A
E1ectrocheMca.l Processes at Liauid Interfaces Pelr Vanysek Northern Illinois University Deprnenl of Chamistry DeKalb, IL 60115
More than 100 years ago, Nernst performed the first experiments that provide the theoretical basis for today's potentiometric and voltammetric studies of interfaces ( 1 ) . As the significance of interfaces has become more widely recognized, new techniques to probe them have evolved. Today, electrochemical experiments provide a better understanding of the significance of the potential difference that occurs a t liquid-liquid (LL)interfaces. These studies may have applications in electroanalytical chemistry, in separation science, and in the electrochemistry of biological membranes. This article will describe some applications of electrochemical techniques for use a t the L-L interface as analytical tools in ion determinations and outline some areas for future development. Equilibrium studies of two immiscible phases in contact are considered a standard part of separation science (e.g., extraction-based separations). Most often, the extracted species are molecular compounds, in which case no charge transport is involved. Any electric potential difference across the boundary between the two solvents can be justifiably ignored. Unfortunately, the practice of disregarding the interfacial potential difference is often extended to salt separations, where the potential effect is important. In salt ex0003-2700/90/0362-827A/$02.50/0
@ 1990 American Chemical Society
traction theory, the balance of charge is usually considered, but the more general theory that includes potential a t or across the interface and the resulting ion repartitioning is often disregarded. lntertace between two Immiscible electrolyte solutions (ITIES) The hebavior oftheITIES firstcame to the attention of electrocbemists when Koryta et al. (2) postulated that this interface should behave similarly to an electronically conductive electrode immersed in a solution. Previously, Blank and Feig (3)suggested that this interface also could serve, a t least to a rough approximation, as a model for one-half of a biological membrane. This concept
spurred the early work of French researchers, especially Gavach a t Montpellier (4). The bioloeical membrane is a selfassembled siructure of phospholipids with polar heads facing the aqueous intracellular and extracellular solutions. The lipophilic chains of the phospholipids form the oil-like inner layer of the membrane. One can consider an ITIES (here, oil and water) as half of the cross section of the membrane. The similarity in behavior between the ITIES and an electrode-electrolyte system will be explained later. A typical solvent pair used for ITIES studies is water and nitrobenzene. Nitrobenzene is the most commonly used
nonaqueous solvent because of its low mutual solubility with water and its high relative permittivity, 6 (34.8 a t 25 "C), which is required to support the dissociation of dissolved salts. To study charge transport across the interface both solutions must be conductive, and this is achieved by adding suitable supporting or base electrolytes. LiCl is a commonly used supporting electrolyte in water; the most commpnly used salt for the nonaqueous phase is tetrabutylammonium tetraphenylborate (TBATPB) because it dissolves in organic solvents. When the two immiscible solvents that contain supporting electrolytes are brought into contact, the supporting electrolyte
salts stay virtually in their original phases. Although some repartitioning occurs, its extent is negligible. The interfacial thermodynamic equilibrium of all the ions present can be described by equating their electrochemical potentials ( E ) in both phases. For each ion in the phases (Y and 13 we can write an equation of the form A&)
+ RT In ai(&) + nFq(a) =
&S) + RTln ai(@)+ nFv(I3) (1) where pp is the standard chemical potential of the ion i in each phase, ui is the activity of the ion, n is its charge, R is the universal gas constant, F is the Faraday constant, T is absolute tem-
ANALYTICAL CHEMISTRY. VOL. 62, NO. 15, AUGUST 1, I990
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perature, and B is the inner potential. If more than two ions are considered, we obtain a transcendental function for the interfacial potential drop & = I&) &)I. Despite the apparent complexity of Equation 1, it leads to derivation of the Nernst equation. Indeed, when partitioning of only one ion is considered, the result is formally equivalent to the Nernst equation. As a first approximation, we can assume that within a certain interfacial potential range, ions of supporting electrolytes do not participate appreciably in the equilibrium. Thus the electrical potential of the interface with only supporting electrolytes is not defined. This does not mean that the potential on the interface does not exist. In principle, its value can be calculated, but calculation of the potentialdetermining ion concentrations is a formidable task. (A similar case is that of a mercury electrode immersed in a deaerated solution of KCI.) An important property of thii interface is that if a potential is applied to it from an extraneous source, the interface will attain the potential of the source. The interface is then “ideally polarizable,” following the terminology that is traditionally applied to electrodes. A different situation occur8 if a salt soluble in both phases is added to the system. For example, lithium picrate is soluble in water and picrate anion is a h highly soluble in nitrobenzene. If the picrate anion is the only shared ion between water and nitrobenzene, the interfacial potential drop is
-
A~%%nzene
RTIF
0 water
= 4 nitrobenzene + nitrobenzene
bplcrate
water /~pimtel
(2)
and the potential is determined by the partition ratio of picrate anion between is the the two phases. standard potential of transfer of the individual ion and, for ITIES studies, its meaning is similar to the standard reduction potential in redox electrochemistry. The known values for A: are tabulated in Reference 5 and, for picrate ion, the standard potential is 47 mV. The potential is defined here as the difference between the aqueous and the oil phase; thus the sign of the potential correspondsto the polarity of the water phase. This simple relationship (2) is fdfiiled simultaneously for all ions that participate in the phase equilibrium. The difficult part is determining the individual ion activities. Equation 2, the Nernst-Donnan equation, describes the basic principle used in potentiometric determinations of ions with ion-selective electrodes (IS&). An important property of the interface follows from further analysis
of the relationship. If the interfacial potential is forced to a value given by the potential of an extraneous source, the ratio of the picrate ion concentration must change. If the water phase is made more positive, more picrate ion will be driven into the organic phase to fulfiil the equilibrium prescribed by Equation 2. As a result, a net electric current flows through the system and can he measured by a galvanometer in series with the completed circuit. This current is the basis for the voltammetric characterization of the interface. Experimentalconsiderations The electrochemical cell that is often used in ITIES polarization studies is shown in Figure 1. That portion of the cell where the interface under investigation is confined in the central, narrower tube is typically made from glass tubing. The center portion can be lined with a Teflon insert, which is sometimes used to define the exact position and flatten the meniscus of the interface. The existence of a meniscus rather than a flat interface is not a major problem, although it causes difficulties in calculations that rely on the area of the interface. A screw-driven plunger attached to the side ann of the cell adjusts the position of the interface. By adjusting the volume of the lower part of the cell, it is possible to adjust the interface position precisely. In any electrochemical experiment, the determination of the phase boundary potential difference requires the use of reference electrodes. Because the aqueous phase usually contains chloride ions, a silver wire coated with AgCl is the simplest choice for a reference electrode. For more demanding poten-
A:”,Ebenzme
828A
~~
-_.__
Flgure 1. Experimental _r ._. voltammetric o( potentiometric studies of the interface between two immiscible electrolyte solutions (ITIES).
ANALYTICAL CHEMISTRY, VOC. 62, NO. 15, AWJST 1. ISSO
tiometric measurements, it is possible to separate the reference electrode from the solution by a salt bridge. The reference potential for the nitrobenzene phase is usually maintained by a secondary water-nitrobenzene interface. This is a constant potential interface, often d e d a reference interface. The reference interface is set in such a way that both phases share a common ion. Typically, because the nitrobenzene already contains TBATPB, the aqueous phase contains tetrabutylammonium chloride. A silver wire coated with AgCl serves as a reference electrode connecting the solution to the outer electric circuit. The potential of the reference interface is determined by the shared ions. If one can assume that the only shared ion is TBA+, the interfacial potential will be given by an equation similar to Equation 2, in which the picrate ion activity is replaced by that of TBA+. The potential of this interface, even if a current of small density passes across it, stays constant during the measurement. For equal concentrations of TBA+ in both phases, the potential assigned to this interface is -248 mV, a value that has been calculated from extraction measurements. Because the potential compares activities of an ion in two dissimilar phases it is not a true thermodynamic constant, but it is useful as a constant potential reference point. The total potential of the electrochemical cell that includes the reference interface contribution is denoted Li in t h i article. The two reference electrodes and the reference interface are sufficient for potentiometric work. In voltammetry, however, it is desirable to use a potentiostat. Because compensation of two resistive solutions is necessary, a special four-electrode potentiostat that requires two additional electrodes in the cell is used. These counter electrodes are typically platinum flags or wires. To keep the current density low, a large surface is desirable. However, one must remember that redox processes are taking place on these electrodes and high overpotential would needlessly increase demand on the compliance voltage of the potentiostat. The redox processes on the counter electrodes involve decomposition of the supporting electrolytes and, sometimes, of the solvents as well. A glasa frit separates the electrodes from the rest of the cell to avoid contamination of the interfacial area by the reaction products. Difference between lTlES chargetransfer and electrode reactions
The similarities between the processes at the ITIES and on metal electrodes
have been emphasized so often (2)that we will concentrate on the differenees instead. Current flow through an electrochemical cell is u s d y associated with a redox process occurring in the cell. However, when the ITIES is polarized by an external potential source, the net current flow doea not result from a redox process at the interface. In fact, with a few rare exceptions, there is no redox reaction taking place at the L L interface. Figure 2 compares a current flow on an electrode and at the ITIES.I f a sufficiently negative potential is applied to the electrode (Figure Za), the system will have high enough energy to cause the reduction of some reducible species (here, Fe3+)in the solvent. An electron will leave the electrode and cause reduction, the reducible species will &appear from the solution, and a reduced species (Fez+)will appear on the surface of the electrode. Application of a positive potential would reverne the process, causing oxidation of the reduced form. The magnitude of the current depends on the rate at which the reducible species crosses the electrode surface, the rate of the charge-transfer reaction, and the rate of diffusion of the reduced species away from the electrode. If the solution is unstirred and contains enough supporting electrolytes 80 that the analyte ions move only by difhion, and if the charge-transfer reaction is faster than the diffusion it-
Flgm 2. Comparison of (a) the interface between an electronicallyconducthre electrode and a solution (reduction of Fe3+)end (b) the ITiES (transport ot picrate) during current flow in a closed electric circuit.
self, the current can he calculated from
Fick’s diffusion equations. If a negative potential is applied to the nonaqueous phase (Figure 2b) of the ITIES cell and the system contains an anion that can be transported (such as picrate), the anion will start crossing the interface upward. If the system is not stirred and enough supporting electrolyte is present to suppress migration of the picrate anion, the total charge flow in the system consists of the anion diffusion to the interface, the anion erossing of the interface, and finally, diffusion of the anion away from the interface. The two counter electrodes complete the loop with the outer circuit, and the overall current can be measured by a meter inserted in the circuit. If the ion c r m s the interface rapidly, diffusion of the ion toward and away from the interface will be the rate-determining stepa. Inverting the polarity causes transport of picrate anion in the opposite direction. A similar argument, but with opposite transport directions, can be made for cations. The major d~ferencebetween the ITIES and a redox electrode is that at the interface, the charge of the ion remains unchanged and its diffusion proceeds in one duection in two different solvents. In electrode kineties two different species (oxidized and reduced) are involved and their diffusion is considered in one solvent only, but in opposite directions. The forms of the diffusion equations for either case are the same; only the values of the diffusion coefficients will be different. An example of even greater similarity is the case of polarographic reduction of Cd*+. One can think of cadmium metal dissolved in mercury as a cadmium ion and two electrons. In that m e there is no difference between the ITIES case and the electrode. Although redox proceeaea generally do not occur at the ITIES,they are not absent from the system. As long as current flows through the cell, oxidation of the solvent and available compounds OCCLVS at the anodic counter electrode and reduction o m at the cathode. Thii situation is the same as in any other use of a potentiostat. If the reaction products are prevented from reaching the interface, the processes can be overlooked. As long as a parallel mechanism describing the behavior of the ITIESand metal+lectrolyte system can be considered, any experimental technique applicable to one system can be used in the other system. Voltamnetry on an ekctmiyle drop This technique is derived from polarography but it involves an electrolyte
and scientific data with either a 1-2-3 OT Symphony spreadsheet where it 1s ready for immedsate analysis, storage. or display. Measure accepts data from IEEE-488and RS-232 instruments and plug-in data
drop rather than a mercury dropping electrode (2, 6). Nitrobenzene is normally used as the stationary phase in the cell (Figure 3; Reference 7). The aqueous drop that forms the interface of interest ascends after it dislodges from the Teflon orifice that restricts the solvent flow. Figure 4 illustrates a voltammetric wave of perchlorate transfer from water to nitrobenzene; crystal violet tetraphenylborate is the nonaqueous supporting electrolyte. The advantage of this arrangement is the renewal of a fresh interface every time a new drop is formed. Problems from adsorption and surface precipitation are thus eliminated. The drop time can also be used as a measure of the interfacial tension (8,9). Two drawbacks are that the experiment is not easy and there is a large electrical resistance in the orifice opening, which complicates the electricalmeasurement. The equation describing the instantaneous current i in amperes is
i = 4.17 x 1 0 - 3 ~ ~ (-~ c) 0 x p-213~l/2m213~1/6
(3)
where p is the density of the dropping electrolyte in g cm-', D is the diffusion
coefficient of the analyte in ern2 s-l, m is the flow of the electrolyte from the capillary in g s-1, t is the time in seconds, co is the analyte concentration in the bulk solution in mol L-I, c is the analyte concentration near the interface, and n is the charge of the ion. For a limiting current, the surface concentration equals zero.
Flgure 3. Apparatus for experiments on ascendlngdrop interfaces. The &op 16 lormed on a Teflon wilice immeraed in me IwnwuBous phase, and elfluem solution accumulates in the cmparmxmt. COunter eiectrodes are separated from the siectmactive sobtions by glass hits. me weter liow hom the re=voir is controiied by a gauge on ltw restrictor. (Adapted horn Reference 7; reprinted bv m i s sion 01 me EiectrachemiceiSociety.)
of perchlorate transfer from water to nitro-
Figure 4. Voltammetric wave
benzene. condnions: [ C w - ] = 5 X 10-'moi L-': supw i n g eiecboiyies: 1 mol L-' LiCi in water and 0.05 mal L-' crystal violet tebaphenyiborate in nitrobenzene: suln rate: 100 mV min-'.
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This equation is nearly the same as the IlkoviE equation for polarography; the only difference is in the explicit inclusion of the solvent density. The numerical constant in the original IlkoviE equation contains the density of mercury implicitly in the numerical constant. The analytical application of ascend ing-drop ITIES is similar to applica tions of polarography. The height 01 the wave is proportional to the concentration, and the half-wave potential corresponds to the potential of transfer of the studied ion, ;;Eknzene. Cyclic voltammetry Voltammetric studies at the interface are generally performed with the cell shown in Figure 1;a four-electrode potentiostat is the preferred instrument for polarization experiments. Figure 5 shows an example of a voltammetri curve at the ITIES. Curve a shows th' response to the transfer of picrate ion from nitrobenzene to water (positive peak) and back. The usual criteria for reversibility, such as forward and reverse peak separation of 58ln mV, independence of peak potentials on the sweep rate, and the ratio of positive to negative peak current equal to one, still
apply. A reversible process in ITIES studies means that the charge transfer a c r m the L L interface is faster than diffusion or, in other words, the current flow is governed entirely by analyte ion diffusion. Actually, it appears that the
Figure 5. Voltammetric reversible transport of picrate ion between walL and nitrobenzene. Cuve a: 0.8 m m i L r ' tetrabutylammnlum pirate in niuobenzene; suppwtlng eiecboiyles: 0.02 m i L-' LiCl in water and TBATB in nihobenzene; scan rate: 100 mV s-'. Cuve b sup porting electrolytes miy.
typical rate constants for ion transfer are too fast to be measured by this technique. The current measured at the peak of the curve (i,) is given by
i,
= 269 c 0 n 3 / 2 ~ 1 / 2 U 1 / 2
(4) The scan rate u is expressed in V 5-1, A is the surface area in cm2,and the other symbols are defined in Equation 3. Curve b is a response to the supporting electrolytes only. Here the difference between the ITIES and electrode behavior is most prominent. There are two reasons for the larger base current. First, the polarization window is always narrower as supporting electrolyte ions participate in the transport near its edges. Second, a small amount of ion repartitioning takes place as the potential is scanned. Current scan technique As a rule, voltammetry at L L interfaces requires use of a potentiostat. Although this technique, improved even more by a positive-feedback circuit, works well in cyclic voltammetry, application to voltammetry with ascending electrolyte drop often leads to overcompensation. An alternative to the voltammetric experiment is voltam-
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ANALYTICAL CHEMISTRY, VOL. 62, NO. 15, AUGUST 1. I990
891 A
INSTRUMENTATION 60-s polarization of a thin organic layer
400 -200
0
2cQ
4w
eo
Figure 8. Currem scan voltammogram
at an ascendingwater eIectrodB. (a) Supporting elecbolyte response on a waterdichlaoehane ihterlafsce. (b) Result 01 transfer of 0.3mmol L-' l-phenylsmeMyi4biliuomacetyl5-pyra2olone anion horn the aqueous phase into the organic phase. (Adapted Irom Reference 8.)
metry with current scan pioneered by Freiser and Sinru (8),Freiser and Yoshida (10, 1 0 , and Kihara et al. (12). The advantage of this approach over the usual potential scan technique is that the correction for uncompensated resistance can be done hy subtracting the iR drop from the measured potential. Figure 6 shows the result from a current scan technique experiment applied in a voltammetric study on an ascending-water electrode (8).
or a small hanging drop in contact with a stirred aqueous solution (14). Under properly chosen experimental conditions, the layer becomes enriched in the analyte. This step is followed bydifferentia1 pulse voltammetry with inverse polarity, in which the anal*. now at much higher concentration in the organic layer, transfers into the aqueous phase. The detection limit so far reported is 1 @molL-l. This value conceivably could he lowered by extraction into thin-gel, solidified, spin-coated nonaqueous layers. Potentiimetry Potentiometric measurements at L L interfaces have been the suhjed of studies in the related field of ISEs for some time. The contributions of ITIES to ISEs are isolation and study of one interfacial potential difference at a time. Interpretation requires detailed analysis of the equilibrium conditions prescribed by Equation 1. This equation, rewritten in terms of equilibrium interfacial potential, A;@; potentials of transfer of individual ions, A@; total analytical concentration of the ion, c: equal volumes of the phases; activity coefficients, y; and charge (with sign) of the species, zi, has the form
ac Voltammetry A technique that detects changes in faradaic impedance in relationship to applied de potential is ac voltammetry. Ordinarily the dc potential is varied linearly on a long time scale (-25 mV s-l), and an ac component (-10 mV, f = 10 Hz to 100 kHz) is superimposed on the signal. The output is a plot of the ac component of the current versus the applied de potential. Because it provides very low detection limits, ac voltammetry is an interesting technique for analytical applications. It is possible to determine analytes at concentrations that are lower than 10 pmol L-I, and analytes in mixtures can he determined simultaneously. Figure I,for example, illustrates that nitrate, perchlorate, and cesium ions can be determined together. Stripping voltammetry Equation 1 implies that an ion can he driven from one phase into another at a sufficiently high potential. A more detailed analysis confirms that this process does not violate eledroneutrality. The excess charge transfer is compensated for by redox reactions on the counter electrodes. The stripping technique involves a
.
interfacial potential dependence as a function of the Concentration of picrate ions in aqueous and
nitrobenzene sotutions. The nitrobenzene phase wntains dissolved tetrabulylammonium piuate and 0.01 mol L-' TBATPB; me aqueous phase w n t a i ~ pioric acid and 0.01 mol L'' LICI. The calcu$tlons are lor T = 25 OC.
trobenzene and picric acid in the aqueous phase. This system is, in essence, a picrate-sensitive electrode. The threedimensional perspective allows illustration of the effect of varying picrate concentration in both water and nitrobenzene. The graph clearly shows the region of the low aqueous picrate concentration in which the potential is no longer a function of [Pi-]. The potential in that redon is eoverned bv the equilibrium of the supporting electroIvtes. At hieh concentrations the levelihg of the-dependence indicates the Donnan failure (16,17). In the Donuan failure region only a single picrate salt participates in the equilibrium, and the potential is no longer dependent on the total picrate concentration. The linear part is sometimesdescribed by Donnan exclusion, a process in which only one ion (here, picrate) determines primarily the net potential. As the picrate concentration goes down, the contribution of the supporting electrolytes to the overall potentiaJ becomes significant and, eventually, the only potential-determining factor. The potential profile in Figure 8 corresponds to those observed for IS&. With an appropriate modification, Equation 5 can be rewritten, for example, in the form of the Nicholsky-Eisenman equation for ISEs and other two-phase ion-containing systems. 1
2F
exp
(Agv - A;d)1] = 0
Lnl
(5)
JJ
This equation, first presented hy Hung (15), allows evaluation of the POtential response of the interface to varying ion concentrations. Figure 8 eives a response for a system containh g tetrabutylammonium picrate in ni-
Figuro 7. Simultaneous determination of nitrate. perchlorate, and cesium ions at the !TIES using ac voltammetry. Favard a. S w e w ram 25 mV s-': I = 55 HI. supporting slecholyns 0 01 mol L-' LI,SO, in ustu and 0.01 mol L-' wsmi v~olet Ieuapwlbaate Concernelionsof NO; and Cs' we 0 2 mmOi L-!, conmntratim ofCiO; is 0.1 mmol L - l Dashed line repasen16 ms sup. paring ei&olytss. (Awl&wiih pemisbn hom Relsnrnaa IS 1
832A * ANALYTICAL CHEMISTRY. VOL. 62, NO. 15, AUGUST 1. 1990
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lons that can be determined The analytical usefulness of the ITIES polarization relies on determining as many different ions as possible. Therefore, for a given electrolyte pair it is usually desirable to choose a solvent pair that has the broadest polarization
potential window. For this purpose water-nitrobenzene is the best choice, followed by water-1,2-dicbloroethane. Besides these solvent pairs, the use of water4lvent mixtures such as nitrobenzene-benzonitrile and nitrobenzene-benzene also has been reported (1419).Although these mixtures provide narrower potential windows, they often exhibit selective solvation properties, which influence rates of chemical reactions, solubilities, and so on. It is expected that choices of solvents could help the understanding of ion transfer or electron transfer mechanisms and could lead to the design of systems of preferential or even selective response. Generally it is possible to determine ions that are less hydrophilic than the ions used as supporting electrolytes in water and, at the same time, that are lea hydrophobic than the ions of the salt used as the nonaqueoussupporting electrolyte. It is also obvious that the ions to be determined must not react with the supporting electrolytes or the solvents. From this perspective the most serious obstacle is the low solubility product of certain combinations of ions. A salt can DreciDitate during a measurement on &e interface eithe;as a result of natural repanitionina or as a result of ion transport caused by current flow from an outside source. The precipitate adsorbs on the interface and modifies its properties. As a solid phase, it also functions as a sink for the transported ions, so reaching the equilibrium condition prescribed by Equation 2 may be impossible. Many of the ions that can be determined by the techniques described for L L interfaces are listed in the box below. An extensive list can be found in Reference 5 and in subsequent reviews (20,21).Lithium and sodium ions cannot be determined directly, and they cannot be determined if LiCl is used as a supporting electrolyte. Na+ and, to
some extent, K+ have characteristics too close to Li+ to exhibit peaks that are distinguishable from the supporting electrolyte background. It is still possible to detect transport of these ions if proper steps are taken to form a more hydrophobic ion complex and, thus. to shift the ion Dotential of transfer within the workLng window of the supporting electrolytes. An example of this process is the facilitated transport of sodium ion in the presence of 18-dibenzo-crown-6 (22)or the case of potassium ion and valinomycin (23).If a specific acceptor is found, the technique even becomes selective to the ligand ion. The appropriate complexing agent can then be used to exoand the number of ions determined. donverselv, the techniaue can be modified for determination of the complexing agents, as in an assay of monensin, which is based on the complexing reaction with sodium (24). More recent work describes determination of Pb*+ using complexing properties of polyethylene glycol 400 (25). Microinletface between two immiscible solutions The DurDose of usine a small orifice to restrict ihe area of The interface is u, take advantage of recent Drwress in the development of ultramieroelectrodes. Ultramicroelectrodes help to overcome complications stemming from a potential shift arising from an iR drop. As the interfacial area becomes smaller, the diffusion geome* takes on the character of a spherically symmetric process. This means that the ratio of faradaic current versus solution resistance is increasing and, in the end, rendering the contribution of the iR drop minimal. ITIES, which uses solutions of low conductivity, can benefit from this phenomenon. Restricting the interfacial area and using a current amplifier is an alternative to a four-electrode potentiostat.
I Selected kns that can be determined by L-L interface polarizatkm ) r l J M k MtlOM ..
Natemsry ammnlum cations
Rb+
(variwsalkylchalnsshwterthanbutyl) %line
I
Zcetylchollne
I
CS+
Fb" (
Na+ (only as a complex) LI* (only LU a complex)
~
1
Figure 9. Voltammetric determini of lauryl sulfate (0.4 m m l L-') on a
L-L microlntetface. Scan rates (mv sP)(: 1 ) IO. (2) 20, (3) 50. (4) 100, (5) 200. (0) 500, and (7) 1000. Suppwllng electrolytes: 0.02mol L-' LiCl In water and 0.02 mol L-' TBATPB In nitrobenzene. Diameter of the h i e Is 130 pm.
Small L L interfaces have been used by Girault and co-workers (26-28)and bv Senda et al. (29).We have also exdored a small interface formed in a Figure hole of a thin glass wall (30,31). 9 illustrates voltammetric curves obtained on such an interface. One complication of the microinterface compared with the solid microelectrode is that it is difficult to keep the interface small. Its size is defined by a small window in a thin plate. Interfacial tension, capillary pressure, uneven hydrostatic pressure, thermal expansion, and even vibrations in the laboratory cause the interface to move away from this areadefining region. We have tried to solve this problem by immobilizing the electrolyte solution inside a tube holder by adding an immobilizing agent. Future directions Ongoing investigations of the electrical behavior of the ITIES follow two main directions. The fist, and so far predominant, is research in theoretical electrochemistry. Here, for example, the recent effort by Berube and Buck (32)to elucidate anomalies in theories of liquid membrane time response to a concentration step should be mentioned. Even though the work is theoretical, the results have an impact on applied work. An extension of results (32)can be applied to understand the observed difference between electrodes and the ITIES, such as the high SUPporting electrolyte current in Figure 5b.
ANALYTICAL CHEMISTRY, VOL. 62. NO. 15. AUGUST I , 1990
8 3 8 ~
Some experimental techniques applicable to the ITIES are also of a more fundamental nature. Fast scan voltammetry, in particular on microinterfaces, can be used for determination of charge-transfer rate constants. Impedance analysis, which is an extension to ac voltammetry, can be used not only to detect analytes, but also to obtain a better understanding of surface phenomena (33) and adsorption (34). A special case of impedance analysis is derived from measurement of noise generated by the electrochemical system (35).This is especially useful for work on microinterfaces that have large absolute resistance. The second direction of interest is in application to new or improved analytical sensors. These can be potentiometric, voltammetric, or based on other appropriate techniques. The L-L interface is actually a component of an ionselective electrode with a liquid ion exchanger. Recent analytical applications have resulted in construction and systematic studies of microinterfaces solidified by gels. The advantage of such a modification is ease of handling (36-39). The immobilization can be extended further to studies of frozen in-
'
terfaces, or even to use of solid electrolytes. Significantly, ITIES theory also applies to interfaces that are encountered in ion-doped, conductive, polymer-coated electrodes. This work and the author's research work described herein were supported in part by the Office ot Saval Research. I thank Elizabeth A. Burton for review of the draft version
(11) Yoshida. Z.: Freiser. H. Inore. Chem.
(12) Kihara, S.; Yoshida, Z.; Fujinaga, T.
Bunseki Kagaku 1982,31, E297-E300. (13) Hundhammer, B.; Solomon, T.; Alemu. H. J . Electroanal. Chem. 1983. 149.179-83. - -(14)-MareEek, V.; Samec, Z. Anal. Lett. 1981,14,1241-53. (15) Hung, L. Q. J . Electroanal. Chem. 1980,115,159-74. (16) Melrov. 0. R.: Buck. R. P. J. Electroanal. chem. 1983,143,23-36. (17) Cosofret, V. V.; Buck, R. P. In Fundamentals and Applications of Drug Sensors; Schuetzle, D.; Haemmerli, R., Eds.; ACS Symposium Series 309; American Chemical Society: Washington, DC, 1986; pp. 362-72. (18) Koczorowski, Z.; Paleska, I.; Geblewicz,G. J . Electroanal. Chem. 1984, 164,201-4. (19) Solomon. T.: Alemu. H.: Hundham. mer, B. J . Electroanal. Chem. 1984, 169, 311-14. (20) Koryta, J. In The Interface Structure
.-
~I
References (1) Nernst, W. 2.Phys. Chem. 1888,2,613"O
31.
(2) Koryta, J.; Vanfsek, P.; Bfezina, M. J. Electroanal. Chem. 1976,67, 263-66. (3) Blank, M.; Feig, S. Science 1963, 95, 561-72. (4) Gavach, C. C. R. Acad. Sci. (Paris)1969, 269,1356-59. (5) Vanfsek, P. Electrochemistry on Liquidlliquid Interfaces; Springer: Berlin, 1 QRii _--I. (6) Koryta, J.; Vanfsek, P.; Bfezina, M. J . ElectroanaL Chem. 1977.75.211-28. (7) Vanfsek, P.; Buck, R. P. 3. Electrochem. SOC.1984,131,1792-96. (8) Sinru, L.; Freiser, H. Anal. Chem. 1987, 59,2834-38. (9) Freiser. H. Abstracts of PaDers. 40th Pittsburgh Conference &d Expokition, Atlanta, GA; Pittsburgh Conference and Exposition: Pittsburgh, PA, 1989; Abstract 217. (10) Yoshida, Z.; Freiser, H. J . Electroanal. Chem. 1984,162,307-19.
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and Electrochemical Processes at the Boundary between Two Immiscible Liquids;Kazarinov, V. E., Ed.; Springer: Berlin, 1987.
(21) Girault, H.H.J.; Schiffrin, D. J. In
Electroanalytical Chemistry;Bard, A. J.,
Ed.; Marcel Dekker: New York, 1989; Vol. 15. (22) Hofmanovl, A,; Hung, L.Q.; Khalil, M. W. J. Electroanal. Chem. 1982, 135, 257-64. (23) Homolka, D.; Hung, L. Q.; Hofmanovl, A.; Khalil, M. W.; Koryta, J.;MareEek, V.;
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Samec, %Sen, S. K.; VanBek, P.; Weber, J.;BZezina,M.; Janda, M.;Stibor, I. Anal.
Chem. 1980,52,160610. (24) Koryta, J.; Ruth, W.;
[email protected], P.; Hofmanovh, A. Anal. Lett. 1982,15,168&992. (25) Sun, Z.; VanBek, P. Anal. Chim. Acta 1990,228,24149.
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(26) Taylor, G.; Girault, H.H.J.J. Elec-
troanal. Chem. 1986,208,17943. (27) Campbell, J. A,; Stewart,A. A.; Girault, H.H.J. J. Chem. Soc., Faraday
Trans.I1989,85,8453.
(28) Campbell, J. A,; Girault, H.H. J. Elec-
troonal. Chem. 1989,266,465-69.
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electrochemical Symposium,Mhtrafike$
Akademiai Kiado: Budapest, 1986; pp. 353-64. (30) VanBek, P.; Hernandez,
I. C. Anal. Lett., in preas. (31) Vanjwk, P.; Hernandez, I. C.; Xu,J. Microchem. J. 1990,41,327-39. (32) Berube, T. R.; Buck, R. P. Anal. Lett.
1989,22,1221-35. (33) Buck, R. P. Ion Select. Electr. Re". 1982,4,3-74. (34) Vanl.sek, P.; Sun, 2. Bioeleetroehem. Bioenerget. 1990,23,177-94. (35) Bezegb, A.; Janata, J. Anal. Chem. 1981,59,494 A-508 A.
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(36) Kakutani, T.;Ohkouchi, T.;Oeakai, T.;Kakiuchi, T.Anal. Sci. 1985, I , 21976
(37) MareCek, V.; Jhcbenovh, H.;Colombini. M. P.; Papoff, P. J. Electrwml. Chem. 1981.217,213-l9. (38) Baum, J. Anal. Lell. 1970,3,10511. (39) Osakg, T.;Kakutani, T.; Senda. M. Bunsekr Kogabu 1984.33, E371-E377.
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