- Email: [email protected]
Observations of the diffusion coefficient of the perhydroxyl ion (HO−2) in lithium hydroxide solutions
OF THE DIFFUSION COEFFICIENT OBSERVATIONS OF THE PERHYDROXYL ION (HO,) IN LITHIUM HYDROXIDE SOLUTIONS E. L. LITTAUER and K. C. TSAI Laboratory, Lockheed Missiles & Spaa Company,
Lockheed Palo Alto Research
CA 94304, U.S.A.
(Receiwd
forpubiicattin 23
January
Inc., Palo Alto,
1979)
Abstract -The diffusioncoefficientofthe perhydroxylion (HO;) was dctermincd in 4.5 M LiOH supporting electrolyte.Over the range of concentration 0.2-0.8 M H,O, it was found to be independent ofconcentration and had a value of 5 * 0.2 x 10m6 cm’/s at 293 K. The experimentaldata were obtained cbronoamperometrically used a rde. At the limiting current, O2 evolution was found to be almost entirely supprcsscd and the current limiting species was therefore assumed to be the HOL ion. The electrochemical mechanism for the perhydroxyl
reduction
reaction
was invoked
to explain the experimental
2HO;
INTRODUCTlON
battery that H, evolution from the anode and O2 evolution from the cathode can be virtually eliminated if the cell is polarized sufficiently. The electrochemistry of the Li anode has been dealt with elsewhere in some detail[l]. This paper presents data obtained as a part of a Iarger program designed to provide insight into the electrochemistry of the cathode reaction. The selection of an electrode at which~minimal gassing occurs with least polarization was the practical goal. Cell tests revealed that operating at near to the limiting cathodic current reduces the O2 evolution rate significantly. It was also observed that, although the extent of cathode polarization was influenced by the electrode composition, under controlled hydrodynamic conditions, the limiting current was independent of this factor and was affected only bv the peroxide concentration in the electrolyte. These observations lead to the belief that H,O, reduction in LiOH solution follows the electrochemical mechanism such as that proposed by Gerischer[Z]. To model the Li-H,O, system requires knowledge ofthe diffusion coefficient of H,O, in alkaline solution and specifically in LiOH. No data have been published on this, presumably because ofthe difficulty in making unambiguous measurements. In acidsolution, H,O, is stable and presents no experimental difficulties but in alkaline solution it is transformed into the perhydroxyl ion HO; : + OH-
* HO;
+ H,O
+ ZOH- + O,t.
(2)
Thus conventional diffusion measurement techniques are confounded. The capillary tube apparatus suffers interference from 0, bubbles and the platinum electrodes used in conductivity cells also become covered with the gas. However, the fact that a clearcut limiting current (iJ is observed indicates that the diffusion coefficient of the current limiting species may be derived from chronoamperograms obtained with a rde. The experiments described here were designed to permit discrimination of the limiting species, ie HO; us 0,. A high concentration of supporting electrolyte (4.5 M LiOH) ensured that migration effects would be avoided.
It has been observed in studies of the Li-H,O,
H,O,
results.
RXPEEUIUIVWAL
The instrumentation, which comprised a Pine Instrument rotating disc apparatus in conjunction with a Princeton ADDlied Research PotentiostatBalvanostat, and the t&t cell have been described pre;iously[3]. Catalytic surfaces of An, Ni, Pd and Ag were obtained by electrodeposition onto a O&i-cm2 Pt rde. The Pt substrate was prefinished by polishing with 1 w diamond paste. Proprietary plating solutions listed below* were used. To ensure that all electrodes posscsscd the same surEace roughness, they were polished using a metal finishing fiber and rinsed with pentene and triply distilled water prior to test. Electrolytes comprised 99.95% LiOH (MCB) dissolved in triply distilled water. Precise amounts of 30”/, H,O, (J. T. Baker electronic grade) were added to 4.5 M LiOH prior to each experiment. The solution and test cell were thermostated at 2O+O.l”C. The chronoarnperograms were obtained at a scan rate of 5 mV/s. At the termination of the cathodic scan, the electrode was allowed to return to OCV. Significant hysteresis was observed as the electrode depolarized indicating that a change in the surface characteristics had occurred. After some minutes the original surface condition was presumably attained because the initial OCV was noted and a second or third polarization
(I)
K M”c = 1.5 x 102, and subsequently within the bulk solution and especially at a catalytic surface:
* Pd - Engelhard ES01 process. Ni - Barrett sulfamate process. Au - W-Rex Temperex HD process. Ag - Silver cyanide bath. see eg Metal Finishing, p. 305, Metals and Plastics Publications, Inc., New York (1977). 681
E. L. LITTAUERAND K. C. THAI
682
-1.25
I
0
20
I 40
1
I 80
I
MI
i (n&cm
IM
I 120
RESULTS
In strongly alkaline solutions, H20, is transformed almost completely into the HO;, perhydroxyl ion, which is then reduced electrochemically. Typical chronoamperograms for the electrochemical reduction of HO; on the catalysts are given in Fig. 1. It was found that at a scan rate of 5 mV/s or less a quasi steady state i-E curve was obtained which was independent of scan rate. The Ag An and Pd electrodes behaved somewhat differently from Ni. These electrodes were superior catalytically as will be noted from their higher OCV’s but they manifested substantial instability at the higher polarizations associated with the limiting perhydroxyl reduction. The electrode instability is believed caused by the onset of an alternate, simultaneous cathodic process. Most likely this process is the reduction of surface oxides. The Ni electrode, despite its poor catalytic properties, showed the most stable behavior. An unambiguous limiting current was always observed. This electrode has been characterized in some detail[5] and it is generally accepted that steady state highly reproducible surEace conditions are attained with it[6]. It should be noted however that all the electrodes did suffer a change in their surface characteristics as a result oftheir polarization at the limiting current. This
I
160
‘)
Fig. 1. Chronoemperograms of HO; reduction at a rde in 4.5 M LiOH = 1500 rev m&l. scan rate = 5 mV/s.
scan was of similar configuration. In fact, by following this technique, reproducibility was such that in most cases the curves would be superimposed on one another. To cover a broad range of experimental conditions, four concentrations of H,O, (0.2, 0.4, 0.6 and 0.8 M) and eight rotation speeds (100,500, 1000, 1500, 2000, 2500, 3000 and 4OOOrev min-I) were investigated. The measurements of solution kinematic viscosity was performed in triplicate using Shoemaker’s viscometer technique[4]. Reproducibility was within & 2%.
I
140
+ 0.2 M H,O,.
T = 293 K, N
appeared as a hysteresis on terminating the scan, but on standing for some minutes at OCV, the electrodes reverted to their original condition. Of particular interest to this investigation is the fact that although the rates of catalytic decomposition of perhydroxyl ion at the various electrodes is known to be vastly dint at OCV[7J the limiting currents were obsexved to be virtually identical. The dependence of the limiting currents on rotation speed with various concentrations ofHO; is plotted in Fig. 2. Table 1 contains the experimentally determined kinematic viscosities. With these data, the Levich equation ii = 0.62nFCbD2i’v~1’b~1i2 9 where
i, = n = F = C, = Y= D =
limiting current density (A/cm’) charge transfer number (es/mole) Faraday (96,500 A s/es.) Bulk concentration (mole/en?) Kinematic viscosity (cm2/s) Diffusion coefficient of the current limiting species (cm’/s) and w = Angular velocity of the disc given by u) = 2nN, N = rev s-l,
may be used to calculate the diffusion coefficient as a function of HO; concentration. These diffusion coefficients are also given in Table 1. Table 1. Kinemsric EHWM 0.2 0.4 0.6 0.8
viscosity
and diffusion coefficient
vhm2/s) 0.0276 0.0279 0.0282 0.0283
data
DUO, x lo6 cm2/s 5.0 4.8 5.2 5.0
Diffusion
coefficient
of (HO; ) in LiOH solutions
683
80cu-
&o.P
40-
20-
0
0
10
20
40
DLSCIJSSION The specific mechanism of the catalytic decomposition of HO; is a subject of some controversy[8]. The problem is further complicated when the catalyst is within an electric field. Moreover, the method of pretreatment of the catalytic surface plays an important role in influencing its behavior. For this reason, great care was taken in this investigation to ensure that each electrode was pretreated identically. Two basic mechanisms have been proposed to expiain the decomposition of peroxide z+ a metal surface. in alkaline solution. In the electrochemical mechanism[2,9], the decomposition is caused by simultaneous electro-oxidation of HO; to O2 and electrorednction of HO; to OH- :
HO;
+ OH-
HO;
+H,O+2e-
2HO;
#O,
* 0,
+ H,O + 2e-
(3)
*30H-
(4)
+ 20H-.
(5)
In the chemical me.chanism[lO, 111, the deeomposition of HO; involves the simultaneous chemical formation and reduction of metal surface oxides : Me + HO;
+ Me(O) + OH-
Me(O)fHO;#MefO,+OH(overall)
ZHO; &O,
60
between limiting current density, rotation speed and concentration of H,O, LiOH. T = 293 K.
Fig. 2. Relationship
(overall)
50
+ 20H-
(6) (7)
.
(5)
For the present investigation we have found. that by using the electrochemical model, our experimental results can be satisfactorily explained. Figure 3 depicts a mixed potential diagram which is constructed based on this model. It comprises the reversible polarization curves for the couples O,-HO,(I) and HO;-OH-(III) as represented by (3) and (4). The
in 4.5 M
reversible potentials @O’s) for these reactioqs are shown. The mixed potential curve obtained from the combination of the two redox couples is shown as (II). It reptesents the rate of formation and decomposition of HO;. It should be pointed out here that the two redox curves are illustrative only. They are not constructed from experimental data because this information is mostly unavailable for the catalysts of concern to this study. However, it is believed that the curve configuration is a reasonable facsimile of the actual situation. From the figure it may be seen that when the cathodic reduction of HO; reaches its limiting,current, curves II and III tend to merge, and the 0, evolution reactions tends to zero current. Thus at the limiting current, 0, gas generation does not perturb the reduction of peroxide and reliable experimental measurements can be obtained. In summary, inspection of Fig. 3 leads to the fol[owing observations : (i) The “better” catalyst which gives a higher decomposition rate also results in a higher OCV. (ii) Since the limiting diffusion current is related to the electrode surface area (A), the bulk concentration of HO;(C,) and the effective diffusion layer thickness (6) by the equation nFDA i, = ~ *C, 6 (where the other terms have their usual significance), the value of i, should be dependent only on the pretreat-ent of the catalyst surface, ie its surface area if 6 and C, are kept constant. In other words, the intrinsic catalytic activity of the electrode material does not play a role in the determination of the diffusion coefficient from the limiting current data. (iii) At large overvoltages, ie q > 3OOmV, the
E. L.
684
LB-fAUER
AND
: hI
HO; +
L I+,
+
OH-- O2+ Hz0 Zql
H20+ Z.- - HO;+ OH-
Fig. 3, Schematic&d
potential diigrm
and Hoi-OH-.
for
decomposition of H,O,
reduction reaction approaches looo/, and the measured limiting current represents that reaction exclusively. Returning to the experimental curves of Fig. 1, it will he noted that the observed limiting current is virtually independent of the composition of the electrode surface and presumably therefore the surface area of each electrode was essentiaIlL identical. The diffusion
coefficient of 5 x 1W6 cm*/s (f 5%) was found to be
independent of peroxide concentration investigated.
bassd
on redo% COUPES
O,-HO;
E, represents the mixed poteatial at OCV.
current efficiency of the HO;
over the range
CONCLUSION Addition of H,02 to alkaline solution results in conversion to the HO; ion, and this then reacts so that oxygen evolution occurs virithin the bulk solution and especially at any catalytic surf& hnmers al in it. Because of this inherent instability in alkal& s&ion, conventional methods for experimentally determining Dxo; _are unsuitable. The observation that by ekctrochenucally reducing the ion at its limiting current will result in almost total inhibition of 0, evolution at the electrode, has provided a means to obtain the’dif&ion coefficient. Chronoamperometric polarization at a rde gave virtually identical i,‘s irrespective of the catalyst used, provided the surfficc smoothuess was the same for all electrc&s. This observation provides new information to the controversy surrounding the mcchanism of hydrogen peroxide reduction in aIk&ne solution. To this end a mixed potential model based on the two operating ralox reactions has been con-
structed to illustrate the situation. Inserting the cxpeinto this mode1 tends to support the electrochemical mechanism.
rimental observations
K. C. Tut
The investigation has provided a value for the di&sion coefficient of fairly concentrated solutions of perhydroxyl ion in LiOH. This coefficient has not, to the best of the authors’ knowledge, been reported elsewhere. A simii value could be anticipated with other strongly alkaline solutions. The observed inde pendence of the d#usion coefficient with perhydroxyl ion concentrations would be expected because excess supporting electrolyte was used. Further work is desirable to co&m the value bf I&,; at lower concentrations than those in this study. REFERENCES
1. E. L! Littauer. K. C..Tsai and R. P. Hotlandsworth, J. electrdtem. Sot. 11, (6) 845 (1978). 2. R. Gerischer and H. G&cl&, Z:phy~. Chem. 6, 178 11936). 3. E. L. Littaucr aud K. C. Tsai, EIectrockim. Acta 2d, 351 (1979). 4. D. P. Shoemaker and C. W. Garland, Exp&ne~~s in Physicul Chemistry, (2nd edn) pp. 28&281. McGrawHill, New York (1967). N. Sato and G. Okarnoto, Ekctrochim. Acta 10, 495 (1965). B. E. Conway, hi. A. Sattar and D. Gilroy, ibid 14,677 (1969). T. Katan, M. Chin arid F. J. Schoeaweis. Nrrture, L.ond. 4951. 1281 11964). N. ,k Shun&va.‘G. V. Zhutaeva, N. D. Mcrkilova, E. I. Krus&cva, G. B. Samoilov, and V. S. Baaotskii, Kinetika i Kataliz 14, {l) 168 (i973j. 9. D. Wiilmau, Z. Elek.trochem 60,731 (1956) 10. V. S. Bngotskii and I. E. Yablokova, Zh.fiz. Khim. 27, 1663 (1953). 11. F. Ha&r &I S. Grinbarg, Z. anorg. allgm. Ch. 18.37 (1938).
Lockheed Palo Alto Research
CA 94304, U.S.A.
(Receiwd
forpubiicattin 23
January
Inc., Palo Alto,
1979)
Abstract -The diffusioncoefficientofthe perhydroxylion (HO;) was dctermincd in 4.5 M LiOH supporting electrolyte.Over the range of concentration 0.2-0.8 M H,O, it was found to be independent ofconcentration and had a value of 5 * 0.2 x 10m6 cm’/s at 293 K. The experimentaldata were obtained cbronoamperometrically used a rde. At the limiting current, O2 evolution was found to be almost entirely supprcsscd and the current limiting species was therefore assumed to be the HOL ion. The electrochemical mechanism for the perhydroxyl
reduction
reaction
was invoked
to explain the experimental
2HO;
INTRODUCTlON
battery that H, evolution from the anode and O2 evolution from the cathode can be virtually eliminated if the cell is polarized sufficiently. The electrochemistry of the Li anode has been dealt with elsewhere in some detail[l]. This paper presents data obtained as a part of a Iarger program designed to provide insight into the electrochemistry of the cathode reaction. The selection of an electrode at which~minimal gassing occurs with least polarization was the practical goal. Cell tests revealed that operating at near to the limiting cathodic current reduces the O2 evolution rate significantly. It was also observed that, although the extent of cathode polarization was influenced by the electrode composition, under controlled hydrodynamic conditions, the limiting current was independent of this factor and was affected only bv the peroxide concentration in the electrolyte. These observations lead to the belief that H,O, reduction in LiOH solution follows the electrochemical mechanism such as that proposed by Gerischer[Z]. To model the Li-H,O, system requires knowledge ofthe diffusion coefficient of H,O, in alkaline solution and specifically in LiOH. No data have been published on this, presumably because ofthe difficulty in making unambiguous measurements. In acidsolution, H,O, is stable and presents no experimental difficulties but in alkaline solution it is transformed into the perhydroxyl ion HO; : + OH-
* HO;
+ H,O
+ ZOH- + O,t.
(2)
Thus conventional diffusion measurement techniques are confounded. The capillary tube apparatus suffers interference from 0, bubbles and the platinum electrodes used in conductivity cells also become covered with the gas. However, the fact that a clearcut limiting current (iJ is observed indicates that the diffusion coefficient of the current limiting species may be derived from chronoamperograms obtained with a rde. The experiments described here were designed to permit discrimination of the limiting species, ie HO; us 0,. A high concentration of supporting electrolyte (4.5 M LiOH) ensured that migration effects would be avoided.
It has been observed in studies of the Li-H,O,
H,O,
results.
RXPEEUIUIVWAL
The instrumentation, which comprised a Pine Instrument rotating disc apparatus in conjunction with a Princeton ADDlied Research PotentiostatBalvanostat, and the t&t cell have been described pre;iously[3]. Catalytic surfaces of An, Ni, Pd and Ag were obtained by electrodeposition onto a O&i-cm2 Pt rde. The Pt substrate was prefinished by polishing with 1 w diamond paste. Proprietary plating solutions listed below* were used. To ensure that all electrodes posscsscd the same surEace roughness, they were polished using a metal finishing fiber and rinsed with pentene and triply distilled water prior to test. Electrolytes comprised 99.95% LiOH (MCB) dissolved in triply distilled water. Precise amounts of 30”/, H,O, (J. T. Baker electronic grade) were added to 4.5 M LiOH prior to each experiment. The solution and test cell were thermostated at 2O+O.l”C. The chronoarnperograms were obtained at a scan rate of 5 mV/s. At the termination of the cathodic scan, the electrode was allowed to return to OCV. Significant hysteresis was observed as the electrode depolarized indicating that a change in the surface characteristics had occurred. After some minutes the original surface condition was presumably attained because the initial OCV was noted and a second or third polarization
(I)
K M”c = 1.5 x 102, and subsequently within the bulk solution and especially at a catalytic surface:
* Pd - Engelhard ES01 process. Ni - Barrett sulfamate process. Au - W-Rex Temperex HD process. Ag - Silver cyanide bath. see eg Metal Finishing, p. 305, Metals and Plastics Publications, Inc., New York (1977). 681
E. L. LITTAUERAND K. C. THAI
682
-1.25
I
0
20
I 40
1
I 80
I
MI
i (n&cm
IM
I 120
RESULTS
In strongly alkaline solutions, H20, is transformed almost completely into the HO;, perhydroxyl ion, which is then reduced electrochemically. Typical chronoamperograms for the electrochemical reduction of HO; on the catalysts are given in Fig. 1. It was found that at a scan rate of 5 mV/s or less a quasi steady state i-E curve was obtained which was independent of scan rate. The Ag An and Pd electrodes behaved somewhat differently from Ni. These electrodes were superior catalytically as will be noted from their higher OCV’s but they manifested substantial instability at the higher polarizations associated with the limiting perhydroxyl reduction. The electrode instability is believed caused by the onset of an alternate, simultaneous cathodic process. Most likely this process is the reduction of surface oxides. The Ni electrode, despite its poor catalytic properties, showed the most stable behavior. An unambiguous limiting current was always observed. This electrode has been characterized in some detail[5] and it is generally accepted that steady state highly reproducible surEace conditions are attained with it[6]. It should be noted however that all the electrodes did suffer a change in their surface characteristics as a result oftheir polarization at the limiting current. This
I
160
‘)
Fig. 1. Chronoemperograms of HO; reduction at a rde in 4.5 M LiOH = 1500 rev m&l. scan rate = 5 mV/s.
scan was of similar configuration. In fact, by following this technique, reproducibility was such that in most cases the curves would be superimposed on one another. To cover a broad range of experimental conditions, four concentrations of H,O, (0.2, 0.4, 0.6 and 0.8 M) and eight rotation speeds (100,500, 1000, 1500, 2000, 2500, 3000 and 4OOOrev min-I) were investigated. The measurements of solution kinematic viscosity was performed in triplicate using Shoemaker’s viscometer technique[4]. Reproducibility was within & 2%.
I
140
+ 0.2 M H,O,.
T = 293 K, N
appeared as a hysteresis on terminating the scan, but on standing for some minutes at OCV, the electrodes reverted to their original condition. Of particular interest to this investigation is the fact that although the rates of catalytic decomposition of perhydroxyl ion at the various electrodes is known to be vastly dint at OCV[7J the limiting currents were obsexved to be virtually identical. The dependence of the limiting currents on rotation speed with various concentrations ofHO; is plotted in Fig. 2. Table 1 contains the experimentally determined kinematic viscosities. With these data, the Levich equation ii = 0.62nFCbD2i’v~1’b~1i2 9 where
i, = n = F = C, = Y= D =
limiting current density (A/cm’) charge transfer number (es/mole) Faraday (96,500 A s/es.) Bulk concentration (mole/en?) Kinematic viscosity (cm2/s) Diffusion coefficient of the current limiting species (cm’/s) and w = Angular velocity of the disc given by u) = 2nN, N = rev s-l,
may be used to calculate the diffusion coefficient as a function of HO; concentration. These diffusion coefficients are also given in Table 1. Table 1. Kinemsric EHWM 0.2 0.4 0.6 0.8
viscosity
and diffusion coefficient
vhm2/s) 0.0276 0.0279 0.0282 0.0283
data
DUO, x lo6 cm2/s 5.0 4.8 5.2 5.0
Diffusion
coefficient
of (HO; ) in LiOH solutions
683
80cu-
&o.P
40-
20-
0
0
10
20
40
DLSCIJSSION The specific mechanism of the catalytic decomposition of HO; is a subject of some controversy[8]. The problem is further complicated when the catalyst is within an electric field. Moreover, the method of pretreatment of the catalytic surface plays an important role in influencing its behavior. For this reason, great care was taken in this investigation to ensure that each electrode was pretreated identically. Two basic mechanisms have been proposed to expiain the decomposition of peroxide z+ a metal surface. in alkaline solution. In the electrochemical mechanism[2,9], the decomposition is caused by simultaneous electro-oxidation of HO; to O2 and electrorednction of HO; to OH- :
HO;
+ OH-
HO;
+H,O+2e-
2HO;
#O,
* 0,
+ H,O + 2e-
(3)
*30H-
(4)
+ 20H-.
(5)
In the chemical me.chanism[lO, 111, the deeomposition of HO; involves the simultaneous chemical formation and reduction of metal surface oxides : Me + HO;
+ Me(O) + OH-
Me(O)fHO;#MefO,+OH(overall)
ZHO; &O,
60
between limiting current density, rotation speed and concentration of H,O, LiOH. T = 293 K.
Fig. 2. Relationship
(overall)
50
+ 20H-
(6) (7)
.
(5)
For the present investigation we have found. that by using the electrochemical model, our experimental results can be satisfactorily explained. Figure 3 depicts a mixed potential diagram which is constructed based on this model. It comprises the reversible polarization curves for the couples O,-HO,(I) and HO;-OH-(III) as represented by (3) and (4). The
in 4.5 M
reversible potentials @O’s) for these reactioqs are shown. The mixed potential curve obtained from the combination of the two redox couples is shown as (II). It reptesents the rate of formation and decomposition of HO;. It should be pointed out here that the two redox curves are illustrative only. They are not constructed from experimental data because this information is mostly unavailable for the catalysts of concern to this study. However, it is believed that the curve configuration is a reasonable facsimile of the actual situation. From the figure it may be seen that when the cathodic reduction of HO; reaches its limiting,current, curves II and III tend to merge, and the 0, evolution reactions tends to zero current. Thus at the limiting current, 0, gas generation does not perturb the reduction of peroxide and reliable experimental measurements can be obtained. In summary, inspection of Fig. 3 leads to the fol[owing observations : (i) The “better” catalyst which gives a higher decomposition rate also results in a higher OCV. (ii) Since the limiting diffusion current is related to the electrode surface area (A), the bulk concentration of HO;(C,) and the effective diffusion layer thickness (6) by the equation nFDA i, = ~ *C, 6 (where the other terms have their usual significance), the value of i, should be dependent only on the pretreat-ent of the catalyst surface, ie its surface area if 6 and C, are kept constant. In other words, the intrinsic catalytic activity of the electrode material does not play a role in the determination of the diffusion coefficient from the limiting current data. (iii) At large overvoltages, ie q > 3OOmV, the
E. L.
684
LB-fAUER
AND
: hI
HO; +
L I+,
+
OH-- O2+ Hz0 Zql
H20+ Z.- - HO;+ OH-
Fig. 3, Schematic&d
potential diigrm
and Hoi-OH-.
for
decomposition of H,O,
reduction reaction approaches looo/, and the measured limiting current represents that reaction exclusively. Returning to the experimental curves of Fig. 1, it will he noted that the observed limiting current is virtually independent of the composition of the electrode surface and presumably therefore the surface area of each electrode was essentiaIlL identical. The diffusion
coefficient of 5 x 1W6 cm*/s (f 5%) was found to be
independent of peroxide concentration investigated.
bassd
on redo% COUPES
O,-HO;
E, represents the mixed poteatial at OCV.
current efficiency of the HO;
over the range
CONCLUSION Addition of H,02 to alkaline solution results in conversion to the HO; ion, and this then reacts so that oxygen evolution occurs virithin the bulk solution and especially at any catalytic surf& hnmers al in it. Because of this inherent instability in alkal& s&ion, conventional methods for experimentally determining Dxo; _are unsuitable. The observation that by ekctrochenucally reducing the ion at its limiting current will result in almost total inhibition of 0, evolution at the electrode, has provided a means to obtain the’dif&ion coefficient. Chronoamperometric polarization at a rde gave virtually identical i,‘s irrespective of the catalyst used, provided the surfficc smoothuess was the same for all electrc&s. This observation provides new information to the controversy surrounding the mcchanism of hydrogen peroxide reduction in aIk&ne solution. To this end a mixed potential model based on the two operating ralox reactions has been con-
structed to illustrate the situation. Inserting the cxpeinto this mode1 tends to support the electrochemical mechanism.
rimental observations
K. C. Tut
The investigation has provided a value for the di&sion coefficient of fairly concentrated solutions of perhydroxyl ion in LiOH. This coefficient has not, to the best of the authors’ knowledge, been reported elsewhere. A simii value could be anticipated with other strongly alkaline solutions. The observed inde pendence of the d#usion coefficient with perhydroxyl ion concentrations would be expected because excess supporting electrolyte was used. Further work is desirable to co&m the value bf I&,; at lower concentrations than those in this study. REFERENCES
1. E. L! Littauer. K. C..Tsai and R. P. Hotlandsworth, J. electrdtem. Sot. 11, (6) 845 (1978). 2. R. Gerischer and H. G&cl&, Z:phy~. Chem. 6, 178 11936). 3. E. L. Littaucr aud K. C. Tsai, EIectrockim. Acta 2d, 351 (1979). 4. D. P. Shoemaker and C. W. Garland, Exp&ne~~s in Physicul Chemistry, (2nd edn) pp. 28&281. McGrawHill, New York (1967). N. Sato and G. Okarnoto, Ekctrochim. Acta 10, 495 (1965). B. E. Conway, hi. A. Sattar and D. Gilroy, ibid 14,677 (1969). T. Katan, M. Chin arid F. J. Schoeaweis. Nrrture, L.ond. 4951. 1281 11964). N. ,k Shun&va.‘G. V. Zhutaeva, N. D. Mcrkilova, E. I. Krus&cva, G. B. Samoilov, and V. S. Baaotskii, Kinetika i Kataliz 14, {l) 168 (i973j. 9. D. Wiilmau, Z. Elek.trochem 60,731 (1956) 10. V. S. Bngotskii and I. E. Yablokova, Zh.fiz. Khim. 27, 1663 (1953). 11. F. Ha&r &I S. Grinbarg, Z. anorg. allgm. Ch. 18.37 (1938).