Electrochemical Reduction of Ozone in Acidic Media D. C. Johnson, D. T. Napp, and Stanley Bruckenstein Chemistry Department, University of Minnesota, Minneapolis, Minn. The stoichiometry of the electrochemical reduction of ozone at a platinum electrode in 1M sulfuric and 0.1M perchloric acid media was studied by classical chemical and electrochemical methods at a rotating gold ring-platinum disk electrode. The reaction was 2H+ 2e + O2 HzO. The established to be O3 diffusion coefficient was calculated from the Levich equation to be 1.53 =k 0.12 X sq cm per second in 1M sulfuric acid and 1.39 =k 0.05 X 10-5 sq cm per second in 0.1M perchloric acid. These values were verified using chronopotentiometry and linear scan voltammetry. From a wave analysis of a current-potential curve obtained for a potential scan from to -, a n o = 0.32, ana = 0.57. while for a potential scan from - to In 0.1M perchloric acid, the corresponding values were 0.38 and 0.59.
+
+
+
+ +,
IN CONNECTION WITH STUDIES involving the anodic behavior of platinum in acid media, it became necessary to establish the half-wave potential for the reduction of ozone, and the stoichiometry for all species involved in its reduction at a platinum electrode. In addition it seemed desirable to investigate the use of the rotating disk electrode for the quantitative determination of ozone in sulfuric and perchloric acid media, and to redetermine the molar absorptivity coefficient of ozone, for which three values are given in the literature (I-3).
The electrochemical reduction of ozone O3
+ mH+ + ne
+
products
(1)
a t various inert electrode materials has been studied in acidic media (4-7). The net reaction was proposed to be 0 3
+ 2H+ + 2e
+0 2
+ HzO
(2)
on the basis of various indirect considerations. However, no study was found giving explicit experimental verification of this proposed reaction. The convective-difiusion limiting current for the electrochemical reduction of ozone has been reported to be proportional t o the ozone concentration (8). This study involved a rotating micro platinum wire electrode and the diffusion coefficient of ozone could not be determined. EXPERIMENTAL
Chemicals. Supporting electrolyte solutions were prepared from Baker Analyzed reagent concentrated sulfuric acid and G. Frederick Smith 70% ,perchloric acid. The water used for all solutions was deionized, followed by three (1) R. S. Ingols, R. H. Fetner, and W. H. Eberhardt, Aduan. Chem. Ser., No. 21, 102 (1959). (2) M. L. Kilpatrick, C. C. Herrick, and M. Kilpatrick, J . Am. Chem. SOC.,78, 1784 (1956). (3) M. G. Alder and G. R. Hill, Zbid., 72, 1884 (1950). (4) N. D. Tomashov and A. Z . Valiulina, Zh. Fiz. Khim., 26, 417 (1952). ( 5 ) K. I. Nosova, A. A. Rakov, and V. I. Veselovskii, Ibid., 31, 349 (1959). (6) A . A. Rakov and V. I. Veselovskii, Zbid., 35, 2297 (1961). (7) M. M. Fliskii, Elekfrokhim., 1, 1377 (1965). (8) K. I. Nosova, A. A. Rakov, and V. I. Veselovskii, Zhur. Fiz. Khim., 33, 349 (1959). 482
ANALYTICAL CHEMISTRY
distillations, with the second being from an alkaline permanganate solution. All other chemicals used were either Baker Analyzed reagent or Baker and Adamson reagent chemicals. Ozone was generated in a Welsbach Corp. Model T-23 ozone generator (maximum efficiency about 15%) using National Gas oxygen. The pressure of ozone in the discharge tube was maintained a t 8 psi and that in the line leading from the ozonator was 2 psi. The ozone-oxygen mixture coming from the discharge tube was dispersed in the electrolysis cell containing 500 ml of supporting electrolyte solution. The concentration of ozone in the cell was varied by adjusting the voltage across the discharge tube. Instrumentation used to study the electrochemical reduction of ozone has been described (9, IO). All voltammetric and chronopotentiometric data were recorded on a Houston Omnigraphic HR-97 X-Y recorder which was calibrated using a standard voltage reference. All potentials were measured and are reported in volts with respect to the SCE. All work was performed a t 25’ C. Electrodes. Geometric parameters of the ring-disk electrodes used in this research are listed in Table I, where i3 = (R3!Rd3 - ( R z ! R ~ ) ~
R1 is the radius of the disk and Rz and R I are the inner and outer radius of the ring, respectively. The collection efficiency, N , was calculated according to Albery and Bruckenstein (11). Standard Analysis. Several methods were used to determine ozone quantitatively. METHODA. The ferrous ammonium sulfate method (I) was employed, in which ozone reacts with Fe(I1) to form Fe(II1) and Oz, followed by back-titration of excess Fe(I1). METHODB. Sodium thiosulfate was used to titrate IZ produced by oxidation of I- by ozone ( I , 12, 13). The end point was determined visually using starch. METHODC. The end point used in Method B was changed t o an amperometric one, using a rotating wire electrode. All three procedures gave results agreeing within 1%. In the experiments involving determination of total acidity, sodium hydroxide was used as titrant and phenolphthalein as indicator. Constant-Current Coulometry. In the coulometric experiments reported below, the ozone-oxygen mixture from the ozonator was dispersed through the electrolysis cell containing 500 ml of supporting electrolyte. In determining n in Equation 1, after dispersiod for about liZhour, the gas flow was stopped and the ozone concentration determined from the diffusion current measured with electrode C at 1050 rpm. Electrode C was immediately removed and a magnetic stirring bar and the large platinum gauze electrode were inserted into the cell. A stopper was fitted into the neck of the cell to retard loss of ozone. The stopper was designed so electrical contact with the platinum gauze was possible.
(9) D. T. Napp, D. C. Johnson, S. Bruckenstein, ANAL.CHEM.,39, 481 (1967). (10) T. Rouse, Ph.D. thesis, University of Minnesota, 1960. (11) W. J. Albery and S. Bruckenstein, Trans. Faraday Soc., 62, 1920 (1966). (12) C. M. Birdsall, A. C. Jenkins, and E. Spadinger, ANAL.CHEM., 24,662 (1952). (13) I. M. Kolthoff and R. Belcher, “Volumetric Analysis” Vol. 111, p. 281, Interscience, New York, 1957.
A 400
200
T Y
0
- 200
I
1
I
1
I
1
I
1.6
1.4
1.2
1.0
0.8
0.6
0.4
1.6
E (SCE)
= 2500
+.
1
I
I
I
I
1.2
1.0
0.8
0.6
0.4
EM E )
Figure 1. I / E curves for electrochemical reduction of ozone atgold and platinum Electrode A , rotation speed min. 3 X M 0 3 Scan - to
I 1.4
rpm. Scan rate 1.0 volt per
Figure 2. Effect of rotation speed upon IIE curve for ozone reduction Electrode B (disk), rotation speeds in rpm, 3.57 X 10-4M03. Scan rate 1.0 volt per min from - to
+
- - - - - Scan + to -
E
The stirring of the solution was stopped immediately upon completion of the electrolysis. The platinum electrode and stopper were then removed and the limiting current and concentration of ozone determined.
=
constant
12 - I + 0.059 __ log ana Z ~
where I is the current at potential E, and ZIis the convectivediffusion limiting current. For the - to potential scan shown in Figure 1, ana equals 0.57. For the to - potential scan, ana equals 0.32. Similar results were obtained for a wave analysis in 0.1M perchloric acid solution, ana having values of 0.59 and 0.38 when the potential was scanned - to and to -, respectively. These results are not surprising, since the state of oxidation of the platinum depends upon its prior electrochemical history-i.e., the starting potential of the potential scan, and the rate and direction of the potential scan. The electrode is least oxidized during potential scan. The value of an for the cathodic the - to sweep agrees well with the value obtained below by linear scan voltammetry. The half-wave potential for the reduction at the gold ring electrode agreed with that on platinum when the potential scan was from - to but shifted by about -0.3 volt if the potential scan was from to -. The anodic processes observed for potentials more positive than 1.3 volts were elec-
+
RESULTS AND DISCUSSION
Current-Potential Curves. Z/E (current-potential) curves for the reduction of ozone at a gold and a platinum electrode were obtained using electrode A, and are given in Figure 1. Ozone is reduced at the platinum disk at potentials less anodic than 1.3 volts. The curve obtained for the cathodic potential sweep appears less reversible than that obtained for the succeeding anodic sweep. The cathodic current increased sharply for potentials more negative than 0.6 volt because of the reduction of oxygen which is present in the gas stream from the ozonator. Oxygen evolution occurred at potentials more positive than 1.3 volts. The value of ( ~ was n obtained from a wave analysis of the Z/E curves shown in Figure 1 for a rotating electrode in a 1M sulfuric acid solution of ozone. The wave equation for the reduction of ozone in this situation is
+
+
+
+
+,
+
Table I. Electrodesa and Parameters Electrode A B C
Ring Au
Pt Pt
Disk Pt Pt Pt
RI,cm
Rz,cm
Ra, cm
PZi3
N
0.3843 0.3869 0.3886 0.383 0.3876
0.3977 0.3981 0.3986 0.399 0.3969
0.4142 0.4051 0.4443 0.409 0.4069
0.245 0.151 0.557 0.088 0.264
0.145 0.090 0.268 0.108 0.101
D AU Au E Pt Pt Electrodes constructed by Pine Instruments Co., Grove City, Pa.
VOL 40, NO. 3, MARCH 1968
483
Table 11. Concentration of Ozone before and after Electrolysis in 1M HB04 Constant cathodic current. 21.00 Expt. A
Ammole
3.51 X 1.13 x 2.38 x 2.32 x 0.43 X 1.95 X 0.097
n
2.0
C*(t =
Omin)
Cb(t = 15 min) 4Ctot Cb
(A Cbef f)oalod ACd
0.05 ma
Expt. B
lW4M 10-4~
10-4~ 10-4~ lV4M
lW4M
3.29 X 1.21 x 2.08 x 2.25 X 0.42 X 1.66 X 0.083 2.4
Expt. C 3.67 x 1.50 X 2.17 x 2.58 X 0.48 X 1.69 X 0.084 2.3
lW4M 10-4~ 10-4~ lW4M lW4M le4M
Expt. D
10-4~ lW4M 10-4~ lW4M lW4M lW4M
3.63 x 1.33 X 2.30 x 2.48 x 0.46 X 1.84 X 0.092 2.1
10-4~ lW4M
10-4~ 10-4~ lW4M lW4M
Table 111. Loss of Ozone Due to Effusion @(t = 0 min)
@(t = 15 min)
2.87 X lW4M 3.31 X lW4M
2.38 x 1 0 - 4 ~ 2.75 X lW4M
C*
AC*eff
ACl Cb/At
2.63 X lW4M 3.03 X lW4M
0.49 X l e 4 M 0.56 x 1 0 - 4 ~
1.24% rnin-' 1 . 2 3 x min-l
trode oxidation and oxygen evolution. The difference in half-wave potentials with potential scan direction is probably connected with the fact that the reduction of the gold-oxygen layer formed during anodization of the electrode at potentials more positive than 1.3 volts occurs at 0.9 volt. The increase in the cathodic current for potentials more negative than 0.5 volt was due to oxygen reduction. The ratio of ring to disk diffusion-limiting currents measured at 0.8 volt is 0.25, while the theoretical value of P z I 3 for electrode A is 0.245. Figure 2 shows the effect of rotation speed to 6400 rpm upon the anodic sweep of the I / E curve for the reduction of ozone. The values of the diffusion current measured at 0.8 volt were directly proportional to the square root of the speed of electrode rotation. It is convenient to report data obtained at a rotating disk electrode in terms of L, which is defined by
where the symbols have their conventional significance (9). The value of L is 73.6 =t 3.8 in 1 M sulfuric acid and 69.1 ==I 1.1 in 0.1Mperchloric acid. Evaluation of n and m in Equation 1. Constant-current coulometry at a large area (-60 sq cm) platinum gauze electrode was used to evaluate n and m in Equation 1. To determine n, a constant cathodic current less than the limiting current for the experimental conditions was passed for a known time through the rapidly stirred solution of ozone. The limiting current for ozone reduction was determined using electrode C, before and after constant-current electrolysis. The concentrations of ozone were then calculated using the value of L. The concentrations before, Cb((t= 0 minute), and after electrolysis, C'(t = 15 minutes), and the total change in concentration, ACbtOt, are given in the first three rows of Table 11. The results for ACb in Table I1 are too large, since some ozone was lost by effusion to the atmosphere even though precautions were taken to decrease such losses. The rate of effusion was determined by duplicating the experimental conditions described above, the only difierence being that no current was passed through the electrode during the 15-minute electrolysis period. The ozone solution was stirred during that period, at a rate approximately equal to the rate used during the electrolysis. The losses due to effusion are given in Table 111. 484
rn
ANALYTICAL CHEMISTRY
dC/dt A C/Cb The relative rate of loss, 7was , approximated as 7
eb
L
AI
where is the average concentration during the experiment and At is equal to 15 minutes. The results of this calculation are also shown in Table 111, column 4, in which ACberris the change in ozone concentration resulting from effusion. Appreciable loss of ozone by the reaction
can occur after relatively long times. As shown below, the loss rate via this process is less than 0.5 of the initial ozone concentration. The concentration data given in Table I1 were corrected to give the net concentration change due to electrolysis ACbel. The loss, due to effusion, (ACeffb)cale, was estimated by multiplying the average concentration, Cb, by the rate of loss and by the total time. The results of these calculations are shown in the fifth and sixth rows of Table 11. The number of milliequivalents of current passed during the 15-minute electrolysis was 0.196. The average number of mmoles of ozone reduced, Ammoles, was 0.089 + 0.008 as calculated from the seventh row of Table 11. The corresponding number of equivalents per mole of reaction is given in the final row. The average value obtained for n was 2.2 A 0.2. To determine the hydrogen ion stoichiometry in the electrochemical reduction of ozone-i.e., m in Equation l-constant-current electrolysis was performed on an ozone-saturated 0.1M NaC104-2.65 X lO-3M HClO4 solution. A solution of sodium bromide was used in the counterelectrode (CE) compartment, so that the pH of the solution would not be altered by the reaction at the CE. The ozoneoxygen mixture from the ozonator was continuously dispersed through the solution during the electrolysis, using a constant current of 20.00 A 0.05 ma for 60 minutes. The total acidity of a IO-ml sample was determined before and after electrolysis by titration with standard sodium hydroxide; 0.75 milliequivalent of current passed during the 60-minute electrolysis and 0.73 milliequivalent of hydrogen ion was consumed; hence the number of hydrogen ions consumed per mole of ozone reduced, m in Equation 1, was
2 O(
300 IO(
a
t
c
Y
-IOC
-2oc
1.8
1.6
1.4
t.2 E
1.0
0.8
0.6
(SCE)
Figure 3. I / E curves obtained at gold electrode for hydrogen peroxide 1.8
Electrode D (disk), rotation speed 400 rpm. Scan rate 1.0 volt per min 7X10-4M
H202
- - - - - Residual
+ 2H+ + 2e
-
+
HZOZ '/zOZ
1.4
1.2
1.0
0.8
0.6
E NE)
Figure 4. I/E curves for ozone in presence and absence of hydrogen peroxide
calculated to be 1.95. A similar experiment was performed in a 1M NaS04 solution to which a known amount of had been added. In the sulfate medium, the number of hydrogen ions consumed per mole of ozone reduced was calculated to be 1.8. From the above-described experimental results, Equation 1 could be written as Equation 2 or 3 , 0 3
1.6
(3)
Voltammetric Investigation with a Ring-Disk Electrode. The rotating ring-disk electrode seemed ideally suited to determine if hydrogen peroxide is a product of the electrochemical reduction of ozone. However, it was necessary to show that hydrogen peroxide could be detected in an ozone solution. Figure 3 shows a I/E curve obtained at the gold disk of electrode D for hydrogen peroxide. A gold disk electrode was selected in preference to a platinum electrode because of its larger anodic potential limit. The residual curve is included in Figure 3. A peak anodic current was obtained during the anodic potential sweep in the hydrogen peroxide solution. No wave was observed corresponding to the reduction of hydrogen peroxide. The IIE curves for ozone in the presence and absence of hydrogen peroxide are shown in Figure 4. The presence of hydrogen peroxide is manifested by an anodic peak current obtained during the anodic potential sweep at potentials more positive than 1.0 volt. Hence, hydrogen peroxide can be detected in the presence of ozone at a gold electrode. Figure 5 shows IIE curves obtained at the gold ring of electrode A (platinum disk) in a solution of 1M sulfuric acid containing ozone. The curves were obtained when no current passed through the platinum disk (open circuit), and for a disk potential of 0.8 volt. Also shown is the residual curve obtained in an ozone-free solution. The portions of the
Electrode D (disk), rotation speed 400 rpm. Scan rate 1.0 volt per min
3 x i o - 4 o3 ~ -3 x 10-4M O3and 7 X 10-4M H202
curves obtained in the ozone solution at potentials more positive than 1 . 3 volts are identical to the residual curve. Thus, hydrogen peroxide is not a product of the electrochemical reduction of ozone, and the reduction as described by Equation 3 is not possible. It can be demonstrated that the reduction of ozone at 0.8 volt produces oxygen. The change in the IIE curve at the ring electrode produced by switching the disk potential from open circuit to 0.8 volt is a function of the reaction at the disk electrode. Z/E curves at the ring were calculated assuming Reaction 2. The ring current at 0.8 volt, when the disk is at open circuit, I, (open, 0.8), is related to the limiting disk current for the reduction of ozone, Id(o.8) (9). Ir (open, 0.8)
= PZ"ld
(0.8)
(4)
For platinum electrode potentials of 0.1 volt, oxygen is reduced by the reaction (14-16) 0 2
+ 4H+ + 4e
-+
2Hz0
The ring current for a ring potential of 0.1 volt is
I, (Open, 0.1) = P2'31d(0.8)
+ 2/32'31d(0.8)f (iz)02
(5)
(14) L. Muller and L. N. Nekrasov, Dokl. Akad. Nauk., 154,437 (1964). ( 1 5 ) D. T. Sawyer and L. V. Interrante, J. Electroanal. Chem., 2, 310 (1961). (16) J. P. Hoar, Proc. Roy. SOC.,A 142,628 (1933). VOL 40, NO. 3, MARCH 1968
0
485
I I
I I I
I
I I
I
1.4
1.2
1.0
0.8
0.6
0.4
0.2
E, W E )
1.6
I
I
I
1
I
I
1.4
1.2
1.0
0.8
0.6
0.4
Figure 6. I / E curves at the ring electrode with the disk electrode at open circuit and a potential of 0.8 volt
E, (SCE)
Figure 5. I / E curves at the ring electrode with the disk electrode at open circuit and a potential of 0.8 volt Electrode A (ring), rotation speed 2500 rpm. Scan rate 1.0 volt per min, 3
x
+
=
1.0 volt
1 0 - 4 ~ 0 ~
Scan - to + - _ - - -Scan + to -
In Equation 5, the terms on the right-hand side have the following significance : p2’33rdO.8)= ozone reduction current at the ring electrode, the same as I,(open, 0.8) 2p2’3Z(0.8) = current from the reduction of oxygen produced by ozone reduction at the ring according to Equation 2 ( i l ) ~ , = current from the reduction of 0 2 present in the bulk of the solution The ring current for a ring potential of 0.8 volt with a disk potential of 0.8 volt, 1,(0.8,0.8) is (9) Zr(0.8, 0.8)
=
/3*’31,j(0.8)- NId(0.8)
=
(p2’3
=
Z,(open, 0.8)
- N)I40.8) - NZ40.8)
(7)
The ring current for a ring potential of 0.1 volt and disk potential of 0.8 volt is Zr(0.8, 0.1) = (pZi3- N)Zd(0.8) 2(p2’3- N)Id(0.8)
+
+
2N140.8) = 3p2’31,i(0.8)- NId(0.8)
+
(ii)02
+ (iz)oz
(8) (9)
In Equation 8, the terms on the right side have the following significance : ( p 2 / 3 - N)Zd(0.8) = current due to ozone in the bulk which reaches the ring [shielding (9)] 2(p2/3- N)I40.8) = current due to oxygen produced by the reduction of ozone at the ring according to Reaction 2 486
Electrode B (ring), rotation speed 400 rpm. Scan rate per min from to -, 3 X 10-4M O3 Disk at open circuit - - - - - Disk at 0.8 volt (SCE)
ANALYTICAL CHEMISTRY
2NId(0.8) = current due to oxygen produced by reduction of ozone at the disk according to Reaction 2 (il)o2 = same as above Hence
NZd(0.8) = Z,(open, 0.1) - 1,(0.8, 0.1)
(10)
Figure 6 shows the I/E curves obtained on the cathodic sweep at the ring of electrode B for disk potential at open circuit and at 0.8 volt. The difference in the I , us. E, curves, when Ed = 0.8, is the same for ring potentials of 0.8 and 0.1 volt. This difference in current was found to be equal to NId(0.8) as predicted by Equation 10. Hence, our initial assumed reaction, Equation 2, is correct and oxygen is a product of the reduction of ozone. Diffusion Coefficient of Ozone. Using the value of n equal to 2 in Equation 1 and the values of L given above, the diffusion coefficient of ozone was calculated as 1.53 f 0.12 x 10-5 sq cm per second in 1 M sulfuric acid and 1.39 =t 0.05 x 10-5 sq cm per second in 0.1Mperchloric acid. Chronopotentiometric Studies. A series of experiments was conducted on 5.6 X 10-4Mozone solutions of 0.1Mperchloric acid with a stationary platinum disk electrode (electrode E ) using the chronopotentiometric technique. The current density was varied from 84.9 to 764 pa per sq cm. To obtain a reproducibly oxidized platinum surface, the electrode was rotated in solution for 5 seconds with no applied potential or current. The rotation was stopped, and constant current passed after 10 seconds, when the solution had become quiescent. All transition times were less than 5 seconds, and deviations from semi-infinite linear diffusion could be ignored.
2
I
t
3
(sac1
Figure 7. Chronopotentiograms for ozone reduction at a stationary electrode
_____
Electrode B (disk) 0.1MHC104,5.6 X 10-4M03, 254.5, pa per sq cm 1M H,S04, 3.6 X 10-4M 03,190.5, pa per sq cm
Figure 7 shows a typical chronopotentiogram which demonstrates the irreversibility of the ozone reduction on a platinum electrode. The transition time for the irreversible reduction of ozone was taken when the electrode potential reached +0.8 volt. A fairly constant value for the transition time constant, ir112/AC= n F d T / 2 ( D 1 l 2 )was , obtained over a ninefold range of current densities-Le., 633 =k 29. Similar chronopatentiograrns are obtained for the reduction of ozone at platinum in 1M sulfuric acid. One such curve is shown in Figure 7 . The transition time constant for ozone reduction was 608 f 57. The calculated diffusion coefficients for ozone were 1.38 X 10-6 sq cm per second in 0.1M perchloric acid and 1.30 X 10-5 sq cm per second in 1M sulfuric acid. Linear Scan Voltammetry. The electrochemical reduction of ozone in 1M sulfuric acid was investigated using linear scan voltammetry. The concentration of the ozone solution was determined from L using electrode E. Since the potential sweep for the voltarnrnetric scan started at a potential in the region of platinum-oxygen film formation, a pretre itment technique was used to ensure uniform surface coverage of the electrode. Following each potential sweep, the potential was switched to 1.40 volts and the electrode rotated at 400 rpm for 5 seconds. The rotation was then stopped and the solution allowed to become quiescent during a 10-second period. A IIE curve was then recorded at the stationary electrode for a cathodic potential sweep. Figure 8 shows the I / E curves obtained at the disk of electrode E for various rates of potential scan. The residual I/E curves are included in Figure 8 . The peak current values, I,, corrected for residual current, I,,,, are tabulated in Table IV as a function of the square root of the potential scan rate, Y1I2.Plotting (I, - I,,,) us. V1i2yields a straight line having a slope equal to 2.16 X 10-I ma volts-l/z sec1/2. The equation relating peak current and scan rate is
I,
=
3.01 X 105n(~n,)1/2AD1/2C*V1'2
where V is scan rate in volts per second, n, is the number of electrons up to and including the rate-determining step, and CY has the usual electrochemical significance. The product an, can be determined from (17) (17) R. S. Nicholson and I. Shain, ANAL.CHEM., 36, 706 (1964).
4 1.4
1.2
0.6
1.0 0.8 EME)
0.4
Figure 8. IIE curves obtained using linear scan voltammetry at a stationary electrode Electrode E (disk), scan from to -, volts per second. 3.59 X 10-4M 03-1M H2S04. Lower set of curves obtained in supporting electrolyte
+
RT -0.0477 ana ffna where Epis the peak current potential and E,lz is the half-peak potential. From the I/E curve obtained for V1I2 equal to 0.409 volt1/*sec-1/2, E, - E,iz was evaluated as -0.16 volt. The value of an, was calculated as 0.30. Using an, equal to 0.30 in Equation 1, and n = 2 , D(03) was calculated as 1.57 X sq cm per second. These results agree well with those reported above. E p - E,/z = -1.857-
~
Table IV. Values of Peak Current as a Function of Potential Scan Rate Electrode F(disk). [03] = 3.59 X lO-'M in 1M H1S04 V, volts per second Vl/Z I p - I,,,, ma 0.0167 0.129 3.25 X 10-2 0.0334 0.183 4.44 x 10-2 0.0500 0.224 5.29 X 10-2 0.0667 0.258 5.92 X 1 0 - 2 0.0833 0.289 6.73 X 10-2 0.1ooo 0.316 7.25 X 10-2 0.1167 0.342 7.85 x 10-2 0.1333 0.364 8.31 X 10-2 0.1500 0.387 8.83 X 10-2 0.1665 0.409 9.15 X 10-9 sq cm per second D ( 0 3 ) = 1.57 X VOL. 40, NO. 3, MARCH 1968
487
Spontaneous Decomposition of Ozone. After using solutions of ozone for some time, in the platinum ring-disk electrochemical studies described above, it was observed that the anodic currents measured for potentials more positive than 1.3 volts were larger than when the solution was fresh. Ozone was removed by bubbling nitrogen through the solution. Z/E curves obtained from the resulting ozone-free solutions were typical of those obtained in solutions of hydrogen peroxide. To establish whether the decomposition of ozone was catalyzed at the platinum electrode surface, ozone was bubbled through 1 M sulfuric acid for 150 minutes, in the absence of any platinum and in the presence of a -6O-sq cm bright platinum gauze electrode. Ozone was removed by bubbling nitrogen through the solutions. The Z/E curves obtained following the experiment showed that the hydrogen peroxide content was the same as that obtained in the absence of the platinum electrode and equaled 2 x lO-5M. The concentration of ozone during the experi-
ment was 5 x 10-4M. Hence, the spontaneous decornposition of ozone to hydrogen peroxide is not appreciably catalyzed by a bright platinum surface. Molar Absorptivity Determination. Because of the large variations of the reported value of the molar absorptivity coefficient for ozone at 258 to 260 mp, it was redetermined with a Cary 15 spectrophotometer. Beer’s law was obeyed in the concentrations range 2 X 10-5 to 4 X 10-4Min a 1-crn quartz cell. A molar absorptivity coefficient of 2500 liter mole-’ cm-1 was calculated, which is in close agreement with the value of 2600 liter mole-’ cm-’ reported by Ingols, Fetner, and Eberhardt ( I ) .
RECEIVED for review November 3, 1967. Accepted December 26, 1967. Work supported by the Space Science Center of the University of Minnesota and the National Science Foundation. Taken in part from the Ph.D. theses of D. C. Johnson (December 1967) and D. T. Napp (December 1967).
Linear Sweep Voltammetry of Silver(l), Cadmium(I I), Lead( I I), Copper( I I), and Indium(lll) in Molten NaN03-KN03Eutectic Gleb Mamantov, James M. Strong, and Fred R. Clayton, Jr. Department of Chemistry, University of Tennessee, Knoxville, Tenn. 37916 Linear sweep voltammetry of silver(l), cadmium(ll), lead(ll), copper(ll), and indium(ll1) in molten sodium nitrate-potassium nitrate (50-50 mole %) was studied at platinum wire and indium pool [in the case of In(lll)] electrodes. Linear dependence of peak current on concentration obtained for the first four cations indicates that metal oxide formation is not appreciable at the experimental temperature of 245’ C. The addition of anhydrous lnCla to the solvent results in a yellow precipitate, possibly InOCI. The importance of the oxidation of reduction products by the nitrate solvent, blocking by the oxide film, and alloy formation is discussed, based on the observed variation of peak current with scan rate, and the calculated n values.
A NUMBER OF ELECTRODE PROCESSES have been studied in molten nitrates (1-3). The electrode reactions of the solvent, NaNOo-KN03eutectic, and the products of the NOo- reduction, NO2- and 0-2,have received considerable attention recently (4-8). The reductions of cations in this melt (1) C. H. Liu, K. E. Johnson, and H. A. Laitinen, “Molten Salt Chemistry,” M. Blander, ed., Interscience, New York, 1964. (2) H. A. Laitinen and R. A. Osteryoung, “Fused Salts,” B. R. Sundheim, ed., McGraw-Hill, New York, 1964. (3) Yu. K. Delimarskii and B. F. Markov, “Electrochemistry of Fused Salts,” Sigma Press, Washington, D. C., 1961. (4) H. S. Swofford and H. A. Laitinen, J . Electrochem. Sac., 110, 814 (1963). (5) H. S. Swofford and P. G. McCormick, ANAL.CHEM., 37, 970 (1965). (6) G. G. Bombi, R. Freddi, and M. Fiorani, Ann. Chim., 56, 759 (1966). (7) L. E. Topol, R. A. Osteryoung, and J. H. Christie, J . Phys. Chem., 70, 2857 (1966). (8) D. Inman and J. Braunstein, Clzem. Cornmum. 1966, 148.
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have been studied at either mercury (9, 10) or platinum (II16) electrodes. Studies at the platinum electrode usually examined the reduction of Ag(1) (11, 13-15), frequently under poorly defined conditions of mass transfer. It was not always realized that a number of cations are sufficiently acidic to undergo partially the reaction (7, 17, 18) (written for a divalent cation) M+2 2N03- =- MO(,, 2N02 SO2. Another complication present with more active metals is the oxidation of the reduced form by the nitrate solvent, probably to the metal oxide (9, 19). In fact, Inman and Bockris (9) have stated that in the cases where metals oxidizable by the nitrate solvent are produced upon electrolysis, a solid electrode cannot be used. Since a platinum wire electrode is considerably simpler to use at elevated temperatures than a mercury electrode, it seemed desirable to re-examine the use of a stationary platinum wire electrode in molten NaN03-
+
+
+
(9) D. Inman and J. 0. M. Bockris, J. Electroanal. Chem., 3,126 (1962). (10) H. S. Swofford and C. L. Holifield, ANAL.CHEM., 37, 1509 (1965). (11) Yu. K. Delimarskii, B. F. Markov, and L. S. Berenblum, Zh. Fiz. Khim., 27, 1848 (1953). (12) Yu. K. Delimarskii and I. D. Panchenko, Ukr. Klzim. Zh., 19, 47 (1953). (13) S. N. Flengas, J . Chem. Sac., 1956,534. (14) N. G. Chovnyk and V. V. Vashchenko, 211. Fiz. Khim.,35,580 (1961). (15) D. L. Manning, Talanta, 10, 255 (1963). (16) R.Narayan and D. Inman, J. Polarog. Sac., 11, 27 (1965). (17) F. R. Duke, “Fused Salts,” B. R. Sundheim, ed., p. 413, McGraw-Hill, New York, 1964. (18) G. G. Bombi and M. Fiorani, Talarzta, 12, 1053 (1965). (19) B. J. Brough and D. H. Kerridge, Znorg. Chem., 4, 1353 (1965).
