Adsorption of Bisulfate and Sulfate Anions on a Pt(111) Electrode

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J. Phys. Chem. B 2001, 105, 2588-2595

Adsorption of Bisulfate and Sulfate Anions on a Pt(111) Electrode A. Kolics and A. Wieckowski* Department of Chemistry, UniVersity of Illinois at Urbana-Champaign, 600 South Mathews AVenue, Urbana, Illinois 61801 ReceiVed: September 28, 2000; In Final Form: January 11, 2001

A modified radioactive labeling method was used to study bisulfate and sulfate anion adsorption on a Pt(111) electrode from aqueous HClO4/H2SO4 media. The highest surface concentration of the anion is 3.2 × 1014 ions/cm2, which corresponds to a coverage (packing density) of 0.21 ( 0.01monolayer. Overall, the coverage data confirm our previous results reported in J. Electroanal. Chem. 1993, 348, 451. We found fast adsorption kinetics and a semilogarithmic anion adsorption isotherm. In the “butterfly” electrode potential range [that is, in the range where anomalous anion adsorption was found on the Pt(111) electrode (Clavilier, J. J. Electroanal. Chem. 1980, 107, 211)], the charge on the anion and, consistently, electrosorption valency vary with the electrode potential, demonstrating progress in bisulfate dissociation to sulfate, with the increase in potential. On the plateau of the surface concentration-electrode potential plot, we conclude that the adsorbate is a partially discharged (or neutralized) sulfate anion, which displays a net charge of -1.7 e. We also conclude that the species coadsorbed with sulfate are predominantly water molecules rather than hydronium cations. A brief review of the theory of anion adsorption on metal surfaces is presented and a model of sulfate-platinum bonding is proposed.

1. Introduction The significance of anion adsorption in the structure and reactivity at the metal/solution interface has been highlighted before, recently in ref 1. Apart from fundamental aspects, adsorbed anions affect the reactivity of the processes of practical importance in electrocatalysis, electrosynthesis, and deposition,1 including underpotential deposition (ref 2 and references therein), and are effective in corrosion inhibition or promotion.3 Anion-electrode interactions are also good templates to study interactions between nanoparticle alloy catalysts and terminal oxy groups of proton conducting polymers, such as sulfonate groups in Nafion,4 of key importance in fuel cell science and technology.5,6 Major effort has been exerted by this group to enhance the fundamental scientific understanding of the anion effect in electrochemical surface science and to delineate the nature of the metal-anion surface chemical bond. The present study is a continuation of this effort. Adsorption of sulfate and bisulfate (until a discrimination between bisulfate and sulfate forms is confirmed by the experiment below, bisulfate and/or sulfate will be denoted as (bi)sulfate) on a Pt(111) electrode is of specific interest to many investigators, and the (bi)sulfate surface concentration (Γ, the number of anions per square centimeter) or coverage (packing density, the number of anions per one Pt surface site) under a variety of experimental conditions should be known exactly.7-13 The (bi)sulfate concentration is quantitative information needed to support scanning tunneling microscopy (STM) data concerning the composition and structure of the Pt(111) electrode in sulfuric acid media7,12,13 and is of general importance in imaging/ spectroscopic work with the “sulfate-type” anionic adsorbates. Other than sulfate and/or bisulfate anions, halide or phosphate anions are also known to chemisorb on platinum.14-17 On the * To whom correspondence [email protected].

should

be

addressed:

e-mail

other hand, the anions that display an alternative behavior are perchlorate and fluoride,18,19 which are not contact-adsorbed; that is, they remain in the diffuse part of the electrical double layer at potentials not significantly more positive than the potential of zero charge for a given electrode material.20 (However, perchlorate can be specifically adsorbed on a platinum electrode modified by thallium depostion.21) Therefore, either perchlorate or fluoride solution provides a perfect medium from which adsorption of specifically adsorbed species can be followed. Perchloric acid was used in this study since it can be directly obtained from a manufacturer in an acceptably clean form. Since surface structure of the (bi)sulfate anion on the Pt(111) electrode is Pt(111)(x3 × x7),7,12,13 the coverage associated with this structure is either 0.2 or 0.4 monolayer (ML). One of the objectives of this work is to reinvestigate the anion coverage by the use of an improved radioactive labeling method, in view of previous controversy discussed before.8 Once the surface concentration of the anion is known, the comparison with the charge associated with the anion adsorption is the right entry for assessing the nature of the anion adsorbate or, at least, the amount of charge carried by the anion upon its adsorption. For the specific case of (bi)sulfate adsorption on a Pt(111) electrode, the adsorption is shifted away from the hydrogen adsorption (desorption) range (ref 22 and references therein). Therefore, the ratio between the charge and the amount of the adsorbate obtained by radioactive labeling is free of possible pseudocapacitive artifacts due to hydrogen coadsorption and is fundamentally meaningful in terms of electrosorption valency,23 as a fundamental qualifier of the (bi)sulfate adsorbate on the Pt(111) electrode. 2. Experimental Section 2.1. Procedure, Chemicals, and Conditions. The Pt(111) electrode 10 mm in diameter and 8 mm in height,

10.1021/jp003536f CCC: $20.00 © 2001 American Chemical Society Published on Web 03/13/2001

Adsorption of Anions on a Pt(111) Electrode oriented in the (111) direction within 0.25°, and polished down to 0.25 µm surface finish was obtained from Mateck-Bu¨ro (Germany). Before the measurements, the crystal was annealed in a hydrogen flame and cooled in a 2:1 mixture of ultrapure argon and hydrogen, saturated with water vapor.24 After being cooled to a moderate temperature, the crystal was immersed in Millipore water and transferred directly to the radiochemical cell, with a water drop isolating the surface from the ambient atmosphere. A modified radioactive labeling method25 was applied to measure the (bi)sulfate surface concentration on Pt(111). The present setup includes new provisions, which enhanced the sensitivity of the technique, as well as the reproducibility of the results. This predominantly includes implementation of a new calibration method versus that previously used in this laboratory9,26 (see below). Also, glass scintillators used in previous research were replaced by a CaF2(Eu) scintillator (Bicron). While the glass phosphor has a light efficiency of 6570% versus anthracene, the light efficiency of calcium fluoride is 140%. A Teflon spacer that guides the electrode to the scintillator15 was also redesigned to allow a more controllable positioning of the Pt(111) substrate. However, as before, all measurements were carried out in a Teflon cell equipped with a working electrode, a platinum wire counterelectrode, and a Ag/AgCl/3 M NaCl reference, separated from the cell by a porous glass-bridge system, which eliminated a possible penetration of chloride into the cell electrolyte. Millipore water (18 MΩ cm), double-distilled perchloric and sulfuric acids (GSF Chemicals) were used to prepare the solutions. 35S-Labeled H2SO4 (molar activity 40 Ci mol-1) was supplied by ICN Biomedicals. Unless otherwise stated, the measurements were performed in 0.1 M HClO4 supporting electrolyte at 23 ( 1 °C. During the adsorption measurements, electrode potentials were chosen randomly in the region of -0.10 to 1.20 V. The surface signal was taken for 50 s in a “squeezed position”,26 and then the gap was opened to 1-2 mm and the electrode was held in this lifted position for 10 s for accumulation of an adsorbate before it was reinstalled in the squeezed position for another adsorption measurement. We performed five such measurements at each electrode potential. Before moving to another adsorption potential, two voltammetric scans between -0.20 and 0.90 V were applied at a scan rate of 0.2 V s-1 to renew the surface. After the adsorption was measured for a few different electrode potentials, the electrode was pulled out to a meniscus position to check the cyclic voltammogram of the Pt(111) electrode. If the voltammetric curve displayed preliminary evidence for surface disorder, the crystal was removed from solution, annealed, and placed back into the cell for further radiochemical measurements. Squeezing efficiency26 was checked each time after the reintroduction of the crystal into the solution or after each set of five adsorption measurements at a given electrode potential. Since the reproducibility of the squeezing efficiency is a critical factor in surface concentration determinations, we also checked the reproducibility of the squeezing efficiency for each series of experiments (Figure 1). As seen from the figure, the “squeeze ratio” was highly reproducible and was equal to 90.3% ( 0.2%. This allowed us to estimate the gap thickness (the distance between the scintillator and the electrode surfaces in the squeezed position26) at 1.5 µm. 2.2. Calibration. Instead of using a theoretically deduced relationship between the measured count rates and surface concentration,26 we employed a calibration procedure with a standard; that is, a

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Figure 1. Reproducibility of the “squeezing efficiency” (cf. ref 26, p 640) obtained with the Pt(111) electrode pressed against the scintillator surface in a 1 mM solution of 35S-labeled H2SO4 in 0.1 M HClO4 at -0.15 V. [No (bi)sulfate adsorption occurs at this potential.]

known amount of labeled (bi)sulfate was introduced between the electrode and the scintillator.25 The key requirement for the validity of such a calibration is that the calibration experimental conditions must closely imitate those for determination of (bi)sulfate adsorption.25 Such conditions are met when the gap thickness is less than 2 µm, since self-absorption in the gap solution is very small, and there is only a small difference in the count rate (intensity) from the surface and from the gap radioactivity. (Under our experimental conditions, the gap intensity and the surface signal differ less than 1%.) To conform to these requirements, we used a liquid rather than a solid source (the latter being most frequently used for radiochemical calibration27,28). A calibrated syringe allowed us to reproducibly apply 20-100 nL of 35S-labeled sulfuric acid solution in 10 nL increments on the polypropylene foil that covered the scintillator. The same platinum single-crystal substrate as used for adsorption measurements was pressed against the foil and the scintillator underneath and the nanoliter droplet of the radioactive sulfuric acid, with the same pressure as in the adsorption measurements.26 Using several aliquots, we managed to control the amount of sulfuric acid in the gap in order to determine the gap intensity (using a dead-time correction). The results provide evidence for the linear relationship between count rates and the amount of radioactivity in the gap in the range of 20-80 nL; see Figure 2. This relationship yielded a conversion factor (η) between the count rates and the (bi)sulfate concentration, which is equal to (4.0 ( 0.2) × 1011 cps·mol-1 (cps denotes counts per second). The surface concentration then can be calculated from

Γ)

(Ntotal - Ngap) ηSRλ

(1)

where Ntotal and Ngap (in cps) are the total and gap intensity in the adsorption experiment [on the surface of the Pt(111) electrode], respectively; S is the geometric surface area in square centimeters; R is the roughness factor (taken as 1); and λ is the decay constant:

λ ) 2-t/T1/2

(2)

where t is the time taken between the calibration and adsorption

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Figure 2. Calibration plot constructed in order to obtain a relationship between the amount of surface radioactivity and radiation intensity: count rates versus the radioactive volume in the electrode-scintillator gap (see text).

Figure 4. (A) Enlarged sections of cyclic voltammograms of the Pt(111) electrode in 0.1 M HClO4 (1) and 0.1 M H2SO4 (2) solutions. Scan rate: 0.1 V·s-1. (B) Demonstration of the baseline used for the voltammetric charge calculation (see text).

Figure 3. Cyclic voltammograms of the Pt(111) electrode in 0.1 M HClO4 + x mM H2SO4: x ) 0 (s); x ) 0.1 (- · -); x ) 0.5 (- - -); x ) 1.0 (···); x ) 5(- ·· -); and x ) 9 (---). Scan rate: 0.1 V·s-1.

measurements and T1/2 is the half-life of 35S. The data obtained by use of this calibration procedure removed remaining doubts from the previous determinations of surface concentrations via the radioactive labeling method in use, which could come from uncertainties as to the exact value of the backscattering factor and radiation absorption coefficient.26 3. Results and Discussion 3.1. Voltammetry and (Bi)sulfate Coverage. We first demonstrate a family of well-known cyclic voltammetric curves for the Pt(111) electrode in mixed sulfuric acid/ perchloric acid medium with a systematically increasing concentration of sulfuric acid in 0.1 M HClO4 (Figure 3). [All measurements were obtained in a meniscus position in which only the “working” Pt(111) plane of the electrode was exposed to solution.] All features of these voltammograms are wellknown, including a reversible “butterfly” envelope,29 correctly interpreted to be due to anion adsorption.30-34 [The term

“butterfly” is frequently used to refer to a current-potential voltammetric profile in the electrode potential range where anomalous anion adsorption was proposed29 and confirmed9 on the Pt(111) electrode.] Notice the fully adequate voltammetric appearance of our data in comparison to the universal voltammetric standard in the field24 and that the voltammetric double layer charging current is essentially the same in all solutions studied (see the solid line parallel to the E axis in Figure 4A, for 0.1 M H2SO4 and for 0.1 M HClO4). A simple calculation of the integral double layer capacity (as the baseline for the butterfly in both the sulfuric acid and perchloric acid solutions, Figure 4A) yields the value of ca. 80 µF cm-2. This is a high value, showing some unusual and not yet fully understood properties of the Pt(111)/solution interface.35 Kinetic measurementssthat is, the determination of coverage vs time relationships by the use of radiolabelsswere conducted as follows. After a cyclic voltammetric (CV) curve characteristic of the Pt(111) electrode was obtained, of the type shown in Figure 3, the electrode was fully immersed in perchloric acid electrolyte containing a desired amount of sulfuric acid. Several CV sweeps at 100 mV/s were then applied, and the sweep was interrupted at an adsorption potential of choice. After a predetermined time (not shorter than 1 s), the electrode was lowered against the surface of the radiation detector (the CaF2 scintillator; see Experimental Section) and the counts due to adsorption were measured for 50 s. Under the present measuring

Adsorption of Anions on a Pt(111) Electrode

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Figure 5. Count rates as a function of time curves for (bi)sulfate adsorption on Pt(111) electrode obtained in 0.1 M HClO4 + 1 mM H2SO4, at 0.60 V (1) and 0.80 V (2).

geometry conditions, no “escape” of the trapped radioactivity could be considered since the diffusion from or into a micrometer-size gap is very slow. After the 50 s counting time, the gap was opened up for yet another predetermined period, and the electrode was squeezed against the detector for the next 50 s counting measurement. Since some extra adsorption may occur from the solution film of ca. 1 µm thick, the kinetic data thus obtained can be considered approximate. However, all measurements are accurate enough to yield the trends discussed below. Cyclic voltammograms were also taken after the end of the adsorption experiment. The data were considered as acceptable if the voltammogram before and after the adsorption measurement were identical; other data were rejected as invalid. The kinetic data indicate that (bi)sulfate adsorption is fast (Figure 5) and that the adsorption is followed by a slow desorption, with its rate strongly dependent on adsorption potential. These observations are in line with some former results.22,36 [By use of impedance spectra, the rate of (bi)sulfate adsorption on platinum was found to be 6605 nmol cm-2 s-1 at 0.5 M H2SO4.36] A slow rearrangement of the anionic adlattice at a constant potential previously identified by STM13 can be connected to the slow desorption shown by radioactive labeling (Figure 5). The radiochemical data also demonstrate that the desorption process is relatively fast at electrode potentials located in the platinum oxidation range. The latter result can be understood by assuming that (bi)sulfate adsorption follows an aging of the oxide film produced on platinum at such positive potentials. In view of the desorption kinetics, the data for adsorption time of 2 s were chosen for further treatment and interpretation to avoid having been affected by desorption. The electrode potential dependence of (bi)sulfate adsorption for several values of the bulk concentration of sulfuric acid in 0.1 M perchloric acid (Figure 6A) indicates that the increasing branch of the Γ-E plot coincides with the development of the butterfly feature21 on a representative voltammogram. We found no evidence as to the occurrence of a two-inflection-point curve reported by chronocoulometry10 and clearly seen for the Pt(110) electrode by using either the radioactive labeling method9,37 or Fourier transform infrared spectroscopy (FTIR).38 A welldeveloped Γ-E plateau appears at 1 mM (Figure 6A), which is less apparent at other concentrations. (For the concentration

Figure 6. (A) Adsorption of (bi)sulfate on the Pt(111) electrode as a function of the electrode potential in 0.1 M HClO4 + x mM H2SO4: x ) 0.1 (1); x ) 0.5 (b); x ) 1.0 (2); and x ) 5.0 (9). (B) Demonstration of the experimental error in the measurements of (bi)sulfate adsorption on Pt(111) as a function of the electrode potential in 0.1 M HClO4 + 1 mM H2SO4. Error bars represent standard deviations of the mean.

of 1 mM, the analysis of the experimental error is shown in Figure 6B. Data for other concentrations were basically similar, with decreasing experimental accuracy at the higher bulk concentrations used for this study, as discussed in ref 26.) The decrease in (bi)sulfate surface concentration with the electrode potential, beyond the adsorption maximum, coincides with the threshold of platinum surface oxidation in the mixed acidic perchlorate/bisulfate medium, which depends on the bulk concentration of sulfuric acid in HClO4 (Figures 3 and 6A). The surface concentration of the anion increases with the increase in the bulk concentration of (bi)sulfate anion, giving rise to a semilogarithmic isotherm, as shown for three electrode potential values in Figure 7. (A deviation from the linear Γ vs log cH2SO4 plot is found at the most positive potential shown in this figure, at low cH2SO4.) Notably, the highest surface concentration obtained in this paper, corresponding to 82 µC·cm-2 of charge (see below), is 3.2 × 1014 ions/cm2, which corresponds to 0.21 ( 0.01 ML coverage, confirming our previous results.9

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Figure 7. Concentration dependence of (bi)sulfate adsorption on the Pt(111) electrode in 0.1 M HClO4 solutions of sulfuric acid. Bulk concentration of H2SO4 varied from 0.1 to 9 mM. E ) 0.50 V (0), E ) 0.60 V (4), and E ) 0.80 V (O).

3.2. Anion Adsorption Charge and Electrosorption Valency. We have chosen voltammetric determination of the charge specific to anion adsorption, since such measurements offer a reliable and easily accessible coulometric baseline (see the results of Figure 4 and compare to the data in Figure 3). An alternative is to refer the anion adsorption charge to the zerocurrent line, which would then contain an uncorrected double layer contribution to the total voltammetric charge measured. The method of the incremental voltammetric charge measurements is illustrated in Figure 8. The anion adsorption charge obtained under voltammetric conditions is clearly separated from the hydrogen stripping charge,22 and the voltammograms demonstrate a suppressionsby the anion adsorptionsof platinum surface oxidation, which otherwise is characteristic of perchlorate solutions. Figure 9 shows a typical plot of surface concentration (Γ) vs voltammetric charge (qads) obtained in this study. The electrode potential associated with this plot was 0.6 V. The charge qads was obtained by integration of the voltammetric currentpotential plot from a lower potential boundary (e.g., from 0.12 V in Figure 8) to the upper limit, 0.6 V located in the “past butterfly” range (above the capacitive baseline, see above). The range of sulfuric acid (predominantly bisulfate) concentration in 0.1 M HClO4 varied from 0.1 to 9 mM. The plot is linear, the straight line passes through the origin (not shown), and when the abscissa (upper axis) is expressed in 10-14 Ne/e cm-2, the slope is electrosorption valency (γ):23

λ)-

NA ∂q F ∂Γ E

( )

Figure 8. Demonstration of the evaluation of a voltammetric charge involved in (bi)sulfate adsorption (qads) on the Pt(111) electrode in 0.1 M HClO4 + x mM H2SO4. Upper plot: a linear potential sweep waveform. Lower plot: a blow-up of a typical “butterfly” on a voltammogram for the same electrode (the positive-going run).

(3)

equal to -1.7 e/ion (where Ne ) NA/F). In eq 3, NA is the Avogadro number, F is the Faraday constant, q is the charge density related to the anion adsorption (coulombs per square centimeter), and E is the electrode potential. The electrosorption valency principle is restated in Figure 10A, showing a complete analysis of the experimental error, with the γ data scattered between -1.6 and -1.7. A question was then posed whether electrosorption valency

Figure 9. Surface concentration (left axis) and coverage (right axis) of adsorbed (bi)sulfate versus voltammetric charge (qads, see Figure 8) on the Pt(111) electrode in 0.1 M HClO4. Bulk concentration of H2SO4 varied from 0.1 to 9 mM, at E ) 0.60 V. Upper abscissa is expressed as 10-14 Ne/e cm-2 (where Ne ) NA/F) for easy determination of electrosorption valency from the plot (see text).

was electrode potential-dependent. Therefore, a partial integration of the voltammetric i-E curves was performed (e.g., the shaded area in Figure 8), producing several electrosorption

Adsorption of Anions on a Pt(111) Electrode

Figure 10. Experimental error in measurements of electrosorption valency (γ) for adsorbed (bi)sulfate ions on Pt(111). (A) 0.1 M HClO4 + 0.1-9 mM H2SO4, at 0.60 V. (B) 0.1 M HClO4 + 1.0 mM H2SO4. Error bars represent standard deviations of the mean.

valency values (from eq 3). Figure 10B shows a relationship between γ values obtained in this way, with the upper integration boundary expressed as the electrode potential, in volts, not time (Figure 8), for one representative concentration of H2SO4, 1 mM (consistently, in 0.1 M HClO4). Since we found that γ vs E curves were basically similar for other bulk concentrations of sulfuric acid, only this particular experiment at 1 mM is demonstrated. Apparently, for all bulk concentrations studied in this project, the data indicate that the electrosorption valency increases with the electrode potential in the butterfly range. This implies that there is a continuous transition in the charge and related composition of the adsorbed anion layer in the butterfly range (also, see below). This increasing trend is followed by the electrode potential-independent section of the electrosorption valency-electrode potential plot (Figure 10B) on the Γ-E plateau (Figure 5), which begins just positive of the familiar spike in the butterfly feature (Figure 3). The γ value is then equal to -1.7, as mentioned above (see also ref 13), and is shown in Figure 10A. Notice that the stability in γ coincides with appearance of an ordered anion adlattice, stable for another 0.2-0.3 V, as indicated by STM.12,13 Also, notice that the high error bar in the low electrode potential region (Figure 10B) is

J. Phys. Chem. B, Vol. 105, No. 13, 2001 2593 due to a reduced accuracy in surface count rate measurements at low surface concentrations (Figure 6B). 15 Using the electrosorption valency value near -1.7, as obtained above, we propose that the anion adsorbed on Pt(111) at positive potentials, past the butterfly spike, is a partially discharged sulfate anion. We also propose, on the basis of STM results obtained by other investigators,12,13 that the surface sulfate anion is coadsorbed with water molecules. These issues, as well as that concerning the anionic composition in the butterfly range, will systematically be discussed below. 3.3. Electrosorption Valency versus the Nature of the Adsorbate. A physical technique suitable for identifying the nature of ionic adsorbates on platinum in sulfuric acid medium (or in mixed sulfuric-perchloric acid solutions) is infrared spectroscopy.20,21,39-46 Apparently, there is a consensus among the FTIR groups that the anionic adsorbate on the Pt(111) electrode is not uniform but contains two components. Some data point at hydronium cation as one of the components, but there is also a controversy with respect to whether bisulfate or sulfate anion is the second component. Iwasita and co-workers41 reported that the S-containing anion on Pt(111) was sulfate (but presented no analysis of a coadsorbate), and Faguy et al.42 also concluded that sulfate was adsorbed but was coadsorbed with hydronium cation. On the other hand, Ito and co-workers believe that the S-species is bisulfate, coadsorbed with hydronium cation.43,44 Notably, if the last model prevailed in the full butterfly electrode potential range, and the ion pair HSO4--H3O+ was adsorbed, the voltammetric current in this range would be of a purely capacitive origin. We believe this is not feasible. On the other hand, the above-proposed option that the structure is actually sulfate and an associated surface-water network could be reconciled with the STM results, as concluded in refs 12 and 13, and has recently found support in the spectroscopic analysis of sulfate adsorption on Au(111),45 where similar butterfly features correlated to sulfate adsorption46 were found. A theoretical evaluation of such (bi)sulfate/water structures on Pt(111) is currently being made by other investigators.47 Our data indicate that the charge on the adsorbate in the butterfly range is electrode potential-dependent. Notice that the butterfly features develop in the range more positive, or close to the potential of zero charge (pzc); the latter in perchloric acid is close to 0.35 V versus RHE.48,49 (See also ref 35, where ca. 100 mV more positive pzc was found.) Below 0.25 V, the prevailing component is either bisulfate (γ ) -1), or a sulfatehydronium ionic pair; the latter difference cannot be resolved by the experimental approach chosen in this work. At the lowest potentials investigated, surface species may even correspond to a more neutral form than bisulfate. While a high experimental error at low surface concentrations does not allow us to reach a more definite conclusion as to the exact nature of such a form (H2SO4?), the trend shown in Figure 10B was found at all bulk concentrations of sulfuric acid in 0.1 M HClO4 studied in this work. As the positive electrode potential is increased, and the butterfly feature develops, the adsorbate increases its negative charge and moves closer to a surface sulfate form at the end of the butterfly (and in the plateau range). The surface dissociation that we believe is observed may correspond to a deprotonation of the weak HSO4- acid on platinum upon its chemisorption. This would then be in agreement with FTIR measurements concerning a variety of weak acids adsorbed at the metal/solution intrface.50 However, the final charge on the anion is -1.7 rather than the -2.0 value expected for a solution sulfate analogue

2594 J. Phys. Chem. B, Vol. 105, No. 13, 2001 (obtained from bisulfate dissociation). The missing -0.3 charge could be due to electron donation from the sulfate to platinum, within the donation-back-bonding adsorption mechanism postulated by Sellers et al.,1 or due to the above-mentioned coadsorption of hydronium cation with the sulfate anion, within the Pt(111) (x3 × x7) structure (see below), or both. Nevertheless, the adsorbate displaying a charge of -1.7 e is much closer to sulfate than to bisulfate, with the coadsorbed species being predominantly water molecules, rather than hydronium ions. 4. Conclusions: Theory and Experiment Quantum mechanical calculations for anion binding to metal surfaces (metal cluster with 50-100 atoms) have recently been reported by Olivera et al.1 The binding energies were obtained on different crystal surfaces and for different adsorbate coordination, including the effect of electrode potential, to determine the most stable adsorption geometries on each crystal face. The key finding was that the nature of anion adsorption was not strongly dependent on the electrode material since the contribution of d electrons to the binding was weak. Therefore, the theory explains experimental observations showing similar surface structure and potential range of adsorption (vs the potential of zero charge) of (bi)sulfate on such diverse metals as platinum, gold, silver, and copper. Both chemical and electrostatic processes contribute to the binding energy of the anions; Pauli repulsion, metal polarization, and adsorbate polarization are the electrostatic interactions, whereas chargetransfer processes are associated with chemical bonding. The latter involves both charge donation to the surface (the predominant binding factor) and back-donation to empty levels in the adsorbate. More specifically, data indicated a backbonding due to interactions between the occupied orbitals of a metal (the case study was silver) and the unoccupied orbitals of free sulfate.1 The increased electron density around sulfur atoms had been found before by core electron energy loss spectroscopy for (bi)sulfate adsorption on Pt(111)11 and Au(111) electrodes.51 On the other hand, the charge donated to the metal originates from the nonbonding electron density localized around the oxygen atoms. Notably, it was shown that the difference in binding energies of bisulfate and perchlorate arises from the different degree of back-bonding, but still it was concluded that that the main contribution to the bonding was electrostatic. While the anion geometry is distorted in the presence of electric fields, it was found close to that of typical salts. This corresponds exactly to the experimental observations.11,51,52 For the isoelectronic series H2SO4, HSO4-, and SO42-, a very good linearity was found between binding energy and square of adsorbate charge.1 That is, the binding energy of sulfuric acid is considerably lower than those of both bisulfate and sulfate. The sulfate ion loses ca. 0.4 electron toward the metal and most of that charge comes from the noncoordinated oxygen atoms. Anion-metal interactions in the zero coverage limit were only considered, and it was noticed that high surface coverage corresponding to the Pt(111) (x3 × x7) structure “is difficult to understand, taking into account the Coulombic repulsions among the sulfate anions”. Therefore, it was concluded that the hydration layer around the anions played an important stabilizing role. Finally, it was concluded that positive electric fields increased the binding energy of the anions and caused a higher charge transfer to the surface. Our data add additional strong credence to the theory. The same adsorbate, sulfate, was found on both Au(111)46 and Pt(111) in this work and displayed a similar potential depen-

Kolics and Wieckowski dence of adsorption with respect to, and beyond the potential of, zero charge of these two metals. As mentioned above, the contribution of back-bonding and the similarity between the electronic structure of (bi)sulfate adsorbate and that of sulfate salts has already been shown experimentally.11,51,52 We found very little evidence for sulfuric acid adsorption, which, according to the theory, should display the smallest binding energy, although some adsorption of a considerably neutralized Scontaining sulfate form at the beginning of the butterfly range cannot be excluded (Figures 3 and 6A). While we found that the sulfate ion loses ca. 0.3 electron, rather than the 0.4 electron predicted, 0.1 electron may be back-donated to Pt(111). If this is so, a quantitative agreement between the theory and experiment could be claimed, excluding a still feasible option for hydronium cation chemisorption. Another token of agreement is the strong evidence for water coadsorption with sulfate (STM and this work), substantiating the thought about “sulfate hydration” in the theory. However, in the Γ-E plateau range (Figure 10B), no increasing degree of charge transfer to the metal was found with the increase in the electrode potential. This will have to be further examined. The theory does not give an explanation of the total charge flow in the butterfly region, because the charge transfer between the adsorbing anions and the metal due to the metal/anion bonding (see above and ref 1) is not sufficient to account for ca. 80 µC·cm-2 charge found in this range at 0.20 ML coverage. In view of the absence of any redox involved in sulfate adsorption (the sulfate adsorbate is electronically similar to a sulfate anion in a sulfate salt),11,51,52 the charge measured in the butterfly region is concluded to be due to attachment of the negatively charged species, first bisulfate and then sulfate to platinum, and a concomitant removal of the conduction band electron from the interfacial region. The electrons (electron density) then move through the external circuit to the counterelectrode and contribute to the anion adsorption current measured. The resulting model of sulfate-platinum bonding is in essence (not in terms of the specific system investigated) similar to that postulated by Lipkowski and co-workers53 for underpotential deposition of copper on Au(111)54 based upon thermodynamic53 and XANES55,56 considerations. Namely, despite a 2e surface redox, Cu+ (a valency equal to 1, not 0) is present as an adsorbate, and the positive charge is compensated for by the negative charge residing on the metal. For the two surface systems, therefore (adsorbed Cu+ and SO42-), a charge remains on the adsorbate and is +1 and -1.7 for copper and sulfate, respectively, and the surface chemical bond has a dipolar character, or sulfate is binding platinum via a polar chemical bond. Acknowledgment. This work was supported by the National Science Foundation under grant No. CHE-9985837, and the Department of Energy Grant DEFG02-96ER45439. References and Notes (1) Olivera, P. P.; Patrito, M.; Sellers, H. In Interfacial Electrochemistry: Experimental, Theory and Applications; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999; p 63. (2) Zhang, J.; Sung, Y.-E.; Rikvold, P. A.; Wieckowski, A. J. Chem. Phys. 1996, 104, 5699. (3) Kolics, A.; Polkinghorne, J. C.; Thomas, A. E.; Wieckowski, A. Chem. Mater. 1998, 10, 812. (4) Herrero, E.; Franaszczuk, K.; Wieckowski, A. J. Phys. Chem. 1994, 98, 5074. (5) Cleghorn, S. J. C.; Ren, X.; Springer, T. E.; Wilson, M. S.; Zawodzinski, C.; Zawodzinski, T. A.; Gottesfeld, S. Int. J. Hydrogen Energy 1997, 22, 1137.

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