read directly. More significant is the fact that information theory predicts that cross-correlation enhances the signalto-noise ratio (11, 22). This allows greater resolution than could be obtained uia Lissajous figures. The instrument described above has been demonstrated to be well suited to its primary task of measuring high impedance systems a t low frequencies using minimal excitation. However, the potential of the correlation method for use in systems where conductivity is a dynamic parameter (11) G . M. Heiftje, Anal. (12) G. M . Heiftje, Anal.
Chem., 44(6), 81A (1972). Chem., 44(7), 69A (1972).
should not be underestimated, particularly since the necessary equipment is available commercially or may be constructed in the laboratory a t low cost. The instrument may be easily modified to function in either the impedance or the admittance mode. In addition, the nature of the instrumental response makes it applicable to a wide variety of problems. Received for review August 24, 1973. Accepted November 28, 1973. This work was supported by National Science Foundation Grant GP-20524.
Electroanalytical Studies of Methylmercury in Aqueous Solution R. C. Heaton’ and H. A. Laitinen S c h o o l of Chemical S c i e n c e s , University of lllinois at Urbana-Champaign, Urbana, ///. 6780 7
The electrochemical reduction of monomethylmercury compounds in aqueous solution has been studied through the use of pulse polarography, cyclic voltammetry, and related techniques. Reduction of these compounds at a mercury electrode occurs in two one-electron steps, the first of which results in the formation of a methylmercuric radical on the electrode surface. This step is reversible under polarographic conditions, but the polarographic wave is distorted because of the involvement of the methylmercuric radical in subsequent chemical reactions. Addition of the second electron results in reduction of the methylmercuric compound to elemental mercury and methane, and gives rise to the second polarographic wave. This reduction is irreversible and the wave is also distorted. The first pulse polarographic wave has analytical utility arising from a linear peak current vs. concentration curve between concentrations of lo-’ and 10-4M. The analytical implications of the reduction mechanism are discussed, with particular attention given to effects of coordinating agents and to detection limits.
Electrochemical reduction of methylmercury was first observed by Kraus (1) in 1913. While other investigators (2-4) repeated Kraus’s experiments and verified his results, little new insight into the reduction mechanism was gained until 1948, when Costa (5) published his studies of the reduction of ethylmercury by electrochemical means. Since that time, mechanistic studies of a variety of organomercury compounds have been reported (6-15), Present address, Hercules Research Center, Wilmington, Del. ( 1 ) C. A . Kraus, J. Amer. Chem. S o c . 35, 1732 (1913). (2) B. G Gowenlock and J . Trotrnan,J. Chem. Soc.. 1957, 21 14. ( 3 ) B. G . Gowenlock, P. P. Jones. and D . W. Ovenall. J. Chem. SOC.. 1958,535. (4) J . L . Maynard and J . C. Howard Jr., J . Chem. Soc.. 1923, 960. (5) G . Costa. Ann. Chim. ( R o m e ) . 38, 655 (1948). (6) V. Vojir. Collect. Czech. Chem. Commun.. 16, 489 (1951) ( 7 ) R. Benesch and R. E. Benesch, J . Amer. Chem. SOC.. 73, 3391 (1951). ( 8 ) R. Benesch and R. E. Benesch. J. Phys. Chem.. 56, 648 ( 1 9 5 2 ) . (9) R . F . Brornan and R. W . Murray, Anal. Chem.. 37, 1408 (1965).
the works of Benesch and Benesch being most often quoted. I t is generally agreed that the polarographic reduction proceeds in two one-electron steps, the first yielding an organomercury radical, the second giving elemental mercury and, in the case of alkylmercury compounds, the alkane. Beyond this point, authors disagree as to the nature of the reduction process. Some investigators (5, 7, 8, 11) regard the first wave as irreversible, while others (6, 1215) consider it to be reversible. Vojir (6) suggested the irreversible dimerization of the organomercury radical might take place following the electroreduction, giving rise to an insoluble film on the electrode surface. Although most workers since that time have included some sort of dimerization or disproportionation reaction in their formulations, little has been done to detail the fate of the organomercury radical, or to determine the nature of any intermediates which may be involved. Several authors have noted that the polarographic waves exhibit irregularities (6, 7, 8, 24) and that prewaves often occur (7-9, 2214), but interpretations of these phenomena vary. The finer details of the reduction mechanism may well depend on the particular organomercury compound in question, the solvent, and the experimental conditions employed. The purpose of this investigation is to determine the details of the electrochemical reduction of methylmercury in aqueous solution, and to assess the feasibility of using electrochemical techniques to determine methylmercury and to study its coordination chemistry.
EXPERIMENTAL Instrumentation. Polarograms were recorded using a Heath polarography module, model EUA-19-2, connected to our own operational amplifier system. Recordings were made using either a Moseley model iOOlAM x-y recorder or a Heath model EUW-20A variable speed servo recorder. Potential scales were calibrated with a Leeds and Northrup volt potentiometer. catalog number 8687. Drop times used in constructing electrocapillarj curve5 were (10) K . P. B u t i n , I . P. Beletskaya, A . N. Ryabtsev, and 0. A . Reutov, Elektrokhimiya. 3, 1318 (1967) (11) V . G . Toropova, M . K . Saikina, and M. G Khakimov, Zh. Obshch Khim.. 37, 47 (1967) (12) C. Degrand and E . Laviron, Bull. Soc. Chim. F r . . 5, 2233 ( 1 9 6 8 ) . (13) C. Degrand and E. Laviron, Bull. SOC Chim. Fr.. 5, 2228 (1968) (14) B. Fleet and R. D . Jee, J. Electroanal. Chem.. 25. 397 ( 1 9 7 0 ) . (15) N. S. Hush and K . B. Oldharn, J. Electroanal Chem , 6, 35 (1963) A N A L Y T I C A L C H E M I S T R Y , V O L . 46, NO. 4 , A P R I L 1974
547
metric flasks and then filling t h e flasks t o t h e marks with t h e appropriate supporting electrolytes. More dilute solutions were prepared by dilution of t h e 1 m M solution.
O.O!
RESULTS AND DISCUSSION
OIC
-
3us
h l E
Y
w
0. x
O.2t
- 1.00
0.03 . .
1.00
log= I
Reduction of the methylmercuric cation in 0.1M perchloric acid solution or of methylmercuric hydroxide in 0.1M sodium hydroxide solution yields two well defined polarographic waves of nearly equal height. Both waves are diffusion controlled as the dependence of the limiting current on the concentration of the electroactive species is linear. In addition, the variation of the limiting current with the square root of the corrected mercury column height is also linear. The polarographic diffusion coefficients, calculated from either standard polarographic data or pulse polarographic data are 1.08 f 0.02 x 10-5 cmz/ sec in 0.1M HClOl solution, and 1.11 f 0.08 x 10-5 cm2/sec in 0.1M NaOH solution. The first wave is asymmetric, being more steep a t the foot of the wave than a t the head of the wave. This asymmetry is reflected in the standard polarographic log plot [E us. log (id - i ) / i ] which exhibits a change of slope near the half-wave potential. This distortion can be adequately explained by postulating the occurrence of a chemical reaction following the electrochemical reduction. The following reactions are proposed: CH,Hg+ -t e-
CH,Hg.
(la )
Figure 1. Dependence of E on log ( i d - i ) / i for the first wave in 0.1M H C I O 4 solution under normal pulse polarographic condi-
tions a ) 9.76 X .10-5M CH3HgL; b ) 4.88 X 1 0 - 5 M CH3HgC; C) 2.44 X 1 0 - 6 M CH3Hg+
measured by allowing t h e falling mercury drops t o interrupt a light beam and electronically timing t h e interval between two such events. Cyclic voltammograms were obtained by t h e use of a Wenking potentiostat built by G. Bank Electronik, No. 7163-61RS. T h e current was determined by measuring t h e voltage drop across a resistor in series with t h e counter electrode. Current-voltage curves were recorded using a Tektronix model 503 x-y oscilloscope equipped with a Tektronix model C-12 camera. Triangular voltage functions were generated by a function generator built in this laboratory. A Princeton Applied Research model 174 Polarographic Analyzer was used to obtain normal pulse and differential pulse polarograms. The Heath servo recorder was used t o record t h e results. All electrochemical experiments were performed in completely enclosed, all-glass cells. Dropping mercury electrodes were of standard construction, using soft glass marine barometer tubing for the capillary. For a hanging mercury drop electrode, a micrometer electrode, obtained from Princeton Applied Research, catalog No. 9302, was used. Potentials were measured relative t o a saturated calomel electrode. R e a g e n t s a n d Solutions. With t h e exception of the methylmercury salts, all reagents were reagent grade, used without further purification. Methylmercuric hydroxide obtained from Alfa Inorganics was recrystallized from pyridine, yielding white crystals. melting range 92-102 "C. I n view of t h e fact t h a t this melting range is significantly lower t h a n t h a t of the oxide (110-137 "C) and considerably higher than t h a t of the hydroxide (62-68 "C), it seems likely t h a t the sample is a mixture of t h e hydroxide and the oxide. This is of little consequence since the oxide is immediately converted t o the hydroxide upon dissolution in water. Titration of t h e sample with perchloric acid in t h e presence of potassium iodide indicates t h a t it is more t h a n 95% pure as monomethylmercury. The equivalent weight determined by t h e titration was used t o calculate solution concentrations. Mercury was scrubbed successively with nitric acid and potassium hydroxide solutions, and then vacuum distilled. Glassware was cleaned with aqua regia and thoroughly washed with double distilled water immediately before use. Millimolar solutions of methylmercuric hydroxide were prepared by introducing weighed amounts of the sample into volu548
A N A L Y T I C A L C H E M I S T R Y , V O L . 46, NO. 4, A P R I L 1974
Because reaction 2 effectively lowers the activity of the methylmercury radical, CH3Hg., the electrochemical process, 1, should be facilitated, increasing the slope of the polarographic wave and shifting the half-wave potential anodically. Also, because reaction 2 is second order in CH3Hg., it should exert less influence on the electrochemical process as the concentration is lowered. Careful perusal of Figure 1 shows that all of these trends do in fact occur. Bonnaterre and Cauquis (16) have developed a unified mathematical treatment of the polarographic wave observed in the general case of electroreduction followed by a chemical dimerization. While the equations were derived with a rotating disk electrode in mind, they can be made to apply to a dropping mercury electrode by writing the diffusion layer thickness in terms of the capillary characteristics and the drop time. Two cases have relevance to this investigation-that of a reversible electrochemical reduction followed by an irreversible dimerization, and that of a reversible electrochemiCa1 reduction followed by a fast reversible dimerization. For the former case the equations obtained are identical to those derived by Hanus (27, Hanus and Koutecky (18), and Mairanovskii (19, 20).
E l / ? = Eo'
RT kdCt + __ 3nF ln- 4
(16) R . Bonnaterre and G . Cauquis. J . E/ectroana/. Chem.. 32, 199 (1971). (17) V . Hanus. Chem Zvesti. 8, 702 (1954). (18) J Koutecky and V Hanus, Coiiect. Czech. Chem. Commun.. 20, 124 (1955). (19) S. G . Mairanovskii. Dokl. Akad Nauk SSSR. 110, 593 (1956). (20) S. G . Mairanovskii. i z v Akad. Nauk SSSR. Otd. Khim. Nauk. 12, 2140 (1961).
0.10
W
0 in
9
-> 3.15 Y
0.20
-8
Figure 2.
Dependence of E on log
(id
- ii'/i for the first wave in 0.1M HC104 solution under normal pulse polarographic conditions
a ) 9 76 X 10-5MCH3Hg+; b ) 4.88 X 1 0 - 5 M C H 3 H g f ;
C)
9.76 X 1 0 - 6 M C H 3 H g f ; d ) 2.44 X 10-6M CH3Hgf
In these equations, q is the average area of the drop, t is the drop time, k d is the chemical dimerization rate constant, and C is the bulk concentration of the electroactive species. The remaining variables have their usual significance. It is clear that a plot of E us. log i2 3/(id - i) should yield a straight line whose slope is -59 mV for each decade increase in depolarizer concentration. It is, in principle, possible to calculate the dimerization rate constant if the formal potential is known. If the chemical dimerization is fast and reversible, and if the equilibrium position strongly favors the dimer, then the following equations apply: =
E"' - -RT In 112 nF
FDli2t-''?K +
K is the dissociation constant for the dimer. I t is assumed that n is equal to unity. A plot of E us. log [(id - i)2/i] should give a straight line with a slope of 30 mV. In addition, the half-wave potential shifts 30 mV in the positive direction for each decade increase in depolarizer concentration. The dissociation constant for the dimer may be calculated if EO' is known. Equations C and D represent a situation in which the diffusion of the reduced species away from the electrode exerts an insignificant effect on the observed polarographic current. Consequently, these equations apply equally well to systems involving either homogeneous or heterogeneous chemical reactions. Figure 2 shows the pulse polarographic data for methylmercuric cations in 0.1M perchloric acid solution plotted according to Equation C. Similar results were observed for the standard polarographic data. At higher concentrations, the curves are linear between 10 and 90% of the wave height, suggesting that Equation C adequately describes the situation. As the concentration is lowered, the wave shifts cathodically and the curves begin to deviate from linearity. This behavior complements that shown by the curves in Figure 1. The standard log plots are nonlinear a t high concentrations but become linear for concen-
trations of methylmercury less than 2.5 x 10-6M in 0.1M perchloric acid solution. This linear behavior at low concentrations is expected when the electrochemical reduction is followed by chemical reactions of orders greater than unity, because the equilibrium position of the chemical reaction favors the monomer more as the concentration is lowered. At very low concentrations, the chemical reaction may be ignored altogether. Data for polarographic reductions of organomercury compounds have been plotted according to Equation A by previous authors ( I O , 11) and the resulting plots have been presented as evidence supporting the hypothesis that reaction 2 is irreversible. Plots of data according to Equation A made during this investigation were always nonlinear when plotted between 10 and 90% of the polarographic wave height. Consequently, it is suggested that the chemical dimerization is better described as a fast and reversible process than as an irreversible process. The half-wave potential for the first wave shifts in the negative direction as the concentration decreases. However, a t very low concentrations (4 x 10-'jM in perchloric acid solution), the half-wave potential becomes independent of concentration. Because this potential a t very low concentrations is not influenced by the following chemical reaction, it should be closely related to the formal potential for the electrochemical process. If it is assumed that the ratio of the diffusion coefficients for the oxidized and reduced species is unity, then the formal potential for this reduction in 0.1M Ferchloric acid solution is -0.143 volts us. SCE. Knowledge of this value permits the calculation of the dissociation constant for the dimer by the use of Equation D. The value so obtained is 10-5.7. This must be regarded as an approximation because the kinetic scheme described thus far is not complete. The second polarographic wave is distorted, having a long drawn out appearance. In addition, significant changes in shape are observed when the supporting electrolyte is changed from perchloric acid to sodium hydroxide. Although the drawn out appearance of this wave makes it appear to be irreversible under polarographic conditions, the electron transfer appears to be fairly fast, as it gives a well defined cyclic voltammetric peak even a t sweep rates of 100 volts per second. The irreversible apANALYTICAL CHEMISTRY, VOL. 46, NO. 4 , A P R I L 1974
549
J 0.3
- 01 I
’-
- 20
I
00
-08
E (L! vs.
J
I
-12
SCE)
- 0.5 1
I
b
-
b
I-
+0.3
I
1
-04
-”
0.0
-0.8
-0.4
-1.2
E (L! vs S L E )
- 0.1
-0.5
1
I
1
I
I
E (V: vs. S.C.E.)
Figure 4. Cyclic voltammograms of 4.49 X 1 0 - 4 M CH3Hg+ in Triton X-100 0 . l M HC104 w i t h 1.2 X E (V. vs. S.C.E.)
Figure 3. Cyclic voltammograms of 4.49 X 10-4M CH3Hg+ in O.1M HCIOI w i t h 1 . 2 X Triton X-100 a) Sweep rate = 0.50 V / s e c . ; b) Sweep rate = 5.0 V/sec.; c ) Sweep rate = 50 V/sec
pearance of the wave under polarographic ‘conditions can be attributed to two factors. First, the reaction is chemically irreversible, meaning that the forward reaction is very fast, but that the reverse rate is negligible. This is due to the fact that the reduction product, methane, is inert under the conditions employed. The second factor affecting the wave shape is the fact that the electrochemical reduction is competing with the chemical dimerization of the radical, giving the wave a drawn out appearance. The half-wave potential was observed to be independent of the supporting electrolyte and the types of buffers used, so that one may conclude that coordination is not a factor in the second reduction. However, the dependence of the half-wave potential on pH is interesting. At high p H values, the wave is independent of pH, hut a t pH values less than 5, the half-wave potential shifts 13 mV in the positive direction for each unit decrease in pH. This behavior may be interpreted in the following way. At high pH, where hydrogen ions are scarce, the organomercury radicals may be reduced to elemental mercury and a carbanion. The carbanion may then abstract a proton from a water molecule to give methane and a hydroxyl ion. The wave is thus unaffected by the hydrogen ion activity and, because the chemical reaction is irreversible, it is not affected by the hydroxyl ion activity either. At low pH, the reduction of the radical may be assisted by attack of a proton at the back side of the methyl group. This should aid the reduction, shifting the half-wave potential to more positive values and should also lead to inversion of the methyl group, which is without observable consequences. As this explanation essentially involves a change in the reduction mechanism with pH, differences in the wave 550
ANALYTICAL CHEMISTRY, VOL. 46, NO. 4 , A P R I L 1974
a ) Sweep rate = 0.50 V j s e c . ; b ) Sweep rate = 5.0 V!sec.; rate = 50 V / s e c
c ) Sweep
shapes obtained in acidic and basic solutions are expected. The reactions for the second wave may be written as follows:
CHJHg.
+ Hf + e-
CHIHg*-t eCH,-
+ H,O
-
-
e CH,-
CH,
+ Hg
+ Hg
CH, f OH-
i
(acidic solution)
(3a) (basic solution)
(3b) (3c)
Cyclic voltammograms of the methylmercuric cation in
0.1M perchloric acid solution are shown in Figure 3. Two characteristics of these curves are of major importance in the postulation of a reduction mechanism. The first is the shape of the anodic peak and the second is the ratio of the anodic peak current and the cathodic peak current. It is clear that the amplitude of the anodic peak is dependent on the sweep rate. It is virtually nonexistent at slow sweep rates, but increases as the sweep rate increases. It is also appafent that the anodic peak does not show behavior typical of a diffusion controlled wave, although the cathodic peak does. The anodic peak shape is typical of a stripping peak-a peak resulting from the destruction of an adsorbed or insoluble film. The implications of this will be discussed shortly. Similar behavior is exhibited by the second cathodic wave (Figure 4 ) . While this wave never entirely disappears, its amplitude is small at slow sweep rates, and it increases as the sweep rate increases. In addition, the peak shape is again typical of a stripping process. The situation is even better described in Figure 5 , which shows the peak currents as functions of the sweep rate. The anodic peak current is less than the cathodic peak current at slow sweep rates, hut greater than the ca-
' P
C
L
150
100
a
A
-
L
50
Figure 5.
Cyclic voltammetric peak currents as functions of sweep rate in 0.1M HCI04 solutions
ai 8.98 X 1 0 - 4 M CH3Hg' i- 0.1M HCIOl -k 1 . 2 X l o - ' % Triton X-100. 1 . first cathodic peak. 2. anodic peak; 3. second cathodic peak. b ) 1.08 X 1 0 - 4 M CH3Hg0.1M HClOn i- 1 . 2 X Triton X-100. 1. first cathodic peak: 2. anodic peak; 3. second cathodic peak
+
thodic peak currents at fast sweep rates. Note that there is nothing wrong with having the anodic peak current larger than the cathodic peak current, since the anodic peak is taller, but narrower, than the diffusion-controlled cathodic peak. As the concentration is reduced, the anodic peak current increases relative to the cathodic peak current. This behavior is typical of a second-order process following the electrochemical reduction. Because the second cathodic process corresponds to the reduction of the methylmercury radical, CH3Hg., while the anodic wave results from the oxidation of the same radical, the peaks corresponding to these processes should behave very much alike. It can be seen that they do. Small quantitative differences are observed because the waves occur a t different potentials. The fact that the anodic peak is small or nonexistent at slow sweep rates but large and well defined a t fast sweep rates, strongly suggests the occurrence of a slow chemical reaction following the electrochemical reduction. The evidence also suggests that the reaction is second order. This results in a dilemma, since the polarographic experiments seemed to indicate that the electrochemical reduction was followed by a fast chemical reaction. This dilemma can be resolved by rewriting reaction 2 in the following way: fast
2CHJHg.
e ( CHJHg
slow
(CH,Hg),
eCH3HgCH, + Hg
(2b)
Reactions 2a and 2b may be used to explain the cyclic voltammetry data in the following way. At very fast sweep rates, reaction 2b may not take place to a significant degree during the course of the experiment, so that the cyclic voltammogram looks like that expected for a mechanism including only steps 1 and 2a. However. a t slow sweep rates reaction 2b may approach equilibrium. If the equilibrium of 2b favors the product, then the transformation of dimethylmercury to dimethyldimercury should be much slower than even the forward reaction. Thus, reaction 2b should appear to be nearly irreversible. Consequently the dimer may be effectively transformed to an
inactive species a t slow sweep rates, so that the anodic peak may be very small, or nearly non-existent. Under polarographic conditions, in which the radical, CHaHg-, is continually being generated by electrolysis, the system behaves much as if reactions 1 and 2a were the only things happening. The effect of reaction 2b is to cause some leakage of the dimethyldimercury from the system, resulting in a larger apparent formation constant for the dimer. The behaviors of the anodic and second cathodic peaks observed during the cyclic voltammetry experiments suggest that the reactions 2a and 2b may be heterogeneous. Further evidence to support this conclusion is the fact that the electrocapillary curve is slightly depressed in the potential interval between the two polarographic waves (Figure 6). That suppression of the interfacial tension only occurs negative of the electrocapillary maximum suggests that the adsorbed species is oriented in such a way as to resemble a dipole with its positive end toward the electrode. The exchange of labeled elemental mercury with organic mercury compounds has been studied in some detail (21-27) and is directly relevant to this investigation. Of particular interest is the study of Butin et a1 (27). It was proposed that the exchange occurs through the following steps. First, the diorganomercurial is adsorbed onto the mercury surface. This adsorption is a function of the structure and concentration of the organomercurial and the nature of the solvent. The second step involves an electron transfer from the diorganomercurial to the mercury metal. This may occur by the formation of a fourcentered cyclic intermediate consisting of two mercury atoms in the electrode surface bridged by the two organic groups, and is believed by Butin and coworkers to be the rate-determining step. The third step involves separation of the intermediate into two adsorbed organomercury radicals. Retention of configuration is retained as the radicals are strongly inserted into the mercury surface. The series of steps just described is clearly the reverse of reactions 2a and 2b. The evidence obtained in this study requires that the transition between the adsorbed A N A L Y T I C A L CHEMISTRY, VOL. 46, NO. 4 , A P R I L 1974
551
Fleet and ,Jee ( 1 4 ) have explained the anodic cyclic voltammetric peak behavior for a number of organic mercury compounds in terms of the stability of the adsorbed radical, They go on to assume that the unadsorbed radical undergoes rapid dimerization and that the concentration dependence of the anodic to cathodic peak current ratio is due to a variation in the proportion of the radical adsorbed. While the first part of this explanation is essentially in agreement with that offered here (except that the present work details the subsequent fate of the radical), the second assumption is a t variance with the observation of a number of investieators (21-27) that the exchange reactions between elemental mercury and organic mercury are heterogeneous, and involve adsorbed dimeric intermediates. In addition. it has been shown above that the concentration dependence of the anodic to cathodic peak current ratio can be adequately explained by the occurrence of reactions 2a and 2b. In contrast to previously published work on a number of other organic mercury compounds, no prewaves were observed in this study of methylmercury. nor were irregularities observed in the curfent-time behavior a t the foot of the first polarographic wave. Polarographic maxima of the first kind often occur a t the head of the first wave, while the maximum a t the head of the second wave satisfies Frumkin's criteria for maxima of the third kind. All of the irregularities observed can be suppressed by the addition of suitable surfactants to the system. and thus constitute no particular obstacle to the analytical chemist.
156
5.2
5 01
E
F
a
e a
4.7,
t
ANALYTICAL IMPLJCATTONS OF THE REDUCTION MECHANISM
4.3
Figure 6.
Electrocapillary curve for
CH3Hg-
in
0 1M H C I O 4
solution a i 0 1 M HCIO, bi 9 59 X 10 ' M CH,Hggram of solution b
+ 0 1 M HCIOl
c) Polaro-
radicals and the intermediate dimer be fast and reversible, and that formation of the dimer be strongly preferred over the radical. The assertion by Butin and coworkers that the rate determining step is the electron transfer step is in perfect agreement with the results of this study. Some care is indicated here, as the electron transfer step ill this case involves the chemical oxidation-reduction between elemental mercury and the diorganomercury compound to give the diorganodimercun compound. and is distinct from the electrochemical reduction of the organomercury cation. To briefly summarize. the proposed reduction mechanism embodies the following reactions:
CH,-f H.0 552
+ e-
-
CH,Hg. (Surf.)
CH,
CH -
+ OH-
+
Hg
1
(3b) Basic Solution (3c)
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 . A P R I L 1974
A number of reports have been made of the use of polarographic half-wave potentials to evaluate complex equilibrium constants for organomercury compounds ( 5 , 7, 10, 1 1 . 14. 15, 28). Of' these studies, the ones by Costa (j), Benesch and Benesch ( 7 ) , and Toropova e t a / . ( 1 1 ) deal primarily with the effects of p H on the half-wave potential, In several cases. difficulties in achieving consistent results were encountered. Although the fact that the methylmercuric cation forms stable complexes with most commonly used buffer anions and with most halide ions ( 2 - 3 1 ) would explain most of the difficulties by itself, other factors may also influence the potential of the polarographic wave. The chemical reactions 2a and 2b. cause a positive shift in the half-wave potential which depends on the concentration. While the magnitude of the shift is predictable and reproducible if the solution conditions are maintained constant. it may change with the solution composition. In any p H study or any study involving different concentra(211 0 A Reutov and G M Ostapchuk D o k l A k a d Nauk SSSR 117 826 11957i (221 3 F pollard and J V Westwood J A m e r Chem Soc 87 2809 (1965) ( 2 3 ) D R Poilard and J V Westwood J 4 m e r Chem Soc 88 1404 (19661 ( 2 4 i M M Kreevoy and E A Walters J Amer Chem SOC 89. 2986 (1967) ( 2 5 ) R A G Marshall and D R Pollard J Arner Chein Soc 92, 6723 (19701 (261 R A G Marshall and D R Pollard J Organometal Chem 27, 149 119711 (271 K P Bu'in A N Kashin A B Ershler V V Strelets I P BeletsKava and 0 A Reutov I Organometa1 Chem 39. 39 (1972) ( 2 8 ) R B Simpson J Amer Chem Soc 8 3 . 4 7 1 1 ( 1 9 6 1 ) ( 2 9 ) T D Waugh H F Walton and J A Laswick J Phys Chem 59, 395 (1955) (301 G Schwarzenbach an? M Schellenberg Chi? Acta 48, 28 119651 ( 3 1 ) P Zanella G Plazzogna and G Tagllavlnl inorg Chim A c t a 2. 340 (19681
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 C h e m i s t r y and t h e A m e s L a b o r a t o r y - U S A E C ,
l o w a S t a t e U n w e r s i t y . A m e s , iowa 5007 0
Evidence is presented showing that an adsorption mechanism which was proposed in earlier work to explain the potentiometric response of the NaxW03 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 NaxW03 electrode by dissolved oxygen. These responses were found to be unique to the cubic NaxW03 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 ( 1 9 7 3 ) . ( 3 ) P. 6. Hahn. Ph.D. Thesis, IowaState University, Ames. Iowa, 1 9 i 3 . A N A L Y T I C A L C H E M I S T R Y , V O L . 46,
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