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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 observation t h a t the titration of a 5 X M o-tolidine solution 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
ANIALYTICAL CHEMISTRY, VOL. 50, NO 13, NOVEMBER
in that neither the ligand nor the metal cation exhibit strong spontaneous adsorption when present alone. In this report, the combination of surface and coordination chemistry that is responsible for the observed behavior is discussed. It might be termed “metal-activated-ligand-induced adsorption”. T h e dependence of the adsorption on the ligand concentration and the electrode potential show many similarities with behavior that has been described by Murray and co-workers (2-5) for t h e halide-induced adsorption of lead(I1) a n d thallium(1) from solutions in which the limit of the metal halide solubility is approached (but not attained). “Crystalline bilayers” can be obtained in such cases in which the total quantity of metal cation present on t h e electrode surface exceeds (by about twofold) that calculated for a close-packed monolayer (3). Several of the thioether ligands utilized in these studies induce the adsorption of comparably large quantities of Pb(I1) in the form of what appears to be a new phase on the electrode surface. T h e formation, properties, and possible structures of these adsorbed films are described in what follows.
EXPERIMENTAL Reagents. The thioether ligands examined in this study are listed in Table I. They were obtained commercially in the form of the parent acid: 2,2’-thiodiacetic, 3,3’-thiodipropionic (Matheson, Coleman and Bell); thiodisuccinic, methylenebisthioacetic, and thioacetosuccinic (Evans Chemetics); ethylthioacetic and diglycollic (Aldrich). The acids were converted to sodium salts and recrystallized repeatedly from water-ethanol until no odor could be detected from the solids. As a final test for the absence of traces of mercaptans, cyclic voltammograms were recorded in 0.1 M solutions of each ligand with a hanging mercury drop electrode which was exposed to the solution for 30 to 60 s. Insufficiently purified samples frequently produced current peaks or sharp spikes so that a smooth, featureless background current between +200 and -600 mV was a requirement for each solution of ligand before it was utilized in adsorption experiments. Metal cations were obtained from reagent grade nitrate salts. Oxygen was removed from solutions with prepurified argon or nitrogen that had been passed through a vanadium(I1) solution. Triply distilled mercury (Bethlehem Instrument Co.) and water were employed. Apparatus and Procedure. A conventional hanging mercury drop electrode (Brinkmann Instruments Co.), a two-compartment cell, and a sodium chloride-saturated calomel electrode were employed. The quantity of each complex adsorbed was determined by double potential step chronocoulometv (6) with a PDP 11/40 minicomputer and an associated electrochemical instrument based on conventional utilization of operational amplifiers. Data acquisition rates and electrode equilibration times were adjusted to achieve adsorption equilibrium. With solutions of Pb(I1) the measured adsorption values were independent (*lo%) of the equilibration time for times from ca. 10 to 100 s in unstirred solutions. However, with several complexes of Cd(I1) the measured adsorption remained time-dependent at all equilibration times from 5 to 60 s. The few values given for Cd(I1) adsorption in Figure 1 were obtained after the electrode was exposed to the solution for 5 s. They were reproducible but are not equilibrium values. At the high coverages obtained with Pb(I1) complexes, the concentration of adsorbed Pb(I1) did not reach its final equilibrium value until late in the reverse potential step. The resulting small errors in the value of the charge consumed by the electrical double layer ( 7 ) were neglected in evaluating the much larger quantities of adsorbed reactant from the intercepts of the chronocoulometric charge-time plots. The double layer capacitance and cell resistance were obtained from ac impedance measurements with an applied ac voltage of 5-mV amplitude and lo3 Hz. The ac signal was imposed via a Princeton Applied Research Model 173 potentiostat and the in-phase and quadrature ac responses were obtained with an Ithaco Model 391-A Lock-in-Amplifier and a Fluke Model 2000A Digital Voltammeter.
1978
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RESULTS AND DISCUSSION Adsorption of Ligands. Data on the adsorption of simple carboxylate anions on mercury are relatively sparse (8- 10). Pospisil and Kuta (9),who examined the adsorption of maleic and succinic acids and the corresponding monoanions, concluded that the netural acids behave similarly to other uncharged organic adsorbates in that their adsorption is maximal near the point of zero charge. T h e acid anions, on the other hand, resemble simple inorganic anions in showing steadily increasing adsorption as the charge on the mercury surface becomes more positive. Quantitative data on the adsorption of thioethers or of molecules bearing both carboxylate and thioether groups have not been reported. We examined the adsorption of several ligands qualitatively by measuring the change in the electronic charge density on the electrode produced by the addition of the ligand to solutions of N a F or NaC104 (0.3 M) adjusted to p H values between 6 and 7 where the ligands were anionic. For example, a t potentials between -100 and 800 mV, the presence of 0.05 M thiodiacetate produced small, positive increases in the charge density similar to those observed with the hydrogenmaleate anion (9). The increase in charge density is much smaller at potentials negative of -800 mV, presumably because anion adsorption becomes insignificant at sufficiently negative surfaces. At potentials positive of -100 mV, larger increases in charge density are produced by the presence of thiodiacetate anions and it seems likely that the sulfur atom of the thioether may begin to participate in the adsorption a t the most positive potentials. Polarograms of solutions containing 1 to 2 m M thiodiacetate anions show no welldefined anodic wave corresponding to the formation of a Hg(I1)-ligand complex. This is not surprising if the formation constant for the Hg(I1) complex is comparable to the known values for Pb(II), Cd(II), and Zn(I1). ca. lo3 M-I. All of t h e ligands listed in Table [ produced charge-potential curves that were qualitatively similar to that of thiodiacetate. T h a t is, none of the ligand anions appears to be adsorbed on mercury as strongly ab, for example, iodide, thiocyanate, or bisulfide. Adsorption of Pb(I1)-Ligand Complexes. Despite the relatively weak interaction between the free ligands and mercury electrodes, the Pb(I1) complexes of all of the ligands in Table I show extraordinarily large adsorption. With TAS and T D S as ligands, as much as 8 X 10 lo mol cm12 of Pb(I1) is attached to the electrode surface. The adsorption can be eliminated by lowering the p H below 2 where the Pb(I1)ligand complexes are unstable. Figure 1 summarizes some representative data for four of the ligands. The other ligands in Table I behaved similarly except for thiodisuccinate which deviated somewhat from the common pattern shown in Figure 1. (The complete data set is available in reference 11). There are several remarkable features in the behavior shown in Figure 1. T h e Pb(I1) complexes of the ligands all exhibit a n adsorption which is rather insensitive to the electrode potential over a significant range with a shallow minimum usually appearing between +50 and -50 mV. However, outside of this central potential range, the adsorption increases both a t more negative and more positive potentials. In many cases, plots of adsorption vs. potential exhibit a common intersection point on the positive side near +200 niV. A similar common intersection point also appears to result on the negative side if the plots are extrapolated to potentials beyond that a t which the reduction of Pb(I1) commences. These common intersection points may signal the formation of particularly stable configurations of the adsorbed films toward which solutions of all compositions converge. However, the significantly larger value of the adsorption a t the intersection point on the negative side indicates that two different stable forms (or
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978
1
*-
TDP
1
I
-
-
1
IOT
I
,
1I
TIS
r;Y e 50-
ol.-.
.
Figure 1. Representative adsorption behavior of Pb(I1) complexes of four ligands. The first plot in each pair gives the potential-dependence and the second plot the ligand concentration dependence of the adsorption. Unless otherwise noted all measurements were made with 0.5 m M Pb(I1) in 1 M NaCIO,. Other experimental parameters: For TDA: first plot: concn. of TDA: (A) 0.8, (B) 1.4, (C) 2.2, (D) 3.0, (E) 6.0, (F) 18.4 mM. Second plot: potential vs. SCE: (W) +200 mV; ( 0 )0 mV; (A)-300 mV. Supporting electrolyte 0.1 M NaCIO,: (0)+200 mV; ( A ) 0 mV; (0) -300 mV. For TDP: first plot: concn of TDP: (A) 0.4, (B) 0.8, (C) 1.2, (D) 1.6, (E) 2.0, (F) 2.8, (G) 4.0 mM. Second plot: potential vs. SCE: (0) +200 mV; (0)0 mV; ( A ) -300 mV. For ETA: first plot: [Pb(II)] = 1.0 mM; concn of ETA: (A) 0.2, (B) 0.6, (C) 0.8, (D) 1.2, (E) 1.5, (F) 2.0, (G) 2.5, (H) 3.5, (I) 5.5, (J) 7.5 mM. Second plot: potential vs. SCE: ( A ) +200 mV; (0) 0 mV; (0)-300 mV. For TAS: first plot: concn of TAS: (A) 0.6, (B) 0.8, (C) 1.0, (D) 1.2, (E) 1.5, (F) 2.0, (G) 2.8 mM. Second plot: potential vs. SCE: (0)+200 mV; (0)-300 mV; [Pb(II)] = 1 mM: (W) $200 mV; (A)0 mV; ( 0 )-300 mV. Lowest lying curve: 1 mM Cd(I1) (no Pb(I1)) at -300 mV
compositions) of the films are involved. Equally noteworthy is the narrow range of ligand concentrations within which the adsorption of Ph(I1) rises from low to very high values that are much less sensitive to further increases in the ligand concentration. T h e sudden increase in adsorption occurs in the range of ligand concentrations where the fraction of Pb(I1) that is complexed undergoes its most rapid increase. However, the adsorption increases much more abruptly than does this fraction. The behavior resembles a phase transition rather than a smooth change in speciation. I t is reminiscent of the behavior reported by Murray and co-workers ( 3 4 , Armstrong e t al. ( 1 2 ) . and Sluyters and co-workers (13) in which unusually large and abruptly changing adsorptions were interpreted as arising from the formation of a new phase on the electrode surface. The solubilities of the lead salts of most of the ligands listed in Table I have not been measured but precipitates do result with several of the ligands if the concentrations of Pb(I1) and ligand are increased to values somewhat larger than were employed in the adsorption measurements. For example, the solubility of Pb(I1) in a millimolar solution of TDA is only ca. 1.4 m M (14). It seemed reasonable, therefore, to examine the possibility that a new surface phase is formed under the conditions corresponding to the sudden increases in adsorption shown in Figure 1. Murray and co-workers were able to correlate abrupt changes in the adsorption of Pb(I1) and Tl(1)
from halide solutions with a reasonably constant (“surface solubility”) product of the bulk concentrations of the anion and metal cation (3-5). No similar correlation was observed with the ligands investigated in this study. Characteristic properties of phase formation on electrode surfaces have been described by Armstrong and Barr (15). They include a sudden decrease in the effective differential double layer capacitance to values well below the background capacitance as well as increases in the apparent cell resistance. Figure 2 shows how the double layer capacitance and cell resistance are affected by the adsorption of Pb(I1) solution containing the ligand TAS. Throughout the potential range (-100 to -350 mV) where t h e adsorption is large, the capacitance is depressed and the cell resistance is enhanced compared with their values in the supporting electrolyte alone. As the adsorption decreases rapidly between -100 and 50 mV the capacitance increases and the resistance falls. (The steep rise in capacitance a t more positive potentials is probably associated with anodic depolarization of the mercury electrode.) At potentials more negative than -400 mV, where any adsorbed film containing Ph(I1) is destroyed by the reduction of Pb(I1) to Pb(Hg), the capacitance increases again to within 10% of the background value. This behavior matches closely that to be expected when initial adsorption gives way to the formation of an insulating phase on the electrode surface (15).
ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978 Table I. Ligands Examined
ligand
abbreviation
log of formation constants for PbUI) complexes (20)
formula CH,CH, I
ethylthioacetate
ETA
S' \
3.7 2.9
CH,CO; CH,CO;
/
thiodiacetate
TDA
S
3.6 \
CH,CO; CH, CH, CO; I
thiodiTDP propionate
2.7
S \
CH,CH,CO; SCH,CO; /
methvlenebischioacetate ~.~
MBTA CH
3.7 \
SCH,CO; CH,CO; /
thioacetato- TAS succinate
...
S \
CHCO; I
CH,CO; CH,CO; CHCO; /
thiodisuccinate
TDS
S \
CHCO; I
CH,CO; I
I
,
,
,
,
TAS
Figure 2. Effect of adsorption on differentialcapacitance and apparent cell resistance. 1 M NaCIO, supporting electrolyte: Differential capacitance (curve 1, left ordinate) and cell resistance (curve 3, right ordinate). 1 M NaCIO, + 0.5 mM Pb(I1) -k 2 mM TAS: Quantity of adsorbed Pb(1I) (curve 5, lefi ordinate),differential capacitance (curve 2, left ordinate) and cell resistance (curve 4, right ordinate)
Effect of Adsorption on the Electronic Charge Density of the Electrode. Under conditions where the quantity of Pb(I1) adsorbed is well below the approximate value calculated for a close-packed monolayer, the change in electronic charge density produced by the adsorption is in fair agreement with that predicted thermodynamically on the basis of the Gibbs adsorption equation. Thus, the potential dependence of the adsorption, (anFT/aE),, can be used to calculate the effect of adsorption on t h e electronic charge density, ( n F /
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2.3R7Naqm/a log C)E. ( p is the chemical potential of the adsorbing species, approximated as 2.3RT times the logarithm of its concentration, and q" is the electronic charge density.) As an example. 30 p C cm-2 of Pb(I1) is adsorbed a t -175 mV from a solution containing 1.0 m M Pb(I1) and 4.0 m M TDA. T h e corresponding values of (l(nFI'/lE),, a n d ( n F / 2.3RT)(Aqm/Alog C), are ca. 0.16 and 0.18 HC cm-2 mV ', respectively. By contrast, under conditions where much sharper potential dependences of the adsorption prevail, the apparent values of these two differential coefficients are no longer in approximate accord. For example, the adsorption of Pb(I1) at -200 mV from a 4 m M solution of T D S increases from 5 to 73 pC cm-2, and the electronic charge on the electrode is decreased by 3 pC cm-2 when the concentration of Pb(I1) is increased from 0.2 to 0.4 mM. T h e value of ( n F / 2 . 3 R T ) ( l q m / Alog C), is, t h w , 0.33 F C cm-2 mV-'. However, the corresponding value of ( A ( n F r ) / A E ) ,is only ca. 0.07 pC cm mV'. The clear implication is that the Gibbs adsorption equation is not applicable in the second case, presumably because the adsorbed film effectively insulates the electrode surface from the solution phase so that its electronic charge does not remain in thermodynamic equilibrium with the solution. This interpretation is consistent with the very large quantity of Pb(I1) that is present on the surface when it appears to become insulated from the solution. I t is also noteworthy t h a t under these conditions the quantity of Pb(I1) present in the adsorbed film shows very little dependence on the electrode potential but remains a strong function of the concentration of Pb(I1) in the bulk of the solution. Therefore, the high plateaus in the plots of adsorption vs. potential do not necesswily represent saturated surface layers. Instead, they may siignal the presence of a compact surface layer containing sites which are effectively insulated from the electric field emanating from the charged electrode surface, but where adsorption of additional Pb(I1) may be induced by increasing its bulk concentration. Very similar behavior has been reported (13) for the thick films that form spontaneously on mercury electrodes in 1 M KC1 solutions in the presence of sufficiently high concentrations of Pb(I1). In (more dilute) bromide and iodide electrolytes, limiting coverages of adsorbed Pb(I1) are obtained that are independent of both the electrode potential and the bulk concentration of Pb(I1) ( 4 ) . Apparently true saturation of the surface layer is achieved in these latter electrolytes. Composition and S t r u c t u r e of the S u r f a c e Films. We were unable to devise a successful method for assaying the ligand content of the surface films. T h e relatively weak interaction of the free ligands with Hg2+ means that no well-defined anodic waves were present which might have been utilized to measure the quantity of ligand present on the surface. Furthermore, the lack of equilibrium between the electrode surface and the solution phase in the presence of the densely packed films prevented the use of conventional thermodynamic methods for relating changes in interfacial tension to the superficial excesses of ligand (or metal cations). However, despite the lack of full analytical data on the composition of the surface films, some speculation on its possible structure may be warranted on the basis of recent crystal structures that have been determined for the acid form of TDA (16),its cadmium(I1) salt and several additional carboxylate amino acid salts of cadmium(I1) (28-21). (No crystal structures for the corresponding lead derivatives are available, perhaps because of the greater difficulties in locating the positions of lighter atoms in the presence of an atom with as high an X-ray scattering cross section as lead.) All of the substances mentioned crystallize from aqueous solutions to give polymeric structures containing layers which are held together by hydrogen-bonded water molecules or bridging
(In,
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ANALYTICAL CHEMISTRY, VOL. 50, NO. 13, NOVEMBER 1978
Figure 3. Possible schematic structure for the adsorbed layer
carboxylate ligands. I t seems reasonable that the same substances might form layered structures on the surface of liquid mercury to which they are spontaneously attracted. A structure something like the highly simplified version shown in Figure 3 suggests itself. T h e atom that binds the array to the electrode seems most likely to be sulfur despite the weak adsorption of the free ligand because the complex of Pb(I1) with oxydiacetate anions, O(CH2CO2)?’, is much less strongly adsorbed than that with thiodiacetate. Molecular models suggest t h a t upon chelation with Pb(I1) the sulfur atom in t h e TDA molecule suffers few steric impediments toward binding to mercury atoms in the electrode surface. T h e interaction between Pb(I1) and the sulfur atom in its complex with TDA is not particularly strong because the formation constant for the analogous, sulfur-free Pb(I1)-oxydiacetate complex is over six times larger (22). Nevertheless, some Pb(I1)-sulfur interaction seems likely because a chargetransfer band appears in the UV spectrum of the TDA complex (14). Furthermore, the Zn(I1) complexes of TDA and TDS, while more stable than the Pb(I1) complexes (22).are not adsorbed on mercury. Thus, in addition to whatever steric advantages the sulfur atom enjoys because of the formation of a metal-ligand complex its propensity for binding to the mercury surface is apparently also altered by interaction with t h e metal cation in the complex. T h e structure shown in Figure 3 is only one of a variety of possible structures which can be constructed from molecular models. One virtue of the structure shown is that it can accommodate the observation that the adsorption remains dependent on the concentration of Pb(I1) under conditions where it has become independent of potential. T h e Pb(I1) sites in the upper-most layer may be only partially occupied by lead cations a t lower bulk concentrations of Pb(I1) while the primary, anchoring layer is fully metallated and insensitive to changes in potential. I n addition, an upper-most layer which exhibits cation exchange behavior a t the carboxylate groups would account for the observation that the adsorption of Pb(I1) from solutions of TDA increases substantially when the concentration of the sodium perchlorate supporting electrolyte is decreased from 1 to 0.1 M (Figure 1). I t is difficult to imagine an alternative explanation for this strong ionic strength dependence of t h e amount of Pb(I1) incorporated into the adsorbed film. If the crystal structure of the C(I1)-TDA hydrate (17) is used to estimate the quantity of a divalent metal cation that might be accommodated in a close-packed monolayer, a value of ca. 90 pC cm-* is obtained. This is greater than the measured values of Pb(I1) adsorbed in TDA solutions under any conditions. However, values as large as 150 pC cm-* were obtained with larger ligands such as TAS and TDS. These higher values may reflect the binding of more than one Pb(I1) by these multidentate ligands although the presence of multilayered films is not excluded. Adsorption of Other Metals. Inasmuch as the ability of the ligands studied to produce the strong adsorption of Pb(I1)-complexes depends upon the presence of a thioether
group in each ligand, it might seem reasonable to expect that every metal cation which forms complexes with the ligands would be comparably adsorbed. In fact, this is not the case. Not all of the ligands in Table I were tested with every metal cation studied but general trends became apparent from experiments with representative ligands. (i) The relative adsorbabilities of the metals tested is Pb(I1) > Cd(I1) > Tl(1) > Zn(I1). No adsorption of Zn(I1)-complexes was detected with any of the ligands investigated. (ii) The adsorption of Cd(I1) does not attain equilibrium as rapidly as that of Pb(I1). Substantial adsorption is obtained but under no conditions does it match the extraordinarily large surface concentrations obtained with Pb(I1). (iii) With most of the ligands Zn(I1) forms stronger complexes than Cd(I1) but none of the Zn(I1)-complexes were adsorbed detectably. (iv) TI(1) forms rather weak complexes with the ligands tested and their adsorption is detectable but weak. The variation in the adsorbability of the structurally similar metal complexes of the same ligand seems likely to be the reflection of two influential factors: the extent of meta-sulfur interaction in a complex and its solubility or tendency to form polymeric solids. The solubility of the metal salts of TDA increase in the same order as their adsorption decreases: P b < Cd < Zn. In addition, Zn(I1) shows less tendency to form polymeric sheets in its salts with ligands such as S-methylcysteinate than does Cd(I1) (and, presumably, Pb(I1)) (21). The need to invoke differences in metal-sulfur interaction in the complexes stems from comparisons of relative adsorptions a t low surface coverages where stabilizing interactions among adsorbed molecules of the complex are minimized and no new surface phase is formed. T h e large and persistent differences in the relative adsorbability of the Pb(II), Cd(II), TI(I), and Zn(I1) complexes under these conditions require an explanation that is specific to the metal involved. The degree of electronic interaction between each metal and the thioether sulfur atom seems the most likely candidate for the needed variable parameter.
CONCLUSIONS T h e thioether carboxylate ligands examined in this study displayed a surprising variety in the adsorption of their metal complexes despite the basic similarities in their structures. T h e thioether group in these ligands shows little tendency toward adsorption on mercury without participation by coordinated metal cations and there are large differences in adsorbability among the cation complexes tested. Thus, the thioether carboxylates do not provide a general means for anchoring metal ions to mercury electrode surfaces. Extremely large quantities of the Pb(I1) complexes of several of the ligands can be adsorbed in the form of a multilayered phase on the mercury surface. This phenomenon appears to correlate with the intrinsic solubility of the Pb(11)-ligand salt but large adsorption can be obtained from solutions which are not close to saturation with respect to the formation of bulk phases. T h e potential dependence of the adsorption matches that of many neutral organic adsorbates which suggests that the lower dielectric constant of the aqueous layer at the electrode surface may play an important role in controlling the adsorption. ACKNOWLEDGMENT Helpful discussions with Pamela Peerce are acknowledged with pleasure. LITERATURE CITED (1) F.C. Anson, Acct. Chem. Res., 8, 400 (1975) and references cited therein. (2) R. W. Murray and D. J, Gross, Anal. Chem., 38, 392 (1966). (3) C. M. Elliott and R. W. Murray, J . Am. Chem. Soc., 96, 3321 (1974).
ANALYTICAL CHEMISTRY, VOL. 50, (4) H. B. Herman, R. L. McNeeiy, P. Surana, C. M. Elliott, and R. W. Murray, Anal. Chem., 46, 1258 (1974). (5) C. M. Elliott and R. W. Murray, Anal. Chem., 48, 259 (1976). (6) F. C. Anson, Anal. Chem., 38, 54 (1966). (7) C. M. Elliott and R. W. Murray, Anal. Chem., 47, 908 (1975). (8) E. Blomgren, J. O'M Bockris, and C. Jesch, J . Phys. Chem., 65, 2000
(1961). (9) L. Pospisii and J. Kuta, Collect. Czech. Chem. Commun.,34,3047 (1969) (10) B. Damaskin, A. Frumkin, and A . Chizhov. J . Nectroanal. Chem., 28. 93 (1970). (11) B. A. Palkinson, Ph.D. Thesis, California Instiuie of Technology,Pasadena, Calif., 1978. (12) R. D. Armstrong, W. P. Race, and H. R. Thirsk, J . Electroanal. Chem., 23, 351 (1969). (13) M. Sluflers-Rehbach,J. Breukei, K. A. Gijsbersen, C. A . Wiznhorst, and J. H. Siuyters, J . Electroanal. Chem., 38, 17 (1972). (14) A. Napoii and P. L. Cignini, J . Inorg. Nucl. Chem., 38, 2013 (1976).
(15) (16) (17) (18)
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R. D. Armstrong and E. Barr, J . Electroanal. Chem., 20, 173 (1969). S. Paul. Acta Crystallogr., 23, 491 (1967). S. H. Whitlow, Acta Crystallogr., Sect B , 31, 2531 (1975). W. Harrison and J. Trotter, J . Chem. Soc., Dalton Trans., 956 (1972);
1923 (1974). (19) M. L. Post and J. Trotter, Acta Crystallogr. Sect. B , 30, 1880 (1974). (20) P. J. Flook, H. C. Freeman, C. J. Moore, and M. L. Scudder, J . Chem. Soc., Chem. Commun., 753 (1973). (21) P. DeMeester and D. J. Hodgson, J . Am Chem. Soc., 99, 6884 (1977). (22) L. G. Sillen and A. E. Martell, "Stability Constants",The Chemical Society, London, 1964.
RECEIVEDfor review May 8, 1978. .4ccepted July 20, 1958. This work, Contribution No. 5783, was supported by a grant from the National Science Foundation.
Determination of Thiols by Conductometric Titration with Mercury(I1) Chloride in Water and in N,N-Dimethylformamide L. M. Doane and J. T. Stock" Department of Chemistry, University of Connecticut, Storrs, Connecticut 06268
A conductometric device that allows simultaneous ac bridge balance, without regard to capacitance effects, and generation of a signal proportional to solution conductance has been constructed and used to determine solution conductance in titrations of thiols with mercury(11) chloride. Quantitative data for determination of submillimolar solutions of 1-propanethiol, 2-propanethiol, and 2-methyl-2-propanethiol in water and in DMF were obtained with an accuracy and precision of fl%. Determinations of aqueous cysteine solutions gave similar results. Chloride, bromide, iodide, or thiocyanate up to concentrations of 1 mM did not generally interfere in the aqueous titrations. Similar titrations in DMF were unsuccessful. The accurate titration of thioglycolic acid in aqueous solution required the presence of bromide, iodide, or thiocyanate ion. With 3-mercaptopropanoic acid, accurate titration was successful with added thiocyanate, but not in the presence of bromide or iodide.
Apart from recent work with high-frequency signals (1,2), conductometric thiol titrimetry with precipitating or complexing reagents has received little attention. Although high-frequency methods permit measurements without solution electrolysis or electrode deactivation effects, the titration curves are rarely linear (3). A great advantage of conventional low-frequency conductometry is the ease and precision of end-point location. Mercury(I1) chloride and silver nitrate have found wide use in thiol titrimetry, since the mercaptides formed have very large formation constants, e.g., for cysteine, log KfiHg(RS)21 43.5 ( 4 ) ; log KspfAgRSI = ca. -20 ( 5 ) . The results obtained in the conductometric titration of thiols with silver nitrate were not very satisfactory (6)and similar titrations with mercury(I1) salts have not been attempted. The present work shows that mercury(I1) chloride is a n effective titrant in thiol conductometry both in aqueous and in DMF solution. T h e use of DMF minimizes electrode deactivation effects because mercaptide precipitation is prevented. 0003-2700/78/0350-1891$01.00/0
One of the problems in the conventional Wheatstone bridge conductance methods is obtaining a sharp "null". One approach is to use a switching circuit that is insensitive to phase shifts (7, 8). T h e present authors have modified this circuit to eliminate losses due to transistor lbias and have added an analog circuit to give fast conductance readouts. Because the two outputs are independent, they can be used simultaneously.
EXPERIMENTAL Reagents. N,N-Dimethylformamide (DMF) was used as received from Eastman Kodak Co. I.-(+)-Cysteine (J. T. Baker Chemical Co.) was dried over P205. All 1hiols were obtained from the Aldrich Chemical Co. D,L-Penicillamine wm used as received. 1-Propanethiol, 2-propanethiol, and 2-methyl-2-propanethiolwere distilled under nitrogen. Thioglycolic acid and 3-mercaptopropanoic acid were distilled under reduced pressure. The thiols were stored under nitrogen. S t a n d a r d Solutions. Fresh standard 0.1 M thiol solutions were prepared daily. To avoid volatilization and oxidation of the liquid thiols, the following procedure was used. A volumetric flask having a Teflon stopper with a small center hole was flushed with nitrogen. The hole was then closed with a short plug of stainless steel rod and the flask was weighed. The appropriate amount of liquid thiol was then introduced by a :,yringe. The stopper was then plugged and unplugged to establish vapor pressure equilibrium, and the flask plus thiol weighed. To minimize loss of thiol, dilutions of the stock solutions were made by use of a syringe. With the exception of cysteine and D,L-peniCilhmine,which were diluted with outgassed water, 95% ethanol was used as diluent. Aqueous mercury(1I) chloride solution was standardized with potassium iodide (9). The standard solution in DMF was prepared by the direct weighing of mercury(I1) chloride that had been assayed with potassium iodide (9). Two-milliliter microburets (Roger Gilmont Instruments, Inc.) were used to deliver the thiol solutions, and the titrant. Instrumental. A circuit diagram of the conductometric device is shown in Figure 1. The device was assembled largely from dual operational amplifiers (National Semiconductor LM 1458 N) and standard electronic component:, (10). In the simplified block diagram in Figure 2, the signal from the bridge at point B is inverted by A3 and then fed into a precision limiter, which simulates ideal diode behavior (11). A similar limiter accepts the signal from point C. Positive half-cycles are thus removed from 0 1978 American Chemical Society
