Anodic Passivation of Tin by Alkanethiol Self ... - ACS Publications

Article pubs.acs.org/Langmuir

Anodic Passivation of Tin by Alkanethiol Self-Assembled Monolayers Examined by Cyclic Voltammetry and Coulometry Barrett C. Worley,† William A. Ricks,† Michael P. Prendergast,† Brian W. Gregory,*,† Ross Collins,‡ John J. Cassimus, Jr.,‡ and Raymond G. Thompson*,‡ †

Department of Chemistry & Biochemistry, Samford University, Birmingham, Alabama 35229-2236, United States Vista Engineering & Consulting, LLC, Birmingham, Alabama 35203, United States

‡

S Supporting Information *

ABSTRACT: The self-assembly of medium chain length alkanethiol monolayers on polycrystalline Sn electrodes has been investigated by cyclic voltammetry and coulometry. These studies have been performed in order to ascertain the conditions under which their oxidative deposition can be achieved directly on the oxide-free Sn surface, and the extent to which these electrochemically prepared self-assembled monolayers (SAMs) act as barriers to surface oxide growth. This work has shown that the potentials for their oxidative deposition are more cathodic (by 100−200 mV) than those for Sn surface oxidation and that the passivating abilities of these SAMs improve with increasing film thickness (or chain length). Oxidative desorption potentials for these films have been observed to shift more positively, and in a highly linear fashion, with increasing film thickness (∼75 mV/CH2). Although reductive desorption potentials for the SAMs are in close proximity to those for reduction of the surface oxide (SnOx), little or no SnOx formation occurs unless the potential is made sufficiently anodic that the monolayers start to be removed oxidatively. Our coulometric data indicate that the charge involved with alkanethiol reductive desorption or oxidative deposition is consistent with the formation of a close-packed monolayer, given uncertainties attributable to surface roughness and heterogeneity phenomena. These experiments also reveal that the quantity of charge passed during oxidative desorption is significantly larger than what would be predicted for simple alkylsulfinate or alkylsulfonate formation, suggesting that oxidative removal involves a more complex oxidation mechanism. Analogous chronocoulometric experiments for short-chain alkanethiols on polycrystalline Au electrodes have evidenced similar oxidative charge densities. This implies that the mechanism for oxidative desorption on both surfaces may be very similar, despite the significant differences in the inherent dissolution characteristics of the two materials at the anodic potentials employed.

■

INTRODUCTION Tin metal finds widespread use as a plating material for protecting a wide variety of metallic surfaces. In the food processing industry, tinplate has been used extensively for coating food processing equipment, kitchen utensils, and storage materials for fruits, vegetables, meat, and dairy products. The most obvious examples of the latter are “tin cans”, which are actually produced from sheet steel and are coated with a thin layer of tinplate. In the electronics industry, the solderability of electronic components and their metal contacts on circuit boards is maintained by a layer of electroplated tin, which prevents oxidation of the substrate metal onto which it is deposited. The high degree of corrosion resistance exhibited by tin − coupled with its malleability/ductility, its ease of solderability, and its nontoxicity − have made tin metal a very attractive candidate for many applications. The passivation characteristics displayed by Sn arise from its ability to form a layer that is compact and insulating (both ionically and electronically) upon oxidation in air, or during anodization in aqueous solution at pHs near neutrality. The © 2013 American Chemical Society

resulting oxide/hydroxide surface layer prevents corrosion and additional anodic reactions from occurring, except at very large voltages where transpassivity and anodic breakdown occur. As a consequence, the barrier layer shields the underlying metal from other processes such as chemical or electrochemical dissolution.1 Though low anodic currents are observed upon the passivation of Sn, large cathodic currents can easily be generated, which result in cathodic breakdown by electrochemically removing the barrier layer. Such diode-like current− voltage behavior is typical of valve metals (such as Al, Ti, and Ta), which are used in various types of electronic components.2,3 The close resemblance of the kinetics of Sn anodization to that for the valve metals has been noted previously.4−7 The general experimental conditions necessary for oxide formation and passivation of the Sn surface can be predicted Received: July 16, 2013 Revised: September 23, 2013 Published: September 24, 2013 12969

dx.doi.org/10.1021/la402703w | Langmuir 2013, 29, 12969−12981

Langmuir

Article

density is exponentially dependent on the field strength, and is usually assumed to occur mainly by the transport of metal ions within the layer toward the solution.5 At low field strengths, the high field model predicts that the current density will exhibit ohmic behavior.1,5 In this article, we present the results of our initial investigations into the general electrochemical behavior of polycrystalline Sn electrodes onto which n-alkanethiol selfassembled monolayers (SAMs) have been formed. In particular, we have been interested in examining the anodic passivation characteristics that such SAMs might impart to the Sn surface, perhaps as a barrier to surface oxide formation. To the authors’ knowledge, the work reported here are the first studies of alkanethiol SAMs on Sn surfaces of any kind, whether single crystal or polycrystalline. Electrochemical studies of alkanethiol SAMs on various metal surfaces are plentiful, and many of them have focused on details associated with reductive desorption and oxidative deposition of the film.34 Other investigations have centered on studying the passivation and corrosion resistance properties resulting from SAM formation on numerous metal and semiconductor surfaces, including Au, Co, Cu, Fe, GaAs, Ge, InAs, Ni, Pd, Zn, and ZnO (see Supporting Information for published work in this area). Cyclic voltammetry and coulometry have been the techniques of choice in our initial studies so that useful comparisons can be made to the earlier electrochemical work on Sn passivation, as well as those studies that have examined the passivation characteristics imparted by SAMs to metal surfaces. The research reported here involves SAMs formed from various medium chain length n-alkanethiols (CnH(2n+1)SH, where n = 8, 10, or 12), where some emphasis is given to differentiating the voltammetric features arising from the SAMs compared to those from the surface oxide of Sn. Data demonstrating the effectiveness of electrochemically prepared SAMs as barriers to oxide growth on Sn will be presented and discussed.

from its potential-pH (Eh-pH, Pourbaix) diagram.7,8 In either highly acidic (pH < 0) or highly basic aqueous solution (pH > 13), active corrosion of the Sn will proceed due to the oxidative formation of various soluble forms of Sn2+- and/or Sn4+containing species, the nature of which depend on both the particular pH and electrode potential. From pHs 3 to 10 and in solutions free from very strong oxidizing agents, Sn transitions from a region of immunity to that of passivity; it is therefore almost entirely corrosion resistant across this pH range, except at very high voltages.8 Given that the immunity and passive regions are not contiguous outside this range (for 0 < pH < 3 and 10 < pH < 13), more demanding oxidative conditions (e.g., more positive applied potentials, stronger oxidizing agents) are required to reach the passive regime. The anodic behavior of tin in aqueous electrolyte solution and the nature of the resulting passive film structure have been extensively investigated.4−30 Tin oxides and hydroxides are amphoteric and exhibit a broad minimum in solubility near pH 8.5.8 Hence, many studies have employed borate or bicarbonate buffers so that the solution pH can be suitably controlled between pH 7 and 12. In addition, the simultaneous formation of both oxides and hydroxides for Sn(II) and for Sn(IV) are likely during anodization, given the similarities in their standard potentials; the presence of highly hydrated layers of SnO and SnO2 are therefore expected.4,8,16,18 Not surprisingly, the growth of the passivating layer on Sn has been found to be a complex process and the layer’s composition and thickness depend on the solution pH, the electrode potential, and the nature of anions comprising the electrolyte.4−30 Although some debate still exists as to the actual processes at work, the electrochemical oxidation of tin has been qualitatively divided into three potential regimes: (1) a narrow active dissolution region that is thought to correspond to formation of approximately one monolayer of SnO and/or Sn(OH)2 via a dissolution−precipitation process; (2) an active-to-passive region that corresponds to solid state conversion of the SnO/ Sn(OH)2 layer to SnO2/Sn(OH)4; and (3) a large passive region that is attained after conversion to SnO2/Sn(OH)4 is complete and over which dehydration occurs to form SnO2 as the primary film constituent.13 Studies using X-ray photoelectron spectroscopy25 and reflection−absorption infrared spectroscopy27 have provided evidence for hydration of the Sn films in the passive state, and that the hydroxylation content increases with pH.27 Using potentiodynamic methods, the initial processes involved with Sn oxidation in the active dissolution and active-to-passive regions in neutral or alkaline media have been explained using the layer pore resistance model developed by Müller and Calandra.4−6,14,15,21,24,31−33 In this model, the first step in film formation consists of the growth of a poorly conducting layer of uniform thickness at the electrode surface, until approximately 99% of the surface is covered. Once this coverage is reached, the rate of current flow sharply decreases and further film growth is limited by the ohmic resistance of the pores remaining within the layer. This model has been used to explain the first two anodic steps in film formation (i.e., regimes (1) and (2) above) as well as electroreduction of the passive layer.4,14−16,21,24 In contrast, under galvanostatic conditions others have explained oxide growth on Sn by a solid state ionic conduction mechanism typical of valve metals when large electric fields are present within the material (>106 V/cm).5,6,15 In this high f ield model, ionic conduction is driven by the electric field within the material such that the measured current

■

EXPERIMENTAL SECTION

Solutions. Solutions for all electrochemical experiments employed either ultrahigh purity water (resistivity = 18.2 MΩ-cm, provided by a Millipore Synergy water purification system) or absolute methanol (Pharmco, 99.9%) as the solvent. Both aqueous and methanolic borate buffer solutions were prepared with 50.0 mM boric acid (H3BO3, Sigma, electrophoresis grade) and 50.0 mM sodium tetraborate decahydrate (Na2B4O7·10H2O, Acros, 99.5%), which were used as received; thus, the total borate concentration in these solutions was 0.25 M. For the aqueous 50 mM borate buffer solutions, measured pH’s typically fell in the range 9.0−9.5. Solutions containing nalkanethiols (CnH(2n+1)SH = CnSH) were freshly prepared in standard volumetric glassware by dissolving a sufficient quantity of octane-1thiol (C8SH, Aldrich, 98.5%), decane-1-thiol (C10SH, Aldrich, 96%), or dodecane-1-thiol (C12SH, Acros Organics, 98.5+%) into methanolic 50 mM borate buffer to yield concentrations near 1.0 mM. This was accomplished using Wiretrol II glass micropipets (Drummond Scientific, Broomall, PA) or a Hamilton 7000 Series 5.00 μL syringe (Hamilton #7105). Prior to use, all volumetric glassware were cleaned in 0.5−1.0% v/v aqueous Hellmanex solutions (Hellma Worldwide) to minimize contamination. After cleaning, the glassware was rinsed with copious amounts of ultrapure water. Prior to solution preparation, a small volume of absolute methanol (5−10 mL) was used to rinse the volumetric glassware to remove any excess water left from the cleaning process. The choice of methanol as the nonaqueous solvent for these experiments was based on several considerations. Medium- and longchain alkanethiols are not readily soluble in aqueous solution, and thus an organic solvent was necessary to achieve the required concentrations (approximately 10−3 M). However, several of the 12970


Langmuir

Article

most common organic solvents used for solution-based alkanethiol SAM studies (e.g., ethanol, acetonitrile) were too nonpolar to solubilize the sodium tetraborate required for borate buffer formation. In addition, given the extensive literature involving electrochemical studies of tin electrodes and their oxidation in aqueous solution, the use of a relatively polar, protic solvent was necessary in order to make useful comparisons with these previous results. Methanol was the optimal choice given the above considerations: as a polar protic solvent, it was able to yield millimolar concentrations of all three alkanethiols studied as well as solubilize both H3BO3 and Na2B4O7 at the concentrations that were employed. Electrochemical Instrumentation and Accessories. All controlled-potential electrochemical experiments were performed with a CH Instruments 620A potentiostat, and employed an electrochemical H-cell to house the working, auxiliary, and reference electrodes. The Sn working electrode was fabricated from a 2-mm-diameter Sn rod (Alfa Aesar, Puratronic, 99.9985%) encased in a solvent-resistant CTFE plastic body (CH Instruments, Austin, TX); the resulting exposed geometric area of the electrode was 0.0314 cm2. A Pt wire was used for the auxiliary electrode, and all potentials were measured against either (1) a Ag/AgCl (3.5 M KCl) reference electrode for aqueous experiments, or (2) a Ag/AgNO3 (1.0 mM AgNO3, 0.10 M NaClO4, in methanol) reference electrode for nonaqueous experiments. Both reference electrodes (aqueous and nonaqueous) employed a porous Vycor glass tip to isolate the interior reference electrode solution from the exterior electrolyte solution. Typical scan rates in all experiments were 0.010 V/s. Prior to use, the Sn electrode was polished with coarse sandpaper, then a 1200 grit Carbimet disk, followed by 1 μm, 0.3 μm, and 0.05 μm alumina powder; the 1 and 0.3 μm alumina polishes were carried out on nylon cloth pads (Buehler, Lake Bluff, IL), whereas the final 0.05 μm alumina polish was performed using a Microcloth pad (Buehler). Afterward, electrochemical pretreatment of the Sn electrode surface was performed in order to remove any vestiges of Sn oxide or hydroxide that remained on the surface following the polishing procedures. This pretreatment typically consisted of cathodic polarization of the electrode for 15−20 min at a potential sufficiently negative to reduce any Sn oxide or hydroxide back to the native metal. In aqueous solution, cathodic polarization was usually performed at −1.4 V (vs Ag/AgCl), whereas in methanolic solutions the pretreatment was typically carried out at either −1.8 or −2.0 V (vs Ag/AgNO3). All solutions were purged with ultrahigh purity Ar gas for 15−30 min prior to each series of experiments in order to remove dissolved dioxygen gas. Normally, this purge was performed concurrently with the cathodic pretreatment process described above. In most instances, purging was halted during voltammetric measurements. While the removal of dissolved O2 from aqueous solutions by Ar purging was relatively straightforward, its removal from methanolic solutions was more difficult, perhaps because of its greater solubility. The presence of dissolved dioxygen in these studies was problematic for several reasons: (1) O2 reduction occurs at potentials < −1.35 V (vs Ag/ AgNO3) at the Sn surface and its reduction current tends to obscure redox features associated with the SAM, which occur in the same potential range. (2) O2 reduction in alkaline solution is known to produce superoxide and peroxide species, which can lead to oxidative removal of the alkanethiol SAMs from the Sn surface.35−38 Therefore, the elimination of dissolved dioxygen gas was critical in these studies. The H-cell employed in this work was found to provide a satisfactory seal against the ambient air and was able to remove dissolved O2 with reasonable efficiency so that the voltammetry associated with reductive desorption and oxidative deposition of the alkanethiol SAMs could be discerned.

■

base and methanolic base solutions. Since the electrode potentials necessary to initiate surface oxide/hydroxide formation on Sn (i.e., SnOx) were expected to be in close proximity to those corresponding to SAM deposition and removal, it was important to determine whether the potentials for the SAM could be differentiated from those for SnOx formation − and if so, whether the former were sufficiently cathodic that the SAM could be formed directly on the metal with little or no interference from competing growth of the surface oxide. Thus, the focus of these precursor studies with bare Sn electrodes was simply to distinguish Sn-specific voltammetric features from those for the SAMs − not to exhaustively characterize the nature of the oxidation processes occurring at the clean Sn surface. Voltammetric studies were first performed on bare Sn in 50 mM aqueous borate buffer (pH 9) in order to confirm its expected electrochemical behavior (as described in the literature) and to ascertain the general cleanliness of the surface. Afterward, the voltammetry of bare Sn was performed in 50 mM methanolic borate buffer so that comparisons could be made to the aqueous data. As described in the Experimental section, methanol was the optimal choice of polar protic solvent for these studies due its ability to solubilize the alkanethiols of interest as well as both components of the borate buffer. It was initially anticipated that voltammetric features for the bare Sn electrode would be similar in appearance in both the aqueous and methanolic borate buffer solutions. Aqueous Studies of Bare Sn Electrodes. Typical cyclic voltammograms (CVs) obtained for a bare Sn electrode in 50 mM aqueous borate buffer (pH 9) are displayed in Figure 1A, with voltammetric peaks labeled using Roman numerals. Each CV was initiated at −1.4 V (vs Ag/AgCl) and the potential was cycled at a scan rate of 10 mV/s to successively higher anodic potential limits, where the anodic limit was increased by 100 mV with each cycle. The inset to Figure 1A displays an expanded view of Peak V, which changes in magnitude and position as a function of the anodic limit. Prior to each voltammogram, cathodic polarization was performed at −1.4 V for 15−20 min while purging with Ar(g) in order to remove any remnants of surface SnOx that may have persisted from a previous voltammogram. Figure 1A displays only those scans within the range −1.4 to 1.0 V, where the upper end coincides with the beginning of Peak IV, a very broad and featureless plateau that extends up to the oxygen evolution region. The oxygen evolution reaction was observed beginning at potentials near 1.6 V, consistent with previous studies.4,14,15,19,29,30 When the potential is scanned more positively than 1.0 V (i.e., across the broad plateau represented by Peak IV), an additional reduction peak appears near −0.95 V whose beginnings can be identified by a weak feature just prior to Peak V in the CV whose anodic limit is 1.0 V. As discussed briefly in the Introduction, the processes associated with Sn anodization shown in Figure 1A are intrinsically complex given the multiple valence states for the metal and the amphoterism of the oxidation products. The appearance of Peaks I−V in aqueous solution and the behavioral dependence of Peak V on the anodic potential limit have been discussed at length elsewhere and are only briefly summarized here.4−30 Most studies indicate that Peak I signals the primary passivation of Sn, which occurs through the formation of a thin layer of SnO and/or Sn(OH)2 in a process that involves the precipitation of an intermediate soluble species onto the metal surface to form the film. The active-to-

RESULTS AND DISCUSSION

Electrochemistry of Bare Sn Electrodes. As a precursor to our study of alkanethiol electrochemistry at polycrystalline Sn electrodes, a necessary first step was an examination of the general redox behavior of bare Sn electrodes in both aqueous 12971


Langmuir

Article

limit (up to 1.0 V) may signify the presence of a single electroactive species, or multiple species whose reduction potentials are sufficiently close that they undergo reduction simultaneously. The association of Peak V with the simultaneous reduction of both Sn(II) and Sn(IV) to Sn(s) has previously been advocated and is consistent with the nearness of their standard potentials.4 With increasing anodic limit, the position of Peak V shifts negatively, which in turn shifts the onset of solvent reduction toward negative potentials since the hydrogen evolution reaction is kinetically hindered on the SnOx surface. However, one should note that regardless of the anodic potential limit, the total reductive charge contained within Peak V is significantly less than the oxidative charge that was passed during the initial anodic sweep. This indicates that other processes are likely operating in parallel with Sn oxidation (e.g., partial SnOx film dissolution, O2 evolution) or that some portion of the SnOx layer is not electroactive and cannot be reduced back to the metal at potentials prior to hydrogen evolution. Some evidence has indicated that electrochemical dissolution of SnOx occurs in the region of Peak III, which is quite broad and superimposed on top of the large steady state current plateau.4 This would be consistent with other studies of passive films, where anodic corrosion and dissolution of the layer has been observed to occur under steady state conditions such that a constant film thickness is maintained.1 Extensions of the high field model for passive films have been used to show that such steady state behavior occurs when the rate of oxidation at the metal/oxide interface equals the rate of dissolution.2 Under these conditions, the field strength within the oxide layer is constant (regardless of the applied potential) and a steady state current arises, the value of which is dictated by the rate of corrosive dissolution at the solution interface. Nonaqueous Studies of Bare Sn Electrodes. CVs obtained for a bare Sn electrode in 50 mM methanolic borate buffer are displayed in Figure 1B, again with voltammetric peaks labeled using Roman numerals. Prior to each CV in this series, cathodic polarization was performed at −2.0 V for 15− 20 min while purging with Ar(g) in order to remove any remnants of SnOx from a previous CV. Only those CVs leading to an anodic limit of −0.5 V (vs Ag/AgNO3) are shown, and peaks are numbered in accordance with their analogues in the aqueous data described above (Figure 1A). As expected, the voltammetry displayed in Figure 1B is similar to that observed in the aqueous borate buffer solutions, even when the potential is scanned more positively than −0.5 V. These similarities include (1) the number, general shape, and relative positions of the labeled peaks; (2) the measured current densities (40−50 μA/cm2) in the plateau region near the appearance of Peak III; and (3) the general potentiodynamic behavior of Peak V as a function of anodic limit. Differences between the aqueous and methanolic Sn voltammetry appear to be largely due to differences in electrochemical reversibility, with larger separations between the forward and reverse peak potentials in methanolic solution. For example, in aqueous solution Peaks I and V are separated by approximately 90 mV in those scans up to the lowest anodic limits, whereas in methanol the separation is closer to 125 mV. In methanol, the relative peak positions for Peaks I and II are more widely separated, and the peak potentials for Peak V shift over a broader range of values as a function of anodic limit. Furthermore, the voltammetric linewidths for Peak V appear slightly larger in methanol than in aqueous solution, particularly as the anodic limit approaches Peak III. It is also worth noting

Figure 1. (A) Cyclic voltammograms (CVs) of Sn in aqueous borate buffer. The inset expands the region around Peak V showing its behavior as a function of potential limit. Zero current density is represented by a horizontal line passing through the middle of the plot. The initial scan direction for all CVs is indicated by an arrow from a starting potential of −1.40 V. Cathodic polarization was performed prior to each scan at −1.40 V for 15−20 min in order to remove any surface oxide. Potentials were measured against a Ag/AgCl (3.5 M KCl) reference electrode, a Pt wire was employed as the counter electrode, and all scan rates were 10 mV/s. (B) CVs of bare Sn in methanolic borate buffer. Zero current density for each run is represented by a horizontal line passing through the middle of each scan. The initial scan direction for all CVs is indicated by an arrow from a starting potential of −2.00 V. Cathodic polarization was performed prior to each scan at −2.00 V for 15−20 min in order to remove any surface oxide. Potentials were measured against a Ag/ AgNO3 (1.0 mM AgNO3 + 0.10 M NaClO4 in methanol) reference electrode, a Pt wire was employed as the counter electrode, and all scan rates were 10 mV/s.

passive transition appears to be completed in Peak II when this SnO/Sn(OH)2 layer is further oxidized to Sn(OH)4, where the large degree of overlap between Peaks I and II reflects the closeness in the standard potentials for the respective processes.16,18 With increasing potential, it is thought that dehydration of the film leads to a more crystalline SnO2 deposit, which continues to grow via a solid state conduction mechanism. XPS evidence indicates that across the passive regime, that portion of the film in contact with the external solution consists of a hydrated mixture of Sn(II) and Sn(IV) compounds, whereas the interior portion contiguous to the metal consists of 1−2 monolayers of SnO whose thickness does not change with potential.28 The primary constituent within the center of the film is SnO2, whose thickness does increase with potential.28 The potentiodynamic behavior of Peak V is intriguing, as the peak potential and reductive current density are highly dependent on the anodic limit (Figure 1 and inset). Its occurrence as a single reduction feature regardless of the anodic 12972


Langmuir

Article

twice that passed in methanol. In contrast, the maximum RCD passed during the reverse sweep for anodic limits in Peak I→III region is virtually identical for both solutions (aqueous = filled squares; methanol = open triangles) and is attained at −0.70 V (methanol) and −0.10 V (aqueous), both corresponding to anodic limits in the region of Peak III. For anodic limits slightly higher than these, the RCD decreases in both solutions and plateaus at nearly the same minimum value; the plateau in each solution’s RCD occurs in the same range of anodic limits that characterizes that solution’s corresponding plateau in the OCD. Thus, the processes associated with Sn oxidation appear to be very similar in both solutions, and both solutions lead to similar quantities of surface oxidation products that are available for reduction when the potential scan is reversed (as characterized by their identical RCD’s). In fact, the quantity of reductive charge from the undissolved portion of the SnOx film that remains after scanning the potential to the valley between Peaks I and II in methanol (350 μC/cm2, Figure 2) is near the value previously attributed the oxidation of a single atomic layer of Sn(0) to Sn(II) in aqueous solution (0.3 mC/cm2).4 The most significant difference in the observed behavior between the two solutions stems from their respective OCDs, which indicate that Sn oxidation occurs to a greater extent in aqueous solution, though both solutions yield similar oxide film thicknesses. This is again consistent with the earlier notion that the expected oxidation products in methanol (polymeric Sn(II)- and Sn(IV)methoxides) are less soluble than their aqueous counterparts.39 A more detailed analysis of the processes involved in Peaks I−V is beyond the focus of this particular article and will not be attempted here. The characteristics of Sn anodization in aqueous and methanolic base solution will be the subject of a future publication. Electrochemistry of SAM-Covered Sn Electrodes: General Observations. When a clean polycrystalline Sn electrode under potential control (to prevent SnOx formation) is exposed to a moderately concentrated alkanethiol solution, oxidative deposition of the alkanethiols can be achieved via a one-electron process by scanning the electrode potential to values just prior to the commencement of surface oxide formation:

that a significantly larger current density is observed for Peak I in aqueous solution than in methanol. Since Peak I is thought to involve precipitation of an intermediate Sn(II)-containing soluble species onto the metal surface, the aqueous data suggests that active dissolution proceeds to a greater extent in water than in methanol. Such a result is consistent with the fact that Sn(II)- and Sn(IV)-methoxides are polymeric and insoluble in organic solvents, perhaps more so than their hydroxyl analogues in water.39 Coulometric Investigations of Bare Sn Electrodes. Coulometric information extracted from the CVs shown in Figure 1A,B also allow one to compare the effects of Sn surface oxidation in these solutions. Figure 2 displays the total oxidative

Figure 2. Plot showing the effect of anodic potential limit on the total oxidative and reductive charge density measured for bare Sn electrodes in aqueous (filled symbols) and methanolic (open symbols) 50 mM borate buffer solution. Roman numerals designate peak positions for the primary oxidation features in Figure 1A and B, where the solid and dotted downward arrows correspond to the methanolic and aqueous data, respectively. Italicized v’s indicate the position of valley minima between Peaks I and II, Peaks II and III, or Peaks III and IV.

and reductive charge densities as a function of the anodic potential limit for bare Sn in borate-buffered methanol (open symbols) and in aqueous solution (filled symbols). Note that in this figure, the peak potentials associated with Peaks I, II, III, and IV are indicated by arrows, where solid arrows refer to their positions in the methanolic data and the dotted arrows refer those in the aqueous data; lowercase italicized v’s indicate the position of valley minima between Peaks I and II, Peaks II and III, or Peaks III and IV. Some general observations are worth noting from the data presented in Figure 2. First, the aqueous coulometry (filled symbols) suggests that within the anodic potential limits investigated, SnOx film growth appears to occur in two stages, each of which culminates in a short plateau in the continually increasing oxidative charge density (OCD) and in a maximum in the reductive charge density (RCD) that coincides or slightly precedes the OCD plateau. Similar behavior is observed in the methanolic coulometry (open symbols) for film growth through the first stage (up through Peak III). Second, as the anodic limit is increased, significantly more oxidative charge is passed in aqueous solution (filled diamonds) than in methanol (open squares), at least in the early stages of the passivation process. Since the anodic limit for the methanolic data only extends up to the beginning of Peak III, useful comparisons can be made by examining data for the two solutions up to this peak. A quantitative comparison up to Peak III shows that the oxidative charge passed in aqueous solution is approximately

CnS− + Sn(s) → CnS − Sn(s) + e−

(1)

In this regard, Figure 3 displays a typical CV for bare Sn in 50 mM methanolic borate buffer compared to those obtained when the electrode was exposed to solutions that contained a different chain-length alkanethiol (C8SH, C10SH, and C12SH). All CVs in this figure originate from the same polycrystalline Sn electrode and extend to the same anodic limit. Each alkanethiol solution was prepared in 50 mM methanolic borate buffer and had a nominal alkanethiol concentration of 1.0 mM. Prior to each CV in this series, cathodic polarization was performed at −2.0 V (vs Ag/AgCl) for 15−20 min while purging with Ar(g) in order to remove any remnants of SnOx from a previous scan. CVs were then initiated from this potential up to 0.0 V at a scan rate of 10 mV/s. Figure 3B−D shows that the presence of the alkanethiols significantly suppresses the initial stages of surface oxide formation, which begin at −1.25 V in the blank CV shown in Figure 3A. Small oxidative peaks can be seen in this region, some of which are associated with oxidative deposition of the alkanethiol film as they are slightly more negative than those associated with SnOx formation (to be discussed below). These small oxidative peaks are obscured in Figure 3B by the 12973


Langmuir

Article

monolayers deposited on polycrystalline Au electrodes produces an almost identical oxidative peak which also coincides with the surface oxide (see Figure S-1, Supporting Information). Like the Sn data, it is only after the C4SH monolayer has been oxidatively desorbed that the features associated with formation of the Au surface oxide reappear. Third, the dependence of the oxidative peak potential on alkyl chain length is qualitatively similar to results obtained previously for alkanethiol40,41 and (n-mercaptoalkyl)biphenyl42 SAMs on Au. For reductive desorption in the alkanethiol-Au system, the reductive peak potential has been observed to shift to more negative values in a linear fashion with increasing chain length (by 15−20 mV/CH2 unit).40,41 This chain length dependence has previously been attributed to one of two possible effects: (1) greater pairwise interactions between nearest-neighbor alkyl chains as the chain lengthens,41,43 or (2) the increasing difficulty as the chain lengthens in establishing a potential gradient within the film sufficient to promote ion migration and penetration through the SAM to support the redox process.40,44 Unfortunately, the variation in oxidative desorption peak potentials with increasing chain length for the CnSH/Au system does not appear to have been investigated to date. For the oxidative desorption of 4-methyl-4′-(nmercaptoalkyl)biphenyl SAMs on Au(111), peak potentials were observed to shift linearly to more positive values for a narrow range of short alkyl spacer lengths (for n = 1−6), where the slope was identical to that measured for their reductive desorption (18 mV/CH2 unit).42 For the CnSH/Sn system under study here, a plot of the oxidative peak potential as a function of the number of CH2 units indeed shows a linear shift toward more positive potentials with increasing chain length. However, this shift is 75−80 mV/CH2 group (Figures S-2 and S-3, Supporting Information), which is significantly larger than that measured previously for the CnSH/Au system. It seems unlikely that the 4- to 5-fold increase in this shift relative to the CnSH/Au system can be explained by either larger pairwise interactions between neighboring alkyl chains or differences in SAM film thickness. The voltammetry shown in Figure 3 suggests the formation of a relatively compact blocking layer (particularly for the C10SH and C12SH films) with few surface Sn atoms available for surface oxide formation. Though larger film thicknesses might lead to larger peak shifts, SAMs in the CnSH/ Sn system can only be thicker than those on Au by at most 12− 13%, given the cosine dependence of the film thickness on the hydrocarbon chain tilt angle and the known chain tilt for CnSH/Au SAMs (θ ≅ 32°, relative to the surface normal).45 Such potentially small relative differences in the SAM thickness between these two systems does not comport with the 400− 500% increase in oxidative peak shift observed for the CnSH/Sn system. The large differences in peak shift observed for SAM oxidative desorption in the CnSH/Sn system compared to SAM reductive desorption in the CnSH/Au system can be reasonably explained as an interfacial resistance problem. The processes in question involve direct redox reactions of the SAM, which are both faradaic in nature; hence, the behavior of the peak potentials for both processes should therefore be described in terms of the film resistance rather than its capacitance. Interfacial capacitance measurements for alkanethiol SAMs are known to exhibit a distinct chain-length dependence,40 where the capacitance reflects the ability to store charge across an interface − or conversely, to leak charge across an interface.

Figure 3. Cyclic voltammograms of a polycrystalline Sn electrode taken at a scan rate of 10 mV/s in the following solutions: (A) 50 mM methanolic borate buffer solution (blank). (B) 1.0 mM C8SH + 50 mM methanolic borate buffer solution. (C) 1.0 mM C10SH + 50 mM methanolic borate buffer solution. (D) 1.0 mM C12SH + 50 mM methanolic borate buffer solution. Zero current density for each run is represented by a horizontal line passing through the middle of each scan. The initial scan direction for all CVs is indicated by an arrow from a starting potential of −2.00 V. Cathodic polarization was performed prior to each scan at −2.0 V for 15−20 min in order to remove any surface oxide.

electrochemical reduction of dissolved O2, which appears as a voltammetric wave extending from −1.8 to −1.2 V on the forward sweep. Suppression of O2 reduction in this region appeared to improve as the alkanethiol chain length increased (Figure 3C,D), allowing one to more clearly discern the alkanethiol voltammetry. Such behavior is consistent with the formation of a dense film that restricts solvent penetration and access of redox-active species to the electrode surface, thus preventing simple faradaic processes from occurring. The blocking characteristics of alkanethiol SAMs are known to improve with film thickness, and for sufficiently thick films faradaic currents are almost completely inhibited.34 The most significant difference between bare Sn and the SAM-covered Sn surfaces involves the large oxidative peak that occurs near −0.80 V for C8SH films (Figure 3B), which shifts positively with increasing chain length (Figure 3C,D). This peak, absent in the bare Sn CV (Figure 3A), is likely due to electro-oxidation of the SAM film. This assignment is based on several observations. First, as the potential is scanned positively from −2.0 V in solutions containing any of these alkanethiols, the features normally attributed to surface oxide formation on bare Sn at −1.2 V are almost completely absent. Only after the potential is scanned past the large oxidative peak in the alkanethiol-containing solution does the large steady state current plateau attributable to SnOx surface passivation appear. These observations indicate that chemisorption of the SAMs inhibits Sn surface oxidation, and that SnOx formation only proceeds once oxidative removal of the SAM has been achieved. Second, the total oxidative charge observed for this feature is comparable in magnitude to that observed for oxidative removal of similar alkanethiol films on polycrystalline Au. For example, the voltammetry of butane-1-thiol (C4SH) 12974


Langmuir

Article

because some portion of the film has been electrochemically removed.40 Thus, at the peak potential, dielectric breakdown of the passivating film has already occurred. Although the most reasonable explanation for the large differences in peak shift originates from considerations of the intrinsic film resistance of the SAM, we cannot rule out the possibility that low solution ionic conductivities may play some role. Several factors relating to the nature of the borate buffer solutions used in this work are expected to lead to relatively low solution conductivities, relative to those solutions employed in previous SAM studies.40−44 First, at the pHs and total borate concentrations employed in this work, distribution diagrams indicate that the supporting electrolyte is expected to consist primarily of neutral boric acid (B(OH)3) and borate anion (B(OH)4−), which together constitute approximately 60% of the total borate-containing species (Figure S-4, Table S-1, Supporting Information). Thus, a significant fraction of the supporting electrolyte consists of neutral species. Second, the tetrahedral borate anion exhibits a large hydrodynamic or Stokes’ radius (261 pm) in aqueous solution at 25 °C, which is comparable in size to the highly hydrated Li+(aq) ion.46 Estimates of the equivalent ionic conductivities and diffusion coefficients of the two species at 25 °C are very similar: B(OH)4− (D = 1.19 × 10−5 cm2 s−1 (pH 7.0); λ0− = 35.3 S cm2 mol−1); Li+(aq) (D = 1.029 × 10−5 cm2 s−1; λ0+ = 38.66 S cm2 mol−1).46−49 Such large effective sizes in aqueous solution lead to relatively sluggish aqueous mobilities, the magnitudes of which are comparable to other large singly charged ions such as iodate.49 Third, aqueous borate ion has been shown to form ion pairs with both Na+(aq) and Li+(aq), where the latter are more highly associated.47,50 Given the composition of the borate buffers used our studies, Na+·B(OH)4− ion pair formation will tend to reduce the equivalent molar conductivity of the solution. Fourth, the distribution diagrams also indicate the presence of small amounts of various polyborate anions (e.g., B3O3(OH)4−, B4O5(OH)42−), and that their speciation and fractional contributions occur in the same general pH range where Sn metal experiences its lowest solubility (pH 8−10) (Figure S-4, Table S-1, Supporting Information).8 Given their sizes, polyborate anions are also expected to exhibit lower ionic mobilities than borate; their formation should therefore decrease the solution conductivity. Oxidative Deposition and Reductive Desorption of Alkanethiol SAMs on Sn. As mentioned previously, oxidative deposition of an alkanethiol SAM can be achieved by scanning the potential of a clean polycrystalline Sn electrode to values just prior to the commencement of surface oxide formation in a moderately concentrated alkanethiol solution. In this regard, Figure 4 displays sequential CVs for bare Sn exposed to solutions containing either C10SH (Figure 4A) or C12SH (Figure 4B). Each alkanethiol solution was prepared in 50 mM methanolic borate buffer and had a nominal alkanethiol concentration of 1.0 mM. Prior to each CV in Figure 4, cathodic polarization was performed in the range −1.8 to −2.0 V (vs Ag/AgNO3) for 15−20 min while purging with Ar(g) in order to remove any remnants of SnOx from a previous scan. CVs were then initiated from these potentials up to −0.800 V at a scan rate of 10 mV/s. Typical CVs obtained for a bare Sn electrode in 1.0 mM C10SH + 50 mM methanolic borate buffer are displayed in Figure 4A. Each CV was initiated at −2.0 V and the potential was cycled to successively higher anodic potential limits, where the anodic limit was increased by 50 mV with each cycle. For all

If a significant amount of charge is leaked, the interface exhibits poor blocking behavior and the charge transfer process should then be characterized in terms of an interfacial resistance (conductance) as opposed to its capacitance. Porter and co-workers previously modeled SAM films on Au as two capacitors connected in series, where one corresponds to the metal−sulfur headgroup region and the other to the hydrocarbon chain region.40 From their modeling of the SAM reductive desorption peak potentials (Ep,red), they demonstrated that the potential gradient within the hydrocarbon portion of the SAM film reaches a constant value (approximately 4 × 106 V cm−1) and becomes chain-length independent when the chain contains more than 7−8 carbons. They estimated the intrinsic ionic conductivity of the hydrocarbon layer to be in the range 10−13 to 10−15 Ω−1 cm−1 and showed that the shift in Ep,red depends on both the film thickness (i.e., chain length) and solution ionic conductivity. They interpreted the shift in the reductive desorption overpotential with increasing chain length − having a slope of 20 mV/CH2 group − as arising from the resistance of the hydrocarbon layer to allow ion transport through the film to support the redox process.40 It is wellknown that the resistance to charge transport through the film should depend on the film thickness (L), the area of the electrode (A), and the intrinsic resistivity of the film itself (ρ): R = ρL/A. Assuming the active electrode surface area remains relatively constant between measurements, the internal resistance of the SAM is therefore expected to exhibit a thickness dependence. Ohmic considerations then dictate that larger film thicknesses, which result in larger internal resistances, should necessitate larger overpotentials. But even more important than this, ohmic behavior leads to even greater overpotentials if the passage of very large currents are required, such as those necessary to drive a redox process. Based on the arguments of Porter and co-workers, we propose that the dramatic difference between our peak potential shifts and theirs arises primarily from the fact that larger currents are required to drive the SAM oxidation process compared to the reduction process. It is clear that SAM oxidative desorption is significantly more complicated than reductive desorption, since oxidation of both the SAM and the Sn surface occur simultaneously. (The same is true for SAM oxidation on Au since both processes occur simultaneously there as well; see Supporting Information, Figure S-1.) The ohmic character associated with charge transfer through the SAM should lead to larger shifts in the peak potential if more current (ion transport) is required to achieve the redox process. This would certainly be the case for oxidative desorption of the SAMs since their oxidation is expected to proceed via a 3- or 5electron process, as compared to the 1-electron event associated with reductive desorption. In fact, the situation is made worse because some portion of the charge is being consumed to produce SnOx from the bare Sn, thereby reducing the efficiency of the SAM oxidation process. An even greater overpotential would therefore be required to promote sufficient charge transport across the interface to support SAM oxidation. Based on these arguments, the observed shifts can be seen to arise purely from an ohmic effect and it is not unexpected that a larger shift in peak potential is observed for SAM oxidation on Sn (75 mV/CH2 group) compared to SAM reduction on Au (20 mV/CH2 group). As Porter and co-workers noted, by the time the peak potential is reached, enough Faradaic charge has passed that the film no longer displays the intrinsically low conductivity it exhibited at the outset of the redox process 12975


Langmuir

Article

V becomes subsumed by the second reduction peak such that the former can no longer be differentiated. The overall behavior of these features is shown quantitatively in Figure 5, which displays the positions of both reduction

Figure 5. Plot showing the effect of the anodic potential limit on the positions of various reduction peaks observed for C10SH-covered Sn electrodes. Note that for the C10SH/Sn system, the behavior of the more positive reduction peak (□) coincides with that of Peak V on clean Sn (⧫), which is associated with SnOx reduction. The invariance in position for the more negative reduction peak (○) with anodic limit suggests that this feature corresponds to reductive desorption of the SAM.

peaks as a function of the anodic limit. For comparison, the position of Peak V (associated with SnOx reduction, Figure 1B) is also shown as a function of the anodic limit. Note that the behavior of the more positive reduction peak at −1.38 V mimics the behavior of Peak V, shifting in position in a similar manner across the same range of anodic limits. Based on these data, the reduction peak appearing at −1.38 V is therefore assigned to the reduction of SnOx. The voltammetric data show that this peak first appears as the potential is swept into the region where oxidative removal of the C10SH SAM occurs (at potentials > −1.10 V), indicating that Sn surface oxidation occurs concomitantly with oxidative desorption of the SAM. The position of the more negative reduction peak near −1.44 V deviates little with anodic limit (Figure 5) and is therefore assigned to reductive desorption of the C10SH SAM (i.e., the reverse of eq 1). Although features associated with SAM reduction are observed, those related to oxidative deposition of the SAM are not. Their absence is likely due to overlap of the O2 reduction wave which occurs in the same potential region. Electrochemical reduction of dissolved O2 is known to lead to superoxide ion O2−, which disproportionates in protic solvents to yield the hydroperoxide ion HO2−.35−38 Such species can lead to chemical oxidation of the SAM from the Sn surface and can explain why a peak attributable to electrochemical oxidative deposition of C10SH is not observed until the potential surpasses an anodic limit of −1.35 V. Typical CVs obtained for a bare Sn electrode in 1.0 mM C12SH + 50 mM methanolic borate buffer are displayed in Figure 4B. Each CV was initiated at −1.8 V and the potential was cycled to successively higher anodic potential limits, where the anodic limit was increased by 50 mV with each cycle. Electrochemical reduction of dissolved O2 can be observed again as a small reductive wave extending from −1.6 to −1.4 V in all scans during the forward sweep. The anodic end of this wave terminates at potentials that are approximately 50 mV more negative than those observed in the C10SH/Sn system,

Figure 4. CVs of a polycrystalline Sn electrode in the following solutions: (A) 1.0 mM C10SH + 50 mM methanolic borate buffer solution. (B) 1.0 mM C12SH + 50 mM methanolic borate buffer solution. The inset indicates the anodic potential limit for each scan. The initial scan direction for all CVs is indicated by an arrow from a starting potential of −2.00 V. Cathodic polarization was performed prior to each scan at −2.00 V for 15−20 min in order to remove any surface oxide. Potentials were measured against a Ag/AgNO3 (1.0 mM AgNO3 + 0.10 M NaClO4 in methanol) reference electrode, a Pt wire was employed as the counter electrode, and all scan rates were 10 mV/ s.

scans, a reductive wave extending from −1.8 to −1.35 V appears on the forward sweep, indicating that electrochemical reduction of dissolved O2 occurs in this region. This assignment is based on the observation that the current density associated with the wave increases when the Ar purge is halted or when no purge is performed during the cathodic polarization precleaning step. Once the forward sweep surpasses an anodic limit of −1.35 V, a small reduction peak appears near −1.44 V on the reverse sweep; this is the only reduction peak observed until the anodic limit extends past −1.20 V. For higher anodic limits, a small oxidation peak appears at −1.22 V on the forward sweep, and is followed by a second small reduction peak at −1.38 V on the reverse sweep. The position of this oxidation peak coincides with that for Peak I during SnOx formation (Figure 1B) and is therefore assigned to the initial stage of oxidation of residual Sn atoms not passivated by the SAM. When the anodic limit is increased beyond −1.10 V, the potential sweep starts to extend into the large oxidation peak near −0.80 V, which was discussed previously as being associated with oxidative removal of the SAM. At these anodic limits, the second reduction peak undergoes a significant increase in current density as well as a shift in its position toward more negative potentials. For anodic limits greater than −0.90 V, the first reduction peak near −1.44 12976


Langmuir

Article

indicating that the C12SH SAMs provide improved blocking characteristics for O2 reduction. Such blocking behavior is further corroborated by the fact that the residual Peak I SnOx feature observed at −1.22 V is significantly smaller in the C12SH/Sn system than in the former. More importantly, a couple of small broadened oxidation peaks appear in the potential region between the O2 reduction wave and the residual Peak I feature (i.e., between −1.40 and −1.22 V) − exactly where voltammetric peaks associated with oxidative deposition of the SAM are predicted to occur. Figure 4B also demonstrates that the greater passivation characteristics of the C12SH SAM push oxidative desorption of the film to anodic limits that are outside the range shown in this figure. Note that even up to −0.80 V, very little SnOx reduction is observed on any of the reverse sweeps. Unlike the C10SH SAMs, however, resolution of the individual reduction features for the SAM and for SnOx was more difficult in the C12SH/Sn system. The initial reduction peak for C12SH/Sn is observed at −1.42 V for anodic limits up to −1.25 V, and appears to be correlated with the small peaks attributed to oxidative deposition of the SAM. This reduction peak shifts very little with anodic limit and exhibits similar current densities compared to the C10SH SAMs (15−20 μA/cm2); it is therefore assigned to reductive desorption of the C12SH SAM. As the anodic limit passes over the residual Peak I feature at −1.22 V, the reverse sweep shows that the SAM reduction peak becomes obscured by an increase in current on its right-hand side due to the reduction of residual SnOx. Given the peak assignments made previously for Peaks I and II, it would be more specific to assign this additional current to the reduction of surface Sn(II)containing species not passivated by the SAM. The total reduction current density on the reverse sweep reaches a relatively constant value for anodic limits in the range −1.15 to −1.05 V, which corresponds to the valley between Peaks I and II in methanolic borate buffer (Figure 2). Afterward, another increase in reduction current is observed for anodic limits near −1.00 V, which corresponds to the position of Peak II. For anodic limits between −1.00 and −0.80 V, the reduction current remains relatively constant. Since Peak II is a significantly broader feature than Peak I, its presence would be difficult to discern in the forward scan if only a residual quantity of surface Sn atoms were involved. Therefore, based on the peak assignments, it is appropriate to assign the second increase in reduction current observed for anodic limits near −1.00 V to the reduction of surface Sn(IV)-containing species not passivated by the SAM. The two increases in reduction current that are observed for scans extending to anodic limits of −1.22 and −1.00 V likely originate from the same set of unpassivated Sn surface atoms that have undergone stepwise oxidation during the forward sweep. When the anodic limit is scanned to potentials greater than −0.80 V, there is a much larger increase in reduction current during the reverse sweep (Figure S-5, Supporting Information). Therefore, for anodic limits between −1.25 and −0.80 V, the C12SH/Sn data indicates that Sn surface oxidation occurs to a much more limited extent than in the C10SH/Sn system. Coulometric measurements support these arguments and are discussed in the next section. Coulometric Studies of SAM-Covered Sn Electrodes. Useful information on these CnSH/Sn systems can also be obtained by examining the charge contained within the peaks that have been assigned to electrochemical deposition or removal of the SAM. In this regard, Figure 6 plots the measured

Figure 6. Plot showing the effect of anodic potential limit on the charge density associated with the oxidation and reduction peaks displayed in Figure 4A and B. Coulometric data measured for Sn electrodes in methanolic 50 mM borate buffer solution containing either 1.0 mM C10SH (⧫ and ◊) or 1.0 mM C12SH (● and ○). Filled symbols represent oxidation data calculated during the forward sweep toward the anodic limit; open symbols represent reduction data calculated from the subsequent reverse sweep.

charge density associated with the oxidation and reduction peaks shown in Figure 4A and B as a function of the anodic potential limit. The charge was calculated directly from the peak area after performing background subtraction across the potential region of interest in order to correct for double layer charging effects. While this was readily achieved using a linear baseline subtraction for the reduction peaks, a background curve was used for the oxidation peaks on the forward sweep, as these features lie just a few hundred millivolts more negatively than the large peak assigned to SAM oxidative desorption. Thus, the oxidation data shown in Figure 6 (filled symbols) do not reflect the charge contained in the SAM oxidative desorption peak, but only that within the small oxidation features between −1.40 and −1.10 V. Several items in Figure 6 are worth noting. First, the charge associated with C12SH oxidative deposition (filled circles) initially appears at −1.40 V, nearly 100 mV more negative than the charge associated with the initial stage of SnOx formation (Figure 2). It is likely that the charge for C10SH oxidative deposition (filled diamonds) occurs in this region as well; however, overlap of the O2 reduction wave in this region precludes its observation. Second, the oxidative coulometry plateaus at a charge density in the range 124 ± 4 μC/cm2 (for C12SH/Sn) to 141 ± 8 μC/cm2 (for C10SH/Sn), where the errors in the numbers represent the 95% confidence limits obtained from the 7 or 8 data points making up the plateau. For the C10SH/Sn system, much of this charge probably results from oxidation of residual Sn atoms not passivated by the SAM. For the C12SH/Sn system which provides better surface passivation, a considerably smaller fraction of the total charge appears to be attributable to the oxidation of residual Sn. Third, the reduction coulometry for C12SH/Sn (open circles) first plateaus at anodic limits near −1.20 V (which corresponds to the location of Peak I) at a charge density in the range 144 ± 6 μC/cm2. Though small, a further increase in reductive charge occurs when the anodic limit nears −1.00 V (which corresponds to the location of Peak II), resulting in another charge density plateau. Afterward, the only significant increase in charge occurs at higher potentials, when the anodic limit shifts into the region where oxidative desorption of the SAM occurs (>−0.80 V). These changes in reduction charge 12977


Langmuir

Article

density near anodic limits of −1.20 and −1.00 V mimic the qualitative behavior described previously for the C12SH/Sn voltammetry and reflect the stepwise oxidation of unpassivated surface Sn atoms during the forward sweep. Thus, the difference in integrated charge density between these plateaus should represent the number density of surface Sn atoms not passivated by the SAM. The significantly larger reductive charge that is observed when the anodic limit extends above −0.80 V indicates that Sn surface oxidation occurs simultaneously with SAM oxidative desorption. In contrast, oxidative desorption of the SAM in the C10SH/Sn system is in close proximity to its oxidative deposition (Figure 4A). This is evident from the fact that the oxidative desorption peak potential occurs nearly 150 mV more negatively than that for C12SH/Sn (see Supporting Information, Figure S-2). Given their relative proximities, the initial stages of SAM oxidative desorption in the C10SH/Sn system significantly overlap with the residual Peak II features, and partly overlap with that of residual Peak I. Based on these considerations, it is not unexpected that a somewhat gradual increase in integrated reductive charge is observed as a function of the anodic limit in the C10SH/Sn system (Figure 6, open diamonds). The coulometric data for this system indicate only a slight break in integrated charge near the first plateau observed for the C12SH/Sn system. The limiting oxidative charge density at the plateau for these SAMs on Sn (120−140 μC/cm2) is significantly less than the total charge measured during SnOx formation, and is even a bit less than that measured through Peak I (530 μC/cm2 in methanol; see Figure 2). This difference can largely be explained by the fact that (1) Sn surface oxidation involves both 2- and 4-electron processes, compared to the 1-electron process associated with oxidative deposition/reductive desorption of alkanethiols; and (2) electrochemical dissolution of some portion of the SnOx layer operates in parallel with Sn surface oxidation, thereby increasing the overall oxidative charge. Such dissolution processes on our polycrystalline electrodes may lead to considerable variation in atomic-scale surface roughness between experiments, making coulometric estimates of the CnSH surface coverage less informative. Nonetheless, assuming a geometrical surface roughness factor of 1.5−2.0, the plateau value itself is comparable to charge densities measured for the reduction of various functionalized alkanethiols adsorbed in a close-packed (m√3 × n√3)R30° arrangement on Au(111) (74 μC/cm2).51−57 Though roughness factors of this magnitude are not unreasonable, no attempt has been made in the work reported here to ascertain the active surface area of our polycrystalline Sn electrodes. Such studies will be the subject of future investigations. Lastly, it is apparent from the CnSH/Sn CVs (Figure 3) that the charge associated with oxidative desorption of the SAMs is considerably larger than that corresponding to their reductive desorption. The relative quantities of charge for these two processes are shown in Table 1, which tabulates the charge densities for oxidative desorption of these SAMs and compares them to the limiting charge densities discussed above for their reductive desorption (or oxidative deposition). The data shown in this table were acquired from the CVs displayed in Figure 3, where the oxidative charge was calculated from the areas underneath the curves after performing a background subtraction using a linear baseline. Measurement of the charge for oxidative desorption is more complicated than that for reductive desorption because the former is convoluted with Sn surface oxidation. Since all CVs in Figure 3 originate from the

Table 1. Coulometry for Oxidative Desorption of CnSH/Sn SAMsa system

Qtot,ox (mC/ cm2)

QSAM,ox (mC/ cm2)b

QSAM,ox/QSAM,red ratioc

bare Sn C8SH/Sn C10SH/Sn C12SH/Sn

3.623 4.773 4.778 4.807

 1.150 1.155 1.184

 9.58 9.63 9.87

a

Data acquired from CVs displayed in Figure 3. bQSAM,ox = Qtot,ox(CnSH/Sn) − Qtot,ox(bare Sn). cLimiting value used for QSAM,red = 120 μC/cm2.

same Sn electrode and extend well past the SAM oxidative desorption peak to the same anodic limit, it has been assumed that the charge passed for Sn surface oxidation is identical in all CVs. Thus, the total oxidative charge for bare Sn (measured in Figure 3A up to 0.00 V) has been subtracted from that measured in the other three CVs to quantitatively determine the charge for oxidative removal of the SAM, Q SAM,ox = Q tot,ox(CnSH/Sn) − Q tot,ox(bare Sn)

(2)

where QSAM,ox is the oxidative charge that can be ascribed solely to SAM oxidative desorption. These values are then compared to the charge for SAM reductive desorption by calculating their ratios relative to the limiting charge densities at the plateau. Using this approach, the average charge ratio QSAM,ox:QSAM,red for our CnSH/Sn SAMs is 9.7 ± 0.4 (Table 1). One expects this value to correspond to the stoichiometric ratio of electrons associated with the two processes if the change in the active surface area during the oxidation process is either negligible or identical for all CVs (including the bare Sn), and if the amount of SnOx dissolution is the same for all systems (up to the same anodic limit). Our data indicates that if surface roughening or SnOx dissolution occurs during SAM oxidation, the extent of such processes is virtually identical for all the SAMs studied here since Qtot,ox(CnSH/Sn) varies by less than 0.4% (RSD) over the course of these experiments; this is certainly well within the error expressed by the charge ratios themselves. These estimates, however, do not take into account any differences in these processes between the bare Sn and SAMcovered Sn data. Assuming that reductive desorption of the SAM is a 1electron process as given in eq 1, our QSAM,ox:QSAM,red charge ratios suggest the occurrence of a more complex oxidation mechanism than that predicted by the simple formation of either alkylsulfinates or alkylsulfonates, which involve 3 and 5 electrons, respectively. RS−Au(s) + 2H 2O → Au(s) + RSO2− + 4H+ + 3e− (3)

, RS−Au(s) + 3H 2O → Au(s) + RSO3− + 6H+ + 5e− (4)

Similar observations have been made in electrochemical oxidative desorption studies of various alkanethiols on Au(111) single crystal electrodes, where QSAM,ox:QSAM,red ratios as high as 11−12 have been measured.51,54,58 These previous studies noted that the high ratios were consistent with (1) more extensive oxidation of the adsorbed alkanethiols, beyond that involving only the sulfur headgroup and/or (2) partial dissolution of the gold oxide during the anodic scan. 12978


Langmuir

Article

Normalized differential infrared reflectance spectra obtained in situ during the oxidative desorption of nonanethiol SAMs on Au(111) have previously evidenced ν asym (CO 2 − ) and νsym(CO2−) vibrational stretching modes at 1546 and 1405 cm−1, respectively, suggesting that oxidation of the S headgroup, C−S bond cleavage, and oxidation of the terminal CH2 moeity are all involved.58 Using a measured QSAM,ox:QSAM,red ratio of 11 ± 1, these authors proposed the following oxidation process:

their oxidative removal has been proposed to involve more than just the sulfur headgroup. Our chronocoulometry data for short-chain alkanethiols on polycrystalline Au electrodes evidenced oxidative charge densities comparable to those measured on the Sn electrodes. This data suggests that the alkanethiol oxidation processes on both surfaces are very similar, despite the significant differences in the inherent dissolution characteristics of the two metals.

■

C8H17CH 2S−Au(s) + 14OH−

ASSOCIATED CONTENT

S Supporting Information *

→ SO4 2 − + C8H17CO2− + 8H 2O + Au(s) + 11e−

Some figures and tables referenced in the text (Figures S-1 to S6; Table S-1). Additional publications concerning the passivation and corrosion resistance properties of SAMs on various metal and semiconductor surfaces can be found as well. This material is available free of charge via the Internet at http://pubs.acs.org.

(5)

On the other hand, several studies have indicated that electrochemical oxidative desorption of SAMs on Au leads to surface roughening, possibly as the result of precipitation of an insoluble product on the electrode surface.59,60 Given the limited work in this area at present, a single, unified description of the operative mechanisms involved in the oxidative desorption of thiol SAMs on Au has yet to be resolved, however. Our own chronocoulometric investigations of C4SH SAMs deposited on polycrystalline Au electrodes have evidenced QSAM,ox values in excess of 1.0 mC/cm2 (Figure S6, Supporting Information). These quantities are comparable in magnitude to those in the CnSH/Sn systems, despite the significant differences in anodic dissolution behavior between the two electrode materials. Our QSAM,ox:QSAM,red ratios for CnSH/Sn SAMs are therefore similar in magnitude to those measured on Au, which provides some indirect support for a similar oxidative mechanism on both surfaces. Only a clear understanding of the intrinsic dissolution characteristics of Sn during anodization and of Sn surface roughening for CnSH/Sn SAMs will help clarify whether the oxidative desorption process reflects an oxidation mechanism that is more extensive than that shown in eqs 3−4.

■

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Author Contributions

Student authors on this article from Samford University are listed in chronological order with regard to their participation in this work. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was partially supported by a summer undergraduate research stipend (B.C.W.) provided by Vista Engineering & Consulting, LLC (Birmingham, AL). Acknowledgment is also made to the Donors of the American Chemical Society Petroleum Research Fund for partial support of this research (B.W.G.). The authors are indebted to Drs. John T. Tarvin (Dept. of Physics, Samford University) and John E. Baur (Dept. of Chemistry, Illinois State University) for useful discussions during the preparation of this manuscript.

■

SUMMARY This work has shown that the potentials for oxidative deposition of medium chain length alkanethiols on polycrystalline Sn electrodes are more cathodic (by 100−200 mV) than those for Sn surface oxidation and that film formation can be performed directly onto the (virtually) oxide-free Sn surface. Furthermore, the data demonstrate that electrochemically prepared SAMs act as a barrier to oxide growth on Sn and that their passivating abilities improve with increasing film thickness (or chain length). Reductive desorption potentials for the SAMs are also in close proximity to those for reduction of the surface oxide, but little or no SnOx formation occurs unless the potential is made sufficiently anodic that the SAMs start to be removed oxidatively. The potentials for oxidative desorption of the SAMs shift to more positive potentials with increasing film thickness, and the measured shift is highly linear with chain length. Within the range of chain lengths investigated, coulometric measurements indicate that the charge passed during reductive desorption or oxidative deposition of the SAM is consistent with the formation of a single alkanethiol monolayer, taking into account surface roughness and heterogeneity effects. However, the charge measured during oxidative desorption of the SAMs is significantly higher than that predicted on the basis of simple alkylsulfinate or alkylsulfonate formation. These results are consistent with prior coulometric studies of alkanethiol SAMs on Au, where

■

REFERENCES

(1) Schultze, J. W.; Lohrengel, M. M. Stability, Reactivity, and Breakdown of Passive Films. Problems of Recent and Future Research. Electrochim. Acta 2000, 45, 2499−2513. (2) Lohrengel, M. M. Formation of Ionic Space Charge Layers in Oxide Films on Valve Metals. Electrochim. Acta 1994, 39, 1265−1271. (3) Electrochemical Dictionary; Bard, A. J., Inzelt, G., Scholz, F., Eds.; Springer-Verlag: Berlin, 2008. (4) Kaputsta, S. D.; Hackerman, N. Anodic Passivation of Tin in Slightly Alkaline Solutions. Electrochim. Acta 1980, 25, 1625−1639. (5) Metikoš-Huković, M.; Omanović, S. Kinetics of Anodic OxideFilm Growth on Tin: Ionic Transport Across a Barrier in the Highfield Limit. Mater. Chem. Phys. 1994, 38, 55−62. (6) Metikoš-Huković, M.; Rešetić, A.; Gvozdić, V. Behavior of Tin as a Valve Metal. Electrochim. Acta 1995, 40, 1777−1779. (7) Ammar, I. A.; Darwish, S.; Khalil, M. W.; El-Taher, S. Anodic Oxide Film Formation on Tin. Corrosion 1990, 46, 197−202. (8) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions, 2nd ed.; NACE Intl.: Houston, 1974. (9) Takeno, N. Atlas of Eh-pH Diagrams: Intercomparison of Thermodynamic Databases; National Institute of Advanced Industrial Science and Technology (NIAIST), 2005; Open File Report No. 419.

12979


Langmuir

Article

Under Ohmic Resistance Control. Electrochim. Acta 1974, 19, 901− 905. (32) Devilliers, D.; Lantelme, F.; Chemla, M. Surface Processes: Effect of Ohmic Polarization on Potentiodynamic V/I Curves. Electrochim. Acta 1986, 31, 1235−1245. (33) Noskov, A. V.; Grishina, E. P. Note on the Theory of ThreeDimensional Growth of the Porous Passivating Layers on Metals. J. Solid State Electrochem. 2008, 12, 203−206. (34) Finklea, H. O. Electrochemistry of Organized Monolayers of Thiols and Related Molecules on Electrodes. In Electroanalytical Chemistry; Bard, A. J., Rubinstein, I., Eds.; Marcel Dekker: New York, 1996; Vol. 19, pp 109−335. (35) Sawyer, D. T.; Seo, E. T. One-Electron Mechanism for the Electrochemical Reduction of Molecular Oxygen. Inorg. Chem. 1977, 16, 499−501. (36) Sawyer, D. T.; Chiericato, G., Jr.; Angells, C. T.; Nanni, E. J., Jr.; Tsuchiya, T. Effects of Media and Electrode Materials on the Electrochemical Reduction of Dioxygen. Anal. Chem. 1982, 54, 1720− 1724. (37) Wallace, G. G. Dissolved Oxygen: The Electroanalytical Chemists Dilemma. Trends Anal. Chem. 1985, 4, 145−148. (38) Šljukić, B.; Banks, C. E.; Compton, R. G. An Overview of the Electrochemical Reduction of Oxygen at Carbon-Based Modified Electrodes. J. Iran. Chem. Soc. 2005, 2, 1−25. (39) Turova, N. Y.; Turevskaya, E. P.; Kessler, V. G.; Yanovskaya, M. I. The Chemistry of Metal Alkoxides; Kluwer Academic Publishers: New York, 2002. (40) Widrig, C. A.; Chung, C.; Porter, M. D. The Electrochemical Desorption of n-Alkanethiol Monolayers from Polycrystalline Au and Ag Electrodes. J. Electroanal. Chem. 1991, 310, 335−359. (41) Imabayashi, S.; Iida, M.; Hobara, D.; Feng, Z. Q.; Niki, K.; Kakiuchi, T. Reductive Desorption of Carboxylic-acid-terminated Alkanethiol Monolayers from Au(111) Surface. J. Electroanal. Chem. 1997, 428, 33−38. (42) Aguilar-Sanchez, R.; Su, G. J.; Homberger, M.; Simon, U. Wandlowski, Th. Structure and Electrochemical Characterization of 4Methyl-4′-(n-mercaptoalkyl)biphenyls on Au(111)-(1 × 1). J. Phys. Chem. C 2007, 111, 17409−17419. (43) Aoki, K.; Kakiuchi, T. Probability Theory of Desorption Kinetics of Self-Assembled Alkanethiols Stabilized with Pair Interaction. J. Electroanal. Chem. 1998, 452, 187−192. (44) Evans, S. D.; Ulman, A. Surface Potential Studies of Alkyl-thiol Monolayers Adsorbed on Gold. Chem. Phys. Lett. 1990, 170, 462−466. (45) Schreiber, F. Structure and Growth of Self-Assembling Monolayers. Prog. Surf. Sci. 2000, 65, 151−256. (46) Corti, H.; Crovetto, R.; Fernández-Prini, R. Properties of the Borate Ion in Dilute Aqueous Solution. J. Chem. Soc. Faraday Trans. I 1980, 76, 2179−2186. (47) Corti, H.; Crovetto, R.; Fernández-Prini, R. Mobilities and IonPairing in LiB(OH)4 and NaB(OH)4 in Aqueous Solutions. A Conductivity Study. J. Solution Chem. 1980, 9, 617−625. (48) Park, J. K.; Lee, K. J. Diffusion Coefficients for Aqueous Boric Acid. J. Chem. Eng. Data 1994, 39, 891−894. (49) Vanýsek, P. Ionic Conductivity and Diffusion at Infinite Dilution. In CRC Handbook of Chemistry and Physics, 85th ed.; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 2005. (50) Rowe, L. M.; Tran, L. B.; Atkinson, G. The Effect of Pressure on the Dissociation of Boric Acid and Sodium Borate Ion Pairs at 25 °C. J. Solution Chem. 1989, 18, 675−689. (51) Calvente, J. J.; Kovácǒ vá, Z.; Sanchez, M. D.; Andreu, R.; Fawcett, W. R. Desorption of Spontaneously Adsorbed and Electrochemically Readsorbed 2-Mercaptoethanesulfonate on Au(111). Langmuir 1996, 12, 5696−5703. (52) Yang, D.-F.; Wilde, C. P.; Morin, M. Electrochemical Desorption and Adsorption of Nonyl Mercaptan at Gold Single Crystal Electrode Surfaces. Langmuir 1996, 12, 6570−6577. (53) Yang, D.-F.; Wilde, C. P.; Morin, M. Studies of the Electrochemical Removal and Efficient Re-formation of a Monolayer

(10) Kapusta, S.; Hackerman, N. Capacitive Studies of the Semiconducting Properties of Passive Tin Electrodes. Electrochim. Acta 1980, 25, 949−955. (11) Kapusta, S.; Hackerman, N. Photoelectrochemical Studies of Passivating Films on Tin in Slightly Alkaline Solutions. Electrochim. Acta 1980, 25, 1001−1006. (12) Kapusta, S.; Hackerman, N. Optical Studies of the Anodic Passivation of Tin. J. Electrochem. Soc. 1982, 129, 1886−1889. (13) Drogowska, M.; Ménard, H.; Brossard, L. Electrochemical Behaviour of Tin in Bicarbonate Solution at pH 8. J. Appl. Electrochem. 1991, 21, 84−90. (14) Brunetti, V.; Teijelo, M. L. Oxide/Hydroxide Films on Tin. Part I: Kinetic Aspects of the Electroformation and Electroreduction of the Films. J. Electroanal. Chem. 2008, 613, 9−15. (15) Brunetti, V.; Teijelo, M. L. Oxide/Hydroxide Films on Tin. Part II: Characterization of the Anodic Growth in Alkaline Solutions. J. Electroanal. Chem. 2008, 613, 16−22. (16) Gervasi, C. A.; Varela, F. E.; Vilche, J. R.; Alvarez, P. E. Investigations of the Electroreduction Process of Anodic Films Grown on Tin in Carbonate-Bicarbonate Solutions. Electrochim. Acta 1997, 42, 537−547. (17) Alvarez, P. E.; Ribotta, S. B.; Folquer, M. E.; Gervasi, C. A.; Vilche, J. R. Potentiodynamic Behavior of Tin in Different Buffer Solutions. Corros. Sci. 2002, 44, 49−65. (18) Alvarez, P. E.; Gervasi, C. A. Characterization of Passive Films on Tin Through Transient Electrochemical Techniques. Corros. Sci. 2004, 46, 91−107. (19) Gervasi, C. A.; Alvarez, P. E. Anodic Oxide Films on Tin in Carbonate-Bicarbonate Buffer Solution. Corros. Sci. 2005, 47, 69−78. (20) Almeida, C. M. V. B.; Giannetti, B. F. Protective Film Growth on Tin in Perchlorate and Citric Acid Electrolytes. Mater. Chem. Phys. 2001, 69, 261−266. (21) Bonilla, S. H.; Giannetti, B. F. Early Stages of Film Formation and Surface Growth on Tin Electrodes at Biocarbonate Medium: A Modelistic Approach. Z. Phys. Chem. 2004, 218, 837−856. (22) Bonilla, S. H.; Rodriguez, J.; Zinola, C. F.; Bello, C.; Giannetti, B. F. Film Formation and Surface Growth on Tin Electrodes in Bicarbonate Solutions: An Impedance Spectroscopy Study. Corros. Sci. 2005, 47, 835−848. (23) Giannetti, B. F.; Almeida, C. M. V. B.; Bonilla, S. H. Electrochemically Formed Film Layers of Importance in Electrocatalysis. In Electrocatalysis: Computational, Experimental, and Industrial Aspects; Zinola, C. F., Ed.; CRC Press: Boca Raton, FL, 2010; pp 191− 208. (24) Šeruga, M.; Metikoš-Huković, M. Passivation of Tin in Citrate Buffer Solutions. J. Electroanal. Chem. 1992, 334, 223−240. (25) Šeruga, M.; Metikoš-Huković, M.; Valla, T.; Milun, M.; Hoffschultz, H.; Wandelt, K. Electrochemical and X-ray Photoelectron Spectroscopy Studies of Passive Film on Tin in Citrate Buffer Solution. J. Electroanal. Chem. 1996, 407, 83−89. (26) Metikoš-Huković, M.; Omanović, S.; Jukić, A. Impedance Spectroscopy of Semiconducting Films on Tin Electrodes. Electrochim. Acta 1999, 45, 977−986. (27) Moina, C. A.; Ybarra, G. O. Study of Passive Films Formed on Sn in the 7−14 pH Range. J. Electroanal. Chem. 2001, 504, 175−183. (28) Keller, P.; Strehblow, H.-H. Passivity of Tin and CuSn Alloys in Alkaline Media Studied by X-ray Photoelectron Spectroscopy. In Passivation of Metals and Semiconductors and Properties of Thin Oxide Layers; Marcus, P., Maurice, V., Eds.; Elsevier: Amsterdam, 2006; pp 137−142. (29) Díaz, R.; Díez-Pérez, I.; Gorostiza, P.; Sanz, F.; Morante, J. R. An Electrochemical Study of Tin Oxide Thin Film in Borate Buffer Solutions. J. Braz. Chem. Soc. 2003, 14, 523−529. (30) Díaz, R.; Joiret, S.; Cuesta, Á .; Díez-Pérez, I.; Allongue, P.; Gutiérrez, C.; Sanz, F. Electrochemically Grown Tin Oxide Thin Films: In Situ Characterization of Electronic Properties and Growth Mechanism. J. Phys. Chem. B 2004, 108, 8173−8181. (31) Calandra, A. J.; De Tacconi, N. R.; Pereiro, R.; Arvia, A. J. Potentiodynamic Current/Potential Relations for Film formation 12980


Langmuir

Article

of Hexadecanethiol Self-Assembled at an Au(111) Single Crystal in Aqueous Solution. Langmuir 1997, 13, 243−249. (54) Exaplandiú, M. J.; Hagenström, H.; Kolb, D. M. Functionalized Self-Assembled Monolayers on Au(111) Electrodes: 1. Surface Structure and Electrochemistry. Langmuir 2001, 17, 828−838. (55) Sumi, T.; Wano, H.; Uosaki, K. Electrochemical Oxidative Adsorption and Reductive Desorption of a Self-Assembled Monolayer of Decanethiol on the Au(111) Surface in KOH+Ethanol Solution. J. Electroanal. Chem. 2003, 550/551, 321−325. (56) Sumi, T.; Uosaki, K. Electrochemical Oxidative Formation and Reductive Desorption of a Self-Assembled Monolayer of Decanethiol on a Au(111) Surface in KOH Ethanol Solution. J. Phys. Chem. B 2004, 108, 828−838. (57) Uosaki, K. Electrochemical Oxidative Formation of Ordered Monolayers of Thiol Molecules on Au(111) Surface. Chem. Rec. 2009, 9, 199−209. (58) Yang, D.-F.; Al-Maznai, H.; Morin, M. Vibrational Study of the Fast Reductive and the Slow Oxidative Desorptions of a Nonanethiol Self-Assembled Monolayer from a Au(111) Single Crystal Electrode. J. Phys. Chem. B 1997, 101, 1158−1166. (59) Musgrove, A.; Kell, A.; Bizzotto, D. Fluorescence Imaging of the Oxidative Desorption of BODIPY-Alkyl-Thiol Monolayer Coated Au Bead. Langmuir 2008, 24, 7881−7888. (60) Smith, S. R.; Guerra, E.; Shepherd, J. L. Investigating Gold Dissolution During Oxidative Desorption of Alkylthiol Self-Assembled Monolayers. ECS Trans. 2010, 28, 249−258.

12981


Anodic Passivation of Tin by Alkanethiol Self ... - ACS Publications

Recommend Documents