Oxidation and Reduction - The Journal of Physical Chemistry (ACS

OXIDATION AND REDUCTION‘ BY HAMILTON PERKINS CADY AND ROBERT TAFT

I. Statement of the Problem Apparently the first introduction into chemical literature of the conception of oxidation and reduction as a change of valence was made by the American chemist Johnson,2 in 1880. According to his definition an oxidizing agent “is one that can increase the number of bonds of some other substance; hence oxidation of one substance must involve the reduction of some other.” It is this definition which is virtually adopted in all American text books of general or inorganic chemistry. The term “bond,” since the introduction of the ionic theory, is now used with the idea of unit charge. Oxidation thus being a gain in positive charge (or charges) or a loss in negative charge (or charges); reduction would be the reverse of the above operation. Some of the most modern chemists define oxidation as a loss of electrons by an element, while reduction is the gain of electrons. In the past few years, this conception of change of valence has been applied with more or less success in the interpretation of organic reaction^.^ As to the actual mechanism of oxidation-reduction processes, whether of chemical or electrochemical origin, no general theory has been satisfactorily established. Nevertheless, a more or less well-defined concept of the process is now almost universally taught in this country. The general form of this concept may be stated as follows: In oxidation, the primary action is the liberation of oxygen, either from the oxidizing agent itself or from water; the liberated oxygen then reacts with the substance which is oxidized. In reduction, the primary action is the liberation of hydrogen, either from the reducing agent or from water, which then causes the reduction of the remaining substance. To make the point more specific, we quote from several of the most widely used texts: Smith4: “Inorganic Chemistry,” third edition (1918). A. “The permanganic acid, with excess of sulfuric acid, tends to undergo . . . . the following changes, provided a substance is present which can take possession of the oxygen that would remain as a balance,” page 320. 2HMn04+2H2S04=zMnS04+3HzO(+50) I.

The work included in this paper has been taken from the thesis presented in January

of 1925 by Robert Taft in partial fulfillment of the requirement for the degree of Doctor

of Philosophy in the University of Kansas. Chemical News, 42, 51 (1883). Falk and Xelson: J. Am. Chem. Soc., 32, 1637 (1910); Fry: “Electronic Conception of Valence and the Constitution of Benzene.” Smith’s ideas on the nascent state should be reviewed in this connection. See page 543.

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HAMILTON P E R X I N S CADY AND ROBERT TAFT

B. In the preparation of NO from ferrous sulfate, sulfuric acid, and nitric acid, the following equation is used: zFeS04 HzS04= Fez(S04)3 (zH) page j z g . C. In the action of nitric acid upon non-metals the following equation is given, page j36. zHT\’03= (Hz0.Nz06)= zNO HzO ( 3 0 ) D. The oxidation of HC1 by potassium dichromate is represented by two equations, the first of which is KzCr20; 8HC1 = 2KClf CrCL 4 H 2 0 (+3O), page 858.

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Holmes: “General Chemistry” (1921). “In the presence of a reducing agent, dilute nitric acid breaks up, yielding three oxygen atoms for every two molecules of the acid,” page 2 4 2 . In speaking of the reduction of nitric acid by zinc, “all of the oxygen is torn away from the nitrogen by the nascent hydrogen and then hydrogen begins to add on,” page 244. 2.

3.

McPherson and Henderson:

“General Chemistry,” 2nd edition,

(1921).

A. In speaking of the final equation representing the action of zinc on nitric acid, “It is unsatisfactory in that it does not . . . . suggest that hydrogen is first formed, and subsequently transformed into water,” page 267. B. ‘rFez(S04)3 Z H (nascent) = zFeS04 HzS04,” page 628. C. “zKMn04 3H2S04= K2S04 zMnS04 3Hz0 50. This action is not very noticeable unless some reducing agent is present to take up the oxygen,” page 687.

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Norris: “Inorganic Chemistry,” (192I). A. “All processes of oxidation may be considered as taking place in two steps-first, the breaking down of the oxidizing agent to furnish oxygen, and, second, the oxidation of the element or compound present.” page 2 j8. Bancroftl also cites a number of illustrations from various texts of this idea in connection with the action of nitric acid upon metals. That these views are not confined to elementary text books alone is shown by the following reference, although many others could be cited: 4.

Nernst : “Theoretical Chemistry” j t h English edition, page 867. “Chemically an oxidizing material is characterized by its power of giving off oxygen, a reducing material by its power of given off hydrogen.” The original literature, of course, is replete with similar suggestions. TO cite again only a few of the many:Kempf2in an article on the “Electrolytic Oxidation of p-benzoquinone” I.

J. Phgs. Chem. 28, 475 (1924). J. prakt. Chem., (2) Folge, 83, 329 (1911).

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OXIDATION AND REDUCTION

inclines to the view that in all cases of electrolytic oxidation the actual oxidizing apent is a metallic peroxide formed by the action of nascent oxygen on the anode. Armstrong and Colgate’ state that ‘%hereis no immediate direct combination between the oxidizing agent and the oxidizedsubstance . . , the oxygen of the former acts merely as a depolarizer”2. Dunstan, Jowett, and Goulding3 in studying the oxidation of iron in the absence of free oxygen, but in the presence of an oxidizing agent and water summarize their theory of the process as follows : Fe H 2 0 = FeO HZ H z O2 (source not stated) = HZOZ 2FeO H20z= Fez02(OH)2 Munn4 believes like Armstrong that H20z is formed (necessarily) before oxidation can take place. Orlov has determined what he thinks are the rates of the following two reactions :5 Mnz05 = 2MnO 50 3Mn207= 6Mn03 3 0 which take place when an acid solution of permanganate oxidizes potassium iodide. Denis6 in discussing the action of permanganate on alcohol uses the following equation to explain the reduction of the permanganate; the hydrogen, which is “atomic” comes from the action of ethylidene on water: K M n 0 4 3H = KOH M n 0 2 H 2 0 These are but a few of the many illustrations, picked more or less at random over a considerable number of years, which could be cited as proof of our assertion. The number of cases where the writer or investigator assumes such an oxidation-reduction process without actually stating in so many words that oxygen and hydrogen are liberated, is much greater-so great that no attempt has been made to cite references. An examination of the literature would soon convince the most sceptical of this fact. The introduction of this explanation is not hard to trace. It originally started from the fact that many metals in contact with acids gave hydrogen, which gave rise to the supposition that all metals in contact with all acids gave hydrogen’. Later, the development of the theory of electrolytic dissociation with the explanation which it offered for the appearance of oxygen and hydrogen a t the electrodes during electrolysis of certain aqueous solutions, together with the subsequent development of the theory of auto-

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J. SOC.Chem. Ind. 32, 391 (1913). Compare also Armstrong and Acw0rt.h: J. Chem. Roc. 32, 56 (1877). J. Chem. SOC.87, 1564 (1905). Z. physik. Chem. 59, 459 (1907). J. Russ. Phys. Chem. SOC.43, 1524-abstracted in Chem. Abs. 6 , 1247 (1912). Am. Chem. J. 38, 564 (1907) Compare Armstrong and Acworth: J. Chem. SOC.32, 56 (1877).

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HAMILTON P E R K I N S CADY AND ROBERT TAFT

oxidation1 which was used to account for the fact that during certain types of reactions hydrogen peroxide was actually obtained during an oxidation, were considered as good grounds for explaining all oxidation and reduction phenomena as involving a primary liberation of oxygen and hydrogen and the subsequent action of these substances, either directly in the so-called nascent state or indirectly (formation of peroxide). It is strange that this concept with practically no experimental proof should have been so long in use. Presumably on account of the difficult nature of the experimental disproof it has been continued in use for want of any better. Objections have, of course, been made to the theory. These in the last few years have grown to considerable volume. We shall attempt to marshal the facts and arguments against the theory in the following summary: I . It is not necessary to assume the intermediate formation of hydrogen and oxygen. Oxidation and reduction phenomena can be satisfactorily explained from both a qualitative and a quantitative standpoint without the assumption of the intermediate stage suggested above. Thus, from a qualitative standpoint, the process may be regarded as a direct transfer of charges, or, if one prefers, of electron transfer. Both of these methods are in use in several standard texts2. Clark3 has developed the quantitative side of the case with considerable care and a t considerable length. In this connection, it is worth while quoting him : ((ASsuggested , , , . we can regard the reversible transformation of ferric to ferrous ion to proceed through say one of a number of possible courses such as the following: 2Fe’ ’ ‘ Hz = 2Fe’ ’ 2H‘ 2Fe’ ’ ’ 2 0 ’ = 2Fe’ ’ 02 Fe”’ e = Fe“

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“If we express the concentrations by means of brackets, the equations expressing the equilibrium condition for the cases mentioned are as follows : (Fe’ “)2(H2)

_____

(Fe’ ‘ ) 2 ( H ’ ) 2 (Fe

(Fe‘ ‘)YO21

:? :f

-

(Fe’ ’ )

‘)z(O’)2

___--

(Fe’ ‘1 ___ ’

=‘&or

Kzor

=

(Fe’ ’ )

f-----; Kz(02)

(Fe‘ ’ ’ ) =

Mellor: L[Treatise”Vol. I , 925, for explanation and references. Bailey and Cady: “Qualitative Analysis,” 8th edition,, p. 64 and Stieglitz: “Qualitative Chemical Analysis,” Vol. I, pp. 262 and 282. Stieglitz (loc. cit.) also gives a very thorough and careful development of the quantitative side of the theory. His complete development is well worth study. See Chaps. 14 and I j. 3 Studies on Oxidation and Reduction, Parts one and two. Reprints No. 823 and 826, Public Health Reports. Quotation is from Part I, page 5 . 1

2

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(Fe‘ ‘ ’ )

(Fe‘ “)(e) -

= KSor

(Fe’ ’ )

___

(Fe‘ ‘ )

-

KB

-(e)

(Fe’ ’ ) ‘

“For any given ratio of (Fe’ ‘ )

“This procedure is capable of indefinite expansion and shows that from the schematic point of view we are at liberty to choose any hypothetical scheme with which to express the equilibrium state.” It should be noted in passing that although any hypothetical scheme may be used to express these equilibrium conditions, “e” (one faraday per mole) is the only one of the quantities experimentally determinable. Clark goes on to show that, no matter which scheme we prefer to use, all lead to the same result in calculating the electrode potentials of any given oxidation-reduction system, and further that “it is possible to express relative oxidation-reduction intensities in terms of electrode potential.” 2 . I n most cases, the liberation of hydrogen and oxygen is purely hypothetical. That is, it is not experimentally possible to determine the actual presence of either substance during the course of the reacti0n.l I n those cases where gases are actually detectable, it remains to be proved whether they are produced during consecutive reactions or as the result of concurrent reactions. 3. If the formation of oxygen and hydrogen is presupposed in electrolytic oxidation and reduction, then the decomposition voltages of oxidation and reduction systems should be not greater than that of an oxygen-hydrogen system, which is I . 79 volts with polished platinum electrodes.2 Bancroft3 has found that certain oxidation-reduction cells have voltages considerably in excess of this value. Thus, the cell Pt - (SnC12.KOH)- NaC1- (KMn04.H2S04)- Pt gave an E. M. F. of 2.061 volts. If hydrogen and oxygen were the actual reducing and oxidizing agents and had been formed a t the electrodes, this should not have exceeded 1.79 volts. &

4. There is considerable experimental evidence that does not support this

oxygen-hydrogen theory, some of which may be cited: A . Sugden4 has shown that the reduction of ferric sulfate by zinc in the presence of sulfuric acid passes through a minimum by varying the concentraCompare B, Objection No. 4 of this paper. Lehfeldt : “Electrochemistry,” page 177 (1904). 2.physik. Chem. 10, 394 (1892). J. Chem. SOC. 119, 236 (1921).

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HAMILTON PERKINS CADY AND ROBERT TAFT

tion of the acid and the salt, a fact which is in direct disagreement with the view that “nascent” hydrogen is the reducing agent. 3. Milligan and Gillettel found no experimental evidence of the liberation of either of these gases in the reduction of nitric acid by certain salts. According to Milligan and Gillette, the reduction of nitric acid by ferrous sulfate proceeds according to the equations: 2FeS04 ”03 HzS04 = F’e2(SO4)3 HNOz HzO 3HT\’02= HNO, 2NO HzO C. Benson2, in studying the rate of reaction of ferrous sulfate, potassium iodide and chromic acid, found that the speed of the reaction could not be accounted for on the basis of the peroxide theory. “The ‘peroxide theory’ was set up before the rate measurements were made and in attempting to explain them, it falls down.” D. Oxidation can take place in the absence of oxygen and reduction can take place in the absence of hydrogen. A few typical cases will suffice:

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I . Anhydrous ferrous chloride can be oxidized to ferric chloride by dry chlorine, the chlorine being reduced to chloride. 2. Manganous compounds can be oxidized in the dry state to manganates by potassium nitrate, the nitrate being reduced to nitrite.

This, of course, may be a case of thermal decomposition of the nitrate, but nevertheless, the nitrogen has been reduced in the absence of hydrogen. 3. Some compounds are capable of being reduced by other reducing agents than hydrogen; carbon and potassium cyanide for example3.

4. Sulfur is reduced as indicated in the following equation, all substances being dry solids4: 2SnOz zNa2C03 9s = 3 S 0 2 aNazSnSs 2C02. In this case, sulfur acts both as an oxidizing agent and a reducing agent; being oxidized to + 4 in sulfur dioxide and being reduced to -2 in the thiostannate. E. Oxidation and reduction are simultaneous processes-one necessarily accompanies the other. If the advocates of the gas theory are consistent, then both hydrogen and oxygen must be assumed to be formed during oxidation and reduction. Very few of the proponents of the theory apparently are wilIing to go this far. Very considerable trouble would be experienced in explaining the differences in oxidizing and reducing reactions in acid and alkaline solutions if this theory were carried out to its logical conclusion6. If both gases are formed, these reactions would of necessity be slow as will be shown by a consideration of the next paragraph.

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J. Phys. Chem. 28, 744 (1924). J. Phys. Chem. 7, 356 (1903). 3 Caven and Lander: “Systematic Inorganic Chemistry” page 88, for many other examples. Treadwell-Hall: [‘Analytical Chemistry” 4th English edition, Vol. I, page 260. Bee explanation offered by Stieglitz: “Qualitative Analysis,” Vol. I, p. 284, et seq. l


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F. As has been pointed out by Dharl, gases react rather slowly with liquids, so that either there should be a sufficient lapse of time such that the free gases would be produced or these oxidation-reduction reactions would proceed slowly. Both of these possibilities are contrary to fact. These are the main arguments against the theory of actual formation of hydrogen and oxygen in oxidation and reduction processes, While some are debatable, we believe that, in the absence of any proof to the contrary, the most convenient and the most generally correct view, is to regard the process as a direct transfer of charges according to the original definition of Johnson, a conception which is also in accord with the present electronic theory of valence. It must be understood that we are advancing no arguments against the formation of intermediate compounds in oxidation and reduction reactions as there is experimental evidence on record in so large a number of cases that this is not a question for argument. Ostwald’s law of successive reactions is aptly applied to these oxidation-reduction reactions. W e are arguing against the actual formation of hydrogen and oxygen during this type of reaction. It will be noted in the cases cited above under objection “D” that they are for the most part reactions which take place a t high temperatures, i.e. reactions which might be possible on account of thermal decomposition. The proof would be much more convincing if the reactions took place a t normal temperatures or under. It was with this view that the present investigation was begun. Our plan, in brief, was to carry out the electrolytic reduction of a substance containing no hydrogen in a solvent free from that element, and to carry out the electrolytic oxidation of a substance containing no oxygen in a solvent containing no oxygen2. It is generally understood that any chemical reaction which involves oxidation and reduction is separable by electrolysis into two reactions, a cathodic process (reduction) and an anodic one (oxidati~n)~. Choice of Solvents After having stated the problem, the question of choice of solvents naturally arises. A solvent which will conform to the above restriction and still be a good conductor upon the addition of oxidizable or reducible salts is a I( rara avis.” Ammonia suggests itself upon the work of Cady4and of Franklin and Krausj as being a suitable medium in which to carry out oxidations. In carrying out reductions, liquid sulfur dioxide was chosen for trial but on account of the complex character of the cathodic processes was abandoned J. Phys. Chem. 29, 142 (19-75) It should be understood that even if it could be shown that the discharged ions of these solvents are the actual reducing and oxidizing agents, our argument would still hold good. as oxidation would take place in t,he absence of oxygen and reduction in the absence of hydrogen. Compare Bancroft: Trans. Am. Electrochem. SOC.8, 33 (1905); 9, 13 (1906). J. Phys. Chem. 1, 707 (1897). Am. Chem. J . 20, 820 (1898): 23. 277 (1900).

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H

5m

HAMILTON PERKINS CADY AND ROBERT TAF'T


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in favor of phosphoric oxychloride1. Our results of these electrolytic experiments in SOz have been recorded and published separately2. During the course of the investigation, a knowledge of the physical properties of the above solvents was required. It is needless t o say that a knowledge of these physical constants is not only useful in the choice of a solvent for a given purpose but also helpful in interpreting results. AS much time was spent in the search of the literature for some of these constants, it was thought worth while to include as complete as list as possible of these values together with the authority from which they were compiled. The constants are given in Table I. Experimental Part

Phosphoric Oxychloride as a Xolvent, This liquid was chosen on account of the fact that not only is its dielectric constant somewhat higher than is that of sulfur dioxide and that its boiling point is much higher (thus increasing the ease of manipulation), but chiefly because neither the liquid itself or its hydrolytic products are reducing agents. The chief objection to its use for our purposes lay in the fact that it readily takes up water and undergoes hydrolysis-the hydrolytic products, where the oxychloride is in excess, being metaphosphoric acid and hydrogen chloride; the presence of these substances vitiating the results of electrolysis. There is an erroneous idea that phosphoric oxychloride hydrolyzes slowly3, If a small amount of water is poured over POcl3,two liquid layers are formed. hydrolysis takes place slowly a t first and then more rapidly until the reaction takes place with semi-explosive violence, the entire reaction not taking more than several minutes. If the water is agitated with the liquid, hydrolysis takes place immediately. It should be noted that the temperature was not kept constant in carrying out the experiment described above; it should also be remembered that under working conditions reactions are not carried out under constant temperature unless for the specific purpose of determining the rate of reaction. I n general, this liquid is a better solvent than is sulfur dioxide; a more varied number of inorganic solids are appreciably soluble than is the case with the first solvent. Walden4was the first to call attention to the solvent properties of this substance. In his paper, he lists as soluble in phosphoric oxychloiide the following substances: alkali, and many of the heavier metallic, iodides; ferric chloride, hydrocarbons, esters and tertiary amines. Walden determined the conductances of some four solutions and concluded that Poc13 was a good ionizing agent for binary salts. This substance is called “phosphorus” or “phosphoryl” oxychloride in text books and in the literature but the name “phosphoric” is preferable to avoid confusion (which is frequent) with “phosphorous” oxychloride, POC1. * J. Phys. Chem. 29, page 1075, (1925). a Roscoe and Schorlemmer: “Treatise on Chemistry,” Vol. I , p. 677, 5th ed. Z. anorg. Chem. 26, 212 (1900).

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As no reference to the solubilities of oxy-salts was found, some qualitative determinations were carried out. Our method was similar to that in the case of sulfur dioxide1 save that smaller amounts of solvent were used. The results of these determinations together with the solubilities of a few other substances are given in Table 11. TABLE I1 Appreciably Soluble

AsBra KC10 4 (slight) KIO4(slight) FeCL CuBr Hg(CN)z(slight)

212

3KC103 3KBr03 3KI03 KLh.07 K2Cr04

Insoluble

KNO3 HgCl &Fe(CNh KzCz04 CUClZ MnClz

In addition to these cases, mention should be made of several others. Potassium permanganate dissolves giving a faint pink color; if this solution be allowed to stand for a day or so out of contact with air, the color is discharged. This is probably due to the action of the permanganate on hydrogen chloride, as even freshly-distilled phosphoric oxychloride contains some dissolved HC1. After standing some months in a sealed tube, the solution precipitates a white gelatinous substance. Whether this is due to hydrolytic action of the solvent or was simply the reduced manganous salt is a t present undetermined. The formation of gelatinous precipitates seems to be quite common in solutions of this solvent. Copper chromate behaves similarly, a reddishcolored solution is first formed which, on standing, changes to a semi-solid, greenish mass. Similar changes were noted in the case of solutions of ferric chloride and of potassium iodate which had been sealed up for some months. Nearly all inorganic salts which dissolve in phosphoric oxychlorjde form colored solutions. I n the case of the alkali iodides, it is stated in the literature that free iodine is liberated upon solution in this solvent. It is doubtful in our judgment if such would be the case if the solvent were entirely free from traces of hydrogen compounds. A sample of potassium iodide which has been sealed up for some six months with the solvent is brown in color, whereas, a similar tube containing free iodine is red. The absorption spectra of the two solutions are quite diff erent-that of potassium iodide showing absorption of all but the red and a small portion of the green whereas that of iodine absorbs only in the violet. As is to be expected, colored solutions of other salts in phosphoric oxychloride give absorption spectra; none of those so far examined have been banded spectra but consist in a more or less complete extinction of the violet end of the spectrum, the absorption rarely extending into the blue. Upon obtaining some knowledge of solubilities in phosphoric oxychloride, preliminary electrolyses were run upon several substances, potassium iodate loc. cit.

* More so than in SO*;

3

Increasing solubility in the order named.

OXIDATION AND REDCCTION

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and ferric chloride being finally chosen as the most suitable for our purposes. It was noted, however, in one of our preliminary trials when electrolyzing potassium iodate that a gas was liberated a t both electrodes. The solvent which we used was freshly distilled and we were a t a loss to account for the rather unexpected result. The idea occurred to us that the formation of the gas might be due to traces of dissolved hydrogen chloride which would be present as the result of partial hydrolysis of the solvent. Metaphosphoric acid, the other hydrolytic product, would be completely removed by the distillation. To test out the correctness of our assumption, a quantity of the freshly distilled liquid was placed in a eudiometer tube and pieces of metallic potassium introduced-after standing for a day, sufficient gas was collected for analysis. Upon analyzing this gas, it was found to consist very largely of hydrogen1. We considered this as proof of our assumption and undertook to remove hydrogen chloride from the liquid. We finally succeeded by distilling the oxy-chloride with metallic potassium. It seems remarkable that such an experienced chemist as Walden2 should have used phosphorus pentoxide in an attempt to "dry" phosphoric oxychloride. Any moisture present would be converted to metaphosphoric acid and hydrogen chloride, the first of which would be removed by distillation, whereas the pentoxide would be without effect on the hydrogen chloride either physically or chemically. I n our first attempts to remove the hydrogen chloride with metallic potassium, we added the potassium to the distilling flask containing the commercial oxychloride. After distilling off about half of the original mixture, the material in the flask would explode, completely shattering the apparatus. The explosions mere finally attributed to the increasing concentration of the glacial phosphoric acid, which, upon reaching a certain value, would react with explosive violence with the potassium a t the somewhat elevated temperature (105-107"). A first distillation without the metallic potassium to remove the glacial phosphoric acid and then a subsequent distillation with potassium in an atmosphere of dry natural gas enabled us to distill the liquid with safety. The same electrolysis tube as was used in our experiments with sulfur dioxide3 was used in our trials with solutions in phosphoric oxychloride. The purified solvent was distilled into a separatory funnel protected with a drying tower of calcium chloride. The liquid was transferred to the electrolysis tube by inserting the stem of the funnel into one of the necks of the tube, closing all openings save the neck and passing a current of dry natural gas through the apparatus during the operation; the solute having first been introduced in the same manner as described under sulfur dioxide. A solution of potassium iodate obtained by purifying and transferring the solvent as just described gave no gas a t either electrode upon electrolysis. Thanks are due to Mr. I. G. Malm for the analysis of the gas. 2. anorg. Chem. 26, 212 (1900). J. Phys. Chem. 29, ro7j (192j).

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A. Electrolysis of Chromates. It was hoped upon finding that the chromates (or dichromates) were soluble in phosphoric oxychloride that electrolysis of their solutions would result in the reduction of the chromium from the hexavalent condition to some lower state of oxidation. Such a change could be followed by the eye as the chromates (or dichromates) form dark red solutions with this solvent. Electrolysis of potassium dichromate did not yield the desired result nor were any of the forms of chromium in the lower states of oxidation detectable in the residue. Furthermore, there was no deposit upon either electrode-very probably gas was produced but the solutions were too dark colored to observe any electrode phenomena. Similar results took place with copper chromate although one would expect that a t least copper would be deposited upon the cathode. None was observed. It should be stated, however, that these results were obtained before our method of purifying phosphoric oxychloride was worked out, freshly distilled solvent being used. Electrolysis was continued for several hours in most cases, which should have been sufficient time to remove any hydrogen chloride present.

B. Electrolysis of Potassium Iodate. Potassium iodate dissolves sparingly in phosphoric oxychloride. The saturated solution is light reddish-brown in color, darkening on long contact with the solvent (several months), and is a poor conductor of the current when compared to aqueous solutions, but a t least a ten times better conductor than is a similar solution in sulfur dioxide. The solvent is without action upon the solute, for upon distilling the substance with a quantity of the solvent, the residue gave no test for iodide or free iodine, nor was any free iodine found in the distillate. Our first trials gave results showing that iodine was set free during the course of the electrolysis, either as a result of the direct reduction from iodate to iodine, or from iodate to iodide with a subsequent reaction between the two t o produce free iodine. The iodide and iodate will react in freshly distilled solvent to give iodine but there is the possibility that the reaction is caused by the presence of traces of hydrogen chloride. Several electrolyses were carried out after we had devised our method for purifying the solvent, in which case most of the cathodic current is used in the deposition of metallic potassium, no iodides being found in the residue but in each case free iodine was present. Our last trial is reported below: Grams of KI03 used.. . . . . . . . I . 166 Time of electrolysis.. . . . . . . 1 2 days Current.. . . . . . . . . . . . . . . o . 01 amp. Voltage.. . . . . . . . . . . . . . . . . . . . . I I O The current varied somewhat from the value given above but the fluctuations were not marked. The cathode became coated with a white crystalline deposit, which proved, upon examination to be metallic potassium. The anode


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(wire) was covered with a thin yellowish cast which did not grow perceptibly during the electrolysis. The solution was stirred by dried gas from time to time to increase the concentration of potassium iodate, an excess of the solid being always present. The stirring also prevented the potassium from bridging across to the anode. At the completion of the electrolysis, the liquid was transferred to a distilling flask and after distillation, the residue was extracted with ether. The residue left after extraction with ether contained no iodide. The iodine obtained by the ether extraction was titrated against a standard thiosulfate solution and found to contain 0.0563 grams of iodine. The formation of this iodine was taken to prove our point that reduction had taken place in the absence of hydrogen. We feel especially confident of this fact due to the formation and continued presence of the metallic potassium, which would certainly insure the complete absence of any hydrogen. It should be stated further, that the solution deepened in color during the course of the electrolysis, givjng additional proof that iodine was accumulating in the solution. Examination was made of the anode deposit, the quantity of which was exceedingly small. Upon immersing the anode in water, the deposit was freed from the wire and dissolved slowly. The aqueous solution gave tests for chlorides and phosphates, which, of course, might be due to adhering solvent. A qualitative test for iodides was made on this aqueous solution but it gave negative results.

C . Electrolysis of Ferric Chloride. Ferric chloride dissolves to a somewhat greater extent in phosphoric oxychloride than it does in sulfur dioxide, the resulting solution is accordingly a better conductor than a solution of the same compound in sulfur dioxide. The chief difficulty in the use of anhydrous ferric chloride lies in its great affinity for water-especial care was therefore exercised in introducing this substance into the electrolysis tube; the method has already been described in connection with our work upon sulfur dioxide. The electrolysis of this compound produced a black, thin, adhering deposit upon the cathode. After the electrolysis had continued several hours, the solution was removed, most of it being subjected to distillation. A portion of the solution was cautiously added to water; potassium ferricyanide was then added to the cold solution, whereupon a dark blue precipitate was formed. Similar results were obtained with the residue left from the distillation-the residue was extracted with water and then treated with potassium ferricyanide, the dark blue precipitate again appearing. This was taken to be proof that (-) = Fe" had taken place. the change Fe' The nature of the deposit upon the cathode was determined as far as the scantness of the material would allow. An acid solution of this substance had strong reducing properties; producing with silver nitrate a precipitate of various colors, finally becoming black. It also decolorized permanganate solution. These properties resemble to some extent those of the hypophosphites. There is the possibility that the positive ion of the solvent itself has ' '

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,

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been reduced and combined with a part of the F e ' ' to form an insoluble substance, According to Walden, phosphoric oxychloride ionizes Poc13 = Po' 3c1' If the reaction PO"' 4(-) = PO' be assumed to take place, the formal resemblance to the hypophosphites becomes apparent, for ]Es3Po2 - HzO = HPO = H ' PO' While there is no meta hypophosphorous acid recorded in the literature , it requircs no great stretch of the imagination to see that, if it were formed in a solution entirely free from water, it would probably have the composition given above. If this is the correct analysis of the situation, it might be possible that it was PO' which reduced Fe' to Fe' ; it is immaterial from our point of view, as in either case, reduction has taken place in the absence of hydrogen. ' '

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' '

'

11. Liquid Ammonia as a Solvent The amount of work which has been done upon the various solvent properties of liquid ammonia is very large as compared to the amount of work which has been done upon the solvent already described. As it was in this solvent that we hoped to carry out the oxidation of substances in the complete absence of oxygen, a knowledge of the solubilities of substances capable of being oxidized was required. A study of the papers of Cady' and of Franklin and Kraus2 together with some preliminary trials led us to confine our work to the electrolysis of solutions of the following substances : thallous iodide, cuprous iodide, hydrazobenzene, methyl and ethylamine hydrochlorides. Our experimental procedure was much the same as that described in the electrolysis of solutions in sulfur dioxides. The apparatus was dried a t 110' over night, dry natural gas freed from COz was passed thru the tube while cooling, and the substance to be electrolyzed, carefully dried, was then introduced. The exit tube "C" was closed with a tube about a meter in length which dipped into mercury at its farther extremity. This change was made to secure some pressure within the tube as an aid in the condensation of the ammonia and, also, to give some indication of the direction of the pressure. The solvent was drawn from a cylinder which had been charged some years previously with metallic sodium to dry it, and was then passed through a drying tower filled with ignited asbestos to remove any mechanical impurities which previous investigators had found were carried by the moving gas. Commercial ammonia, vaporized rapidly below its boiling point by bubbling natural gas through the liquid, was used as a refrigerant to condense the dried ammonia. J. Phys. Chem. 1, 707 (1897). *Am. Chem. J. 20, 820 (1899)~23, 277 (1900). J. Phys. Chem. 29, 1075 (1925).


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Upon completion of the electrolysis, the Dewar flask containing the refrigerant was removed and the solvent allowed to vaporize through “C,” the adjustment of the height of the mercury being used to regulate the pressure in the tube. Pressure was a t all times kept greater within the electrolysis tube than out, so that any flow by diffusion would be outwards, this serving as an additional precaution in keeping the solution entirely free from moisture. The electrode “D” served in these experiments as the anode.

A . Electrolysis of Thallous Iodide. This salt was prepared by adding the requisite amount of potassium iodide to a hot solution of chemically pure thallous sulfate. The filtrate from the insoluble thallous iodide continued to give tests for sulfates after many washings. The residual sulfate was finally removed by suspending the thallous iodide in several liters of distilled water, shaking for sometime, allowing the iodide to settle out, and then pouring off the supernatant liquid. After several such treatments, no sulfate could be detected. The salt, which is a bright yellow solid, was then dried a t 110’ over night, powdered and then redried over another night. The yellow solid dissolves quite freely in liquid ammonia giving a colorless solution. Upon electrolyzing this solution, black amorphous thallium was deposited with exceeding rapidity upon the cathode-the wire cathode “growing” to a t least ten times its original size in the course of a few minutes. This necessitated breaking the current and stirring with dry natural gas to prevent bridging between the electrodes. At best, the electrolysis could be continued only a short time, A trial using .j 4 4 grams of thallous iodide gave a current of 0.6 amperes, the voltage being 110. A gas was given off a t the cathode. After fifteen minutes, the electrolysis was stopped and the solvent distilled off. The cathode was removed and the residue extracted with ammonia water and filtered. This would leave as a residue thallic hydroxide if any of the thallous ion had been oxidized to thallic, thallous hydroxide being quite soluble in water. A considerable residue was formed; this was dissolved in a small quantity of hydrochloric acid and the solution then neutralized with ammonium hydroxide. A voluminous white precipitate resembling aluminum hydroxide was formed. This was filtered, washed and dissolved again in a small quantity of acid. Sufficient ammonium sulfide reagent was then added to neutralize the acid and precipitate the thallic ion as thallous sulfide, as thallic ion is reduced by this reagent to the thallous state. An abundant black precipitate was formed, as thallous sulfide is black. This was considered as proof that the reaction T1’ 2 (+) = T1‘ ’ had taken place. Liquid ammonia itself is without oxidizing action upon thallous iodide, as was shown by the fact that a small quantity of thallous iodide, when dissolved in liquid ammonia, gave no test for T1’ after removing the solvent.

+

’

’ ’

B. Electrolysis of Cuprous Iodide Some cuprous salts are soluble in liquid ammonia, the chief difficulty from our point of view being the fact that it is difficult to free them from traces of

1072


cupric ion in liquid ammonia solution. Sloan' prepared an ammonate of cuprous nitrate but was unable to obtain a colorless solution. Franklin2 cites a number of instances of cuprous compounds, all of which give colored solutions in liquid ammonia, thus showing the presence of the cupric ion. He was able to prepare a colorless solution of cuprous amide from solutions of considerable age. Cuprous iodide appeared to be the best choice among the copper compounds, as it is soluble in liquid ammonia and contains no oxygen. It was prepared by adding together the calculated amounts of blue vitriol and potassium iodide, the excess iodine being removed by passing sulfur dioxide through the solution. It was washed and dried in the absence of air, ground and redried. The dry salt had a grayish appearance. Upon dissolving this salt in liquid ammonia, a blue solution resulted. Another trial in which metallic copper was allowed to stand in contact with the solution of cuprous iodide, gave similar results-although the solution was not as dark colored as before. This solution upon electrolysis apparently became darker blue in color, although it was difficult to judge by the eye the extent of the color change. It was hoped that the cuprous solution might be obtained colorless, for any color change which occured could be easily followed. There is a possibility then, that the change C u ' (+) = Cu' has taken place.

+

'

C. Electrolysis of Hydraxobenxene White and Knight3 have called attention to the solubility of this compound in liquid ammonia. As the change from hydrazobenzene to azobenzene is a typical oxidation, the electrolytic oxidation of the substance would offer considerable evidence in favor of our concept. The sample of hydra~obenzene~ used was clear white, melting a t 131'. It dissolved in liquid ammonia, giving a pale yellow solution, the solution, however, was an exceedingly poor conductor and other solutes were added to carry the current. I n aqueous solution, hydrazobenzene can be electrolytically oxidized in the presence of alkali, By analogy, since amides give alkaline solutions in liquid ammonia, it appeared possible to electrolytically oxidize hydrazobenzene in the presence of an amide in liquid ammonia. However, the simple addition of sodamide to a solution of hydrazobenzene produced a red coloration, which, after studying the paper of White and Knight6, was ascribed to the presence of metallic sodium in the sodamide (sodium and hydrazobenzene give according to these investigators a red solution in liquid ammonia.) Another explanation of the red color thus produced might be found in the possible presence of the ammonia analogue of sodium peroxide in the sodamide. Additional evidence on this last explanaticn is furnished by the fact that sodium peroxide and hydrazobenzene in liquid ammonia give a red solution. J. Am. Chem. SOC.r'2, 972 (1910). Private communication. J . Am. Chem. Sor. 45, 1780 (1923) Thanks are due to Mr. Edwin C. Wise for the preparation of this compound. J. Am. Chem. Soc. 45, 1780 (1923).


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Electrolysis in neutral solution was then tried, sodium chloride being used as the carrier of thercurrent. This gave a red solution a t the cathode which was without doubt due to the reaction between the hydrazobenzene and the sodium formed a t the cathode, as suggested by White and Knight. Finally, electrolysis in acid solution was carried out, dried ammonium chloride being added to produce the acidity. As this attempt gave positive results, the following trial was made: Weight of NH&1 used.. . . . . . . . . . . . 0 . 6 8 grams Weight of C6H6NHHNC6H,.. , . . . . . , . 0 . 2 3 4 4 ', During the electrolysis, the color of the solution deepened and after passing the current through the liquid for an hour the current was shut off, the solvent evaporated and the residue extracted with toluene. From the toluene solution, long blade-like crystals, dark red in color, were obtained. This amounted to 0.1878 grams and had a melting point of 65". A second crystallization gave a product melting sharply a t 67". The recorded melting point of azobenzene is 68". The agreement was considered satisfactory, taking into consideration the difference between the melting points of the azo and hydrazo compounds. We have shown, therefore, that eighty per cent. of the hydrazobenzene had been oxidized to azobenzene. A blank trial, (i.e., a similar solution without electrolysis) gave no evidence of the formation of any azo benzene. The cathode, after the completion of the electrolysis, was found to be covered with a black substance, insoluble in toluene, alcohol, or water.

D. Electrolysis of the A m i n e Hydrochlorides In alkaline aqueous solutions, ethyl alcohol can be oxidized in the presence of iodine to iodoform. A somewhat analogous process in liquid ammonia, would be the oxidation of the amines in the presence of iodine to iodoform, On paper, for example, the following reaction appears possible: CHSNHZ z I= ~ CHI3 NHJ the oxidation consisting in replacing the positive hydrogen by the negative iodine, Le., the valence of carbon changes from - 2 to + z . No iodoform could be detected when the electrolysis was carried out in the presence of potassium iodide and sodamidel, but upon electrolyzing a solution of methyl ammonium chloride and ammonium iodide, i.e., an acid solution, a trace of iodoform was detectable2 by its odor in the residue left after the water extraction of the contents of the cell. Several trials did not give a greater yield. Since the oxidation of ethyl amine to iodoform would be more strictly analogous to the oxidation OC ethyl alcohol in aqueous solution, attempts were made to oxidize this compound in the form of the hydrochloride, but neither in alkaline or acid solutions were any traces of iodoform observable.

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+

Rather than use the free amine, the hydrochloride from Eastman's purified base was used. Identified by three independent observers.

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HAMILTON P E R K I N S CADY AND ROBERT TAFT

The main product of the reaction appeared to be an iodine substituted amine, probably C Z H ~ N I Z . 111. Discussion of Results The number of cases where positive evidence of our main thesis has been experimentally shown is not large. It must be remembered, however, that limitations were placed upon our problem by the nature of the problem itself and by our method of attack, A review of our work would show that these limitations were : A. The number of substances available for experimental purposes was limited. That is, we restricted ourselves to those compounds which contained no oxygen in a non-oxygen solvent and vice-versa. B. The number of compounds available with the restriction placed in (A) was still further reduced by the fact that they were required to be appreciably soluble in the solvent under consideration. C. The very marked difference in the electrolytic behavior of aqueous and non-aqueous solutions. I n the light of these limitations, we believe that we have succeeded in proving our contention that oxidation can take place in the absence of oxygen and that reduction can take place in the absence of hydrogen. These results are summarized in Table 111. I n addition, some evidence has been produced showing that the following reactions take place: CU' (+) = CU" and CH3NHz 2 1 2 = CHI3 NHJ

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TABLE I11 Substance

Solvent

KI03 Poc13 FeCla Poc1, T1I 3" CsHSNHHNCsHS 3"

Product

0.0563 grams

12

Fe'

Analysis

'

T1' ' ' (C6H5N)2

Qualitative Qualitative 80% yield.

Summary Objections to the conception that oxygen and hydrogen are actually I. formed during oxidation-reduction actions have been pointed out. It has been shown experimentally that oxidation can take place in the 2. absence of oxygen and that reduction can take place in the absencevof hydrogen. 3 . The qualitative solubilities of a number of substances in phosphoric oxychloride have been examined. 4. Some experimental facts have been added to our knowledge of the electrochemistry of solutions in phosphoric oxychloride and liquid ammonia. Chemical Laboratory, Uniiiersity o j Kansas, Lawrence, Kansas, March, 1986.

Oxidation and Reduction - The Journal of Physical Chemistry (ACS

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