J. W. Turrentine, "Reversed electrolysis" (1908)

8/2/2019 J. W. Turrentine, "Reversed electrolysis" (1908)

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REVERSED ELECTROLYSIS

BY J. W. TURRENTINE" In the case of electrolysis the only specific action w ~ i c h we have to attribute to the current is that it tends to setfree the anions at the anode and the cations at the cathode."'At the anode this action may result in the formation of newcations due to the corrosion of the anode by the liberatedanions. Employing a different terminology, these facts maybe expressed as the diminution of electric charges on thecations at the cathode and the increase of like charges at theanode. The electric charges of an ion may be looked uponas the manifestation of valence if we regard valence as po-tential combining capacity. I f we define oxidation or reduc-tion as a change in valence and keep in mind the theory ofvalence changes, from either view point it is seen that at thecathode we have a reduction and at the anode, an oxidation.So, we may adopt as a rule that electrolysis in aqueous solu-tion invariably leads to oxidation at the anode and reductionat the cathode.

Oxidation at the anode may be fully explained by thesimple chemical action of the discharged anion, and reductionat the cathode by the action, likewise, of the dischargedcation. Anions being electronegative in character, on beingdischarged at the anode they appear as strong oxidizingsubstances and either escape as such or corrode the anodeor react with some oxidizable substance in the electrolyticbath. In an analogous manner may be explained the corre-sponding reduction at the cathode.

In pure chemistry we understand oxidation as meaningthe addition of electronegative substances, as oxygen, or thesubtraction of electropositive substances, as hydrogen, anda reduction as the reverse. Since either process may manifestitself as a change in valence or the number of electric changes

1 Bancroft: Trans. Am. Electrochem. Soc., 8, 33 {1905).


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Reversed Electrolysis 449of the ion acted upon, the chemical and electrochemicaloxidation-reduction phenomena are pnlctically identical.

The reduction of one substance results in the oxidationof the reducing agent-which fact is as well-known as thelaw of the conservation of energy. The products of a reduction, however, may appear to be more potent reducing agentsthan the agent. which induced the reaction leading to theirformation. A parallel statement may be made in regardto the products of the oxidation of hydrazine by numerousoxidizing agents1 yet the ratio of hydrogen to nitrogen ishigher in ammonia than that in hydrazine, but the formeris a weaker reducing agent than the latter on account of itsgreater stability. Likewise, the oxidation of phosphine byoxygen, 2 when the two gases, highly diluted by some neutralgas, are allowed to mix slowly by diffusion, results in theproduction of free hydrogen, according to the equation:

2PH3 + 202 = 2HP02 + 2 H ~ yielding a product more highly reduced than the reducingagent itself, though, as in the previous case, more stable thanthe reducing agent.

Potassium permanganate and hydrogen peroxide acttowards each other in acid solution as mutual reducing agents,each giving up a part of its oxygen. s

Depending on- conditions, hydrogen peroxide may actboth as an oxidizing and as a reducing agent, in either casegiving up one-half of its oxygen.

Such irregularities may be explained on the basis of anintermediate reaction, though the existence of such a reactioncannot always be demonstrated. Secondary reactions maytake place which effectively conceal the primary reaction sothat the ultimate result may appear as a reduction by anoxidizing agent or an oxidation by a reducing agent.

1 Browne and Shetterly: Jour. Am. Chem. Soc., 29, 1305 (1907); 30 1 53(HJ08).2 Van der Stadt: Zeit. phys. Chem., 12 , 322 (1R93)., Brodie: Jour. Chem. Soc., 7, 304 (1855); Lunge: Zeit. angew. Chern.,

J89


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/450 ]. W. Turrentine

In the electrolysis of aqueous solutions a reversal of theusual order of things, to which we have given the designationof reversed electrolysis, must always be considered as theresult of secondary reactions since, from definition, the primaryreaction is always one of oxidation at the anode and of reduc-tion at the cathode. Such cases are numerous. Luther 1shows that when Fehling's solution is electrolyzed, cuprouSoxide is deposited on the anode as well as metallic copper onthe cathode, a reduction resulting at both poles. The directand primary action of the current is the oxidation of thetartaric acid to formic acid, an oxidation product but a lessstable substance which reduces the cupric salt and precip-itates on the anode cuprous oxide. In the same way goldcan also be reduced and deposited, as, also, in methyl alcoholsolution, if alkaline, when aldehyde is produced at the anode.

In an ammoniacal solution of potassium bromide,potassium pemianganate is reduced to the green man-ganate at the anode. This result is brought about throughthe oxidation of bromide at the anode to hypobromite;this oxidizes the ammonium hydroxide to hydroxylaminewhich in turn reduces the permanganate, purple, to themanganate, green. Also, when a gold .salt, or a pennanganate,is present in a saturated solution of an alkaline . carbonate,upon electrolysis, reduction occurs at the anode due to theformation there of percarbonates which, breaking down inthe aqueous solution, yield as a product hydrogen peroxide.This acts as the reducing agent towards the gold and thepermanganate.

Yet more striking is the reversal obtained when a dilutesolution of nitric acid containing potassium iodide is elec-trolyzed-iodine being liberated at the cathode as well asat the anode. Nitric acid on electrolysis is reduced to nitrousacid. The latter reacts at once with potassium iodide, oxidiz-ing it and liberating iodine against the cathode.Tommasi,2 on electrolyzing a solution of chloral hydrate,

1 Zeit. Electrochem., 8, 645 (1902).2 Electrochimie, 741.


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Reversed Electrolysis 451obtained chlorine at the cathode. Wood and Jones 1 obtained a deposit of metallic copper simultaneously at bothpoles from a solution of the double carbonate of copper andpotassium. This we may explain as we did the depositionof gold from the carbonate solution as due to a reduction byhydrogen peroxide which had been produced by the decompositions of percarbonates.

Kraus, 2 in his interesting. work on solutions of metalsin liquid ammonia, furnishes us with the most perfect instanceof reversed electrolysis, though hardly comparable to the othercases cited as it does not have to do with aqueous solutions.When a solution of sodium in liquid ammonia is electrolyzedbetween lead electrodes lead is dissolved from the cathodeand appears in the solution as negatively charged ions ofPb 2 The observed loss in weight of the cathode agrees almost perfectly with the value calculated from the electrochemical equivalent of lead.

Hydrogen at the AnodeBeetz1 has described experiments in which he obtained

an evolution of hydrogen from the anode when he electrolyzeda solution of magnesium sulphate between magnesium electrodes. Thin magnesium wire electrodes of the size of knit-ting-needles were projected through the bottom of a glassvessel. The anode, he noted, became covered by a blacksubstance, supposed by the author to be a suboxide of magnesium, which slowly dissolved with an evolution of hydrogen and which, also, filled the solutions in the anoderegion, causing turbidity. As these phenomena would indicate a disintegration of the anode with subsequent decomposition of water by the finely divided magnesium (as theauthor himself deemed possible), a similar experiment wasperformed in connection with this paper, as follows:

1 Proc. Cam. Phil. Soc., 14, 11 , 171 (1907).2 Jour. Am. Chern. Soc., :z9, 1556 (1907).1 Phil. Mag., 3:2, 269 (1866); .Pogg: Ann., 1:27, 45 Cf. also Elsasser:

Ber., 9, 1818 (1876); u , 587 (1878).


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',/452 ]. JV. Turrentine

A solution of sodium sulphate (5 grams to 250 cc H:P)was electrolyzed between a platinum cathode and a magnesium anode. The anode consisted of a bu'nch of magnesium ribbon sealed in the end of a glass tube by means ofcerusine. The tube was bent back upon itself so that theanode in its end could be thrust into an inverted burettewhich served as eudiometer. An oxy-hydrogen gas coulometer was joined in series. The accumulated gas was withdrawn and analyzed over alkaline pyrogallol in a Hempelpipette. The residue, after being subjected to qualitativetests, was taken as hydrogen. Under the conditions of theexperiment the current density at the anode could not beestimated nor maintained constant, so the different measurements obtained can hardly be compared with each other.I t is seen from Table I that the evolution of gas from theanode was vigorous and, from the analysis, was almost purehydrogen.

TABLE IGases from Magnesium Anode

No .

I .2 .34 56.7

T.V .ccJ1.218.81J.89412.7

'444 8

o. H,cc ccI .0 JO. 20 .6 18.20 .6 13.2o.8 8.6o.6 I 2 . I0 .4 14 .00 .2 4 6

Gases from CoulometerT. V. o. H,

- -- - - -cc cc cc

28.2 9 4 18.826 . J 8 . 7 J7.630 5 10 . 2 20.340.0 13 3 26.i40-5 1.) . 5 27.015 .0 5 0 10 .0

While the results shown in this table verified Beetz'sobservations as to the evolution of hydrogen, no indicationwas obtained of the formation of a black compound or cf thedisintegration of the anode. The corrosion of the anode wasin every case rapid and led to the appearance of a voluminous white precipitate of magnesium hydroxide in theanode region. Bubbles of gas occasionally came off from the


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Reversed Electrolysis 453precipitate, but seemed to be merely the escape of the en-tangled hydrogen. The magnesium anode remained whiteand became covered with scales of magnesium hydroxide,periodically dislodged by the vigorous evolution of gas. Onbreaking the current the cessation of the evolution was notsharp, which fact would indicate the formation of some com-pound which gradually decomposed water. Such a sub-stance could only be metallic magnesium or a compound ofmagnesium in which the effective valence of the magnesiumis less than two. I f the latter, it would be acting analogouslyto cuprous sulphate which, in cold solutions, breaks downand deposits metallic copper, and to chromous salts whichreact with water and evolve hydrogen. 1

The absence of a visible change in the surface of theanode and yet a free evolution of hydrogen, persisting for ashort time after the current had been broken, would make itappear that this decomposition of water is quite analogousto that by chromous salts. Instead of the product of theelectrolysis appearing as a solid, as noted by Beetz, theseexperiments would rather indicate the actual solution ofthe anode as magnesious sulphate which at once reacted withthe water of the electrolytic bath and formed magnesic sul-phate and hydroxide, according to the equation

Mg2SO + 2H20 = MgSO. + Mg(OH)2 + H2.This reaction would account for the evolution of hydrogenand for the voluminous precipitate of magnesium hydroxidein the immediate region of the anode. The scale of mag-nesium hydroxide surrounding the anode, by allowing slowdiffusion outward of the unstable salt, would afford an ex-planation of the continuation of the evolution of hydrogenafter the interruption of the electrolysis.

Beetz made a series of quantitative measurements of thehydrogen evolved from the two poles and of the magnesiumdissolved from the positive pole. In every case he foundthat the combined volumes of the gas evolved from the two

1 llfer: Liebig's Ann., 11:1 1 ]02 (IXSIJ).


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454 ]. JV. Turrentineelectrodes was approximately equivalent to the weight of

the magnesium dissolved from the anode.These measurements are not inconsistent with our

assumption of the occurrence under these conditions of Mgwith a lower valence, but they can not, from the nature of thecase, prove the existence of such a form. Indeed, in theexample mentioned above of the oxidation by water of anunstable compound, in which the metal is apparently oxidizedfrom a lower to a higher valence, we may be dealing solelywith a question of linkage of molecules and not with one of achange of valence. Mercurous nitrate, we know from thework of Ogg, 1 must be considered as HgN0 3.HgN03 Ifwritten H g - N0 3 , an actual difference in valence between

Hg -NOsthe mercury in this molecule and that of the mercury in

/NOsHg"'. does not exist.NOSTo bring the structural formula for ferric chloride intoagreement with certain of its reactions it must be considered

Cr -so.as Fe2Cl 8 Chromous sulphate may then be IC r - so.

Mg Mg-Mg"'.magnesium sulphate rnay be )so. or SO.( /SO .Mg Mg-MgBut the matter of linkage is entirely aside from the questionof valence as considered from the electrochemical stand-point. Linkage may exist. There may be no actual changein the valence of a metal. But there is an effective valence,based on Faraday's law and the electrochemical equivaleDts,which manifests itself in every electrolytic corrosion. It iswith the effective valence that we have to do in this paper.

This belief in the lower valence of magnesium is sub-stantiated by the work of others. Christomanos 2 obtained a1 Zeit. phys. Chern., 7.71 28s (189H).2 ner. chem. G e ~ . nerlin, 36, 2076 (1


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Reversed E /ectrolysis 455gray powder when he suddenly cooled a magnesium flamewhich, from its analysis, appears to have the formula Mg 80 6 , but which he regarded as a mixture of magnesium and magnesium oxide. Baborovsky 1 considers the suboxide obtained by Christomanos identical with that noted by Beetz.Considerable light could be thrown on this point if the compound prepared by Christomanos were placed in water; anevolution of hydrogen would strongly indicate the correctnessof Baborovsky's contention.

White, 2 in his study of the action of solutions of bleaching powder on metals, observed that magnesium and aluminumare distinctive in that, while other metals studied, as iron,copper, nickel, etc., evolved oxygen, these two evolved hydrogen, the magnesium producing as much as soo cc in twelvehours.

Luther and Schilow3 employ the hypothetical mag-nesium of a valence of one to explain the reaction which takesplace between iodine, methyl alcohol and magnesium, inwhich hydrogen is evolved. Their supposition of an inter-mediate reaction in which is formed a magneswus iodide isrepresented by the equations:

1\lg + I I 2 12 = 1\fg IMg I + CH30H = Mg IOCH3 + I /2 H2

In short, while the evidence adduced above does not establishthe existence of monoyalent magnesium, yet the chemistryand the electrochemistry of the metal demand the assumptionof such a form.

Hydrogen at an Aluminum AnodeIn the chlorination of organic compounds where the socalled ' 'carriers'' are employed, aluminum chloride, ferric

chloride, and stannic chloride, all behave practically identi-1 Ber. chem. Ges. Berlin, 36, 2i19 (1903).2 Jour. Soc. Chem. Ind., :u , 132 (I)OJ).3 Zeit. phys. Chern., 46, Ho3 (1903).4 "Data on the Chemical Role of Catalytic Agents," Jour. Rus.


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]. l'V. Turrentinecally. Ferric chloride and stannic chloride are both susceptibleof a reduction to a lower chloride, thus furnishing an explanation of their r6le as catalytic agents in chlorinationprocesses on the ground of intermediate reactions in whichthe ferric or stannic compound, in conveying chlorine intothe molecule of the body to be chlorinated, is reduced to theferrous or stannous compound, respectively, and is subsequentlyreoxidized, the oxidation taking place instantaneously. Theclose analogy between aluminum and iron and tin as carriers ofchlorine made it seem probable that the analogy extended to thematter of valence, our lack of knowledge of aluminum compounds, in which the aluminum appears with a valence otherthan three, being absolute because of the great instabilityof such compounds. The position of aluminum in the periodic system, furthermore, being a close neighbor of bothcarbon and boron, makes it appear odd that this elementdoes not exhibit the property of linkage as do both carbonand boron. The evidence obtained as to the dual nature ofmagnesium from the behavior of that metal when marleanode led to the belief that a clue could also be gotten in thesame way to the existence of aluminum with an effectivevalence lower than three. Accordingly, actuated by theseconsiderations alone, we made aluminum anode in a solutionof sodium chloride containing 10 grams of the salt. to 200 ccdistilled water. The same apparatus was employed here asthat used in the like experiment with the magnesium anode.A strip of aluminum foil was sealed, by means of cerusine,in a glass tube so bent that the exposed end of the foil, whichwas to- serve as anode, could be projected into t h ~ mouth ofan inverted burette to serve as eudiometer. Each face of theanode was about a square centimeter in dimensions. Forcathode was employed a strip of platinum foil. A gas coulometer was joined in series with the electrolytic cell. Theaccumulated gases were withdrawn from the eudiometer intoa Hempel gas burette, were measured and were then analyzedfor oxygen by shaking with alkaline pyrogallol in a Hempelpipette. The residual volume, which gave affirmative qualita-


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Reversed Electf'olysis 457tive tests, was taken as hydrogen. The results are given inTable II.

No.

I234

TABLE I I----- ---- - ~ ~ - - ~ ~ ~ ~ - - - - ~ - - ~ - - - - ~ Gas from Al Anode

T.V. 02 Hs .cc cc cc5-0 0. 2 4854 o. 2 5-25.8 0. I 5-710.6 0. I 10.5

il!III

II

Gas from ConiometerT.V.cc40.050.0450

02ccIJ.J16.615.0

20.6334JU.O

The anode was rapidly corroded as a high current densitywas used; and, as in the case of the magnesium, the solutionbecame clouded with a voluminous precipitate of aluminumhydroxide of the characteristic gelatinous consistency. Tlfeevolution of hydrogen persisted for a few minutes after theelectrolysis had been discontinued.Unfortunately, electrochemical methods and efficiencymeasurements do not admit of the determination of theeffective valence at which a metal is dissolving from the anodewhen the entire corrosion does not take place in accordancewith that figure. When the metal has the choice of a numberof possible valencies still greater are one's difficulties. Thisis readily seen from an illustration. I f the aluminum dissolved quantitatively on the trivalent basis, stable aluminum trichloride would be formed and no hydrogen evolved. I fit dissolved quantitatively to aluminum dichloride, the reaction with water could be represented thus:

6 A l C l ~ + 6H20 = 4AlCI8 + 2Al(OH)8 + 3H2 ;i f to the monochloride, the equation,

3AlCl + 6H20 = A1Cl1 + 2Al(OH)1 +3H2,would represent its reaction with water. On the above basisfrom the ratio of the aluminum chloride in solution or of theprecipitated Al(OH) 8 to the hydrogen evolved could be determined readily enough whether the aluminum dissolved as


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]. W. Turrentinethe monovalent or as the divalent form. However, it is impossible to show from such analytical data, when we are nothypothesizing a quantitative corrosion as one definite elec-trochemical equivalent, that the corrosion has not taken placeat any one of the numerous possible percentage combinationsof the three valencies.

If, however, the yield of hydrogen at the anode couldbe shown to increase with increasing current density andcould be made to approach asymptotically a certain definitevalue, a clue could be obtained as to the form of aluminumwith which we are dealing which is not based on analogyand supposition.The striking similarity of behavior between the magnesium and the aluminum is further instanced by the likebehavior of the two in bleaching powder solution, 1 wherethey both dissolve with the evolution of hydrogen.

Wohler and Buff, in an article entitled "A Compoundof Silicon with Hydrogen," published some 50 years ago,2made note of the evolution of hydrogen at an aluminumanode. The subject has since been discussed by Nordenwho produces evidence also from mineralogical and othersources to substantiate the hypothesis of the probable existence of aluminum of a lower valence.

Oxygen at the CathodeThe statement of Victor Meyer' that hydrogen, if shaken

with a solution of potassium permanganate, is absorbed andthat an equivalent (one-half the volume) amount of oxygenis evolved led to the belief that such a reduction with accompanying secondary reactions could be duplicated electrolytically. Accordingly, a solution of potassium permanganate,containing 5 percent by weight of KMnO, and 2.5 percentby volume of concentrated sulphuric acid, was electrolyzed

1 White: Loc. cit.2 Liebig's Ann., IOJ 1 218 (1857).1 Zeit. Elektrochemie, 6, 159 (189


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Re'Versed E lectf'olysis 459between platinum electrodes in a specially constructed cell.The cell consisted of a beaker of 200 cc capacity in which,to serve as cathode compartment, was inverted a 100 ccgraduated tube, drawn out at its upper end into a capillarytube and closed with rubber tubing and screw clamp. Theplatinum plate to serve as cathode, about 3 em 2 on a face,was fastened to a short wire which was fused in the end of anarrow glass tube; the tube was so bent that the cathodecould be projected up into the inverted, graduated tube.Connection was made with the cathode by filling the narrowtube, in which it was sealed, with mercury. To prevent thed ~ u s i o n of the anode gases into the cathode compartment,the anode was carefully enclosed in a parchment envelope.Connection was then completed through a Riihstrat resistance frame and milammeter, with the storage cells. Thesolution was placed in the cell and drawn up into the eudiometer until it had displaced the air therein; a current was thenpassed until a convenient volume of gas had been collectedin the eudiometer, upon which it was withdrawn into a gasburette, was measured and was analyzed for oxygen by shaking with alkaline pyrogallol in a Hempel pipette. The residuefrom the absorption after being subjected to qualitativetests, was in each case taken as hydrogen.

There was obtained in every case, as shown by TableIII, a marked evolution of oxygen from the cathode.

No. T.V.cc-- ---I 42.02 25.63 26.64 52.05 so.66 10.87 12.7

TABLE IIICathode GasesHs 02cc cc- ---

299 I 2. 1179 77I9.0 76435 8.s42.0 8.6s.o s.87.0 57

- -

Milliamps Time,hrs.- - - - - - - - ~ - -I25.0 I.756s.o I 525.0 s.o250.0 1.01.0 18.oso.o 1.5so.o l . 2


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] . V. TurrentineThe cathode became heavily encrusted with a hard and

compact coating of manganese dioxide which cra,cked andfell off, only to be refonned. The solution, also, becamethick with a black precipitate which, without analysis, wastaken to be manganese dioxide.

Morse/ replying2 to the paper of Meyer, showed that aspontaneous decomposition of potassium permanganate occursin solutions of this compound' in the presence of manganesedioxide, with an evolution of oxygen and that the reactioninvolved is accelerated by an increase in the amount of manganese dioxide present. With this fact in mind a number ofblank experiments were run with solutions of the same concentration as above and contained in the same apparatusas that in which the electrolysis was conducted. As it wasobserved that the spontaneous evolution of oxygen increasedup to certain limits with the age of the solution, due, nodoubt, as stated by Morse, to the accumulated manganesedioxide in the solution, the blank experiments were run bothbefore and after the experiments in which electrolysis wasemployed. On the basis of the average rate of evolution ofoxygen thus procured were made corrections in the calculations which followed. In Table IV may seen the results ofa number of the blank experiments.

No.I .02 .0J.O

Time,hrs.

10 .012 . 024 .0

TABLE IV--- ... --- - - - - - -

II o,ccr - ~ ~ ~ I 9-018.o

I No. Time,lus.4 -0 : 12 .05.0 I 48.06.o I 36.o

o,cc2.5

34-021 .0

No.7-08.og.o

Time,hrs. 02cc---1 - ~ - - -

IJ.O10.02 i .0

I f corrections are applied to the values shown in Table IIIon the basis of the figures in Table IV it is evident that the

1 Morse, Hopkins and Walker: Am. Chern. Jour., t8, 401 (1896).2 Ber. chem. Ges. Berlin, 30, 48 (I8


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Reversed Electrolysisevolution of hydrogen at the cathode has been greatly accelerated by the electrolysis.

When a solution similar to that used in the above experiments was electrolyzed in a single-ann gas coulometer,the anode and cathode gases being caught together, thefollowing figures were obtained which show that the yield of

TABLE V

No. T.V. H2 02 No. T.V. HI 02cc cc cc cc cc ccI '----I 15.6 I .0 14.6 I 5 I2 .8 1.5 I I. 3

2 44 1. 4 J.O 6 I ] . 6 I . 2 16.43 IO.O I .0 9 0 7 1].4 1.5 IS 94 s .8 O .J 55 I 8 1 ] . 2 I .0 16.ohydrogen was practically nil, or as shown by Meyer, and alsoby Jones/ which seems more probable, that the hydrogen wasabsorbed by the permanganate. Yet more probable doesit seem that the reduction in the evolution of hydrogenfrom the cathode is a matter explainable by the proximityof the anode to the cathode, the oxygen there evolved actingpossibly as a depolarizer at the cathode.

To detennine the cathode efficiency in hydrogen andoxygen, so that a ratio could be obtained between the oxygenevolved at the cathode and t}Je hydrogen which failed to beevolved, an "oxy-hydrogen gas" coulometer, filled with dilutesulphuric acid, was placed in the circuit in series with theelectrolytic cell. A number of these results, chosen at randomfrom a large number obtained, are given in Table VI.

The ratios, however, obtained on the following basis,varied widely; the results could not be duplicated even thoughall the conditions within one's control were maintainedconstant.In attempting to explain the anomaly of the ~ v o l u t i o n of oxygen from the cathode the hypothesis of a secondaryreaction, from analogy within such cases, was adopted. And

1 Jour. Chern. Soc., 33, 95 (1878).


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J. W. Tu"entineTABLRVI

C a t ~ o d ~ - G : ~ s from M n ~ ~ ~ - ~ a s e s from CoulometerNo. 1 T.V.

I2345678910

II12131415I6I718

'I5.016.222.623.224.296I6.8 1I7.820.415.814.024.2I8.37-512 .8II .617 .oI8.4

IJ.O.O I 5-010.4. 5.815.6 I 7.0I6 .7 6 .5I9'.215-08. I I . 51I3.2 J.614.6 3. 2 .

17.2 3-.2 i14.0 I.8110.0 4-020.3 3-91I4.2 4-I5-5 2 .0 'II .4 I 4

10.8 o:81I34 34I4.8 3-6

- --- - ----;--T.V. , H 1i

I75-0 :40.0 i42.0 I40.0 I6o.o41.050.039-041.028.035-645-040.040.040.040.040.040.0

50.026 . ]28.026.740.027-333426.027-318.723.830.026.726 . ]26.726.726.726.7

o.25.0133I4.0I3320.0I37I6.6I3.0I3.893II.8I5.01J.313-3IJ . JI3-313-3IJ . J

Ratio of o:z:ygenevolved to hydrogenabsorbed (on basis0 1 - 2H.)approximate10:4012: I6I4: IJI3: 10IO: 213= I97= 106: 126:104= 508: I48:108: 124= 213: 152: 167: 137: I2

I :43=4I:I:I : 2I :62:31 :21 :24=54=74=52:3I: 5I: 5I:81:21:2

in the remaining pages will be given, as briefly as possible,a description of the experiments performed with the viewswhich led to their performance, in our attempts to arrive atan understanding of the underlying chemistry of the elec-trolytic reductions of potassium permanganate. Morse1 ex-plained the catalytic action of manganese dioxide on potassium permanganate on the ground that the former behavesas a carrier of oxygen. In substantiation, he showed byanalysis that manganese dioxide on standing in air lost oxygenand on being replaced in a solution of potassium permanganate regained oxygen. To use this fact as an explanationof our observations, the acceleration of the cathode reactionby the current must be due to the formation of a particularlyreactive form of manganese dioxide, Mn02, which was oxidizedby the Mn20 7 to a higher, instable oxide, as M ~ O , , this at

1 Loc. cit.


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Reversed Electrolysisonce breaking down with the evolution of oxygen. Such areactive form of Mn0 2 might be a soluble form of tetravalentmanganese as Mn(S04) 2 , 1 a substance prepared chemicallyby reducing potassium permanganate with manganous sulphate, and stable only in strongly acid solutions, hydrolyzingto manganese dioxide in weakly acid solutions. No proofcould be adduced, however, of a reaction between the Mn02and Mn20 7 by adding the one to the other, as the concentratedacid necessary to hold the tetravalent manganese in solutionwas sufficiently strong to cause alone a slow evolution of oxygenfrom the heptavalent manganese. The two substances werethen brought together in an oxidation-reduction cell, constructed as follows :Two small porous cups were filled respectively with asolution of potassium permanganate containing 5 percent byweight of K.Mn04 and 15 percent by volume of concentratedH 2SO, and a solution of Mn(S04) 2 Each was tightly closedby means of a two-holed rubber stopper. Through one holein each stopper projected a glass tube through which, bymeans of a platinum wire s e a l ~ d in its end, connection wasmade with a wide coiled platinuf!I foil electrode. Throughthe other hole a capillary tube led from the interior of thecup to a eudiometer dipping beneath water. The two cupswere then placed in a vessel filled with 15 percent concentrated sulphuric acid. The electrical circuit was completedoutside the cell through a millivoltmeter. The course of thecurrent through the external circuit was from the potassiumpermanganate compartment to that occupied by the manganese disulphate. The electromotive force generated pro-duced a difference of potential between the poles of the cell ofo.1 volt.

I f Mn02 reduced Mn20 7 , according to our hypothesis,Mn,O, should be the point of equilibrium between the two.Then in the cathode compartment of the cell we should havea reduction as represented by the equation,

Mn20 7 + 2H2 = M ~ + 2H20,- - - - - - - - - - -1 Zeit. Elektrochemie, n , 853 (1905).


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J. W. Tuf'f'entineoccurring; and in the other compartment, the oxidation, asrepresented thus,

2Mn,O, + O, = 2Mn20 6As is represented by these equations, the oxygen consumedat one pole is equivalent to the hydrogen consumed at theother pole. We then have formulated per equivalent ofelectricity one part of Mn,06 by the reduction and two partsby the oxidation. From the decomposition of the Mn20,to Mn,O ,, we should then have evolved from the cathodecompartment one volume of oxygen to two volumes from the anode compartment.From the relative amounts of oxygen evolved from thetwo poles, however, no evidence could be obtained to showthe occurrence of a reaction between the two manganesecompounds.

The reduction of Mn,07 by electrolysis directly to theinstable Mn 20 5 would account for the presence of oxygen inthe cathode gases. Numerous arguments for the existenceof pentavalent manganese can be found in the periodic system.The proximity of this element to chlorine, and bromine, etc.,would lead one to expect the existence of Mn 20 1, analogous toCl,O, and Br,O,, less stable than Mn 20 7 just as Cl20 6 is lessstable than Cl 20 7 Four molecules of hydrogen would reducetwo molecules of Mn 20 7 to Mn 20 6 , according to the equation,

2Mn 20 7 + 4H:a = 2Mn 20 6 + 4H202Mn 20 6 = 2M11zO, + O,2H 2 : 0 2 = 2 : r.

Two molecules of Mn 20 5 breaking down to Mn,O, would. evolve one molecule of oxygen, making the ratio of oxygenevolved to hydrogen absorbed 1 : 2 , a ratio frequently obtainedexperimentally.The reduction of Mn,07 to Mn 20 11 and the subsequent de-

composition of this oxide to the dioxide, according to theequations,Mn20 7 + H2 = Mn,O, + H,O

Mn20, = Mn20, + 0 12H:a: o, .. I: 2


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Reversed Electrolysiswould yield a ratio of oxygen evolved to hydrogen absorbedof 2: r, a ratio in no instance realized experimentally. Be-sides no reaction between manganate and permanganatecould be found in which oxygen is evolved; the decompositionof manganates results in the formation of manganese dioxideand permanganate, as shown by the equation,

3Mn20, = Mn:.04 + 2Mn:.07an irreversible reaction in acid solution, the mat1ganese dioxideprecipitating and the reaction therefore running to an end.With the tetravalent manganese present in a soluble form asMn(S04) 2 , the reaction should be reversible.The cathode conditions in the electrolysis were probablyfavorable to the production of manganates as the solutionin direct contact with the cathode was no doubt alkalinefrom the formation there of KOH. The enveloping coat ofmanganese dioxide, playing the r6le of a diaphragm, wouldadmit of slow diffusion of a manganate into the outer solutionof acid permanganate. However, when these conditions wereduplicated as closely as possible chemically and mechanically,no evolution of oxygen could be obtained.

Thus were practically eliminated, assuming the accuracyof our observations, the hypothesis of a reduction by Mn02capable of an acceleration by the electrolysis and of a directreduction of Mn20 7 , at the cathode to either Mn20 5 or Mn20 41

I t now seemed evident that the elusive reaction must be afunction of the diaphragm formed by the heavy, compactcoat of manganese dioxide enveloping the cathode. Thecomplications incident thereto were accordingly eliminatedby the use of a rapidly rotating cathode.

The cathode was attached to a rotating device! Itsstem projected through a mercury seal into a porous cup.The cup was tightly closed with a rubber stopper containingthree holes. Through a central hole projected a glass tubethrough which worked the shaft of the rotating cathode;surrounding this tube was the annular shaped, mercury cup

1 Cf. L3b: Zeit. Elektrochemie, 7, uS (x


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]. W. Turrentineformed from a glass tube of larger diameter, supported by aclosely fitting rubber stopper and blown out at its upper endinto a hemisphere with inturning edges. Dipping into themercury cup from above was another glass tube sealed at itsupper end to the stem of the cathode. By this means thecathode was allowed to rotate freely through an air-tight,mercury joint, capable of withstanding any pressure depending on the height of the mercury column. Through a second hole in the stopper projected a glass tube for introducingsolutions into the cup, bearing a stop-cock and, on its upperend, a thistle-top funnel. Connection was made through thethird hole by means of a capillary tube with a eudiometer.As anode was employed a broad platinum plate surroundingthe porous cup. A gas coulometer was joined in series andafter the cell had been partially filled with sulphuric acidof the desired strength, the current was turned on, withcathode rotating, to test the tightness of the cathode compartment. The required amount of potassium permanganatein saturated solution was then added through the funnel tubein sufficient amount to bring the concentration of the cathodesolution up to the desired point and the electrolysis was continued.

With rotating cathode which precluded the possibilityof the formation of a coating of manganese dioxide, even athigh current densities no gas was evolved, showing a quantitative reduction of the permanganate to a stable form. Manganese dioxide was formed as before. With stationary cathode,under the same conditions of current density, gas was evolved,thus establishing the correctness of the hypothesis that themanganese dioxide diaphragm was responsible in some wayfor the anomalous evolution of oxygen from the cathode.I t seemed quite possible that the manganese dioxidediaphragm might act like a bipolar electrode; but when thispaper was presented at the May meeting of the AmericanElectrochemical Society, Mr. F. A. Lidbury suggested that hy-drogen peroxide might be formed by the electrolytic reductionof manganese dioxide. This hydrogen peroxide would be un-


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Reversed Electrolysisstable and would evolve oxygen. Subsequent experimentsby Mr. Wilkinson have shown that the explanation offeredby Mr. Lidbury is the correct one.At various times during the work with KMnO4 the odorof ozone was noticed. This odor being particularly strong inthe gas from the enclosed cathode compartment used in thelast experiment, a sample of this gas was subjected to qualitative tests for ozone by the method of Keiser and McMaster.1Potassium iodide was instantly oxidized by it. The absenceof peroxides which could scarcely be evolved from a solutionof KMnO4 was demonstrated by the non-appearance of ablue color on bubbling the gas through a solution of potassiumferricyanide and ferric chloride.This work was suggested by Professor Bancroft. Itscompletion was made possible amid numerous distractionsonly by his constant i n t e r ~ t and assistance. The authortakes this opportunity of expressing to him his thanks andappreciation.

Cornell University,April, ItJo8.

1 Am. Chern. Jour., 39, 96 (1908).

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J. W. Turrentine, "Reversed electrolysis" (1908)

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