Electrochemically Modulated Liquid Chromatography and the Gibbs

(3) Chattoraj, D. K.; Birdi, K. S. Adsorption and the Gibbs Surface Excess; Plenum. Press: New ... 2002, 500, 722-740. ... on the stationary phase.19,...
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Anal. Chem. 2005, 77, 7399-7407

Electrochemically Modulated Liquid Chromatography and the Gibbs Adsorption Equation David W. Keller and Marc D. Porter*

Ames LaboratorysUSDOE, Institute of Combinatorial Discovery, and Department of Chemistry, Iowa State University, Ames, Iowa 50011

Electrochemically modulated liquid chromatography uses a conductive material like porous graphitic carbon (PGC) as a chromatographic stationary phase and a working electrode. This dual functionality enables manipulation of separations by changes in the potential applied (Eapp) to the packing. Thus, by monitoring the retention factor (k′) with respect to Eapp, a chromatographic tool for examination of electrosorption processes can be devised. This novel capability is developed herein by examining the retention of charged aromatic compounds at PGC. The Gibbs adsorption equation and related formulations (e.g., the Lippmann equation) are used to determine interfacial excesses (Γ) of these solutes in different supporting electrolytes, changes in interfacial tension (dγ), the charge on the electrode (qM), and the potential of zero charge (PZC). Values of the PZC were also determined from plots of ln k′ versus Eapp. In this case, the dependence of ln k′ reveals a shift in the PZC to more cathodic values as the strength of specific adsorption by anions as the electrolyte concentration increases. Together, these results provide insights into the retention mechanism and, more generally, to adsorption at electrified carbon electrodes. Extensions of this strategy as a probe of electrified interfaces with respect to mobile-phase composition, temperature, and pressure are briefly described. The electrical double layer is central to areas spanning colloidal stability, ion exchange at biological membranes, adsorption at chromatographic surfaces, and rates of heterogeneous electron transfer.1-6 Thus, the development of methods that probe the nature of the electrical double layer are of fundamental and technological importance. Traditional electrochemical approaches * Corresponding author. E-mail: [email protected]. Phone: 515294-6433. (1) Bockris, J. O. M.; Khan, S. U. M. Surface Electrochemistry: A Molecular Level Approach; Plenum Press: New York, 1993. (2) Bard, A. J.; Faulkner Electrochemical Methods: Fundamentals and Applications, 2 ed.; Wiley: New York, 2000. (3) Chattoraj, D. K.; Birdi, K. S. Adsorption and the Gibbs Surface Excess; Plenum Press: New York, 1984. (4) Conway, B. E. In Encyclopedia of Suface and Colloid Science; Somasundaran, P., Ed.; Marcel Dekker: New York, 2002; pp 1658-1681. (5) Lyklema, J. Fundamentals of Interface and Colloid Science: Liquid-Solid Interfaces; Academic Press: London, 1995. (6) Cline, K. K.; McDermott, M. T.; McCreery, R. L. J. Phys. Chem. 1994, 98, 5314-5319. 10.1021/ac051176s CCC: $30.25 Published on Web 10/20/2005

© 2005 American Chemical Society

have relied largely on measurements of potential, current, and surface tension to unravel connections between thermodynamics, electron-transfer rates, and double layer structure. Breakthroughs in spectroscopy, scanning probe microscopy, and other techniques have also been invaluable in the characterization of these interfaces.7-15 This paper harnesses a recently created tool in analytical separationsselectrochemically modulated liquid chromatography (EMLC)sas a new strategy to examine the electrical double layer. EMLC combines electrochemistry and high-performance liquid chromatography.16 It employs a conductive packing (e.g., porous graphitic carbon (PGC)) as both the stationary phase and working electrode in a chromatographic column designed to function as a three-electrode electrochemical cell. Through changes in the electrical potential applied (Eapp) to the packing, solute retention can be manipulated in a manner described by electrosorption processes.16,17 Most work with EMLC to date has focused on its utility in separation science.16-42 Herein, we take a different tactic (7) Lipkowski, J.; Ross, P. N. Imaging of Surfaces and Interfaces; Wiley: New York, 1999. (8) Tsionsky, V.; Daikhin, L. I.; Urbakh, M.; Gileadi, E. J. Electroanal. Chem. 2004, 2004, 1-99. (9) Schindler, W.; Hugelmann, M.; Hugelmann, P. Electrochim. Acta 2005, 50, 3077-3083. (10) Kolb, D. M. Surf. Sci. 2002, 500, 722-740. (11) Tian, Z. Q.; Ren, B. Annu. Rev. Phys. Chem. 2004, 55, 197-229. (12) McCreery, R. L. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1991; Vol. 17, pp 221-374. (13) Gileadi, E., Ed. Electrosorption; Plenum Press: New York, 1967. (14) Lipkowski, J. Can. J. Chem. 1999, 77, 1163-1176. (15) Abruna, H. D., Ed. Electrochemical Interfaces; VCH: New York, 1991. (16) Harnisch, J. A.; Porter, M. D. Analyst 2001, 126, 1841-1849. (17) Nikitas, P. J. Electroanal. Chem. 2000, 484, 137-143. (18) Deinhammer, R. S.; Ho, M.; Anderegg, J. W.; Porter, M. D. Langmuir 1994, 10, 1306-1313. (19) Deinhammer, R. S.; Porter, M. D.; Shimazu, K. J. Electroanal. Chem. 1995, 387, 35-46. (20) Deinhammer, R. S.; Shimazu, K.; Porter, M. D. Anal. Chem. 1991, 63, 18891894. (21) Deinhammer, R. S.; Ting, E.-Y.; Porter, M. D. J. Electroanal. Chem. 1993, 362, 295-299. (22) Deinhammer, R. S.; Ting, E.-Y.; Porter, M. D. Anal. Chem. 1995, 67, 237246. (23) Deng, H.; VanBerkel, G. J.; Takano, H.; Gazda, D.; Porter, M. D. Anal. Chem. 2000, 72, 2641-2647. (24) Ting, E.-Y.; Porter, M. D. Anal. Chem. 1997, 69, 675-678. (25) Ting, E.-Y.; Porter, M. D. J. Chromatogr., A 1998, 793, 204-208. (26) Ting, E.-Y.; Porter, M. D. J. Electroanal. Chem. 1998, 443, 180-185. (27) Ting, E.-Y.; Porter, M. D. Anal. Chem. 1998, 70, 94-99. (28) Ting, E.-Y.; Porter, M. D. J. Chromatogr., A 1998, 793, 204-208.

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by employing EMLC as a tool to examine fundamental properties of electrified liquid-solid interfaces. As will be shown, the chromatographic component of EMLC enables the direct determination of solute adsorption via the retention factor, k′, a situation that exploits the inherently large surface area of chromatographic packings with respect to that of electrodes in conventional electrochemical cells. The electrochemical component of EMLC, on the other hand, permits an investigation of the relationship between k′ and Eapp. By combining these characteristics, EMLC studies can yield insights into how Eapp modifies solute adsorption and, ultimately, several physical descriptors of the electrical double layer structure. Factors that influence retention in EMLC have been reviewed from a chromatographic perspective.16,17,43 In short, Nernstian behavior describes changes in retention through the redox activity of electroactive analytes, mobile-phase components, or coatings on the stationary phase.19,36,38,39,44 For electrochemically inert systems, donor-acceptor interactions can be applied to assess the relationship between k′ of a charged eluite (e.g., aromatic sulfonates, protonated pyridines, and inorganic anions) and Eapp.16,23,45-47 The trends in these systems often reflect the influence of an electric field on a Boltzmann’s distribution of ions. More recent work has used submolecular dipole, steric, and hydrophobic parameters to assess the effects of substituents on the retention of both charged and neutral analytes.22,26 This paper is part of a series of investigations that examine the fundamental factors that impact adsorption by modeling of the electrified interface in EMLC by using Gouy-Chapman electrical double layer theory.42 The work herein develops a basis for applying EMLC as a chromatographic tool for the examination of electrosorption processes17 by monitoring the retention factor (k′) with respect to Eapp. It is demonstrated by examining the retention of charged aromatic compounds in different supporting electrolytes (KCl, NaF, LiClO4) at PGC via the Gibbs adsorption equation and related formulations (e.g., the Lippmann equation). The approach results in readily adaptable procedures to determine solute interfacial excess (Γ), changes in interfacial tension (dγ), the charge on the electrode, and the potential of zero charge (PZC) in each electrolyte. Values of the PZC were also determined (29) Ting, E.-Y.; Porter, M. D. J. Electroanal. Chem. 1998, 443, 180-185. (30) Wang, S.; Porter, M. D. J. Chromatogr., A 1998, 828, 157-166. (31) Wallace, G. G.; Ge, H.; Teasdale, P. R. J. Chromatogr. 1991, 544, 305-316. (32) Nagaoka, T.; Fujimoto, M.; Nakao, H.; Kakuno, K.; Yano, J.; Ogura, K. J. Electroanal. Chem. 1994, 364, 179-188. (33) Shibukawa, M.; Unno, A.; Miura, T.; Nagoya, A.; Oguma, K. Anal. Chem. 2003, 75, 2775-2783. (34) Knizia, M. W.; Vuorilehto, K.; Schrader, J.; Sell, D. Electroanalysis 2003, 15, 49-54. (35) Takano, H.; Porter, M. D. Proc. Electrochem. Soc. 1999, 99-5, 50-60. (36) Fujinaga, T.; Kihara, S. CRC Crit. Rev. Anal. Chem. 1977, 6, 223-254. (37) Ge, H.; Wallace, G. G. J. Liq. Chromatogr. 1990, 13, 3245-3260. (38) Hern, J.; Strohl, J. H. Anal. Chem. 1978, 50, 1954. (39) Ghatak-Roy, A. R.; Martin, C. R. Anal. Chem. 1986, 58, 1574. (40) Ponton, L. M.; Porter, M. D. J. Chromatogr., A. 2004, 1059, 103-109. (41) Ponton, L. M.; Porter, M. D. Anal. Chem 2004, 76, 5823-5828. (42) Keller, D. W.; Porter, M. D. J. Chromatogr.. A 2005, 1089, 72-81. (43) Porter, M. D.; Takano, H. In Encyclopedia of Separation Science; Wilson, I. D., Adlar, E. R., Cooke, M., Poole, C. F., Eds.; Academic Press: London, 2000; pp 636-646. (44) Ge, H.; Wallace, G. G. Anal. Chem. 1989, 61, 2391-2394. (45) Knox, J. H.; Kaur, B.; Millward, G. R. J. Chromatogr. 1986, 352, 3-25. (46) Tanaka, N.; Tanigawa, T.; Kimata, K.; Hosoya, K.; Araki, T. J. Chromatogr. 1991, 549, 29-41. (47) Ross, P. LCGC 2000, 18, 14-25.

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from plots of ln k′ versus Eapp. In this case, the dependence of ln k′ reveals a shift in the PZC to more cathodic values as the strength of the specific adsorption by the anions of the electrolyte increases. Collectively, the findings indicate that the energetics of adsorption in EMLC are closely related to those observed in traditional electrochemical cells, yielding intriguing insights into the retention mechanism in EMLC and, more generally, adsorption at electrified carbon electrodes. Further extensions of this strategy as a probe of electrified interfaces are briefly described. THEORETICAL BACKGROUND The Gibbs adsorption equation permits a first principle assessment of the thermodynamics that govern retention in EMLC. Historically, this equation was used to describe adsorption in electrocapillary experiments in which the interfacial tension of a mercury electrode was followed as a function of Eapp and solvent composition48-52 via changes in droplet size and shape. These data were employed to calculate key descriptors of the electrified interface, most notably electrode charge and interfacial excess. Assuming an ideally polarized electrode-solution interface and the absence of junction potentials elsewhere in the cell, a specific form of the Gibbs adsorption equation can be derived upon application of the Gibbs-Duhem relationship.

-dγ ) qM dEapp +

∑Γ du i

i

(1)

i

In eq 1, dγ is the change in interfacial tension, qM is the electrode charge, and Γi and dµi are the respective surface excesses and changes in chemical potential associated with all ions i in solution. When the solution composition is fixed, eq 1 reduces to the Lippmann equation.

[-dγ/dEapp]µi ) qM

(2)

The Lippmann equation shows that the changes in interfacial tension with respect to Eapp can be used to determine qM. In an analogous fashion, the classical form of the Gibbs adsorption isotherm can be obtained by fixing Eapp in eq 1 and rearrangement,

[-dγ/dui]Eapp ) Γi

(3)

The Gibbs adsorption isotherm is useful for the determination of the interfacial excesses of each component in the solution. Electrosorption at solid-liquid interfaces can be interrogated by a variety of techniques.7-15 EMLC represents an intriguing addition to this toolbox through its ability to monitor the potential dependence of solute retention on a stationary phase, partially mimicking earlier work that employed fluidized bed reactors and other strategies with high surface area electrodes.53-55 Importantly, (48) Bockris, J. O. M.; Reddy, A. K. N. Modern Electrochemistry, an Introduction to an Interdisciplinary Area; Macdonald: London, 1970. (49) Grahame, D. C.; Whitney, R. B. J. Am. Chem. Soc. 1942, 64, 1548-1552. (50) Grahame, D. C. Chem. Rev. 1947, 41, 441-501. (51) Devanathan, M. A. V.; Tilak, B. V. K. S. R. A. Chem. Rev. 1965, 65, 635. (52) Delahay, P. Double Layer and Electrode Kinetics; Interscience: New York, 1965. (53) Alkire, R. C.; Eisinger, R. S. J. Electrochem. Soc. 1983, 130, 93-101. (54) Soffer, A.; Folman, M. J. Electroanal. Chem. 1972, 38, 25-43.

Table 1. Structures of the Mono- and Divalent Ions Employed as Test Solutes

Γi is closely related to the chromatographic retention factor of solute i, k′i. As defined, k′i describes the migration rate of solute i in a column as a function of its distribution between the mobile and stationary phases.

k′i ) ns/nm

(4)

In eq 4, ns and nm represent the number of moles of solute adsorbed onto the stationary phase and in bulk solution, respectively. Similarly, Γi is defined as the difference in concentration of a species present in the interfacial region and its concentration in bulk solution,

Γi ) (1/A)(ns - nm)

(5)

where A is the surface area of the stationary phase. Since the total amount of analyte is fixed by the volume and concentration of sample injected onto the column, the relationship between k′i and Γi is described in eq 6 by combining eqs 4 and 5,

Γi )

ntot (k′i - 1) A (k′i + 1)

(6)

where ntot is the moles of injected analyte (ntot ) ns + nm). Taken together, eqs 1-6 permit calculations of qM, Γi, and dγ from the experimentally observed values of k′i. In other words, Γi can be determined from the retention time of the analyte and subsequently used to calculate dγ and qM. EXPERIMENTAL SECTION Chemicals and Reagents. The analytes and their acronyms are given in Table 1. Disodium 1,3-benzenedisulfonate (1,3-BDS), sodium benzenesulfonate (BS), sodium p-toluenesulfonate (TS), phenyltrimethlyammonium p-toluenesulfonate (PTMA), dibromomethane, and lithium perchlorate were purchased from Aldrich. Sodium p-chlorobenzenesulfonate (CBS) was from TCI America. (55) Niu, J.; Conway, B. E. J. Electroanal. Chem 2003, 546, 59-72.

Disodium 1,5-naphthalenedisulfonate (1,5-NDS) was acquired from Eastman Kodak. Sodium fluoride, potassium chloride, and HPLCgrade acetonitrile were from Fisher. All chemicals were used as received. All aqueous solutions were prepared with water purified by a Millipore Milli-Q system. Before use, solutions were passed through a 0.5-µm filter (Osmonics) and thoroughly sparged with helium. EMLC Column. The design of the column has been detailed elsewhere27 and employed 5-µm PGC (ThermoHypersil) particles as the stationary phase. PGC is devoid of any detectable oxygencontaining functional groups, as determined by X-ray photoelectron spectroscopy (estimated detection limit, 0.2 atom %).22 The manufacturer has specified a nominal pore diameter for PGC of 250 Å. The surface area of PGC packed in the column, based on BET measurements (120 m2/g) and the amount of material loaded into the column (0.25 g), is 30 m2. Instrumentation. Chromatographic experiments were performed using a HP 1050 series module equipped with solvent cabinet, autosampler, quaternary pumping system, and diode array detector (Hewlett-Packard). The HP 1050 module was interfaced to a Pentium III 600 MHZ computer equipped with HP Chemstation software for control of injection sequences, data acquisition, and pumping parameters. The potential of the working electrode was controlled using a model 174A polarographic analyzer (Princeton Applied Research). Mode of Operation. In examining the effect of supporting electrolyte on the separation of aromatic sulfonates, the mobilephase composition was controlled by mixing of a water/acetonitrile (95:5) solution on-line with a water/acetonitrile (95:5) solution that contained the electrolyte. After a change in either Eapp or the mobile phase, the system was allowed to reach a steady state under the new conditions. The achievement of a steady state was assumed if the baseline remained stable for 10 min after the change in operational conditions. Injection volumes were set at 10 µL, and injections containing mixtures of analytes were repeated five times at each condition. Individual injections of each analyte were performed in triplicate for peak identification and for retention time determination when elution bands overlapped. Under conditions where PTMA and TS Analytical Chemistry, Vol. 77, No. 22, November 15, 2005

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coelute, the retention of PTMA was determined by subtracting the response obtained when only TS was injected from that for injections of mixtures containing both compounds. An injection of water was performed to determine the dead volume of the column (0.44 mL) for calculation of k′. The absorbance of all analytes was monitored at 212 nm. To compensate for peak tailing, retention times were determined from the first statistical moment56-58 analysis of all chromatographic peaks. The flow rate of the mobile phase was 0.4 mL/min. All experiments were performed at room temperature. RESULTS AND DISCUSSION System Definitions. We have paralleled the approach used in the classical work50,52 in applying the Gibbs adsorption equation to EMLC, which reflects equilibrated analyte adsorption as demonstrated by a nondetectable changes in k′ with respect to the flow rate. Thus, the shorthand notation for an EMLC column that employs a Ag/AgCl (saturated NaCl) reference electrode, sodium perchlorate as the supporting electrolyte, an acetonitrilewater (ACN-H2O) mobile phase, and a negatively charged solute can be written as

Cu′/Ag/AgCl/NaCl, H2O//NaClO4, H2O,

Figure 1. Chromatograms showing the influence of supporting electrolyte concentration at +50 mV vs Ag/AgCl (saturated NaCl) reference electrode. NaF concentrations in a 95:5 H2O-ACN mobile phase: (A) 40, (B) 80, (C) 120, and (D) 160 mM.

ACN//NaClO4, H2O, ACN, NaA/C/s-s/Cu′′ (7)

where Cu′ and Cu′′ represent the electrical leads between the potentiostat and reference and working electrodes, s-s denotes the stainless steel frit used to make electrical contact with the carbonaceous stationary phase C, and NaA symbolizes the monosodium salt of analyte A-. The Gibbs equation for this cell can be simplified by applying the standard assumptions regarding interfacial polarization51 and the Gibbs-Duhem relationship.59 The Gibbs adsorption equation under these conditions with our EMLC setup becomes

-dγ ) qM dEapp + ΓA-(H2O) duNaA + ΓACN(H2O) duACN + ΓNa+(H2O) duNaClO4 + ΓClO4-(H2O) duNaClO4 (8)

The term ΓX(H2O) is the relative surface excess of species X in solution, where X is the analyte (A-), ACN, or supporting electrolyte (Na+ or ClO4-). By fixing the ACN level in the mobile phase and the analyte concentration, eq 8 reduces to a form that is identical to the classical solution for an ideal polarized electrode, i.e., eq 1. Effect of Eapp and Electrolyte Composition on Retention. Chromatographic separations for mixtures of the six aromatic ions in Table 1 were performed at multiple values of Eapp and varied concentrations of three different monovalent (1:1) supporting electrolytes (NaF, KCl, LiClO4). Several sets of chromatograms, (56) Hewlett-Packard In Understanding your Chemstation; Hewlett-Packard: Wilmington, 1994; Vol. 1, pp 185-201. (57) McNair, H. M.; Cook, W. M. Am. Lab. 1973, 5, 12-18. (58) Grushka, E.; Myers, M. N.; Schettler, P. D.; Giddings, J. C. Anal. Chem. 1969, 41, 889-892. (59) It is noted that by assuming only the stationary mobile-phase interface is polarized, a slight overestimate in the calculation of charge on the stationaryphase surface may arise due to small potential drops at the liquid junctions.

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for example, were obtained using NaF as the supporting electrolyte at 40, 80, 120, and 160 mM, a mixed solvent mobile phase (5% acetonitrile-95% water), and values of Eapp equal to +50, 0, -50, -100, and -150 mV. Similar sets of experiments were carried out using KCl and LiClO4. Representative sets of the chromatograms obtained with NaF are shown in Figures 1 and 2 at +50 and -150 mV, respectively. At +50 mV, the components of the mixture are fully resolved at all electrolyte concentrations, with elution times of 21-26 min. Moreover, a change in electrolyte concentration alters retention in a manner that is dependent on the magnitude and sign of the charge on the solute. For example, k′ for the singly charged analyte BS (2) undergoes a decrease from 2.49 in 40 mM NaF to 2.21 in 160 mM NaF. This change corresponds to an 11% decrease in k′ and is comparable to those found for the singly charged anions TS (3) and CBS (4). The same concentration span yields a larger relative decrease in k′ for the dianions 1,3-BDS (1) and 1,5-NDS (6); the respective decreases are 25 and 18%. The retention dependence of the monovalent cation PTMA (5), however, differs from the anionic solutes. An increase in NaF concentration actually augments k′ by ∼7%, changing from a value of 1.49 to 1.61 in moving from the lowest to highest concentration of NaF. The separations in Figure 2 are markedly different from those in Figure 1. As is apparent, an Eapp of -150 mV not only changes elution order and time but also the retention dependence on the concentration of supporting electrolyte. At +50 mV, the elution order is 1,3-BDS < PTMA < BS < TS < 1,5-NDS < CBS. The order at -150 mV, however, is 1,3-BDS < BS < 1,5-NDS < TS < PTMA < CBS. The overall effect is a reduction in elution time by as much as 12 min, depending on electrolyte concentration. Moreover, the impact of electrolyte concentration on retention in Figure 1 reverses in Figure 2. In Figure 2, an increase in electrolyte concentration decreases the retention of the positively

Figure 2. Chromatograms showing the influence of supporting electrolyte concentration at -150 mV vs Ag/AgCl (saturated NaCl) reference electrode. The insets show an expansion of the elution profile for PTMA. NaF concentrations in a 95:5 H2O-ACN mobile phase: (A) 40, (B) 80, (C) 120, and (D) 160 mM.

charged PTMA but increases the retention of the five aromatic anions. When the NaF concentration is raised from 40 to 160 mM, k′ for PTMA decreases by 31%. In contrast, k′ increases for the monoanions BS, TS, and CBS by 15-23% and the dianions 1,3BDS and 1,5-NDS by 33-36%. The next section develops an interpretation of these trends. Capacity-Potential Plots. To more fully delineate the role of supporting electrolyte, data from experiments such as those in Figures 1 and 2 for the three electrolytes were analyzed by constructing plots of ln k′ versus Eapp.21,22,27,42 A portion of these findings, which are indicative of the results for all six solutes, is presented for TS in Figure 3. The plots yield four interesting observations. First, all the plots exhibit a linear dependence of ln k′ on Eapp (R2 values are greater than 0.95). Second, all slopes diminish as electrolyte concentration increases. Third, the plots for each electrolyte converge and intersect within a narrow window (∼10-40 mV) of Eapp. Fourth, and perhaps most importantly, the point of convergence is dependent on electrolyte identity. In discussing these data, the point of intersection in Figure 3A serves as an effective starting point. Linear regression analysis for each set of data indicates an intersection at -20 ( 12 mV. It signifies that the adsorption of TS at -20 mV is independent of NaF concentration. We ascribe the intersection potential to the PZC of the packing as probed by TS in this mixed solvent mobile phase. This assertion reflects the fact that the charge on an ideally polarizable electrode is zero at the PZC. As a consequence, the electrostatic driving force for adsorption of charged species is zero. It follows that the PZC corresponds to the potential in which retention is invariant with respect to electrolyte concentration. The value of the PZC found using this method is consistent with other reported values for carbon electrodes (+100 to -200 mV (vs Ag/

Figure 3. Plot of ln k′ vs Eapp (vs Ag/AgCl saturated NaCl) for TS at multiple concentrations of (A) NaF, (B) KCl, and (C) LiClO4. Inset is an expanded view (1.5×) showing the intersection of the plots. Each data point represents the average of five replicate injections with error bars on the order of the size of the points.

AgCl/saturated KCl)).60 The validity of this treatment, along with issues related to effect of surface oxygen groups, will be examined in further in a subsequent section. A purely electrostatic argument insinuates that a charged solute should not be retained at the PZC, eluting rather with the solvent band. Retention at the PZC therefore reflects the strength of the existence of specific interactions between TS and the high surface area packing, which is then modulated by the electrostatic interactions at departures of Eapp from the PZC. That is, changes in Eapp alter the magnitude of electrostatic interactions between the solute and electrode and between the electrolyte components (60) Kinoshita, K. Carbon Electrochemical and Physicochemical Properties; Wiley: New York, 1988.

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Table 2. Chromatographically Determined Values for the PZC of 1,3-BDS, BS, TS, CBS, 1,5-NDS, and PTMA in Electrolytes Containing NaF, KCl, or LiClO4a PZC (mV vs Ag/AgCl satd NaCl) analyte

NaF

KCl

LiClO4

1,3-BDS BS TS CBS 1,5-NDS PTMA average

-80 ( 28 -40 ( 11 -20 ( 12 -20 ( 12 -25 ( 13 -10 ( 23 -30 ( 17

-100 ( 35 -245 ( 25 -290 ( 20 -290 ( 18 -280 ( 21 -50 ( 28 -205 ( 25

-185 ( 40 -260 ( 20 -340 ( 17 -350 ( 29 -325 ( 24 not detectedb -285 ( 26

a The ranges provided were estimated based on the variability of the slopes in plots such as those shown in Figure 3. b PTMA was buried in the background signal with LiClO4 as the supporting electrolyte.

and electrode. Note that the latter interactions are in direct competition with the former interactions. This prediction is evident in the experimental data. For example, when the NaF concentration increases from 40 to 160 mM, the sensitivity (i.e., the slope of a ln k′ vs Eapp plot) for TS to changes in Eapp decreases from 5.2 to 3.5 mV-1. This dependence arises from the increase in electrolyte concentration, raising the probability for electrolyte ions to act as counterions in the electrical double layer. The slopes of the plots also reveal a dependence on the identity of the supporting electrolyte. At a given ionic strength, the sensitivity of TS to a change in Eapp increases as LiClO4 < KCl < NaF. This trend is similar to that found for the strength of specific adsorption of the electrolyte anions on carbon and other solid electrodes.61-66 Since the three electrolytes are composed only of monovalent ions, the difference in the sensitivity of retention to Eapp reflects the increase in competition by the more specifically adsorbed electrolytes for adsorption on the packing. Therefore, the amount of solute adsorbed on the surface and the extent in which Eapp influences retention will be reduced as the strength of the specific adsorption by the electrolyte components increases. Studies also indicated that differences in the identity of the cation have a negligible influence on solute retention. Furthermore, the identity of the cations used herein has little (if any) affect on the PZC at mercury electrodes.67 We conclude that observed differences in the PZC originate from the identity of the electrolyte anions. Table 2 summarizes the results from graphical determinations of the PZC for each analyte in the three electrolytes. As is evident, the PZC is strongly dependent on the composition of the electrolyte, but much less so on the identity of the solute. The trend for electrolytes follows the dependence expected from the (61) Paik, W.; Genshaw, M. A.; Bockris, J. O. M. J. Phys. Chem. 1970, 74, 42664275. (62) Mattson, J. S.; Mark, H. B. Activated Carbon: Surface Chemistry and Adsorption from Solution; Marcel Dekker: New York, 1971. (63) Bode, D. D.; Andersen, T. N.; Eyring, H. J. Phys. Chem. 1967, 71, 792797. (64) Perkins, R. S.; Andersen, T. N. In Modern Aspects of Electrochemistry; Bockris, J. O. M., Conway, B. E., Eds.; Plenum Press: New York, 1969; Vol. 5, pp 203-280. (65) Elkafir, C.; Chaimbault, P.; Dreux, M. J. Chromatogr., A 1998, 829, 193199. (66) Conway, B. E. Electrochim. Acta 1995, 40, 1501-1512. (67) Graham, D. C. J. Electrochem. Soc. 1951, 98, 343-350.

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observed eluent strength. That is, the greater the eluent strength of the electrolyte, which is dictated primarily by its extent of specific adsorption, the more negative the value of the PZC. There are also intriguingly subtle changes in the PZC with respect to the identity of the solute. These differences point to contributions from both nonspecific (e.g., 1,3-BDS and BS) and specific interactions (e.g., BS and CBS) of varying degrees. Nevertheless, efforts to unravel the role and magnitude of the contributions to specific adsorption by drawing, in part, on the use of submolecular (e.g., dipole and hydrophobic)22,26 parameters will require investigations of a much broader set of structurally varied solutes. The differences in the PZCs provide a means to quantify elution strength with respect to the electrolyte. This treatment normalizes the system to the weakest eluent, NaF, and calculates the relative free energy of adsorption (∆Gads ) for the two other electrolyte anions as,2

∆Gads ) zF ∆Ψ

(9)

where ∆Ψ is the difference in the PZC observed for KCl and LiClO4 relative to NaF. Using the average PZC values for each electrolyte in Table 2, the shift in the PZC in going from NaF to KCl is ∼-175 mV, whereas that from KCl to LiClO4 is ∼-70 mV. With these averages, eq 9 indicates that adsorption of Cl- is thermodynamically favored over that of F- by -16.8 kJ/mol-1, while ClO4- adsorption is -23.5 kJ/mol-1 more exergonic than F- adsorption. These values of ∆Gads therefore reflect the differences in the specific interactions of these anions. Gibbs Adsorption Analysis of the Chromatographic Data. As eq 3 shows, the Gibbs adsorption isotherm describes the influence of system parameters such as electrolyte concentration and Eapp on Γi and, therefore, k′. This section presents a Gibbsian analysis of the results that were partially summarized in Figures 1 and 2. The analysis employs the molar approximation between dµi and the change in the mean activity a (dµi) 2RT d ln a) such that the slope of a plot of Γi versus 1/ln[SE] (SE is supporting electrolyte) yields an estimate of dγ. Values for k′ were used to calculate Γi via eq 6; both are listed in Table 3 for TS as the solute and NaF as the electrolyte. Plots of ΓTS versus 1/ln[NaF] are given in Figure 4 at several values of Eapp. The plots contain three key points. First, the plots follow a linear dependence of ΓTS with respect to 1/ln[NaF]. This dependence validates the use of the Gibbs adsorption isotherm in dissecting the retention data via changes in electrostatic interactions. Second, at a constant NaF concentration, ΓTS increases as Eapp becomes more positive. This behavior is consistent with the trends in the chromatographic data in that an increase in retention corresponds to an increase in ΓTS. Third, the slopes of the plots become increasingly positive as Eapp moves to more positive values. This trend shows that the sensitivity of retention to electrolyte concentration is dependent on Eapp and parallels the observation of the convergence points in Figure 3. The inversion of the slope in Figure 4, which occurs between 0 and -50 mV, reveals the existence of an intermediate value of Eapp in which ΓTS is independent of electrolyte concentration. Importantly, this results agree well with the value of Eapp for the point of convergence in Figure 3A. The Lippmann equation predicts this occurrence, which corresponds to the value of the

Table 3. Retention Factors (k′)a and Interfacial Excesses (Γi)a for a 10-µL Injection of 50 mM Toluenesulfonate as a Function of Applied Potential (Eapp) and NaF Concentrationb NaF concentration (mM) Eapp

40

80

120

k′

Γ (10-16 mol cm-2)

k′

Γ (10-16 mol cm-2)

k′

Γ (10-16 mol cm-2)

6.78 ( 0.09 5.31 ( 0.06 4.29 ( 0.02 3.55 ( 0.01 2.89 ( 0.01

9.8 ( 0.4 9.0 ( 0.1 8.2 ( 0.1 7.4 ( 0.2 6.4 ( 0.1

6.61 ( 0.07 5.27 ( 0.03 4.36 ( 0.01 3.74 ( 0.02 3.09 ( 0.01

9.7 ( 0.3 9.0 ( 0.1 8.3 ( 0.1 7.6 ( 0.2 6.7 ( 0.1

6.50 ( 0.11 5.19 ( 0.04 4.34 ( 0.03 3.74 ( 0.02 3.17 ( 0.01

9.7 ( 1.3 8.9 ( 0.7 8.2 ( 0.5 7.6 ( 0.2 6.9 ( 0.1

(mV)

k′

Γ (10-16 mol cm-2)

50 0 -50 -100 -150

7.14 ( 0.09 5.46 ( 0.07 4.27 ( 0.04 3.27 ( 0.01 2.54 ( 0.01

9.9 ( 0.3 9.1 ( 0.1 8.2 ( 0.1 7.0 ( 0.2 5.7 ( 0.2

a

160

Average values from five replicate injections. b Mobile phase: 90% water-10% acetonitrile.

Figure 4. Plot of Γi versus 1/ln [NaF] for TS at multiple Eapp (versus Ag/AgCl saturated NaCl). Each data point represnts the average first statistical moments for the TS elution profile determined in five replicate injections with error bars on the order of the size of the data points. Slopes and the correlation coefficients: [, (1.19 × 10-16 mV-1, R2 ) 0.968); 9, (7.36 × 10-17 mV-1, R 2 ) 0.953); 2, (-3.54 × 10-17 mV-1, R 2 ) 0.751); 0, (-2.68 × 10-16 mV-1, R 2 ) 0.877); X, (-4.77 × 10-16 mV-1, R 2 ) 0.918).

PZC with TS as the probe molecule. As noted earlier, this result is consistent with the location of the PZC (+100 to -200 mV (vs Ag/AgCl/saturated KCl)) for carbon electrodes in earlier reports.60 However, the breadth of the PZC values in the carbon literature is clearly influenced by the presence and type of surface oxygen groups (e.g., carboxylic acids, alcohols, quinones, and ketones). Available correlations indicate that the PZC at carbon shifts to more positive values the greater the amount of surface oxygen.60 Since PGC is devoid of detectable levels of oxygen, studies designed to examine directly the retention of ionic species that interact only by electrostatics (e.g., F-) represent an intriguing means for determining the effect of oxygen groups by systematic surface modification procedures.12 The slopes of the plots also demonstrate the role played by the anionic and cationic components of the electrolyte in determining ΓTS. When positive of the PZC (e.g., +50 mV), competition from F- will drive a decrease in ΓTS as [NaF] increases. On the other hand, an increase in ΓTS with increasing [NaF] is observed when negative of the PZC (e.g., -150 mV). In these cases, the solution side of the double layer is composed of cations that counter the excess negative charge on the electrode. The observed increase in ΓTS with [NaF] therefore reflects the presence of anions such as TS in response to the local excess of

cations. There may also be a contribution from ion pairing between TS- and Na+ due to the propensity of aromatic sulfonates to weakly associate in bulk solution with alkali metals, especially Na+.68 Finally, these slopes can be analyzed via eq 3 to determine -dγ as a function of Eapp, which yield the contribution of the analyte to the net charge on the interface through eq 2. The values for dγ resulting from an analysis of plots such as those in Figure 4 are given in Table 4 for each solute at multiple values of Eapp in each electrolyte. The values for dγ are, as expected, larger for 1,5-NDS and 1,3-BDS than for the monovalent solutes. However, this difference does not quantitatively correlate with the difference in valency.40 Furthermore, the values of dγ are ∼10-12 J cm-2, which are ∼10-5 smaller than expectations.69 To account for the large discrepancy, we hypothesized that the sensitivity of this technique to track dγ is, in part, defined by solute concentration. In other words, the values of Γi are based on the injection of a few picomoles of solute, which would have a negligibly small contribution to the overall surface excess of ions with respect to the 40-160 µmol of supporting electrolyte. To assess this possibility, the experimental values of k′ were used to estimate a value of dγ for an injection of 100 µmol of solute. This calculation, which assumed a constant distribution coefficient for the adsorption of ions at the higher solute concentration, predicts values of dγ ((0.1-1) × 10-7 J cm-2) that are consistent with expectations.69 As noted, variations in the slopes of plots in Figure 4 indicate that dγ changes with Eapp. This dependence is evident in Table 4 by the decrease in dγ for the aromatic sulfonates and an increase in dγ for PTMA that occur as Eapp becomes more cathodic. Interestingly, the values of Eapp at which the sign of dγ changes correspond to the convergence points in the capacity-potential plots and follow the trends in Table 2. The inversion in the sign of the slope in Figure 4 and of dγ in Table 4 can therefore be interpreted as crossing at the PZC. These correlations demonstrate the thermodynamic basis for using the convergence points for PZC determination. The PZC values in Table 2 can be employed in a rational potential scale for application of the Lippmann equation. The Lippmann equation states that for a given solution composition, temperature, and pressure, the charge on the electrode (qM) can (68) Boyd, G. E. In Ion Exch. Process Ind., Pap. Conf.; Hall, G. R., Ed.; The Society of Chemical Industry: London, 1970; pp 261-269. (69) Manciu, M.; Ruckenstein, E. Adv. Colloid Inerface Sci. 2003, 105, 63-101.

Analytical Chemistry, Vol. 77, No. 22, November 15, 2005

7405

Table 4. Change in Interfacial Excess Due to Changes in Supporting Electrolyte Concentration at Multiple Potentials (vs Ag/AgCl satd NaCl) for the Charged Analytesa supporting electrolyte NaF analyte

Eapp

1,5-NDS

50 0 -50 -100 -150 -200 -250 -300 -400 50 0 -50 -100 -150 -200 -250 -300 -400 50 0 -50 -100 -150 -200 -250 -300 -400 50 0 -50 -100 -150 -200 -250 -300 -400 50 0 -50 -100 -150 -200 -250 -300 -400 50 0 -50 -100 -150 -200 -250 -300 -400

CBS

TS

BS

1,3-BDS

PTMA

KCl

dγ (pJ/cm2)

charge (pC/cm2)

1.3 0.7 -0.1 -1.5 -2.3

-15.7 -20.5 -4.0 -22.7 -19.8

0.2 0.1 -0.1 -0.7 -1.3

0.6 0.4 -0.2 -1.3 -2.4

1.3 0.7 -0.1 -1.5 -2.3

3.9 2.5 0.5 -0.8 -2.1

-0.5 -0.5 0.5 1.5 3.5

-3.4 -7.9 -3.5 -8.4 -9.5

-8.1 -15.9 -6.5 -17.2 -18.6

-15.7 -20.5 -4.0 -22.7 -19.8

-32.4 -36.3 -23.7 -28.1 -26.2

8.52 51.7 26.0 16.1 24.6

LiClO4

dγ (pJ/cm2)

charge (pC/cm2)

1.7

-7.3

1.7

-8.0

1.9 0.5 0.0 -0.2 -0.2 -2.5

-8.7 -5.8 -4.4 -1.5 -2.3 -12.0

1.5

-13.5

0.7

-54.9

0.1 -3.2

-1.1 -16.7

1.0

-3.4

1.4

-4.1

0.9 0.8 0.6 0.5 -0.3 -1.0

-5.2 -7.1 n5.5 -0.1 -0.3 -9.0

0.6

-2.4

0.7

n4.5

1.1 -0.9

-21.8 -17.9

-9.0

2.6

-7.4

1.8 2.0 1.5 0.6 0.0 -1.5

-9.7 -4.2 -16.5 -14.6 -3.2 -13.3

0.9

-3.6

0.6

-3.8

0.2 -1.6

-30.2 -30.1

1.7

-7.3

1.7

-8.0

1.9 0.5 0.0 -0.2 -0.2 -2.5

-8.7 -5.8 -4.4 -1.5 -2.3 -12.0

1.5

-13.5

0.7

-54.9

0.1 -3.2

-1.1 -16.7

1.0

-10.3

-0.7

-3.7

0.8 -0.4 -0.5 -1.5 -1.5 -2.0

-8.9 -7.9 -5.0 -9.8 -7.4 -6.6

-2.5 -1.2 -1.2 -1.4

-29.4 -79.3 -10.2 -6.3

-1.9

38.1

0.0 0.1 0.7 1.1 1.6 2.0

0.6 0.5 4.9 5.7 6.4 5.6

2.61

dγ (pJ/cm2)

charge (pC/cm2)

a Values show