AMI. them. 1991, 63,1158-1164
1158 (18) (10) (20) (21)
Otaka, K.; Tambe, S.; Ando, T . J . -tow. 1987,396,350-354. Fujwara. 5.; Honda, S. Ann/. Chem. 1888, 58, 1811-1814. Robe, D. J.; JWgWISOn, J. W. AMI. chsm.1888, 80, 642-648. schwa&, n. E.; hiera, M.; B o w n h , R. G. J . "a*.1868, 480, 129-139.
RECEIVED for review December 26,1990. Accepted March 1,
1991. This work is supported in part by Grant DE-FGOB86ER13487 from the U.S. Department of Energy, Office of Energy Research, by Grant CHE-8901382 from the U.S. National Science Foundation, and by the cooperative agreement between the University of Tennessee and Oak Ridge National Laboratory.
Effects of Surfactants on Cathodic Stripping Determination of Iodide A. R.Harman and A. S.Baranski* Department of Chemjstry, University of Saskatchewan, Saskatoon, Saskatchewan, Canada SN7 0 WO
The d f ~ &of Trtton X-100 On kdkle m k n etltckncy, the shape of the strlpping peak, and the multlng callkatbn pbb were studied under cycllc vokanwnetrlc coditlam. Thdynamk parameters of mercurous iodide adwrptlon were determined in the presence and in the absence of the wrfactant by a new method based on the convolutbn transform. The rebulk rhow that the M#ltkn of lrtton X-100 affects only thermodynamic propertles of the system In the interfacial regbn. The pmence of the surfactant Improves signlfkantly the detection ilmit of the stripping determlnation of Iodide because the efficiency of lodlde deposition Ir increased due to a stronger adsorption of mercurous iodide on mercury electrode8 (a synergk effect), the stripping peak Is narrower (thereforehlghw and easier to measure) because the energy of admrptbn I s h affectedby potonilal, and thebackground current Is reduced due to the lower capacitance of the doubk layer.
INTRODUCTION The analysis of halides, especially at trace and ultratrace levels, is becoming increasingly important in various areas of human activities. For example such analysis is often undertaken in the food industry (1-3), where levels of iodide are of particular concern, and in the analysis of environmental samples such as natural waters and particulate air pollution (4-8).
The methods most often employed for the determination of halides, and in particular iodide, include gas chromatography (9),liquid chromatography (IO),ion-selective electrodes ( 2 , 6 , 9 ) ,flow injection spectrophotometry (11,121, and differential pulse polarography (13). When it is necessary to analyze for halides at the ultratrace level, the method usually employed is neutron activation analysis (NAA). Although this technique offers both the sensitivity and precision needed (14), it suffers from the disadvantage of the very high capital expense associated with the required facilities. This has led a number of researchers to evaluate the use of cathodic stripping voltammetry (CSV) as an alternative method, incorporating either linear sweep (151, differential pulse (I, 15,17), or square wave (16) techniques, with working electrodes made of either silver (17-19), mercury (1, 15, 16,
* To whom correspondence should be addressed.
20),or copper amalgam (21). Such methods of analysis have
proven to have the advantage of experimental simplicity, sensitivity, and low unit cost of equipment. In a recent publication, Luther et al. (16) reported that cathodic stripping analysis carried out in the presence of low concentrations of the surfactant Triton X-100(TX) enhanced the stripping peak of mercurous iodide under square wave voltammetric conditions. The use of this method enabled determinations of iodide in environmental samples a t concentrations more than 1 order of magnitude below previously reported detection limits for either electrochemical techniques (22) or ion chromatography (23). The discovery made by Luther et al. (16) is extremely interesting because in most cases the adsorption of organic molecules on mercury electrodes causes a decrease in sensitivity due to inhibition of charge-transfer processes. In order to better understand the action of these surfactants in cathodic stripping analysis, a detailed study was undertaken. In this publication, we report on the possible mechanism by which cathodic stripping peak enhancement is possible by the presence of adsorbed surfactant on the surface of a mercury working electrode.
EXPERIMENTAL SECTION All experiments were performed by using a conventional three-electrode arrangement. The reference electrode used was either a commercial saturated calomel electrode (SCE) or a laboratory fabricated Pb/PbSO,(,, electrode; however all potentials reported in this work are vs SCE. The auxiliary electrode consisted of a platinum wire with surface area ca. 0.5 cm2. With the exception of the hanging mercury drop electrode (HMDE) (Metrohm Model 6.03351,all other electrodes used in this work were laboratory fabricated, as described in a previous publication (24.
The equipment used for cathodic stripping has been described elsewhere (24). Altemating current impedance experiments were performed by using a EG&G PAR Model 273 potentiostat coupled to a Model 5301 lock-in amplifier. In measurements requiring high sensitivity, a built-in current follower of the Model 273 potentiostatwas bypassed and replaced with a EG&G Model 181 current-sensitivepreamplifier. Data acquisition and analysis was performed by using an AT&T Model 6300 microcomputer. All computer programs for data acquisition, processing, and numerical simulations were devised in this laboratory (copies of numerical simulation programs can be obtained from authors). All chemicals used were of analytical grade. The stock solutions (0.1 M) were prepared by using water distilled in a Corning Mega-Pure system, and all dilute solutions were made just prior to their use. The surfactants employed in this study were obtained
0003-2700/9 110383-1158$02.50/0 0 1991 American Chemical Soclety
ANALYTICAL CHEMISTRY, VOL. 03, NO. 11, JUNE 1, 1991
1158
T
-400 -100
0
100
200
300
Deposition Potential /mV
1. Effect Of Trlton X-100 addtkn On ~ c u o l kdlde ~ s cdlectkn efflclency. Experimental conditions: HMM (0.0087 cm2), 00-s deposition with stirring, stripping rate 50 mVls, soiutlon 0.1 M KNO containing 2.0 x 10-7 M I- ( w r w A-C, ~ ieft+~nd scale) OT 1.0 x i odl M I- (curve D, right-hand scale). Concentration of the surfactant: (0) (*) 2.4 X lod M. 0, (A)8 X
from Sigma and used without further purification. All analyte solutions were degassed for 20 min with oxygen-free argon prior to measurements. During all experiments, a blanket of argon was maintained over the analyte solution. Blank experiments run in the absence of the depolarizer showed no appreciable contamination of either the supporting electrolyte or the surfactant solutions.
RESULTS It has been recently shown by Luther et al. (16)that addition of TX to the analyte solution increases the stripping peak of mercurous iodide. Although the presented results were only for the determination of iodide, it was expected that similar results should be observed for the other halides. Experiments conducted in this laboratory confirmed Luther et al.'s observations for iodide; however, a similar increase in peak height was not observed under similar conditions for the other halides, thiocyanate, or any sulfur-containingcompounds when stripped from either a mercury or silver working electrode. Interestingly, the effect disappeared even for iodide when the mercury electrode was replaced by one containing copper amalgam. In order to further study the effect of surfactants in cathodic stripping voltammetry, experiments were performed in which the electrical double-layer capacitance of the working electrode was measured as a function of applied potential both with and without the addition of various surfactants. Under the experimental conditions used in this work, two tensammetric peaks caused by the adsorption/desorption of the organic molecule were observed at -1600 and +25 mV, respectively. These mark the potential range over which the surfactant is adsorbed on the surface. The fact that, in the case of TX, the surfactant desorbs at +25 mV could explain why no increase in the stripping peak was observed for either bromide or chloride, since the deposition of these anions must be carried out at potentials more positive than +25 mV. Similar measurements performed with a silver working electrode showed that both the deposition and stripping potentials of the silver halides lie outside the adsorption range of all surfactants studied. As a result of this, no further work was conducted with these electrodes. Additional double-layer measurements were carried out to check for surfactants with the adsorption range extending over more positive potentials than that of TX. Although a number of neutral and anionic surfactants were investigated, all were found to desorb from the mercury surface at potentials too negative for the deposition of either mercurous bromide or chloride.
-200
0
200
400
Potential /mV
400
Figure 2. Dmerentlai double-layer capacitance of the HMDE in 0.1 M KNOS solution contalnlng 4 X M Triton X-100. Experimental c o n d i t h phase sensitive ac voltammetry, frequency 1084 Ht,a m p i M e 7.5 mV, sweep rate 0.04 V I S . Deposition: 10 s at 0.250 V vs SCE without stirring. IodMe concentration: (A) 0, (B) 1.0 X lod, (C) 2.0 X lod M. Ail curves recorded from posltive to negative potentials.
Other surfactants used in this study, notably the Tween and Brij series, showed similar effects to TX for the cathodic stripping of iodide from mercury electrodes. Effects of Potential on Deposition Efficiency. As already stated, the efficiency of halide deposition on mercury electrodes depends on potential. In Figure 1, both the effect of potential and the effect of surfactant concentration on the electrical charge associated with the stripping peak are shown. Experiments were carried out at the HMDE in a solution containing a small concentration of iodide (2 X lo-' M for curves A-C) and 0.1 M KNOBas the supporting electrolyte. A 60-s deposition was followed by a stripping cycle at the rate of 50 mV/s. The charge of the stripping peak was determined by numerical integration. Charges associated with c w e s A-C were smaller than 15 pC/cm2, which indicates that the electrode surface coverage with the iodide (0,) was below 0.2. Curve D represents the results of a similar experiment with all conditions the same as in the previous series except for the concentration of iodide, which this time was 1.0 X lo4 M; consequently, was close to one monolayer. Curve D shows no effect of surfadant; evidently the surfactant is not present on the surface because the entire surface is occupied by mercurous iodide. However, some results for small coverages were rather unexpected; the surfactant affects the deposition even at potentials much more positive than +25 mV, i.e., well above potentials where desorption of TX from a mercury electrode should take place. The puzzle was solved by running ac voltammetric experiments in a solution containing TX at the HMDE partially covered with iodide. The peak of TX desorption moved to more positive potentials as 8,was increased (see Figure 2). Apparently, the energy of TX adsorption decreases in the presence of mercurous iodide, but to explain the results in Figure 1,curve D, one should assume that this energy is always higher than the adsorption energy of the iodide. Therefore, TX is completely removed from the surface when enough mercurous iodide is generated by the electrode process. Effects bf Surfactanton the Shape of StrippingPeaks. Figure 3 shows examples of stripping peaks at similar surface coverages in the absence (A) and in the presence (B) of TX. The solution contained 1.0 X 10" M I- in 0.1 M KN03, the sweep rate was 0.1 V/s, and the deposition was carried out at a potential of 0.0 V without stirring, for 20 s in the presence of the surfactant and for 150 s in its absence. As is evident from these results, in the presence of the surfactant, the stripping peak is narrower and occurs at a more negative potential. Both the peak potential (E,) and the peak width at half-height (WlIz) were found to depend on 01. The results
1160
ANALYTICAL CHEMISTRY, VOL. 63, NO. 11, JUNE 1, 1991 0.01 T
. 4
-0.01 O I
2T
-.TI""
w
t4
-0.03 -0.04
-800
-6
B
-400
-200
D \ v /
4
-8 -600
I
-600
-/I
tI
0
I
-400
Potential /mV
Flgure 3. Effect of Triton X-100 eddiion on the stripping peak of mercurous lodlde. Experimental conditions: HMDE (0.0087 cm9, deposlbn without stirring at 0.0 V vs SCE, strlpping rate 100 mVls, solution 1.0 X lod M I- in 0.1 M KNO* Concentratbn of surfactant and depositton tlme: (A) 0, 150s; (B) 8 X lo-' M, 20 s. The electrode coverage with mercurous iodide was 0.076 and 0.061 for curves A and B, respectively. 180
-200 Potential /mV
0
200
Flgurs 5. Stripping peaks of mercurous iodide as a function of the deposition time. Experimental conditions: HMDE (0.0087 cm2), stripping rate 1 V/s, solution 1.0 X 10" M I- in 0.1 M KNO,, deposition at 0.0 V vs SCE without stkring: tlme (A) 10, (B) 40,(C)120, (D) 300 5.
T
-2
I
-3 -800
-600 -400
-200
0
-600 -400
-200
0
200
Potential /mV
80 0.01
F I p m 6. Cyclic vdtammogams at HMDE (0.0087 cm2). Sweep rate 10 VIS, supporting electrolyte 0.1 M KNO (A) 2.0 X lo4 M I- and 4 X lo-' M Triton X-100. (B) 1.0 X 10'M I-, no surfactant. The electrode was held at -0.6 V for several seconds prior to the experiment.
.
0.1
1
Initial 0,
Flgurs 4. Dependence of the peak width at haif-height on electrode coverage with mercurous iodide. Experimental conditions: HMDE (0.0087 cm?,varlable depositkn time (15-300 s)at 0.0 V vs SCE with stlrrlng, stripping rate 100 mV/s, solution 1.0 X lod M I- in 0.1 M KNO,. (A) No surfactant, (B) 4 X lo-' M Triton X-100.
in Figure 4 indicate that W1,2for the reduction of mercurous iodide in the absence of TX is about 160 mV and is almost independent of 81, whereas in the presence of surfactant WIl2 changes from about 95 mV at e1< 0.14 to about 160 mV for 81 >0.7. Note that, in agreement with the results in Figure 1, the effect of TX disappears when eIbecomes large. A relation between the peak potential, E,,, and the electrode coverage can be represented by a linear equation: E, = a + b log eIwith a = -163 f 3 mV, b = 121 f 3 mV for solutions without TX and a = -155 f 5 mV, b = 74 f 4 mV in the presence of the surfactant (experimental conditions as in Figure 4). Here again the difference between E, in solutions with and without TX decreases when eIincreases (102 f 11 mV at eI = 0.01 and 8 f 6 mV at BI = 1). It should be strongly emphasized that when the coverage of the mercury electrode with the iodide does not exceed a monolayer, the deposit behaves as an imperfect two-dimensional gas and no nucleation phenomena are observed. is formed following the However, a t 81 > 1, a bulk Hg212(B) nucleation process. The stripping peak of bulk Hg212,,! is sharper (because the activity of a solid phase is an intensive property) and occurs at a more positive potential. Figure 5 shows an early stage of the multilayer deposition; from that experiment, the electrical charge required to form a monolayer was determined as 72 pC/cm2, which is close to 83 pC/cm2
reported by Propost (15) for the formation of a monolayer of Hg2I. In another series of experiments, the effect of the sweep rates on the peak potential was studied using the HMDE as the working electrode in solutions containing 1.0 pM I- and 0.1 M KN03. The sweep rate was varied from 0.1 to 10 V/s and eIfrom 0.1 to 0.9. In all cases, a linear relation between the peak potential and the logarithm of the sweep rate was observed. In the absence of the surfactant, E, moved by about -58 mV with a 10-fold increase in the sweep rate, and this shift was independent of In the presence of TX, the shift was -40 mV/decade at eI= 0.1 and about -55 mV/decade a t OI = 0.4. In Figure 6, cyclic voltammograms recorded at the sweep rate of 10 V/s are shown. Both curves with and without TX indicate a reversible charge-transfer process. The ratio of cathodic to anodic peak current is less than 1because the rate of the cathodic process is controlled by diffusion of I- ions from the bulk solution, whereas the anodic process involves stripping from the surface of the electrode. Note that in the absence of surfactant the charging current is much larger; therefore, a higher concentration of the depolarizer has been used to obtain well-defined faradaic currents. Calibration Curves. Figure 7 illustrates the effect of the addition of surfactant on the resulting calibration curves for the cathodic stripping of mercurous iodide from the HMDE. When the deposition step is carried out at relatively negative potentials (which is always done in analysis of real samples to avoid interferences from the codeposition of other halides (25)),then the calibration curve at low concentrations has a concave shape. This means that the sensitivity of the method (Le., the signal/concentration ratio) decreases with a decrease
ANALYTICAL CHEMISTRY, VOL. 63,NO. 11, JUNE 1, 1991
0
0.5 1 Concentmtion/pM
1.5
Flgure 7. Callbration curves for strlpplng of mercurous lcdlde. Experimental conditions: H M M (0.0087 cm2), BO-s deposition wlth stkring, stripping rate 100 mVls, supporting electrolyte 0.1 M KNO,. Depositkn potential (A) -100, (6) -50, and (C) -50 mV wlth 4 X lo-' M Triton X-100 added.
in the analyte concentration. This, of course, has an adverse effect on the detection limit. However, as is evident from the results in Figure 7, deposition done at a similar potential in the presence of the surfactant produces a linear calibration plot in the low concentration region.
DISCUSSION In order to explain the effect of TX addition on the stripping peak of mercurous iodide, let's first answer the question as to whether the surfactant affects the kinetics of the electrode process and/or the thermodynamic properties of the system. Luther et al. (16) emphasized the kinetic effects. The mechanism by which the surfactant caused the increase in peak size was explained assuming an increase in the depition efficiency brought about by the ability of the surfactant to block the electrode surface, thus effectively preventing diffusion of Hg+ : into bulk solution and thereby increasing the amount of Hg212available for subsequent stripping. The diffusion of Mn+ions (where M is the electrode material) into bulk solution has previously been put forward in order to explain a potential dependence of deposition efficiency of halides at both mercury and silver electrodes. According to this concept, Mn+ions, which are produced in large quantities at a sufficiently positive potential, diffuse from the electrode surface into the solution, causing a precipitation of MX(,) in the bulk solution rather than on the surface of the electrode (26, 27). The explanation suggested by Luther et al. (16) implies selective blocking of electrode by the surfactant, Le., when the transport of Hgt+ is inhibited; the movement of I- to the electrode is expected to occur without change. Such an action of TX cannot be explained by taking into account the properties of the surfactant. Also, it can be shown that when the concentration of iodide is less than about 5 X lo4 M (i.e., a concentration range typically used in cathodic stripping determination of iodide) Hg2120will never be formed, either on the electrode or in the solution, regardless of the Hg:+ concentration. This arises from the thermodynamic instability of Hg212(.)at low iodide concentration due to a disproportionation reaction: Hg212(,) Hg, + Hg12(,. The equilibrium constant for this reaction is K&K, = 2.5 X lo4 (where K, is the solubility product of Hg212(1),O2 is the stability constant of Hg12(w),and K is the equilibrium constant of Hgz2+e Hg(1) + Hg2+). However, on mercury electrodes submonolayer deposits of mercurous iodide can be formed, even at very low concentrations of iodide, because of a contribution from the adsorption energy. As it was shown by Baranski and Galus
*
1161
(28), the potential dependence of the halide deposition efficiency can be quantitatively explained by considering the complex formation equilibria. Unfortunately, the paper (28) was printed with mixed up figures and equations. A very similar explanation can be provided in the case of an Ag electrode where the formation of AgX,,) and at least two soluble complexes (AgX(,) and Ag2X+(w)) should be considered. The results in Figure 6 show that the electrode process is very fast both with and without the surfactant. The same conclusion was derived from previously reported experiments involving the cathodic stripping of iodide from ultramicroelectrodes at sweep rates of 10-700 V/s (24). These results are in agreement with the work of Osteryoung and co-workers (B), who concluded, on the basis of normal pulse polarographic studies, that the formation/dissolution of mercurous iodide can be described as a reversible one-electron-transfer process: Hg I- s 'l2Hg2I2 e-. All experimental results clearly indicate that the surfactant effect is independent of the rate at which the stripping process is carried out, but it depends on the electrode coverage with the iodide and TX. Clearly, the effect must be thermodynamic, not kinetic, in nature. A possibility that the mechanism involves complex formation, of some sort, between the halide in solution and the surfactant, which results in adsorptive accumulation, was ruled out. A brief study indicated that an addition of TX did not affect the equilibrium potentials of either Hg(l)(Hg2+nor Hg(l)lHg212(311electrodes. At this point, it becomes obvious that only thermodynamic properties at the mercury-solution interface are significantly affected by a small addition of TX to the solution. In order to describe the interfacial interaction, adsorption parameters for the adsorption of mercurous iodide with and without TX were determined. For simplicity we will refer to adsorbed mercurous iodide as HgI(,d). Therefore, the electrode reaction leading to the formation of this species will be treated as a one-electron process.
+
+
This is in agreement with Osteryoung et al.'s conclusion (29). In reality, the product formed at a low electrode coverage may be Hg21td),as suggested by Propost (15). From the viewpoint of thermodynamical formalism, there is no essential difference between treating the product as HgI(,d), HgzI(,d), or Hg&d) (if only the number of electrons in the electrode reaction is the same as the number of iodine atoms). It should also be stressed that mercurous iodide adsorbed on mercury electrodes is chemically different from solid Hg212,and perhaps, it could be better described as a chemisorbed iodine. Only the lowenergy electron diffraction (LEED) studies at solid mercury electrodes could reliably establish the structural arrangement of the adsorbed molecules. Efforts were made to determine the adsorption isotherm of mercurous iodide under experimental conditions similar to those taking place during the cathodic stripping experiment. A relation between the electrode coverage, e,, and the activity of HgI(,d) can be written as aHgI
= F(eI)
(2)
The electrode coverage at a potential E can be determined from the electrical charge associated with reduction of HgI(d)
where Q, is the charge associated with the reduction of the monolayer, Qt is the total charge associated with the stripping peak, and i(t) is the faradaic reduction current expressed as a function of time.
1162
ANALYTICAL CHEMISTRY, VOL. 63, NO. 11, JUNE 1, 1991
The activity of the product on the electrode surface can be determined from the Nernst equation, if the electrode process is reversible (which was already proven) aHII
= ~1
"
exp[ =(E - E O )
1
1
' T
0.8
(4)
where f1 is the activity coefficient of I- ions in solution (for 0.1 M KNOB,f~ = 0.78 was assumed), QIis the concentration of I- at the electrode surface, and Eo is the standard potential for reaction 1. The value of Eo depends on the selection of the standard state for HgI(4). It is convenient to set Eo as equal to the standard potential of the Hg212(,)/Hgo) redox couple (-0.283 V vs SCE),so when ab1 reaches 1,formation of a multilayer (bulk) Hg212(s)deposit becomes possible. Therefore, the standard state of HgI(4)is defined here as one at equilibrium with Hg212(o. The concentration of iodide ions on the electrode surface can be calculated by using the convolution method (30-32). In the case of a spherical symmetry of diffusion, the convolution expression can be written as (33,34)
0.6 010 init 0.4
0.2
0
-16
-12
-8
-4
-400 -300 -200 -100
0
0
Potential /mV
(5)
Flgure 8. Electrode coverage with mercurous iodide plotted as a function of the adsorbate acbMty (A) and electrode potential (B). Data were obtained from a series of cyclic voltammetric experiments at sweep rates (a)0.1, (b) 0.2, (c) 0.3, (d) 0.5, (e) 1, (f) 2, and (9) 3 VIS. Solution: 6.0 X lod M I- and 4 X lo-' M Triton X-100 h 0.1 M KNO,, 40-s deposition at 0.0 V vs SCE without stirring. The initial electrode coverage was 0.78 f 0.05.
where CoIis the bulk concentration of iodide, F is the Faraday constant, V is the volume of a spherical electrode, and IOut is defined by the following integral
where g describes intermolecular interactions between the surface adsorbed species (g > 0 attractive or g < 0 repulsive), and /3 depends on the adsorption energy:
Iout CBI = C O I - 3FV
P = eXP(--)
where D is the diffusion coefficient of I-, r is the electrode radius, and i(t)is the faradaic current (negative for reduction). Methods of numerical calculations of eq 6 are discussed elsewhere (34). The most critical part in the computations of expressions 3 and 6 is the background current correction. This was done by fitting the background current before and after the stripping peak (at least 150 data points on each side) into a third-order polynomial expression (a, + alx + a2x2+ a3x3)by the least-squares method. Then the background current under the peak was numerically evaluated and subtracted. The described method was used to obtain isotherms from cathodic stripping experiments run at different sweep rates in solutions with and without TX. The reason for running the experiment at several sweep rates was to obtain the potential dependence of the adsorption parameters. When the sweep rate is increased, a higher concentration of I- ions is generated and the reduction process is shifted into more negative potentials. As an example, Figure 8 shows a series of adsorption isotherms obtained for the initial surface coverage, BIht = 0.75, in the presence of TX. This family of curves was obtained by varying the sweep rate from 0.1 to 3 V/s. To avoid errors associated with the ohmic drop, electronic iR compensation was used during the measurements. Each data point consists , O1. In Figure 8A, eI of three quantities: potential, U H ~ and is plotted vs am,and in Figure 8B 81 as a function of potential is shown. The curves in Figure 8A do not have the familiar sigmoidal shape because together with aHg1 the electrode potential is changing. In order to determine the adsorption parameters, the experimental results were compared with the Frumkin adsorption isotherm (35) @HgI
=
81
exp(-2geI)
(7)
AGoad
The standard free energy of adsorption, AGoad, is expected to depend on potential AGo,d = AGO0 alF(E - E O ) + u ~ [ F ( E E0)I2 (9)
+
where u1 and u2 are constants independent of potential and AGO, is the standard free energy of adsorption at the standard potential of the Hg212(8)/Hgo) couple. From the physical point of view, it perhaps would be better to measure potentials with respect to the zero charge potential, however, such a reference point would create a serious complication because the zero charge potential certainly depends on 81 and the electrode coverage with TX. The interaction parameter g may also depend on potential (36);here however, for sake of simplicity, such a dependence was neglected. By combining eqs 7-9, we obtain aHgI
F(E - E o ) RT
=
+
where 0, is a constant independent of eI and potential. When eIis fixed, eq 10 can be written as In aHgI = constant + u ~ F ( E- E o ) / R T uZ[F(E- E 0 ) / R n 2 (11)
+
Thus, by fitting In ab1 vs E into a second-order polynomial expression, u1 and u2 can be determined. However, in all casea that were analyzed, only a linear relation between In aw and E was observed; therefore, only ul was determined. Figure 9 shows such analysis done for the family of isotherms given in Figure 8. For systems without TX, u1 is practically independent of 6.However, in solutions containing the surfactant, the dependence of u1on the initial value of 81 is quite large and it can be represented by a linear relation: u1 = 0.22 + 0.198I,hit. The u1 parameter depends only on the initial electrode coverage with mercurous iodide, because the stripping process is relatively fast (0.07-2 s) and there is not enough
ANALYTICAL CHEMISTRY, VOL. 63. NO. 11, JUNE 1, 1991
-1
Table I. Thermodynamic Parameters of Mercurous Iodide Adsorption on Mercury Electrodes solution
100
adsorpn param
PO" dependence of PO on %init Q1
-100
J -12
i
-10
-8
-6
-4
-2
0
Figure 0 . Andy& of mercurous kd#e adsorption isathenn at constant Qaccordlng to eq 11. Plots for systems with Triton X-100 (4 X lo-' M in solution): = 0.04 (LTX) and = 0.78 (HTX). Plots for = 0.12 (L) and = 0.7 (H). systems without the surfactant:
time to populate sites released by HgI molecules with molecules of the slowly diffusing surfactant. Values of both g and Bo parameters were obtained by comparing a family of adsorption isotherms at a fixed potential. When E is constant, eq 10 can be written as In
(uHgI)
1163
- In (1:eI) U&E
=
- EO)/RT - In (bo)- 2ge1 (12)
The quadratic term in eq 11was omitted because experiments show that for this system u2 = 0. By plotting the left side of eq 12 against el,g and Bo were determined from the slope and the intercept, respectively. When the analysis is repeated at various potentials, the potential dependence of g can be established. However, the standard deviation for the slope was too large and the potential range was too narrow to conclude whether or not g in this system is potential dependent. Parameter g was found to be practically independent of the initial 01 in both solutions with and without TX. However, in solutions containing the surfactant, the Bo parameter changes following approximately the relation In Bo = 6.89 - l.681,init. All adsorption parameters for solutions with and without TX are listed in Table I. The value of the Bo parameter is given at the standard potential of the Hg212w/Hgo,couple. The random errors for u1 and Bo are less than lo%, but the precision of g is rather poor. A comparison of the adsorption parameters obtained for solutions with and without TX is very interesting. Despite the competition between TX and mercurous iodide for the electrode surface, the adsorption of HgI is up to 10 times stronger in the presence of TX. This surprising synergic effect is in agreement with a negative shift in the stripping peak potential caused by an addition of surfactant (Figure 3) and with evidently stronger adsorption of T X in the presence of mercurous iodide (Figure 2). The effect disappears, as expected, when the electrode coverage with TX is low, i.e., at high values of Another major difference appears in the parameter ul, which describes the potential dependence of the adsorption energy. This parameter is about 2 times smaller in the presence of the surfactant. The adsorption of organic molecules increases the thickness of the Helmholtz layer, which in turn significantly decreases the interfacial potential gradient. Since only a fraction of the interfacial potential difference acts on small HgI molecules, the potential dependence of the adsorption energy for these molecules is reduced. However, such a situation exists only when the electrode coverage with TX is large and eIis small. With an increase in el,organic molecules are removed from
0.1 M KNOB +
0.1 M KNOS,
0.4 fiM TX
no TX
80f3
920 f 35 In & = 6.83 - l.681,ht
-0.5 f 0.2 0.44 f 0.04 (at = 0.12)
-3 f 0.8 0.23 f 0.01 (at Bitinit = 0.04) u1 = 0.22 o.igeiat
+
dependence of Q1 on h l i t a At the standard potential of the Hg2I2(.)/Hg(l)couple.
the interface and the potential distribution becomes similar to one existing in the absence of the surfactant. One would expect that the potential dependence of B should become the same as in solutions without TX, and indeed this is observed experimentally. There are significant differences in the values of the g parameter between systems with and without TX. The large negative value of this parameter in the presence of the surfactant is consistent with a broadening of the stripping peak observed for large coverages of the electrode with mercurous iodide (Figure 4). The value of this parameter may to some extent reflect the energy needed for compressing or rearranging TX molecules on the surface to allow the increase of 0I. The observed variation of u1 and Po with 0 1 ; ~and ~ , a very negative value of the interaction parameter,g, seem to indicate that the Frumkin isotherm is too simple for this complicated system, but if five empirical parameters are allowed (Bo, ul, g, dBo/dOI,int,and dul/dOI,inJ, this isotherm can be forced to represent almost any data. Of course, another solution is to restrict considerations to systems with OI;mt < 0.1; then the variation of Bo and u1 with the coverage can be neglected. It should be stressed that in the presence of TX the stripping peaks of mercurous iodide can be measured accurately at coverages as low as 5 X The adsorption parameters listed in Table I were used in numerical simulations to predict the potential dependence of the deposition efficiency. The theoretical model was similar to one used by Baranski and Galus (28) except that the activity of HgI(,d). was calculated from the Frumkin adsorption isotherm using parameters determined in this work. The simulation was based on the following approach. It was assumed that a change in electrode coverage, A€+, during a period At can be calculated from a difference between the flux of I- ions from the bulk solution to the electrode and the flux of Hg(11)-iodide complexes from the electrode to the solution. For simplicity, it was assumed that diffusion coefficients of all species are equal and the thickness of the diffusion layer is fixed (Le., the solution is stirred during the deposition). The following equilibria were considered [Hg22+]= exp[(E - Et0)2F/RT] [Hg2+] =
K[
Hg22+]
(13) (14)
[Hg22+lIX-1 = KsoQHg12
(15)
[HgXn2-"] = @,[Hg2+][X-]"
(16)
aHgI
=
mI)
(17)
where E t 0 = 0.544 V is the standard potential of the Hgo,/ Hgz2+redox couple, K = 0.0115 is the equilibrium constant, K, = 3.5 X is the solubility product constant, p,, is the complex stability constant (only two constants were used: Bl
1164
ANALYTICAL CHEMISTRY, VOL.
63,NO. 11, JUNE 1, 1991
formation, and smaller quantities of soluble Hg(II)I,2-" complexes are produced. However, the improvement in the detection limit arises also from other factors. Adsorption of large organic molecules reduces the potential gradient across the interface, and consequently, the energy of mercurous iodide adsorption is less dependent on potential. As a result, the stripping peaks are narrower and, therefore, higher and easier to measure. Finally, the presence of surfactant reduces the charging current, causing an additional increase in the signal to background noise ratio. It is very advantageous from the viewpoint of analytical applications that all beneficial effects of TX addition increase when the electrode coverage with mercurous iodide decreases.
0 -200
0
200
400
Deposition Potential /mV
Fburo 10. Numerical simulation of mercurous iodide deposition as a functlon of depositlon potential for systems wlth (6)and without (A) Triton X-100. Adsorptbn parameters from Table I were used: iodide concentrationwas set to 2 X lo-' M, the depositkn time to 60 s, and the thickness of the diffusion layer to 20 pm.
= 12.87 and o2 = 23.79), aHgI is the activity at the adsorbed mercurous iodide, and F(eI)is the adsorption isotherm. All equilibrium constants are from ref 37. The electrode coverage, €31,i,for a given time, iAt (where i > 0 is the iteration number), was calculated numerically by using the algorithm
and
= 01,i-l + Ae1.i-1
(19) In eq 18, SI,^-^ and SQ-~are fluxes of iodide ions and iodide complexee, respectively, for i - 1iterations. 8, is the electrical charge needed to form a monolayer. The period At was set small enough so POI was always smaller than 5 X lo4. Fluxes can be obtained from eqs 13-17 0I.i
=
D
SIC*- K,0[F(01,i-1)]2exp[(E'O
- E)2F/RT]J (20)
and %.i-1
=
where C* is the bulk concentration of I-, D is the diffusion coefficient, and b is the thickness of the diffusion layer. The results of such simulations for solutions with and without TX are shown in Figure 10. The calculations (2000 iterations) were carried out for adsorption parameters listed in Table I, C* = 0.2 p M , and DAt/d = 3 X lo-' cm. Clearly, the addition of the surfactant improves the deposition efficiency, which is in qualitative agreement with the experimental results shown in Figure 1. Due to a synergic effect, the adsorption HgI is stronger in the presence of the surfadant. This shifts the equilibria at the interface in the direction of HgI(,,
LITERATURE CITED Holek, W. Anal. Chem. 1967, 59, 2218-2221. Van Staden, J. F. Anal Lett. 1966, 19, 1407-1419. Cherklewski, F. L.; Rldllngton, J. W. J . Assoc. Anal. CY". 1967, 70, 924-928. Laitlnen, H. A.; Boyer, K. W. Anal. Chem. 1972, 44, 920-927. Z O k , W. He;Qordon, (3. E. Anal. Chem. 1970, 42, 257-282. a m ,N. A.; Marshan, G. B. AMI. m.A- 1907,202.49-59. Nakahara, T.; Wase, T. Anal. Sei. 1986, 4 , 223-225. Miyazakl, A.; Bansho, K. S p e W " . Acta, PaH 8 1967, 428, 227-233. Miles, P. J . A m . Anal. Chem. 1976, 61. 1388-1371. Hvst, W. J.; Snyder, K. P.; Martin. R. A. J . Uq. &o".1963, 6, 2087-2075. Carbon, A.; Lundstrom. U.; Olln, A. Talenta 1967, 34, 815-819. Anfait, T.; Twengstrom. S. Anal. chkn.Acta 1986, 179, 453-457. RKHnpson, D.; Lee, S.; Allen, R. J . Assoc. Anal. Chem. 1963, 6 , 2087-2075. Mahrano, R.; Bwrlgoll, 0.; Scariatini, M.; Cenderelli, G.; Gondolfl, C.; Grasso, P. Anal. Chem. 1972. 44, 201-209. PrOPOSt, R. C. Anal. Chem. 1977, 49, 1199-1205. Luther, 0. W.; Swartz, C. B.; Ullman, W. J. Anal. Chem. 1966, 60, 1721-1724. schdz, F.; NltsdJte, L.; Henrlon, G. J . E k m n a l . Chem. 1967. 224, 303-306. shein, I.; P e r m , S. P. Anal. chwn.1961, 33, 325-329. Laitinen, H. A.; Lin, 2. F. Anal. Chem. 1963, 35, 1405-1410. Mandandher, K.; Pietcher, D. Ta&nta 1977, 24, 387-390. BUewlc~,R.; KUblik, 2. Anal. chkn.Acta 1985. 171, 205-213. Wong, 0. T. F.; Takayanagl, K.; Todd, J. F. h k r . Chem. 1965, 17, 177- 183. Hen, K.; Koch, W. F.; Pratt, K. Anal. Chem. 1967, 59. 731-738. Herman, A.; Baranskl, A. S. Anal. chkn. Acta 1990, 239, 35-44. cdovoe, 0.; Wllson, 0. S.; Moyers, J. L. Anal. Chem. 1974, 46, 1045-1050. Kemule, W.: Kublk. 2.; Taraszewska. J. h#". J . S m . 1962. 1, 856-882. h v , W. W.; Zhikaharev, T. N.; Khaedakova, N. I. Zh.Anal. Chem. 1962, 37, 1944-1947. Beranski, A. S.; Gelus, 2. TBlenta 1972, 19, 781-788 (errata. 1973, 20. I). T G , J. A.; Able, R. H.; Osterywng, R. A. Anal. Chem. 1975, 47, 1343-1347. S a v ~ n tM. , J.; TeSSk, D. J . E k t r ~ a n a l .Chem. 1975, 65, 57-84. &ennew. M.; Oldham, K. B. Anal. Chem. 1972, 44, 1121-1129. Imbeaux, J. C.; Saveant. J. M. J . Electroanal. Chem. 1973, 44, 189-187. Myland, J. C.; Oldham, K. 0.; Zoski, C. 0. J . Electroanel. Chem. 1985, 193, 3-22. Herman, A.; Baranskl, A. S. Can. J . U".1966, 66, 1038-1048. Frumkln, A. N. 2.pnys. 1926, 35, 792-802. Frumkln, A. N.; Damaskin, B. B. Mdsm Aspects of E A s & " b b y ; Bodtris, J., Conway, B., E&.; Butterworths: London, 1984; Vol. 3, pp 149-223. StabMy Constants;Special Publicetkm No. 17; The Chemlcel Society: London, 1984; Bwllngton House, W.l, Part I, Complkd by Sllkn, L.G.
- _
RECEIVED for review September 19,1990. Revised manuscript received February 11,1991. Accepted February 15,1991. The financial support of the Natural Sciences and Engineering Research Council, Canada (NSERC), through the operating grant, is gratefully acknowledged.
