tions of complexing agents, the solution conditions are not constant, so that the chemical reaction, whose rate may depend on the structure of the electrode-solution interface, may cause different potential shifts a t different concentrations of complexing agents. Consequently, the results of such studies should be treated with some caution. It is possible to obtain unambiguous results by using concentrations low enough so that the chemical reaction does not exert a significant effect on the half-wave potential. This requires concentrations of lO-6M or less, so that differential pulse polarography must be used. It is important to demonstrate that the half-wave (or peak) potential is independent of the concentration in order to be certain that the chemical reactions do not cause changes in the potential of the wave. The second major factor in the determination of the half-wave potential, that of coordination with the buffer anions as well as hydroxide ions, is entirely predictable ( 1 4 ) . Assume that there are three ions in the solution which form stable complexes with the electroactive species. If the three ions are represented by X, Y , and Z and the respective formation constants by K,, K , , and K,, then the half-wave potential may be written as follows:
It is assumed that the diffusion coefficients for the oxidized, reduced, and complexed species are equal, and that the activity coefficients are constant. In addition, the concentrations of the complexing agents must be at least 100 times larger than the concentrations of methylmercury in order to minimize changes in concentration at the electrode surface during electrolysis. In a real experiment, X,
Y, and 2 might represent the two buffer anions and the hydroxyl ion. Use of a supporting electrolyte with coordinating tendencies toward methylmercury would add a fourth variable to the equation. I t is clear that the coordinating properties of all the ions in the solution must be considered when evaluating the significance of the halfwave potential. The fact that methylmercury forms complexes with most common buffer anions t o about the same degree as hydroxide (28-31) attests to the futility of attempting p H studies without accounting for this behavior. The differential pulse polarographic peak current u s . concentration curve for the first wave is linear from 10-7 to 10-4M methylmercury. The extreme length of this linear range suggests that this technique may have analytical utility. Two improvements over previous procedures may be realized by this technique. First, sensitivities of lO-7M or 20 pg/l. in the solution polarographed can be achieved. If a small degree of concentration can be effected during the required extraction procedures, it should be possible to accurately measure methylmercury concentrations less than 1 pg/l. in the original sample. Second, one may construct synthetic systems and study coordination chemistry at concentrations less than 10-6M. It is not possible t o do this type of study with most of the normally used techniques. The main drawback to the use of polarographic means for measuring methylmercury concentrations is that the peak potential is dependent upon the solution conditions. However, this problem can be eliminated if sufficient care is taken to devise extraction procedures which yield reproducible solution conditions in the final solution. Received for review June 6, 1973. Accepted November 29, 1973.
Mixed-Potential Mechanism for the Potentiometric Response of the Sodium Tungsten Bronze Electrode to Dissolved Oxygen and in Chelometric Titrations P. B. Hahn,' D. C. Johnson, M. A. Wechter,2 and A. F. Voigt D e p a r t m e n t of Chemistry and the A m e s Laboratory-USAEC,
l o w a State Unwersity. Ames, iowa 5007 0
Evidence is presented showing that an adsorption mechanism which was proposed in earlier work to explain the potentiometric response of the N a x W 0 3 electrode in alkaline solution is not correct. The potential response to dissolved oxygen in alkaline solution and the potential shift observed at the equivalence point in titrations of metal ions in ammoniacal solution with EDTA are explained by a mixed-potential mechanism. The potential is established as a result of the spontaneous oxidation of the N a x W 0 3 electrode by dissolved oxygen. These responses were found to be unique to the cubic N a x W 0 3 among all the highly conducting alkali metal tungsten bronzes.
Sodium tungsten bronzes, highly conducting nonstoichiometric compounds of formula Sa,W03 (0.5 < x < 0.9), have recently been demonstrated t o be useful as po-
tentiometric indicating electrodes in alkaline solution for the determination of dissolved oxygen and for indicating the equivalence point in chelometric titrations of many metals with EDTA (1-3). The response to dissolved oxygen in the 0.2-8 ppm concentration range in 0.1M KOH1mM EDTA was determined to be a linear function at' the log of 0 2 concentration with an unexpected large s l o p of approximately 120 mV/decade. For chelometric titrations in ammoniacal solution, negative potential shifts of 301 Present address, Radiation Management Carp.. Philadelphia, P a . 19104. Present address, Southeastern Massachusetts University. North Dartmouth. Mass. 02747.
(1) P. B. Hahn, M . A . Wechter, D . C. Johnson, and A F. Volgt, A n a / . Chem.. 4 5 , 1016 ( 1 9 7 3 ~ . ( 2 ) M . A. Wechter, P. 6. Hahn, G . M Ebert. P. R . Montoya, and A F Voigt, A n a / . Chem , 45, 1267 (1973). ( 3 ) P. 6. Hahn. Ph.D. Thesis, IowaState University, Ames. Iowa, 1 9 i 3 .
ANALYTICAL CHEMISTRY, VOL. 46,
NO. 4 , APRIL 1974
553
150 mV were observed a t the equivalence point when 0.01-2.5 millimoles of metal ion were titrated with EDTA in the presence of dissolved oxygen. Moreover, substantial potential shifts were obtained in the titration of such electroinactive species a s calcium(II), magnesium(II), and zinc(11) without the presence of a n electroactive metal cation. Additional details concerning the response of the tungsten bronze electrodes may be found in References 1-3.
These observations cannot be explained in terms of simple oxidation-reduction reactions. The Nernst equation predicts either a 15 or 30 mV/decade response for the reduction of oxygen to water or to hydrogen peroxide, much less than the 120 mV/decade response experimentally observed. The apparent response to electroinactive species in the chelometric titrations also cannot be explained on the basis of simple redox reactions. Very early in the course of these investigations, a correlation was made between the unique behavior of the NaxW03 electrode and the presence of hydroxyl ion. The electrode potential is a function of hydroxyl ions exhibiting a negative shift with increasing concentration. At low pH, there is negligible potentiometric response to oxygen and to reducible metals ( I , 2). A mechanism was proposed for the response of the NaxW03 electrode on the basis of adsorption of the negatively charged hydroxyl ion ( I , 2 ) . Numerous examples of adsorption mechanisms can be cited for the response of ion-selective potentiometric electrodes developed in recent years ( 4 , 5 ) . A similar mechanism would explain the response of the NaxW03 electrode if uncharged oxygen molecules are strongly adsorbed a t the electrode displacing hydroxyl ions and causing a positive potential shift. The success of EDTA titrations could be similarly explained only if the metal cations titrated are adsorbed displacing hydroxyl ions. Observations from subsequent experiments, including voltammetric studies of the NaxW03 and other alkalimetal tungsten bronze electrodes in alkaline solution, place serious doubt on the adsorption mechanism. Spontaneous oxidation of the NaxW03 electrode and reduction of dissolved oxygen was demonstrated a t potentials in the same potential region as the potentiometric response, These observations are consistant with those of Straumanis (6, 7) who demonstrated the oxidation of NaxW03 to W042-(tungstate ion) during prolonged contact with alkaline solutions of oxidizing agents such as oxygen, sodium peroxide, and silver(1). The potential of the YaxW03 electrode is now concluded to be a mixed potential resulting from the spontaneous reaction of the electrode with the oxidizing agents in the solution. Mixed potential phenomena have been discussed by various authors (8-12) and are considered to arise when a nonequilibrium state exists involving two or more electrode processes. Such a state is associated with a spontaneous change a t the electrode surface where simultaneous oxidation and reduction occur. The algebraic sum of the "Glass Electrodes for Hydrogen and Other Cations," G. Eisenman, Ed.. Marcel Dekker, New York, N Y , 1967. "Ion-Selective Electrodes." R . A . Durst. Ed., National Bureau of Standards Special Publication 314, U.S. Government Printing Office, Washington, D.C.. 1969. M. E. Straurnanis, J . Amer. Chem. Soc.. 71, 679 (1949). M E Straumanis and S. S. Hsu, J Arner. Chem. Soc., 72. 4027 (1950). C. Wagner and W. Traud. 2. Electrochem., 44, 391 (1938). J. Koryta. J. Dvorak, and V. Bochackova. "Electrochemistry," Methuen and Co. Ltd.. London, England, 1970. pp 318-21 D. Gray and A . Cahill, J . Electrochem SOC.,116, 443 (1969) J M . Herbeiin, T. N. Anderson, and H. Eyring, Electrochim. Acta. 15. 1455 (1970). J. O'M. Bockris and A K. N. Reddy, "Modern Electro-Chemistry," Voi. 2, Pienum Press, New York N . Y . . 1970.
ANALYTICAL CHEMISTRY, VOL. 46,
NO. 4 , APRIL 1974
partial currents at the electrode, icathodic + ianodic, must be equal to zero and, since the partial currents are related to electrode potential, the mixed potential satisfying the condition of zero net current is established. If the voltammetric polarization curves of the cathodic and anodic processes can be individually determined, the value of the mixed potential can be calculated graphically by the method of Wagner and Traud (8) or mathematically if representative equations can be written for the rates of the partial cathodic and anodic processes (10, 11).
EXPERIMENTAL Materials a n d Apparatus. The tungsten bronze crystals used for electrodes were cubic sodium tungsten bronze, XaxWOs ( x = 0.65 and 0.79); cubic lithium tungsten bronze, LixWOj ( x 0.35); tetragonal potassium tungsten bronze, K x W 0 3 (x N 0.5); and hexagonal potassium, rubidium, and cesium tungsten bronzes ( x = 0.3). Details of the bronze preparation (13) and electrode fabrication (14) are presented elsewhere. All solutions were prepared from reagent grade chemicals and deionized water. Gases used to establish a given oxygen concentration in solution were dry 99.995% nitrogen, 99.6'70 oxygen. M a theson "Zero" grade air. and specific oxygen-nitrogen mixtures (10.12. 3.27, 0.99, 0.35. and 0.10% 0 2 by volume) prepared and analyzed (f270relative) by Matheson Gas Products. Potential measurements were made with a Beckman Zeromatic SS-3 p H meter using a saturated calomel electrode as reference electrode. Voltammetric measurements were made with a Leeds and Northrup. Electro-Chemograph Model E polarograph using a large area saturated calomel reference electrode ( 1 5 ) . Procedure. The response to metal ions in the 10-6 t o 10-2M concentration range was performed in 1.OM ",OH by diluting 5-gl to 10-ml aliquots of the appropriate 0.1Mor 0.01M metal solution to 100-200 ml. The titration of l millimole of calcium(I1) with 0.1M disodium ethylenediamine tetraacetate (EDTA) and 0.3 millimole of EDTA with 0.1M CaC12 was also performed in 1M KHIOH. The response t o dissolved oxygen was determined in 0.1M K O H - l m M EDTA by equilibrating the solution with air or one of the five pre-mixed gases and measuring the potential of the bronze electrode u s . a SCE electrode after the potential reached a steady-state value (approximately 10 m i n ) . The temperature of the solution was maintained at 25 f 1 "C using a constant temperature bath. Solutions were stirred magnetically for all potentiometric measurements. Voltammetric studies were performed with rotating bronze electrodes at 300 r p m . Voltage scans were from positive to negative at 0.200 V/min. The concentrations of dissolved oxygen were established as described for potentiometric measurements using the appropriate gas mixture.
RESULTS AND DISCUSSION Figure 1 illustrates the potentiometric response of several alkali-metal tungsten bronze electrodes to dissolved oxygen in 0.1M KOH-lmM EDTA solution. The cubic Na0.65W03 electrode exhibited the greatest response (- 120 mV/decade) and a linear potential-logjC(02)1 relationship. The hexagonal Rbo.3WO3 and K0.3W03 electrodes showed virtually no response and the cubic Li0.35WO3 exhibited a marginal 30 mV/decade response to dissolved oxygen. Significant response (60-80 mV/ decade) was observed for the hexagonal Cs0.3WO5 and tetragonal Ko.5W03 electrodes but these electrodes responded much rlore slowly and much less reproducibly than the Nao.65W03 electrode. The potentiometric response of the ?Ja0,6~W03electrode to a number of metals in 1M NH40H is presented in Figure 2. Response was found to be limited to those species which are easily reduced in aqueous solution. Response to calcium(I1) and other electroinactive species was extremely small while the responses to silver(1) (80-150 (13) H. R . Shanks, J. Crystai Growth. 13-14. 433 (1972). ( 1 4 ) M . A. Wechter, H. R . Shanks, G . Carter, G . M . Ebert, R . Guglielmino, and A. F. Voigt, Anal. Chem., 44, 850 (1972). (15) A . I Vogel. "Quantitative Inorganic Analysis," John Wiley and Sons, New York N.Y., 1961
I - , , O r -* --. -----t
i
RbG3 w03
.i I
-300 __..__
3
-4ooL
/
/'
-800-
%02 IN EOUICIBRATING GAS
Figure 1. Oxygen response of hexagonal Rbo.aWOs, Ka.3W03, and Cs0.3W03 electrodes, and cubic Lio.35W03,tetragonal Ko.sW03 and c u b i c Nao.ssW03 electrodes in 0.1M KOH-lmM
EDTA
CONCENTRATION (M)
FORMAL
Figure 2. Potentiometric response of the Nao.65W03 electrode
to Several species in 1M N H 4 O H '
501
'
1-7 '
mV/decade), ferricyanide (180-300 mV/decade), and copper(I1) ( - 125 mV/decade) were very large, paralleling that for oxygen in alkaline solution. The responses of other alkali-metal tungsten bronzes to oxidizing agents in 1M NHlOH were less than for NaxW03. The responses to copper(I1) were 30-40 mV/decade. Serious doubt was placed on the adsorption mechanism proposed in earlier work following the observations that the sodium tungsten bronze electrode responded only to oxidizing agents. The absence of response to calcium(I1) and other electroinactive species was not consistent with the proposal that these species were causing the desorption of hydroxyl ion at the electrode surface and the observed potential shift at the equivalence point of chelometric titrations. Current-potential (I-E) curves obtained at a Nao 65W03 electrode in 0.1M KOH-1mM EDTA for several oxygen concentrations are presented in Figure 3. The curves were recorded scanning E from -0.4 to -1.5 V us. SCE. The I-E curve exhibits the expected cathodic wave for the reduction of HzO to Hz at approximately -1.5 V us SCE. Also observed is an anodic wave beginning at -1.0 to -1.2 V us SCE, becoming extremely large at potentials more positive than -0.5 to -0.6 V us. SCE. The anodic wave is attributed to the oxidation of the Nao 65W03 electrode to tungstate ion (W042-). When oxygen was present, a cathodic wave was observed with a half-wave potential in the same potential range (-0.6 to -1.0 V) where the Nao 65Wo3 oxidation takes place. Similar voltammetric curves were observed for a Nao 79WO3 electrode. The existence of the anodic (NaxW03 oxidation) wave in oxygen-free solutions at potentials more negative than the cathodic ( 0 2 reduction) wave provides strong evidence that spontaneous oxidation of the NaxWO3 electrode is occurring in the presence of oxygen in alkaline solution
'
1
POTENTIOMETRIC ,RESPONSE,
t-
02
1
L
- 50
-0 5
1
1
4
N2
I
I
I
1
,
P O T E N T I A L vs
SCE
,
, I 1 -I 5
-1 0
(V)
Figure 3. Voltammetric curves of Na0.65W03electrode in 0.1M KOH-lmM EDTA as a function of oxygen concentration. .k =
open circuit potential in a nitrogen-saturated solution and that the potential established at the electrode is a mixed potential. The effect of varying the concentration of dissolved oxygen on the mixed potential is schematically illustrated in Figure 4. Represented in this figure are the anodic wave for the oxidation of the NaxW03 electrode in 0.1M KOH1mM EDTA and a series of cathodic waves for the reduction of oxygen at various concentrations (C, < CZ < C3 < C4). The anodic wave is that obtained in a nitrogen-saturated solution while the cathodic oxygen waves were estimated by substracting the anodic wave from the net I-E curves obtained in the presence of oxygen (Figure 3). The condition of equal and opposite anodic and cathodic currents (zero net electrode current) to establish the mixed ANALYTICAL CHEMISTRY, VOL. 46, NO. 4 , APRIL 1974
555
-1-4 5
ra: 1-50;
-12
-11
-I0
-09
-08
-0.7
-06 - 0 5 - 5 4
POTENTIAL v s S C E !VI
Figure 5. Results of wave analyses for anodic and cathodic partial currents illustrated in Figure 4
I
x = cathodic wave for solution equilibrated with 3.3% 02-96.7% N2 mixture; 0 = anodic wave
/
II
I
OXIDATION OF
Na0.65w3 I
/
I
Figure 4. Schematic diagram illustrating the establishment of the mixed potential as a function of oxygen concentration Per cent
O2in gas mixture: C,,
1%; C2, 3 3%; C3, l o % , and Cd, 100%
potentials is represented by the vertical dashed lines for the various concentrations of dissolved oxygen. The shift of the mixed potential toward positive values with increasing dissolved oxygen concentration is clearly illustrated. Assuming the back reactions are negligible for the anodization of the tungsten bronze electrode and the reduction of dissolved oxygen, the net electrical current is given by Equation 1. inet
E
icathodic
+
ianodic
= n,FAk,Co2 exp{*E}
- naFAk, e x p PanaF { r E } In Equation 1, the subscripts c and a denote that the designated quantities apply to the cathodic and anodic half-reactions, respectively, and cyc and Pa are empirical coefficients. The surface concentration of dissolved oxygen is Co,. At inet = 0, E = E m and
Solving Equation 2 for E m ,
E,
2.3 R T PF
2.3RT lOg~C0,l
= ___
where
P
= Pana
+ acne
(3) (4)
Equation 3 is of the Nernstian form and a t 25 "C
E,,,
=
E,
+ 0.0591 -loglCo,J P
If the cathodic partial current is negligible in comparison to the limiting current for the reduction of dissolved oxygen, Co, equals the bulk concentration of oxygen, C O , ~A . 556
ANALYTICAL CHEMISTRY, VOL. 46, NO. 4, APRIL 1974
plot of E , us. log(C0,b) is expected to be linear with a slope of O.O591/p mV/decade. Analyses of the anodic wave for the electrode in oxygenfree solution and the cathodic wave for dissolved oxygen, corrected for charging current, are shown in Figure 5. A plot of log ia us. E for the anodic reaction is linear with a slope @,&/0.0591 a t 25 "C. A plot of log ic/(i,,c - zc) us. E for the cathodic reaction, where il,c is the limiting cathodic current, is linear with slope -a,n,/0.0591 a t 25 "C. From the measured slopes. 4.13 and -4.88, pan, is 0.244 and acne is 0.289. The value of p in Equation 3 is 0.533 and the coefficient to logiCo,/ in Equation 4 is 111 mV/ decade. This is in excellent agreement with experimentally observed responses. The large potentiometric response observed in the determination of oxygen, copper(II), and ferricyanide is consistent with other results reported in the chemical literature. Herbelin et al. (11) encountered similarly large potentiometric responses (125-250 mV/decade) for cerium(1V) and iron(II1) when studying a mixed potential system involving the cerium(IV)/cerium(III) and iron(III)/ iron(I1) redox couples. Responses larger than observed in this study could be obtained under unique conditions. If the cathodic current for the reduction of dissolved oxygen were diffusion limited, it would be directly proportional to the concentration of oxygen and independent of the electrode potential. A tenfold increase in the oxidant concentration would result in a tenfold increase in the cathodic partial current which in turn would require a tenfold increase in the anodic partial current to satisfy the condition of zero net current. It is evident from the wave analysis for the anodic wave (Figure 5) that this increase in anodic current would require a 250-mV shift in the mixed potential of the electrode. The mixed-potential mechanism also provides excellent explanations for several secondary observations associated with the oxygen response of the T\;axW03 electrode (3). Any factor influencing either the rate of NaxW03 oxidation or the rate of oxygen reduction will subsequently alter the mixed-potential response of the electrode. An enhanced response in the presence of trace silver(1) is explained in terms of trace silver depositing on the electrode and catalyzing the oxygen reduction. The poor oxygen response of NaxW03 electrodes with low x value in LiOH is thought to be due to an adsorbed LixW03 layer inhibiting the bronze oxidation reaction. The day-to-day variations in electrode response to dissolved oxygen and the dependence of the potential on stirring rate can also be understood in terms of the mixed-potential mechanism. Changes in electrode surface characteristics may result
1 350 r T T - I - 1
I
I
I
I
I
I
'
t
600
LlLLLLJ
L
- 500
-0 5
-15
-I 0
P O T E N T I A L vs SCE ( V I
Figure 8. Voltammetric curves for the Na0.65W03 electrode in 1M NHIOH showing effect of added EDTA
650 -
i 4
5
6 rpl
E!
7
9
O2
12 0
,011
C I M ED-A
1
ml o
1
2 IM Go3'-
Figure 6. Titration of 1.0 millimole Caz+ with 0.1M EDTA and 0.3 millimole EDTA with 0.1M Ca2f with zero (x) and 5 X l O W 5 M( 0 )added C u ( l l ) 32
+-
u2
POTENTIOMETRIC RESPONSE
LL--15
IO
05
-i
05
00
POTENTIAL v s SCE
1J
I
IO
3
(VI
Figure 9. Voltammetric curves of cubic LIO35wo3 electrode in oxygen and nitrogen-saturated 0 1M KOH-lmM EDTA solutions
*=
open circuit potential in nitrogen-saturated solution
5 0 - , , 1
.50;ildL_LL' , , , -0 5
00
I
I
,
I
1
"
, ,
i
,
r-
-r-l-
1
I
-I 0
POTENTIAL vs SCE ( V )
Figure 7. Voltammetric curves of Na0,65W03electrodes in 1M NH40H as afunction of Cu(1l) concentration
O,* = open circuit potentials
from the slow dissolution of NaxWOs and possibly cause alterations of the anodic and cathodic reaction rates and the oxygen response. The shift in potential toward more negative values a t slow stirring speeds is a result of a decrease in the rate of mass transfer of the electroactive species and the consequential decrease of the cathodic partial currents. The substantial positive shift in the potential of the NaxW03 electrode a t the equivalence point in EDTA titrations of electroinactive species can also be explained by the mixed-potential mechanism. Evidence was found that trace copper(I1) or some other oxidizing agent with a large formation constant for its EDTA complex would indeed act as a potentiometric indicator in these titrations. The addition of 1.3 X lO-5M copper(I1) resulted in significantly larger potential breaks in both the titration of calci-
PGTEY'ICL
YS
SCE
J
Figure 10. Voltammetric curves of hexagonal Cso 3 W 0 3 electrode in oxygen and nitrogen-saturated 0 1M KOH-lmM EDTA sol utlons
*=
open circuit potential in nitrogen-saturated solution
um(I1) with EDTA and the titration of EDTA with CaClz (Figure 6). Voltammetric curves of the Ka0.65W03 electrode in 1M NHIOH a t several copper(I1) concentrations are presented in Figure 7. A positive shift in the mixed = 0) with increasing copper(I1) concentrapotential (& tion is demonstrated. In the complexometric titrations of calcium(II), copper(I1) and other oxidizing agents were absent. The shift in the mixed potential a t the equivaANALYTICAL CHEMISTRY, VOL. 46, NO. 4, APRIL 1974
557
r-r-
%.-,Nz ~
50 I5
"
10
"
0 5
POTENTIMvlETRlC RESPONSE
00
PGTENTAL vs SCE
05
i -10
-1
5
1
(VI
Figure 11. Voltammetric curves of hexagonal
Rbo.sWO3 electrode in oxygen and nitrogen-saturated 0 . 1 M K O H - l m M E D T A -05-
SOIUtlOn
* = open circuit potential
00
in
nitrogen-saturated solution
lence point is actually due to the presence of EDTA. Figure 8 presents voltammetric curves for the Na0.6~W03 electrode in nitrogen and air-saturated 1M NH40H showing the effect of EDTA a t a concentration of 1.3 X lO-5M. The presence of trace EDTA shifts the anodic polarization wave for the oxidation of the Nao 65W03electrode to more negative potentials by 120-140 mV. This shift in the anodic wave results in a similar shift in the mixed potential established from the spontaneous oxygen reductionNaxWOs oxidation reactions and explains the potential break observed at the equivalence point in the chelometric titrations. Both the large oxygen response and the large potential break in EDTA titrations are unique to the cubic sodium tungsten bronze electrodes. Voltammetric curves for several other highly conducting alkali metal tungsten bronze electrodes in 0.1M KOH-lmM EDTA are presented in Figures 9-12. Immediately apparent is the absence of any bronze oxidation wave at potentials between -0.5 and -1.0 V us. SCE in oxygen-free solutions. The stability of these bronzes toward oxidation by dissolved oxygen is consistent with their poor performance in responding to dissolved oxygen and in detecting the end point in chelometric titrations. The observations that cubic Lio 35wo3 and hexagonal CSO3W03 exhibit some response and hexagonal KO3W03
558
A N A L Y T I C A L C H E M I S T R Y , V O L . 46, N O . 4, APRIL 1974
-I 0
-0 5 POTENTIAL
VS
SCE (V)
Figure 12. Voltammetric curves of hexagonal K0.3W03 electrode in oxygen and nitrogen-saturated 0 1 M K O H - l m M E D T A solutions
* = open circuit potential in nitrogen-saturated solutions
and Rbo,3WO3 exhibit essentially no response to dissolved oxygen can be rationalized from their voltammetric waves and the mixed-potential mechanism. For both the Li0.35WO3 and Cs0.3W03 electrodes, the cathodic oxygen wave begins at more positive potential than the open circuit potential in oxygen-free solution as required for response. However, the anodic wave is not in the region of the mixed potential. For the hexagonal K0.3WO3 and Rbo.aWO3 electrodes, the oxygen waves begin a t more negative values of potential than the open circuit potential in oxygen-free solution. In these instances, dissolved oxygen is not a sufficiently strong oxidizing agent to participate in establishing a mixed potential and no response is noted.
ACKNOWLEDGMENT The authors thank Howard Shanks of the Ames Laboratory for providing the tungsten bronze crystals used for construction of electrodes in this study. Received for review August 20, 1973. Accepted Kovember 14, 1973.
