Electrochemical and Spectroscopic Studies of Hydroxide Adsorption at

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J. Phys. Chem. B 1999, 103, 682-691

Electrochemical and Spectroscopic Studies of Hydroxide Adsorption at the Au(111) Electrode Aicheng Chen and Jacek Lipkowski* Guelph-Waterloo Center for Graduate Study in Chemistry, Guelph Campus, UniVersity of Guelph, Guelph, Ontario N1G 2W1, Canada ReceiVed: September 8, 1998; In Final Form: NoVember 30, 1998

The adsorption of hydroxide ions at a Au(111) single-crystal electrode has been investigated quantitatively using chronocoulometry and subtractively normalized interfacial Fourier transform infrared spectroscopy (SNIFTIRS). By thermodynamic analysis of the charge density data, the Gibbs excess, Gibbs energy of adsorption, and number of electrons flowing to the interface per one adsorbed hydroxide ion at a constant electrode potential (electrosorption valency) were determined. The electrosorption data indicate that the adsorption of OH- has a three-state character. The adsorbed OH- forms quite a polar surface bond at a negatively charged surface, while the polarity of the surface bond is significantly decreased at positive charge densities. Oxide formation begins at higher charge densities. Infrared spectroscopy shows that oxide formation takes place when the surface concentration of hydroxide ions exceeds one-third of a monolayer. The integrated infrared intensity of the O-H stretching band correlates very well with the Gibbs excess of hydroxide determined by chronocoulometry.

1. Introduction Electrochemical oxidation of noble metal electrodes has been extensively studied.1-6 It has been postulated that this reaction involves a consecutive transfer of two electrons. The adsorption of hydroxide ion constitutes the first step of this reaction. The specifically adsorbed OH- is then oxidized to form MOH, and MOH is oxidized further at more positive potentials to produce MO. Nguyen et al.7 applied electroreflectance spectroscopy to investigate the oxidation of gold electrodes. Their electroreflection spectra gave the first direct evidence of the chemisorption of hydroxide ions. However, they were unable to identify the oxidation state of adsorbed OH and described it as an “incipient oxidation of gold”. Potential-dependent surface-oxygen vibrational bands have been measured by Weaver et al. using surfaceenhanced Raman spectroscopy (SERS) in the potential range where gold undergoes oxidation.2 The submonolayer formation of OH on a gold electrode surface has also been postulated on the basis of cyclic voltammetric studies in aqueous acid solution by Angerstein-Kozlowska et al.3,4 Recently, Faguy et al. applied real-time polarization modulation infrared spectroscopy to study water and hydroxide adsorption at a copper electrode surface.8 Despite all these efforts, very little is known about the energetics of hydroxide adsorption at the gold electrode surface. On the other hand, it has been found that the presence of chemisorbed OH- anions plays a key role in determining the catalytic properties of Au(hkl) surfaces for O2 reduction9 and electrooxidation of hydrogen4 and small organic molecules.10,11 For these reasons, it is quite important to acquire a more complete description of the adsorption behavior of OH- ions. In the present work, chronocoulometry and a thermodynamic analysis of the charge density data, described in our previous work,12,13 have been used to investigate the energetics of OHadsorption at the Au(111) electrode surface. These electrochemical studies have been complemented by subtractively normalized interfacial Fourier transform infrared (SNIFTIRS)

measurements14-16 used to acquire additional molecular-level information concerning the properties of OH- at the gold surface. The objectives of this study are (1) to determine the energetics of hydroxide adsorption at the gold(111) surface, (2) to compare the adsorption of OH- to the adsorption of other anions, (3) to investigate the dependence of the integrated IR intensity on the surface concentration to cross-check the surface concentration data, and (4) to determine the electrode potential of the onset of oxide formation. 2. Experimental Section The electrochemical experimental procedures and instrumentation were described in refs 12, 13, and 17. The working electrode (WE) was a Au(111) single-crystal rod that was grown, cut, and polished in our laboratory. The counter electrode (CE) was a gold coil. Before each experiment both the WE and CE were cleaned by flame annealing and then quenched with Millipore water. The reference electrode was a saturated calomel electrode (SCE) connected to the investigated solution through a salt bridge. The electrochemical experiments were performed using a PAR model 173 potentiostat controlled by a computer, with all data acquired via a plug-in acquisition board (RC Electronics model IS-160). The experimental strategy involved characterization of the surface by recording cyclic voltammetry and differential capacity and determination of the electrode charge density from chronocoulometric experiments. The charge density data were subsequently used to determine the relative Gibbs excesses, and the Gibbs energies of adsorption. The data treatment procedures have been described in refs 12, 13, and 17. Water was purified in a tandem of the Milli-Q and Milli-Q plus UV systems (18.3 MΩ cm). The supporting electrolyte was 0.1 M KClO4 purified according to the procedure described in our previous paper.12 Suprapure potassium hydroxide monohydrate (MERCK) was used without further treatment. An Orion Research model EA 920 Expandable Ion Analyzer was used to

10.1021/jp9836372 CCC: $18.00 © 1999 American Chemical Society Published on Web 01/09/1999

Hydroxide Adsorption at the Au(111) Electrode

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measure the pH of the investigated solutions. All solutions were deaerated with argon before measurements, and argon was passed over the top of the solution during the experiment. All measurements were conducted at room temperature (20 ( 2 °C). The electrochemical measurements were conducted in 0.1 M KClO4 + x M KOH solutions. The concentrations of KOH shown in this paper were as follows: 3.16 × 10-5, 5.62 × 10-5, 1.00 × 10-4, 1.78 × 10-4, 3.16 × 10-4, 5.62 × 10-4, 1.00 × 10-3, 1.78 × 10-3, 3.16 × 10-3, 5.62 × 10-3, 1.00 × 10-2 mol dm-3. The SNIFTIRS technique was employed to record the IR spectra. The SNIFTIRS experimental procedures and instrumentation were described in refs 16 and 17. A syringe-type IR cell with a CaF2 prism beveled at 60° was used in this study. The in situ FTIR experiments were carried out on a Nicolet 20SX/C FTIR apparatus equipped with a MCT-B detector cooled by liquid nitrogen. The sample compartment of the FTIR apparatus was purged throughout the experiment using CO2 and H2O free air provided by the Puregas Heatless Dryer. The electrode potential was controlled by a PAR 173 potentiostat. The spectra were recorded using a multiple potential step (MPS) procedure in which the electrode potential was stepped periodically between the reference potential E1 and the sample potential E2. The acquisition was delayed for 30 s after each potential change to allow the interface to reach thermodynamic equilibrium at that potential. The change of the electrode potential was synchronized with the acquisition of the interferograms by connecting the external trigger port of the PAR 173 potentiostat to the communication port of the DX 486 computer of the FTIR instrument. The interferograms were added, Fourier transformed, and used to calculate a relative change of the electrode reflectivity, which is defined as

∆R/R ) [R(E2) - R(E1)]/R(E1)

(1)

where R(E1) and R(E2) are the electrode reflectivity at potentials E1 and E2, respectively. The spectra were recorded with a resolution of 4 cm-1. The IR incidence beam was normal to the prism. The SNIFTIRS experiments were carried out in 0.1 M KClO4 + 5.6 × 10-3 M KOH and 0.1 M KClO4 + 5 × 10-3 M KCl solutions. 3. Electrochemical Studies of OH- Adsorption 3.1. Cyclic Voltammetry and Charge Density Plots. Figure 1a shows cyclic voltammograms (CVs) recorded in the presence of 0.1 M KClO4 solution as the supporting electrolyte (dotted line) and with the addition of KOH, 1.00 × 10-4 M, 1.00 × 10-3 M (dashed lines), and 1.00 × 10-2 M (solid line). The shapes of these curves are consistent with the voltammograms found in the literature.18,19 A broad hump can be seen in the CV, which is followed by a sharp peak with a shoulder present on the negative side. When the bulk concentration of KOH is increased, the height of the sharp peak increases and the position of the peak moves in the direction of negative potentials. This sharp (irreversible) peak may correspond to lifting of the (1 × 23) reconstruction of the Au(111) surface.19-21 At potentials more positive to the sharp peak, the current density increases significantly. This feature indicates that oxide may be formed on the surface at these higher potentials. Therefore, hydroxide adsorption studies were restricted to potentials lower than 0.35 V/SCE. In Figure 1a, the current density curves for the various hydroxide concentrations merge with the CV curve for the supporting electrolyte at potentials lower than -0.4 V. This behavior suggests that, in this region of potentials, OH- anions

Figure 1. (a) CV curves recorded at a Au(111) electrode in 0.1 M KClO4 (dotted line) and in the supporting electrolyte upon the addition of KOH, 1.00 × 10-4 M, 1.00 × 10-3 M (dashed lines), and 1.00 × 10-2 M (solid line) at a sweep rate of 10 mV/s. (b) Charge density versus electrode potential curves for the Au(111) electrode in 0.1 M KClO4 (dotted line) and 0.1 M KClO4 + x M KOH solutions.

are completely desorbed from the Au(111) electrode surface. We selected E ) -0.75 V as the final potential in the potential step experiments, performed to determine the charge density at the electrode surface. Figure 1b shows the absolute charge density-potential plots that were determined for the Au(111) electrode in the presence and absence of KOH. These curves correspond to a potassium hydroxide concentration ranging from 3.16 × 10-5 to 1.00 × 10-2 M. The σM versus E curves for solutions containing KOH merge with the curve for the supporting electrolyte when the potential is lower than -0.4 V. This behavior is consistent with the shape of CVs shown in Figure 1a. It shows that the OHanions are desorbed from the Au(111) electrode surface and justifies our choice of E ) -0.75 V as the final potential for the potential step experiments. Adsorption of OH- apparently causes a positive charge to flow to the metal side of the interface. The charge density curves for hydroxide solutions display multiple regions. The first, observed at E < 0.1 V, is characterized by a very small slope. The second corresponds to the fast rising segment of the curve, which correlates well with the position of the large peak on the CV curve. At more positive potentials, the charge density increases steeply, apparently due to the formation of oxide. In electrochemical experiments, we can independently measure three related physical quantities, current density, differential capacity, and charge density. It is interesting to compare these three independent measurements. In Figure 2, the differential capacity determined from the ac impedance measurement is compared to the capacity calculated from the positive-going section of the CV and to the capacity calculated by differentiation of the charge density data determined from chronocoulometric experiments. These curves can be divided into three sections corresponding to E < 0.1 V (section I), 0.1 V < E
0.25 V (section III). In section I, all three independent measurements give the same capacity values. In section II (the region of the tall peak), the directly measured, single-frequency differential capacity is much lower than capacities determined by the two other techniques. The differences between capacities determined from CV and by differentiation of charge densities are less pronounced. Section III corresponds to the region of oxide formation. The capacity calculated from the charge density is much larger than that determined from the CV in this section. The inset to Figure 2 presents the comparison of charge densities measured by chronocoulometry with the charge densities determined by integration of the differential capacity or cyclic voltammetry curves. The inset shows that charge densities determined by integration of the differential capacity curves are much lower than those directly measured by chronocoulometry. The charge densities calculated by integration of the CV are close but still somewhat smaller than the chronocoulometric data. For these reasons, cyclic voltammetry and differential capacities were used for qualitative characterization of the electrode surface and chronocoulometric experiments were carried out in order to determine quantitative data for the adsorption process. Consequently, further quantitative data analysis was based on the charge density derived from chronocoulometric experiments. 3.2. Gibbs Excess Data. In the present series of measurements, the adsorption of hydroxide took place from a solution with an excess of the supporting electrolyte. As shown in refs 13 and 22, the electrocapillary equation for a gold electrode in contact with that electrolyte is given by

-dγ ) σM dE + ΓOH-RT d ln cOH-

(2)

where γ is the interfacial tension, E is the electrode potential measured with respect to SCE, and ΓOH- is the Gibbs excess of hydroxide ions. To calculate the Gibbs excess of OH-, the charge density curves were integrated using the procedure described in refs 13 and 17 to give the relative interfacial tension (γ). The relative interfacial tension was then plotted versus ln cKOH at a constant potential E, and the resulting curves were differentiated to give ΓOH-. Independently, the Gibbs excess at constant charge was calculated by plotting the Parsons function

Figure 3. Plots of the Gibbs excess of OH- against (a) electrode potential and (b) electrode charge density for 0.1 M KClO4 + x M KOH solutions.

ξ ) σME + γ 23 and differentiating the relative ξ versus ln cKOH plots at a constant σM. Figure 3a shows the Gibbs excess versus electrode potential plots determined using E as the independent electrical variable. The Gibbs excess plots show three characteristic sections within which the coverage changes with potential in a quasilinear fashion. The first section is seen at the most negative potentials. It corresponds to low coverage and is characterized by a very small slope. The second section at intermediate coverage is steep. At higher coverages, the change of the surface concentration with potential initially slows down but later increases sharply. This behavior is different from that displayed by other anions. The surface concentration of three halides (Cl-, Br-, and I-) and sulfate attains a limiting value at the most positive potentials.17,24 This result shows that, in the case of hydroxide adsorption, oxide formation may take place at the most positive potentials. Figure 3b shows the Gibbs excess data plotted against the charge density. Three characteristic sections within which the Gibbs excess changes with charge density in a quasilinear fashion can also been seen in Figure 3b. They correspond to (1) a negatively charged surface, (2) moderate positive charge densities (10 < σM < 40 µC cm-2), and (3) high charge densities (σM > 40 µC cm-2), respectively. This behavior shows that the values of the Gibbs excess determined at constant potential are consistent with the Gibbs excess calculated using charge density as the independent electrical variable. 3.3. Gibbs Energies of Adsorption. The Gibbs energies of sulfate,13 chloride,22 bromide,25 and iodide24 adsorption at the Au(111) electrode surface have been determined using the “square root” isotherm. To determine the Gibbs energies of hydroxide adsorption, according to Parsons’s suggestion we have fitted the surface pressure data Φ of hydroxide adsorption to the equation of the “square root” isotherm:26,27

ln(kTcKOH) + ln β ) ln Φ + BΦ1/2

(3)

Hydroxide Adsorption at the Au(111) Electrode

J. Phys. Chem. B, Vol. 103, No. 4, 1999 685

Figure 4. Fit of the adsorption data to the equation of the square root isotherm (a) at constant potential, and (b) at constant charge density σM.

where β ) exp(-∆G/(kT)) is the adsorption equilibrium constant, B is a constant, and Φ is the surface pressure. When analysis was carried out at a constant potential, Φ was set to (γθ)0 - γθ), where γ is the interfacial tension. In contrast, Φ was set to (ξθ)0 - ξθ) when charge was used as the independent electrical variable. The subscripts θ and θ ) 0 denote the values of the interfacial tension and the Parson’s function in the presence and absence of KOH in the supporting electrolyte, respectively. Parts a and b of Figure 4 show plots of the square root of the surface pressure versus ln((kTcKOH)/Φ). The plots are fairly linear and when extrapolated to zero surface pressure give an intercept with the ln((kTcKOH)/Φ) axis equal to

lim ln((kTcKOH)/Φ)Φ)0 ) -ln β

(4)

from which the Gibbs energies of adsorption can be calculated. Note that in equations 3 and 4 the potassium hydroxide concentrations are multiplied by the term kT so that in the limit of low coverage the film pressure is described by Henry’s law Φ ) kTβcKOH as explained in ref 26. The Gibbs energies of adsorption, determined using this procedure, are plotted against potential and charge in parts a and b of Figures 5, respectively. The standard state is an “ideal” Γ ) 1 ion cm-2 for the adsorbed species and an “ideal” cKOH ) 1 mol dm-3 for the bulk species. We emphasize that the square root isotherm is an empirical isotherm, and although it is called “virial type isotherm”, eq 3 cannot be derived from the equation of state for the virial isotherm.27 The isotherm is not based on a well-defined physical model and hence should be viewed as a convenient procedure to linearize experimental data only. Although the fit to the square root isotherm gives linear plots, the extrapolation to zero surface pressure covers a relatively large distance. Consequently, even a small change in the slope may lead to a significant error in the intercept. The constant charge and the constant potential analyses give plots of a different slope. Therefore, it is useful to compare Gibbs energies determined at the constant E and at the constant σM to check the consistency of their values. For that purpose, the ∆G data determined from the constant charge analysis were plotted also against the electrode potential in Figure 5a together with the ∆G values determined from the constant E analysis. The charge density versus potential curve for the supporting electrolyte was used to convert the Gibbs energy data from ∆G versus σM onto a ∆G versus E scale. The ∆G values determined by the two methods differ by a factor of 2-3 kJ mol-1, and this difference shows the order of magnitude

Figure 5. (a) Plot of the Gibbs energy of adsorption versus electrode potential (O) determined from the φ1/2 versus ln((kTcKOH)/Φ) plots at constant E in Figure 4a and (b) from the Φ1/2 versus ln((kTcKOH)/Φ) plots at constant σM in Figure 4b. (b) Plot of the Gibbs energy of adsorption versus electrode charge density determined from the φ1/2 versus ln(kTcKOH)/Φ) plots at constant σM in Figure 4b.

of the systematic errors that may be included in our data analysis due to the long extrapolation procedure that was used. The Gibbs energy plots are nonlinear and may be formally decomposed into two segments of different slopes corresponding to E < 0.2 V (before oxide formation) and E > 0.2 V (at which oxide formation starts to take place). This behavior reflects a different character of OH- adsorption at the Au(111) surface. It is also useful to note that the ∆G versus E plot is concave while the ∆G versus σM curve is convex. 3.4. Model of the Inner Layer. Although all the thermodynamic quantities that we have described so far are very helpful to understand the adsorption behavior of hydroxide, they provide little information about the potential and charge distribution at the interface unless interpreted with the help of a physical model. Many models of the electrical double layer were developed to describe the structure of the metal solution interface in the presence of the specific adsorption of anions.28 Among these models, the Grahame-Parsons model of the inner part of the double layer29 has been used to describe a large number of adsorption data. Recently, this model has been applied to describe sulfate,13 chloride,22 bromide,25 and iodide24 adsorption at the Au(111) electrode surface. The model assumes a total separation of charge at the interface and assumes a specific location for the adsorbed species (inner Helmholtz plane at a distance x1 from the metal surface). To compare the adsorption of hydroxide to the adsorption of other ions, we will discuss the thermodynamic data describing hydroxide adsorption at the Au(111) electrode using the Grahame-Parsons model. We emphasize that the conclusions of this discussion will depend entirely on the validity of the model assumed and the validity of the model cannot be proved on the basis of electrochemical measurements alone. The capacity of the inner layer (Ci) may be calculated from the overall electrode capacity (C) as shown in Figure 2 determined by differentiation of the charge density curves in Figure 1b and using the theory of the diffuse layer, with the help of the formula26

(Ci)-1 ) (C)-1 - (1 - F(∂ΓOH-/∂σM))(Cd)-1

(5)

where Cd is the capacity of the diffuse layer and ∂ΓOH-/∂σM is the slope of the plots in Figure 3b. We have compared the adsorption behavior of the three halides at the Au(111) electrode

686 J. Phys. Chem. B, Vol. 103, No. 4, 1999

Chen and Lipkowski

Figure 6. Inner layer capacities determined for the Au(111) electrode in 0.1 M HClO4 + 10-3 M K2SO4 (O), 0.1 M KClO4 + 10-3 M KCl (0) and 0.1 M KClO4 + 10-3 M KOH (b) solutions.

surface using the Grahame-Parsons model in refs 17 and 24. We have found that the adsorption strength of these three halides increases from chloride to iodide. Hence, in this study, we will compare the adsorption of hydroxide to chloride and sulfate ions. Figure 6 shows the inner layer capacities plotted versus the charge density on the metal for these three anions. They display a large peak at charge densities σM > 0. The intensity of the peak increases and the position of the peak moves in the positive direction by moving from SO42- to OH-. This trend corresponds to the order of increasing adsorption strength. At a charge density greater than 40 µC cm-2, the inner capacities in the presence of sulfate and chloride adsorption decrease slowly. In contrast, the inner capacity of hydroxide adsorption increases steeply with increasing charge density in this region. This behavior may be due to oxide formation. The inner layer capacity is a function of two variables, charge on the metal and the amount of adsorbed anion. It may therefore be expressed in terms of two components, capacity of the inner layer at a constant charge (ΓC) and capacity of the inner layer at a constant amount adsorbed (σC), as described by the formula26,30

(Ci)-1 ) (σC)-1 - F(∂ΓOH-/∂σM)(ΓC)-1

(6)

The capacity ΓC may be determined from the slope of the plot of the potential drop across the inner layer, ∆φM-2 ) E - Epzc - φ2, versus the charge of adsorbed hydroxide, -FΓOH-, at a constant σM (φ2 is the outer Helmholtz plane potential). These plots are shown in Figure 7; they are quite linear. A similar feature was found for bromide25 and iodide24 adsorption at the Au(111) surface. However, we note that, in the case of sulfate13 and chloride22 adsorption at the Au(111) surface, the experimental points on the ∆φM-2 against -FΓOH- plot deviated from linearity at the lower limit of the anion concentration. The capacities at constant charge may be considered as integral capacities described by the formula: ΓC

) /(x2 - x1)

Figure 7. Plot of the potential drop across the inner layer ∆φM-2 versus the charge of adsorbed OH- (-FΓOH-) at constant charge density at the metal side of the interface as indicated in the figure. The dotted line and open points represent the potential drop corresponding to the condition of zero diffuse layer charge.

(7)

where  is the permittivity and x2 the thickness of the inner layer. The capacities at constant amount of adsorbed anion were then calculated from capacities Ci and ΓC using eq 6. Independently, ∆φM-2 was plotted versus σM at constant amount of adsorbed hydroxide ions and σC values were calculated from slopes of these plots. The capacity at constant amount adsorbed

Figure 8. Components of the inner layer capacity (a) at a constant charge and (b) at constant amount adsorbed for the Au(111) electrode in (O) SO42-, (0) Cl-, and (b) OH- solutions.

may also be considered as the integral capacity given by σC

) /x2

(8)

Parts a and b of Figure 8 show the components of the inner layer capacity at constant charge and at constant amount adsorbed plotted against the charge density on the metal for SO42-, Cl-, and OH-. For charge densities larger than 10 µC cm-2 and smaller than 50 µC cm-2, the two components of the inner layer capacity increase in the order SO42- < OH- < Cl-. This trend is not preserved for small charges, where capacities ΓC and σC display a maximum, and for large charges, where oxide formation takes place. The interpretation of ΓC curves is difficult, since they may be affected by both the change of the permittivity and the position of the inner Helmholtz plane x1. The interpretation of σC curves is easier, particularly if one assumes that the thickness of the inner layer x2 does not change

Hydroxide Adsorption at the Au(111) Electrode

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Figure 9. Plots of the surface dipole moment of adsorbed anion at the Au(111) surface for (O) SO42-; Cl- (0), and OH- (b). Inset: the electrosorption valency versus electrode charge density plot determined from the ratio of σC to ΓC.

with charge. The shape of σC displays how its permittivity changes with the charge on the metal. The changes of the permittivity may be correlated with the orientation of water dipoles at the electrode surface. As a result of the dielectric saturation, the permittivity is expected to be low at either very negative or very positive potentials and to display a maximum at small charge density where the solvent molecules reorient and where the disorientation of the water dipoles is at its maximum. The changes of the inner layer capacity at the constant amount adsorbed should correlate well with the ability of the solvent molecules to screen the surface dipole formed by adsorbed anion. The ratio of the inner layer capacities at a constant amount of adsorbed anion and at a constant charge is equal to the electrosorption valency:29

γ′ ) σC/ΓC ) (x2 - x1)/x2

(9)

Electrosorption valencies determined from the inner layer capacities with the help of eq 9 are shown in the inset to Figure 9. The magnitude of the electrosorption valency and its dependence on charge may be discussed in terms of the thickness ratio as described by eq 9. However, Parsons and coworkers31-33 and later Schmickler and Guidelli34-37 have demonstrated that the dipole moment of a dipole formed by adsorbed anion and the image charge in the metal may be calculated from the electrosorption valency with the help of the formula

µs )

ze0((1 - y′)z) σC

(10)

where e0 is the unit charge,  is the permittivity, and σC is the capacity of the inner layer at the constant amount of adsorbed anion. The advantage of this approach is that the magnitude of the surface dipole is a direct measure of the polarity of the bond formed between the adsorbed anion and the metal. The dipole moment is given by the product of the charge on the adsorbed anion and the distance from its image charge in the metal. As in any other area of physical chemistry, the knowledge of the magnitude of the dipole moment is sufficient to describe the

polarity of the electrosorption bond. There is no need to speculate about the magnitude of the charge on the adsorbed anion and the electrosorption bond length, when these quantities cannot be directly measured. The dipole moments calculated for sulfate, chloride, and hydroxide at the Au(111) surface are plotted against the charge on the metal in Figure 9. The permittivity was taken as equal to the permittivity of a vacuum (0 ) 8.85 × 10-12 C2 J-1 m-1) in these calculations. The dipole moments for adsorbed anions have a negative sign; however, in the following section we will discuss their absolute values only. Overall, at the positively charged surface, the values of the dipole moment increase in the order OH- < Cl- < SO42-, indicating that the polarity of the dipole moment increases by moving from OH- to SO42-. Apparently, the polarity of the electrosorption bond is a strong function of the charge on the metal. For adsorbed OH-, the polarity is large at the negatively charged surface but decreases significantly at positive charges. In contrast, for adsorbed SO42-, the polarity of the bond becomes quite large at positive charges. These results indicate that the adsorption behavior of SO42- is significantly different from the adsorption behavior of Cl- and OH-. The absolute values of all dipole moments shown in Figure 9 are much smaller than the values of -21 D (1 D ) 3.336 × 10-30 C m) for SO42-, -8.69 D for Cl-, and -7.34 D for OH-, expected for a dipole formed by a charged sphere of charge z and an ionic radius ri adsorbed on a perfect conductor (equal to the product ze0r 35). The surface dipoles attain a minimum at small charge densities at the metal (near the zero charge). The minimum values of the surface dipole are -0.02 D for OH-, -0.08 D for Cl-, and -0.25 D for SO42-. These results indicate that the dipole formed by the adsorbed anion and its image charge in the metal is significantly screened by the solvent and the metal. Schmickler and Guidelli31-33 described the screening as a combination of the spill over of electrons from the metal and the polarization of the surface solvent molecules. The spill over shortens the ion-image distance. The adsorbed ion polarizes the solvent molecules in its vicinity such that their dipoles

688 J. Phys. Chem. B, Vol. 103, No. 4, 1999

Figure 10. SNIFTIRS spectra in the 3100-1500 cm-1 region for hydroxide adsorbed at the Au(111) electrode from 0.1 M KClO4 + 5.6 × 10-3 M KOH solutions using (s) p-polarized and (‚‚‚) s-polarized infrared beams. For each spectrum, the reference potential E1 was equal to -0.75 V/SCE and the value of E2 is indicated in the figure.

assume an orientation opposite to the adsorbate dipole. The dipole on the solvent in turn depolarizes (screens) the dipole of the adsorbate. This model gives the calculated values of the surface dipoles, which are in excellent agreement with experiment. Figure 9 shows that for σM > 10 µC cm-2, the dipole of adsorbed SO42- significantly increases with the charge on the metal and attains a constant value of -4.5 D at high charge densities. This result indicates that the screening effect diminishes with charge. Following Schmickler and Guidelli,34-36 we can explain this behavior in terms of the solvent displacement by adsorbed sulfate. At higher SO42- coverage, the solvent molecules are squeezed out of the surface and hence the screening of the adsorbed dipole diminishes as well. For Cl-, the surface dipole also increases with the charge, however, in contrast to sulfate, its absolute value remains small, -0.3 D, even at very high charge densities. However, for OH-, the surface dipole decreases when the charge density is greater than 35 µC cm-2. The very small dipole moment for Cl- and OHmay be a result of charge redistribution due to the mixing of electronic states of the adsorbed anion with electronic states of the metal. In fact, our SNIFTIRS spectra will show that for hydroxide adsorption, oxide formation occurs at charge densities greater than 35 µC cm-2. 4. FTIR Studies of OH- Adsorption The above thermodynamic studies have shown that, in the region of high potentials, the hydroxide adsorbed at the Au(111) electrode surface behaves differently from other anions. It is necessary to apply other techniques to investigate OHadsorption in order to cross-check the thermodynamic quantities determined from the chronocoulometry. Hydroxide adsorption at the Au(111) electrode surface may be studied using FTIR spectroscopy.8,14 Figure 10 shows a series of SNIFTIRS spectra measured for hydroxide adsorption from 0.1 M KClO4 + 5.6 × 10-3 M KOH solutions with E2 ) 0.15 V (a) and 0.25 V (b) using p-polarized (solid line) and s-polarized infrared beams (dotted line). The SNIFTIRS spectra in Figure 10 and all the other electroreflectance spectra reported in this study were

Chen and Lipkowski acquired using the reference potential E1 ) -0.75 V(SCE), where hydroxide is totally desorbed from the electrode surface. Two positive bands can be observed in the spectra using p-polarized infrared beams. The band centered at 1618 cm-1 can be assigned to a bending mode of surface H2O since this band is either very weak or has a negative sign in the SNIFTIRS spectra using s-polarized infrared beams. The electric field of an s-polarized photon is equal to 0 at the electrode surface, and hence adsorbed H2O is optically inactive for this polarization.16,38-41 The positive sign of the band recorded using a p-polarized beam may be explained by the fact that at E1 ) -0.75 V(SCE) hydroxide is desorbed from the Au(111) electrode surface. As a result, water molecules move toward the electrode surface. Therefore, this band displays chiefly the properties of surface water molecules at E ) -0.75 V(SCE). This interpretation is in agreement with recent FTIR studies of water at a Au(111) surface by Ataka et al.,42 who demonstrated that at negative potentials the frequencies of the bending mode of surface water are comparable to the frequencies of a dimer of water molecules observed in the range 1600-1620 cm-1. The very broad band in the region from 2600 to 3100 cm-1 can be seen using either p-polarized or s-polarized infrared beams. This band may be assigned to the O-H stretching mode. The positive sign of this band indicates that the IR absorption of the O-H stretching at E1 is stronger than that at E2. Our thermodynamic studies have shown that hydroxide is totally desorbed from the electrode surface at the potential E1 ) -0.75 V(SCE). The question is whether this band is due to the O-H stretching of pure water molecules or due to the presence of aquated hydroxide. To answer this question, we will compare SNIFTIRS spectra measured for 0.1 M KClO4 + 5.6 × 10-3 M KOH and 0.1 M KClO4 + 5.0 × 10-3 M KCl solutions. Our results show that adsorption of hydroxide at the Au(111) electrode surface is only slightly stronger than that of chloride.22 We will assume that the same number of water molecules is displaced from the electrode surface at potential E2, when an equal amount of Cl- or OH- anions adsorbes at the gold surface. If the broad O-H stretching band is due to the change of the surface water, similar SNIFTIRS spectra should be observed in the chloride and the hydroxide solutions. If the broad band is not seen in the spectra recorded for the chloride solution and is observed only for the hydroxide solution, we can assign this band to the aquated hydroxide ions. A comparison of the Gibbs excess of hydroxide (open square) and chloride (open circle) taken from ref 22 is presented in Figure 11A. To compare the SNIFTIRS spectra of hydroxide and chloride adsorption, E1 was set to -0.75 V(SCE), where both hydroxide and chloride are totally desorbed from the Au(111) electrode surface. For hydroxide solution, two values of E2 were chosen: (a) 0.15 V, corresponding to a moderate coverage, and (b) 0.25 V, representing a high coverage of hydroxide. For chloride, the values of E2 were selected as equal to (a) 0.22 and (b) 0.42 V (SCE), where the Gibbs excesses of chloride are equal to hydroxide. Figure 11B shows the comparison of the SNIFTIRS spectra for hydroxide (solid line) and chloride (dotted line), recorded using p-polarized infrared beams. The positive band centered at 1618 cm-1 can be seen in all spectra shown in Figure 11B. The integrated IR intensity of this band has a comparable value for the hydroxide and the chloride solutions. This behavior indicates that comparable amounts of surface water molecules are indeed exchanged when the potential is stepped between E2 and E1 in the chloride and the hydroxide solutions. In contrast, the broad positive O-H stretching band can be seen only in the spectra recorded for

Hydroxide Adsorption at the Au(111) Electrode

J. Phys. Chem. B, Vol. 103, No. 4, 1999 689 than the hydrogen bond between water molecules. The intensity of this band is proportional to the hydroxide concentration, and it is assigned to the O-H stretching of a H2O molecule hydrogen bonded to the oxygen atom of the hydroxide ion.43-51 In Figure 10, a negative band centered at 1635 cm-1, due to the H2O bending vibration, appears in the SNIFTIRS spectrum for hydroxide adsorption at E2 ) +0.25 V(SCE) using spolarization. This negative band at 1635 cm-1 is absent in the spectrum for hydroxide adsorption at a lower potential E2 ) +0.15 V(SCE). The appearance of this band may be explained assuming that oxide formation takes place at E2 ) +0.25 V.

Figure 11. (A) Gibbs excess versus electrode potential plots for the Au(111) electrode in 0.1 M KClO4 + 5.6 × 10-3 M KOH (0) and 0.1 M KClO4 + 5.0 × 10-3 M KCl (O) taken from ref 21. (B) Comparison of SNIFTIRS spectra for (s) hydroxide adsorbed at the Au(111) electrode from 0.1 M KClO4 + 5.6 × 10-3 M KOH solutions and (‚‚‚) chloride from 0.1 M KClO4 + 5.0 × 10-3 M KCl using (A) p-polarized. For each spectrum, the reference potential E1 was equal to -0.75 V/SCE and the value of the E2 is indicated in the figure.

Au + OH- f AuOHad + e-

(11)

AuOHad + OH- f AuOad + H2O + e-

(12)

The consequences of reactions 11 and 12 are the consumption of OH- species and production of H2O in the thin layer between the Au(111) electrode surface and the CaF2 window at potential E2. In contrast, when the electrode potential is stepped back, a reduction of the oxide formed at E2 occurs at E1 ) -0.75 V (SCE).

AuOad + H2O + 2 e- f Au + 2 OH-

Figure 12. SNIFTIRS spectra in the range 3200-1500 cm-1 for hydroxide adsorption at the Au(111) electrode using an s-polarized infrared beam. For each spectrum, the reference potential E1 was equal to -0.75 V and the value of E2 is indicated in the figure. The inset shows the CV curves recorded in 0.1 M KClO4 (‚‚‚) and in 0.1 M KClO4 + 5.6 × 10-3 M KOH (s) at a sweep rate of 10 mV s-1.

the hydroxide solution. It is absent in the spectrum for the chloride solution. These features are consistent with IR43-46 and Raman47-51 studies of aqueous hydroxide solutions. It has been well established that the O-H stretching band, which in pure water is observed in the range 2900-3600 cm-1, is significantly red-shifted in the presence of hydroxide and is observed in the spectral range 2200-3000 cm-1. This shift indicates that the hydrogen bond between OH- and water molecules is stronger

(13)

This reduction consumes H2O molecules and produces OHspecies. Both the oxide formation at E2 and the oxide reduction at E1 should result in the negative band centered at 1635 cm-1 for H2O and the positive broad band for hydroxide in the spectrum. We note that the frequency at the center of this band is somewhat lower than the frequency of 1648 cm-1 measured for water in transmission.42 This shift may indicate that the band at 1635 cm-1 results from an interference between the negative band due to water produced/consumed in reactions 12 and 13 and a positive band seen for the chloride solution and most likely is due to a change in ionic solvation. To determine the potential for the onset of oxide formation, one has to find the potential of the initial appearance of the negative band centered at 1635 cm-1 in the spectrum recorded for the solution of hydroxide. Figure 12 shows a family of SNIFTIRS spectra in the region from 3200 to 1500 cm-1, measured using an s-polarized IR beam with E2 varied from -0.30 to +0.35 V (SCE). The negative peak at 1635 cm-1 appears obviously at the potential +0.20 V, which indicates that oxide formation begins at this potential. The intensity of this peak increases significantly with each increment in the potential. The inset to Figure 12 shows the cyclic voltammetry curves recorded at the Au(111) electrode in 0.1 M KClO4 solution (dotted curve) and in 0.1 M KClO4 + 5.6 × 10-3 M KOH solution (solid curve). These results indicate that oxide formation starts at potentials more positive than the potential of the large peak in the CV curve. The surface concentration of hydroxide at the potential E ) +0.2 V (SCE), shown in Figure 11A is equal to 4.3 × 1014 ions cm-2. For an ideal Au(111) surface, the surface concentration of gold is 1.39 × 1015 atoms cm-2. Therefore, the surface concentration of hydroxide at +0.2 V (SCE) corresponds to the coverage of 1/3 monolayer. It is interesting to compare the integrated IR intensity of the positive broad band shown in Figure 12 with the surface concentration of adsorbed hydroxide determined by chronocoulometry. The SNIFTIRS spectra in Figure 12 were acquired using the reference potential E1 ) -0.75 V (SCE) where hydroxide anions are totally desorbed from the electrode surface. Under these conditions, the measured changes of the reflectivity (∆R/R) represent the difference between the absorption spectrum

690 J. Phys. Chem. B, Vol. 103, No. 4, 1999

Chen and Lipkowski

Figure 13. Comparison of the surface concentration of hydroxide and the integrated IR intensity of the O-H stretch band determined from Figure 12. The inset presents the integrated absorptivity of the O-H stretch band.

of the ΓOH- anions that are in solution at potential E1 and the spectrum of the hydroxide species that are adsorbed at the electrode surface at potential E2. When linearly polarized light is used in the experiment, the Beer’s law is given by IR/IO ) exp(-2.3(E1)Γ cos2 θ)52 and the measured (∆R/R) can be expressed as

(∆R/R) ) 2.3Γ[cos2 θ(E1)(E1) - cos2 θ(E2)(E2)] (14) where R is the electrode reflectivity, θ is the angle between directions of the electric field of the photon and the direction in which the dipole moment of the molecule changes,  is the molar absorption coefficient, and Γ is the surface concentration of the adsorbed species. Using s-polarized light, the electric field of the photon at the surface is equal to 0 and hence ads(E2) is equal to 0 as well. Equation 14 simplifies then to

∆R/R ) 2.3Γdes(E1)

(15)

IR intensity measured from Figure 12 on the electrode potential for the O-H stretching mode (closed circles). The changes in the integrated band intensity follow the changes of the surface concentration determined from chronocoulometic studies. The integrated intensities and the Gibbs excess data may be used to calculate the integrated absorptivity A for adsorbed hydroxide species

A)

1 2.3Γ

dν ∫band ∆R R

(17)

The integrated absorptivity is plotted against the electrode potential presented in the inset to the Figure 13. Its value is essentially independent of potential and is approximately equal to 2.1 × 105 cm mol-1. Overall, the data presented in Figure 13 demonstrate an excellent agreement between the chronocoulometric and IR experiments. 5. Summary and Conclusions

and

∫(∆R/R) dν ) 2.3Γ∫des(E1) dν

(16)

Therefore, the absorption band is positive and the band intensity is proportional to the product of the surface concentration of hydroxide adsorbed at potential E2 and the integrated molar absorption coefficient of hydroxide in the bulk of the investigated solution. Figure 13 shows a plot of the surface concentration versus potential (open squares) and the dependence of the integrated

The adsorption of hydroxide at the Au(111) electrode surface has been investigated by the subtractively normalized interfacial Fourier transform infrared spectroscopy and electrochemical methods such as cyclic voltammetry, differential capacity, and chronocoulometry. The surface tension, Gibbs excess, and Gibbs energy of adsorption were also determined from thermodynamic analysis of the charge density data. The main results of our work are summarized in Figure 14 which shows CV, surface concentration, and electrosorption valency plots. Electrochemical studies indicate that hydroxide adsorption on the Au(111)

Hydroxide Adsorption at the Au(111) Electrode

J. Phys. Chem. B, Vol. 103, No. 4, 1999 691 References and Notes

Figure 14. (a) CV curves recorded at a Au(111) electrode in 0.1 M KClO4 (dotted line) and in 0.1 M KClO4 + 10-3 M KOH solutions (solid line). (b) Gibbs excess versus electrode potential curves for the Au(111) electrode in 0.1 M KClO4 + 10-3 M KOH solution. (c) Electrosorption valency versus electrode potential plot determined from the inset of Figure 9. The charge density curve for the 0.1 M KClO4 + 10-3 M KOH solution was used to plot the electrosorption data as a function of the electrode potential.

surface commences at -0.4 V/SCE. The surface concentration plot indicates that OH- has a three-state character. At E < 0.1 V, the surface concentration changes slowly with potential, while at higher E it increases steeply. The changes of the surface concentration correlate well with the change of the electrosorption valency. The electrosorption valency is small for E < 0.1 V and increases steeply at E > 0.1 to reach unity. At E ≈ 0.2, the electrosorption decreases somewhat to rise again at the positive limit of potentials. Our IR experiments indicate that this increase correlates well with the onset of the oxide formation. The changes of the electrosorption valency were explained in terms of the change of the surface dipole formed by the adsorbed hydroxide and its image charge on the metal. The surface dipole is quite large at the negatively charged surface, indicating that in this range hydroxide forms quite a polar bond with the gold surface. The polarity of the bond significantly decreases at positive charge densities due to either a very strong screening of the charge on the anion by the charge in the metal or a significant charge transfer from the adsorbed ion to the gold surface. The adsorption of OH- has been compared to the behavior of SO42- and Cl- at the Au(111) electrode surface. Our results show that the adsorption strength increases in the order SO42- < Cl- < OH-. Our present work demonstrates excellent agreement between the results of thermodynamic and spectroscopic investigations of OH- adsorption. Acknowledgment. This work was supported by a grant from Natural Sciences and Engineering Council of Canada. The authors express their gratitude to Dr. D. E. Irish and Dr. B. Pettinger for helpful discussions.

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