Thermodynamic studies of bromide adsorption at the Pt(1 1 1) electrode surface perchloric acid solutions: Comparison with other anions

Journal of

Electroanalytical Chemistry Journal of Electroanalytical Chemistry 591 (2006) 149–158 www.elsevier.com/locate/jelechem

Thermodynamic studies of bromide adsorption at the Pt(1 1 1) electrode surface perchloric acid solutions: Comparison with other anions Nuria Garcia-Araez a, Victor Climent a, Enrique Herrero a, Juan Feliu Jacek Lipkowski a,b a

a,*

,

Departamento de Quı´mica Fı´sica, Universidad de Alicante, Facultad de Ciencias Apdo. 99, E-03080 Alicante, Spain b Department of Chemistry and Biochemistry, University of Guelph, Guelph, Ont., Canada N1G 2W1 Received 11 November 2005; received in revised form 23 March 2006; accepted 4 April 2006 Available online 24 May 2006

Abstract The thermodynamics of the so-called perfectly polarizable electrode was employed to analyze the voltammograms of a Pt(1 1 1) electrode in KBr solutions with an excess of a supporting electrolyte (0.1 M HClO4 and 0.1 M KClO4 + 103 M HClO4 + xM KBr where x varied between 5 · 104 and 1 · 102). The surface Gibbs excess, the Gibbs energy of adsorption and the charge number at a constant electrode potential and a constant chemical potential have been determined. The effect of pH on the magnitude of these parameters has been evaluated. The thermodynamic parameters for bromide have been compared to the parameters determined from previous thermodynamic studies of (bi)sulfate, chloride and OH adsorption at the Pt(1 1 1) electrode surface. As expected, bromide adsorption is stronger than for the other anions and halide adsorption seems to be limited by close packing of the adlayer. The calculated charge number values suggest that Br adsorption involves a full charge transfer, as in the case of chloride.  2006 Elsevier B.V. All rights reserved. Keywords: Pt(1 1 1) electrode; Bromide adsorption at Pt; Gibbs thermodynamics of a Pt electrode

1. Introduction This paper is a part of a broader project on thermodynamic studies of anions adsorption and their co-adsorption with cations at platinum single-crystal electrodes. Previous works from this series described adsorption of sulfate/ bisulfate at Pt(1 1 1) [1] and stepped surfaces [2], and hydrogen and OH [3], chloride [4] and chloride co-adsorption with hydrogen [5] at a Pt(1 1 1) surface. Adsorption of halides is a classical subject explored in specific adsorption studies on Hg, Ag and Au electrodes [6,7]. Here we report on bromide adsorption at the Pt(1 1 1) electrode and we compare it with the adsorption of sulfate/bisulfate, OH and chloride.

*

Corresponding author. Fax: +34 965 903 537. E-mail address: [email protected] (J. Feliu).

0022-0728/$ - see front matter  2006 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2006.04.008

Bromide adsorption at a Pt(1 1 1) surface was extensively studied by Hubbard’s group. Their experiments involved transfer of the sample from the electrochemical cell to UHV, where Auger spectroscopy and LEED were used to characterize the adlayers [8,9]. The structural information extracted from these studies was used for the interpretation of the characteristic features observed on a cyclic voltammogram (CV). In neutral or acidic solutions, the CV has two peaks: one at lower potentials whose envelope contains most of the charge, and a sharp spike that appears at higher potentials. While the main peak can easily be assigned to the adsorption of bromide and the consequent desorption of adsorbed hydrogen, it was suggested that the sharp spike corresponds to a phase transition from a (4 · 4) to a (3 · 3) hexagonal close packed Br adlayer, driven by the increase of the electrode potential and bromide coverage. In alkaline solutions, bromide adsorption is weakened by the compe-

150

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

tition with the adsorption of OH and the spike is no longer seen on the voltammograms at pH > 12. Itaya and co-workers studied bromide adsorption in situ using high resolution STM [10]. They found that Br forms an ordered (3 · 3) superlattice at potentials higher than the potential of the sharp spike but lower than the potential of the onset of OH adsorption. Both hexagonal and asymmetric type (3 · 3)-4Br structures were identified. However, the STM images obtained at potentials lower than the sharp spike revealed that the Br adlayer was not the (4 · 4) structure proposed in Ref. [9], but an incommensurate structure. Independently, Orts et al. studied the properties of strongly adsorbed bromine monolayers formed either by bromide or by bromine adsorption on a Pt(1 1 1) surface, using cyclic voltammetry, CO charge-displacement method and ex situ STM imaging [11]. From the charge displaced during CO adsorption at 0.50 V (RHE), a maximum Br coverage of 0.46 Br/Pt was calculated (in this calculation the double layer contributions to the measured charge were neglected). The same coverage was obtained regardless the adlayer was formed from adsorption of bromine or bromide. Ex situ STM images revealed an ordered uniaxially commensurate adlayer with preferential ordering along the directions of the substrate densest rows. This interpretation was consistent with independent in situ X-ray scattering studies by Lucas et al., who demonstrated that Br forms an hexagonal close-packed structure that is aligned but incommensurate, and that undergoes compression with increasing potential [12]. Gasteiger et al. applied the rotating ring-disk electrode with a Pt(1 1 1) single crystal in the disk position to determine the adsorption isotherm for bromide [13]. They found maximum bromide coverage of 0.42 Br/Pt. In addition, they determined that the electrosorption valency of bromide adsorption is essentially 1. In the present paper, we describe a thermodynamic method to study bromide adsorption on a Pt(1 1 1) electrode. For this system, we have determined the Gibbs excess, Gibbs energy of adsorption and charge numbers at a constant electrode potential and a constant chemical potential. We have performed measurements for two series of solutions. First, we used 0.1 M HClO4 as the supporting electrolyte. In these solutions, the substitution of hydrogen by bromide at the Pt(1 1 1) surface takes place in a very narrow potential range, as described previously [6–11]. The corresponding voltammetric peaks are very sharp, and the integration of such curves is subjected to a non-negligible error. Therefore, a second series of experiments was performed using solutions with a higher pH (pH  3), where the voltammetric peaks related to the competitive hydrogen desorption and bromide adsorption processes are broader. For these solutions, integration of the cyclic voltammetry curves can be performed with higher precision and the thermodynamic quantities can be more accurately determined. In addition, the comparison of both results gives insights on the effect of pH on bromide adsorption at the Pt(1 1 1) surface.

2. Experimental The Pt(1 1 1) working electrode was prepared from a single crystal platinum bead oriented, cut and polished with diamond paste down to 0.25 lm as described elsewhere [14]. Before each experiment, the electrode was annealed and cooled down in H2 + Ar reductive atmosphere (N50, Air Liquide) as described in [15]. Electrochemical experiments were performed at room temperature (22 ± 2 C) in a classical two-compartment electrochemical cell equipped with a large area platinum counter-electrode and a silver–silver chloride reference electrode in a KCl saturated solution, separated from the main compartment through a Luggin capillary. For experiments performed with solutions with 0.1 M KClO4, this reference electrode was substituted by a calomel reference electrode in CsCl saturated solution. By using CsCl instead of KCl, the precipitation of KClO4 is avoided at the liquid junction between the solution of the reference electrode and the investigated solution, where concentration of KClO4 is close to saturation. The solutions were de-aerated by purging with Ar (N-50, Air Liquide) and a blanket of argon was kept over the solution during the experiment. All potentials were converted to the standard hydrogen scale (SHE) and are reported here vs. SHE, unless otherwise stated. Water from an ELGA PURELAB ultra system (18.2 MX cm resistivity) was used to rinse the cell and to prepare the solutions. The solutions were prepared from KClO4 (Merck pro Analysis), HClO4 (Merck Suprapure) and KBr (Merck Suprapure) reagents. The electrochemical experiments were performed using the Autolab model PGSTAT 30 potentiostat controlled by a computer. The numerical treatment of the data has been performed with software packages Microcal Origin 6.1 and Mathcad Professional 2001. 3. Results 3.1. Cyclic voltammetry The cyclic voltammetry curves (CVs) of a Pt(1 1 1) electrode were measured in 0.1 M HClO4 solutions, pure and with eight additions of KBr. The sweep rate was selected slow enough in order to achieve almost symmetrical curves with respect to the potential axis. This assures that the adsorption equilibrium was reached at each potential on the time scale of the experiment. Sweep rates as low as 10–1 mV/s for bromide containing solutions and 50 mV/s for the pure perchloric acid solution were used. The concentration of KBr was progressively increased by spiking the 0.1 M HClO4 supporting electrolyte with a concentrated solution of KBr. Differential pseudocapacity curves were calculated from the CVs, by dividing the voltammetric current j by the sweep rate m (C = j/m), and some representative examples are shown in Fig. 1A. Similar measurements have been repeated for solutions with a higher pH (pH 3.11), using 0.1 M KClO4 + 103 M

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158 6000

151

4000

5000

3000 4000

2000

3000

-2

(C = j / ν ) /μF cm

(C = j / ν) /μF cm

-2

2000 1000 0 -1000

1000

0

-1000

-2000

-2000 -2

60

-4000

j / μA cm

j / μA cm

-2

-3000

50

-3000

0

-50

-5000

-60

-4000 0.0

-6000

0.1

0.2

0

-0.2

0.3

0.0

0.2

0.4

0.6

0.8

E /V vs SHE

-0.1

0.0

0.1

E /V vs SHE

E /V vs SHE

-5000 -0.2

0.0

0.2

0.4

0.6

0.8

E /V vs SHE

Fig. 1A. Differential capacity calculated from cyclic voltammograms recorded at the Pt(1 1 1) electrode in 0.1 M HClO4 + xM KBr where x = 0, 5 · 104, 1.5 · 103, 3.75 · 103 and 7.5 · 103, using a scan rate ranging from 1 to 10 mV/s (50 mV/s for the pure supporting electrolyte). Arrows indicate directions in which the current increases with the bromide concentration. The inset shows the overlapping of the voltammograms in the low potential region when the scan rate is 50 mV/s.

Fig. 1B. Differential capacity calculated from cyclic voltammograms recorded at the Pt(1 1 1) electrode in 0.1 M KClO4 + 103 M HClO4 + xM KBr, where x = 0, 4 · 104, 7.9 · 104, 2 · 103 and 6 · 103, using scan rate of 10 mV/s (50 mV/s for the pure supporting electrolyte). Arrows indicate directions in which the current increases with the bromide concentration. The inset shows the overlapping of the voltammograms in the low potential region when the scan rate is 50 mV/s.

HClO4 + xM KBr where x varied between 5 · 104 and 1 · 102 electrolyte, with nine additions of KBr. Fig. 1B plots selected differential capacity curves from this series of experiments. The increase of pH shifts hydrogen and OH adsorption regions towards lower potentials. In contrast, Br adsorption is affected by pH only indirectly as a result of the competition or co-adsorption with adsorbed hydrogen or OH. The differential capacity curves for the pure 0.1 M HClO4 and 0.1 M KClO4 + 103 M HClO4 solutions display a sharp peak at E  0.72 V (pH  1) and E  0.60 V (pH  3), and a flat current region at E < 0.2 V, indicating that the Pt(1 1 1) surface was almost defect free [14]. The differential capacity curves for the bromide containing solutions display a high and narrow peak at E  0.1 V (pH  1) and E  0.0 V (pH  3), followed by a sharp spike at E  0.2 V. The fact that the sharp spike at E  0.2 V is situated nearly at the same potential for both pH’s suggests that hydrogen adsorption is significantly suppressed (hydrogen coverage is reduced to a negligible value) by adsorbed bromide at this potential. In situ STM images indicated that this sharp spike is due to a phase transition from an incommensurate Br adlayer to the (3 · 3)-4Br structure [10]. Since that feature is almost independent on pH, this transition does not involve displacement of adsorbed hydrogen as already noted in [9].

For the 0.1 M HClO4 + 102 M KBr solution, the charge density calculated by integration of the differential capacity curve between 0.01 and 0.25 V amounts to 274 lC cm2, in good agreement with the value 269 lC cm2 reported elsewhere [11,13]. The CO-displacement experiments showed that 40% of that charge corresponds to bromide adsorption, while the remaining 60% is due to the adsorption of 2/3 monolayers of hydrogen [11]. The capacitance peaks are smaller and broader in solutions with pH  3, because competition with hydrogen adsorption is smaller than at pH  1. One assumption required to apply the thermodynamic analysis described in the next section is the existence of a potential, E*, where the charge state of the interphase is independent of the anion concentration. This assumption is fulfilled by selecting E* low enough in order to guarantee that specifically adsorbed Br is completely desorbed; the excess of supporting electrolyte assures that the amount of not-specifically adsorbed Br is negligible as well. The coincidence of the differential capacities obtained without and with different concentrations of an anion is usually taken as a proof that the anion is not adsorbed at these low potentials. For the differential capacity curves in Fig. 1A and B, this overlap is masked at first sight, because of a very small reduction current due to the presence of

152

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

residual O2. This interference is more significant when slow scan rates are used to record the voltammograms. However, this current has no effect on the results of the thermodynamic analysis because it is cancelled when the positive and negative scans are averaged. On the other hand, the voltammograms measured with a scan rate of 50 mV/s (inset of Fig. 1A and B), overlap nicely at E < 0.05 for pH  1 and E < 0.07 for pH  3, because in this case, the contribution of the reduction current due to residual O2 is comparatively much smaller. With the help of these figures, we have chosen the bromide-concentration independent potential of E* = 0.05 V for the experiments performed at pH  1, and E* = 0.07 V for those of pH  3. 3.2. Gibbs excess data The electrocapillary equation at constant temperature for a Pt electrode in bromide -containing solutions is described by [16–19]:

size that the knowledge of the charges Q is not necessary because p can also be determined by integration of the E vs. DQ curves if one can identify a potential E* where bromide is completely desorbed and hence the charge density is independent of the bromide concentration [20]: Z Q Z DQ p ¼ nh¼0  nh ¼ ðEh¼0  Eh Þ dQ ¼ ðEh¼0  Eh Þ dDQ 0

Q

ð5Þ Q *,

Q*

E*

where DQ = Q  is the charge at and subscripts h = 0 and h denote the values of the Parsons function and the potential measured in the pure supporting electrolyte and the bromide containing solutions, respectively. The total charge density Q = DQ + Q* plots, calculated from the integration of the cyclic voltammograms, are shown in Fig. 2A (pH  1) and B (pH  3). As discussed in the previous section, E* = 0.05 V for the experiments performed at pH  1, and E* = 0.07 V for those of pH  3. Although only charge differences are needed in

 dc ¼ Q dE þ CBr RT d ln aBr þ ðCH  COH ÞRT d ln aHþ ð1Þ 150

-2

50 0

ð2Þ

-50

When experiments are carried out in solutions of a constant pH and a constant ionic strength, Eq. (1) can be simplified to:

-100

(In a solution of a constant ionic strength the activity coefficient is constant and hence d ln aBr ¼ d ln cBr .) Eq. (3) shows that the Gibbs excess CBr can be determined by differentiation of c with respect to RT ln cBr at a constant potential. Alternatively, one can use the Parsons function, n = Q Æ E + c, to calculate the Gibbs excess when charge is considered as the independent electrical variable [19]. Previous studies demonstrated that the Gibbs excesses calculated at constant charge and a constant potential are not too different, but a somewhat better fit and a smaller spread of the data are observed when charge is used as the independent variable [4]. In that case the electrocapillary equation is: dn ¼ E dQ þ CBr RT d ln cBr

-150 150

ð3Þ

ð4Þ

Eq. (4) shows that the Gibbs excess CBr can be determined by differentiation of n with respect to RT ln cBr at a constant charge. The absolute value of the Parsons function is not known. However, one can calculate the surface pressure of adsorbed anion at a constant charge, p = nh=0  nh, by integration of the E vs. Q curves. It is useful to empha-

B

100 -2

dc ¼ Q dE þ CBr RT d ln cBr

Q / μC cm

Q ¼ r  F CH þ F CBr þ F COH

A

100

Q / μC cm

where c is the surface energy; Q is the total charge density at the electrode surface; CBr, CH and COH are the relative Gibbs excesses of specifically adsorbed bromide, hydrogen and OH species; and aBr and aHþ are the bromide and proton activities. The total charge density includes both the free charge (r) and the charge transferred during the adsorption processes:

50 0 -50 -100 -150 -0.1

0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

E /V vs SHE Fig. 2. (A and B) The total charge density of Pt(1 1 1) plotted against the electrode potential, calculated from the integration of the voltammograms obtained in 0.1 M HClO4 + xM KBr where x = 0, 5 · 104, 103, 1.5 · 103, 2.5 · 103, 3.75 · 103, 5 · 103, 7.5 · 103 and 102 (A) and in 0.1 M KClO4 + 103 M HClO4 + x M KBr, where x = 0, 4 · 104, 6 · 104, 7.9 · 104, 1.2 · 103, 2 · 103, 3 · 103, 4 · 103, 6 · 103 and 8 · 103 (B), assuming that the charge at E* = 0.05 V (A) and E* = 0.07 V (B) is bromide concentration independent. Arrows indicate the direction of increasing bromide concentration.

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

450

400

75 μC cm

350

-2

/ ions cm-2

-2

6 5

-14

100 μC cm

The Gibbs excesses of bromide are plotted against the total charge on the Pt(1 1 1) surface in Fig. 4A (pH  1) and B (pH  3). The curves in these two figures display similar features. At low charge densities the Gibbs excesses depend on the bulk bromide concentration. This region corresponds to the lower potential voltammetric peak, responsible for the first step on the charge potential plots in Fig. 2A and B. The dependence decreases with charge and when Q > 88 lC cm2 (pH  1) and for Q > 51 lC cm2 (pH  3) the curves corresponding to different bulk bromide concentrations merge into one line. This is a consequence of the degree of the polynomial used to calop culate ðRT o ln Þ , although we cannot rule out a small cBr Q dependence below the experimental error at high charge densities. In Fig. 5A and B, the same sets of data are plotted against the electrode potential. To construct the Gibbs excess vs. potential plots, for each bulk bromide concentration, a C vs. Q plot from Fig. 4 was combined with a corresponding Q vs. E plot shown in Fig. 2. The C vs. E plots display two steps. The first has a limiting value of C  (5.0 ± 0.1) · 1014 ions/cm2. It is followed by a second,

3

Γ · 10

order to apply the previously described analysis, to facilitate a comparison of the results of the present study with other data, values of Q are presented in Fig. 2, calculated from the values of the potential of zero total charge (pztc) (Epztc = 0.256 V (SHE) for 0.1 M HClO4 and Epztc = 0.162 V (SHE) for 0.1 M KClO4 + 103 M HClO4 solutions [21]). Fig. 3 plots p vs. RT ln cBr for selected charge densities for experiments performed at pH  1. The points represent experimental data and the lines show the fit of the data to a polynomial. A second order polynomial was used when the plot had a significant curvature. However, when the curvature was smaller than the experimental error, the use of a second order polynomial only introduced a higher error, especially for the points corresponding to the higher and lower concentration values. In those cases, we chose to use a first order polynomial. For charge densities of less than 88 lC cm2 (pH  1) and 51 lC cm2 (pH  3) the data fit well to a second order polynomial. For higher charge densities a first order polynomial gives the best fit. In the previous study of chloride adsorption on Pt(1 1 1) in 0.1 M HClO4 [4], we also found that the best fit to the data was obtained with the first order polynomial for Q > 88 lC cm2 and the second order polynomial for Q < 88 lC cm2. Once the polynomial is determined, it is differentiated to calculate the Gibbs excess.

4

2

0

250

6

-2

-2

50 μC cm

200

150

100

0 μC cm

50

0

Γ · 10

-14

25 μC cm

-2

/ ions cm

-1

A

1

300

π / mN m

153

-2

-25 μC cm

-2

-50 μC cm

-2

B

5 4 3 2 1 0

-19

-18

-17

-16

-15

-14

-13

-12

-11

-1

RT lncBr- / kJ mol

Fig. 3. Surface pressure, p, plotted against RT ln cBr for selected total charge density values using the measurements performed in 0.1 M HClO4 + xM KBr solutions. Points are the experimental data and lines shows the fit to a second order polynomial (for Q < 88 lC cm2) and to a straight line (for Q > 88 lC cm2).

-100

-50

0

Q / μC cm

50

100

-2

Fig. 4. (A and B) The Gibbs excesses of bromide calculated at constant charge density for 0.1 M HClO4 + xM KBr (A), and 0.1 M KClO4 + 103 M HClO4 + xM KBr (B), where x values are the same as in Fig. 2A and B. Arrows indicate the direction of increasing bromide concentration.

154

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

Γ · 10

-14

/ ions cm

-2

6

A

5 4 3 2 1 0

Γ · 10

-14

/ ions cm

-2

6

B

5 4 3 2 1 0 -0.1

0.0

0.1

0.2

0.3

0.4

0.5

On the other hand, the present result is also 10% smaller than the maximum bromide coverage of 0.44 ML determined by the CO displacement method at E = 0.43 V [11]. However, the agreement between charges measured in Ref. [11] and in the present work is very good. The difference between the bromide coverages may be due to the fact that the estimation from the CO displaced charge requires a correction for the difference in the free charge density of the electrode, covered by CO and free of CO. According to Weaver [22], the CO-covered electrode is negatively charged at E = 0.43 V and taking into account this correction, the bromide coverage is decreased to 0.40 ML. Finally, the present result agrees within the limits of the experimental error with the coverage of 0.38 ± 0.02 ML determined at E = 0.44 V, by the rotating disk electrode technique [13]. Overall, the agreement between the maximum Br coverage determined by various methods is within 10%. For 103 M solution of a specifically adsorbed anion in 0.1 M HClO4, the Gibbs excess-potential plot for bromide is compared to the Gibbs excess plot for chloride [4] and (bi)sulfate [1] in Fig. 6. The Gibbs excess plots for hydrogen and OH determined in a pure 0.1 M HClO4 [3] solution are also included in this figure. The potential region in which the adsorption of each anion takes place increases in the order: bromide, chloride, (bi)sulfate and OH. All curves display a two-step (two state) character of the anion

E /V vs SHE Fig. 5. (A and B) The Gibbs excesses of bromide shown in Fig. 4A and B, plotted as a function of potential. Arrows indicate the direction of increasing bromide concentration.

8

Cl 7

Br

H

OH

5

4

-14

/ ions cm

-2

6

Γ · 10

smaller but more abrupt step to the maximum value of C  (6.0 ± 0.1) · 1014 ions/cm2. A comparison of Fig. 5A and B shows that the first step is steeper in solutions of pH  1. In this case, the adsorption of bromide starts at a higher potential due to a stronger competition with hydrogen adsorption than at pH  3, but it ends at approximately the same potential. The maximum Gibbs excess of bromide is equivalent to 0.40 ± 0.01 ML coverage of the Pt(1 1 1) surface by adsorbed anions (where one monolayer (ML) corresponds to the surface concentration of Pt atoms at an ideal Pt(1 1 1) surface equal to 1.5 · 1015 atoms/cm2). This value is about 10% smaller than the theoretical value of 0.44 ML corresponding to the hexagonal (3 · 3)-4Br structure observed by LEED [8,9] and STM experiments [10,11]. It also is 7.5% smaller than the coverage 0.43 ML calculated from the structure determined by SXS experiments at E = 0.44 V [12]. It should be stressed, however, that the coverage calculated from the dimensions of a unit cell of a two-dimensional structure is based on the assumption that the surface is covered by a perfect, defect free adlayer but in reality the adlayer is not perfect. Indeed, in situ STM images revealed not only the presence of defects, but also the co-existence of two different (3 · 3)-4Br structures: the hexagonal and the asymmetric one [10].

3

SO4 2

1

0

-1 0.0

0.2

0.4

0.6

0.8

1.0

E /V vs SHE Fig. 6. Comparison of Gibbs excesses of bromide (solid line), chloride (dashed line [4]) and (bi)sulfate (dotted line [1]) in 0.1 M HClO4 with 103 M of the anion. The hydrogen and OH Gibbs excesses (dotteddashed line [3]) calculated in pure 0.1 M HClO4 are included for comparison.

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

3.3. Charge numbers per adsorbed anion Cross differentiation of the electrocapillary equation gives two charge numbers related to the number of electrons flowing to the electrode per one adsorbed anion [26]:       1 oQ 1 ol 1 oDG l¼ ¼ ¼ ð6Þ F oC E F oE C F oE C and

    1 oQ 1 ol n ¼ ¼ F oC l F oE Q 0

ð7Þ

where l is the charge number at a constant electrode potential and it is usually known as electrosorption valency (IUPAC recommends the name ‘‘formal partial charge number’’ [26]) and n 0 , the charge number at a constant chemical potential, is the reciprocal of the Esin–Markov coefficient. In Eqs. (6) and (7), Q is the total charge density, C is the Gibbs excess of adsorbed anion and l is the chemical potential of the anion in the bulk of the electrolyte. Fig. 7 plots the total charge density as a function of the Gibbs excess for selected electrode potentials using the data obtained at pH  1. The plots are linear and their slopes give the charge numbers at constant E. By including the charge values measured in the pure supporting electrolyte (when CBr = 0), the uncertainty in the determination of the slope is reduced. These charge numbers and those calculated from the data obtained at pH  3, are plotted as a

140 120

0.2V

100

0.175V

80

0.15V

-2

60

Q / μC cm

adsorption at the Pt(1 1 1) surface. The curves for bromide, (bi)sulfate and OH are steeper than for chloride. It should be pointed out that this behavior cannot be directly related to differences between lateral interaction between adsorbed particles, because it can be dramatically influenced by coadsorption phenomena. For instance, Monte Carlo simulations have demonstrated that the sharpness of the voltammetric profiles of a Pt(1 0 0) electrode in bromide containing acidic solutions is due to the competition between bromide and hydrogen for the surface sites and it is not related to attractive lateral neighbor interactions [23]. On the other hand, the maximum Gibbs excess is similar for chloride [(7.3 ± 0.16) · 1014 ions/cm2, equivalent to a coverage of 0.49 ± 0.01] and OH [(7.5 ± 0.14) · 1014 ions/cm2, h = 0.50 ± 0.01], while for bromide it is (6.0 ± 0.1) · 1014 ions/cm2 (h = 0.40 ± 0.01) and for (bi)sulfate (3.0 ± 0.1) · 1014 ions/cm2 (h = 0.20 ± 0.01). The bigger size of bromine atoms compared to chlorine can explain the difference between the maximum coverage of these ions. The limiting surface concentrations of Br and Cl expected for a closed packed hexagonal monolayer of these ions, calculated from the van der Waals radii of ˚ and rCl = 1.80 A ˚ [24]), are CBr  these anions (rBr = 1.95 A 7.59 · 1014 ions/cm2 (hBr  0.506) and CCl  8.91 · 1014 ions/cm2 (hCl  0.594). The small maximum coverage ofpffiffiffi(bi)sulfate is due to the formation of an open pffiffiffi ð 3x 7ÞR19:10 ordered adlayer, identified from STM images [25].

155

40 20

0.125V 0 -20 -40 -60

0.1V

-80

0.075V

-100 0

1

2

Γ 10

3

4

-14

/ ions cm

5

6

7

-2

Fig. 7. Plot of the charge density vs. Gibbs excess of adsorbed bromide at a constant electrode potential in 0.1 M HClO4 + xM KBr solutions.

function of the electrode potential in Fig. 8. For comparison, the charge numbers l of chloride [4], (bi)sulfate [1] and OH [3] are also included in this figure. The values much lower than 1 are observed at the lower potential region. Since the oxidation state of adsorbed bromide may vary between 0 and 1 only, the values of charge number lower than 1 can be explained by the effect of hydrogen on the adsorption of the specifically adsorbed anion:   oQ Fl ¼ oCBr E       oQ oQ oCH ¼ þ oCBr E;CH oCH E;CBr oCBr E   oCH ¼ F ðlBr ÞCH  F ðlH ÞCBr ð8Þ oCBr E Since lHþ  1 and lBr  1, at potentials where hydrogen adsorption takes place, the values of l < 1 indicate that oCH ðoC Þ < 0. [For example the value of l  3 suggests that Br E oCH ðoC Þ  2 or that two hydrogen atoms are displaced Br E from the electrode surface by adsorption of one Br atom. This is a reasonable number.] In turn, this condition implies that the adsorption of hydrogen and bromide has a competitive character. The steeper decrease of the charge number value at pH  1, compared to pH  3, is due to the higher overlap between bromide adsorption and hydrogen desorption processes at this pH. When the potential increases, hydrogen is desorbed from the surface and the charge number l increases up to lBr  1 at E = 0.2 V.

156

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

positive than for bromide, and the overlap with hydrogen adsorption is less significant in that case. As a result, lCl  1 in a relatively broad range of potentials 0.3 6 E 6 0.65 V. Finally, the charge number l of (bi)sulfate and OH also depend on the potential. However, these adsorption processes occur at too positive potentials to be influenced by adsorption of hydrogen. Therefore, the explanation of this dependence should be based on double-layer effects. The charge number at constant chemical potential defined by Eq. (7) (inverse of the Esin–Markov coefficient) may be calculated by differentiation of the C vs. Q plots in Fig. 4A and B, and from the slope of E vs. RT ln cBr at constant Q. The two independently determined derivatives (oC/oQ)l and (ol/oE)Q are plotted against each other in the inset of Fig. 9. The experimental data are scattered randomly along the line with a unity slope indicating that the values of C are free of major data-processing errors. The charge numbers n 0 of bromide, chloride [4], (bi)sulfate [1] and OH [3], are plotted against electrode potential in Fig. 9. The results are in very good agreement with those shown in Fig. 8.

Cl -1.0

OH

l / electron per ion

-1.5

-2.0

Br

SO4

-2.5

-3.0

-3.5

-4.0

-4.5

-5.0 0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

E /V vs SHE Fig. 8. Electrosorption valencies of bromide (pH  1: black circles, pH  3: black squares), chloride (open up-triangles [4]), sulfate (grey down-triangles [1]) and OH (open diamonds [3]) adsorption determined from the slope of the difference of charge vs. Gibbs excess plots at constant potential.

where z = 1 is the charge of the bromide ion, e0 is the permittivity of vacuum, er is the relative permittivity of the media and QC is the capacity at the constant amount adsorbed. The surface dipole moment is a measure of the polarity of the electrosorption bond. However, its magnitude depends also on the space charge and dipole–dipole coupling effects. A detailed discussion of effects contributing to the magnitude of the surface dipole moment is given in [28,31]. In the present case, the value of lBr  1 indicates that the surface dipole is close to zero and that adsorption of Br at Pt(1 1 1) involves a significant redistribution of charge. Effectively, the adsorbed species may be considered as adsorbed bromine atoms. Similar explanation may be applied to discuss the changes of the charge number l for chloride ion. The values of l lower than 1 at the lower potential region indicate competitive hydrogen adsorption. The potential window at which chloride adsorption takes place is more

Br -2.0

-1

SO4

-2.5

6 5

12

ð9Þ

OH

-1.5

(dΓ / dQ )μ · 10 / mol μC

ze0 er ð1  lBr =zÞ ls ¼ QC

-1.0

n' / electron per ion

The charge number lBr may be used to calculate the surface dipole moment ls, which represents the dipole formed by an adsorbed anion and its image charge in the metal, using the formula [27,28]:

Cl

-3.0

-3.5

4 3 2 2

3

4

5

6

12

-(dE / dμ )Q · 10 / mol μC 0.0

0.1

0.2

0.3

0.4

0.5

0.6

-1

0.7

E /V vs SHE Fig. 9. Charge numbers at constant chemical potential (or inverse of the Esin–Markov coefficient), of bromide (pH  1: black circles, pH  3: black squares), chloride (open up-triangles [4]), sulfate (grey downtriangles [1]) and OH (open diamonds [3]), plotted against the electrode potential in 0.1 M HClO4. In the inset, the comparison of bromide Esin– Markov coefficients determined from the slope of E vs. l plots and from the slope of the CBr vs. Q plot at a constant bulk bromide concentration is shown.

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

3.4. Standard Gibbs energies of adsorption The standard Gibbs energies of adsorption are usually determined from a fit of the Gibbs excess or surface pressure data to the equation of an adsorption isotherm. In the case of specific anion adsorption the surface pressure data are usually fitted to a ‘‘square root’’ isotherm [29]: ln ðkTcBr Þ þ ln b ¼ ln p þ Bp1=2

ð10Þ

DGo

where b ¼ expð kT Þ is the adsorption equilibrium constant, B is a constant and p is the surface pressure. Fig. 10 shows plots of the square-root of the surface pressure vs. lnðkTcBr =pÞ at a constant charge with the data obtained at pH  1. The plots are fairly linear and the extrapolation to zero surface pressure give an intercept with the lnðkTcBr =pÞ axis equal to: lim ½ln ðkTcBr =pÞ ¼  ln b

ð11Þ

p!0

from which the standard Gibbs energies of adsorption can be calculated. The standard state is an ‘‘ideal’’ C = 1 molecule cm2 for the surface species and an ‘‘ideal’’ cBr ¼ 1 mol dm3 for the bulk species [29]. Note that the bromide concentrations are multiplied by the kT term so that, in the limit of low coverage, the film pressure is described by Henry’s law p ¼ kT bcBr as explained in [29]. The standard Gibbs energies of bromide adsorption have been calculated at constant potential as well as at constant charge for both pH’s. Then, the results have been

157

plotted against potential in Fig. 11. The charge vs. potential plot for the corresponding pH, bromide-free, solution was used to convert the charge densities to potentials. The agreement between DG values determined at both pH’s from the constant potential and the constant charge analyses is quite good, with the exception of the data for solution at pH  1 analyzed at constant charge. It indicates that the DG data may be considered free from major errors, in spite of the long extrapolation procedure that was necessary to use. For comparison, the standard Gibbs energies of chloride [4] and (bi)sulfate [1] adsorption are also included in Fig. 11. The adsorption of all three anions starts at the potential in which the standard Gibbs energy of adsorption is about 100 kJ/mol. This large absolute value is due to the selected standard state, and corresponds to about eads  25 kJ/mol when the standard state is one particle in the solution, and one adsorbed particle at the electrode surface [30]:   Cs DG0 ¼ eads  kT ln ð12Þ cs l where Cs is the number of particles in the adsorbed standard state (Cs = 1 molecule cm2), cs is the number of particles in the solution standard state (cs = 1 mol dm3), and l is the thickness of the inner layer, assumed to be the van ˚ ). der Waals diameter of Br (l = 3.9 A

Cl 7.0

100 μC cm

-100

-2

6.5

SO4

6.0

75 μC cm

-2

-120

50 μC cm

-2

Δ G ˚ / kJ mol

5.0

-1

5.5

4.0

25 μC cm

3.5 3.0

-2

0 μC cm

π

1/2

1/2

/ μJ cm

-1

4.5

-140

Br -160

-2

2.5

-25 μC cm

2.0

-2

-50 μC cm

1.5 1.0

-180

-2

-75 μC cm

-2

0.5

0.0

0.2

0.4

0.6

0.8

E /V vs SHE -50

-49

-48

-47

-46

-45

-44

ln(kTcKBr /π) Fig. 10. Fit of the adsorption data to the square root isotherm at constant total charge in 0.1 M HClO4 + xM KBr solutions.

Fig. 11. Plot of the Gibbs energy of bromide (closed symbols), chloride (open symbols [4]) and (bi)sulfate (grey symbols [1]) adsorption at Pt(1 1 1) vs. electrode potential determined from the p1/2 vs. ln(kTcanion/p). Plots at constant potential (pH  1: circles, pH  3: stars) and at constant charge (pH  1: squares, pH  3: triangles).

158

N. Garcia-Araez et al. / Journal of Electroanalytical Chemistry 591 (2006) 149–158

The slope of the DG vs. E plots reflects the average value of the charge number at a constant E (see Eq. (6)), which is of 1.8 electrons per ion for bromide (that has important contributions from competitive hydrogen adsorption), 1.0 for chloride and 1.6 for bi(sulfate) adsorption. Finally, the standard Gibbs energies of adsorption are more negative for bromide than for the two other anions indicating a stronger bond between this anion and the platinum surface. 4. Conclusions Bromide adsorption at a Pt(1 1 1) electrode in 0.1 M HClO4 and 0.1 M KClO4 + 103 M HClO4 solutions has been studied by a thermodynamic method. The surface Gibbs excess, Gibbs energy of adsorption and charge numbers at a constant electrode potential and a constant chemical potential were determined. The maximum Gibbs excess of adsorbed Br was found to be (6.0 ± 0.1) · 1014 ions/cm2 (0.40 ± 0.01 Br/Pt), for both pH’s, in agreement with previous studies [8–11,13]. At low potentials, the charge numbers are much lower than 1 as a result of the competition with hydrogen adsorption. However, they approach 1 at sufficiently positive potential where hydrogen adsorption does not take place, indicating that the adsorbed species are almost neutral. In the potential region in which the sharp spike related to the formation of the (3 · 3)-4Br structure is seen on CVs, the charge numbers are very close to 1. This number suggests that this phase transition involves only bromide ions and does not involve displacement of hydrogen. The Gibbs energies of adsorption at constant potential are pH independent. They are much lower than Gibbs energies of adsorption of chloride and (bi)sulfate. This behavior reflects a stronger bond of bromide with the Pt(1 1 1) surface. Acknowledgements Financial support from the MCyT (Spain) through project BQU2003-04029 is gratefully acknowledged. J.L. acknowledges Natural Sciences and Engineering Council of Canada for a financial support and Canada Foundation of Innovation for the Canada Research Chair Award. N.G. thanks the MECD (Spain) for the award of a FPU grant. V.C. acknowledges financial support from the MEC under the Ramon y Cajal program.

References [1] E. Herrero, J. Mostany, J.M. Feliu, J. Lipkowski, J. Electroanal. Chem. 534 (2002) 79. [2] J. Mostany, E. Herrero, J.M. Feliu, J. Lipkowski, J. Phys. Chem. B 106 (2002) 12787. [3] J. Mostany, E. Herrero, J.M. Feliu, J. Lipkowski, J. Electroanal. Chem. 558 (2003) 19. [4] N. Garcia-Araez, V. Climent, E. Herrero, J. Feliu, J. Lipkowski, J. Electroanal. Chem. 576 (2005) 33. [5] N. Garcia-Araez, V. Climent, E. Herrero, J. Feliu, J. Lipkowski, J. Electroanal. Chem. 582 (2005) 76. [6] D.C. Grahame, Chem. Rev. 41 (1947) 441. [7] A. Hamelin, in: B.E. Conway, R.E. White, J.O. Bockris (Eds.), Modern Aspects of Electrochemistry, vol. 16, Plenum Press, New York, 1985, pp. 1–101 (Chapter 1). [8] J.L. Stickney, S.D. Rosasco, G.N. Salaita, A.T. Hubbard, Langmuir 1 (1985) 66. [9] G.N. Salaita, D.A. Stern, F. Lu, H. Baltruschat, B.C. Schardt, J.L. Stickney, M.P. Soriaga, D.G. Frank, A.T. Hubbard, Langmuir 2 (1986) 828. [10] S. Tanaka, S.L. Yau, K. Itaya, J. Electroanal. Chem. 396 (1995) 125. [11] J.M. Orts, R. Go´mez, J.M. Feliu, A. Aldaz, J. Clavilier, J. Phys. Chem. 100 (1996) 2334. [12] C.A. Lucas, N.M. Markovic, P.N. Ross, Surf. Sci. 340 (1995) L949– L954. [13] H.A. Gasteiger, N.M. Markovic, P.N. Ross, Langmuir 12 (1996) 1414. [14] J. Clavilier, D. Armand, S.-G. Sun, M. Petit, J. Electroanal. Chem. 205 (1986) 267. [15] A. Rodes, K. El Achi, M.A. Zamakhchari, J. Clavilier, J. Electroanal. Chem. 284 (1990) 245. [16] A.N. Frumkin, O.A. Petrii, Electrochim. Acta 20 (1975) 347. [17] H.D. Hurwitz, J. Electroanal. Chem. 10 (1965) 35. [18] E. Dutkiewicz, R. Parsons, J. Electroanal. Chem. 11 (1966) 100. [19] R. Parsons, Proc. R. Soc. Lond. Ser. A 261 (1961) 79. [20] W. Savich, S.G. Sun, J. Lipkowski, A. Wieckowski, J. Electroanal. Chem. 388 (1995) 233. [21] N. Garcia-Araez, V. Climent, E. Herrero, J. Feliu, J. Lipkowski, Electrochim. Acta 51 (2006) 3787. [22] M.J. Weaver, Langmuir 14 (1998) 3932. [23] N. Garcia-Araez, J.J. Lukkien, M.T.M. Koper, J.M. Feliu, J. Electroanal. Chem. 588 (2006) 1. [24] D.R. Lide (Ed.), Handbook of Chemistry and Physics, 74th ed., CRC Press, Boca Raton, FL, 1993. [25] A.M. Funtikov, U. Stimming, R. Vogel, J. Electroanal. Chem. 428 (1997) 147. [26] S. Trasatti, R. Parsons, Pure Appl. Chem. 58 (1986) 437. [27] K. Bange, B. Straehler, J.K. Sass, R. Parsons, J. Electroanal. Chem. 229 (1987) 87. [28] W. Schmickler, J. Electroanal. Chem. 249 (1988) 25. [29] R. Parsons, Trans. Faraday Soc. 51 (1955) 1518. [30] M.T.M. Koper, J. Electroanal. Chem. 450 (1998) 189. [31] R. Guidelli, W. Schmickler, in: B.E. Conway, C.G. Vayenas, R.E. White, M.E. Gamboa-Adelco (Eds.), Modern Aspects of Electrochemistry, vol. 38, Kluwer, New York, 2005, p. 303.