+} + 2 ⊖ ⇄ U^{4+} + 2 H_{2}O. For the condition of equilibrium (zero potential), it follows that

[UO_{2}^{2+}] × [H^{+}]^4 / [U^{4+}] = K_{equil.} (3)

The value of this constant, at 18°, is found, by calculation,[574] to be approximately 1 / 10^{24}. Now, the uranyl-ion UO_{2}^{2+} may be assumed to have the power of ionizing, with the aid of water, to a very slight degree into ions U^{6+} and HO^{−}, according to

UO_{2}^{2+} + 2 H_{2}O ⇄ U(OH)_{4}^{2+} ⥃ U^{6+} + 4 HO^{−}. (4)

For the ionization of an extremely weak base of this character, we have, further, [U^{6+}] × [HO^{-}]^4 / [UO_{2}^{2+}] = k_{base}. And, since [HO^{−}] = K_{HOH} / [H^{+}], we also find, by substitution and by solving for U^{6+},

[U^{6+}] = [UO_{2}^{2+}] × [H^{+}]^4 × k_{base} / (K_{HOH})^4. (5)

[p287]

In other words, ‹we may substitute› [U^{6+}] ‹and a constant factor› K_{HOH}^4 / k_{base} for [UO_{2}^{2+}] × [H^{+}]^4 in the first term (numerator) of the oxidation-reduction equation (3), derived from Luther's quantitative work. We thus obtain:

[U^{6+}] / [U^{4+}] = K_{equil.} × k_{base} / K_{HOH}^4 = K, (6)

which must agree just as well with the quantitative data,[575] as does the original equilibrium equation (3). It follows, that we may write the chemical equation, for the action in acid solutions, simply U^{6+} ⇄ U^{4+}, exactly as we have Fe^{3+} ⇄ Fe^{2+} (p. 269). [U^{6+}] cannot be measured, as yet, but in the analogous case of Fe^{3+} ⇄ Fe^{2+}, where both terms of the equilibrium equation are accessible to direct measurement, the experimental evidence distinctly favors[576] the views expressed.[577]

«Permanganic Acid, Chromic Acids, etc., as Oxidizing Agents.»—We may extend the same views to the oxidizing power of such important agents as permanganic, dichromic, and nitric acids. In each case we may assume that the oxidizing component is a highly charged positive ion, ‹e.g.› in the case of KMnO_{4} a septavalent manganese-ion, Mn^{7+}, whose oxidizing power will depend on its tendency to discharge part of its heavy positive electrical charge and whose efficiency, in accordance with the law of equilibrium, will also be proportional to the concentration, in which the highly charged ion is present. Permanganate is used as a favorite oxidizing agent in the laboratory, for instance, in the oxidation of ferrous to ferric salts in quantitative analysis; the action proceeds quantitatively and rapidly in acid solution, and the end of the action is recognized by the fact that the intense pink color of the permanganate is not destroyed (‹exp.›). [p288]

If we bring permanganate, against potassium iodide, into the beakers of the chemometer (p. 253), we find that it is a much more vigorous oxidizing agent than is arsenic acid, and again we find that the addition of acid (sulphuric) to the permanganate solution enormously increases the potential (‹exp.›) and therefore its oxidizing power. The addition of an acid would, obviously, enormously increase the concentration of a positive septavalent ion, if permanganic acid is assumed to be, to a slight extent, ‹base forming› and therefore amphoteric:

H^{+} + MnO_{4}^{−} ⇄ (HO)MnO_{3}

(HO)MnO_{3} + 3 HOH ⇄ Mn(OH)_{7} ⥃ Mn^{7+} + 7 OH^{−}.

Similar experiments may be made with ferrous sulphate against permanganate.

The oxidation of ferro-ion, or of iodide-ion, may be represented, most simply, by the equations:

2 Mn^{7+} + 10 Fe^{2+} → 2 Mn^{2+} + 10 Fe^{3+}

and

2 Mn^{7+} + 10 I^{−} → 2 Mn^{2+} + 5 I_{2}.

Each heptavalent manganese ion is derived from a salt, such as MnX′_{7} or Mn_{2}Y″_{7}, and, consequently, when two manganese ions Mn^{7+} are reduced, ten univalent negative ions X′, or five bivalent ions Y″, are liberated and become available for salt formation with the ferric ions, produced, or with the hydrogen ions (from hydroiodic acid) set free by the oxidation of the iodide ions to iodine.

Thus, the oxidation of ferrous sulphate by permanganate, in the presence of sulphuric acid, may be represented, in greater detail, by the equations:

2 KMnO_{4} + H_{2}SO_{4} ⇄ K_{2}SO_{4} + 2 HMnO_{4} (1)

2 HMnO_{4} + 7 H_{2}SO_{4} ⇄ Mn_{2}(SO_{4})_{7} + 8 H_{2}O (2)

Mn_{2}(SO_{4})_{7} ⇄ 2 Mn^{7+} + 7 SO_{4}^{2−} (3)

2 Mn^{7+} + 7 SO_{4}^{2−} + 10 Fe^{2+} + 10 SO_{4}^{2−} ⇄ 2 Mn^{2+} + 1 OFe^{3+} + 17 SO_{4}^{2−} (4)

⇄ 2 MnSO_{4} + 5 Fe_{2}(SO_{4})_{3} (5)

Analogous results are obtained with potassium chromate or dichromate against potassium iodide, ferrous sulphate, hydrogen sulphide, and other reducing agents.

«Nitric Acid.»—It is characteristic of nitric acid that, ionized as an acid, it is not a powerful oxidizing agent; that is, the nitrate-ion, NO_{3}^{−}, is not the oxidizing component. For instance, a nitrate, such as potassium nitrate, in aqueous solution, does not appreciably oxidize ferro-ion or iodide-ion (‹exp.›). Concentrated nitric acid, or a mixture of a large excess of concentrated [p289] sulphuric acid with a nitrate, are far more effective in oxidizing the substances mentioned, as shown by experiments with the chemometer and with the mixtures described. It is significant that, in concentrated nitric acid, and in the presence of concentrated sulphuric acid, ‹the basic ionization of nitric acid›

H^{+} + NO_{3}^{−} ⇄ (HO)NO_{2}

(HO)NO_{2} + 2 HOH ⇄ N(OH)_{5} ⥃ N^{5+} + 5 HO^{−}

‹would be facilitated› by the high concentration of hydrogen-ion.

«Summary.»—In every instance, then, we find that it is possible to produce an electric current by using one of these common, powerful oxidizing agents in aqueous solution, just as it is possible with oxygen and with the simpler agents discussed earlier. Whatever the theory of the formation of the current, whether we are dealing, in the last cases considered, with acids which have the capacity to ionize minimally as bases, forming highly charged positive ions, whose discharge involves an oxidation of the substance receiving the discharged electricity, or whether we accept the view that oxygen is liberated by them in exceedingly concentrated form and tends to form in aqueous solutions the hydroxide-ion, HO^{−}, unquestionably the closest possible relation has been established between oxidation-reduction reactions and electrical relations, as formulated on the basis of the theory of ionization. In all cases, at a constant temperature, as demanded by the law of mass action, the net result of an action will depend on ‹constant factors›, measuring the tendencies of atoms to assume or lose electric charges (electrons), and on ‹variable factors›, the concentrations of the reacting components.

«Oxidation of Organic Compounds.»—In conclusion, it may be asked whether ‹the oxidation of organic compounds› may also be brought into relation with electrical charges and interpreted on the basis of the theory which has been presented. For this purpose we may study the oxidation of one of the simplest of organic compounds, formaldehyde CH_{2}O, which forms the active component in solutions of ‹formalin›. Formaldehyde is also intimately related to some of the most important food products, the carbohydrates, whose oxidation is utilized in the animal economy. It is formed[578] in the green leaves of plants by the [p290] reduction of carbon dioxide, absorbed from the atmosphere, with the aid of the energy of the sun's light: CO_{2} + H_{2}O → CH_{2}O + O_{2}. It is readily condensed, forming glucose (6 CH_{2}O → C_{6}H_{12}O_{6}), cane sugar C_{12}H_{22}O_{11}, and still more complex carbohydrates, such as starch.

Formaldehyde, like other aldehydes, is readily oxidized. A favorite reagent, used in oxidizing it, is an ammoniacal solution of silver nitrate (‹exp.›), the separation of silver from such a solution being a characteristic reaction of aldehydes. The reagent is rendered still more sensitive by the addition of sodium or potassium hydroxide.[579] We may ask how we would interpret, from the point of view of the electric theory of oxidation and reduction, the oxidation of formaldehyde and the reduction of silver nitrate to silver, under these conditions. According to the theory, the oxidizing agent in silver nitrate is the silver-ion, the discharge of which gives positive electricity, which the oxidized substance, the formaldehyde, must absorb. But in a silver nitrate solution there is a far larger concentration of the silver-ion than in an ammoniacal solution (p. 220), containing the same total concentration of silver. The complex silver-ammonium-ion Ag(NH_{3})_{2}^{+}, it may be recalled, is a rather stable one,[580] and, consequently, the addition of ammonia to silver nitrate should decidedly ‹weaken› its oxidizing power. Still, the practical use of ammonia, especially in combination with sodium hydroxide, is found to be most effective. We are led to suspect that, in spite of the untoward effect of ammonia on the oxidizing power of the silver compound, an ‹alkaline› solution is desirable for the sake of the effect of the alkali on formaldehyde, the reducing substance involved. To follow up this conclusion, we must next consider, in some detail, the nature of formaldehyde; we shall presently find that the conclusion, which we have just reached, as to the probably favorable effect of alkali on the ‹reducing power› of formaldehyde, will be verified by experiments, which the consideration of formaldehyde will suggest.

The oxidation of formaldehyde may most clearly be formulated on the basis of views, developed by Nef, on the formation of methylene[581] derivatives, containing bivalent carbon atoms. A solution [p291] of formalin contains formaldehyde in a variety of forms, in a very complex condition of equilibrium. Of these compounds, the aldehyde, CH_{2}O, probably exists in two forms, which have the same composition and molecular weight, but which differ in the arrangement of the atoms in the molecules (in the ‹structure› of the molecules); we probably have CH_{2}=O ⥃ CH(OH), the former of which (CH_{2}=O) is, most likely, by far the more stable and the chief one of these two substances, present under ordinary conditions. The second compound CH(OH) may be present in traces only. One difference, we note, lies in the position of one of the hydrogen atoms in the respective molecules; the second form contains a hydroxide group (OH), which gives it the properties of an acid and renders it capable of forming salts CH(OMe) with bases. But the molecule of this second form also would ‹contain a carbon atom, only two of whose valences are satisfied› (by H and OH), two of the ordinary four valences of a carbon atom being thus left free or ‹unsaturated›. We may indicate the two free carbon valences in the formula =CH(OH). ‹Such an unsaturated, bivalent carbon atom› =C ‹would be particularly sensitive to oxidation›.[582]

Besides these two forms, a formalin solution also contains a polymerized form, probably (CH_{2}O)_{2}, which in dilute solution, or under the influence of heat, slowly breaks down into formaldehyde, (CH_{2}O)_{2} ⇄ 2 CH_{2}O.

The addition of alkali to the mixture probably leads to the formation of the salt =CH(OMe), thus disturbing all the conditions of equilibrium and ‹leading to the transformation of a very much larger part of the aldehyde into a compound containing the characteristic unsaturated (bivalent) carbon›, than was originally present. The aldehyde will thus become ‹more susceptible to oxidation› as a result of the enormous increase in the concentration of the oxidizable component. We may assume this to be either the salt, =CH(OMe), or its negative ion, =CH(O^{−}), or both, or some analogous derivative. Further, the two free valences of a bivalent carbon atom may be considered to consist of a positive and a negative charge of electricity, either actual or potential,[583] and ‹the oxidation› will consist [p292] ‹primarily in the absorption of two positive charges, from the oxidizing agent, to convert the negative charge on the carbon atom›, say in ±CH(ONa), ‹into a positive charge›.[584] If the oxidizing agent is alkaline silver nitrate solution, we may formulate the successive actions as follows:

(NaO)HC± + 2 Ag^{+} → (NaO)HC^{2+} + 2 Ag ↓.

The two positive silver ions correspond to two negative ions, ‹e.g.› hydroxide ions HO^{−}, which are set free by the discharge of the silver ions, and which, in turn, will combine with the oxidized carbon atom holding the two positive charges:

(NaO)HC^{2+} + 2 HO^{−} → (NaO)HC(OH)_{2} → (NaO)HC:O + H_{2}O.

The salt formed, HCO_{2}Na, is ‹sodium formate›, which is the first isolated product of the oxidation of formaldehyde.

It would appear, from this point of view, that the ‹alkaline nature› of the silver nitrate mixture is advantageous primarily because a base is required by the formaldehyde, the reducing agent, to convert it into some readily oxidizable form. And the proved efficiency of the alkaline mixture (see above) makes it appear probable that the advantage gained by this result ‹more than offsets the loss in oxidizing power›, suffered by the silver nitrate following the suppression of its real oxidizing component, the silver-ion, when, in the presence of ammonia, the latter is converted largely into the ion, Ag(NH_{3})_{2}^{+}. Ammonia, in turn, is employed in the oxidizing mixture, essentially with the object of preventing the precipitation of the silver-ion, as silver oxide, by the hydroxide-ion of an alkaline mixture. These conclusions, as well as, in particular, the ‹main conception› «that in the oxidation of formaldehyde there is an actual transfer of electrical charges», may be fully confirmed with the aid of the chemometer.[585]

EXP. A small beaker, containing a platinum electrode, which is connected with the positive post of the voltmeter, is half filled with a solution of silver and sodium nitrates. A similar small beaker, containing a platinum electrode leading to the negative post of the voltmeter, is charged with a solution of sodium nitrate (to render the solution a good conductor) and with some formalin. The solutions in the two beakers are connected by means of a salt-bridge containing sodium nitrate. [p293]

Only a ‹very slight› current is produced under these conditions; the potential between silver nitrate and formaldehyde is found to be ‹extremely› small. If, now, sodium hydroxide is added to the formalin mixture, an ‹enormous increase› in potential is observed, proving, unmistakably, that the addition of the alkali to the formalin solution ‹enormously increases the concentration of the reacting, oxidizable component›.[586]

When some ammonia is added to the silver nitrate mixture, we find, as anticipated, that the ‹oxidizing power of the silver solution is greatly reduced›, the silver-ion being converted into the complex ion, Ag(NH_{3})_{2}^{+}; but the potential is still very much ‹greater› than the potential between silver nitrate and formalin without any alkali—which shows that the advantage of using alkali with the formaldehyde greatly outweighs the disadvantage of using ammonia with the silver nitrate.

An electric current may also be readily obtained by combining alkaline formaldehyde with other oxidizing agents—for instance with an ‹oxygen› electrode (p. 279). We find (‹exp.›) that the oxidation proceeds with remarkable ease under these conditions. Permanganate, dichromate, etc., may be substituted for oxygen, with the same general result.

«Summary.»—It is clear, then, that the oxidation of an ‹organic substance› may readily be interpreted as consisting, ultimately, in a transfer of electrical charges, of exactly the same nature, as is found in the other oxidation reactions which we have considered, [p294] and that the conditions for producing a maximum current, as investigated with the aid of the chemometer, give us, again, important guidance in following the course of the reactions. It is needless to say that the oxidation of other organic compounds, such as glucose, alcohol, etc., may be profitably studied from the same point of view.

It also follows, from the conclusions reached, that, under proper experimental conditions, electricity, in the form of a current, must be capable of effecting the oxidation, or the reduction, of organic as well as inorganic compounds (p. 252). Extended investigations have, indeed, shown that electric currents belong to the most important and efficient agents for this purpose, because the oxidation, or the reduction, of the organic compound becomes susceptible to the most exact control through the regulation of the potentials used.[587]

«Tables and Summaries.»—In the first table the equilibrium (solution-tension) constants of a number of metals and non-metallic elements are given. The table is followed by brief explanations of its meaning and a summary of some of its more important applications.

TABLE[A] OF EQUILIBRIUM (SOLUTION-TENSION) CONSTANTS (IN MOLAR TERMS) AND OF POTENTIAL DIFFERENCES BETWEEN ELEMENTS AND THEIR IONS IN UNIMOLAR AQUEOUS SOLUTIONS.

Element, Ion. E.P._{El.,Ion.} K_{Ion.}

K, K^{+}[B] (−2.92) 6E50 Na, Na^{+}[C] −2.44 2.5E42 Ba, Ba^{2+} (−2.54) 2.1E88 Sr, Sr^{2+} (−2.49) 4.0E86 Ca, Ca^{2+} (−2.28) 2.0E79 Mg, Mg^{2+} (−2.26) 4.1E78 Al, Al^{3+}[D] −0.999 ? 1.3E52 Mn, Mn^{2+} −0.798 5.7E27 Zn, Zn^{2+} −0.493 1.4E17 Cd, Cd^{2+} −0.143 9.5E4 Fe, Fe^{2+}[E] −0.122 ? 1.8E4 Co, Co^{2+}[F] +0.0138 ? 0.3314 Ni, Ni^{2+}[G] +0.108 ? 1.8E−4 Sn, Sn^{2+} <+0.085 <1.1E−3 Pb, Pb^{2+} +0.129 3.3E−5 H_{2}, H^{+}[H] +0.277 1.52E−5 Cu, Cu^{2+} +0.606 8.3E−22 As, As^{+++} <+0.570 <2.7E−30 Bi, Bi^{3+} <+0.668 <1.4E−35 Sb, Sb^{3+} <+0.743 <1.7E−39 Hg, Hg^{+} +1.027 1.38E−18 Ag, Ag^{+} +1.048 6E−19 Pt, Pt^{4+} <+1.140 5E−80 Au, Au^{3+} <+1.356 <1.8E−71 F_{3}, F^{−}[H] (+2.24) 9.0E88 Cl_{2}, Cl^{−}[H] +1.694 3.16E29 Br_{2}, Br^{−} +1.270 1.23E22 I_{2}, I^{−} +0.797 ? 7.26E13 O_{2}, HO^{−}[I] +0.698 1.36E12

TABLE NOTES:

A. The table is based on Wilsmore's compilation of solution-tension potentials, ‹Z. phys. Chem.›, «36», 91 (1901).

B. Values in parentheses have been estimated by indirect measurements.

C. G. N. Lewis, ‹J. Am. Chem. Soc.›, «32», 1467 (1910).

D. Values marked with? are uncertain.

E. Calculated from the data of Richards and Behr (‹Z. phys. Chem.›, «58», 301 (1907)), who found the potential of iron against 0.5 molar FeSO_{4} to be −0.15 volt. The degree of ionization of 0.5 molar FeSO_{4} is taken as 22%. [Λ = 25.8 (Kohlrausch and Holborn, ‹loc. cit.›, p. 152) and Λ_{∞} is taken as 117, as for ZnSO_{4} (‹ibid.›, p. 200).] On account of the doubtful value for the degree of ionization, the values in the table are marked?, but the value found by Richards and Behr appears to be quite accurate.

F. Calculated from the data of Schildbach (‹Z. für Elektroch.›, «16», 967 (1910)). The same uncertainty as to the degree of ionization exists as that discussed in the previous footnote.

G. Calculated from the data of E. P. Schoch (‹Am. Chem. J.›, «41», 208 (1909)). The same uncertainty as to the degree of ionization exists as that discussed in footnote 5, p. 294.

H. The values for ‹gaseous› elements refer to the gases under one atmosphere pressure.

I. The potential of oxygen at 18°, 760 mm., against an alkaline solution in which [HO^{−}] = 1. K_{Ion} refers to the concentration of HO^{−}, with which oxygen under atmospheric pressure would be directly in equilibrium, at 18°.

The potential differences, given in the table, are based on the assumption that the ‹absolute zero› of potential is at such a point, that the so-called standard normal calomel electrode has a value of +0.56 volt relative to this zero (‹cf.› Ostwald, ‹Z. phys. Chem.›, «36», 97 (1901)). The exact determination of this value is a very difficult matter. Recently Palmaer (‹ibid.›, «59», 129 (1907)), located the absolute zero at a point 0.04 volt more positive than the above, making the absolute potential of the normal calomel electrode, approximately, +0.52 volt. To refer potentials, given in this book, to this new zero, one would subtract 0.04 volt from all positive potentials and add 0.04 to the numbers representing negative potentials (‹e.g.› E. P._{Zn,Zn^{2+}} would become −0.569 in place of −0.529 volt). Since the ‹equilibrium (solution-tension) constants› are calculated from the potential differences referred to the absolute zero (p. 259), any change in the zero involves corresponding changes in the values of the equilibrium constants, as calculated for this book. However, it should be noted that all potential differences would be corrected by the same constant quantity (0.04 volt for Palmaer's zero): ‹all the equilibrium constants for univalent metallic ions would be increased proportionally to a constant factor› ‹c› (‹c› is very nearly equal to 5, for Palmaer's zero), the equilibrium constants for bivalent metallic ions would be increased proportionally to ‹c›^2, etc. ‹The equilibrium ratio› for two metals and their ions ‹would in no wise be changed› by these alterations: ‹e.g.› for the equilibrium between zinc and copper and [p296] their ions (p. 267), K_{equil.} = K_{Zn^{2+}} / K_{Cu^{2+}}; the factor ‹c›^2 would be introduced into both terms of the ratio and would not affect the value of the latter. For the condition of equilibrium between silver and copper and their ions (p. 267) K_{equil.} = K_{Ag^{+}}^2 / K_{Cu^{2+}}, and since (‹c›)^2 = ‹c›^2, this equilibrium ratio would also not be affected. For elements, which produce negative ions, the corresponding correction factors would be 1 / ‹c›, 1 / ‹c›^2, etc., and the equilibrium relations between two such elements and their ions likewise would remain unchanged. Since these ‹equilibrium relations› are the ‹significant› ones in this work, and since our conclusions have been based on them, it is clear that a change in the absolute zero would not affect the conclusions reached.

On account of the uncertainty attaching to the determination of the absolute zero of potential, it is preferred, in practice, to report the experimentally determined potentials as measured against a constant, well-defined electrode (such as the calomel electrode or a hydrogen electrode) and thus to eliminate the variation, which a change in the determination of the zero potential would make necessary. However, for an elementary discussion of oxidation-reduction reactions, from the same viewpoint as is used in considering all other reversible chemical actions, the idea of the absolute potential has certain advantages, making a uniform treatment possible.

1. ‹Meaning of K_{Ion}.› Under K_{Ion} is given, for each element, the ‹concentration of its ion›, with which the element would be directly in equilibrium at the ordinary temperature (see p. 258). The constants for ‹gaseous› elements represent the constants of the gases under atmospheric pressure.

2. ‹The Condition for Equilibrium between Two Elements and Their Ions.› The condition of equilibrium in a system of two elements and their ions may be found with the aid of the constants K_{Ion}, as follows: For Zn ↓ + Cu^{2+} ⥂ Zn^{2+} + Cu ↓ we have for the condition of equilibrium (see p. 267)

[Zn^{2+}] / [Cu^{2+}] = K and K = K_{Zn^{2+}} / K_{Cu^{2+}} = 1.4E17 / 8.3E−22 = 1.7E38.

Zinc-ion must be present in enormous excess in the condition of equilibrium and zinc will precipitate copper from solutions of cupric salts until this relation is established. The ‹suppression of the cupric-ion›—by precipitation in the form of insoluble salts or by conversion into very stable complex ions—makes [Cu^{2+}] exceedingly small and makes it increasingly difficult for zinc to precipitate copper, and, under certain conditions, the ‹ordinary› course of the action may be ‹reversed› (p. 268).

For Cu ↓ + 2 Ag^{+} ⥂ Cu^{2+} + 2 Ag ↓, we have (p. 267)

[Cu^{2+}] / [Ag^{+}]^2 = K and

K = K_{Cu^{2+}} / K_{Ag^{+}}^2 = 8.3E−22 / (6E−19)^2 = 2.3E15.

3. ‹Potential Differences Calculated with the Aid of K_{Ion}.› For ‹metallic elements›, which send out ‹positive ions›, in contact with an aqueous solution containing the ion in concentration [C], the potential difference is (see p. 261)

ε_{‹Me, Ion›^{‹v›}} = (0.0575 / ‹v›) ‹log›([C] / K_{Ion}) ‹volts›,

[p297]

and for elements which form negative ions (see footnote 1, p. 261),

ε_{‹Elem., Ion›^{‹v›}} = (−0.0575 / ‹v›) ‹log›([C] / K_{Ion}) ‹volts›.

In these equations ‹v› represents the valence of the ion. It is clear that for the condition of equilibrium, in which [C] = K_{Ion}, the potential is 0. Further, for the potential difference between copper and a cupric salt solution in which [Cu^{2+}] = 1, we would have

ε_{Cu, Cu^{2+}} = (0.0575 / 2) log(1 / 8.3E−22) = 21.08 × 0.0575 / 2 = +0.606 volts.

4. ‹Meaning of E.P._{Element, Ion}.› Under E.P._{Element, Ion} the table gives the ‹potential difference in volts›, calculated for the ‹element› named and ‹an aqueous solution of its ion in unit concentration› (one gram-ion per liter). For instance, for zinc and [Zn^{2+}] = 1 (65.4 grams zinc-ion per liter), we have a potential E.P._{Zn, Zn^{2+}} = −0.493. The signs used, in accordance with the convention adopted (p. 261), indicate the ‹character of the charge› on the ‹element electrode› (which is named ‹first› in the subscript to E.P.). For instance, zinc in a solution in which [Zn^{2+}] = 1 would acquire a ‹negative› charge (p. 266), the potential difference E.P._{Zn, Zn^{2+}} being −0.493 according to the table; silver, immersed in a solution in which [Ag^{+}] = 1, would acquire a ‹positive charge›, the potential difference E.P._{Ag, Ag^{+}} = +1.048.

The potentials given for the ‹gaseous› elements represent the potentials of the gases under 760 mm. pressure.

5. ‹Potential Differences Calculated with the Aid of E.P._{Element, Ion}.› The potential corresponding to any concentration [C] of a ‹metal ion› may be found from the equation[1]

ε_{El., Ion^{‹v›}} = E.P._{El., Ion} + (0.0575 / ‹v›) log[C] ‹volts›,

and the potential for any concentration [C] of the ions of elements forming ‹negative ions› is found[588] according to

ε_{El., Ion^{‹v›}} = E.P._{El., Ion} − (0.0575 / ‹v›) log[C] ‹volts›.

6. ‹The Condition for Equilibrium between Two Metals and Their Ions, Calculated with the Aid of E.P._{Element, Ion}.› The condition for equilibrium in a system of two metals and their ions is determined by the fact that the potential of the system must be 0 when equilibrium is established. We have, for instance for the two metals zinc and copper and their ions, Zn^{2+} and Cu^{2+}, for Zn ↓ + Cu^{2+} ⇄ Cu ↓ + Zn^{2+} the condition for equilibrium that ε_{Cu, Cu^{2+}} − ε_{Zn, Zn^{2+}} = 0. According to the equation given in § 5, we have, then, for the condition of equilibrium,

E.P._{Cu, Cu^{2+}} + (0.0575 / 2) log[Cu^{2+}] − E.P._{Zn, Zn^{2+}} + (0.0575 / 2) log[Zn^{2+}] = 0.

[p298]

Then

(0.0575 / 2) log([Zn^{2+}] / [Cu^{2+}]) = E.P._{Cu,Cu^{2+}} − E.P._{Zn,Zn^{2+}} = +0.606 − (−0.493) = 1.099.

From the last relation we find log([Zn^{2+}] / [Cu^{2+}]) = 38.2261, and therefore, for the condition of equilibrium, [Zn^{2+}] / [Cu^{2+}] = 1.7E38.

7. ‹Equilibrium Constants for Elements with Variable Concentration.› The concentration of a pure metal at a given temperature may be considered a constant, except in the case of extremely thin films of the metal (p. 258). The concentration of hydrogen, and of the non-metallic elements given in the table, is variable, and K_{Ion} has a definite value only when the concentration of the element is defined (see the preceding table, footnotes 3, 4, p. 295). For certain estimations the ‹equilibrium constants›, which show the relation between the two variables, namely the concentration of the element and that of its ion, are very helpful (see pp. 274 and 275). In the following table some of the more important equilibrium constants of this nature are given.

TABLE OF EQUILIBRIUM CONSTANTS.

Element. K_{equil.}.

Hydrogen: [H^{+}]^2 : [H_{2}] 5.6E−9 Oxygen: [HO^{−}]^4 : [O_{2}] 8.2E49 Chlorine: [Cl^{−}]^2 : [Cl_{2}] 2E60 Iodine[A]: [I^{−}]^2 : [I_{2}] 5.6E29

TABLE NOTE:

A. The value of the constant, as given, is only an approximate estimation (p. 273).

The significance of the constants is indicated by the ratios given in the table. The relation of these constants to those given in the first table may be seen from the following illustration. For hydrogen we have H_{2} ⇄ 2 H^{+}. The first table tells us that hydrogen, at 18° under atmospheric pressure, is in equilibrium with its ion when the concentration of hydrogen-ion is 1.52E−5 (under K_{Ion}). Now, a mole of hydrogen at 18° occupies 22.4 × 291 / 273 = 23.9 liters under atmospheric pressure, and its concentration (per liter) is therefore 1 / 23.9 mole. Then equilibrium exists, when [H_{2}] = 1 / 23.9 and [H^{+}] = 1.52E−5 and K_{equil.} = [H^{+}]^2 : [H_{2}] = (1.52E−5)^2 × 23.9 = 5.6E−9.

FOOTNOTES:

[553] These constants are calculated from data given in Wilsmore's tables (‹loc. cit.›) on the solution tension of hydrogen. Hydrogen, at 18° under one atmosphere pressure, produces a potential of ε_{H_{2}, H^{+}} = +0.277 (see p. 261, in regard to the sign) against a solution containing hydrogen-ion in a concentration [H^{+}] = 1 (see the table at the end of this chapter). Now, there must be some concentration of hydrogen-ion, which we will call [C], with which hydrogen at 18° and 760 mm. pressure is directly in equilibrium, with the potential 0. For any concentration of hydrogen-ion [H^{+}], ‹other than› [C], a potential is produced according to ε_{H_{2}, H^{+}} = 0.0575 log([H^{+}] / [C]). If we insert into this equation the values [H^{+}] = 1 and the potential ε = +0.277, and if we solve the equation for [C], we find [C] = 1.52E−5. That is the concentration of H^{+}, with which hydrogen of one atmosphere pressure at 18° is directly in equilibrium. Since under these conditions of temperature and pressure [H_{2}] = 1 / 23.9 mole, we have for the condition of ‹equilibrium› [H^{+}]^2 / [H_{2}] = K : (1.52E−5)^2 : (1 / 23.9) = 5.55E−9 = K.

[554] Experimentally the relations for an "oxygen electrode" are much more complicated than for a hydrogen electrode, as a result, apparently, of the oxidation of the metal (‹e.g.› platinum), with the aid of which the electrode is prepared. For a critical review and summary of the more recent results on this point, ‹vide› Schoch, ‹J. phys. Chem.›, «14», 665 (1910). For the purposes of this book it will be sufficient to limit our discussion to the behavior of an ideal oxygen electrode.

[555] The bivalent oxygen ions, O^{2−}, combine with hydrogen ions (formed, for instance, by the ionization of water) and form the more stable hydroxide ions (p. 246): O^{2−} + H^{+} + HO^{−} ⇄ 2 HO^{−}, or simply, O^{2−} + H^{+} ⇄ HO^{−}. Then, [O^{2−}] × [H^{+}] / [HO^{−}] = k and [O^{2−}] = k × [HO^{−}] / [H^{+}]. But since we have [H^{+}] × [HO^{−}] = K_{HOH} for the ionization of water (p. 176), we also have:

[H^{+}] = K_{HOH} / [HO^{−}] and [O^{2−}] = (k / K_{HOH}) × [HO^{−}]^2.

By substituting this value for the concentration [O^{2−}] of the oxide-ion in equation (1), equation (2) is obtained. The constant K_{2} includes then the constants k and K_{HOH}.

[556] The constants are calculated from the estimated potential of the oxygen-hydrogen cell, +1.231 volt, at 18°. (‹Vide› G. N. Lewis, ‹Z. phys. Chem.›, «55», 465 (1906); Nernst and Wartenberg, ‹ibid.›, «56», 534 (1906); Brönsted, ‹ibid.›, «65», 91 (1908); and a summary and discussion by Schoch, ‹loc. cit.›) At 18° oxygen, under one atmosphere pressure, gives an estimated potential ε_{O_{2}, HO^{−}} = +1.508 against an ‹acid› solution, in which the concentration of the ‹hydrogen-ion› [H^{+}] = 1 (see the table at the end of this chapter). Since at 18° [H^{+}] × [HO^{−}] = 0.81E−14, the value for [HO^{−}] in this acid solution is 0.81E−14. Now, for oxygen, at 18° and 760 mm., there must be some concentration of hydroxide-ion, which we will call [C], at which the tendency of the oxygen to ionize is exactly balanced by the tendency of the hydroxide-ion to form oxygen—at this point the potential is 0. For any concentration [HO^{−}] of the hydroxide-ion, ‹other than› [C], a potential will exist ε_{O_{2}, HO^{−}} = 0.0575 log([C] / [HO^{−}]). Since for [HO^{−}] = 0.81E−14, we have a potential ε_{O_{2}, HO^{−}} = +1.508, these values can be introduced into the equation and the latter solved for [C]. We find thus [C] = 1.36E12, and ‹oxygen, at 18° and› 760 mm. ‹pressure, would be directly in equilibrium with a solution in which› [HO^{−}] = 1.36E12. At 18° and 760 mm. pressure a liter of oxygen contains 1 / 23.9 mole, and thus we have for the condition of equilibrium [HO^{−}]^4 : [O_{2}] = K : (1.36E12)^4 : (1 / 23.9) = 8.2E49 = K.

[557] The most convenient form of electrode for this purpose consists (see Fig. 14) of a cylinder (about one inch long) of platinum gauze, which is fused to a glass tube and connected with a wire leading through the tube to some mercury, held in a small branch tube, fused into the main tube near its upper end. The gas is easily conducted to the platinum gauze electrode through such a tube. The cylinder of platinum gauze may be made by joining the ends of rolled gauze with pieces of molten glass. It is coated with platinum black.

[558] ‹Cf.› footnote 1, p. 278.

[559] ‹Cf.› footnote 1, p. 279.

[560] See Ostwald, ‹Z. phys. Chem.›, «11», 521 (1893), Arrhenius, ‹ibid.›, «11», 805 (1893), and Nernst, ‹ibid.›, «14», 155 (1893), for a detailed discussion of oxygen-hydrogen gas cells. For more recent work, ‹vide› G. N. Lewis, ‹J. Am. Chem. Soc.›, «28», 158 (1905), where references to other recent investigations are given.

[561] ‹Cf.› footnote 1, p. 261, and pp. 277–279.

[562] See pp. 42, 252, for the expression of the changes as ‹transfers› of ‹electrons›.

[563] Fredenhagen, ‹Z. anorg. Chem.›, «29», 424 (1902), has brought interesting experimental evidence of the charging of an electrode with ‹gaseous oxygen›, when ferric-ion is the oxidizing agent in aqueous solutions. Whether the oxygen, which is liberated by the action of the ferric-ion on water, 4 Fe^{3+} + 4 HO^{−} ⇄ 4 Fe^{2+} + O_{2} + 2 H_{2}O, is always ‹the intermediate product and the direct oxidizing agent› in aqueous solution, can hardly be considered decided by the experiment—it may well be the product of a ‹parallel action›, which must take place to a certain extent, according to the laws of equilibrium, in a system containing both Fe^{3+} and HO^{−} ions. The result hardly proves that oxygen must be the intermediate product in the ‹main action›, when ferric ions act as the oxidizing agent. We may consider, for instance, a solution containing an iodide and a ferric salt: iodide ions have a far smaller affinity for their negative charges (electrons) than hydroxide ions have, and, consequently, will transfer their negative charges (electrons) more readily to the ferric ions than the hydroxide ions would. The action, if oxygen were ‹first› liberated, would lead to the same ultimate result, but the observation made by Fredenhagen would not prove that the main action would not nevertheless go by the shorter direct path rather than through an intermediate formation of oxygen.

[564] According to the theory, that arsenic acid is an oxidizing agent because it gives up oxygen of a definite pressure, ‹this pressure would be the more effective, the more completely the opposing hydroxide-ion is suppressed› by added acid (p. 280; see also p. 272, on the action of ferro-ion on silver-ion in the presence and in the absence of fluorides).

[565] Only the simplest form of basic ionization of arsenic acid is considered. Intermediate ionization into As(OH)_{4}^{−}, As(OH)_{3}^{2−}, etc. (see p. 249), is, of course, to be assumed in any complete investigation of the subject.

[566] Arsenious acid As(OH)_{3}, or HAsO_{2}, as well as its anions AsO_{3}^{3−} and AsO_{2}^{−}, may have their own characteristic tendencies to assume positive charges and be oxidized to arsenic acid and its derivatives (see footnote 1, p. 270). In ‹alkaline› solutions these tendencies, and the potentials corresponding to them, might well be more important factors in determining the course of an action, than the tendency of As^{3+} to form As^{5+}. The discussion in the text, which deals primarily with acid solutions, does not exclude such relations.

[567] A millivoltmeter is used for this experiment.

[568] See footnote 1, p. 284, in regard to the interpretation of this experiment on the basis of the theory of liberation of oxygen by arsenic acid.

[569] ‹Cf.› Smith's ‹General Inorganic Chemistry›, p. 712.

[570] The essential feature of this point of view was first published by Abegg, ‹Z. anorg. Chem.›, «39», 330 (1904), and ‹Z. phys. Chem.›, «43», 385 (1903); ‹vide› also Stieglitz, ‹Am. Chem. J.›, «39», 51 (footnote) (1908), and ‹Qualitative Analysis Notes›, University of Chicago (1905). Abegg's view has recently received support from J. J. Thomson in his ‹Corpuscular Theory of Matter›, p. 118.

[571] H^{+} does not change its valence (charge) in the action, and yet it appears as an essential component in both styles of the current equations for the oxidation-reduction reaction.

[572] A somewhat similar development of these relations, for arsenic acid, has been found in Abegg's ‹Anorg. Chem.›, III, 3, p. 552 (1907).

[573] Luther and Michie, ‹Z. für Elektroch.›, «14», 826 (1908).

[574] The calculation is based on the results obtained by Luther and Michie.

[575] The proportion of UO_{2}^{2+} converted into U^{6+} would be so minute, that in an experimental determination of the concentration [UO_{2}^{2+}], the concentration [U^{6+}] would no doubt be a negligible quantity.

[576] Peters, ‹Z. phys. Chem.›, «26», 193 (1898).

[577] In other instances, the action of the hydrogen-ion in facilitating and accelerating chemical actions (often called its "catalytic effect") has been explained, in a similar fashion, as being based on ‹salt formation, followed by the ionization of the salts formed›, the ‹active› components being the ions (‹vide›, for instance, Bredig, ‹Z. für Elektroch.›, «9», 118 and «10», 586 (1904) and Stieglitz, ‹Report of the Congress of Arts and Sciences›, St. Louis, «4», 276 (1904) and ‹Am. Chem. J.›, «39», 29, 415 (1908) and later articles). In many of these cases, the ion concentrations of the reacting components have not yet been accessible to direct measurements, but the viewpoint has been sustained by quantitative studies of analogous reactions, which were selected for study, because the factors involved could be measured (‹cf.› Stieglitz, ‹loc. cit.›).

[578] Usher and Priestley, ‹Proc. Roy. Soc.›, B, «77», 369 (1905); «78», 318 (1908).

[579] Tollens, ‹Ber. d. chem. Ges.›, «15», 1635 (1882).

[580] [Ag^{+}] × [NH_{3}]^2 / [Ag(NH_{3})_{2}^{+}] = 1 / 10^7.

[581] Methylene CH_{2}, itself, has never been isolated, but derivatives of it are known, such as the cyanides, C(NH), C(NK) (see pp. 66, 237).

[582] Potassium cyanide, =C(NK), is a powerful reducing agent (see p. 89).

[583] Pp. 65, 66.

[584] If a negative charge is an electron, a positive charge the absence of an electron in an atom, the bivalent carbon atom ‹loses two electrons›, when it is oxidized.

[585] See Stieglitz, ‹Science›, «27», 774 (1908).

[586] The oxidation occurs essentially in the same manner as described (p, 292) for the action of formaldehyde on ammoniacal silver solution, when they are brought together in a single vessel. In the present case, where the action is used to produce an electric current, there is a ‹migration› of ‹negative ions› into the formalin solution through the salt-bridge (p. 254). For every two silver ions discharged on the electrode in the silver nitrate solution, two hydroxide ions are liberated in the formaldehyde solution, as a result of this migration, and they combine with the oxidized carbon atom. The oxidation may be expressed, then, simply as follows:

(NaO)HC± + 2 ⊕ + 2 HO^{−} → (NaO)HC^{2+} + 2 HO^{−} → (NaO)HC:O + H_{2}O.

By comparison with the equations given on p. 292, it is evident, that the only difference lies in the fact that the ‹positive charges›, in the present case, are carried to the formaldehyde salt through ‹metal wires› and a ‹metal electrode›, while previously they were discharged ‹directly by silver ions on the formaldehyde salt›.

[587] ‹Cf.› Nernst, ‹Theoretical Chemistry› (1904), p. 739, and the applications mentioned there.

[588] Le Blanc, ‹Elektrochemie,› p. 215; ‹v› is the valence of the ion.

[p299]