Underpotential Deposition of Mercury on Au(111): Electrochemical

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Langmuir 1997, 13, 4446-4453

Underpotential Deposition of Mercury on Au(111): Electrochemical Studies and Comparison with Structural Investigations E. Herrero and H. D. Abrun˜a* Department of Chemistry, Baker Laboratory, Cornell University, Ithaca, New York 14853-1301 Received February 3, 1997. In Final Form: May 1, 1997X Electrochemical studies of the UPD of mercury on Au(111) electrodes have been carried out in sulfuric acid and in perchloric acid alone as well as in the presence of sulfate and chloride anions and in acetate/ acetic acid buffer with emphasis on the processes taking place in the presence of (bi)sulfate anions. These voltammetric studies have been compared to recent in-situ AFM, STM, and surface X-ray scattering studies in an effort to correlate voltammetric features with surface structures and their transitions. The different processes taking place during mercury UPD on A(111) in sulfuric acid can be summarized as follows: (i) At potentials higher than the first mercury UPD peak, there exists an ordered (bi)sulfate adlayer. The onset of mercury deposition triggers an order/disorder transition that gives rise to the first set of UPD peaks; C1/A1. (ii) Mercury deposition is governed mainly by mercury-gold interactions. As monolayer deposition reaches completion a disorder/order transition takes place giving rise to the peaks C2/A2. The adlayer is likely composed by Hg2SO4 with the coadsorption of additional H3O+ cations, as an analysis of the coulometric charges appears to indicate. (iii) At +0.816 V there appears a small voltammetric feature (peak C3) which corresponds to the disappearance of the ordered coadsorbed structure. (iv) Reduction of Hg2+ to Hg2+ 2 in solution occurs at +0.54 V. (v) The final process is the completion of the mercury monolayer (peak C5) followed by the formation of a mercury-gold amalgam at +0.50 V that leads to the roughening of the electrode surface. These interpretations are consistent with both voltammetric and recent in-situ surface X-ray results as well as with previous electrochemical and STM investigations and provide a very detailed microscopic picture of the processes taking place.

1. Introduction The underpotential deposition (UPD) of metals onto foreign metal substrates has been widely studied over the last two decades. Although the UPD process is primarily controlled by the interactions between the depositing metal and the foreign substrate,1,2 other interactions, such as metal-anion and substrate-anion,3-8 and the surface structure of the electrode substrate itself4,9 can play important roles. The use of single-crystal electrodes in conjunction with in-situ surface techniques, including STM,9-12 AFM13,14 and surface X-ray based techniques,5,6,14-17 has allowed the investigation and to a certain extent the untanglement of the different interacX

Abstract published in Advance ACS Abstracts, July 15, 1997.

(1) Kolb, D. M. In Advances in Electrochemistry and Electrochemical Engineering; Gerischer, H.; Tobias, C. W., Eds.; Wiley: New York, 1978; Vol. 11, p 125. (2) Adzic, R. In Advances in Electrochemistry and Electrochemical Engineering; Gerischer, H.; Tobias, C. W., Eds.; Wiley: New York; Vol. 13, p 159. (3) Zei, M. S.; Quiao, G.; Lehmfuhl G.; Kolb, D. M. Ber. Bunsenges. Phys. Chem. 1987, 91, 349. (4) Markovic, N.; Ross, P. N. Langmuir 1993, 9, 580. (5) Yee, H. S.; Abrun˜a, H. D. Langmuir 1993, 9, 2460. (6) Go´mez, R.; Yee, H. S.; Bommarito, G. M.; Feliu, J. M.; Abrun˜a, H. D. Surf. Sci. 1995, 335, 101. (7) Moller, F. A; Magnussen, O. M.; Behm, R. J. Phys. Rev. B 1995, 51, 2484. (8) Varga, K.; Zelenay, P.; Wieckowski, A. J. Electroanal. Chem. 1992, 330, 453. (9) Matsumoto, H.; Inukai, J.; Ito, M. J. Electroanal. Chem. 1994, 379, 223. (10) Hachiya, T.; Honbo, H.; Itaya, K. J. Electroanal. Chem. 1991, 315, 275. (11) Sashikata, K.; Furuya, N.; Itaya, K. J. Electroanal. Chem. 1991, 316, 361. (12) Dietterle, M.; Will, T.; Kolb, D. M.; Surf. Sci. 1995, 342, 29. (13) Chen, C.-H.; Vesecky, S. M.; Gewirth, A. A. J. Am. Chem. Soc. 1992, 114, 451. (14) Chen, C.-H.; Kepler, K. D.; Gewirth, A. A.; Ocko, B. M.; Wang, J. J. Phys. Chem. 1993, 97, 7290. (15) Yee, H. S.; Abrun˜a, H. D. J. Phys. Chem. 1993, 97, 6278. (16) Wang, J. X.; Adzic, R. R.; Ocko, B. M. J. Phys. Chem. 1994, 98, 7182.

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tions present in UPD processes. These techniques have provided information about the surface structure of the metal adlayer (with, in some cases, the participation of anions present in the supporting electrolyte), charge of the adsorbed/deposited metal, and metal-anion interactions. However, a thorough and detailed understanding of UPD processes will only be achieved through the combination of data from different in-situ surface techniques as well as electrochemical measurements. The present study deals with the UPD of mercury onto Au(111) electrodes. The first voltammetric studies of mercury UPD on gold were carried out using polycrystalline electrodes.18-23 In those studies it was shown that adsorbed mercury was only stable in the monolayer regime. Once the first monolayer was completed, mercury adatoms would diffuse into the bulk forming an amalgam. The first attempt to determine the surface structure of mercury UPD on Au(111) electrodes and the influence of anions on it was carried out by Gewirth and co-worker using AFM.24 They found that the surface structure of mercury adlayers was greatly influenced by the supporting electrolyte anion. More recently, Itaya and co-workers studied the same system in sulfuric and perchloric acid media using STM.25 They found that in sulfuric acid mercury forms an ordered lattice which they proposed to be composed of mercury and sulfate anions. Recent X-ray (17) Toney, M. F.; Gordon, J. G.; Samant, M. G.; Borges, G. L.; Melroy, O. R.; Yee, D.; Sorensen, L. B. J. Phys. Chem. 1995, 99, 4733. (18) Schadewald, L. A.; Lindstrom, T. R.; Hussein, W.; Evenson, E. E.; Johnson, D. C. J. Electrochem. Soc. 1984, 131, 1583. (19) Shay, M.; Bruckenstein, S. Langmuir 1989, 5, 280. (20) Romeo, F. M.; Tucceri, R. I.; Posadas, D. Langmuir 1990, 6, 839. (21) Salie´, G.; Bartels, K. J. Electroanal. Chem. 1988, 245, 21. (22) Salie´, G. J. Electroanal. Chem. 1989, 259, 315. (23) Salie´, G. Bartels, K. Electrochim. Acta 1994, 39, 1057. (24) Chen, C.-H.; Gewirth, A. A. Phys. Rev. Lett. 1992, 68, 1571. (25) Inukai, J.; Sugita, S.; Itaya, K. J. Electroanal. Chem. 1996, 403, 159.

© 1997 American Chemical Society

UPD of Mercury on Au(111)

Langmuir, Vol. 13, No. 16, 1997 4447

results from our laboratory26 have confirmed this observation, establishing that in the mercury-(bi)sulfate adlayer, mercury is bonded directly to the electrode surface. We also identified two additional hexagonal structures at potentials prior to the completion of the first monolayer.27 In these structures, mercury retained part of its charge, in agreement with previous reports on polycrystalline gold electrodes.21-23 In this article, we examine the voltammetric behavior of Hg UPD on Au(111) in light of the previous studies. The influence on the voltammetric profile of the mercury and sulfate concentrations has allowed us to characterize the metal-substrate and metal-anion interactions. The use of chloride and acetate, other strongly adsorbing anions, has provided further insight into this issue. With the known structural information, we have been able to correlate the voltammetric response to the adlayer surface structure. Finally, the nature of the adsorbed species after the first UPD peak is discussed. 2. Experimental Section Electrochemical apparatus and cells have been described elsewhere.28 All potentials were measured versus a Ag/AgCl electrode in 3 M NaCl. A large area coiled gold wire was used as a counter electrode. All experiments were carried out at room temperature. A Au(111) electrode (1 cm, diameter) was grown from the melt at the Materials Science Center at Cornell University. It was cut and polished with a miscut of less than 0.3°. Prior to any electrochemical measurement, it was flame annealed and quenched with ultrapure water29 and transferred to the electrochemical cell. Solutions were prepared by using ultrapure water (18 MΩ Millipore Milli-Q water). Aqueous acid solutions were prepared from high-purity sulfuric acid and perchloric acid (Baker Ultrex). Acetate buffers were prepared from high-purity acetic acid and sodium acetate. Mercuric ion solutions were prepared by dissolution of high-purity HgO (99.999% Alfa) in the appropriate supporting electrolyte solution.

3. Results 3.1. Underpotential Deposition of Mercury in 0.1 M H2SO4. Figure 1 shows the voltammetric behavior at 1 mV‚s-1 of a Au(111) electrode in contact with a 0.10 M H2SO4 + 1.0 mM Hg2+ solution over the potential range from +0.70 to +1.00 V. As can be seen and was reported previously,25,26 there are two sets of very sharp deposition and stripping peaks with peak potential values of +0.932 and +0.919 V for deposition (C1 and C2 in Figure 1) and +0.937 and +0.930 V for stripping (A1 and A2), respectively. The deposition and stripping peaks C1 and A1 are very sensitive to the long-range order of the electrode surface and the presence of impurities in solution. Electrodes with a large number of defects did not show these peaks. When the positive potential limit is increased to +1.50 V, the Au(111) electrode surface is oxidized, yielding a voltammetric response that is virtually identical to that obtained in the absence of mercury ions in solution. This fact indicates that the aforementioned peaks correspond to the first deposition/stripping process of mercury on the Au(111) electrode and that no additional process overlaps the oxidation/reduction of the gold surface itself. There are also additional, albeit very small, deposition processes at +0.816 V (C3) and +0.752 V (C4), respectively. (26) Li, J.; Abrun˜a, H. D. J. Phys. Chem. B 1997, 101, 244. (27) Li, J.; Abrun˜a, H. D. J. Phys. Chem. B 1997, 101, 2907. (28) White, J. H.; Abrun˜a, H. D. J. Electroanal. Chem. 1991, 330, 521. (29) Clavilier, J.; Armand, D.; Sun, S. G.; Petit, M. J. Electroanal. Chem. 1986, 205, 267.

Figure 1. Voltammetric profile of a Au(111) electrode in 0.10 M H2SO4 + 1 × 10-3 M Hg2+. Scan rate: 1 mV s-1.

At potentials below +0.70 V, two new voltammetric processes are evident (Figure 2). The first one at a formal potential of +0.54 V has a diffusional profile, and, as has been previously reported,25 the maximum peak current is proportional to the square root of the scan rate, indicating that it corresponds to a diffusional process. By examining the redox potentials of mercury species, this process can be assigned to:

2Hg2+ + 2e- T Hg2+ 2

E° ) +0.698 V vs Ag/AgCl (1)

The last set of peaks, at +0.517 V (C5) for the reduction and +0.520 V (A5) for the oxidation, induce a modification of the electrode surface, as indicated by the change in the voltammetric profile in the region between +0.90 and +1.00 V after a potential excursion to +0.50 V (Figure 2A, solid line). The charge underneath these peaks is 50 µC‚cm-2. Previous investigations have failed to observe (or make reference to) the existence of A525 since it is extremely sensitive to the potential scanning limits. If the upper potential limit is set at +0.45 V, A5 is absent (Figure 2A, dotted line) and the modification of the voltammetric profile at potentials higher than +0.70 V is even more pronounced. The same kind of modification can be achieved slowly by cycling the electrode potential continuously between +0.50 and +1.05 V. As Figure 2B shows, peaks C5/A5 and C1/A1 disappear upon cycling and the previously sharp peaks C2/A2 broaden significantly. The modification of the voltammetric profile observed is irreversible, suggesting that it is due to a modification of the surface properties of the electrode itself. The peak potentials of all the aforementioned process are dependent on the Hg2+ concentration. Table 1 summarizes the peak potential values for the different process over the range of Hg2+ concentration from 10-4 to 10-2 M. As would be anticipated, all peak potentials shift toward more positive values with increasing mercury concentration. It can also be noticed that the potential difference between A1 and A2 and C1 and C2 diminishes with increasing mercury concentration, and at 10-2 M, A1 and A2 merge into a singular peak. At high Hg2+ concentrations (above 5 × 10-3 M), the large currents associated with the 2Hg2+ + 2e- T Hg2+ 2 reaction overwhelm the response due to C5/A5, making an accurate determination of their peak potentials impossible.

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Herrero and Abrun˜ a

Figure 3. Voltammetric profile of a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M Hg2+. Scan rate: 1 mV s-1.

Figure 2. Modification of the voltammetric profile of a Au(111) electrode in 0.10 M H2SO4 + 1 × 10-3 M Hg2+ after a potential excursion to +0.50 V. (A) First scan (full line) and comparison with the stable voltammetric profile over the range from +0.70 to +1.05 V (dashed line). The dotted line represents the voltammetric behavior observed between +0.45 and +0.55 V when the lower limit is set at +0.45 V. (B) Evolution of the voltammetric profile upon cycling. Scan rate: 5 mV s-1. Table 1. Effect of the Mercury Concentration on the Peak Potentials (vs Ag/AgCl) for the Different Process in the UPD of Mercury on Au(111) in 0.10 M H2SO4a [Hg2+]/M

EC1

EA1

EC2

EA2

1 × 10-4 5 × 10-4 1 × 10-3 5 × 10-3 1 × 10-2

+0.905 +0.919 +0.932 +0.953 +0.958

+0.919 +0.926 +0.937 +0.958

+0.882 +0.901 +0.919 +0.943 +0.948

+0.908 +0.918 +0.930 +0.956 +0.962

a

EC3

EC5

+0.763 +0.479 +0.790 +0.506 +0.816 +0.517 +0.857 +0.875

The peak nomenclature is shown in Figures 1 and 2.

3.2. Influence of (Bi)Sulfate Anions on Mercury UPD Processes. It is now well established that UPD processes are very sensitive to the anions present in solution,3-7 often giving rise to additional voltammetric peaks and shifts in peak potentials. In some instances, metal adatoms form mixed adlayers with anions, as in the case of Cu UPD on Au(111) electrodes in sulfuric acid solutions.30,31 From recent X-ray scattering results, we have shown the formation of a bilayer coadsorption (30) Shi, Z.; Lipkowski, J. J. Electroanal. Chem. 1994, 364, 303.

structure of Hg2SO4 at potentials between +0.81 and +0.88 V.26 In order to study the influence of (bi)sulfate on the UPD processes of Hg2+, experiments were carried out in 0.10 M HClO4 + 1 × 10-3 M Hg+2 as well as in 0.10 M HClO4 + 1 × 10-3 M Hg+2 + x M Na2SO4, where x was in the range 1 × 10-4 to 1 × 10-2. Figure 3 shows the voltammetric response for the UPD of mercury on Au(111) electrodes in perchloric acid medium. As can be seen, the voltammetric profile in perchloric acid presents only one pair of broad peaks in the region between +0.70 and +1.05 V, at potentials which are, within experimental error, the same as for C2/A2 in sulfuric acid. For this reason we refer to these peaks as C2′/A2′. These peaks are symmetrical and have a quasiGaussian shape, indicating that the adsorption follows a Langmuir-type isotherm. In fact, the width at half-height for the peaks is 43 mV, a value which is very close to the theoretical value for adsorbates following a Langmuir isotherm (90.6/n mV),32 assuming that the number of electrons exchanged (n) is 2. The addition of sulfate to the solution causes the appearance of peaks C1/A1 and C2/A2. Figure 4 shows a clear example of the changes that take place in the voltammetric profile when sulfate is added to the solution. The changes are more evident in the stripping process than in the deposition process. In the deposition process, the shape of the voltammogram is very similar to that in perchloric acid with the appearance of peak C1 at the onset of the mercury deposition and peak C2 at the final stages. The stripping process shows a very sharp peak A2, with a potential and current that are almost identical to those found in 0.10 M H2SO4. The evolution of peaks C2/A2 and C1/A1 with the sulfate concentration follows two different patterns. The peak potentials for C2/A2 do not change significantly with the sulfate concentration. Although sulfate in solution is necessary for the appearance of these peaks, this result indicates that these peaks are not governed by interactions of sulfate with the mercury adlayer or the gold surface. Instead, they seem to be dominated by mercury-gold interactions. On the other hand, peaks C1/A1 exhibit a strong dependence on the sulfate concentration (Table 2). (31) Toney, M. F.; Howard, J. N.; Richer, J.; Borges, G. L.; Gordon, J. G.; Melroy, O. R.; Yee, D.; Sorensen, L. B. Phys. Rev. Lett. 1995, 75, 4472. (32) Laviron, E. J. Electroanal. Chem. 1979, 101, 19.

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Langmuir, Vol. 13, No. 16, 1997 4449

Figure 4. Voltammetric profile of a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M Na2SO4 + 1 × 10-3 M Hg2+ (full line) and comparison with the voltammetric profile of a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M Hg2+ (dashed line). Scan rate: 1 mV s-1.

Figure 5. Voltammetric profile of a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M NaCl + 1 × 10-3 M Hg2+. Scan rate 1 mV s-1.

Table 2. Effect of the Sulfate Concentration on the Peak Potentials (vs Ag/AgCl) for the Different Process in the UPD of 10-3 M Hg2+ on Au(111) in 0.10 M HClO4a [HSO4-]/M 0 1 × 10-4 1 × 10-3 1 × 10-2 1 × 10-1

EC1

EA1

EC2

EA2

+0.918b

+0.923b

+0.917b +0.957 +0.959 +0.904 +0.948 +0.953 +0.921 +0.932 +0.937 +0.919

EC3

EC5

+0.517 +0.921b +0.515 +0.920 +0.848 +0.521 +0.926 +0.844 +0.519 +0.930 +0.816 +0.517

a The peak nomenclature is shown in Figures 1-3. b Potentials for peaks C2′/A2′.

As the sulfate concentration increases, these peaks shift toward more negative potentials. As will be discussed later, this is a clear indication that these peaks are governed mainly by the (bi)sulfate-gold surface interaction. When the lower potential limit for the voltammetric scan is adjusted to +0.50 V in perchloric acid media containing no sulfate, a voltammetric peak is observed at the same potential as that for C5. The peak potential for this process remained constant over the entire range of sulfate concentrations used in this work, indicating that this process is independent of the anions present in solution. 3.3. Influence of Chloride on Mercury UPD. Chloride is a more strongly adsorbing anion than (bi)sulfate and the voltammetry of mercury UPD in its presence can provide additional insights to the processes involved. Figure 5 shows the voltammetric profile for a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M NaCl + 1 × 10-3 M Hg2+. The voltammetric response obtained in this case is very similar to that observed in 0.10 M H2SO4 (Figure 1). The potentials for the C1/A1 and C2/A2 processes are almost identical to those found in the presence of (bi)sulfate. This indicates that both anions, (bi)sulfate and chloride, play comparable roles in the mercury UPD processes. Unfortunately, higher concentrations of chloride anion could not be employed, since the onset of chlorine evolution overlaps with the first stages of the UPD process of mercury, interfering with them. The similarity in the behavior of mercury UPD in sulfate and chloride solutions differs form observations of copper UPD on Au(111), where the cyclic voltammetry is

Figure 6. Voltammetric profile of a Au(111) electrode in 0.10 M HClO4 + 1 × 10-3 M NaCl + 1 × 10-3 M Hg2+ after a potential excursion to +0.45 V. Scan rate: 5 mV s-1.

very dependent on the anion in solution.33 The different behavior has been explained in terms of the dominant interactions in the adlayer, i.e., the copper UPD process in the presence of chloride is governed by chloridechloride and copper-chloride interactions, whereas in the presence of sulfate, the dominant interaction is the copper-surface.34 This fact suggests that in the case of mercury UPD in both (bi)sulfate and chloride containing electrolytes, the mercury-gold interaction is the dominant one. The main difference in the UPD of mercury on Au(111) in the presence of chloride relative to sulfate media is found at potentials below +0.60 V (Figure 6). In this case, peaks C5/A5 are replaced by a deposition/dissolution process controlled by the diffusion of species to the electrode surface. Holding the electrode potential at the lower limit (+0.50 V) or sweeping the potential toward negative values increases the charge corresponding to the dissolution (stripping) process. However, this did not alter (33) Shi, Z.; Wu, S.; Lipkowski, J. Electrochim. Acta 1995, 40, 9. (34) Matsumoto, H.; Inukai, J.; Ito, M. J. Electroanal. Chem. 1994, 397, 223.

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Figure 7. Voltammetric profile of a Au(111) electrode in 0.10 M CH3COOH + 0.10 M CH3COONa + 1 × 10-3 M Hg2+. Scan rate: 1 mV s-1.

the voltammetric response of the electrode in the region between +0.70 and +1.05 V, suggesting that, in this case, there is no modification of the electrode surface (as was the case in the presence of (bi)sulfate). Taking into account that this voltammetric process overlaps with the oxidation/reduction wave ascribed to the reaction 2Hg2+ + 2e- T Hg2+ 2 , and the low solubility of Hg2Cl2, we ascribe the surface process to the deposition/ dissolution of Hg2Cl2 according to the reaction:

2Hg2+ + 2Cl- + 2e- T Hg2Cl2 E° ) +0.495 V vs Ag/AgCl (2) 3.4. Influence of Acetate on Mercury UPD. The voltammetric profile for mercury UPD in 0.10 M CH3COOH + 0.10 M CH3COONa is shown in Figure 7. Mercury UPD in this solution only gives a pair of irreversible peaks. Even at very low scan rates, the potential for the deposition process is 50 mV more negative than the one for dissolution. Also, deposition and dissolution take place at potentials much more negative (over 100 mV) than in the other media studied in this work. The hindering of the mercury electrodeposition can be the result of either a strong anion (acetate) adsorption on the electrode surface or the formation of mercuric complexes. Acetate and chloride have similar adsorption strengths on gold electrodes, so it is unlikely that this would give rise to a displacement of more than 100 mV in the deposition peak. On the other hand, Hg2+ and acetate anions form stable complexes in aqueous solutions (up to a 1:4 stoichiometry).35 The formation of such complexes would stabilize Hg2+ species in solution giving rise to a displacement in the deposition potential as well as to sluggish kinetics. No additional voltammetric peaks are observed until bulk deposition takes place at +0.39 V (Figure 8). In this medium, solution Hg2+ species are not reduced to Hg2+ 2 since the formation of Hg2+-acetate complexes stabilizes 2+ Hg2+ species (no Hg2 -acetate complexes have been described). Bulk mercury deposition causes irreversible changes in the voltammetric profile that are likely the result of a surface modification. (35) Martell, A. E.; Smith, R. M.; Critical stability constants; Plenum Press; New York, 1977; Vol. 3.

Herrero and Abrun˜ a

Figure 8. Voltammetric profile of a Au(111) electrode 0.10 M CH3COOH + 0.10 M CH3COONa + 1 × 10-3 M Hg2+ after a potential excursion to +0.38 V (full line) and comparison with the initial voltammetric profile (dashed line). Scan rate: 5 mV s-1.

4. Discussion 4.1. Nature of Peaks C1/A1 and C2/A2. Based on the results presented above, we believe that peaks C1/A1 are linked to the presence of adsorbing anions in the supporting electrolyte, whereas peaks C2/A2 are mainly related to interactions between mercury and the electrode surface. We shall now examine these processes in some detail, starting with C1/A1. Very sharp voltammetric peaks in adsorption/desorption processes, such as C1/A1, are normally associated with phase transitions in an adlayer. For example, such is the case for sulfate adsorbed on a clean Au(111) electrode where the sharp spike observed at around +0.80 V is related to an order/disorder transition of the sulfate adlayer. In essence, at potentials below +0.80 V, adsorbed sulfate is completely disordered, whereas at potentials above +0.80 V, it forms an ordered adlayer with a (x7×x3) surface structure, as demonstrated by STM.36,37 Our results show that at potentials above +1.00 V the electrode surface is completely free of mercury since it behaves as a clean Au(111) electrode. Therefore, the sulfate adlayer structure should be the same as that observed in pure 0.10 M H2SO4 electrolyte solution. In our recent in-situ X-ray study a (bi)sulfate layer was found at potentials higher than +0.95 V,26 in agreement with previous studies. In fact, the structure of this (bi)sulfate adlayer in the presence of mercury in solution has been recently observed by STM.25 This implies that there is an ordered (bi)sulfate adlayer on the electrode surface prior to the onset of mercury deposition on the Au(111) surface. Peaks C1/A1 show a strong dependence on the sulfate concentration, and in fact, they only appear when the sulfate concentration in the supporting electrolyte is above 1 × 10-3 M. At lower sulfate concentrations, the ordered sulfate adlayer is not formed. Thus peaks C1/A1 are linked to the presence of an ordered (bi)sulfate adlayer at positive potentials. At potentials negative to the C1/A1 process, mercury UPD begins and no ordered structure would be expected until completion of a mercury UPD layer. Thus, (36) Magnussen, O. M.; Hagebo¨ck, J.; Hotlos, J.; Behm, R. J. Faraday Discuss. Chem. Soc. 1992, 94, 329. (37) Edens, G. J.; Gao, X.; Weaver, M. J. J. Electroanal. Chem. 1994, 375, 357.

UPD of Mercury on Au(111)

we ascribe peaks C1/A1 to an order/disorder transition in the sulfate adlayer induced by the onset of mercury UPD. This is also consistent with the dependence of peaks C1/A1 on sulfate concentration. High sulfate concentrations will tend to stabilize the ordered (x7×x3) structure, hindering, to some extent, the onset of mercury deposition. The initial mercury deposition then triggers the phase transition which takes place rapidly. As a result, peaks C1/A1 should shift toward more negative potentials and become shaper with increasing sulfate concentration, as was indeed observed. Results obtained in the presence of chloride provide additional support to the assignment of peaks C1/A1 to a phase transition of an anion overlayer. From surface X-ray scattering experiments, Ocko and co-workers have demonstrated that for concentrations above 1×10-3 M, chloride forms an ordered adlayer at positive potentials (E > +0.7 V).38 One would then anticipate the presence of voltammetric peaks associated with an order/disorder transition at this chloride concentration, as was observed (Figure 5). In acetate-acetic acid medium, the absence of sharp peaks at the onset of mercury deposition would indicate the absence of an ordered anion adlayer on the Au(111) electrode surface at these potentials. Peaks C2/A2 have a completely different origin. Their peak potentials are not altered by changing the concentration of sulfate (or chloride) in the supporting electrolyte. However, the peaks become sharper with increasing sulfate concentration. Changes in the voltammetric peak shape indicate that the deposition/dissolution kinetics change in the presence of sulfate in the supporting electrolyte. This could be the case if an ordered mercury/ (bi)sulfate adlayer is formed at potentials lower than those of peaks C2/A2. The presence of an ordered (x19×x3) adlayer associated with peaks C2/A2 has been demonstrated by in-situ STM25 and grazing incidence X-ray diffraction.26 In this structure the mercury species (Hg2+ 2 ) are bonded to the surface and (bi)sulfate ions are, in turn, bonded to the mercury, forming an adlayer similar to the honeycomb structure of copper and sulfate on Au(111) electrodes.31,39 The assumption that only mercury cations are in contact with the Au(111) surface is consistent with experimental results, which showed that peaks C2/A2 are controlled by the Hg/Au interaction without significant influence by the sulfate or chloride concentrations. The formation of an ordered structure in the presence of sulfate stabilizes the adsorption of mercury on the Au(111) electrode surface, and consequently, the final stages of mercury deposition will take place faster, giving rise to a sharp peak. This way, the main UPD peaks for mercury deposition (peaks C1/A1 and C2/A2) on Au(111) in sulfuric acid and chloride media can be ascribed to three different process: (i) an order/disorder transition in the anion ((bi)sulfate or chloride) adlayer triggered by the onset of mercury deposition, (ii) the deposition of a mercury-(bi)sulfate (-chloride) adlayer, and (iii) a disorder/order transition in the mercury-(bi)sulfate (-chloride) adlayer at the final stages of the deposition. These three different processes can be readily observed in Figure 4, where a relatively low sulfate concentration was used. Moreover, deconvolution of the deposition peaks in 0.10 M H2SO4 (Figure 9) gives three Gaussian peaks, in good agreement with the proposed model. Additional work is in progress to study, in detail, the kinetics of all these processes involved in the UPD of mercury on Au(111). (38) Magnussen, O. M.; Ocko, B. M.; Adzic, R. R.; Wang, J. X. Phys. Rev. B 1995, 51, 5510. (39) Huckaby, D. A.; Blum, L. J. Electroanal. Chem. 1991, 315, 255.

Langmuir, Vol. 13, No. 16, 1997 4451

Figure 9. Deconvolution of the main deposition peak for mercury UPD on Au(111) in 0.10 M H2SO4 + 1 × 10-3 M Hg2+. (•) Experimental values; (s) fitted values.

The origin of peaks C3 and C4 is still not clear, but they are probably related to small changes in the adlayer, either in composition or in structure. In fact, they only appear when ordered structures are present on the electrode surface, i.e., in the presence of bisulfate or chloride anions in the solution. Peak C3 (which appears at +0.816 V) is likely associated with an order-disorder transition, since the (x19×x3) structure is only observed between +0.81 and +0.88 V. The lower potential limit for the stability of this structure coincides with peak C3. At this point, any additional mercury deposition would destroy the ordered structure. Further work is in progress in order to fully understand the nature of these small features. 4.2. Nature of Peaks C5/A5. Peak C5 exhibits behavior which is completely different from that of the previous ones. First, the concentration of (bi)sulfate affects neither the potential nor on the shape of the peak. Second, when the potential scan is reversed at potentials beyond C5 the cyclic voltammetry of the Au(111) electrode changes significantly, indicating a modification of its surface. Finally, in the presence of chloride anions, Hg2Cl2 is formed on the electrode surface at the same potentials and peaks C5/A5 are essentially absent. Moreover, there appear to be no changes in the surface properties of the electrode in the presence of chloride anions in the supporting electrolyte. The formation of a gold-mercury alloy layer on the electrode surface can explain the behavior observed for peak C5. Metallic mercury is known to form alloys with gold in all proportions (solid solution). When mercury is stripped from Hg/Au alloy electrodes the electrode surface, which was initially smooth, roughens with pits of 2-5 nm in depth.40 On Au(111) electrodes it has been shown that mercury forms ordered metallic adlayers at potentials below +0.68 V. Namely, there are two hexagonal structures at potentials between +0.68 and +0.63 V and between +0.63 V and +0.50 V, respectively.27 Holding the potential at values lower than +0.63 V for extended time periods results in the formation of a liquid monolayer of mercury or an Hg-Au intermetallic compound.27 Unfortunately all these processes overlap with the diffusional wave for the reduction of Hg2+ to Hg2+ 2 , and no additional voltammetric peaks can be observed for them. (40) Yang, X.-M.; Tonami, K.; Nagahara, L. A.; Hashimoto, K.; Wei, Y.; Fujishima, A. Surf. Sci. 1995, 324, L363.

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The voltammetric behavior described above suggests that an alloy is formed at potentials below peak C5. However, sharp peaks, such as C5 and A5, are generally associated with phase transitions and not with alloy formation. We thus tentatively assign peak C5 to a phase transition that induces the formation of an alloy on the electrode surface. Alloy formation will occur after the phase transition and the extent of its formation appears to depend on the value of the lower potential limit and the time spent at potentials negative of peak C5. Thus, when the scan is reversed just after peak C5, the alloy formation has taken place to a very limited extent, so that the corresponding anodic peak (A5) is still observed. On the other hand, if the lower potential limit is set at +0.45 V, there appears to be extensive alloy formation precluding the phase transition that gave rise to peak C5. Therefore, peak A5 is absent. Dissolution of the alloy will take place on the anodic sweep at positive potentials. Dissolution of the alloy disrupts the surface structure of the Au(111) electrode, giving rise to a roughening of the surface, as reflected in Figure 2. Such roughening of a Au(111) surface following alloy desorption has been previously observed with STM.25 In acetate-acetic acid solutions no sharp peaks are observed in this range, but the onset of bulk deposition also causes the disruption of the surface order. The effects of alloy formation are evident from the voltammetric behavior. In the previous section, we discussed the influence of ordered adlayers on the pair of peaks C1/A1 and C2/A2 around +0.90 V. If the electrode surface is or becomes disordered, such transitions do not occur and, therefore, no sharp peaks are expected. Moreover peaks C1/A1 should disappear in disordered surfaces, since sulfate only forms ordered adlayers on well ordered Au(111) surfaces. In addition, peaks C2/A2 should become broader and rounded since no ordered mercury/ sulfate adlayer is formed. Such behavior was observed in these studies and is consistent with the proposed mechanism. 4.3. Nature of the Species Adsorbed on the Electrode Surface over the Potential Range from +0.81 to +0.85 V. The underpotential deposition of a species is often accompanied by the adsorption/desorption of anions. In order to make an accurate determination of the adatom coverage from the voltammetric peaks, the charge associated with the adsorption/desorption of anions has to be known. However, an accurate determination of this contribution is difficult so that there is always an uncertainty in the coverage values determined from coulometric charge measurements alone. In the present case, we have made use of the adatom coverage obtained by X-ray diffraction in order to estimate the contribution of the anion adsorption/desorption process to the charge of the UPD processes. These are, in turn, used to determine which anion, sulfate or bisulfate, is involved in the process. Recent in-situ surface X-ray scattering measurements have shown that over the potential range between +0.81 and +0.85 V an ordered surface adlayer is formed which contains mercury and sulfate species.25,26 The coverage for sulfate species was found to be 0.375, whereas for mercury it was 0.75. In-plane diffraction and CTR measurements indicate that Hg2+ 2 is the mercury species in this adlayer. However, the nature of the anion adspecies is not clear. It could be either sulfate or bisulfate. Coulometric data of these processes can provide some insight into this issue. The charge under C1 and C2 (or A1 and A2) after subtraction of the apparent double layer is 105 ( 5 µC cm-2, a value that has been previously reported.25 Starting

Herrero and Abrun˜ a

from +1.05 V the processes that should take place to form the ordered structure and which must account for the net charge transferred are the following: (i) Desorption of the adsorbed sulfate species present at +1.05 V. Thermodynamic measurements have shown that the species adsorbed on a clean Au(111) electrode are sulfate anions (SO42-).41 These sulfate anions are accompanied by the adsorption of some H3O+ cations.37,41 However, it is still not clear whether the amount of adsorbed H3O+ is stoichiometric with the amount of sulfate in order to form SO42-‚H3O+, as has been proposed for Pt(111) electrodes.42 In any case, the charge necessary to remove all the sulfate species adsorbed on the electrode surface at this potential has been determined by iodine displacement measurements to be equal to 60 µC cm-2,43 lower than the theoretical value required to desorb a SO42adlayer with a 0.2 coverage41 (88 µC‚cm-2). (ii) Possible reduction of the Hg2+ cations to Hg2+ 2 or to Hg°. The charge needed to reduce Hg+2 to Hg2+ 2 to give a 0.75 coverage value of mercury adspecies is 167 µC cm-2. If the adsorbed species were Hg°, the charge required would be 333 µC cm-2. If these charges have to be added to that for process (i), the net charges would be 227 and 393 µC cm-2, respectively, much higher than the value found experimentally. Thus, we conclude that during the processes that give rise to peaks C1 and C2 there appears to be no reduction of mercury. Hence, the mercury species that participate in the formation of the mercury/sulfate ordered adlayer must be already present in solution. As mentioned in the Experimental Section, mercury solutions were prepared by dissolving HgO, yielding Hg2+ species. However, there are always Hg2+ 2 cations (at low levels) present in equilibrium with Hg2+ cations. Both species can participate in the formation of the mercury/sulfate adlayer, since both can be readily adsorbed directly from solution. From the 2:1 ratio of mercury to anions and the size limitation in fitting adsorbates to the in-plane lattice obtained by in-situ X-ray measurements,25,26 we believe that the mercury adspecies present on the electrode surface are Hg2+ 2 cations. (iii) Adsorption of the mercury and sulfate species. Having established that the adsorbed mercury species are Hg2+ 2 cations, two different sulfate species can be adsorbed: either SO42- or HSO4-.44 Since in the adsorbed structure the number of mercury ions is twice that of sulfate, the adsorbed species must be either Hg2SO4 or Hg2HSO4+. In the first case, the charge necessary to adsorb Hg2SO4, a neutral species, should be 0 µC cm-2. Adsorbing Hg2HSO4+ at a coverage of 0.375 would require a charge of 83 µC cm-2. (iv) Reorganization of the double layer structure. In any process where there is a change in the species adsorbed on the electrode surface, there can be a reorganization of the double layer that could lead to charge transfer. Namely, two different aspects have to be considered: a change in the Epzc induced by the adsorbed species and a change in double layer capacitance. This latter contribution is very difficult to evaluate and it is generally estimated and, in many cases, even neglected since it typically represents a very small contribution. Taking into account all the contributions and the experimentally measured charge, only two process could (41) Shi, Z.; Lipkowski, J.; Gamboa, M.; Zelenay, P.; Wieckowski, A. J. Electroanal. Chem. 1994, 366, 317. (42) Faguy, P. W.; Marinkovic, N. S.; Adzic, R. R. Langmuir 1996, 12, 243. (43) Herrero, E.; Feliu, J. M.; Wieckowski, A.; Clavilier, J. Surf. Sci. 1995, 325, 131. (44) The charge needed to adsorb HSO4- and SO42-‚H3O+ species are the same and therefore they cannot be distinguished from charge measurements.

UPD of Mercury on Au(111)

give rise to C1 and C2. The first one would involve the desorption of the (bi)sulfate adlayer and adsorption of a Hg2SO4 adlayer, which would require a charge transfer of 60 µC cm-2 plus the charge required to reorganize the double layer. The second would involve the desorption of the (bi)sulfate adlayer and adsorption of Hg2HSO4+, with a charge transfer of 143 µC cm-2 (60 µC cm-2 for (bi)sulfate desorption and 83 µC cm-2 for the Hg2HSO4+ adsorption) plus double layer contributions. Neither of the two possibilities is in good agreement with the measured charged, which is higher than the first and lower than the second. This would suggest either that the double layer reorganization plays an important role (which is unlikely) or that the adsorbing adlayer is neither Hg2SO4 nor Hg2HSO4+ but a mixture of the two. The adsorption of sulfate on clean Au(111) also poses a similar problem to the one found here. The experimentally measured charge for the process of sulfate adsorption lies between the theoretically calculated values for sulfate and bisulfate adsorption.37 Thermodynamic studies and FTIR experiments37,41 have shown that there is an additional adsorption of H3O+ with sulfate anions and this can account for the difference in charge. This could also be the case for the mercury-sulfate adlayers. Additional H3O+ cations could be adsorbed with the sulfate anions increasing the charge required to form the Hg2SO4 adlayer from 0 to 45 µC cm-2, which upon addition to the charge due to sulfate desorption (60 µC cm-2) would give rise to a value of 105 µC cm-2, which is very close to the experimentally measure one. Therefore the adsorbed adlayer at +0.85 V appears to be Hg2SO4 with an additional adsorption of some H3O+ cations. 5. Conclusions The different processes taking place during mercury UPD on A(111) in sulfuric acid media and their relationship to the voltammetric responses can be summarized as follows:

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(i) At potentials higher than the first mercury UPD peak, there exists an ordered (bi)sulfate adlayer. The onset of mercury deposition triggers an order/disorder transition that gives rise to peaks C1/A1. (ii) Mercury deposition is mainly governed by mercurygold interactions. As the monolayer deposition reaches completion an disorder/order transition takes place (peaks C2/A2). Previous X-ray results indicate that a (x19×x3) mercury-(bi)sulfate bilayer is form at these potentials. The adlayer is likely composed by Hg2SO4 with the coadsorption of additional H3O+ cations, as an analysis of the coulometric charges appears to indicate. (iii) At +0.816 V, the ordered coadsorbed structure disappears (peak C3). (iv) Reduction of solution Hg2+ to Hg2+ 2 occurs at +0.54 V. The presence of this diffusion controlled reduction wave masks the structural transformations that take place in the adlayer; i.e., two different hexagonal structures that have been identified at these potentials via in-situ surface X-ray diffraction experiments.27 (v) The final process is the formation of a mercury-gold amalgam (after peak C5) at +0.50 V, which leads to the roughening of the electrode surface. Mercury UPD in the presence of chloride anions in solution exhibits a behavior that is very similar to the one observed in sulfuric acid media. The only significant difference is the absence of amalgam formation and the apparent formation of a Hg2Cl2 layer. In acetate-acetic acid medium, the formation of Hg2+-acetate complexes appears to hinder the mercury UPD processes. Acknowledgment. This work was supported by the National Science Foundation and the Office of Naval Research. E.H. acknowledges support by a fellowship from the Ministry of Education and Science of Spain. Discussions with Dr. J. Li are gratefully acknowledged. LA970109T