Jan., 1959
CHANGE POTENTIAL ON SILVER ELECTRODE IN ALXALINE SOLUTION
107
THE EFFECT OF CONTINUOUSLY CHANGING POTENTIAL ON THE SILVER ELECTRODE rN ALKALINE SOLUTIONS BY THEDFORD P. DIRKSEAND DALEB. DE VRIES Department of Chemistry, Calvin College, Grand Rapids, Mich.igan Recebed Auoust IS, 1968
When a silver wire is oxidized by continuously increasing its otential in alkaline solutions, four reactions take lace. others are due to the formation of Ag20, of Ag8 and of oxygen. The first and second reactions are indistinguishable in solutions having a high p H . The results show that the formation of Ag20 is not due to the action of oxygen on silver, and the formation of Ago does not involve the HOz- ion.
It is suggested that the first is due to the formation of AgOH. !he
Introduction When silver is electrolytically oxidized in alkaline solutions, AgzO is formed, then, a t a higher voltage level Ago is produced. As the voltage rises still higher oxygen is evolved. There is doubt whether Agz03is formed in such a treatment. Most of the work that has been carried out in an attempt to understand the mechanism of these processes has been done by the use of a constant current technique. l--4 The work reported here was carried out in an effort t o shed more light on these reactions. This was done by applying a continuously increasing potential to the silver electrode. So far as could be determined, no such work on these reactions has been reported in the literature. Such results cannot be subjected to strict quantitative interpretations because of a variety of variables present, e.g., convection and changing electrode surface area. However, significant qualitative conclusions can be drawn from such data. Experimental The circuit employed in this work was similar to that in polarography. A Sargent Model 111 polarograph was used as the voltage control. It was equipped with a 0.5 r.p.m. motor to increase the voltage continuously and uniformly. The current, instead of passing through the galvanometer, was fed into a precision resistor and the voltage drop across this resistor was measured by a Brown recording potentiometer. Only relative current values were necessary in this work. An H type cell contained the electrodes. In one branch there was a saturated calomel electrode with a plug of agar and KNOa solution in the cross piece. The other electrode was a silver wire about 2 cm. long fitted in a polystyrene holder. This electrode was the anode in all theoruns. Measuraments were made a t room temperature and 2 . I n addition to temperature, the other variables studied were: agitation, potassium hydroxide concentration and dissolved oxygen. Since rather high concentrations of potassium hydroxide were used, corrections were made for junction potentials.5
Results On Fig. 1 are shown results that are typical of all those that were obtained. One of the first things noted was the presence of another oxidative stage in addition to that of the formation of Ag20, Ago and oxygen. This is indicated by peak a on Fig. 1. The formation of AgtO takes place a t peak b, of Ago at c, of oxygen at d. The presence of Ag20 and Ago at the peaks indicated was confirmed by e.m.f. (1) R. Luther and F. Pokorny, 2. anarg. allgem. Chem., 51, 290 ( 1908). (2) I. A. Denison, Trans. Electrochem. Soc., 90, 387 (1946). (3) Hiokling and D . Taylor, Disc. Faraday Sac., No. 1, 277
A.
(1947). (4) P. Jones, H. Thirsk and W. F. K. WynneJones, Trans. Faraday Soe., 62, 1003 (19513). (5) T.P. Dirkse, 2. p h y s i k . Chem., N . F . , 6 , 1 (1955).
measurements. The evolution of oxygen was observed visually. For each reaction, the rise in current indicates an increase in the rate of the reaction and a decrease in the current represents a decrease in the rate of the electrode reaction. For reactions a and b the decrease in rate of reaction is rather gradual and this may correspond to the gradual shutting off of the reaction due to the formation of a film of the reaction product. With reaction c the decrease in rate of reaction was very sudden. This was accompanied also by the formation of a visible film of gas (oxygen) on the surface of the electrode. As soon as this film was broken and bubbles of gas were evolved from the electrode surface, the current rose again, peak d. I n some of the runs a vigorous stream of air was passed through the solution during passage of current. However, in general, the nature of the curves was unaffected by such agitation. The effect of dissolved oxygen was observed by carrying out a run after purified nitrogen had been passed through the electrolyte for some time. Then a similar run was made after air had been bubbled through the electrolyte. I t appeared that none of these phenomena was affected by the presence or absence of dissolved oxygen, see Table I. There is a little variation in the potential at which oxygen is evolved but these potentials are a t best hardly reproducible in a given system. Thus it appears that none of the reactions noted involves dissolved oxygen. This indicates, e.g., that in the formation of Ag20 it is not the action of dissolved or liberated oxygen on the silver that is responsible. TABLE I EFFECTOF DISSOLVED OXYGEN ON POTENTIALS FOR OXIDATION OF SILVER IN KOH SOLUTIONS AT ROOM TEMPERATURE mKOH
Soln. satd. with
1.0 1.0 1.6 1.6 4.9 4.9 13.2 13.2
Air Nitrogen Air Nitrogen Air Nitrogen Air Nitrogen
-Minimum potential for---FormaFormaFormaFormation of tion of tion of tion of AgOH Agio Ago OP
0.08 .08 .12 .12 .09 .08
...
...
0.16 .16 .18 .18 .15 .15 .16 .15
0.43 -43 .46 .46 .41 .41 .42 .42
0.68 .66 .69 .68 .61 .65 .57 .60
The effect of KOH concentration is shown on Fig. 2. The point of interest here is that the lines representing points a’, b’ and c’ on Fig. 1 are parallel. Except at the highest concentration these lines are also approximately parallel to that €or oxy-
THEDFORD P. DIRKSE AND DALE B. DE VRIES
108
I
I
I
’ I
I
C
250 -
200(II
5
0 190E 100-
50-
a‘ 01
I
I
I
02
03
I
0 4 05 E vs S C E
I
I
I
06
07
08
Fig. 1.-Variation of current with voltage for 8 silver electrode in 1.7% KOH at.2’; no agitation; voltage varied at a rate of 100 mv. per minute.
I
I
0.2-
1
I
I
1
I
I
0 ‘
-
+
-. -
b’
0‘
I
2
I
4
I
6
I
8
I
10
I
i2
~ K O H .
Fig. 2.-Effect of KOH concentration on the anodic processes of silver a t room temperature. The letters refer to Fig. 1.
gen evolution. Thus, all of these reactions have the same dependence on hydroxyl ion concentration. This dependence can serve as a means for indicating the mechanisms for these reactions. However, it is slight except a t the lower pH values. The reaction for the evolution of oxygen does involve hydroxyl ions in these alkaline solutions. It should be noted that the points on Fig. 2 represent the minimum voltage a t which these processes or reactions begin under these conditions. Since these reactions appear to have the same dependence on the hydroxyl ion concentration, the mechanisms likely also are similar. Reactions I-IV show such similarity and (11),(111) and (IV) are generally accepted as the equilibrium reactions. If these were the rate controlling electrochemical reactions one might expect a greater hydroxyl ion 2Ag 2Ag Ag20
+ 2 0 H - +2AgOH + 2e + 20H- ---+ Ag20 + HzO + 2e + 20H- +2Ag0 + H 2 0 + 2e 2 0 H - + 1/*O2 + HzO + 2e
The reaction at peak a appears to be something hitherto unreported. It was accompanied in each case by a slight discoloration of the silver electrode. It disappeared a t the highest KOH ‘concentration used, or it occurred at a potential so close to that of the formation of Ag20 that it was not discernible. T o investigate this further, the applied voltage was held at that of peak a for some 30 minutes in 5 m KOH, and then the potential of the silver electrode was measured against the S.C.E. with a potentiometer. The results are shown in Table 11, together with the potentials of the other peaks. Whatever is formed a t peak a does give rise to a distinct potential value. TABLE I1 POTENTIAL OF VARIOUE SILVERPRODUCTS I N ~ M K O H ~J substance
Ag wire Peak a Ag@ Ago
o o
(1) (11) (111) (IV)
dependence than is actually observed. However, in the most dilute solutions there is such dependence and as the concentration of KOH increases the number of hydroxyl ions is so large that they are no longer a limiting factor and an increase in concentration does not materially add to the number ‘Lavailablel’ for reaction. Arguing by analogy, ie., similarity in variation of voltage with hydroxyl ion concentration, reaction I likely represents the reaction taking place a t a in Fig. 1.
Vol. 63
E
UB.
S.C.E., V.
+O. 028 .052 .067 .311
+ + +
An attempt was also made to determine the Xray diffraction pattern of this substance. A larger (1 X 4 cm.) sheet silver electrode was held a t a voltage of +0.32 v. vs. a Hg-HgO electrode in 23% KOH. This voltage corresponded to that of peak a. The Hg-HgO electrode was used in order to eliminate the anions other than OH-, and to eliminate the resistance due to the fritted glass disk and agar plug. Because of the larger silver electrode a larger cell was used and the electrodes were farther apart. Even with this arrangement the current passed was only a few microamps. The silver electrode was held a t the potential mentioned for 2 days. During this time the electrode became discolored and the current decreased from 2 microamps to 0. The electrode then was rinsed quickly in distilled water, dried with filter paper and X-rayed. The results, however, were the lines for silver and only silver. In fact, the silver pattern was a better one than that obtained before this treatment. The discolored area also had a smoother appearance under the microscope than did the untreated area. It appears that the “high” points on the silver surface were attacked to form Ag+ ions or AgOH, and the AgOH dissolved in the electrolyte. The discoloration of the electrode, then, was not due so much to a deposit of AgOH on the electrode as t o the “smoothing” of the electrode surface by this action. On the basis of the evidence it seems likely that peak a corresponds to reaction I. The similarity in dependence on hydroxyl ion concentration between this reaction and reactions I1 and I11 is thus accounted for. The potential set up by the reaction at peak a is then the potential between the silver in the electrode and the dissolved AgOH in the electrolyte, whereas that at peak b is the potential between the Ag and the Ag20 both in the electrode. The AgOH likely exists in solution as Ag(0H)z- or Ago-. As soon as the film of electrolyte next to the surface of the electrode has a certain amount of this species dissolved in it the reaction slows down and the current decreases.
.
CHANGE POTENTIAL ON SILVER ELECTRODE IN ALKALINE SOLUTION
Jan., 1959
Since KOH solutions absorb COz very readily it is conceivable that a film of AgzCOs could be formed on the electrode. However, it is unlikely that this happens at peak a because the voltage needed for the formation of such a film in basic solutions is about 130 mv. greater than that required for the formation of Ag20.6 Furthermore, the presence of such a film probably would have been detected in the X-ray diffraction patterns. If peak a were due solely to the charging of the double layer it is difficult to account for the discoloration of the electrode. The discussion above, in terms of reactions 1-111, has dealt with equilibrium conditions and the potentials for these conditions. The mechanisms of these processes, however, need not correspond to these reactions. The fact that the hydroxyl ion dependence for each of these processes is similar to the others indicates a similarity of mechanisms. The simplest mechanism would involve a single hydroxyl ion. As the potential of the silver becomes more positive it attracts the hydroxyl ions to it and may form AgOH or Ago- and H+, the latter immediately combining with another hydroxyl ion to form H20. As the potential of the silver becomes still more positive, not only are hydroxyl ions attracted to the silver but the charge on the silver is sufficient to separate the oxygen from the hydrogen in the hydroxyl ion and the H+ can then unite Ag
Ag
+OH-+
AR
)O+H+
Ag
(V)
I
12
I
13
I 14 PH.
I
15
109
I
16
Fig. 3.-Potential-pH diagram for silver and its oxides at 25’: 0,voltages at which formation of AgZO begins; 0 , voltages at which formation of Ago begins; (3, voltages at which oxygen is evolved.
with an OH- ion to form H20. This reaction continues until the surface is covered with AgzO. The The temperature effect over the range studied is addition of O= ions into the silver lattice also re- small, but most pronounced a t lower KOH conduces the positive charge on the surface. Conse- centrations. This may be due to the greater visquently, in order to continue the introduction of cosity of such solutions a t the lower temperature. oxygen into the lattice the potential of the electrode As the concentration increases, the viscosity also (or the surface) must be increased still further. increases, but no appreciable change in voltage is This can be done by, in effect, removing a second at the higher necessary to start these reactions electron from each silver atom. This then, by the concentrations. This may be due to the fact that same type of action as described above, introduces the effect of higher viscosity is offset by the inanother O= ion into the lattice forming a surface creased number of hydroxyl ions. Since only a layer of Ago. As the potential becomes increassurface film of the oxides is formed in this method ingly positive the reaction continues but since no more oxygen atoms or ions can be accommodated in there are sufficient hydroxyl ions a t the surface for the electrode lattice under these conditions, the reaction a t these higher KOH concentrations and liberated oxygen is evolved. It is possible also little migration of these ions is necessary. Further information regarding the mechanisms that when the potential of the electrode surface is sufficiently high the HO2- ion may be formed and of these reactions can be obtained by reference to Fig. 3, which is a modification of data published decompose to produce the oxygen. The formation of a layer of Ag20 also increases earlier.* The lines plotted here represent reversithe electrical resistance of the electrode surface.’ ble potentials for the reactions indicated. The pH As this layer is built up theacurrent passed through values were determined by using the data of Akerlof and Bender.g The data for the H02- ion were the system a t a given potential decrease. While the effect of agitation was very slight, it obtained from ref. 6 , p. 45. The two lines for the did modify the voltage necessary to begin all these evolution of O2 represent the variation of potential reactions very slightly, but only in the most dilute as the pressure of 0 2 varies from 0.2 to 1.0 atm. At the lower pH values the formation of Ag20 solutions used. I n those solutions, vthere the concentration of the hydroxyl ions is so small, the agi- from Ag takes place a t potentials lower than that tation may disturb the migration of these ions to the for the formation of oxygen. Hence the producelectrode, necessitating a slightly higher voltage for tion of AgzO here is not by means of the action of causing these reactions to begin. This was the ef- oxygen on silver. At higher pH values the formafect observed. (8) P. Delehay, M. Pourbaix and P. Van Rysselborghe, J . Electro(6) W. M. Latimer, “Oxidation Potentials.” Prentioc-Hall, Inc.. New York, N. Y.,1952,2nd Ed., p. 181. (7) M. Le Blanc and H.Srchse, Physik. Z.,92, 887 (1931).
chem. Soc., 98, 65 (1951). (9) G. C. Akerlof and P. Bender, J . Am. Chem. SOC.,TO, 2366 (1948).
110
LEVERNEP. FERNANDEZ AND LOREN G. HEPLER
tion of AgzO does take place a t potentials above that for the formation of oxygen but there is no evidence that the mechanism here is different from that a t the lower pH values. Similar reasoning shows that the formation of Ago from AgzO does not involve the HOz- ion.
Vol. 63
However, the mechanism of the formation of oxygen may involve this ion, but this is a subject that will be dealt with in a later paper. Acknowledgment.-The authors wish t o express appreciation t o the Research Corporation for a . grant which made this work possible.
THERMODYNAMICS OF AQUEOUS BENZOIC ACID AND THE ENTROPY OF AQUEOUS BENZOATE ION BY LEVERNEP. FERNANDEZ AND LORENG. HEPLER Contribution from Cobb Chemical Laboratory, University of Virginia, Charlotlesville, V a . Received August l.&1968
Calorimetric investigations of the heats of solution of sodium benzoate in water and dilute acid have been carried out a t 10,25 and 40". Heats of ionization of aqueous benzoic acid a t each of these temperatures have been calculated from the results of these calorimetric experiments. The heats of ionization have been used to calculate the change in heat. capacity on ionization. Results of earlier investigations of the ionization constant of aqueous benzoic acid have been combined m t h our data t o give the entropy of ionization and thereby have made it possible for us t o derive a thermodynamic equation for the ionization constant of aqueous benzoic acid as a function of temperature. From solubility data in the literature we have calculated the free energy, heat and entropy of solution of benzoic acid. These data have been used with the above mentioned data to calculate the free energies and heats of formation and entropies of aqueous benzoic acid and benzoate ion and the heat of formation of crystalline sodium benzoate (all in the standard state a t 298'K.).
Jones and Parton' have reported thermodynamic ionization constants of aqueous benzoic acid a t several temperatures and have calculated heats of ionization ( f150 cal./mole) a t several temperatures. Heats of ionization of aqueous benzoic acid also have been determined calorimetrically by Cottrell, et aLI2 a t 10, 20 and 30". These heats are in only fair agreement with those reported by Jones and Part0n.l The change in heat capacity on ionization (ACPo)as calculated from these heats2 is not in agreement with ACPogiven by Jones and Parton. Jones and Parton have also calculated (with incorrect results) from solubility data in the literature the entropy of solution of benzoic acid and have combined this entropy of solution with their entropy of ionization and an old value for the entropy of crystalline benzoic acid to obtain the entropy of aqueous benzoate ion. We have determined calorimetrically the heat of ionization of aqueous benzoic acid at 10, 25 and 40". These heats have been used with data from the literature to calculate several free energies and heats of formation and entropies and to derive a thermodynamic equation for the ionization constant of aqueous benzoic acid in the temperature range 10-40". Experimental The solution calorimet,er used in this investigation has been described in detai1.3J Because of the small heats with which this investigation was concerned, particular care was used in determining the heat effects associated with stopping the calorimeter stirrer for a few seconds (known +0.5 second) and the simultaneous breaking of the sample bulb attached to the stirre?. (1) A. V. Jones and H. N . Paston, Tmnlr. F a r a d a y Soc.. 48, 8 (1952). (3 T. L. Cottrell, C. W. Drake, D. L. Levi, K. S . Tully and J. A. Wolfenden, J. Chem. Soc., 10113(1948). (3) C. N. Muldrow. Jr., and L. G . Hepler, J . Am. Chem. SOL.,19, 4045 (1957). (4) R. L. Graham and L. G. Heyler, ibid., 78, 4846 (1956).
Sodium benzoate, Baker and Adamson U.S.P. grade, was recrystallized at least twice from ethanol-water mixtures. Recrystallized sodium benzoate was dried a t 150' and stored in a desiccator. Analysis of this sodium benzoate was by titration with perchloric acid in glacial acetic acid with methyl violet used as indicator.6 Perchloric acid for this purpose was standardized by titration against potassium acid phthalate in glacial acetic acid.6 Various samples of sodium benzoate were found to require for neutralization 99.5-99.7% of the calculated amount of HClOl. Hydrochloric acid solutions were prepared and standardized by common methods. All of the calorimetric exoeriments were carried out with 950 ml. of solution in the cdorimeter vessel at 10.0 f 0.4', 25.0 f 0.3"01-40.0 & 0.5'.
Results and Calculations Heats of reactions 1 and 2 at 10, 25 and 40" were measured over a range of concentrations in dilute solution.
+ Na+(aq) (1) + Na+(aq) (2)
C6H6COONa(c) = CBHd3OO-(aq) CeHd3OONrt(c) H+(aq) = CeI1&OOH(aq)
+
AH1 AH2
Heat and concentration data for the 25" experiments are given in Tables I and 11. The standard heats of reaction 1, AHI0,a t 10, 25 and 40" were determined graphically (AH1plotted us. m''2) by extrapolating the experimental heats to zero concentration dissolved sodium benzoate. I n making these extrapolations we were guided by experimental heats of dilution of other electrolytes and by the Debye-Huckel limiting law as well as by our own data. The standard heats of reaction 2, AHzO, a t 10, 25 and 40" also were determined graphically by extrapolating t o zero concentration. All of these reactions were carried out in such dilute solutions that all heats of dilution were small. The values we have found for AHlo and AHd' a t 10, 25 and 40" are given in Table 111. Uncertainties listed in Table 111are our estimates of the un(5) r. Kashima, BuEE. Natl. Hyg. Lab. (Tokyo), 72, 145 (1954): C. A . , 49, 7186g (1955). (6) W. Seaman and E. Allen, A n d . Chem., 93,692 (1951).
