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tool for biochemists to perform quantitative, potentiometric, preferred method with a faradaic reactant determining the or characterization studies of known and unknown intact current level. biological systems. In addition, the technique allows for ACKNOWLEDGMENT control of redox levels by applications of charge injections We hereby acknowledge t h e kind assistance of John S. continuously or incrementally a t nanoequivalent quantities. Rieske in the preparation of the mitochondrial particles. T h e technique can readily be interfaced with other physical methods to study such systems. In particular, it should be LITERATURE CITED most advantageous in combination with those where t h e cell (1) E. C. Slater, Ann. Rev. Biochem., 46, 1015-1025 (1977). is not readily manipulated or requires remote operation. In (2) R. W. Estabrook, and M. E. Pullman, Methods Enzymol, IO, (1967). many newer techniques, the low signal level may require signal (3) P. Nichdls, and B. Chance, in “Molecular Mechanism of Oxygen Activatlf” 0. Hayaishi, Ed., Academic Press, New York, N.Y., 1974, p 479. averaging for extended periods of time. ICT control of such (4) J. W. Depiene, and L. Ernster, Ann. Rev. Biochem., 48, 201-262 (1977). systems would be most advantageous because of its ease in (5) B. G. Malmstrom, Q . Rev. Biophys., 6, 389 (1973). (6) P. L. Dutton, D. F. Wilson, and C. P. Lee, Biochemistry, 9, 5077 (1970). repetitive operation. (7) M. Erecinska, and D. F. Wilson, Arch. Bbchem., Biophys., 174, 143-157 Hendler et al. (9) have recently automated an electro(1976). chemical method to intact systems. They have pointed out (8) M. Erecinska, D. F. Wilson, and Y. Miyata, Arch. 8iOChem. Biophys., 177, 133-143 (1976). t h e advantages of using electrochemistry with computer (9) R. W. Hendler, D. Songco, and T. R. Clern, Anal. Chem., 49, 1908-1913 control for obtaining potentiometric information. Unfor(1977). (IO) J. J. Eiler, J. E. Goyan, L. D. Tuck, and C. C. Collins, J . Am. Pharm., tunately, their “current” pulse method has some serious ASSOC.,47, 497 (1958). limitations which could lead to erroneous and detrimental (1 1) B. Ke, F. M. Hawkridge, and S. Sahu, Proc. Natl. Acad. Sci. USA, 73, 2211 (1976). results. T h a t is, unless a sufficiently large amount of h n o ~ ~ n (12) D. F. Wilson, P. L. Dutton, M. Wagner, in “Current Topics in Bioenergetics”, depolarizer (could be a mediator) is present in the solution D. R. Sanadi, Ed., Vol. 5, Academic Press, New York, N.Y., 1973. to control the potential during the current pulses, the potential (13) D. F. Wilson, P. L. Dutton, M. Erecinska, J. G. Lindsay, and N. Sato, Acc. Chem. Res., 5 , 234 (1972). would increase (e.g., for an oxidation) to where the faradaic (14) F. M. Hawkridge and T. Kuwana, Anal. Chem., 45, 1021 (1973). reaction may involve undesirable components. The preferred (15) J. L. Anderson, T. Kuwana, and C. R . Hartzell, Biochemistry, 15, 3847 119761. method of coulometry such as in ICT, is by controlling the (16) k. Szentirmay, and T. Kuwana, Anal. Chem., 49, 1348 (1977). applied potential of the working electrode (vs. a reference (17) D. Yates, R. Szentirmay, and T. Kuwana, manuscript submitted for electrode) to a knoun, constant value where the electrolysis publication. (18) R. Szentirmay, Ph.D. Thesis, The Ohio State University, Columbus, Ohio, current is carried by a known depolarizer, e.g., mediator ti1978. trant. T h e problems of electrode potential memory and (19) R. Szentirmay, P. Yeh, and T. Kuwan,n, “Electrochemical Studies of Biological Systems”, D. T. Sawyer, Ed., ACS Symp. Ser., 38, 143 (1977). “instability” a t the working electrode and the incomplete (20) T.Kula, R. Szentirmay, E. Stellwagon, and T. Kuwana. presented at the optical absorbance changes for certain components in Annual Meeting American Chemical Society for Microbiology (1 978). Hendler’s work (9) could have been due to the “uncontrolled” (21) P. L. Dutton. and D. F. Wilson, Biochim. B i o. ~ .h v s Acta. . 346, 165-212 (1974). electrolysis conditions. These problems are accentuated when (22) B. Chance, R. Oshino, T. Sugano, and N. Oshino, Biochim. Bbphys. Acta, t h e current density is relatively high and the solution does 368, 298 (1974). not contain readily electrolyzable, faradaic reactants, such as (23) D. Wilson, and Y. Myata, Biochim. Biophys. Acta, 461, 218-230 (1977). (24) D. Yates, .DI ? ! Thesis, The Ohio State Universrty, Columbus, Ohio, 1976. mediators. T h e advantages of computer automation, as described by Hendler ( 9 ) , can also be applied to the ICT method, as conducted in our laboratory for several years (16. RECEIVED for review May 23, 1978. Accepted July 27, 1978. 27,24). I t cannot be overemphasized to those planning to use Work supported by funds from U S PHS Grant GM 19181. electrochemical methods to change redox levels of biocomPortions of t h e instrumentation utilized in this work were ponents that potential control of the working electrode is the provided by NSF Grant CHE76-81591.
Reexamination of Potentiometric Studies of the Oxidation of o-Tolidine M. A.
Ditzler and W. F. Gutknecht“
D e p a r t m e n t of Chemistry, Duke University, Durham, North Carolina 27706
Several studies have shown that optimizing conditions for the formation of the intensely-colored intermediate observed during the oxidation of o-tolidine leads to a reduction in the ESR signal normally associated with this oxidation process. This observation is in conflict with an earller theory based on potentiometric data, that is, that the intensely-colored intermediates observed during the Oxidation of this and similar compounds (i.e., benzidine analogues) are radical-cations. Here the earlier data are reevaluated and new potentiometric data are presented. It is shown that these data are consistent with a dimeric intermediate, which Is in agreement with the ESR results.
o-Tolidine and related compounds such as benzidine and 0003-2700/78/0350-1883$01.00/0
o-dianisidine (Le., [l,l’-biphenyl]-4,4’-diamines) have been widely used in analytical chemistry. ‘They have recently been used as model compounds in spectrclelectrochemical studies (1,2),as indicators in trace element detection schemes ( 3 , 4 ) , in the production of ion-selective electrodes ( 5 ) , and for chemically modifying electrode surfaces (6). Their analytical utility is due, in part, to the fact t h a t these compounds can be reversibly oxidized to give intensely colored products. If the oxidation is carried out in a weakly acidic solution, two products are observed. One product is at a maximum concentration after one equivalent of an oxidizing agent has been added, while a second product reaches a maximum concentration after two equivalents of the oxidizing agent have been added. A survey of the literature shows that there is a disagreement over the identity of the first of these products. 0 1978 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978
Oldfield and Bockris have made potentiometric studies of o-tolidine and related compounds and concluded that the first oxidation product is a semiquinone radical-cation and the final product is a quinonediimine ( 7 ) (Equation 1).
VHP
"2
Recently this conclusion was supported by further potentiometric studies by Ghimicescu and Dima (8). However it has been shown for o-tolidine that, if the concentration of the intermediate species is optimized by adjusting t h e p H , the ESR signals disappear (9,lO). This observation led to dilution studies by these same workers which suggest that the observed intermediate is dimeric. Furthermore, the presence of a strong, broad absorption band a t 630 n m suggests that the dimer is a charge transfer complex, which is proposed to exist in equilibrium with the reduced and oxidized forms of o-tolidine as shown in Equation 2 (9). ci, "r; w; I,
c '3 121
indicates t h a t a dimeric intermediate is unlikely ( 7 ) . 2. The index potentials (difference in potential at one-half and either one-fourth or three-fourths completion of the titration) observed in the titrations of o-tolidine with aqueous bromine a t 25 "C are greater than 14.1 mV. I t was reported than an index potential greater than 14.1 mV (the theoretical value for a two-electron oxidation) establishes the presence of a semiquinone radical-cation intermediate ( 7 ) . M 3. A graph of the potentiometric titration of 5 X o-tolidine with S2082-shows two vertical regions. T h e first corresponds to the addition of sufficient Sz02-to remove one electron per o-tolidine molecule and the second corresponds to removal of two electrons per molecule. This, it was argued, supports the existence of a stable radical-cation (8). 4. It was reported that the shapes of the titration curves obtained when o-tolidine is titrated with aqueous bromine are independent of the initial concentration used, that is, they are superimposable. If the intermediate is a dimer, the curves would not be superimposable ( 7 ) . I t can be shown that the first three observations actually are not inconsistent with the existence of a 1:l complex (Le., the dimer shown in Equation 2 ) . We have made measurements similar to those which served as the basis for the fourth observation and found results that agree with those expected for a 1:l complex, in conflict with those reported in the earlier study. An examination of the four arguments follows. 1. A potentiometric titration curve is symmetric about the midpoint if the following condition is met:
tH, "2
This apparent contradiction (radical-cation us. dimer) has not been explained. In an attempt t o clarify this matter, previously reported data and arguments have been reexamined; in addition, new potentiometric data have been collected via the oxidation of o-tolidine by bromine. In the following discussion we show that the potentiometric data previously used to support the existence of a radical-cation intermediate are also consistent with the existence of a dimeric intermediate, e.g., a 1 : l complex of the type shown in Equation 2.
where E is the measured potential and the subscripts on E represent the moles of sample titrated. Here p represents the moles of sample initially present and y is any number from zero to 1/2p. This condition can be shown to hold for the case when a 1:l complex is present during the titration. T h a t is, since this type of complex ties u p equal quantities of both the oxidized and reduced species, would be equal to E? and thus Equation 3 can be rewritten as
EXPERIMENTAL Materials. o-Tolidine (Eastman) solutions were prepared in spectral grade acetone (Fisher). The buffer solutions were prepared from acetic acid (B & A reagent grade) and sodium acetate (Fisher ACS certified) so as to have a pH of 3.5 and a concentration of 0.5 M in acetate ion. Aqueous bromine (Fisher reagent grade) solutions were made 1.0 M in KBr (Merck, reagent grade) to lower the Brz volatility. Asz03 (Mallinckrodt, reagent grade) solutions were prepared for use as primary standards. A spiraled platinum wire and a saturated calomel electrode served as the electrodes. Potentials were measured with a Beckman SS-3 pH meter. Procedures. Bromine solutions were standardized by addition of excess KI (MCB, reagent grade) followed by titration of the resulting iodine with As203(11). o-Tolidine solutions were diluted with sufficient buffer to make solutions that were 90% aqueous-10% acetone. Portions of these solutions, 100 mL, were titrated with bromine solutions that were of appropriate concentration to give only a 170volume change. Titrant was added in quantities which allowed at least 10 potential measurements before the end point. Since the oxidized form of o-tolidine is unstable in weakly acid solutions (9,I O ) , care was taken to ensure that all titrations were performed uniformly and rapidly (about 2 min per sample).
This same property allows us to express E(,:,,, as
RESULTS AND DISCUSSION In t h e potentiometric studies there were four arguments presented in favor of t h e intermediate being a semiquinone radical-cation rather than a dimer. 1. T h e titration curve obtained when o-tolidine is titrated with aqueous bromine is symmetric. I t was argued that this
(4)
Eu;zp+qi= En +
0.059 log
-
n
(
+
and
'
-
m,
(6)
5p-q-m I
where u: and m are, respectively, the moles of complex present after (1/2p - y) and (1/2p + q ) moles of sample have been titrated. If we substitute these expressions into Equation 4 and simplify, we get 1 1 -p+q-w -p+y-m 2 2 1I
-p-q-w 2
1
-p-q-m 2
From inspection it is obvious that this equality holds for all possible values of p and q (the requirement for symmetry) when w = m , that is, when the moles of complex present at (1/2p - y) and (1/2p + q ) are equal. I t can be shown that u: = m by examining K,, expressions. Using the conditions present a t (1/2p - y) and (1/2p + y) and assuming the
ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978
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equilibrium in Equation 2 holds, K,, for formation of the complex can be written as either
(W)(H+)
Keq =
(iP .)( -
q-
fp
+ 4 - .)
E (mV, P t vs SCE)
or K,, =
( f P + 4 - m)( ;P - 4 - m ) These expressions are written in terms of moles rather than concentrations since the volume terms cancel. Setting the two Kq expressions equal, cross multiplying, and collecting terms gives
+
Equation 8 implies that either w = m or (1/2p - q)(1/2p q ) = um. However, the second possibility is not physically
significant. This becomes apparent when it is realized that the moles of both species involved in complexation, u and m, must both be less than (1/2p - q ) since K,, (complex formation) < m . Since w ,m, and p represent moles of a species present they must be nonnegative; therefore, if (1/2p - q)(1/2p + q ) = w m ,(1/2p + q ) must be less than (1/2p - q ) , which would imply that q is negative. However, this is not possible since the preceding argument is based on the earlier definition t h a t 0 < q < 1/2p. An analogous argument based on 0 < q < -1/2p would lead to the contradiction that q be positive. Therefore, u; and m must be equal. This implies that Equation 7 holds and thus the potentiometric titration curves would be symmetric even if the 1:l complex is reversibly formed during the titration. Thus, the observed symmetry is not inconsistent with the formation of a 1:l complex as has been reported ( 7 ) . 2 . If the 1:l complex forms, the index potential for otolidine can be expressed as
where again, E is the measured potential and the subscripts on E represent the moles of the o-tolidine that have been titrated; here y represents the moles of the complex present. Since, as was explained earlier, = Eo even with the complex present, Equation 9 can be simplified to
If the amount of complex present (.y) is greater than zero, but less than 1 / 4 p (as it must be), the term in the parentheses would be greater than three, and hence E(hder) would be greater than 14.1 mV. Thus the observation that the index potential for o-tolidine is greater than 14.1 mV at 25 “C does not establish the presence of semiquinone radical-cation as has been reported (7). 3. Some years past Michaelis and Schubert carried out theoretical studies of the redox titrations of analogous systems (12). In one case they assumed that reactant and product formed a 1:l complex; for this case they subsequently predicted that if the mathematical product of K & (the apparent formation constant for the complex) and the initial reactant
figure la
figure lb
Figure 1. A comparison of experimental ( l a ) and calculated ( l b ) potentiometric titration curves for the oxidation of o-tolidine with Br,. Curves shown for M (M), M (O), M (A)o-tolidine at pH 3.5. Titrant concentrations were such that volume changes were negligible. Calculated potentials were based on a Keqof 11 for the formation of a 1: 1 complex between the oxidized and reduced forms of o-tolidine.
concentration was greater than four, an additional vertical region would appear in the titration curve at the point where one half of the reactant had undergone a two-electron oxidation. Phenomenologically this vertical region corresponds first to a rapid decrease in “free” reactant as it is both oxidized and complexed by product, and subsequently, a rapid increase in “free” product as no “free” (uncomplexed) reactant remains to complex with product. The midpoint of this region corresponds then to the concentration of the complex being a t a maximum. A value for K’eqof 1.03 x lo5 ( p H 4) has been
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978
reported for the o-tolidine system (9). Thus it would be predicted that a titration curve of an o-tolidine sample at p H 4 with an initial concentration greater than 3.88 X M would contain the additional vertical region. So too, the M o-tolidine soobservation t h a t the titration of a 5 X lution produces two vertical regions does not necessarily indicate that there is a free semiquinone radical-cation formed, as has been suggested (8). 4. We have carried out potentiometric titrations of lo4, and M solutions of o-tolidine with aqueous bromine using Pt and SCE electrodes. Plots of the potentials measured during the titrations are shown in Figure la. As in the earlier study, the titrant concentrations were such that volume changes were insignificant. I t is apparent that our results do not agree with those reported earlier in which the titration curves were found t o be superimposable regardless of the initial concentration of the o-tolidine ( 7 ) . The p H of the solutions used in the earlier study was not reported, but from the midpoint potentials of the titrations shown and a graph of p H vs. the midpoint potentials presented in the earlier study, it is apparent that the p H was less than 2.0. This may explain the discrepancy since the intermediate species is either absent or present only in low concentrations at such a low p H (9,lO). The solutions represented in Figure l a were buffered at a p H of 3.5, which is within the range over which the intermediate form is observed ( I O ) . Figure 1b shows the theoretical titration curves for aqueous o-tolidine solutions of analogous concentrations and p H , assuming the 1:l complex is formed according to the reverse of Equation 2 and using a previously reported K,, of 11 (the absolute formation constant) (9). In preparing the theoretical curves, it was assumed that the solubility of the complex is M not exceeded. In reality this is not the case for the solution ( 7 ) ,so a 10% acetone-90% water solvent was used in all cases to ensure complete solubility throughout. T h e difference between the theoretical and experimental curves due to this difference in solvents should be small since it has been reported that the concentration of the intermediate species is only slightly affected by switching from water to a water-acetone solvent (7). The difference between the expected point of intersection of all curves a t 50% (where E
= Eo, even with complex present) and the experimental 58% is probably due to the fact that the oxidized form of o-tolidine gradually decomposes under the experimental conditions used ( 9 , I O ) . This would cause the oxidized and reduced forms of o-tolidine to be of equal concentrations later in the titration; hence, the potentials would be independent of the initial concentrations a t a later point, and thus the titration curves should intersect a t a later point. The curves are not superimposable after 100% oxidation as might be expected since the oxidizing solutions contained different bromine to bromide ratios. Taking these points into consideration, it is apparent that the change in shapes of the curves resulting from a variation of the initial o-tolidine concentration corresponds closely to that predicted on the basis of the 1 : l complex forming. In conclusion, we have shown that the results of the potentiometric and ESR studies of the colored intermediate observed during the oxidation of o-tolidine are not inconsistent. Of the four potentiometric measurements that appeared to be inconsistent with the ESR data, three have been shown to actually be consistent with the ESR data, and the fourth measurement becomes consistent when repeated a t a higher pH.
LITERATURE CITED (1) T. P. DeAngelis, R. W. Hurst, A . M. Yacynych, H. B. Mark, Jr., W. R. Heineman, and J. S. Mattson, Anal. Chem., 49, 1395-1398 (1977). V. E. Norvell and G. Mamantov, Anal. Chern., 49, 1470-1472 (1977). C. Ghimicescu, M. Stan, and 6.Dragomir, Taknra, 20, 246-247 (1973). J. Mal? and H. Fadrus, Analyst(London), 99, 128-136 (1974). M. Sharp, Anal. Chirn. Acta., 61, 99-114 (1972). J. F. Evans and T. Kuwana, Anal. Chem., 49, 1632-1635 (1977). L. F. OldfieM and J. O'M. Bockris, J. Phys. ColloidChern.,55, 1255-1274 (1951). (8) C. Ghimicescu and F. Dima, Talanta, 23, 67-69 (1976). (9) T. Kuwana and J. W. Stro' k Discuss. Faraday Soc., 45, 134-144 (1968). (IO) J. D. Johnson and R. &e;by, Anal. Chem., 41, 1744-1750 (1969). (1 1) A. I . Vogel, "Quantitative Inorganic Analysis", 3rd ed.,Longmans, Green & Co. Ltd., London, 1961, p 355. (12) L. Michaelis and M. P. Schubert, Chem. R e v . , 22, 437-470 (1938). (2) (3) (4) (5) (6) (7)
RECEIVEE for review May 22,1978. Accepted August 28, 1978. One of us (M.D.) was supported by an ACS Analytical Division Fellowship sponsored by the Society for Analytical Chemists of Pittsburgh.
Adsorption and Polymeric Film Formation at Mercury Electrodes by Solutions of Lead(I1) and Chelating Ligands Containing a Thioether Group Bruce A. Parkinson and Fred C. Anson" Arthur Amos Noyes Laboratov, California Institute of Technology, Pasadena, California 9 1 125
The adsorption on mercury of the complexes of several chelating carboxylate ligands bearing thioether groups with Pb(I1) and some other di0 metal cations Is examined. The extraordinarily large adsorption observed with a number of complexes is attributed to the formation of new phases on the mercury electrode surface. The structure of the adsorbed films may resemble the polymeric crystals formed by several metal salts of the same ligands.
As part of continuing studies of the surface and coordination chemistry attending t h e adsorption of metal cations on 0003-2700/78/0350-1886$01 .OO/O
mercury electrodes in the presence of adsorption-inducing ligands ( I ) , we have examined the ability of a series of carboxylic ligands containing thioether groups to induce the adsorption of lead(I1) and some other d" metal cations. The idea was to exploit the expected propensity of the sulfur atom in the thioether group to adsorb on the mercury surface while the carboxylate arms of the ligand chelated the metal cation. Somewhat to our surprise we found that the presence of the thioether group does not result in particularly strong adsorption of the ligands. However, the complexes of such ligands with heavy metal cations, especially lead, are strongly attracted to mercury surfaces. The observed behavior is novel
sa 1978 American Chemical Society
