Iodide adsorption at the Au(111) electrode surface

Journal of Electroanalytical Chemistry 467 (1999) 342 – 353

Iodide adsorption at the Au(111) electrode surface Aicheng Chen a, Zhichao Shi a, Dan Bizzotto a, Jacek Lipkowski a,*, Bruno Pettinger b, Christoph Bilger b b

a Guelph-Waterloo Center for Graduate Work in Chemistry, Guelph Campus, Uni6ersity of Guelph, Ontario N1G 2W1, Canada Fritz-Haber-Institute der Max-Planck-Gesellschaft, Department of Physical Chemistry, Faradayweg 4 – 6, D-14195 Berlin, Germany

Received 28 July 1998; received in revised form 10 November 1998; accepted 13 November 1998

Abstract The adsorption of iodide at a Au(111) single crystal electrode has been investigated quantitatively using chronocoulometry. By analyzing the charge density data thermodynamically, the following parameters were determined: the Gibbs excess, Gibbs energy of adsorption, the number of electrons flowing to the interface per one adsorbed iodide ion at a constant electrode potential (electrosorption valency), and at a constant chemical potential. The thermodynamic data for iodide adsorption were compared to the results for bromide and chloride adsorption. All the three halides form a chemisorption bond with the gold surface. The bond is quite polar at the negatively charged surface, however, its polarity drops significantly at the positively charged surface. At low charge densities and coverages, the bond polarity is determined by the ability of free electrons to screen the dipole formed by the adsorbed anion and its image charge in the metal. At high charge densities and coverages, the chemisorption bond has a predominantly covalent character. The strength of the halide adsorption and the covalent character of the chemisorption bond increase progressively by moving from chloride to iodide. © 1999 Elsevier Science S.A. All rights reserved. Keywords: Adsorption; Au(111) electrode; Iodide; Chronocoulometry

1. Introduction The objective of this work is to describe thermodynamics of iodide adsorption at a Au(111) electrode surface. Iodide adsorbs strongly at the gold surface and forms well ordered overlayers, whose structure and composition have been investigated extensively by many ex situ and in situ techniques such as low-energy electron diffraction (LEED), Auger electron spectroscopy (AES) and X-ray photoelectron spectroscopy (XPS) [1,2], surface-enhanced Raman scattering (SERS) [3], scanning tunneling microscopy (STM) [4 – 11], surface X-ray scattering (SXS) [12,13] and electrochemical quartz crystal microbalance [14,15] techniques. There

 Dedicated to Jean Clavilier on the occasion of his retirement and in recognition of his contribution to electrochemistry. * Corresponding author. Fax: +1-519-7661499. E-mail address: [email protected] (J. Lipkowski)

has been significant interest in I − adsorption at the Pt(111) electrode [16–20] and Ag(111) surface [21] as well. Early ex situ LEED [2] and in situ STM [4–6,8,9] studies reported that iodide forms ordered commensurate overlayers at the Au(111) electrode surface. More recent and more precise SXS investigations by Ocko et al. [12,13], demonstrated that the adsorbed iodide forms two distinct incommensurate ordered phases. The first is observed at coverages between 0.37 and 0.41 ML (where the coverage of one monolayer (ML) corresponds to the density of atoms at the ideal 1×1 Au(111) surface equal to 1.39× 1015 atoms cm − 2). It has a rectangular, uniaxially incommensurate (p× 3) structure. At coverages higher than 0.41 ML this structure undergoes a first order transformation to a rotated-hexagonal, phase. The two phases are electrocompressible. The results of SXS studies were confirmed by STM measurements in Weaver’s laboratory

0022-0728/99/$ - see front matter © 1999 Elsevier Science S.A. All rights reserved. PII: S 0 0 2 2 - 0 7 2 8 ( 9 8 ) 0 0 4 3 7 - 9

A. Chen et al. / Journal of Electroanalytical Chemistry 467 (1999) 342–353

[7], published almost at this same time, and more recently by both ex situ LEED and in situ STM experiments performed by Itaya and co-workers [1,10,11]. In contrast to the detailed knowledge concerning the structure of iodide overlayers, very little is known about energetics of iodide adsorption at gold surfaces. In this work, we have employed cyclic voltammetry, differential capacity, and chronocoulometry to determine surface concentrations, Gibbs energies of adsorption and electrosorption valences for iodide at the Au(111) surface. We have combined these thermodynamic data with the structural information available in the literature to give a comprehensive description of iodide adsorption at the gold electrode.

2. Experimental The experimental procedures and instrumentation were described in [22,23]. The working electrode (WE) was a Au(111) single-crystal rod which was grown, cut, and polished in our laboratory, and 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 (RE) 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 used subsequently to determine the relative Gibbs excesses, the Gibbs energies of adsorption, and the electrosorption valency. The data treatment procedures have been described in [24 – 26]. Water was purified in a tandem of the Milli-Q and Milli-Q plus UV systems (18.3 MV cm). The supporting electrolyte was 0.1 M KClO4 purified according to the procedure described in our previous paper [27]. Suprapure potassium hydroxide monohydrate (Merck) and potassium iodide (Aldrich, \99.99%) were used without further treatment. An Orion Research, model EA 920 expandable ion analyzer was used to measure the pH of the investigated solutions. The KOH was added to the KClO4 solution to suppress the hydrogen evolution reaction and to extend the double layer region of the gold electrode to more negative potentials. All solutions were deaerated with argon; during the experiment argon was passed over the top of the solution. All measurements were conducted at room temperature (2092°C).

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3. Results and discussion

3.1. Cyclic 6oltammetry and differential capacity Cyclic voltammetry was used for qualitative characterization of the electrode surface. Fig. 1(A) shows cyclic voltammetry curves (CVs) recorded for the Au(111) electrode in 0.1 M KClO4 supporting electrolyte, in 0.1 M KClO4 + KOH, and 0.1 M KClO4 + KOH+ 10 − 3 M KI solutions (pH 11). The shape of these curves agrees well with the shape of the CVs reported in the literature [28–30]. The CV recorded for the neutral supporting electrolyte (dotted line) shows that, in this solution, the double layer region extends from − 0.8 to +0.6 V. For potentials more negative than − 0.8 V hydrogen evolution on a Au(111) surface is observed. For potentials more positive than +0.6 V oxidation of the gold surface takes place. Upon the addition of potassium hydroxide (dashed line), the onset of gold oxidation shifts to ca. +0.3 V. In addition,

Fig. 1. (A) Cyclic voltammograms recorded at a Au(111) electrode in 0.1 M KClO4 (dotted line), 0.1 M KClO4 +KOH (dashed line), and 0.1 M KClO4 +KOH + 10 − 3 M KI solutions (solid line) at a sweep rate of 10 mV s − 1; (B) Differential capacities determined using an ac perturbation of 25 Hz frequency, 5 mV r.m.s. amplitude at a sweep rate of 5 mV s − 1.

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a tall peak seen at +0.235 V for the positive sweep, may be assigned to the adsorption of OH − ions at the Au(111) surface. The dashed and the dotted lines merge at EB −0.3 V indicating that the OH − is desorbed from the electrode surface at these negative potentials. The symmetric, quasi rectangular shape of the section of CVs between −0.8 and −0.1 V indicates that the two supporting electrolyte solutions were free of oxygen and that no creeping of the solution onto the electrode walls occurred in this region. When KI is present in the solution, the shape of the CV (solid line) changes significantly. Three pairs of peaks are seen on this curve. The sharp (irreversible) peak at E : −0.48 V (SCE) corresponds to the iodide-induced lifting of the (1 × 23) reconstruction of the Au(111) surface [4,10– 13]. The second pair of reversible peaks located at − 0.38 V (SCE) reflect the formation of an ordered (p × 3) iodide adlayer [12,13]. The third pair of peaks near 0.2 V (SCE) corresponds to the phase transition between the (p× 3) and the hexagonal structure of the iodide adlayer. These peaks display a strong dependence on the bulk iodide concentration. Their height increases and the position of the peaks shifts in the negative direction with an increase in the bulk iodide concentration. We note that for iodide solutions the CVs were recorded only up to E = 0.3 V (SCE), to prevent polyiodide formation and oxidation of iodide at more positive potentials [4]. Fig. 1(B) shows differential capacity curves for the Au(111) electrode in the three electrolyte solutions, recorded using a slow positive going voltage sweep. The shape of the differential capacity curves is in general consistent with the shape of CVs. However, due to the irreversibility of the surface reconstruction and slow kinetics of the phase transition in the iodide adlayer, the peaks corresponding to these surface processes are less pronounced on the differential capacity curves than on CVs. The differential capacity curves show clearly that the total desorption of iodide takes place at potentials more negative than − 1.0 V and that KOH has to be added to the solution in order to investigate the adsorption of iodide at E B − 0.8 V(SCE). The differential capacity curves also show that adsorption of iodide apparently suppresses the adsorption of OH − and that the presence of OH − ions in the solution should not affect the determination of the surface concentration of iodide. The cyclic voltammetry and the differential capacity curves will be used below to select the experimental conditions for the potential step experiments

3.2. Electrode charge density Potential step experiments were performed to determine the charge density (sM) at the electrode surface. The procedure used to measure sM in the presence of

adsorbed ions (or molecules) requires that the electrode potential be stepped from a potential at which anions are adsorbed to a potential where they are desorbed totally from the electrode surface [22–27]. In the present case, the potential of total desorption corresponds to EB − 1.0 V. We have observed quite a pronounced ‘creeping’ of the electrolyte onto the walls of the gold single crystal electrode at these negative polarizations. The ‘creeping’ was particularly significant when the electrode was held at these negative potentials for a prolonged period of time and could be minimized or eliminated by shortening the negative polarization time. We note that similar problems with ‘creeping’ in alkaline solutions were reported by Hamelin et al. [31]. To overcome the problem with ‘creeping’ two series of potential step experiments were performed for each concentration of iodide. In the first series, the electrode was held at potential E, a value which varied from − 0.75 to + 0.275 V(SCE) for a period of time long enough to reach adsorption equilibrium (up to 3 min). The potential was then stepped to E1 = − 0.8 V(SCE) where most of the iodide adsorbed at potential E was desorbed from the electrode surface. In the second series of measurements, the electrode potential was held at a value −0.75 V(SCE) for a period of time long enough to establish the adsorption equilibrium. The potential was then stepped to a value E2, which varied from −0.75 V(SCE) to −1.05 V(SCE). The potential was held at E2 for a period of 1 s, which was sufficiently long so that a new state of adsorption equilibrium could be established, but short enough to prevent ‘creeping’ of the electrolyte onto the electrode walls. The potential was then stepped to the value E3 = −1.1 V(SCE) where iodide desorbs totally from the electrode surface. The difference between the charge densities at potentials E2 and E3 (DsM2) was measured during a very short time interval of 0.2 s. Next, the charges DsM1 and DsM2 were added and the difference between charge densities at potentials E and the potential of total desorption E3 (DsM) was calculated. Similar experiments were performed for the pure supporting electrolyte (0.1 M KClO4 + KOH) and for a neutral 0.1 M KClO4 solution. For the 0.1 M KClO4 solution the absolute charge densities sM were calculated from the measured DsM with the help of the independently determined value of the potential of zero charge (pzc). At E= −0.8 V(SCE), the charge density sM in 0.1 M KClO4 + KOH solutions (pH 11.0) has the same value as in a neutral 0.1 M KClO4 solution, since at this potential OH − ions are desorbed totally from the electrode surface. Hence, for the alkaline solution of the supporting electrolyte, the value of sM at E= −0.8 V(SCE) was used to calculate the absolute charge densities at all other potentials. Finally, the charge densities for iodide solutions are equal to that for the alkaline supporting electrolyte at E= − 1.1 V(SCE), since io-

A. Chen et al. / Journal of Electroanalytical Chemistry 467 (1999) 342–353

Fig. 2. Charge density versus electrode potential curves for the Au(111) electrode in pure 0.1 M KClO4 solution (dotted line), 0.1 M KClO4 + KOH (dashed line), and 0.1 M KClO4 + KOH + xM KI solutions. The concentration of KI in mol dm − 3 was: 1.00 × 10 − 5, 1.78× 10 − 5, 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, and 1.00× 10 − 2. The inset shows sections of the relative charge density curves measured in 0.1 M KClO4 + 10 − 5 M KI solutions with () and without () the addition of KOH.

dide is desorbed completely from the electrode surface at this potential. The value of sM at E = −1.1 V(SCE) was therefore used to convert the measured DsM for the iodide-containing solutions into the absolute charge densities. Fig. 2 shows the absolute charge density versus potential curves for the neutral electrolyte (dotted line), alkaline supporting electrolyte (dashed line), and the solutions with 13 different concentrations of iodide (solid lines). The dotted line merges with the dashed line for EB − 0.3 V. Likewise, the charge density curves measured for iodide solutions merge with the curve for the alkaline supporting electrolyte for EB − 1.0 V. This behaviour is consistent with the measurements of CV and differential capacity discussed earlier. Adsorption of I − apparently causes a positive charge to flow to the metal side of the interface. The charge density curves for iodide solutions display multiple sections. The first observed at E B −0.55 V is characterized by the presence of a small step. The position of the second fast rising segment of the curve correlates well with the position of the sharp peak and the large broad peak on the CV curve shown in Fig. 1(A). The last section is seen at the more positive potential and corresponds to a gradual, quasi-linear change of the charge with potential. It is interesting to compare the potentials of zero charge for the Au(111) surface in different solutions. The pzc shifts from +0.29 to + 0.16 V on addition of KOH to 0.1 M KClO4 solution (pH ca. 11.0). In solution with 10 − 5 M KI the pzc shifts further towards −0.35 V. Apparently, the ad-

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sorption of I − is much stronger than that of OH − ion, and hence the presence of OH − in the solution has little effect on the iodide adsorption. This point is further illustrated by the inset to Fig. 2, which shows two series of relative charge densities determined for 0.1 M KClO4 + 10 − 5 M KI solution with and without addition of KOH. Apparently, even for the lowest concentration of KI investigated in this work, the charge densities are independent of the concentration of OH − in the supporting electrolyte. Fig. 3 compares results of cyclic voltammetry, differential capacity, and charge density measurements. The main section of Fig. 3 shows the differential capacity curves calculated by numerical differentiation of the chronocoulometric charge density data, from the positive going section of the cyclic voltammogram and measured using the single frequency ac impedance technique. The inset to Fig. 3 compares the charge density data determined by the chronocoulometric technique to charge densities calculated by integration of the CV and the differential capacity curve. The results show quite good agreement between the data determined by chronocoulometry and cyclic voltammetry. However, charge densities calculated by integration of the CV curve are somewhat smaller than charge densities measured by chronocoulometry. Apparently, the adsorption equilibrium was achieved during the waiting time of 3 min, used in chronocoulometric experiments, while absorption equilibrium was not completed fully when a sweep rate of 10 mV s − 1 was employed to record the CV curve. Clearly, the data obtained from the single frequency ac impedance measurements deviate signifi-

Fig. 3. Comparison of differential capacity curves for a Au(111) electrode in 0.1 M KClO4 +10 − 3 M KI solution determined from: ( — ) ac impedance experiment using an ac perturbation of frequency 25 Hz and amplitude 5 mV rms at a sweep rate of 5 mV s − 1; (· · · ·) CV shown in Fig. 1(a); (---) differentiation of the charge density curve shown in Fig. 2. Inset: comparison of charge density determined by: ( — ) integration of the single frequency differential capacity; (· · · ·) integration of the CV; (---) from chronocoulometry.

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cantly from results of the two other techniques. The CV results are therefore significantly distorted due to kinetic limitations. Therefore, further quantitative data analysis was based on the data derived from chronocoulometric experiments.

3.3. The interfacial tension and Gibbs excess data In this work, the adsorption of iodide took place from a solution with an excess of supporting electrolyte. The electrocapillary equation for a gold electrode in contact with that electrolyte is given by [32]: − dg = sMdE +GI − RTd ln cKI

(1)

where g is the interfacial tension, E is the electrode potential measured with respect to SCE, and GI − is the Gibbs excess of iodide ions. In the presence of an excess of an inert electrolyte, GI − may be considered as equal to the Gibbs excess of specifically adsorbed anion [32]. In order to calculate the Gibbs excess of iodide ions, the charge density curves in Fig. 2 were integrated to give the relative interfacial tension plots. Next, the relative values of g at a constant E were plotted versus ln cKI and the resulting curves were differentiated to give GI − . Fig. 4(a) 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 coverages and is characterized by a very small slope. The second section at intermediate coverages is very steep. It is followed by the third section, which corresponds to a quasi-plateau, within which the surface concentration changes slowly with E. At potential around −0.3 V(SCE) an ordered, solidlike overlayer of adsorbed iodide is formed [1,12,13,18,19]. Our surface concentration data indicate that the transition from the mobile to the solid-like overlayer takes place without any significant increase of the surface coverage of iodide. We note that the limiting surface concentration, calculated for a closed packed hexagonal monolayer of I − , from the van der Waals radius (2.15 A, [13]) is approximately equal to 6.3×1014 ions cm − 2. The maximum surface concentration shown in Fig. 3 is only slightly higher than this theoretical value. Apparently, by moving the potential from the negative limit in the positive direction, the surface concentration of iodide changes from zero to the maximum value corresponding to the close packed monolayer of iodide ions (iodine adatoms). For 0.01 M KI solution the surface concentrations determined from the chronocoulometric measurements may be compared to the packing densities calculated from the SXS experiments and STM images of the ordered iodide adlayer. The inset to Fig. 4(a) shows a

Fig. 4. Plots of the Gibbs excess of I − against (a) electrode potential and (b) electrode charge density at the Au(111) electrode. Inset to (a): comparison of the surface coverage of iodide determined from (“) chronocoulometry and () SXS. The two black squares mark the coverages determined from STM images in [7]. Inset to (b): comparison of the Gibbs excess of I − at (“) the Au(111) surface and () a Hg surface [36].

comparison of surface coverages determined by the thermodynamic method to the coverages obtained from the X-ray diffraction studies by Ocko et al. [12,13] and from the STM images determined in Weaver’s laboratory[7]. The Gibbs excesses calculated from the chronocoulometric data are somewhat higher than the packing densities obtained by SXS or STM. However, within the potential range of the (p× 3) structure the difference is less than 10% and it increases to about 15% at more positive potentials where the hexagonal structure is formed. The coverages obtained from SXS experiments were calculated assuming that the whole surface is perfectly flat and that the overlayer is perfectly ordered. The chronocoulometric experiments correspond to a real surface which has many defects and has some microroughness and represents a real overlayer which is never perfectly ordered. Therefore, it is expected that coverages determined from chronocoulometric experiments are somewhat higher than coverages calculated from the diffraction or STM data. We note

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also that the surface concentrations presented in this work are in reasonable agreement with results of recent quartz microbalance measurements [15]. The latter results are somewhat too low, because the negative limit of potentials employed in [15] was ca. − 0.6 V(SCE). Fig. 4(a) shows that iodide is already adsorbed at these potentials. Since zero iodide coverage at the negative potential limit was assumed in [15], the data reported in Lei et al. [15] contains a negative error. Independently, the Gibbs excesses were also determined at constant charge by plotting the Parsons function j= sME+g [33] and differentiating the relative j versus ln cKI plots at a constant sM. Fig. 4(b) shows two surface concentration versus charge density curves determined for the highest and the lowest concentration of KI in the bulk. The G versus sM plots depend weakly on the bulk KI concentration and hence for the sake of clarity only two curves are shown in Fig. 4(b). Weak dependence of the Gibbs excess versus charge density plots on the bulk anion concentration was also observed in the previous studies of Br − [26] and SO24 − [24] adsorption at the Au(111) electrode. This behavior indicates that the derivative (( ln cKI/(sM)G is large. The coverage versus charge density plots are fairly linear. Their slope is known as the Esin – Markov coefficient. The reciprocal of the Esin – Markov coefficient gives the charge number per one adsorbed ion at the constant chemical potential of iodide. This charge number depends only weakly on the bulk concentration and on average is equal to about −0.96 electron ion − 1. Similar values of this charge number were observed for adsorption of Br − [26] and Cl − [25] at the Au(111) electrode. The inset to Fig. 4(b) compares the G versus sM plots for iodide adsorption at the Au(111) and Hg electrodes. The curve for Hg has a somewhat larger slope and is shifted significantly towards positive charges. This behavior indicates that iodide adsorption is weaker and the absolute value of the charge number at the constant chemical potential is smaller at mercury than on gold.

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Fig. 5. Fit of the adsorption data to the equation of the square root isotherm: (a) at constant potential, (b) at constant sM.

Fig. 5(a and b) show plots of the square root of the surface pressure versus ln (kTcKI/F). All plots are fairly linear. Extrapolation to zero surface pressure gives an intercept with the ln (kTcKI/F) axis equal to lim ln (kTcKI/F)F = 0 = − ln b

(3)

from which the Gibbs energies of adsorption can be calculated. Note that in Eqs. (2) and (3) the potassium iodide concentrations are multiplied by the term kT so that in the limit of low coverage the film pressure is described by Henry’s law F= kTbcKI as explained in [34]. The Gibbs energies of adsorption, determined using this procedure, are plotted against potential in Fig. 6(a) and versus charge in Fig. 6(b). The standard state is an ‘ideal’ G= 1 ion cm − 2 for the adsorbed species and an ‘ideal’ cKI = 1 mol dm − 3 for the bulk species.

3.4. Gibbs energies of adsorption The Gibbs energies of adsorption were determined by fitting the surface pressure data F to an equation of a ‘square root’ isotherm [34,35]: ln (kTcKI)+ ln b = ln F+BF 1/2

(2)

where b =exp − (DG/kT) is the adsorption equilibrium constant, B is a constant and F is the surface pressure. At a constant potential, F was set to F= gu = 0 − gu and at a constant charge, F= ju = 0 −ju. The subscripts u and u =0 denote the values of the interfacial tension and the Parson’s function in the presence and absence of iodide ion in the supporting electrolyte, respectively.

Fig. 6. (a) Plot of the Gibbs energy of adsorption versus electrode potential ( ) determined from the F 1/2 versus ln (kTcKI/F) plots at constant E in Fig. 5(a) and () from the F 1/2 versus ln (kTcKI/F) plots at constant sM in Fig. 5(b). (b) Plot of the Gibbs energy of adsorption versus electrode charge density determined from the F 1/2 versus ln (kTcKI/F) plots at constant sM in Fig. 5(b).

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The square root isotherm is an empirical isotherm and although it is called a ‘virial type isotherm’, Eq. (2) cannot be derived from the equation of state for the virial isotherm [35]. In addition, the isotherm is not based on a well defined physical model and hence it should be viewed only as a convenient procedure to linearize experimental data. Although the fit to the square root isotherm gives linear plots, the extrapolation to zero surface pressure is very long. Consequently, even a small change in the slope may lead to a significant error in the intercept. The constant potential and the constant charge analysis gives plots of a different slope. For this reason, it is useful to compare Gibbs energies at constant E and at constant sM to check the consistency of their values. Therefore, the DG data determined from the constant charge analysis were also plotted against potential in Fig. 6(a). The charge density versus potential curve for the iodide-free electrolyte was used to construct this plot. The agreement between the DG values determined from the constant charge and the constant potential analysis is not as good as in the case of adsorption of sulfate, chloride, or bromide [24 – 26]. For iodide, the reported values of Gibbs energies of adsorption contain some systematic errors. The first derivative of DG versus E is equal to the electrosorption valency g%. Independently, the electrosorption valency can be determined from the slope of the charge density versus Gibbs excess plots: (4) Fig. 7 shows a plot of sM versus GI − for various electrode potentials. The plots for lower potentials (EB −0.4 V) are shown in the main section of Fig. 7, and the plots for higher potentials are presented in the inset. The data display a linear relationship between sM and GI −; at higher potentials, a relatively narrow range of Gibbs excesses was available. Fortunately, the plots included the values of sM measured in the pure supporting electrolyte, and hence the uncertainty of their slopes was somewhat reduced. The electrosorption valences calculated from the slopes of the sM versus GI − plots, and by differentiation of the DG versus E curves, are plotted as the number of electrons flowing to the interface per one adsorbed iodide ion in Fig. 8(a). Although, the two curves display a similar shape the numerical differences between electrosorption valences determined by the two methods are significant. These differences indicate that the data may be affected by systematic errors. For this reason it is useful to compare the electrosorption valences for iodide to the values of this parameter determined earlier for chloride and bromide. For the three halides, the electrosorption valences calculated from the slope of the charge versus the Gibbs excess plots are compared in Fig. 8(b). The three curves have quite a similar sigmoidal shape. The curves

Fig. 7. Plots of the electrode charge density versus Gibbs excess of I − taken at constant values of the electrode potentials given in mV versus (SCE). Inset: sM versus G at more positive electrode potentials.

may be seen as composed of two segments separated by a step. The first at negative potentials consists of g% values significantly lower than unity. This corresponds to adsorption of anions at a negatively charged surface. The second which consists of g% values closer to unity corresponds to adsorption at a positively charged surface. Apparently, adsorption of halides has a different character at a negatively and a positively charged surface. Significantly, the absolute value of the electrosorption valency increases progressively moving from chloride to iodide.

3.5. Esin–Marko6 coefficients Cross-differentiation of Eq. (1) gives the expression for the Esin–Markov coefficient: (5) The first derivative in Eq. (5) is given in terms of a directly measured quantity such as sM and experimental variables such as potential and bulk potassium iodide concentration. The second derivative is given in terms of the Gibbs excess, whose calculation involves one integration and one differentiation step. Therefore, Eq. (5) may be used to check if major errors were included in the calculation of GI − . Fig. 9(a) shows a plot of E versus − RT ln cKCl for various electrode

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charge densities in the range −10 to 60 mC cm − 2; all plots are linear and their slopes give the Esin – Markov coefficients. Independently, the GI − versus sM plots from Fig. 4(b) were fitted by a polynomial and differentiated numerically to give [(GI − /(sM]ln c KI. The two independently determined derivatives are plotted against each other in Fig. 9(b); the black points are the experimental data, and the dotted line is a line with a unit slope. The experimental points scatter randomly around the dotted line, indicating that our GI − are free of major errors of the data processing.

3.6. Model of the inner layer– surface dipole formed by adsorbed I − ion and its image charge in the metal. In this section we will employ the Grahame – Parsons [36,37] model of the inner layer to give a physical interpretation to some of the thermodynamic quantities described in the preceding sections. Details of this model and the data analysis procedures are described in previous papers [24– 26]. To avoid repetition, here we will include only the material which is needed for the clarity of our presentation. We will discuss parameters that characterize the structure of the inner layer in the presence of iodide adsorption by comparing them to the parameters determined earlier for chloride and bromide. The inner layer capacity C i is calculated from the overall electrode capacity C determined by differentiation of the charge density curves and using the theory Fig. 9. (a) Esin – Markov plots. (b) Comparison of the Esin –Markov coefficients determined from the slope of the E versus ln cKI plots in Fig. 9(a) and by differentiation of the GI − versus sM plots shown in Fig. 4(b).

Fig. 8. (A) Electrosorption valences () determined from the slope of the charge versus the Gibbs excess plots and ( ) from the Gibbs energy versus electrode potential plots. (B) Comparison of the electrosorption valences of the three halides.

of the diffuse layer as described in [24–26,35]. The inner layer capacities determined for iodide are plotted against the charge on the metal in Fig. 10 together with the capacities determined earlier for chloride and bromide. These curves are independent of the bulk concentration of the corresponding anion. The three curves display quite a similar shape. They consist of a tall peak located close to the zero charge density and a broad peak or a shoulder at high positive charge densities. The two peaks grow taller and their position shifts towards negative charges by moving from chloride to iodide. The inset to Fig. 10 shows the inner layer capacities for the three halides adsorbed at a mercury electrode, taken from the literature [36,38,39]. For mercury, the trend for the change of the inner layer capacity with the nature of the anion is quite similar to that for gold. However, the inner layer capacities for gold are more than an order of magnitude larger than those for mercury. The inner-layer capacity is a function of two variables, the charge on the metal and the amount of

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Fig. 10. Comparison of inner layer capacities of the Au(111) electrode for iodide, bromide, and chloride. Inset, comparison of the inner layer capacities of the Hg electrode for the three halides.

adsorbed anion. It may therefore be expressed in terms of two components, the capacity of the inner layer at a constant charge (GC) and the capacity of the inner layer at a constant amount adsorbed (sC). The capacity GC may be determined from the slope of the plot of the potential drop across the inner layer (Df M − 2 =E− Epzc −f2) versus the charge −FGI of adsorbed iodide at constant sM (f2 is the outer Helmhlotz plane potential). These plots are shown in Fig. 11. They are linear and this feature indicates that the inner layer capacities are independent of the iodide coverage within the narrow range of coverages used to construct these plots. The capacity at a constant amount adsorbed may then be calculated from C i and GC using the formula: (C i) − 1 =(sC) − 1 −F((GI/(sM)(GC) − 1

(6)

Fig. 11. Plot of the potential drop across the inner layer DfM − 2 versus the charge of adsorbed I − (–FGI –) at constant charge density at the metal side of the interface as indicated in the figure.

Fig. 12. Comparison of the component of the inner layer capacities for I − , Br − , and Cl − (A) at constant charge and (B) at constant amount adsorbed.

The components of the inner layer capacity at a constant charge and at a constant amount adsorbed are plotted against the charge on the metal in Fig. 12(a and b), respectively. These capacities may be considered as integral capacities described by C= o/(x2 − x1)

G

(7)

and C= o/x2

s

(8)

where x1 and x2 are the locations of the inner and outer Helmholtz plane, respectively. The two components of the inner layer capacity increase by moving from iodide to chloride. This trend correlates well with the decreasing size of the anion. The interpretation of GC curves is difficult, since they may be affected by both a change of the permittivity and the position of the inner Helmoltz plane. The interpretation of sC is easier, particularly if one assumes that the thickness of the inner layer x2 does not change with the charge. In that case, the shape of sC displays the change of the permittivity. For chloride and bromide the sC curve has a maximum at small charge densities. The sC curve for iodide is essentially featureless. The changes of the permittivity may be interpreted in terms of a change in the orientation of surface water molecules. Specifically, the maximum seen at small charge densities for chloride and bromide may be explained in terms of an increase of the permittivity due to the disorientation of water dipoles. The absence of a maximum on the sC curve for iodide may therefore indicate that this anion has a weak influence

A. Chen et al. / Journal of Electroanalytical Chemistry 467 (1999) 342–353

on the structure of the surface water. The solvent reorientation phenomena can be studied today with the help of surface spectroscopies. Recently, Ataka and Osawa [40,41] employed surface enhanced infrared absorption spectroscopy to describe the structure of surface water in the presence of adsorbed sulfate and perchlorate ions at a Au(111) electrode surface. For sulfate, the changes in the structure of the surface water observed by IR spectroscopy correlate well with the shape of the inner layer capacity at the constant amount adsorbed. One may hope that the spectroscopic information concerning the orientation of surface water molecules will soon be available for the Au(111) surface in the presence of adsorbed halides. Quite pronounced differences between the inner layer capacities measured for these anions suggest that the structure of the surface water may vary significantly in the presence of these anions. These data offer an interesting opportunity to test the validity of the Grahame–Parsons model of the inner layer. 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 [37]: g%= z(sC/GC)= z(x2 −x1)/x2

(9)

The electrosorption valency and the capacity at the constant amount adsorbed may be used further to calculate the dipole formed by an adsorbed anion and its image charge in the metal-surface dipole (ms) [42– 45]: ms = zeoo(1− g%/z)/sC

(10)

The surface dipole is a direct measure of the polarity of the electrosorption bond. Fig. 13 shows the surface dipole for three halide ions. The inset to Fig. 13 plots the electrosorption valency for iodide determined from the ratio of the inner layer capacities at constant amount adsorbed and at constant charge with the help of Eq. (9). The electrosorption valences agree quite well with the values of this parameter calculated earlier from the slope of the charge versus coverage plots and presented in Fig. 8. Apparently, the values of the surface dipole increase in the order I − BBr − B Cl − indicating that the polarity of the chemisorption bond increases by moving from iodide to chloride. The polarity of the electrosorption bond is a strong function of the charge on the metal. The dipoles are quite large at the negatively charged surface, but decrease significantly at the positively charged surface. The polarity goes through a minimum at small absolute values of the charge density. The minimum shifts in the direction by moving from chloride to iodide. For iodide the minimum value of the dipole moment is equal to zero within the limits of the experimental error. This behavior suggests that the adsorbed anion is discharged effectively and that

351

the adsorbed species is essentially the iodine adatom. The adsorption of iodide involves therefore a significant charge redistribution.

3.7. Electronic structure of the surface The adsorbate-induced changes of the electronic structure of the metal may be studied conveniently using second harmonic generation spectroscopy (SHG) [46–51] (see also electroreflectance studies of the metal electronic structure [52]). There are two major contributions to the measured SHG signal. The first comes from interactions of the electromagnetic fields of the incident photons with the free electrons in the metal and involves excitation of free electrons to a virtual state. The second arises from coupling of the optical fields of the incident photons with the interband transitions and has a resonant nature. It involves electronic transitions between two real electronic states. The photons emitted by these two mechanisms are emitted at different times and consequently are phase shifted [49–51]. The mixing of the electronic states of the adsorbate with electronic states of the metal enhances the contribution of the interband transitions to the measured signal and hence changes the phase angle [46–48]. The measurement of the phase angle provides therefore a convenient means to study the adsorbate-induced changes of the metal electronic structure. Therefore, we have employed the interference second harmonic generation spectroscopy (ISHG), developed recently by Pettinger and Bilger [53] to measure the potential and charge dependence of the phase angle at the Au(111) surface in the presence of the three halides.

Fig. 13. Plots of the effective dipole moment formed by I − , Br − , and Cl − ion adsorbed at the Au(111) surface versus the charge density on the metal. Inset: electrosorption valency versus electrode charge density plots determined from the ratio of sC to GC.

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Fig. 14. Plot of the difference between the phase angle of the SHG photons generated at the Au(111) surface and in the quartz lamella versus potential (top panel), and versus charge density (bottom panel) in 0.1 M NaClO4 (“) and 0.1 M NaClO4 + 10 − 3 M NaCl( ), NaBr(), NaI (2).

The top panel in Fig. 14 shows how the phase angle changes with the electrode potential for pure 0.1 M NaClO4 and sodium perchlorate solution with the addition of millimolar concentrations of the three halides, respectively. The phase angle displays quite a strong dependence on the electrode potential and the nature of the anion. These changes suggest that interband transitions contribute to the observed SHG signal. In the case of the Au electrode and the energy of the input photons 1.17 eV, the interband transitions involve the upper edge of the d band located 2 eV below the Fermi level. The potential dependence of the phase angle apparently increases in the order ClO4− BCl − B Br − BI − , which corresponds to the sequence of the increasing adsorption strength. For bromide we used this change of the phase angle as evidence of the mixing of the electronic state of the adsorbed anion with the electronic states in the metal [26]. However, if the same phase angle data are plotted against the charge on the metal (bottom panel in Fig. 14), the dependence of the phase angle on the nature of the specifically adsorbed anion is to a large extent removed. For charge densities lower than 60 mC cm − 2, all phase angle data obey one common relationship. For charges higher than 60 mC cm − 2, the phase angle data corresponding to different anions diverge. This behavior correlates quite well with

the fact that the anions form disordered (mobile) adlayers at sM B 60 mC cm − 2 and ordered (2D solid-like) adlayers at higher charge densities. Apparently, for the mobile adlayer, the changes of the phase angle are due rather to the change of the field at the interface, than to the change of the electronic structure of the metal. In contrast, the electronic structure of the interface is changed when an ordered adlayer is formed. For bromide and chloride, no apparent change of the electrosorption valency was observed at the potential of the order–disorder phase transition. Therefore, no extra charge is transferred to the interface when new electronic states are formed by the ordered adlayer. For iodide, the step on the complex phase angle plot, seen between sM of 60 and 80 mC cm − 2, correlates to some degree with the decrease of the electrosorption valency observed for these charge densities in the inset to Fig. 13. However, we have mentioned earlier that calculations of the electrosorption valency for iodide may involve systematic errors. Therefore, in that case, the correlation between the change of the phase angle and the change of the electrosorption valency should be viewed with due caution. In contrast to our earlier conclusion [26], the present SHG data do not indicate that the adsorption of anions involves mixing of the electronic states of the adsorbate and the substrate, or charge transfer from the adsorbate to the metal. This discussion illustrates the difficulties in determining the degree of charge transfer involved in the formation of the electrosorption bond and confirms once more that the best way to describe the polarity of the electrosorption bond is to use the surface dipole.

4. Conclusions The adsorption of iodide at a Au(111) single crystal electrode has been investigated quantitatively using chronocoulometry. The experiments were performed in pH 11.0 solutions in order to extend the double layer region to a more negative potential. It was found that the adsorption of I − on the Au(111) surface is much stronger than that of OH − ions. By thermodynamically analyzing the charge density data of iodide adsorption, the Gibbs excess, Gibbs energy of adsorption, number of electrons flowing to the interface per one adsorbed iodide ion at a constant potential (electrosorption valency), the Esin–Markov coefficient, and the surface dipole were determined. The thermodynamic data for iodide adsorption were compared to the results for bromide and chloride adsorption. All three halides form a chemisorption bond with the gold surface. The bond is quite polar at the negatively charged surface, however its polarity drops significantly at the positively charged surface. At low charge densities and coverages the bond polarity is determined by the ability of free

A. Chen et al. / Journal of Electroanalytical Chemistry 467 (1999) 342–353

electrons to screen the dipole formed by the adsorbed anion and its image charge in the metal. At high charge densities and coverages, the chemisorption bond has a predominantly covalent character. The strength of the halide adsorption and the covalent character of the chemisorption bond increase by moving progressively from chloride to iodide.

Acknowledgements This work was supported by a grant from the Natural Sciences and Engineering Research Council of Canada. J.L. thanks Alexander von Humboldt-Stiftung for a Research Award.

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