The Oxygen Electrode as a Quasi-Quantitative Instrument

THE OXYGEN ELECTRODE AS A QUASI-QUANTITATIVE INSTRUMENT* BY WILLIAM T. RICHARDS

Before 1915 the thermodynamic value of the hydrogen-oxygen chain had been definitely fixed as 1.23 volts by indirect methods', and the attempt to obtain reversible potentials with the platinum-oxygen gas electrode had been abandoned. Only recently has interest in this electrode been revived from a more practical point of view, and little or no theoretical discussion has accompanied it. Furman2 and Rideal and Goarda and others4 have, however, demonstrated that the instrument is of some practical value. It is with the hope of extending the range of usefulness of this electrode in analytical work for the determination of hydrogen-ion activities and the partial pressure of oxygen in gaseous mixtures that the present communication is presented. The results given herein are essentially a condensed version of one-fifth of a thesis submitted to the Department of Chemistry of Harvard University in partial fulfillment of the requirements for the degree of Doctor of Philosophy in May 1924, and were for the most part obtained in the Wolcott Gibbs Laboratory under the supervision of Professor T. W. Richards. They were intended to substantiate a theorem which it was found necessary to abandon, and have not been previously published because the writer did not consider them of general interest. No priority is claimed for that part of the work which has been subsequently and much more completely investigated by others, I n order to prevent the paper from assuming the proportions of a Gargantua without the vitality of that engaging character, it has been necessary to suppress both a lengthy historical introduction and many numerical data. Any one to whom these may be of use may obtain them either from the files of Harvard University or by personal communication with the author. References to the literature will be confined only t o those most closely touching the work in hand and, to avoid duplication and consequent waste of space, be quoted largely in the final explanatory section. Many which have been necessarily omitted may be found in the papers of Schoch5or Furrnan.6 * Contribution from the Chemical Laboratory of Princeton University.

' Kernst and von Wartenberg: Z. physik. Chem., 56,534(1906);Sachr. kgl. Ges. Wiss., Gottingen, 1905,35;Lowenstein: Z. hysik. Chem., 54,715 (1905); yon Wartenberg: ibid., 56, 513 1906); Langmuir: J. Am. 8hem. SOC.,28, 1357;Lewis: 139 (1906); Bottger: Z. physik. Ahem., 46, 521 (1903);Noyes and Kohr: 42, 336 (1903);Bronsted: 65, 84, 744 Taylor and Hulet6: J. Phys. Chem., 17,565 (1913); Lewis: J. Am. Chem. Soc., 28, (1909); 158 (191%);Z.physik. Chem., 55,465 ( I @ ) . *Furman: J. Am. Chem. SOC., 44,2685 (1922);Trans. Am. Electrochem. Soc., 43, 79 (1923). Goard and Rideal: Trans. Faraday SOC.,19,740 (1924). Britton: J. Chem. Soc.,125 1572 (1924);127,1896(1925);130,147(1927);Smith and Tiliey and Rahton: Trans. Am. Electrochem. Giesy: J. Am. Pharm. Assoc., 12,855 (1923); SOC., 44,31 (1923). Schoch: J. Phys. Chem., 14,665 (1910). Furman: J. Am. Chem. SOC.,44,2685 (1922);Trans. Am. Electrochem. Soc.,43,79 (1923).

THE OXYGEN ELECTRODE

1.

991

The Preparation of the Electrode Surface

Although a systematic experimental summary of the conditions affecting potential of the oxygen-platinum electrode have not been made, investigators are unanimous in ascribing great importance in the behaviour of the electrode to factors of preparation and environment. It has been found by the present writer that, for the obtainment of reproducible results, certain details of preparation must be observed; these are briefly enumerated below. i. The electrode surface should be as large as possible. Whereas the dissociating power of platinum for hydrogen is large, and a small electrode gives reproducible potential values, its dissociating power for oxygen is small, and a very large electrode is advisable. I n this investigation electrodes of sheet platinum I X z cm. mere used. They were heavily coated with platinum black by the usual electrolytic process, using platinum chloride free from lead acetate, a high current density, and an e.m.f. of about I O volts. They were heated orange red and replatinized when they began to give irregular values due to poisoning. ii. After platinization and washing, the electrodes were transferred to a dilute oxysalt (preferably borate or phosphate) solution, and alternately made anode and cathode in this solution to remove traces of chloride ion due to the initial plating. A great difference was observed if the electrode was finally made anode or cathode in this process. I n the former case, the potential of the oxygen-hydrogen chain may be raised to 1.5 volts when the electrode is subsequently connected to a gas-cell chain, the magnitude of this effect being approximately proportional to the current density and to the length of time of electrolysis. 4 rise in electrode potential of this kind may appropriately be termed a “superpotential” to indicate that it is consequent upon and the inverse aspect of overvoltage. I n electrolyte solutions the superpotential of the oxygen electrode was quickly discharged. If the electrode was dried immediately after the purifying electrolysis it retained, however, its excess charge for many days but not indefinitely. If last made cathode in the purifying electrolysis, the electrode caused the hydrogenoxygen chain to have a very low initial potential, which quickly rose to a value about 1.0volts followed by a gradual fall toward the same limit as that approached asymptotically by the over-charged electrode. An electrode which has reached this final, virtually constant, value will be spoken of as a “seasoned” electrode in the subsequent text, and will be used in all measurements, unless otherwise specified, as the only reliable oxygen electrode. “Seasoning” may also be effected, although not as safely, by discontinuing electrolysis, with the electrode as the anode, a t the first instant that gas bubbles are seen to appear on its surface with a low current density. iii. The surface immersed of the platinum electrode must be maintained as constant as possible, a sudden large increase of immersed surface causing a sudden fall in potential. The simplest method of maintaining adequate constancy of this condition is to have a stream of oxygen gas bubbling through the solution sufficient to keep wet the entire surface of the electrode. This

992

WILLIAM T. RICHARD6

is also made necessary by the fact, brought out in Section 4, that the potential of the electrode is dependent upon the partial pressure of oxygen gas above the solution. iv. The type of solution in which measurements are made has also a profound effect on the constancy of the oxygen electrode potential. Furman' mentions a greater fall per unit time in alkaline than in acid solutions, and earlier investigators have mentioned specifically the effect of sulphate ions.2 If, as seems incontestible, the potential of the electrode is due at least in part to a n oxide of platinum, the oxidizing power of the medium in which it is immersed should influence its magnitude. This was, indeed, found to be true. Nitrate solutions gave a continual fall in potential, sulfate solutions a somewhat less rapid fall, and only phosphate and borate solutions-both excessively hard t o reduce-gave results approaching constancy. 2. Apparatus: A Simple Cell with a Flowing Junction The Hildebrand type of hydrogen electrodes showed, in comparative tests, considerably less ability to give quickly a sharp reproducible hydrogen potential than did the enclosed type of electrode where the partial pressure of the gas is maintained constant at one atmosphere over a saturated solution. This type of electrode was therefore abandoned by analogy for oxygen. An electrode-chain with a flowing junction' u+g only small portions of solution was devised in its place which, although familiar in principle, is perhaps sufficiently convenient to warrant description. It is pictured in Fig. I . A represents the hydrogen electrode vessel with its bubbling jet. B represents a normal calomel electrode of convenient design fitted with a pressure bulb H which is level with another pressure bulb Z, also connected with the chain. C and D are three-way stopcocks of the T type connected with their arms in series, their stems being oppositely directed. The tube J serves as an outlet for the dropping bridge. The pressure bulb I filled with the solution under measurement in A , is level with the potassium chloride pressure bulb H . The thermostat level is kept at TT' thus immersing at constant temperature all of the apparatus save only the junction stopcocks and tubing. If the room temperature is approximately constant this device assures reproducibility of the junction potential. It is convenient to bend the apparatus normal to the paper a t KK' thus allowing the outlet tube J to drop outside the thermostat into a conveniently placed vessel. The only point a t which this apparatus can claim the least originality is a t the stopcock C. One arm of the bore of this stopcock, that nearest the hydrogen electrode when all three passages are opened, is tightly plugged with absorbent cotton. Thus a clear passage is left between I and D,but the flow of solution Z into A is effectually prevented, while, at the same time, electrical conduction 1

Furman: J. Am. Chem. SOC.,44,2685 (1922);Trans. Am. Electrochem. SOC., 43, 79

(1923).

* Bose: Z.physik. Chem., 38, I (1901);Z. Elektrochemie, 7, 817 (1901);Wilsmore: Z. physlk. Chem., 35, 291 (1g00). * Hildebrand: J. Am. Chem. SOC.,35,847 (1913). 4 Lamb and Larson: J. Am. Chem. SOC., 42,229(1920).


993

is freely maintained by the saturation of the plug with the solution under measurement. Because of this treatment it is best to use in stopcock C a solid bored plug rather than a hollow blown plug. The procedure in using the apparatus is as follows: After filling the electrode vessels A and B and t'he pressure bulbs Z and H with suitable solutions, enough liquid is allowed to run through the tube system, by adjustment of C and D and the stopcocks controlling the contents of the pressure bulbs, so that a liquid junction is effected a t D. D is then completely

closed and the potential measurements carried out until a satisfactory constancy has been observed. During the time necessary for this process the value of the junction difference has altered, and the total potential of the chain therefore no longer represents its reproducible value. By a suitable adjustment of D , and of the pressure bulb stopcocks, a gentle dropping of both liquids is then secured, and a changing liquid junction surface is formed at D and removed through J. The potentiometer reading is then taken, and its definitive value recorded. If the bulbs H and Z are of 2 5 cc. capacity several such readings may be taken without refilling. The performance of this electrode chain has been very satisfactory. Potentials about the neutral point are reproducible within two tenths of a millivolt and, with acid or alkaline solutions of considerable concentration, within, at worst, a millivolt. The error avoided owing to junction potential change is, of course, directly read upon the establishment of the dropping junction, and is almost invariably as large as a millivolt in neutral solutions.

994

WILLIAM T. RICHARDS

Such an arrangement makes unnecessary a collodion membrane' or the glass bulb device recommended by Bovie and Hughes*. KO difficulty from the diffusion of calomel into the hydrogen vessel was ever experienced. The system has the advantage of the old-fashioned cotton plug junction method of preventing intercirculation of the solutions under measurement, without its intolerable disadvantage of obscuring the junction effect. The hydrogen electrode used below had the form which is illustrated in Fig. I . A ground-glass stopper F was fitted to A which supported the spiral of platinized platinum wire E which served as electrode, and provided an escaping passage for the bubbled hydrogen gas through G and the stopcock M . The platinum wire which served as electrode was made continuous from F to L , being sealed in the glass tube in both places; electrical connection was made by copper wire wrapped around an end left exposed a t L. This apparatus was, however, modified for use with the oxygen electrode, G both because of the large surface of electrode used, and because a quick change from one solution to another in A was often necessary. Hence duplicate bridge systems of the type illustrated in Fig. I were constructed which might be alternately connected with B by a ground joint a t K', the latter being held in place during measurement by rubber bands. The oxygen electrode vessels were of the type illustrated in Fig. 2 where A represents a two-ounce wide-mouth bottle. B a four-hole rubFIQ.2 ber stopper (the fourth hole containing A convenient form of an escaping jet for bubbled gas is not oxygen-gas cell. shown in the diagram), C a centrally located tube connecting the electrode E to the potentiometer by a mercury column contained in the rubber tube G which fits C closely, F a connection to the stopcock C in Fig. I , and a bubbling jet for maintaining oxygen flow. In this way great flexibility was secured and the advantages of the flowing junction a t the same time retained. A four-finger mercury-toluene regulator was used to control the thermoThe oxygen used in the cell chain was stat temperature within 0.1' of 25.0'. taken from a compressed cylinder, needing apparently no further purification. 1 2

Fales and Stammelman: J. Am. Chem. SOC.,45,1271(1923). Bovie and Hughes: J. Am. Chem. SOC.,45,1904 (1923).


995

Hydrogen was obtained from a battery of four amalgam generators, being purified by Gay-Lusaac towers containing 1/10 normal sodium hydroxide and distilled water in the order named. The calomel electrodes in the cell chains were carefully prepared although, in the work below, their constancy and not their absolute value is alone of significance. The potentiometric system was of the usual type, with a sensibility of greater than 0.1m.v. 3. Constant Oxygen Electrode Potentials The first step in the standardization of the oxygen electrode is to obtain potentials which, although not reversible in the thermodynamic sense, are at least stable and dependable. Table I indicates that this was accomplished. A single typical determination with a nitrate solution is also appended for comparison. Table I illustrates three effects. i. That in nitrate solutions there is a perpetual gradual fall in potential which approaches apparently no asymptotic value, and contains none of the “resting points” observed by Lorenzl under similar circumstances with a bright platinum electrode. ii. That constant potentials, of sufficient stnbility to excuse the conclusion of pioneer workers with the electrode that a reproducible value had been obtained, may readily be established in borate solutions. The existence of a constant but not reversible potential of this kind raises an interesting theoretical point which will be dealt with in Section 6 .

TABLE I Potential of the Oxygen-Hydrogen Chain a t 2 5.0°C in Various Solutions KO, Composition I 2 3 4 j 26 days I KaN03,0.5N 1,071 1,047 1.027 1.019 - 0.965 11 H3B03, 0.05hT 1.089 1 . 1 0 2 1.102 I. I02 - H3B03, 0.s&7 1.102 1.101 1.101 1.101 __ 111 H3B03, 0.05x 1.089 1.089 1.093 1.095 1.095 H3B03, o.gN 1.087 1.088 1.088 1.088 -IV H3B03, o . o ~ N 0.981 0.999 1.001 1.001 1.001 H3B03, 0.5x 0.975 0.995 0.999 0.999 0.999 -

...

That, in borate solutions, the average of all values for the concentrated and ail for the dilute solutions gives a difference of - .003 volt for the concentrated. This effect is fully investigated and discussed in Section 5 . 111.

4. Variation of Potential with Oxygen Pressure ‘ I t has been mentioned in Section I, iii, that the potential of the oxygen electrode is not independent of the partial pressure of the oxygen gas surrounding it. This observation, because of the indication it gives of the nature of the oxygen electrode potential as discussed in Section 6, has been confirmed by a series of measurements summarized in Table 11. 1 Lorenr and Spielmann: 2. Elektrochemie, 15,293,349; Lorenx: 661; Lorenr and Lauber: 157, 206 (1909);Spielmann: Trans. Fa-aday Soc., 5, 88 (1910).

996

WILLIAM T. RICHARDS

These measurements were obtained with the flowing junction apparatus and carefully seasoned oxygen electrodes which had been brought to constancy of potential with a constant oxygen flow a t one atmosphere in a .os N borax solution. At o time an equal flow of nitrogen gas (100 cc. per minute, enough to give vigorous stirring) was turned on by a set of stopcocks which automatically discontinued the oxygen flow : thus the oxygen was gradually swept from the solution with a corresponding fall in potential until the oxygen partial pressure became zero and the nitrogen one atmosphere. At the indicated time the nitrogen flow was discontinued and the oxygen resumed at the same rate of flow; this caused the potential to rise to its former value. The process was quantitatively repeated many times, two characteristic readings only being reproduced. Since, within limit of error of measurement, these two series are congruent, except in the absence of oxygen, they have been indicated graphically in Fig. 3 as a type curve, points from both of the series being set down without distinction.

TABLEI1 The Potential of the Oxygen Electrode in the Presence and Absence of Oxygen Gas. Determination I Time (mine) Hn-0% Potential (volts)

I5

0.979 0.970 0.961 0.938 0.898 0.880

23

0.851

0

I 2

5 I2

38 0.811 53 0.791 Oxygen turned on 0 I 2

3 5 IO

15 20

Determination I1 Time (mina) H1-02 Potential (volts) 0

I 2

5 I2

I7

0.979 0,970 0.962 0.938 0.900 0.878 0.838

28 56 0.801 68 0.798 Oxygen turned on 0 0.797 2 0.895 0.937 3 4 0.947 6 0.963 8 0.969 0.978 15 0.980 27

At the first o time nitrogen was turned on, and oxygen simultaneously discontinued. At the second o time the oxygen flow was resumed, and the nitrogen simultaneously discontinued. They indicate that the oxygen electrode, suitably prepared and standardized, may be used as an analytical instrument in determining the partial pressures of oxygen in a gaseous mixture, an observation which may have a certain application in biology. The effect of increasing the oxygen pressure


997

above one atmosphere was also investigated, but since a paper by Tammann and Rungel has recently appeared treating this matter much more fully the results will not be published. The type of experiment detailed above does not seem to have been attempted by these authors, who were interested chiefly in an attempt to secure the reversible potential of oxygen ions by raising the pressure and temperature. It should be noted that, at the moment before turning on oxygen gm, although its partial pressure is effectively zero, and although nitrogen is not made electromotively active by platinum, the potential of the oxygen elec-

FIQ.3 The variation of the potential of the oxygen electrode with partial Feeawe of oxygen gas. At A a current of nitrogen o about IOOOC per minute is turned on, the oxygen Bow being simultaneously discontinued. At B all oxygen, with the exception of insignificant traces, has been expelled from the solution. At C the flow of oxygen is resumed and the nitrogen discontinued. At D the potential of the oxygen electrode is constant at !ts initial value, and the cycle is repeated.

trode has by no means disappeared but is both finite and constant. This effect may safely be attributed to a n irreversible oxidation of the electrode surface. The remaining aspects of the phenomenon are considered in Section 6.

5. The Quantitative Measurement of Hydroxyl Ion Activity Using the apparatus and experience gained in the preceding sections it was found possible to develop a method of measuring indirectly hydroxyl (and hence hydrogen) ion activity by the oxygen electrode with a precision of a few millivolts. It is possible that Barendrecht* has devised a somewhat similar method but, from his published results, he cannot have worked it out in anything like the detail given below. The method is described at some length in the hope that bio-chemists may find it useful in dealing with oxyhaemoglobin solutions. 1

G, Tammann and F. Runge: 2. anorg. allgem. Chem., 156, 85 (1927).

* Barendrecht: Kon. Akad. Wet. Amsterdam, 22, 2

(1919).

998

WILLIAM T. RICHARDS

I n order to establish the method the following procedure was adopted. Two modifications A and B of the same solution were adjusted to different hydrogen ion activities by the suitable addition of small quantities of acid. Their hydrogen electrode potentials were then carefully measured at 25' in a thermostat. An oxygen electrode was seasoned in solution A , and then removed from the cell (Fig. 2 ) , rinsed with a previously oxygen-saturated portion of solution B, and placed in another portion of B which had also been oxygen saturated and brought to the temperature of the thermostat. Measurements of potential change with time were then made a t suitable intervals for 2 0 minutes, a t the end of which period the potential had almost invariably ceased significant change. The electrode was then removed, rinsed with solution A saturated with oxygen, and placed again in the cell chain in a warmed and oxygen-saturated portion of A , and the potential again observed for twenty minutes. The process of rinsing, transference, and timed measurement was repeated with both E and A a second time. Two sets of values were thus obtained for the oxygen electrode potential difference of the two solutions, one by the subtraction of the first resting potential in A from the average of the first and second resting potentials in B , and the other by the subtraction from the second resting potential in B of the average of the first and second resting potentials in A . Values obtained in this way are independent of any drift upward or downward resulting from irreversible change of the potential of the oxygen electrode. They express, with a degree of accuracy which is at once an indication of the reliability of the process, the difference in hydroxyl ion activity of the two solutions A and B. Comparison of the values obtained shows that the increase in oxygen electrode potential is almost exactly equal to the decrease in hydrogen electrode potential in the more acid of the two solutions. On the awumption, wrely justifiable, that [H+][OH-] = k in any given solution the difference of hydroxyl ion activity between the two solutions may thus be quantitatively measured. Borate solutions were chosen for most of these tests for obvious reasons, the solutions being 0.05 normal in borate as this concentration is as small as is compatible with the minimum conductivity necessary for the obtainment of accurate potentiometer readings. I n order to test the behavior of the electrode in more concentrated solutions a 6 N solution of ammonium sulfate was similarly measured. Somewhat acid solutions were preferably used for two reasons. I n the first place, i t is much easier, in slightly acid solutions, to obtain invariable hydrogen electrode potentials without careful exclusion of atmospheric carbon dioxide; in the second place, as Furman' has pointed out, the oxygen electrode shows greater potential stability in acid solution. A summary of the results obtained is given in Table 111. All the measurements are at 25.0' and all potentials are against a normal calomel electrode except, of course, those in columns 6 and 7 . I n the headings of the columns of this table the symbol T has its usual significance, potential. The table 'Furman: J. Am, Chem. SOC.,44,2685 (1922);Trans. Am. Electrochem. SOC., 43, 79 (1923).

999


contains all information necessary to calculate the desired quantity, but 1s in consequence exceedingly condensed. The notation may be expounded by taking a typical case where solutions A and B are under measurement. The electrode, seasoned in A, is transferred to B and observed for twenty minutes. Its potential after two minutes is the first entry in column I for the series of measurements in this particular solution. After twenty-one minutes, its potential having become virtually constant, a reading with flowing junction is made and entered in column 5 . Solution A is then similarly measured, its oxygen electrode potential after two minutes being the third entry in column I. The second measurement of B after nineteen minutes goes in the second line of column 3, and the second of A on the fourth line. Thus in any given determination from column I through column 5 the actual order of the oxygen electrode measurements is first, third, second, fourth, of the values given. The interrupted line subdividing horizontally each determination shows which of the columns gives results applicable to only one of the two solutions, and which gives values common to both. Column 8 gives the alkalinity of the solution.

TABLE I11 Measurements of Hydroxyl Ion Activity with the Oxygen Electrode. Column No. I 2 3 4 5 6 7 8 Change d a0 with Time, in v. and mins. Av. ?ro Av. rrH Detn. No. Composition z 3 19 20 21 flow differ. M e r . rH I Boric acid0.581 0.580 0.574 0.574 0.5745 Borax, o.osN 0.572 .571 .567.567 0.053 0.050 0.4943 .622 .623 .624 .624 .6243 Ditto. .61z .613 .615 .615 .6163 2

Ditto. Ditto. 3

.496 .496 .496 .496 e4961 .491 490 s o 9 .199 .200 .206 .206 . 2 0 5 7 .197 ,198 ,211 .ZII .ZIIO

Sodium Phosphate, o.ogN Boric acidBorax, o.ogN

0.285

0.282

0.428 0.703

.170 .17o .166 .166 .1667

.154.156.189 ,190.1933 0.102 0.111 0.711 -

4 Ammonium Sulfate, 5.9N

Ditto.

.061 .062 .o85 .075 .076 ,082

-

.224

.227

.085 .os2

.0855 .o811

.227

.2278

.220 .z21 .224 .224 .2258 0.233 - - -

.454 ,454 .459 .459 $4591 ,449 a450 -457 .4575 *4575

0.822

0.241 0.624 0.383

The results given in Table I11 show from a comparison of columns 6 and 7 that in acid solution there is a close quantitative correspondence, both in

dilute and concentrated solutions, between the difference of the oxygen

1000

WILLIAM T. RICHARDS

electrode potentials and the difference of the hydrogen potentials of two solutions of different hydrogen ion activity. This indicates that the oxygen electrode in spite of its irreversibility is a reliable measure of relative hydroxyl ion activity. The procedure in determining the hydroxyl ion activity of an unknown solution is precisely similar except that, of course, the checking hydrogen electrode potential of this solution will not be obtainable. The comparison solution should be adjusted to nearly the same pH, and have approximately the same vapor pressure as the unknown solution. I t should preferably contain sodium tetraborate and boric acid in appropriate proportions. The oxygen electrode is seasoned to constancy in this solution, and the procedure above outlined followed. The method of computation of the pH from these results is so obvious and well known that it will not be discussed.

It has been obsemed in Section 3, iii, that a concentrated solution of borax and boric acid gives a value for the hydrogen-oxygen-chain which is less than that in the dilute solution. Smale' found a similar effect, although the solutions in which his measurements were made were never very concentrated. The method just described obviously gives a means of studying this effect rapidly and with considerable accuracy, since it makes long waits for the stable potential to be attained unnecessary and, by eliminating the effect of irreversible change in the potential, brings a greater number of anions into the realm of possible investigation. Only certain concentrated solutions may, however, be so investigated. For example they must not contain as an impurity elements such as lead, arsenic, or antimony, which are capable of poisoning the platinum electrode surface; probably all metals less electropositive than hydrogen should for this reason be excluded. Because of their oxidizability, solutions of the alkali halides are obviously unsuitable since, as was pointed out by Lewis: the halide ion would be oxidized until its potential against free halogen was equal to that of the oxygen electrode. Nitrates, on the other hand, although reduced catalytically by hydrogen electrode, may be made to give satisfactory results by extrapolating to find the initial potential of the hydrogen electrode; a method which has been carefully justified both theoretically and practically. A similar reduction of considerably less importance is observed in concentrated solutions of sodium acetate; it may be dealt with in similar fmhion. Solutions of alkali perchlorates, were also found suitable if the sodium chloride normal electrode was substituted for the potassium. Sulfate solutions, after many preliminary tests, were found most suitable of all for this purpose, and were therefore given most attention. Table IV, while not as detailed as its immediate predecessor, gives a summary of the results obtained, and should be self-explanatory. Smale: Z. physik. Chem., 14, 577 (1894).

* Lewis:J. Am. Chem. Soc., 28, 158 (1906);Z.physik. Chem., 55,465(I@),


1001

TABLE IV The Effect of Increasing Electrolyte Concentration on the Potential of the Hydrogen-Oxygen Chain Hydrogen DitTerence Oxygen between Dilute Chain, and Concentrated, volts volts

Average Corrected Oxygen Potential

H drogen eLtrode Potential

0.208

0.719

,214

.72I

0.927 .935

0.008

.498

,510

.402 .406

.goo .916

0.016

.498 ,516

,431 .429

.920

505

.43I .494

.936 .957

.409 .413

.838

.SI7 6

,458 .344

.393 .587

.851 .931

7

.458

,383 .406

.841

Detn. NO.

I

Composition of Solution

2

3

4

'

.465 .429

5

*

8

502

-0.116

-0.038 9

*

I37

.282

6.

0.025

.945

'930

.g08

.931 .934

.815

.623 .624

.760

.896

.g06

0.023

0.092

0.080

0.067

0.081 0.146

Discussion

The theory of the oxygen electrode at present accepted is that its irreversibility is due to a coating of oxide of variable and uncertain composition. With this conclusion the present communication can find no fault. Many writers believe, however, that the measured potential of the electrode is eutirely oxidic. This conception originated with Lorene,' was declared final by Schoch,2 and has apparently been wholeheartedly accepted in the more recent work by Foerster,s Grube,' and even Tammann and Runge.6 While 'Lorens and Spielmann: 2. Elektrochemie, 15 293, 349 (I ); Spielmann: Trans. Faraday Soc., 5, 88 (1910); Lorenz: Z. Elektrochekie, 15, 661; E r e n z and Lauber: 157, 2 0 6 (w9). * Schoch: J. Phys. Chem., 14, 665 (~910). Foerater: 2. physik. Chem., 69, 236 (1909). 'Grube: 2.Elektrochemie, 16, 621 (1910). 6 G. Tammann and F. Runge: 2.anorg. allgem. Chem., 156, 85 (1927).

I002

WILLIAM T. RICHARDS

the measurements set forth above are by no means a complete refutation of this hypothesis, they furnish evidence which, in conjunction with the work of others, makes it seem highly probable that the electromotive activity of oxygen gas itself, a t the platinized surface of the electrode, plays a considerable part in the total effect. A number of phenomena emphasized above are difficult or even impossible to explain on the totally oxidic theory. The first of these is the superpotential, or high residual voltage after anodic treatment, of the electrode mentioned in Section I , ii. If this is considered due to a higher unstable oxide of platinum it must, owing to the character of its over-voltage phenomena, be soluble in the electrode in all proportions and decompose according to the law of a monomolecular reaction, a combination of conditions which has few analogies in polyvalent oxidations. Moreover the oxide PtOa is necessary to account for the potential of 1.9volts observed by Foerster for the hydrogen-oxygen chain, since Grube has found the potential of P t 0 3 to be only 1 . 5 volts. Pt04by analogy with RUO3 and RuOl should be considerably more stable than PtOs, whereas it is generally agreed that the superpotential of the electrode is increasingly unstable with increasing magnitude, the function being, in fact, an exponential one. If, on the other hand, it is considered that oxygen gas dissolved in the electrode is responsible for its temporarily elevated potential no similar difficulty is encountered. Although never, apparently, specifically enunciated in this connection, Bose* has evidently assumed a similar hypothesis, and the measurements of Harding and Smith2 make such a conclusion almost incontestible, These authors, in a paper the importance of which from this point of view can hardly be overestimated, have proved that the electrical resistance of a palladium wire decreases suddenly on being made cathode in electrolysis due to the occlusion of hydrogen. An exactly similar decrease is observed when the wire is made anode, differing only in magnitude from that observed during cathodization, and followed, after the discontinuance of the electrolysing current, by a permanent increase of resistance due to oxide formation.' Hydrogen occluded under these conditions is generally accepted as being uncombined.* The extension to oxygen by analogy, and hence to the phenomenon of the superpotential seems a t present, therefore, incontestible in view of the feeble character of the opposing evidence. Here, therefore, is a case in which it seems wise to attribute to oxygen gas a t least a part of the total electromotive activity. It may perhaps be brought out, in passing, that no cogent reason is a t hand for the belief that oxygen and not hydroxyl ions are crowded into the electrode by the high current density; while unconventional this view is not entirely fantastic. Boee: 2.Elektrochemie, 7, 673 (1901);Zeit. anorg. Chem., 30,406 (1902).

* Cf.Smith andothem: J. Am. Chem. Soc., 38,2577 (1916);40,1508 (1918). * T.W.Richards and W. T. Richards: J. Am.Chem. Soc., 46,8g (1924).


roo3

It is difficult also to understand, without attributing electromotive activity to oxygen gas, the often observed phenomenon that an electrode which has not been made anode at once assumes, in the presence of vigorously bubbling oxygen, a potential comparable to that of seasoned anodized electrode under the same conditions. I n other words the superpotential and possible oxide formation produced by anodic treatment have no necessary part in the almost instantaneous production of a potential of like magnitude by platinum and oxygen gas alone. Moreover, the quantitative and immediate response of the electrode to variations in the partial pressure of oxygen gas brought out in Section 4 make necessary according to the oxidic theory the reversible formation of high oxides of platinum which Woehler,l working under much more favorable conditions, has shown to be formed only during many days. It may be mentioned in passing that an attempt was made in this connection to measure directly the ionization of oxygen gas due to its passage over a large quantity of platinized asbestos. A specially constructed and very sensitive radioactive electroscope was used for this purpose. While the ions so produced would have had to traverse a space of only 2 cm. in about one second not the slightest ionization was detected. The failure of this measurement, while disappointing, has obviously no disqualifying effect upon the argument set forth above. I n view of these considerations there seems no valid reason for denying a certain electromotive activity to oxygen gas. That oxide formation is also a factor is hardly contestible, especially in view of the measurements recorded in Section 4. We are then forced to picture the measured potential as a compromise between a low oxidic potential and the only partially attained 0 2 - 20‘ potential. Although theoretically the highest of the several possible potentials in a concentration-chain should alone be manifest, actually a number of such “compromise potentials” are known. For example the halfcells Fe/FeS04, Fe~(S04)a;Pt.Hn/FeS04, Fen(S04)s;and Fe/FeS04, HzSOl have been shown to exhibit a similar effect.2 The writer intends, in the near future, to discuss a number of these potentials from the point of view of the velocities of the concomitant reactions involved. Failing such quantitative justification the conception of the “compromise potential” must be accepted as the only solution, however disagreeable, for the situation in hand. This hypothesis of duality being granted only one further assumption is necessary to account for all apparently anomalous effects which are a t hand, namely, that with increasing oxidation of the blacked platinum surface its activating power for oxygen is diminished. This assumption cannot be substantiated by any direct evidence, but seems in itself highly probable, and is in accord with many observations. For example it was observed that an electrode which had given a stable potential in a solution gave, after having been treated electrolytically only l

Woehter: Ber., 36, 3475 (1903).

* T. W. Richards and W. T. Richards:

J. Am. Chem. SOC.,46,89 (1924).

I 004

WILLIAM T. RICHARDS

as cathode in boric acid, a potential considerably above its former value in the same solution. There can be very little question from the measurements of Harding and Smith that a film of oxide undergoes reduction upon cathodization. The rise in potential is simply accounted for by the increased electromotive activity of oxygen gas due t~ a removal of obstructing oxide from the platinum surface. Again, the constancy of potentials in borate solutions noted in Section 2 is made plausible by this explanation. Since both an oxide of platinum and electromotive oxygen gas may reasonably be supposed to be together responsible for the oxygen electrode potential, it does not seem excessively fantastic to suppose that a balance may be reached between the formation of oxide, which is irreversible and in the nature of a poisoning, and the activation of oxygen gas. Since both these processes are weak it might be expected that, with a solution of sufficiently inactive constituents, a metastable balance between the two would ultimately result which would be manifested by constancy of the potential. Any extremely stable solution should give similar results and, in fact, observations of a similar character were made in less investigated but comparable phosphate solutions, whereas in nitrate solutions, where the lowering of the potential due to oxide formation should be much greater, no constancy was observed. One further phenomenon is elucidated by this interpretation of the oxygen electrode potential, namely, the lowering of the value for the oxygen-hydrogen chain with increasing electrolyte concentration emphasized in Section 5 . Several hypotheses may at once be abandoned for this purpose. It is evident that neither a decrease in the ion product of water nor in the activity of dissolved oxygen can be responsible for this result, since both of these values are differently defined thermodynamically. Passivity and poisoning of the oxygen electrode manifest themselves irreversibly and in a very different manner from that observed, and both may be unequivocally discarded as explanations. It has, however, been found convenient to infer a considerable dependence of the dissociating power of platinum black for oxygen on the condition of the platinum surface. With increasing oxysalt concentration an increased oxidation of the surface, with a concomitant loss of dissociating power for oxygen, and a consequent temporary fall in potential is t o be expected. It is possible, as an alternate hypothesis to account for this effect, that the hydration of the ions in the concentrated solutions would alone be sufficient. Data are not a t hand to treat the matter conclusively.

It is hardly necessary, although agreeable, to acknowledge indebtedness t o Professor Theodore W. Richards for his interest in this investigation. To Professor F. G. Donnan of University College, London, the most cordial thanks are also extended for his hospitality under unusual circumstances. SllmmaY

The preparation of an oxygen-platinum electrode capable of giving reproducible results is discussed. A convenient cell chain having a flowing junction and requiring only 2. small volumes of liquid is described. I.


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3. It has been demonstrated that, under suitable conditions, oxygenelectrode potentials constant to a millivolt for many days may be obtained, although these are far below the thermodynamic value for the oxygen-hydrogen chain. 4. The potential of the oxygen electrode has been shown to vary quantitatively but not thermodynamically with the partial pressure of oxygen gas between o and I atmosphere. 5 . A method is described for measuring hydroxyl ion activities with the okygen electrode which is capable of considerable accuracy. 6. I n order to explain these and other phenomena it has been necessary to attribute electromotive activity to oxygen gas, and to consider the measured potential of the electrode a compromise between this and the oxidic potential which is currently believed to produce the total effect. It is hoped that these findings will prove useful in biochemical work. Princeton, Xew Jersey.

The Oxygen Electrode as a Quasi-Quantitative Instrument

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