Anal. Chem. 1984, 56, 1884-1890
1884
(v/v) HzS04at least 1h while maintaining the external buffer flow. This does not affect the electrode performance. It should also be possible to design this electrode in an autoclavable mode according to the principle described earlier (12). Registry No. Glucose, 50-99-7.
LITERATTJRE CITED (1) Clark, L. C., Jr.; Lyons, C. Ann. N . Y . Aced. Sci. 1982, 102, 29-45. (2) Guilbault, G. G. Enzyme Microb. Techno/. 1980, 2 , 258-264. (3) Barker, A. S.; Somers, J. P. In "Topics In Enzyme and Fermentation Biotechnology";Wlteman, A., Ed.; Wiley: New York, 1978; Vol 2, pp 120-151. (4) Kuiys, J. J.; Pesliakiene,M. V.; Samalius, A. S. Bioelectrochem. Bioenerg. 1981, 8 , 81-88. (5) Tsuchlda, T.; Yoda, K. Enzyme Microb. Techno/. 1981, 3,326-330. (6) Romette, J.-L.; Froment, B.; Thomas, D. Clin. Chim. Acta 1979, 95, 249-253.
(7) Bertermann, K.; Elze, P.; Scheller, F.; Pfeiffer, D.;Janchen, M. Anal. Lett. 1982, 15, 397-404. (8) Mandenius, C. F.; Danielsson, B.; Mattiasson, B. Biotechnol. Lett. 1981, 3,629-634. (9) Enfors, S.-0. Enzyme Microb. Techno/. 1981, 3,29-32. (10) Severinghaus, J. W. J . Appl. Physiol.: Respir., Envlron. Exercise Physiol. 1981, 5 1 , 1027-1032. (11) Johnson, M. J.; Borkowski, J.; Engblom, C. Biotechnol. Bioeng. 1984, 6, 457-468. (12) Cieiand, N.; Enfors, S.-0. Eur. J . Appl. Microbiol. Biotechnol. 1983, 18, 141-147. (13) Enfors, S.-0.; Cleland, N. In "Chemical Sensors"; Proceedings of the International Meeting, Fukuoka, Japan, Sept 19-22, 1983; Seiyama, T., Ed.; Elsevier: Amsterdam, 1983; pp 672-675. (14) Linek, V.; Benes, P.; Sinkule, J.; Hoiecek, 0.; Maiy, V. Biotechnol. Bioeng. 1980, 2 2 , 2515-2536.
RECEIVED for review November 14,1983. Accepted April 19, 1984.
Voltammetry of Halide Ions on Mercury Electrodes in Acetonitrile Marek Wojciechowski and Janet Osteryoung* State University of New York at Buffalo, Department of Chemistry, Buffalo New York 14214
Anodic oxidation of mercury In the presence of chloride, bromide, or iodide in acetonkriie was investigated by DC, NP (normal pulse), and RP (reverse pulse) polarography. For all of these ions the first anodic wave is dlffuslon controlled In halide and well sulted for analytical purposes. The product on the dlffuslon plateau is HgX,-, and up to one monolayer Is strongly adsorbed on the electrode surface. The second process involves formation of both HgX, and Hg,X,(s) whlch are in chemical equilibrium descrlbed by the reaction Hg 4HgX2 Hg2X~(s).
*
In nonaqueous solvents such tts dimethylformamide, nitriles, or acetone the halide ions cause depolarization of a mercury electrode involving complex processes of mercury oxidation which are basically different from those in aqueous solutions. While the formation of insoluble mercury(1) halides dominates in water, soluble mercury(I1)-halide complexes are produced in nonaqueous solvents causing the appearance of two polarographic waves. The more positive one is better defined and has much more negative half-wave potential than that in water (1). The differences between polarographicbehavior of halide ions in water and in acetonitrile (AN) are due mainly to the following interrelated factors: (a) AN is only very weakly acidic (pK, = 25 (2)),(b) solvation energies of Hg(1) and Hg(I1) cations are relatively lower and their activities are higher in AN than in water, (c) Hg(1) halides are much less soluble in AN than in water; pKSo= 37.3 (17.7), 37.4 (22.2), and 39.3 (28.3) for mercurous chloride, bromide, and iodide, respectively, in AN (HzO) ( 3 ) ,(d) Hg(I1) forms more stable complexes with halide ions in AN than in water; for example log @ H ~ = x ~41.3 - (14.0), 42.9 (19.7), and 45.9 (27.6) for X = C1, Br, and I in AN (H20) ( 3 ) . In a previous paper (4)discussing the mechanism of the first polarographic wave of chloride ions, we described the experimental evidence for product adsorption involved in that process. Quantitative analysis of this absorption phenomenon by means of double potential step chronocoulometry has been 0003-2700/84/0356- 1884$01.50/0
Table 1. Adsorption and Diffusion Data for Halide and Trihalogenomercurate Ions in AN (5)
iol0r,, X
mol[ cm
C1 Br I
2.6 2.8 3.1
1osDHgx3-, io5&, cm2/s cm2/s
1.8 2.1
2.4 2.1
2.3
3.4
mol2 s/dms
108t,Cx-2, mol2 s/dm6
0.295 0.293 0.328
1.99 2.05 2.00
Wt,C2HgX3-,
presented in the most recent publication (5). It was shown that oxidation of mercury in the presence of halide ions was coupled with the adsorption (up to a monolayer) of product (HgX,-). In solutions containing HgX3- ions the adsorption was proven to be diffusion limited. Formation of adsorbed and nonadsorbed HgX3- at potentials corresponding to a limiting plateau of the first anodic wave (DC mode) was found to be controlled by diffusion of halide ions. It was noted there that a t potentials more positive than -0.9 V reduction of adsorbed HgX, is slow (not completed in tens of milliseconds), while at -1.0 V it is fast (completed in less than 1ms). Because the covered electrode behaves much differently from the uncovered mercury surface, it is important in discussion of the results to known the coverage for each set of experimental conditions. Using the maximum surface coverages (Fm, mol/cm2) and diffusion coefficients, one can calculate the time necessary to produce a monolayer of HgX3(tm). If HgX3- ions are present in a solution, no faradaic reaction occurs in the potential range between -0.35 and -0.25 V and
tmCHgXa-2= 7rm2/4DHgX3-
(1)
where CQX, is the bulk concentrationof HgX,. For solutions of X- in the same potential range, HgX3- is generated on the electrode according to the diffusion-controlled reaction Hg 3XHgX3-(ads) + 2e(2)
+
-
and
tmCx-2= 97rm2/4Dx0 1984 American Chemical Soclety
(3)
ANALYTICAL CHEMISTRY, VOL. 56, NO. 11, SEPTEMBER 1984
1885
Table 11. NPP and DC Characteristics of the First Anodic Wave of Bromide Ions in AN NPP, t, = 10 ms
NPP, t, = 1 ms CB~, mM 0.05 0.10 0.20 0.50 1.00
-El 2, md 666 666 667 661 655
-Ell29
-iIirn,
-*
him,
PA
INppn
mV
PA
1.39 2.57 4.50 12.4 22.6
1.6 1.5 1.4 1.4 1.3
601 606 600 565 522
0.42 0.84 1.52 3.76 7.24
Table I contains diffusion and adsorption data for halides and their mercuric complexes ( 5 ) and the resulting constants of eq 1 and 3. In the first part of this paper we discuss the polarographic behavior of the first anodic process of mercury in the presence of bromide and iodide ions in AN. The analytical characteristics of the anodic wave are given and analytical and mechanistic applications suggested. As mentioned previously ( 4 ) ,for chloride or bromide ions the second anodic wave of mercury has an unusual shape and irregularities on the limiting plateau. In the next part of this publication we describe the phenomena involved in that process when chloride, bromide, or iodide ions are present in solution. Formation of HgX, or insoluble Hg2X2,as a final product, is considered under various time-concentration conditions. From thermodynamic studies of Coetzee and co-workers (3) it is known that HgX2 undergoes reaction with mercury (4) The equilibrium concentration, [HgX&, calculated from K of this reaction is 13,20, and 160 pM for chloride, bromide, or iodide, respectively. Thus assuming electrochemical generation of HgX2, one can predict whether formation of insoluble mercurous halide is possible under given experimental conditions. The discussion of electrochemical behavior of halide ions in AN is based on detailed studies in solutions containing halide ions alone and in mixtures with mercuric ions by means of normal pulse (NP), reverse pulse (RP), and sampled DC polarography over a wide range of time scales and concentrations. (Investigations by means of constant potential coulometry (large scale) and double potential step chronocoulometry are described elsewhere (6).)
EXPERIMENTAL SECTION Acetonitrile was Spectro-grade from Kodak. Tetraethylammonium chloride (TEAC), tetraethylammonium bromide (TEAB), and tetraethylammonium iodide (TEAI) (from Kodak) were recrystallized from absolute ethanol and dried under vacuum at 70 "C. Solutions 0.05 M in TEAC or TEAB and 0.04 M in T E N in AN were standardized by using the Volhard method and served as a source of halide ions. Mercuric halides (HgX2),ACS grade, were dried under vacuum at 70 "C. Solutions 0.05 M in HgC12,0.04 M in HgBr, or 0.005 M in HgIz were standardized by complexometrictitration employing MgClz,ammonium buffer (pH lo), ErioT, and EDTA solutions (the substitution method). Solutions of HgXzserved ~ t a9 source of mercuric ions. Since HgX, solutions in AN react vigorously with mercury, forming insoluble Hg2X2,we did not investigate solutions with CHg(II,/CX-> 0.5. Solutions of equimolar amounts of HgX, and TEAX were used in experiments with trihalogenomercurate ions (HgX,-). Tetraethylammonium perchloride (TEAP) used in this work as the supporting electrolyte (0.1 M solution in AN) was repurified and stored as described earlier (4). All experiments with DC, NP, or RP polarographywere carried out by using the same equipment and the same procedures as used before (4). A dropping mercury electrode (DME) had the following characteristic: drop time ( t d ) = 7.85 s (natural, measured in AN at open circuit and 84 cm height of the Hg column) and flow rate
INPP'
1.5 1.5 1.4 1.4 1.3
DC,t d = 0.5 S NPP, t, = 100 ms -EI/z, iIirn, -El 21 -ilirn, PA IDCb md PA INPP' mV 559 548 517 476 491
0.19 0.32 0.59 1.49 2.53
2.2 1.8 1.7 1.7 1.5
552 531 499 484 510
0.10 0.19 0.38 0.97 1.82
2.6 2.5 2.5 2.5 2.4
(m)= 0.813 mg/s for controlled t d = 0.5 s. All potentials were measured and are reported against an Ag/(O.Ol M AgN03,0.1 M TEAP in AN) electrode. The electrode was connected with the cell by a 0.1 M TEAP/AN bridge. Potentials for polarographic data were corrected for iR drop according to the spherical symmetry calculations. Voltammetric curves are displayed with the current scale increasing upward and with the potential scale increasing to the right. Therefore all anodic waves appear above the zero-current line and grow from left to the right. Analyte solutions were deareated with argon purified by passing it through copper catalyst (BASF R3-11, Chemical Dynamics Corp.) and Drierite (W. A. Hammond Co.) columns. All experiments were carried out at 25.0 rt 0.2 "C.
RESULTS AND DISCUSSION DC, NP, and RP Polarography of the First Anodic Process of Bromide and Iodide Ions. DC, NPP, and RPP experiments were performed in solutions ranging from 0.01 to 1.5 mM in concentration of halide. Drop times in the DC mode were 0.25,0.5,1.0, and 2.0 s. A 0.5-9 delay time (td)and pulse times (t,) in the range of 1-100 ms were used in NPP or RPP. Since there was no influence of the initial potential (Em)on the position and shape of the NPP curves (when EIN was sufficiently separated from the wave), we used EIN = -0.9 V or -1.1 V in NPP experiments with bromide or iodide ions, respectively. In RPP, EIN = -0.3 V was used. The most characteristic results of the DC and NPP experiments (first wave) in bromide and iodide solutions are presented in Tables I1 and 111. Concentration Dependence. Figure 1 illustrates changes in the DC polarographic response of iodide with concentration of iodide ions. In the lower concentration range the more negative wave shifts to more positive potentials with increasing concentrations of iodide. For the overall reversible reaction Hg
+ pX-
-
HgX,(P-2)-
+ 2e-
(5)
the Nernst equation predicts for p = 3 a -59 mV shift of E l j 2 per decade increase in [X-1. The values of E l j 2in Tables I1 and 111show an opposite trend for DC and NPP waves. But in Figure 2 the dependence of Ell, on log [X-] has three regions. The linear part for low concentrations and short pulse times has a slope of 19 or 22 mV per log [X-] for iodide or bromide, respectively. This range corresponds to formation of up to ca. 2% of a monolayer of HgX3-. For larger values of tC2 the plot is nonlinear. In the third part of the plot diffusing HgX3- forms on the fully covered surface of the mercury electrode ( t > t, at the given value of C) and the curves approach asymptotically the Nernst relation (curve 11 on Figure 2). This part has slope -58 to -61 mV per decade increase of [X-1, as predicted for reaction 5 using thermodynamic data and p = 3, and experimental values of E l I 2are relatively close to those calculated. The Nernstian characteristic of these waves should be observed only after surface equilibrium is reached so that the surface process .does not affect the potential-current response. The oxidation of Hg at more negative potentials under submonolayer conditions is an underpotential deposition of HgX3-
1886
ANALYTICAL CHEMISTRY, VOL. 56, NO. 11, SEPTEMBER 1984
Table 111. NPP and DC Characteristics of the First Anodic Wave of Iodide Ions in AN NPP, t, = 2 ms CI, mM
-EI/z,
0.008 0.037 0.057 0.087 0.187 0.487 0.687 0.988 1.49
828 829 828 823 814 797 788 764 715
mV
-ilirni
fiA
0.24 0.75 0.99 1.52 3.23 8.76 12.2 18.2 29.9
NPP, t, = 4 ms
INPP' 2.4 1.7 1.4 1.4 1.4 1.5 1.4 1.5 1.6
-EI/z,
NPP, t, = 10 ms
-E,
-ilirnr
mV
pA
816 809 805 802 791 762 740 700 650
0.16 0.51 0.68 1.15 2.36 5.96 8.69 13.4 21.7
I ~ p p ~
2.3 1.6 1.4 1.5 1.5 1.4 1.5 1.6 1.7
2,
DC,t d = 0.5 -EI/z,
-'
-ilirn,
md
? :
ZNppa
mV
pA
797 785 781 776 762 707 664 625 599
0.10 0.36 0.46 0.70 1.53 4.03 5.94 8.87 12.9
2.6 1.8 1.5 1.5 1.5 1.5 1.6 1.6 1.6
717 708 699 681 612 576 578 586 594
0.02 0.08 0.10 0.16 0.37 0.93 1.22 1.74 2.65
aZNpp = -ilimt,1i2/CI(mtd)2/3, pA mM-l sl/' mg-2/3, bIDC = -ilirn/CIm2/3td116, pA mM-'
IDcb 3.2 2.8 2.3 2.4 2.5 2.5 2.3 2.3 2.3
mg-2/3,
-840
-780 -720 -660
;
-600
LOG[X-I
Figure 1. DC poiarograms of iodide ions in 0.1 M TEAP/AN: (a) current after background (BKG) subtraction; (b) as (a),normalized to 1 mM concentration of I-. Iodide concentration (mM): (1) 0.05, (2) 0.07, (3) 0.10, (4) 0.20, (5) 0.50, (6) 0.70, (7) 1.0, (8) 1.5. t , = 0.5
Figure 2. Half-wave potential of the first anodic wave vs. log (halide concentratlon): (a) bromide ions, (b) iodide ions; NPP, t , = 0.5 s, t , (ms) (1) 1, (2)2,(3) 5, (4) 10, (5)20,(6) 40, (7) 100; DC, t , (S) (8) 0.5, (9) 1.0, (10) 2.0; curve 11, the reversible E,,, (formatlon of HgX,-).
The limiting current (ifim)of the first anodic wave depends linearly on the bulk concentration of X- and surprisingly is analogous to underpotential deposition of metals on foreign not affected significantly by formation of the monolayer as substrates. shown in Figure 3 and the data in Tables I1 and 111. The slope On the basis of the theory recently published for deposition of the log ilimvs. log [X-] plot varies from 1.05 to 1.13 of mercurous iodide on mercury from aqueous solutions (7) (least-squares fitting) over a wide range of time-concentration for low coverages, anodic formation of adsorbed HgX,@-2)conditions. The mean value of I N P p (Tables I1 and 111)is 1.55 should be governed by f 0.22 pA sl/z/(mM mg2I3)for bromide and 1.63 f 0.30 pA s1/2/ (mM mg213)for iodide ions, which gives diffusion coefEll2 = ficients of 2.5 X cmz/s and 2.8 X cm2/s, respectively. E0Hg2+/Hg - (RT/nF) In (22-prXppr,[X-]p-1/(~~t)1/2) The limiting current constant of the DC wave, IDc, has the (6) mean value of 2.50 f 0.06 pA/(mM s116 mg213)for bromide and 2.51 f 0.29 pA/(mM s1/6mg213)for iodide ions, which where yx is a surface activity coefficient of X- (in cm3/mol). corresponds to the diffusion coefficient 2.8 X cmz/s. This According to this relation E I I z(for p 1 2) should shift to more is in good agreement with the values of D x obtained from negative values as bulk concentration of X- increases (note double potential step chronocoulometric data. that in ref 7 the conclusion regarding the dependence of El12 All of the above shows that polarographic (DC or NP)waves on bulk [X-] derived from eq 20 was mistakenly inverse4 eq of bromide and iodide ions in AN can be used analytically. 20 is correct, but eq 19 contains minor errors). On the conAlthough the wave shifts with concentration, its limiting trary, E I I zshifts to more positive values with increasing bulk current is fully controlled by diffusion of halide ions and is concentration of halide. However Eljz changes with log (pulse proportional to halide concentration. time) as predicted by eq 6 as discussed below. Thus in our A much more complex picture of the first halide process case factors which cause lower coverage, i.e., shorter times and was obtained in RPP experiments. Figure 4 shows the effect lower bulk concentrations of X-, yield more negative values of concentration of iodide ions on RPP curves (the RPP beof El/%This agrees with the intuitive prediction that the havior of bromide was substantially the same). A t -0.3 V oxidation occurs more readily for lower coverages. S.
ANALYTICAL CHEMISTRY, VOL. 56, NO. 11, SEPTEMBER 1984
1887
Er V
Figure 5. NPP curves of 1.0 mM Br- in 0.1 M TEAPIAN: td = 0.5 s; t , (ms) = (1) 1, (2) 2, (3) 5, (4) 10, (5)20,(6) 40, (7) 100. -2 -,$\8
I
I
I
I
-4.2
-3.3
-3.6
-3.9
I
-4.5
I
I
4-13 -2.1
LOG (X-I Figure 3. Logarithmic dependence of limiting current &A) cf the first anodic wave on halide concentration (M): (a) bromide ions, (b) iodide ions; NP, t , = 0.5 s, t , (ms) (1) 2, (2) 4, (3) 1 0 DC, t , = 0.5 s (curve 4). Solid lines are linear least-squares fits.
a
I-
Z
w h a
11: 3 U
I
-1. I
I
I
I
-8.9
-0.7
-8.5
-0.9
I
E, V Figure 4. RP polarograms of iodide ions in 0.1 M TEAPIAN: iodide concentration (mM), (1) 0.05, (2) 0.10, (3) 0.20, (4) 0.50, (5) 1.0, (6) 1.5; E,, = -0.3 V; t , = 0.5 s; t , (ms) = (a) 10, (b) 2: current, after subtraction of the background, normalized to 1 ms pulse and to 1 mM concentration of I-.
HgX, is adsorbed up to a monolayer (5). Thus assuming fast and total reduction of adsorbed Hg13- a t potentials more negative than, e.g., Ellz of the DC polarographic wave, we should see no RP current in experiments with t d shorter than t,. Independently oft,, the RP curve should trace the DC curve if, of course, t d (DC) = t d + t , (RP). Instead of this, we obtained curves exhibiting a wide maximum with a peak potential of about -0.8 V (the requirement t d < t, was fulfilled for curves 1 and 2 in Figure 4). For t d > t, (curves 3 to 6) another maximum is present at more positive potentials. Here prior to pulse application we had a fully covered electrode surface and some defined con-
centration profile of diffusing Hg13- in the vicinity of the electrode. Assuming again a fast surface reduction, we should obtain RP waves close in heights and positions to NP waves. The experimental RP curves are close to predictions only at potentials more negative than -1.0 v (e.g., when t d < t,, the current drops to zero). The observed maxima could indicate slow (on the time scale of milliseconds) electron transfer in the surface reduction of HgI,. Moreover, it seems that slow kinetics of surface reduction strongly affects the rate of reduction of diffusing HgIC ions. The peak shape of the curves arises from the dependence on electrode potential of the rate of reduction (5, 6). The same maxima were recorded in the experiments involving mixtures of Hg(I1) and halide ions (no faradaic reaction a t -0.3 V), as discussed below. Very similar, peak-shaped curves have been reported also for ethylenediaminetetraacetate (8). Peak-shaped curves in pulse voltammetry are commonly associated with reactant or product adsorption (reversible). Theoretical treatments, however, have not predicted such very large maxima (9, IO). Thus the above observations must be related to the slow kinetics of reduction of HgX, from the surface. More detailed study of this phenomenon was beyond the scope of this work. Pulse Width Dependence. Figure 5 illustrates the effect of variation of pulse width (t,) on the NPP responses for oxidation of mercury in the presence of 1.0 mM Br-. Similar effects were observed for bromide and iodide ions over a wide range of concentrations. The limiting current of the first wave was strictly proportional to tC1l2over the range of concentrations of halide ions investigated (0.01-1.5 mM), and to the surface area of the electrode (DME). The latter was checked by varying the delay time before pulse application. Plots of log ih vs. log t, gave straight lines with slope about -0.5 which shows the reaction is diffusion controlled. The most striking feature of Figure 5 and the data in Tables I1 and I11 is the shift of Ell2to more negative values as the pulse width is decreased. This was not, however, the case for measurements done under full monolayer conditions, where E,,Z was basically independent of t,. According to the theory for diffusion controlled adsorption a t low coverage (cf. eq 6) the value of Ell2 should shift to more negative potentials for shorter t, (ca. 15 mV per decade of t,). For very low coverages (a few percent) in the case of bromide or iodide, the plot of Ell2 vs. log t, was linear with a slope of ca. 42 mV. By use of the concept of partial electron transfer in which a nonintegral number of electrons (y) is considered in a surface reaction, the n in eq 6 is replaced by y (7). This treatment for the above slope yields y = -0.70 (y is negative for oxidations)
ANALYTICAL CHEMISTRY, VOL. 56, NO. 11, SEPTEMBER 1984
1888
I
1 -3.d
d
9.e-
e
p. 3 0
4.0 -1.1
0
I
I
-8.5
I
-8.7
I
-0.5
-8.5
E.
-8.1
Ell
8.
v
Figure 7. DC polarographlc curves obtained in solutions containing
(1) 0, mercuric and iodide ions: C,-= 0.2 mM; t , = 2.0 s; CHBz+/C,(2)0.25, (3) 0.33, (4) 0.50. Current normallzed to 1 mM I-.
-11.1
-1.0
-8.S
-8.0
-8.7
-8.8
-8.S
-8.4-8.5
I
E? V
Figure 6. RPP curves obtained in solution of bromide ions. , C (mM): (a) 0.05, (b) 1.0. All other parameters are the same as those In Figure 5. Current was normalized to 1 ms pulse and 1 mM concentration of Br-.
which seems to be a reasonable value. With this value of y and the slope of the El,2 vs. log [X-] plot (19-22 mV), for very low coverages Ell2depends on log [X-I1l4instead of log [X-I2; cf. eq 6. The effect of pulse duration on the RPP curve will be discussed using bromide as an example, keeping in mind the concentration dependence discussed above. Figure 6 illustrates how the pulse time influences the shape of the RP polarograms at two concentrations of bromide ions. Here again we have to consider whether during the electrolysis at EIN, a full monolayer of HgBr,- was formed; i.e., what is the relation between t d and t,. A 0.05 mM bromide solution (Figure 6a requires t, = 8.2 s, assuming constant surface of the electrode, while t, for a 1.0 mM solution (Figure 6b) is only 0.02 s. Thus in the latter solution pulses were applied when the electrode surface was fully covered and a relatively large amount of nonadsorbed HgBr, was produced. Therefore two overlapping peaks in Figure 6b can be attributed to a slow reduction of both diffusing and adsorbed HgBr,-. For the same time conditions only a single peak is present in Figure 6a, which corresponds to a slow process associated with reduction of the adsorbed species. (In order to exclude supposition that a growing mercury drop (DME) might have some effect on surface equilibration, particularly before the pulse was applied, and therefore might influence the shape of RP curves, we tested a few solutions of halides ions using a static mercury drop (SMDE). The results showed that such an effect is negligible.) The RP current at potentials where there is no adsorbed HgX3- (e.g., -1.1 V), agrees with predictions. It is partically equal to zero in solution (a) and close to the NP limiting current in solution (b). Moreover, for larger concentrations
and longer pulse times the RPP curve becomes closer to what is predicted for a fast reduction process (cf. curve 7 of Figure 6a and curves 6 and 7 of Figure 6b. (Due to the fact that the RP current in Figure 6 was normalized to unit pulse time (ms) and bromide concentration (mM) without subtracting the background current, the residual current and DC component of the RP current are somewhat distorted. Figure 4 is free of these distortions because the background currents were subtracted prior to normalization.) Influence of Mercury(I1) Ions. If Hg(I1) ions are added to a solution of halide ions in AN, three main complexes are formed (HgXPb2)-,where p = 2,3, or 4). Using the stability constants determined by Coetzee et al. (3), we calculated complex distribution curves for bromide and iodide complexes of mercury(I1) and determined the ranges of concentration over which the HgX; complex dominates in solution. All of our experiments with Hg(I1) were done in those ranges. Figure 7 illustrates the influence of Hg(I1) ions on the DC polarographic response to iodide ions. The shape of the curve and values of Ellzremain practically unchanged for various CHg/CI ratios while currents become more cathodic as CHg increases. The complex distribution in bulk solution was aI= 24.3,0.54, and O.OO%, a2= 0.04, 1.82, and loo%, a3 = 97.0, 98.1, and O.OO%, and a4 = 2.97,0.06, and O.OO%, respectively, in solutions 2,3, and 4 of Figure 7 (q= lOO[I-]/CI- and a,, = 100[HgIp@-2)-]/C~,,,). Thus in solution 2 we had 24.3% of total iodide in the free form while almost all mercury existed as Hg13-. As a result the first wave is in part cathodic (reduction of HgI,) and in part anodic (oxidation of mercury due to the presence of free iodide ions). In solution 4 all mercury and iodide was in the form of Hg12. Therefore, in that solution the more positive wave corresponds to 3Hg12
+ 2e-
-
2Hg1,-
-
+ Hg
(7)
while on the plateau of the more negative wave HgIz + 2e-
Hg
+ 21-
(8)
The concentration of Hg12in that solution was smaller than [HgI2leq.Thus Hg12 did not undergo chemical reaction with mercury (reaction 4). The same set of solutions was tested by NP polarography. When EN was -1.1 V and more positive potential pulses were applied, the experiment was a traditional NPP experiment, but in solutions containing Hg(I1) reduction occurred at the initial potential, so it was an RPP experiment. When mercuric complexes were reduced at EIN,the presence of free ligand could be seen on positively scanned polarograms. On the other
ANALYTICAL CHEMISTRY, VOL. 56, NO. 11, SEPTEMBER 1984
c
3
0
-1.1
4.9
4.7
4.5 E9
-8.3
-0.1
0.1
8.3
V
Flgure 8. Pulse polarographic curves obtained in SOlutiOnS Containing mercuric and M M e ions. Concentration parameters are the same as those given in Figure 7: t , = 0.5 s; t , = 5 ms; EIN(V) (a) -1.10, (b) -0.30.
hand, when Em was -0.3 V and the scan was in the negative potential direction, we were performing an NPP experiment in the case of Hg13- ions in solution (RPP in the case of iodide ions alone). The NPP(a) and RPP(b) curves shown in Figure 8 are substantially identical (within each set) in the shape and position on the potential scale and differ only in the current at the initial potential. The latter was, as expected, due to differing DC components of the measured current arising from changes in the ratio CHf(II)/Cp Another test of the influence of Hg(I1) ions employed constant potential coulometry of halide ions (in fact, a titration of halide ions with coulometrically generated mercuric ions), which is discussed elsewhere (6). All of the above clearly manifests the identity of the product of the anodic wave and the reactant of the cathodic wave, which is the HgX3- ion. It also shows reversibility of the process on the time scale of DC polarography. DC, NP, and RP Polarography of the More Positive Process of Chloride, Bromide, and Iodide Ions. In the publication discussing the polarographic behavior of chloride ions in AN (41,we mentioned that a second, more positive wave was observed in solutions of halide ions. In the case of chloride and bromide that wave exhibited a maximum and some irregularities on the limiting plateau. Although that wave presented no interest from the analytical point of view, we decided to put more light on its complicated mechanism, in order to have a more complete picture of the electrochemical behavior of halide ions in AN. The most characteristic observations collected in our study of the second wave of halide are as follows: (1)The wave appears only when t d (in DC polarography) or t, (in NPP) is longer than t,. Thus the monolayer of HgX3- is necessary for the process giving rise to the wave to take place (Figures 1 and 5 ) . (2) A maximum and distortion on the plateau of the wave are observed only with chloride or bromide ions. The iodide wave is well-shaped in every solution tested (up to 1.5 mM
1889
I-). Since similar distortions of polarographic waves are known in the literature to be associated with deposition of insoluble product on the surface of the electrode, the above observation could indicate formation of solid mercurous halides (chloride and bromide). Assuming electrochemical generation of HgzXz, that effect should be also seen in the case of iodide because the solubility of Hg212in AN is even lower than that of Hg2C12 or Hg2Brz(3). On the other hand, if HgX2 is formed in the electrochemical step and a follow-up reaction with mercury occurs (reaction 4, it becomes explainable why the iodide wave is free of distortion. Since [HgBr& or [HgC121eqis about 1 order of magnitude lower than [HgL& (see above) it was possible to exceed this value in solutions of chloride or broide under conditions of the polarographic experiment, but not to exceed it in solutions of iodide. (3) The maximum, which on some polarograms has a very sharp form, shifts to more negative potentials with increasing drop time in DC or pulse time in NP polarography (Figure 5) as well as with increasing concentration of chloride or bromide ions. (4) The position on the potential scale and the shape of the maximum (DC polarography in chloride solutions) are not affected significantly by the nature or concentration of the supporting electrolyte (TEAP, TBAP, and LiC104 in concentrations ranging from 0.05 to 0.25 M were tested). (5) Raising temperature caused the maximum to shift by about 16 mV per 10 "C. At the same time the limiting current of the wave (measured at the roughly shaped plateau) increased by about 2% per 1 OC. The latter indicates that, although the wave is deformed by formation of insoluble Hg2ClZ,its limiting current is controlled by diffusion of chloride ions. Another factor supporting this statement is that the limiting current of the DC wave increased proportionally with td1/6 and concentration of chloride ions. The wave recorded in iodide solutions was proportional to the iodide concentration (Figure 1). Obviously, that was observed only if t d (DC) or t, (NPP) was larger than t,. (6) Addition of mercuric ions to halide solutions did not change the shape or position of the wave. Only, as expected, the curve was shifted on the current scale toward more cathodic currents. This proved to be true for iodide (Figure 7) as well as for bromide and chloride ions. (7) Cyclic voltammograms obtained by using a HMDE electrode in solutions containing chloride or bromide ions showed a cathodic stripping peak due to reduction of Hg2X2 deposited on the electrode surface. The peak was not present if the scan was reversed at the foot of the wave. Mechanism. To summarize this study of the electrochemical behavior of halide ions on mercury electrodes in acetonitrile, the following general scheme is proposed. The process occurringat more negative potentials (the first wave) is Hg
+ 3X- 'slow HgX3- (ads. u p t o a monolayer) + 2e-
then Hg(m1r)
+ 3X- e HgXC + 2eslow
where Hg(m1r) represents mercury surface covered by a monolayer of HgX3-. The process occurring at more positive potentials (the second wave) corresponds, dependent on the timebulk concentration and potential-surface concentration conditions, to reaction Hg(m1r) or 2Hg(mlr)
-
+ 2X-
+ 2X-
HgX2
+ 2e+
Hg2X2(s) 2e-
or in the case of trihalogenomercurate ions in the hulk
1890
Hg(m1r)
+ 2HgX3-
or
4Hg(mlr) + 2HgXy
-
-
Anal. Ch8m. 1084, 56, 1890-1898
3HgXz
+ 2e-
3Hg,Xz(s)
+ 2e-
If solid mercurous halide is formed (chloride, bromide) structural changes occur on the electrode surface. Formation of aggregates of halide crystals, probably involving the HgX3monolayer, opens up the mercury surface, which alters significantly the electrochemical conditions. Consequently, the polarographic curve exhibits maxima, and irregularities of current are observed on the limiting plateau of the wave. Independently of whether mercuric or mercurous halide is formed in the electrochemicalstep, equilibrium 4 is in effect. The first wave is controlled by diffusion of halide ions and its limiting plateau is well shaped; thus it is suitable for analytical purposes. It can be used for determination of halide ions (individually or as the sum) or, indirectly, for determination of chlorinated organic compounds which can be chemically or electrochemically reduced with release of free halide ions (11). It may also be used in mechanistic studies of electrochemical reduction of organic halides in which cleavage of the halogen-carbon bonds occurs (12). Registry No. HgI,, 19964-11-5;HgBrB, 21388-05-6;HgCl;, 14988-07-9;Hg, 7439-97-6;C1-, 16887-00-6;Br-, 24959-67-9;I-, 20461-54-5.
LITERATURE CITED (1) Coetzee, J. F.; McGuire, D. K.; Hedric, J. L. J . Phys. Chem. 1963, 6 7 , 1814-1820. (2) Cram, D. J. "Fundamentalsof Carbanion Chemistry"; Academic Press: New York, 1965; p 12. (3) Coetzee, J. F.; Campion, J. J.; Liberman, D. R. Anal. Chem. 1973, 45, 343-347. (4) Wojclechowski, Marek; Osteryoung, Janet Anal. Chem. 1982, 5 4 , 1713-1 719 (5) Wojciechowskl, Marek; O'Dea, J. J.; Osteryoung, Janet J. Phys. Chem. 1983, 8 7 , 4725-4730. (6) Wojciechowski, Marek; Osteryoung, Janet J . Phys. Chem., in press. (7) Brumleve, T. R.; Osteryoung, R. A.; Osteryoung, Janet Anal. Chem. 1983, 5 5 , 698-704. (8) Stojek, 2.: Osteryoung, Janet J . Nectroanal. Chem. 1961, 127, 67-74. (9) Flanagan, J. 8.; Takahashi, K.; Anson, F. C. J . Nectroanal. Chem. 1977, 8 5 , 257-266. (IO) Van Leeuwen, H. P.; Siuyters-Rehbach, M.; Holub, K. J . Electroanal. Chem. 1962, 135, 13-24. (1 1) Wojciechowski, Marek; Osteryoung, Janet, SUNY at Buffalo, unpubllshed work, 1983. (12) Wojciechowski,Marek; Osteryoung, Janet Presented at the 186th Natlonai Meeting of the American Chemical Society, Washington, DC, Aug 1983.
RECEIVED for review December 30,1983. Accepted April 17, 1984. This work was supported by the National Science Foundation under Grant No. CHE7917543 and CHE8305748. It was presented in part at the 12th Northeast Regional Meeting of the American Chemical Society, Burlington, VT, June 1982.
Voltammetric Evaluation of the Effective Acidities (pK,') for Brmnsted Acids in Aprotic Solvents William C. Barrette, Jr., H. W. Johnson, Jr., and Donald T. Sawyer* Department of Chemistry, University of California, Riverside, California 92521
Llnear sweep voltammetry at a platinum electrode has been utlllzed to determlne the effectlve acldlty (pK,' values, wlth K,' the thermodynamlc (proton-actlvity) dlssoclatlon constant relative to that for (H,O)CIO, In water) for a group of protlc substrates [(H,O)CIO,, (pyrH)CIO,, 2,4-( NO,),PhOH, (NH,)CIO,, (Et,NH)CIO4, benzolc acld, phenol, p-ethoxyphenol, and H20] In acetonltrlle (MeCN), dlmethylformamlde (DMF), dlmethyl sulfoxide (Me,SO), pyrldlne (pyr), and water. I n MeCN the pK,' values range from -8.8 for (H,O)CIO, to 30.4 for H,O. The correspondlng values In Me,SO are 2.6 and 36.7, respectively. The resultant solvent-Independent pK,' scale provldes a measure of the dmerentlal proton-actlvlty In these medla and of thelr avallablllty to catalyze proton-lnduced reactlons. I n these solvents, the rate constants ( k M ) for the mutarotation of 2,3,4,6-tetra-O-methyI-~-glucoseand for the decomposltlon of dlphenyldlazomethane are proportional to the effectlve acldlty (proton avallablllty) tor Br~nsted aclds (log k,, = apK,'). The relative thermodynamic dlssbdation constant for a protlc substrate (HA) In an aprotlc solvent, KlHr(W),Is equal to Its classlcal dlssaclatlon constant (K,HiW) muitlplled by the transfer actlvlty coefflclent for the solvent [7H+(w/H20pwlth a reference state of unity for water]. On the basis of the pK,' values for (H,O)CIO, in the varlous solvents and with pKa[(HSopo,lassumed to be 0.000 In each, the estimated values for log YH+(~/H,O) are as follows: MeCN, +8.8; H,O, 0.0; DMF, -0.7; Me,SO, -2.6; pyr, -4.6. Hence, a glven acid Is about 11 orders of magnitude more acidic in MeCN than in Me,SO.
The transfer of protons from Bransted acids to bases is one of the most important and general reactions in chemistry. Consequently, considerable effort has been expended to measure proton-transfer equilibria and thermodynamics in solution and in the gas phase (1-8). The dissociation constants of acids (K,) in aqueous and semiaqueous media can be determined by a variety of means (9-17). For nonaqueous solvents the most common methods are spectrophotometry (18-24), potentiometry (25-30), and conductivity. These methods suffer from a number of drawbacks (31,32),but the principal limitation is that pK, values are only a measure of the degree of dissociation for an acid. Hence, within a given solvent they provide relative proton activities for a series of Brernsted acids. However, the pK, values for a given acid in different solvents do not provide a valid measure of relative proton activity. Popovych (33,34) presents an elegant discussion of this dilemma. An early study (35)discusses the polarographic reduction of Bransted acids (HA) to hydrogen and conjugate bases (A-) in acetonitrile at a dropping mercury electrode. The half-wave reduction potentials, Eq2, become increasingly negative as the acidity of the B r ~ n s t e dacid decreases, and the acidity order is consistent with that in other solvents. That the reactivity of a Bransted acid with an electron at a platinum surface is related to its acidity is in accord with Usanovich's general acid-base theory (36-38). The goal of this investigation has been the development of a linear-sweep voltammetric system (at a platinum electrode) to provide a semiquantitative measure of the relative proton
0003-2700/84/0356-1890$01.50/00 1984 American Chemical Society
