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Electrochemistry of oxygen and hydrogen peroxide at clean and at halide-covered polycrystalline silver
J. Electroanal. Chem., 150 (1983) 97~-109 Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands
97
ELECTROCHEMISTRY OF OXYGEN AND HYDROGEN PEROXIDE AT CLEAN AND AT HALIDE-COVERED POLYCRYSTALLINE SILVER
E. STEVEN BRANDT Research LaboratorieS, Eastman Kodak Company, Rochester, N Y 14650 (U.S.A.) (Received 27th July 1982; in revised form 14th December 1982)
ABSTRACT The effect of halide adsorption on the rate of hydrogen peroxide decomposition in alkaline solution has been studied. The reduction of oxygen at clean silver in acidic, alkaline, and unbuffered electrolytes has also been investigated for comparison. On prereduced electrodes the amount of hydrogen peroxide observed during the reduction of oxygen is related to pH and adventitious contamination. Specific adsorption by halides decreases the efficiency of the "four-electron" reduction of oxygen with concomitant formation of hydrogen peroxide. With iodide, the reduction of oxygen to hydrogen peroxide is nearly 100% efficient. Coverages of lad s approaching 0.2 monolayer are sufficient to create an energy barrier for breaking of the oxygen-oxygen bond and effectively poison the catalytic surface of silver.
INTRODUCTION
Considerable interest in the electrochemical reactions of oxygen and hydrogen peroxide at silver electrodes has been generated from fuel-cell development and from gas-phase heterogeneous catalysis by supported silver catalysts. In addition, reactions of hydrogen peroxide and oxygen at various electrode materials, including silver, are of general fundamental importance in electrocatalysis. Of particular importance in both fundamental and applied areas is the formation of diatomic oxygen (or dioxygen [1]) intermediates during the reaction of oxygen at silver surfaces. In fuel cells these intermediates, particularly hydrogen peroxide, can attack the electrodes oxidatively, leading to premature failure of the cell [2]. However, at least equally important for fuel-cell applications is the fact that the presence of stable dioxygen intermediates is an indication of intrinsic cell inefficiency in the reduction of oxygen at the cathode. On the other hand, the presence of dioxygen species is desired in the catalytic conversion of ethylene to ethylene oxide over silver, since these species are thought to increase catalytic specificity for the desired product [3]. In both systems the rate of oxygen-oxygen bond breaking determines the concentration of stable dioxygen intermediates. In solution the efficiency of this process can be followed by observing the formation of hydrogen peroxide during the reduction of oxygen using a rotating ring-disk electrode [4-7]. In the present study 0022-0728/83/$03.00
© 1983 Elsevier Sequoia S.A.
98 the reduction of oxygen has been investigated on clean silver electrodes, which favor the "four-electron" reduction process [8], and on silver surfaces that have been exposed to halide adsorbates, which significantly reduce the catalytic activity of the silver surface, thus producing conditions which favor formation of stable dioxygen intermediates. EXPERIMENTAL Voltammetric measurements were made using conventional commercially available instrumentation. For experiments with ring-disk electrodes, an in-house-built potentiostat was interfaced directly to the three-electrode potentiostat to provide independent control of two working electrodes [9]. Unless otherwise noted, the disk electrode had a measured area of 0.458 cm 2. The dimensions of the Pt-ring/Ag-disk electrode were taken as r 1 (disk diameter) 0.382 cm, r 2 (ring i.d.) 0.399 cm, and r 3 (ring o.d.) 0.422 cm, which translates to a disk area of 0.458 cm 2 and a ring area of 0.059 cm 2. F r o m these dimensions N, the theoretical collection efficiency, was calculated by the treatment of Albery and Bruckenstein [10] to be - 0.183. Potentials are reported vs. the H g / H g 2 S O 4 (sat. K2SO4) reference electrode, which had a measured potential of 0.371 V vs. SCE in a saturated K2SO 4 electrolyte. Before use, all glassware was cleaned in a 1 : 1 mixture of concentrated nitric and sulfuric acids. When not in use, the cells were also stored in this solution. Vacuum surface analysis of silver samples was performed on a DuPont 650B ESCA [11]. This instrument has a nonmonochromatic Mg K , (1253.6 eV) source, which is regulated at 300 W. System vacuum during these experiments was kept at < 1 × 10 -7 Torr. For ESCA analysis, Teflon-shrouded 6-mm-diameter silver disks were used as electrodes. After a given experiment the sample shroud was removed, and the sample, suspended in distilled water, was transferred to the ESCA. Before each use, silver samples were polished concentrically to 0.25 /~m on a h o b b y lathe with diamond pastes. Cross-contamination of ring and disk electrodes was thereby minimized. After they were polished, silver electrodes were rinsed throroughly with purified water and then cycled between background potential limits several times. While the electrode was under potential control at the lower limit, the electrolyte was aspirated. With the electrode open circuited, the cell was rinsed and refilled with fresh electrolyte. This process was continued until satisfactory i/E curves for the background and for the reduction of 02 at silver (see Results Section) were obtained. According to the terminology adopted by Sepa and co-workers, these samples are referred to as "prereduced" in the text [12]. All chemicals were reagent grade, except for K2SO 4 and K O H , which were A l f a / V e n t r o n ultrapure grade. Solutions were prepared with distilled water that had been passed through a carbon filter train and then redistilled in glass. Water prepared in this fashion was judged acceptable for use, according to published criteria [13].
99 RESULTS
Reduction of oxygen and hydrogen peroxide at clean silver electrodes Effect of pH T h e r e d u c t i o n of 02 a n d H 2 0 2 at p r e r e d u c e d r o t a t i n g silver disk electrodes at different p H ' s is shown in Fig. 1. T h e b a c k g r o u n d currents in H e - d e o x y g e n a t e d electrolytes are shown for c o m p a r i s o n (note c u r r e n t scales). A t p H 12.8 a n d in u n b u f f e r e d 0.05 mol 1-1 K z S O 4 ( p H 5.9), the wave associated with H202 r e d u c t i o n on silver occurs at a higher p o t e n t i a l t h a n the r e d u c t i o n of oxygen, in qualitative a g r e e m e n t with their respective t h e r m o d y n a m i c s t a n d a r d potentials. However, at p H 1.5, the r e d u c t i o n of H202 occurs at the s a m e potential, o r even at p o t e n t i a l s slightly negative relative to O z reduction.
Ring-disk experiments T h e r e d u c t i o n of oxygen on p r e r e d u c e d silver was p e r f o r m e d in 0.05 mol 1 1 K 2 S O 4 while the p r o d u c t s of that r e a c t i o n were m o n i t o r e d at the p l a t i n u m ring. F i g u r e 2a shows that two limiting c u r r e n t p l a t e a u s are o b s e r v e d in the ring
o
0 E / V ~I
::
" i
-10 ;
.
.
"
f
b
f 0
E / V "-,
.
I
/ %0 ~
.
l .
.
.
.
200~A I .............. IOOp.A
Fig. 1. Rotating silver disk voltammograms for 02 and H202 reduction on prereduced silver electrodes; (. . . . . . ) background in He-deoxygenated electrolyte; ( - ) O2 (saturated); ( . . . . . . ) - 10- 3 mol l - 1 H202. (a) 0.05 mol 1-1 H2SO4; (b) 0.05 tool 1 I K2SO4; (c) 0.1 mol 1- l KOH. All scans: 50 mY/s, 900 rpm.
100 a).
ir with disk at E d I I
<
÷ 1.0 I
0
J
I
-liO
,,
>
b, wx
/
~,
E/v
-I-/~5/~A
/
j
b.)
/ 2
J, \\
12OO/.L A -'i° ""-.
,
......~'.\
,
E/V d
"D"""""." |os A
]20~A"5\ . . . . . . . .
x
Fig. 2. (a) i r / E r curves obtained during the constant reduction of oxygen at Ed 'adjusted for ilaim. (b) Ring profiles obtained with E, at potentials a and b (note sensitivities). Conditions: O2-saturated 0.05 mol 1- l K2SO4, scan rate 50 mY/s, 900 rpra.
ir/E d
voltammogram while oxygen is being continually reduced at E o (limiting current density). Assuming n is equal for disk and ring processes, the limiting current observed at + 0.87 V (ilrim) is ca. 60% of the theoretical collection efficiency of 0.183 predicted by the electrode geometry (cf. Experimental Section). A much smaller current is observed at + 0.20 to - 0 . 2 5 V. When the ring electrode is potentiostated at these potentials and the disk electrode is scanned through the oxygen reduction wave, a continuous wave that mirrors the oxygen wave and a wave that goes through a m a x i m u m at - 1.12 V are observed with the ring at +0.87 V (potential a) and - 0 . 1 2 V (potential b), respectively (Fig. 2b). The ring current at - 0 . 1 2 V can be shown to be due to the formation of hydrogen peroxide by standard addition methods. Identification of the electroactive species responsible for the current at + 0.87 V is more difficult. This current might be due to the oxidation of O H - via 4 O H - ~ 02 + 2 H 2 0 + 4 e - , which would explain its magnitude, since this is the reverse of the four-electron oxygen reduction. This species is not observed in either strongly acidic or strongly basic electrolytes where the solutions are well-buffered. The relative amount of H202 produced at p H 12.8 is similar to that produced in unbuffered solutions and also goes through a maximum. In both electrolytes, the peak currents are somewhat irreproducible, since the amount of peroxide produced
101
~ 200/.~A
dc
Ed/V
(
i
I
)
-LO
Ir o
flOrA
/ Fig. 3. Hydrogen peroxide production during the reduction of oxygen in O2-saturated 0.05 rnol 1-1 H2SO4; all scans 50 m V / s , 900 rpm. Dashed horizontal line indicates predicted i/im based on results obtained in 0.1 tool 1-1 KOH.
appears to be related to the extent of contamination of the cell, in agreement with the results obtained by Fischer and Heitbaum [6]. Levels of H202 as low as 8 × 10 -6 mol 1-1 and as high as 2.4 x 10 -4 tool 1-1 have been recorded under apparently similar conditions. In acid solution (0.05 mol 1- l H2SO4) the background evolution of hydrogen overlaps the oxygen and hydrogen peroxide reduction waves at the disk (Fig. 1). However, the magnitude of i r shows that peroxide production is relatively efficient at this pH (Figs. 3 and 2b). This result demonstrates that the decomposition of peroxide generated during the reduction of oxygen at prereduced silver electrodes is much less efficient at low pH's than at alkaline pH or in unbuffered electrolyte. Reduction of oxygen at halide-covered silver electrodes
Mercury salts [6] and organic surfactants [6,7] have been used to demonstrate the effect of adsorption on the rate of peroxide production during the reduction of
102
oxygen at silver. However, anions in general (and the halides in particular) are more desirable candidates to explore the role of specific adsorption on cathodic faradaic processes at silver because (1) in general, metals adsorb anions much more readily than cations or neutral species; (2) in gas-phase catalysis work, halide-containing hydrocarbons are known to improve the selectivity of silver for production of ethylene oxide from ethylene, probably through increasing the surface concentration of dioxygen species [3]. Chloride The addition of KC1 to final concentrations of 1.0 mol l- 1 to a 0.1 mol 1-1 K O H electrolyte saturated in O: causes an increase in ir, attributed to the formation of H202 (Fig. 4). At C1- concentrations below 10-1 mol l-1 relatively minor increases of H202 (compared with prereduced electrodes) occur during the reduction of oxygen at the silver disk. Actually, at least two waves are observed at the ring electrode; however, both waves disappear when the solution is deoxygenated (Fig. 4b).
I/. ~"
...........
~'°
J
//
o),
.
] 2oo~A
//
gdc
" ~0.0 Ed/V < ir 0 "%'
,..,°..
b).
< ~'dc
Ed / V < ,~?.0 Ir 0
~"--L
'
If
-I.O
'
>
15 A Fig. 4. (a) Reduction of oxygen in 0.1 mol 1- l K O H only (1), and in 0.1 mol 1-1 K O H + 1 mol 1-1 KCI (2). (b) Background in deoxygenated electrolyte as used to obtain curve 2. All scans: 50 m V / s , 900 rpm.
103
Bromide The quantity of H202 produced during the reduction of oxygen increases systematically with B r - concentration and reaches a m a x i m u m near 5 × 10 -3 mol 1-1 B r - (Fig. 5). Phase AgBr, which is formed at the upper limit, is reduced before the onset of oxygen reduction and hence H 2 0 2 production, as evidenced by the small cathodic surface waves at ca. - 0 . 4 0 V. However, as shown by both ring and disk voltammograms, the effect of Brad s persists at the more cathodic potentials where these processes occur. At 10 - 3 mol 1-1 B r - (curve 6, Fig. 5) multiple inflections are observed in the oxygen wave at the disk. T h e most pronounced inflection, centered at ca. - 1.1 V, corresponds to the maximum production of H 2 0 2 . Higher concentrations of bromide cause only minor changes in the shape or position of the waves or the quantity of H202 monitored but contribute significantly to the dissolution of the silver disk
/ [! Ed/V < 'ro~iO'O
,
-:/~//////
200ffA I
///
i
' ..
I,0
.-~- >
~ \/,// //' / 2offAl,f
~//.a "~ / / /
/ •
Fig. 5. Effect of B r - on the a m o u n t of peroxide formed during the reduction of oxygen in 0.1 mol 1-1 K O H . Curves 1-5, successive additions of KBr (2 × 10-5 tool per addition); curves 6, to 4 × 10-3 M final concentration. Dotted line, background of prereduced silver in clean O2-saturated 0.1 mol l - 1 KOH. All scans: 50 m V / s , 900 rpm.
104
electrode when the potential is scanned towards positive potentials beyond the foot of the oxygen wave. lodide Iodide shows by far the most pronounced effect of the halides on the rate of reduction of oxygen at silver electrodes. Figure 6 was obtained by successive additions of KI to a 0.1 mol 1-1 K O H electrolyte. The maximum current, / m a x is observed at a much lower concentration (ca. 5 x 10 -6 mol 1-1) of iodide than with bromide or chloride. A t v e r y low concentrations, approaching 10 -8 mol 1-~, the effect of iodide is seen as a minimum in the ilim plateau near the point of zero charge (pzc) of silver. This minimum in the diffusion-limited mass-transport region suggests competition between two rate-determining processes, such as the formation of a passivating layer of lad s at the expense of a catalytically active surface layer of oxygen-containing species. With increasing iodide this minimum deepens and broadens on the anodic side until, at 5 x 10 - 6 mol 1- l I-, two distinct waves of approximately equal height are formed in the disk voltammogram. To eliminate the possibility of solution iodide participating in the reactions of hydrogen peroxide [14], the prereduced electrodes were exposed to iodide at open circuit in a cell separate from the analytical cell containing the purified oxygensaturated electrolyte. Under these conditions approximately one monolayer of lad s is formed, as measured by ESCA [15], over a wide range of concentrations (Table 1).
200ffA
// Ed/V <
ido
,
~'I.0
~
'r 0
~
~
)
" ~
I 20ffA Fig. 6. Successive additions of KI to an O2-saturated 0.1 mol 1- i K O H electrolyte to a final concentration of - 5 x 10 -6 mol 1- J KI. All scans: 50 m V / s , 900 rpm, negative-going sweep after addition.
105 TABLE 1 5-rain exposure in purified H 2 0 containing KI, while the silver sample was rotated at 900 rpm [ K I ] / m o l 1-1
I / A g × 100 a
1 10 - l 10 - 2 10 -3 10 - 4 10 -5 10 - 6
9.20 11.5 10.0 9.82 9.29 10.1 9.56
Average
9.83 _+0.83
a I 3d5/2 and Ag 3d5/2 corrected for cross sections by using the tables of Scofield [16].
The v o l t a m m o g r a m s obtained f r o m these electrodes were essentially the same as observed with iodide in solution, but after reversal at the lower limit the positive-going wave returns to the same c u r r e n t - v o l t a g e wave shape as observed on a prereduced electrode. ~"" ..............
!/o .I
,
,
Ed/V
,
0.o
I0
o/
J o/
.-
o ..........
- 1.0 o
o
s~'"~o
o
o_ a ~6
~
o
2 0
\ 0°~o
Fig. 7. Upper: Comparison between voltammograms for 02 reduction at clean (dotted line) and at externally iodide-treated silver electrodes. Iodide exposure, 5 min at o.c.; unbuffered. Electrode Area: 0.29 cm 2. Lower: (I 3 d s / 2 / A g 3d5/2)×100, corrected for cross sections; from ESCA spectra taken at indicated potentials (see text for experimental details).
106 i/mA. cm -2 C//.~F. cm-2
60 I I
\ W. c\ \
I0
40
'X
8 6
\ 20
4 2
r.p. ,
0
i
0.5
1.0
1,5
E/V
Fig. 8. Correlation of pzc with faradaic current observed during the reduction of oxygen on iodide-treated polycrystalline silver. (a) Reduction of oxygen at silver in O2-saturated 0.1 mol 1 i KOH after preexposure to 10 - 3 mol 1-1 KI for 5 min (unbuffered). (b) Capacity vs. potential curve obtained in Ar-deoxygenated 0.02 mol 1-1 NaF at 10 Hz, AEac = 10 mV peak to peak. (c) Capacity vs. potential curve obtained in Ar-deoxygenated 0.01 mol.1- l KOH at 10 Hz, AEac = 10 mV peak to peak. C vs. E scans, 10 mV/s, 0 rpm; i vs. E scans, 50 mV/s, 900 rpm; electrode area 0.29 cm2.
The coverage of lad s as a f u n c t i o n of potential was followed b y using E S C A in c o n j u n c t i o n with removal of Ida s b y potential step in a cell c o n t a i n i n g a large v o l u m e of purified 0.1 mol 1-1 K O H . T o ensure that steady state had b e e n achieved after the potential step, the electrode was potentiostated for 30 min, d u r i n g which time the c u r r e n t vs. time behavior was recorded. As m e a s u r e d b y the increase in current, the steady state is reached rapidly, o n the order of milliseconds, although some slow d e s o r p t i o n c o n t i n u e s up to 1 min. T h e isotherm o b t a i n e d b y E S C A from electrodes that had been removed while still u n d e r potential control is shown in Fig. 7 along with the p o i n t - b y - p o i n t steady-state results. Coverages of Iads decrease with increasing cathodic potentials u n t i l ca. - 1.1 V, where the a m o u n t of lad s r e m a i n i n g is _< 0.2 m o n o l a y e r , which is the practical detection limit of the ESCA technique. A t this potential, the c u r r e n t for the reduction of oxygen begins to increase, f o r m i n g the second wave observed in the
107 voltammogram, corresponding to the reduction of peroxide. The ratio of the two waves is 1.1 : 1, the first wave being the larger, a ratio that is invariant at sweep rates < 100 mV/s. Using an alternating current technique similar to that described by Clavilier [17], Valette and Hamelin [18], we recorded the differential capacity of polycrystalline silver in dilute electrolytes of 0.02 mol 1- l NaF and 0.01 mol 1 1 N a O H and compared the results with the faradaic process of oxygen reduction on an iodidetreated silver electrode (Fig. 8). Good agreement exists between the experimentally determined pzc, which, according to the G o u y - C h a p m a n - S t e r n double-layer model [19], is reflected by the capacitance minimum observed at ca. - 1 . 3 V and the potential at which i lim for a four-electron wave is reached on iodide-treated polycrystalline silver electrodes. DISCUSSION
Oxygen reduction at clean (prereduced) silver electrodes At prereduced silver electrode surfaces, the reduction of hydrogen peroxide occurs at more positive potentials than the reduction of oxygen in unbuffered and alkaline electrolytes (cf. Fig. 1). This observation is in agreement with thermodynamic calculations which give the potentials for these reactions as 0.46 V vs. SHE for oxygen reduction and 0.97 V vs. SHE for HO 2 reduction at p H 13. Therefore, on the basis of these considerations, the lifetime of any hydrogen peroxide formed during the reduction of oxygen should be very short. By contrast, in acidic electrolyte a relatively large overpotential for the reduction of hydrogen peroxide exists; therefore peroxide production is substantially more efficient at low pH. Additional routes for the consumption of peroxide at silver are through mechanisms that involve catalytic disproportionation of H202 by chemical steps that may involve the participation of Ag + as postulated by Abel [20] and Weiss [21] via H202 + Ag o
~
Ag++ O H ' + O H -
(la)
O H - + H202
~
HO 2 + H20
(lb)
HO 2 + ag +
~
HO~ + Ag o
(lc)
O H ' + HO 2
~
H20 + O2
(ld)
Net:
~
2 H20 + O2
2 H202
or the local-cell or mixed-potential model proposed by Gerischer and Gerischer [22]: ~
H202
E ° / V s n E = 0.682
(2a)
HzO 2 + 2 e - + 2 H +
~
2 H20
E°/VsHE ----1.76
(2b)
Combining:
~
2 H 2 0 + 02
AG°ell = - 208 kJ m o l - 1
O 2 -'F
2 H++ 2 e2 H202
The mechanism of consumption of H202 produced during the reduction of oxygen at silver is unquestionably favored by silver surfaces free of adsorbates, as
108 evidenced by the relative levels of H202 recorded for prereduced and adsorbatecovered silver electrodes. These observations imply that routes involving catalytic participation of the electrode surfaces for decomposition of H202, which may involve specifically adsorbed oxygen-containing species, are necessary in considering possible mechanisms for reduction of oxygen at silver. Oxygen reduction at silver electrodes covered with adsorbed halides The stability of H202 at silver surfaces is greatly enhanced by a variety of species which "cling" to the electrode surface by either chemisorption or physisorption. With strong specific adsorption, as observed with iodide in the present work, the overall four-electron reaction is slowed at the step involving decomposition of H202 so that two two-electron steps are observed. The relative heights of the two waves show that H202 is produced at nearly 100% efficiency during the reduction of oxygen at silver preexposed to iodide in much the same way as when oxygen is reduced at mercury [23]. This conclusion is verified by ring-disk experiments (cf. Fig. 6). Radiotracer experiments have shown that during the reduction of 02 to H202, the O - O bond is not broken [24]. However, final reduction of H202 to H 2 0 (or O H - ) must involve oxygen atom dissociation. For this process to occur in the absence of a catalyst, 160-190 k J / m o l is required [25]. In the presence of iodide (and to a lesser extent with weaker adsorbates), the catalytic surface of the silver is effectively poisoned at potentials positive of the pzc of silver; therefore, the amount of energy required to break the oxygen-oxygen bond at the electrode surface is significantly increased. This increase in energy is reflected in the large negative overpotential (ca. - 0 . 9 V; i.e., AG, = 170 k J / m o l ) observed between reduction of H202 on prereduced and lads-covered silver. By contrast, rapid reduction of 02 to H202 begins as soon as the initially continuous, passiva_ting layer of AgI is reduced. The ESCA-measured. coverage of Iads necessary to produce the maximum amount of H202 during the reduction of oxygen at silver is in surprisingly good agreement with the finding by Kilty et al. that coverages of Clad s of - 0.25 monolayer provide the highest yield of ethylene oxide in the gas-phase catalytic oxidation of ethylene over silver treated with halogen-containing hydrocarbons [26]. For both the gas-phase and solution reactions, the presence of adsorbed halogens promotes the formation of dioxygen species by poisoning the catalytic surface of silver. Ordering of the halides according to their influence on the decomposition of hydrogen peroxide in the solution experiments can be explained by the influence of competitive adsorption for active surface sites by both solvent and dioxygen species. REFERENCES 1 J. Wilshire and D.T. Sawyer,Acc. Chem. Res., 12 (1979) 105. 2 E. Yeager, Progressin Batteriesand Solar Cells, Vol. 3, JEC Press, Cleveland, 1980, p. 238. 3 X.E. Verykios,F.P. Stein and R.W. Coughlin, Catal. Rev. Sci. Eng., 22 (1980) 197.
109 4 5 6 7 8
N.A. Shumilova, G.V. Zhutaeva and M.R. Tarasevich, Electrochim. Acta, 11 (1966) 967. H.S. Wroblowa, Y.C. Pan and G. Razumney, J. Electroanal. Chem., 69 (1976) 195. P. Fischer and J. Heitbaum, Electroanal. Chem., 112 (1980) 231. N.I. Dubrovina and L.N. Nekrasov, Elektrokhimiya, 8 (1972) 1503. (a) J.O'M Bockris and S. Srinivasan, J. Electroanal. Chem., 11 (1966) 350. (b) A.J. Appleby and M. Savy, J. Electroanal. Chem., 92 (1978) 15. 9 D.F. Untereker, Ph.D. Thesis, State Univ. of New York at Buffalo, 1973 (Univ. Microfilms No. 73-19, 247). 10 W.J. Albery and S. Bruckenstein, Trans. Faraday Soc., 62 (1966) 1920. 11 J.D. Lee, Rev. Sci. Instr., 43 (1972) 1291. 12 D.V. Sepa, M. Vojnovic and A. Damjanovic, Electrochim. Acta, 15 (1970) 1355. 13 B.E. Conway, H. Angerstein-Kozlowska, W.B.A. Sharp and E.E. Criddle, Anal. Chem., 45 (1973) 1331. 14 H.A. Liebhafsky, R. Furnichi and G.M. Roe, J. Am. Chem. Soc., 103 (1981) 51, and references therein. 15 See, for example: T.A. Carlson, Photoelectron and Auger Spectroscopy, Plenum Press, New York, 1975, Ch. 5, and W.M. Riggs and M.J. Parker in A.W. Czanderna (Ed.) Methods of Surface Analysis, Elsevier, New York, 1975, Ch. 4 and references therein. 16 J.H. Scofield, Report No. UCRL-5136, Lawrence Livermore Laboratory (1973); also in J. Electron Spectrosc. Relat. Phenom., 8 (1976) 129. 17 J. Clavilier, C.R. Acad. Sci., Ser. C, 263 (1966) 191. 18 G. Valette and A. Hamelin, J. Electroanal. Chem., 45 (1973) 301. 19 P. Delahay, Double Layer and Electrode Kinetics, Wiley-Interscience, New York, 1965, Ch. 3. 20 E. Abel, Monatsh. Chem., 83 (1952) 422. 21 J. Weiss, Advances in Catalysis and Related Subjects, Vol. 4, Academic Press, New York, 1952, p. 343. 22 R. Gerischer and H. Gerischer, Z. Phys. Chem. (Wiesbaden), 6 (1956) 178. 23 R. Cornelissen and L. Gierst, J. Electroanal. Chem., 3 (1962) 219. 24 M. Davies, M. Clarke, E. Yeager and F. Hovorka, J. Electrochem. Soc., 106 (1959) 56. 25 V.I. Vedeneyev, L.V. Gurvich, V.N. Kondrat'yev, V.A. Medvedev and Ye.L. Frankevich, Bond Energies, Ionization Potentials and Electron Affinities, Edward Arnold, 1966, p. 76. 26 P.A. Kilty, N.C. Rol and W.M.H. Sachtler in W. Hightower (Ed.), Proc. 5th Int. Congr. Catal., North-Holland, Amsterdam, 1973, p. 929.
97
ELECTROCHEMISTRY OF OXYGEN AND HYDROGEN PEROXIDE AT CLEAN AND AT HALIDE-COVERED POLYCRYSTALLINE SILVER
E. STEVEN BRANDT Research LaboratorieS, Eastman Kodak Company, Rochester, N Y 14650 (U.S.A.) (Received 27th July 1982; in revised form 14th December 1982)
ABSTRACT The effect of halide adsorption on the rate of hydrogen peroxide decomposition in alkaline solution has been studied. The reduction of oxygen at clean silver in acidic, alkaline, and unbuffered electrolytes has also been investigated for comparison. On prereduced electrodes the amount of hydrogen peroxide observed during the reduction of oxygen is related to pH and adventitious contamination. Specific adsorption by halides decreases the efficiency of the "four-electron" reduction of oxygen with concomitant formation of hydrogen peroxide. With iodide, the reduction of oxygen to hydrogen peroxide is nearly 100% efficient. Coverages of lad s approaching 0.2 monolayer are sufficient to create an energy barrier for breaking of the oxygen-oxygen bond and effectively poison the catalytic surface of silver.
INTRODUCTION
Considerable interest in the electrochemical reactions of oxygen and hydrogen peroxide at silver electrodes has been generated from fuel-cell development and from gas-phase heterogeneous catalysis by supported silver catalysts. In addition, reactions of hydrogen peroxide and oxygen at various electrode materials, including silver, are of general fundamental importance in electrocatalysis. Of particular importance in both fundamental and applied areas is the formation of diatomic oxygen (or dioxygen [1]) intermediates during the reaction of oxygen at silver surfaces. In fuel cells these intermediates, particularly hydrogen peroxide, can attack the electrodes oxidatively, leading to premature failure of the cell [2]. However, at least equally important for fuel-cell applications is the fact that the presence of stable dioxygen intermediates is an indication of intrinsic cell inefficiency in the reduction of oxygen at the cathode. On the other hand, the presence of dioxygen species is desired in the catalytic conversion of ethylene to ethylene oxide over silver, since these species are thought to increase catalytic specificity for the desired product [3]. In both systems the rate of oxygen-oxygen bond breaking determines the concentration of stable dioxygen intermediates. In solution the efficiency of this process can be followed by observing the formation of hydrogen peroxide during the reduction of oxygen using a rotating ring-disk electrode [4-7]. In the present study 0022-0728/83/$03.00
© 1983 Elsevier Sequoia S.A.
98 the reduction of oxygen has been investigated on clean silver electrodes, which favor the "four-electron" reduction process [8], and on silver surfaces that have been exposed to halide adsorbates, which significantly reduce the catalytic activity of the silver surface, thus producing conditions which favor formation of stable dioxygen intermediates. EXPERIMENTAL Voltammetric measurements were made using conventional commercially available instrumentation. For experiments with ring-disk electrodes, an in-house-built potentiostat was interfaced directly to the three-electrode potentiostat to provide independent control of two working electrodes [9]. Unless otherwise noted, the disk electrode had a measured area of 0.458 cm 2. The dimensions of the Pt-ring/Ag-disk electrode were taken as r 1 (disk diameter) 0.382 cm, r 2 (ring i.d.) 0.399 cm, and r 3 (ring o.d.) 0.422 cm, which translates to a disk area of 0.458 cm 2 and a ring area of 0.059 cm 2. F r o m these dimensions N, the theoretical collection efficiency, was calculated by the treatment of Albery and Bruckenstein [10] to be - 0.183. Potentials are reported vs. the H g / H g 2 S O 4 (sat. K2SO4) reference electrode, which had a measured potential of 0.371 V vs. SCE in a saturated K2SO 4 electrolyte. Before use, all glassware was cleaned in a 1 : 1 mixture of concentrated nitric and sulfuric acids. When not in use, the cells were also stored in this solution. Vacuum surface analysis of silver samples was performed on a DuPont 650B ESCA [11]. This instrument has a nonmonochromatic Mg K , (1253.6 eV) source, which is regulated at 300 W. System vacuum during these experiments was kept at < 1 × 10 -7 Torr. For ESCA analysis, Teflon-shrouded 6-mm-diameter silver disks were used as electrodes. After a given experiment the sample shroud was removed, and the sample, suspended in distilled water, was transferred to the ESCA. Before each use, silver samples were polished concentrically to 0.25 /~m on a h o b b y lathe with diamond pastes. Cross-contamination of ring and disk electrodes was thereby minimized. After they were polished, silver electrodes were rinsed throroughly with purified water and then cycled between background potential limits several times. While the electrode was under potential control at the lower limit, the electrolyte was aspirated. With the electrode open circuited, the cell was rinsed and refilled with fresh electrolyte. This process was continued until satisfactory i/E curves for the background and for the reduction of 02 at silver (see Results Section) were obtained. According to the terminology adopted by Sepa and co-workers, these samples are referred to as "prereduced" in the text [12]. All chemicals were reagent grade, except for K2SO 4 and K O H , which were A l f a / V e n t r o n ultrapure grade. Solutions were prepared with distilled water that had been passed through a carbon filter train and then redistilled in glass. Water prepared in this fashion was judged acceptable for use, according to published criteria [13].
99 RESULTS
Reduction of oxygen and hydrogen peroxide at clean silver electrodes Effect of pH T h e r e d u c t i o n of 02 a n d H 2 0 2 at p r e r e d u c e d r o t a t i n g silver disk electrodes at different p H ' s is shown in Fig. 1. T h e b a c k g r o u n d currents in H e - d e o x y g e n a t e d electrolytes are shown for c o m p a r i s o n (note c u r r e n t scales). A t p H 12.8 a n d in u n b u f f e r e d 0.05 mol 1-1 K z S O 4 ( p H 5.9), the wave associated with H202 r e d u c t i o n on silver occurs at a higher p o t e n t i a l t h a n the r e d u c t i o n of oxygen, in qualitative a g r e e m e n t with their respective t h e r m o d y n a m i c s t a n d a r d potentials. However, at p H 1.5, the r e d u c t i o n of H202 occurs at the s a m e potential, o r even at p o t e n t i a l s slightly negative relative to O z reduction.
Ring-disk experiments T h e r e d u c t i o n of oxygen on p r e r e d u c e d silver was p e r f o r m e d in 0.05 mol 1 1 K 2 S O 4 while the p r o d u c t s of that r e a c t i o n were m o n i t o r e d at the p l a t i n u m ring. F i g u r e 2a shows that two limiting c u r r e n t p l a t e a u s are o b s e r v e d in the ring
o
0 E / V ~I
::
" i
-10 ;
.
.
"
f
b
f 0
E / V "-,
.
I
/ %0 ~
.
l .
.
.
.
200~A I .............. IOOp.A
Fig. 1. Rotating silver disk voltammograms for 02 and H202 reduction on prereduced silver electrodes; (. . . . . . ) background in He-deoxygenated electrolyte; ( - ) O2 (saturated); ( . . . . . . ) - 10- 3 mol l - 1 H202. (a) 0.05 mol 1-1 H2SO4; (b) 0.05 tool 1 I K2SO4; (c) 0.1 mol 1- l KOH. All scans: 50 mY/s, 900 rpm.
100 a).
ir with disk at E d I I
<
÷ 1.0 I
0
J
I
-liO
,,
>
b, wx
/
~,
E/v
-I-/~5/~A
/
j
b.)
/ 2
J, \\
12OO/.L A -'i° ""-.
,
......~'.\
,
E/V d
"D"""""." |os A
]20~A"5\ . . . . . . . .
x
Fig. 2. (a) i r / E r curves obtained during the constant reduction of oxygen at Ed 'adjusted for ilaim. (b) Ring profiles obtained with E, at potentials a and b (note sensitivities). Conditions: O2-saturated 0.05 mol 1- l K2SO4, scan rate 50 mY/s, 900 rpra.
ir/E d
voltammogram while oxygen is being continually reduced at E o (limiting current density). Assuming n is equal for disk and ring processes, the limiting current observed at + 0.87 V (ilrim) is ca. 60% of the theoretical collection efficiency of 0.183 predicted by the electrode geometry (cf. Experimental Section). A much smaller current is observed at + 0.20 to - 0 . 2 5 V. When the ring electrode is potentiostated at these potentials and the disk electrode is scanned through the oxygen reduction wave, a continuous wave that mirrors the oxygen wave and a wave that goes through a m a x i m u m at - 1.12 V are observed with the ring at +0.87 V (potential a) and - 0 . 1 2 V (potential b), respectively (Fig. 2b). The ring current at - 0 . 1 2 V can be shown to be due to the formation of hydrogen peroxide by standard addition methods. Identification of the electroactive species responsible for the current at + 0.87 V is more difficult. This current might be due to the oxidation of O H - via 4 O H - ~ 02 + 2 H 2 0 + 4 e - , which would explain its magnitude, since this is the reverse of the four-electron oxygen reduction. This species is not observed in either strongly acidic or strongly basic electrolytes where the solutions are well-buffered. The relative amount of H202 produced at p H 12.8 is similar to that produced in unbuffered solutions and also goes through a maximum. In both electrolytes, the peak currents are somewhat irreproducible, since the amount of peroxide produced
101
~ 200/.~A
dc
Ed/V
(
i
I
)
-LO
Ir o
flOrA
/ Fig. 3. Hydrogen peroxide production during the reduction of oxygen in O2-saturated 0.05 rnol 1-1 H2SO4; all scans 50 m V / s , 900 rpm. Dashed horizontal line indicates predicted i/im based on results obtained in 0.1 tool 1-1 KOH.
appears to be related to the extent of contamination of the cell, in agreement with the results obtained by Fischer and Heitbaum [6]. Levels of H202 as low as 8 × 10 -6 mol 1-1 and as high as 2.4 x 10 -4 tool 1-1 have been recorded under apparently similar conditions. In acid solution (0.05 mol 1- l H2SO4) the background evolution of hydrogen overlaps the oxygen and hydrogen peroxide reduction waves at the disk (Fig. 1). However, the magnitude of i r shows that peroxide production is relatively efficient at this pH (Figs. 3 and 2b). This result demonstrates that the decomposition of peroxide generated during the reduction of oxygen at prereduced silver electrodes is much less efficient at low pH's than at alkaline pH or in unbuffered electrolyte. Reduction of oxygen at halide-covered silver electrodes
Mercury salts [6] and organic surfactants [6,7] have been used to demonstrate the effect of adsorption on the rate of peroxide production during the reduction of
102
oxygen at silver. However, anions in general (and the halides in particular) are more desirable candidates to explore the role of specific adsorption on cathodic faradaic processes at silver because (1) in general, metals adsorb anions much more readily than cations or neutral species; (2) in gas-phase catalysis work, halide-containing hydrocarbons are known to improve the selectivity of silver for production of ethylene oxide from ethylene, probably through increasing the surface concentration of dioxygen species [3]. Chloride The addition of KC1 to final concentrations of 1.0 mol l- 1 to a 0.1 mol 1-1 K O H electrolyte saturated in O: causes an increase in ir, attributed to the formation of H202 (Fig. 4). At C1- concentrations below 10-1 mol l-1 relatively minor increases of H202 (compared with prereduced electrodes) occur during the reduction of oxygen at the silver disk. Actually, at least two waves are observed at the ring electrode; however, both waves disappear when the solution is deoxygenated (Fig. 4b).
I/. ~"
...........
~'°
J
//
o),
.
] 2oo~A
//
gdc
" ~0.0 Ed/V < ir 0 "%'
,..,°..
b).
< ~'dc
Ed / V < ,~?.0 Ir 0
~"--L
'
If
-I.O
'
>
15 A Fig. 4. (a) Reduction of oxygen in 0.1 mol 1- l K O H only (1), and in 0.1 mol 1-1 K O H + 1 mol 1-1 KCI (2). (b) Background in deoxygenated electrolyte as used to obtain curve 2. All scans: 50 m V / s , 900 rpm.
103
Bromide The quantity of H202 produced during the reduction of oxygen increases systematically with B r - concentration and reaches a m a x i m u m near 5 × 10 -3 mol 1-1 B r - (Fig. 5). Phase AgBr, which is formed at the upper limit, is reduced before the onset of oxygen reduction and hence H 2 0 2 production, as evidenced by the small cathodic surface waves at ca. - 0 . 4 0 V. However, as shown by both ring and disk voltammograms, the effect of Brad s persists at the more cathodic potentials where these processes occur. At 10 - 3 mol 1-1 B r - (curve 6, Fig. 5) multiple inflections are observed in the oxygen wave at the disk. T h e most pronounced inflection, centered at ca. - 1.1 V, corresponds to the maximum production of H 2 0 2 . Higher concentrations of bromide cause only minor changes in the shape or position of the waves or the quantity of H202 monitored but contribute significantly to the dissolution of the silver disk
/ [! Ed/V < 'ro~iO'O
,
-:/~//////
200ffA I
///
i
' ..
I,0
.-~- >
~ \/,// //' / 2offAl,f
~//.a "~ / / /
/ •
Fig. 5. Effect of B r - on the a m o u n t of peroxide formed during the reduction of oxygen in 0.1 mol 1-1 K O H . Curves 1-5, successive additions of KBr (2 × 10-5 tool per addition); curves 6, to 4 × 10-3 M final concentration. Dotted line, background of prereduced silver in clean O2-saturated 0.1 mol l - 1 KOH. All scans: 50 m V / s , 900 rpm.
104
electrode when the potential is scanned towards positive potentials beyond the foot of the oxygen wave. lodide Iodide shows by far the most pronounced effect of the halides on the rate of reduction of oxygen at silver electrodes. Figure 6 was obtained by successive additions of KI to a 0.1 mol 1-1 K O H electrolyte. The maximum current, / m a x is observed at a much lower concentration (ca. 5 x 10 -6 mol 1-1) of iodide than with bromide or chloride. A t v e r y low concentrations, approaching 10 -8 mol 1-~, the effect of iodide is seen as a minimum in the ilim plateau near the point of zero charge (pzc) of silver. This minimum in the diffusion-limited mass-transport region suggests competition between two rate-determining processes, such as the formation of a passivating layer of lad s at the expense of a catalytically active surface layer of oxygen-containing species. With increasing iodide this minimum deepens and broadens on the anodic side until, at 5 x 10 - 6 mol 1- l I-, two distinct waves of approximately equal height are formed in the disk voltammogram. To eliminate the possibility of solution iodide participating in the reactions of hydrogen peroxide [14], the prereduced electrodes were exposed to iodide at open circuit in a cell separate from the analytical cell containing the purified oxygensaturated electrolyte. Under these conditions approximately one monolayer of lad s is formed, as measured by ESCA [15], over a wide range of concentrations (Table 1).
200ffA
// Ed/V <
ido
,
~'I.0
~
'r 0
~
~
)
" ~
I 20ffA Fig. 6. Successive additions of KI to an O2-saturated 0.1 mol 1- i K O H electrolyte to a final concentration of - 5 x 10 -6 mol 1- J KI. All scans: 50 m V / s , 900 rpm, negative-going sweep after addition.
105 TABLE 1 5-rain exposure in purified H 2 0 containing KI, while the silver sample was rotated at 900 rpm [ K I ] / m o l 1-1
I / A g × 100 a
1 10 - l 10 - 2 10 -3 10 - 4 10 -5 10 - 6
9.20 11.5 10.0 9.82 9.29 10.1 9.56
Average
9.83 _+0.83
a I 3d5/2 and Ag 3d5/2 corrected for cross sections by using the tables of Scofield [16].
The v o l t a m m o g r a m s obtained f r o m these electrodes were essentially the same as observed with iodide in solution, but after reversal at the lower limit the positive-going wave returns to the same c u r r e n t - v o l t a g e wave shape as observed on a prereduced electrode. ~"" ..............
!/o .I
,
,
Ed/V
,
0.o
I0
o/
J o/
.-
o ..........
- 1.0 o
o
s~'"~o
o
o_ a ~6
~
o
2 0
\ 0°~o
Fig. 7. Upper: Comparison between voltammograms for 02 reduction at clean (dotted line) and at externally iodide-treated silver electrodes. Iodide exposure, 5 min at o.c.; unbuffered. Electrode Area: 0.29 cm 2. Lower: (I 3 d s / 2 / A g 3d5/2)×100, corrected for cross sections; from ESCA spectra taken at indicated potentials (see text for experimental details).
106 i/mA. cm -2 C//.~F. cm-2
60 I I
\ W. c\ \
I0
40
'X
8 6
\ 20
4 2
r.p. ,
0
i
0.5
1.0
1,5
E/V
Fig. 8. Correlation of pzc with faradaic current observed during the reduction of oxygen on iodide-treated polycrystalline silver. (a) Reduction of oxygen at silver in O2-saturated 0.1 mol 1 i KOH after preexposure to 10 - 3 mol 1-1 KI for 5 min (unbuffered). (b) Capacity vs. potential curve obtained in Ar-deoxygenated 0.02 mol 1-1 NaF at 10 Hz, AEac = 10 mV peak to peak. (c) Capacity vs. potential curve obtained in Ar-deoxygenated 0.01 mol.1- l KOH at 10 Hz, AEac = 10 mV peak to peak. C vs. E scans, 10 mV/s, 0 rpm; i vs. E scans, 50 mV/s, 900 rpm; electrode area 0.29 cm2.
The coverage of lad s as a f u n c t i o n of potential was followed b y using E S C A in c o n j u n c t i o n with removal of Ida s b y potential step in a cell c o n t a i n i n g a large v o l u m e of purified 0.1 mol 1-1 K O H . T o ensure that steady state had b e e n achieved after the potential step, the electrode was potentiostated for 30 min, d u r i n g which time the c u r r e n t vs. time behavior was recorded. As m e a s u r e d b y the increase in current, the steady state is reached rapidly, o n the order of milliseconds, although some slow d e s o r p t i o n c o n t i n u e s up to 1 min. T h e isotherm o b t a i n e d b y E S C A from electrodes that had been removed while still u n d e r potential control is shown in Fig. 7 along with the p o i n t - b y - p o i n t steady-state results. Coverages of Iads decrease with increasing cathodic potentials u n t i l ca. - 1.1 V, where the a m o u n t of lad s r e m a i n i n g is _< 0.2 m o n o l a y e r , which is the practical detection limit of the ESCA technique. A t this potential, the c u r r e n t for the reduction of oxygen begins to increase, f o r m i n g the second wave observed in the
107 voltammogram, corresponding to the reduction of peroxide. The ratio of the two waves is 1.1 : 1, the first wave being the larger, a ratio that is invariant at sweep rates < 100 mV/s. Using an alternating current technique similar to that described by Clavilier [17], Valette and Hamelin [18], we recorded the differential capacity of polycrystalline silver in dilute electrolytes of 0.02 mol 1- l NaF and 0.01 mol 1 1 N a O H and compared the results with the faradaic process of oxygen reduction on an iodidetreated silver electrode (Fig. 8). Good agreement exists between the experimentally determined pzc, which, according to the G o u y - C h a p m a n - S t e r n double-layer model [19], is reflected by the capacitance minimum observed at ca. - 1 . 3 V and the potential at which i lim for a four-electron wave is reached on iodide-treated polycrystalline silver electrodes. DISCUSSION
Oxygen reduction at clean (prereduced) silver electrodes At prereduced silver electrode surfaces, the reduction of hydrogen peroxide occurs at more positive potentials than the reduction of oxygen in unbuffered and alkaline electrolytes (cf. Fig. 1). This observation is in agreement with thermodynamic calculations which give the potentials for these reactions as 0.46 V vs. SHE for oxygen reduction and 0.97 V vs. SHE for HO 2 reduction at p H 13. Therefore, on the basis of these considerations, the lifetime of any hydrogen peroxide formed during the reduction of oxygen should be very short. By contrast, in acidic electrolyte a relatively large overpotential for the reduction of hydrogen peroxide exists; therefore peroxide production is substantially more efficient at low pH. Additional routes for the consumption of peroxide at silver are through mechanisms that involve catalytic disproportionation of H202 by chemical steps that may involve the participation of Ag + as postulated by Abel [20] and Weiss [21] via H202 + Ag o
~
Ag++ O H ' + O H -
(la)
O H - + H202
~
HO 2 + H20
(lb)
HO 2 + ag +
~
HO~ + Ag o
(lc)
O H ' + HO 2
~
H20 + O2
(ld)
Net:
~
2 H20 + O2
2 H202
or the local-cell or mixed-potential model proposed by Gerischer and Gerischer [22]: ~
H202
E ° / V s n E = 0.682
(2a)
HzO 2 + 2 e - + 2 H +
~
2 H20
E°/VsHE ----1.76
(2b)
Combining:
~
2 H 2 0 + 02
AG°ell = - 208 kJ m o l - 1
O 2 -'F
2 H++ 2 e2 H202
The mechanism of consumption of H202 produced during the reduction of oxygen at silver is unquestionably favored by silver surfaces free of adsorbates, as
108 evidenced by the relative levels of H202 recorded for prereduced and adsorbatecovered silver electrodes. These observations imply that routes involving catalytic participation of the electrode surfaces for decomposition of H202, which may involve specifically adsorbed oxygen-containing species, are necessary in considering possible mechanisms for reduction of oxygen at silver. Oxygen reduction at silver electrodes covered with adsorbed halides The stability of H202 at silver surfaces is greatly enhanced by a variety of species which "cling" to the electrode surface by either chemisorption or physisorption. With strong specific adsorption, as observed with iodide in the present work, the overall four-electron reaction is slowed at the step involving decomposition of H202 so that two two-electron steps are observed. The relative heights of the two waves show that H202 is produced at nearly 100% efficiency during the reduction of oxygen at silver preexposed to iodide in much the same way as when oxygen is reduced at mercury [23]. This conclusion is verified by ring-disk experiments (cf. Fig. 6). Radiotracer experiments have shown that during the reduction of 02 to H202, the O - O bond is not broken [24]. However, final reduction of H202 to H 2 0 (or O H - ) must involve oxygen atom dissociation. For this process to occur in the absence of a catalyst, 160-190 k J / m o l is required [25]. In the presence of iodide (and to a lesser extent with weaker adsorbates), the catalytic surface of the silver is effectively poisoned at potentials positive of the pzc of silver; therefore, the amount of energy required to break the oxygen-oxygen bond at the electrode surface is significantly increased. This increase in energy is reflected in the large negative overpotential (ca. - 0 . 9 V; i.e., AG, = 170 k J / m o l ) observed between reduction of H202 on prereduced and lads-covered silver. By contrast, rapid reduction of 02 to H202 begins as soon as the initially continuous, passiva_ting layer of AgI is reduced. The ESCA-measured. coverage of Iads necessary to produce the maximum amount of H202 during the reduction of oxygen at silver is in surprisingly good agreement with the finding by Kilty et al. that coverages of Clad s of - 0.25 monolayer provide the highest yield of ethylene oxide in the gas-phase catalytic oxidation of ethylene over silver treated with halogen-containing hydrocarbons [26]. For both the gas-phase and solution reactions, the presence of adsorbed halogens promotes the formation of dioxygen species by poisoning the catalytic surface of silver. Ordering of the halides according to their influence on the decomposition of hydrogen peroxide in the solution experiments can be explained by the influence of competitive adsorption for active surface sites by both solvent and dioxygen species. REFERENCES 1 J. Wilshire and D.T. Sawyer,Acc. Chem. Res., 12 (1979) 105. 2 E. Yeager, Progressin Batteriesand Solar Cells, Vol. 3, JEC Press, Cleveland, 1980, p. 238. 3 X.E. Verykios,F.P. Stein and R.W. Coughlin, Catal. Rev. Sci. Eng., 22 (1980) 197.
109 4 5 6 7 8
N.A. Shumilova, G.V. Zhutaeva and M.R. Tarasevich, Electrochim. Acta, 11 (1966) 967. H.S. Wroblowa, Y.C. Pan and G. Razumney, J. Electroanal. Chem., 69 (1976) 195. P. Fischer and J. Heitbaum, Electroanal. Chem., 112 (1980) 231. N.I. Dubrovina and L.N. Nekrasov, Elektrokhimiya, 8 (1972) 1503. (a) J.O'M Bockris and S. Srinivasan, J. Electroanal. Chem., 11 (1966) 350. (b) A.J. Appleby and M. Savy, J. Electroanal. Chem., 92 (1978) 15. 9 D.F. Untereker, Ph.D. Thesis, State Univ. of New York at Buffalo, 1973 (Univ. Microfilms No. 73-19, 247). 10 W.J. Albery and S. Bruckenstein, Trans. Faraday Soc., 62 (1966) 1920. 11 J.D. Lee, Rev. Sci. Instr., 43 (1972) 1291. 12 D.V. Sepa, M. Vojnovic and A. Damjanovic, Electrochim. Acta, 15 (1970) 1355. 13 B.E. Conway, H. Angerstein-Kozlowska, W.B.A. Sharp and E.E. Criddle, Anal. Chem., 45 (1973) 1331. 14 H.A. Liebhafsky, R. Furnichi and G.M. Roe, J. Am. Chem. Soc., 103 (1981) 51, and references therein. 15 See, for example: T.A. Carlson, Photoelectron and Auger Spectroscopy, Plenum Press, New York, 1975, Ch. 5, and W.M. Riggs and M.J. Parker in A.W. Czanderna (Ed.) Methods of Surface Analysis, Elsevier, New York, 1975, Ch. 4 and references therein. 16 J.H. Scofield, Report No. UCRL-5136, Lawrence Livermore Laboratory (1973); also in J. Electron Spectrosc. Relat. Phenom., 8 (1976) 129. 17 J. Clavilier, C.R. Acad. Sci., Ser. C, 263 (1966) 191. 18 G. Valette and A. Hamelin, J. Electroanal. Chem., 45 (1973) 301. 19 P. Delahay, Double Layer and Electrode Kinetics, Wiley-Interscience, New York, 1965, Ch. 3. 20 E. Abel, Monatsh. Chem., 83 (1952) 422. 21 J. Weiss, Advances in Catalysis and Related Subjects, Vol. 4, Academic Press, New York, 1952, p. 343. 22 R. Gerischer and H. Gerischer, Z. Phys. Chem. (Wiesbaden), 6 (1956) 178. 23 R. Cornelissen and L. Gierst, J. Electroanal. Chem., 3 (1962) 219. 24 M. Davies, M. Clarke, E. Yeager and F. Hovorka, J. Electrochem. Soc., 106 (1959) 56. 25 V.I. Vedeneyev, L.V. Gurvich, V.N. Kondrat'yev, V.A. Medvedev and Ye.L. Frankevich, Bond Energies, Ionization Potentials and Electron Affinities, Edward Arnold, 1966, p. 76. 26 P.A. Kilty, N.C. Rol and W.M.H. Sachtler in W. Hightower (Ed.), Proc. 5th Int. Congr. Catal., North-Holland, Amsterdam, 1973, p. 929.