1964
Anal. Chem. 1986, 58, 1964-1969
(13) Christie, J. H.; Jackson, L. L.; Osteryoung, R. A. Anal. Chem. 1976,
(17) Miller, J. C.; Miller, J. N. Statistics for Analytical Chemlsfry; Wiiey: New York, 1984. (18) Galvez, J. Anal. Chsm. 1985, 57, 585. (19) Oldham. K. B.; Bond, A. M. J . Electroanal. Chem. 1983, 758, 193.
48, 242.
(14) Garden, J. S.;Mitchell, D. G.; Milis. W. N. Anal. Chem. 1980. 52, 2310. (15) Larsen, I. L.; Hartmann, N. A,; Wagner, J. J. Anal. Chem. 1973, 45,
1511. (16) Franke, J. P.; de Zeeuw, R. A,; Hakkert, R. Anal. Chem. 1978, 50, 1374.
RECEIVED
for review August 27, N85. Accepted March 20,
1986.
Polarographic Reduction and Determination of 1,2=Dibromoethanein Aqueous Solutions Robert Tokoro,l Renata Bilewicz? and Janet Osteryoung* Department of Chemistry, State University of New York at Buffalo, Buffalo,New York 14214
The electrochdcal reduction of 1,P-dlbromoethane (EDB) In aqueous solutlon has been studied with dc, normal pulse, and reverse pulse polarography, vdtammetry, and coulometry. The reductlon mechanlsm Involves slow two-electron transfer wrth fast d l o p l a c W of bromide leading to formation of ethene. Resutts of detalled RPP lnvestlgatlons In neutral and slightly acid solutions suggest that the flnal produds are formed through a carbanbn Intermediate. EDB can be detemdned dkectty from the llmlthg redudion current or lndkectly from the llmltkrg curren! lor the anodk oxklatbn of mercury In the presence ofthe redudkn product, bromkle. The detection Ihtt Is ca. 1 pM. The method Is used to determine the OdUMltty of EDB In water, whlch was found to be 3.1 mg/g of H,O.
In recent years there has been concern for the environmental and toxicological effects of 1,2-dibromoethane (EDB). As it was found to cause cancer, birth defects, and genetic disorders in test animals, recommended acceptable levels in grain and food have been established (I). Widespread use of EDB as a fumigant and a scavenger added to leaded gasoline focused the attention on its determination in foodstuff (2-4) and polluted atmosphere (5-7). A number of analytical procedures have been presented with the determinative step via gas chromatography coupled with electron capture detection, electrolytic conductivity, or plasma emission spectrometry (8, 9). Of these methods the electron capture detector yields the best detection limits. Difficulties arising in this technique are associated with the long times used in the chrnmatographic steps and interferences from solvent impurities and coextractables. Though retention times may be reduced, it is difficult to perform a sufficiently rapid analysis with good resolution from the solvent front ( 2 ) . The purity of solvent seems to be critical as the impurity peaks are close to EDB retention times (10)and coextractables may interfere with the confirmation by mass spectrometry (8). The lack of any electrochemical procedure for the determination of EDB prompted the present work. Electrochemical Permanent address: Instituto de Quimica USP B-843,Caixa postal 20780, Sa0 Paul0 SP OlOOO, Brazil. Permanent address: Department of Chemistry, University of Warsaw, 02093 Warsaw, Pasteura 1, Poland. 0003-2700/86/0358-1964$01.50/0
studies of EDB were first carried out by von Stackelberg (11). Meites (12) and Feoktistov (13) have reported a few studies of EDB since that time. Our work was performed to study the electrochemical behavior of 1,2-dibromoethanein aqueous solutions and then to establish the best electrochemical conditions for analytical purposes.
EXPERIMENTAL SECTION Reagents. 1,2-Dibromoethane (Fisher Scientific Co. E 173 Certified) was of 99.5% purity (density 2.179Z4%g/cm3). Ethanol was supplied by US. Industrial Chemicals Co. (U.S.P. specifications). Distilled water was purified further by a Millipore MilliQ system (resistivity > 100 Mi2 cm). Sodium bromide (Mallinckrodt Analytical Reagent), tetraethylammonium perchlorate (TEAP, Eastman Kodak), tetrabutylammonium perchlorate (TBAP, Eastman Kodak), sodium chloride (Baker Analyzed Reagent), sodium perchlorate, purified (Fisher Scientific), and lithium perchlorate, anhydrous (Baker reagent analyzed), were used as supporting electrolytes. Catalyst R-3-11 (Chemical Dynamics Corp.) was used to remove oxygen from argon (99.998%, Liquid Carbonic). Perchloric acid (Fisher Scientific) and mercury (Bethlehem Apparatus) were used as supplied. Instruments and Accessories. The following were used: EG&G PARC Model 174 polarographic analyzer and conversion of module (14,151 with EG&G PARC Model 174/70 drop timer assembly 174/70; Tacussel Electronique Polaropulse PRGB with homemade converter to PARC 174/70; EG&G PARC Model 173 potentiostat/galvanostat with Model 179 digital coulometer; Metrohm Calomel reference electrode, EA404, NaCl 3 M; Eppendorf micropipeta, 10-100 rL. In some normal pulse (NP) and reverse pulse (RP) polarographic experiments a computer-controlled pulse-voltammetric instrument based upon a Digital Equipment Corp. PDP-8/e minicomputer and homemade interface were used. The system has been described elsewhere (16). Gaseous products were identified by use of gas chromatography with a Perkin-Elmer SIGMA 3-B chromatograph and a 180 cm X 3 mm column packed with 20% Apiezon L on 60-80 mesh Chrom-W (column inlet temperature 150 "C, column temperature 40 "C, detector temperature 200 "C). Procedures. All potentials are reported vs. the 3 M NaCl calomel electrode. Measurements were carried out on two dropping mercury electrodes with the following characteristics: (1)natural drop time ( t d ) = 7.20 s, flow rate ( m )= 0.9 mg/s, (2) t d = 6.17 s, m = 1.14 mg/s. Voltammetric experiments were carried out on a HMDE, Kemula and Kublik type (17) with surface area 0.024 cm2. A Luggin capillary with Vycor tip containing 0.1 M TEAP was used to avoid contamination of the working solution with NaCl from the reference electrode. The solutions were thermostated at 25 0.2 "C. Oxygen was removed from solutions by sparging with purified argon (18). Pure EDB
*
0 1986 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 58, NO. 9, AUGUST 1986
was sparged thoroughly with argon, and the same operation was performed with ethanol. The working solution of EDB was a 1% (v/v) solution of EDB in ethanol prepared from deaerated substances. The following technique was used to prepare oxygen-free solutions of EDB. The aqueous supporting electrolytesolution was thoroughly deaerated by sparging with argon for 20 min. After this an aliquot of 10.00 mL of supporting electrolytesolution was transferred to the polarographic cell, which had been flushed with argon, and sparged again with argon, gently for 2-3 min, to avoid splattering and evaporation. A blanket of argon over the surface of the solution was maintained while obtaining the base line for the supporting electrolyte. A small volume (on the order of 10 pL) of EDB diluted with ethanol was added by micropipet. The flow of argon was shut off and the contents of the polarographic cell were stirred for 3 min to saturate the chamber with vapors of EDB. The amount of EDB required to saturate the chamber corresponds roughly to 22% of a 10-pL aliquot (see below). For subsequent additions of EDB the cell was opened as necessary. The amount of oxygen that entered the cell under these circumstances was negligible. However, great care must be taken in handling samples of EDB to avoid serious loss by volatilization. Coulometry at controlled potential was carried out at a mercury pool cathode with area -4 cm2. The anode, platinum foil, was separated from the cathode compartment by using an H cell with a sintered glass frit. The gaseous products of electrolysis were collected in a cold trap at liquid nitrogen temperature and analyzed by gas-phase chromatography. Identity of peaks was established by comparison with authentic samples.
RESULTS AND DISCUSSION Preliminary Studies. Due to different solubility of EDB in organic solvents and aqueous media, some preliminary studies were done to select a suitable solvent and supporting electrolyte. Polarograms in acetonitrile, dimethyl- or diethylformamide, or water with perchlorate supporting electrolytes (TBAP, TEAP, LiC104, NaC104) exhibited qualitatively reasonable shapes with half-wave potentials in the range -1.4 to -1.7 V. Dioxane cannot be used because the background current begins to be appreciable at ca. -1.2 V. It might be possible to find a mixture of dioxane and water that would be a good solvent (11), but examination of mixed solvent systems was not included in this work. In dimethylformamide (0.25 M TBAP) background currents increased at -2.0 V. The addition of 0.4 mM EDB resulted in a wave with Ellz = -1.53 V. EDB is only slightly soluble in water, but a well-defined wave was obtained at -1.6 V. The most convenient supporting electrolyte seems to be TEAP due to its large solubility in water. The tetraethylammonium ion is reduced a t a more negative potential than are alkali cations like lithium and sodium ions. Thus TEAP gives a greater negative potential range than NaClO., or LiC10,. Since it was necessary to avoid the presence of halides in this study, the perchlorate salt was chosen. During the preliminary study aqueous solutions saturated with EDB were used. Under these conditions both dc and N P polarograms had a peak before the limiting current. Usually such a peak in the N P mode is due to adsorption of reactant (19). At lower concentrations of EDB (-5 mM) no maximum was observed, even in the dc mode. At these concentrations water seemed to be a suitable medium for determination of EDB. Solubility of EDB in Aqueous Media. EDB is quite soluble in most organic solvents but has limited solubility in water. This led us to determine the solubility of EDB in water. A saturated aqueous solution of EDB was filtered and samples of 50, 100, or 200 pL were diluted in 0.1 M TEAP solution in water or acetonitrile. A cathodic voltammogram on the HMDE or sampled dc polarograms were recorded. Quantitation was done by standard addition with 0.01 M EDB in ethanol serving as a standard. The solubility obtained is 3.12 0.04 (standard deviation, n = 12) mg/g of H,O. A value
*
1965
Table I. Evaluation of DEDefrom Voltammetrya 103n~1'2,
1050,
solution
E,, V
i,, FA
cm/s1/2
cmZ/sb
0.1 M TEAP, HzO 0.1 M LiCIOI, H20 0.1 M TEAP, DMF 0.1 M TBAP, DMF
-1.420 -1.670 -1.530 -1.632
3.0
6.67 6.73 6.99 7.14
1.1 1.1
a
2.7 3.4 3.4
1.2 1.3
CEDB= 0.4 mM,u = 100 mV/s. *Assuming n = 2.
found in the literature is 3.6 mg/g of HzO (20). Dependence of Current on the Concentration of EDB. Currents for reduction of EDB were measured a t a potential of -1.8 V, corresponding to the limiting current region in the dc mode, for successive additions of EDB in the range 0.1-0.6 mM. The resulting calibration curve had slope 3.79 pA/mM, intercept -0.082 FA, and correlation coefficient r = 0.9997. Using the slope and the aliquot size, one can calculate the current increment per aliquot to be 0.436 PA. For the second and succeeding points of the curve the current increment per aliquot was 0.435 f 0.026 pA. However the current due to the first aliquot was only 0.338 PA. Thus we infer that 22% of the first aliquot was lost due to volatilization into the argon atmosphere above the solution in the closed cell. Because the equilibrium EDB(so1n) EDB(g) is then established, subsequent loss is too small to observe. This conclusion is supported by the fact that the negative intercept is 19% of the current increment per aliquot. Thus correction of the data assuming 22% loss of the first aliquot yields a calibration curve with zero intercept within experimental error. The efficacy of the closed polarographic cell was verified by running a polarogram, allowing the solution to stand quietly for 1h, and then recording another polarogram, identical limiting current values were obtained. Nature ofthe Limiting Current of EDB. Both N P POlarograms and dc polarograms of EDB show a limiting current. To find out which rate-determining step limits the current, NP polarograms at different pulse widths were recorded. Measurements were carried out over the time range 16.8-74.0 ms (effective current measurement, t,, 13.1-70.3 ma) in 0.549 mM EDB (0.1 M TEAP) with 1 s drop time. At the shortest pulse width the half-wave potential was -1.45 V whereas a t the longest value it was -1.38 V. Logarithmic analysis gave slopes a log i/a log t, = -0.49 (r = 0.99941, -0.43 (r = 0.99961, and -0.32 (r = 0.998) a t potentials -1.8, -1.6, and -1.5 V, respectively. At -1.8 V, which is well in the limiting current plateau, plots of iL vs. l/t,1/2 were linear with zero intercept but plots at less negative potentials showed distinct curvature. Thus the limiting current is strictly diffusion controlled, but currents at more positive potentials are kinetically controlled. One can evaluate the diffusion coefficient of EDB using the Cottrell equation from the slope of iL vs. l/tm1/2. Experimentally the value of the slope is 1.375 pA s-llz, where A = 7.68 X cm2 and n = 2 (as will be shown below). Thus the value of D is (0.9 f 0.1) X cm2/s. Assuming that the current is also diffusion controlled in the dc mode, one can evaluate D from the Ilkovic equation, id = 708 nD1/2Crn2/3td1/6 where rn = 1.14 mg/s, td = 2 s, and td,nat= 6.3 s. Diffusion coefficients evaluated from staircase polarograms recorded with the computer system in 0.1 M TEAP ( O . O s O . 2 mM EDB) give D = (1.00 0.05) X 10" cmz/s (n = 2,5 measurements). The values of D obtained from cathodic voltammetric peaks in four different solutions are shown in Table I. The experiments carried out by three independent methods gave the diffusion coefficient of EDB in water (1.0 0.1) X cmz/s. Determination of Bromide as a Reduction Product of EDB. Normally the cathodic reduction of alkyl halides produces the halide anion as a product (21). Bromide was
-
*
le66
ANALYTICAL CHEMISTRY, VOL. 58, NO. 9, AUGUST 1986
i-
I
oi
Flgure 1. Reverse pulse polarogram of EDB 0.1 15 mM EDB in 0.1 M TEAP, t , = 1 s, t , = 70.3 ms, scan rate = 10 mVls; (a) background current, €, = -0.2 V; (b) €, = -1.8 V. +0.3
+0.2
to.l
I
I
03
05
,
I
>
07
[EDB], rnM
Flgure 3. Dependence of (1) RP current and (2) dc reductbn current on concentration of EDB: E,= -1.8 V, scan rate = 10 mVls, t , = 1 s, t , = 57 ms, t , = 16.7 ms. See Flgwe 1 for deflnitbn of currents.
0.0
[ W ] , mM Figure 2. NP polarograms for the anodic oxidation of mercury in the presence of bromide: E ,= -0.2 V, scan rate = 10 mVls, f d = 1 s, t , = 57 ms, t , = 16.7 ms; (0)+0.3 V and (0)+0.17 V [Br] (1)0.02, (2) 0.06, (3)0.10, (4) 0.14, (5) 0.18, (6) 0.22, (7) 0.26, (8) 0.30 mM.
identified as a product of the reduction of EDB by carrying out a reverse pulse (RP) polarogram with initial potential in the limiting current region for reduction of EDB as shown in Figure 1. The wave for oxidation of mercury in the presence of bromide appears at ca. +0.15 V. The NP polarogram with initial potential -0.2 V shows no wave for free bromide; anodic current only becomes appreciable for the discharge of mercury to Hg2+at about +0.4 V. In order to use the bromide wave quantitatively, it was necessary to establish some aspects of the anodic reaction under our experimental conditions. From the data of Revenda (22) and Kolthoff (23) it appears that the reaction for the anodic oxidation of mercury is Hg + Brl/zHgzBrz+ e-. The N P anodic wave for bromide is shown more clearly in Figure 2. The maximum occurring on the second wave at about +0.22 V is perceptible only above a concentration of 0.1 mM Br-. It is likely that f i i formation is responsible (24). Concentrations up to 0.02% of Triton X-100 did not suppress the maximum but shifted it to more positive values of potential. The dependence of the anodic current on bromide concentration is also shown in Figure 2. The current depends linearly on bromide concentration in both the f i t and second limiting current regions. Both the slope and intercept are somewhat greater for the current measured at +0.3 V, but the quality of the relationship is similar in the two cases. (These calibration curves have slope 6.83, 6.85 kA/mM, intercept -0.07, +0.14 FA, correlation coefficient 0.9998, 0.9988, and standard deviation about the line 0.01,0.03 FA at +0.170 and +0.300 V, respectively.)
-
The time dependence of the current was examined over a range of potentials. The current a t +0.17 V appears to be strictly diffusion-controlled,whereas the currents on the rising portion of the wave (at +0.090 V) and in the limiting current region of the post wave increase a t a rate faster than l/t,1/2 a t shorter times. Logarithmic analysis of the data obtained for a solution containing 0.1 mM NaBr gave slopes 8 log i/a log t , of 0.57, 0.50, and 0.58, respectively, for E = -0.090, -0.170, and -0.300 V. Similar experiments carried out in the presence of 0.1 mM EDB with initial potential -1.8 V yielded similar results. The current appears strictly diffusion controlled at +0.170 V. Deviations appear at +0.300 and +0.090 V. Using the data of Figure 2 at +0.170 V, one can evaluate from the Cottrell equation the value of the diffusion coefficient for bromide ion in this medium. The value obtained is 1.50 X cm2/s, which may be compared with those reported for bromide of 1.86 X and 2.28 X cm2/s in NaN03 and H2S04,respectively (12). Quantitative Determination of Bromide Released during the Reduction of EDB. By use of the RP technique the dependence of anodic current at +0.170 V on concentration of EDB was found to be linear up to 0.6 mM EDB as shown in Figure 3. At higher concentrations of EDB the wave of bromide became ill-defined. As the same effect appeared in NaBr solutions and no deviations from linearity for dc limiting currents for EDB were observed (Figure 3, curve 2)) this may be due to high coverages of the electrode surface by solid HgzBr2 attained at high concentrations of EDB or bromides in the solution. The data of Figure 3 can be treated point-by-point using the slope of the calibration curve a t +0.17 V Figure 2 to calculate the concentration of bromide. This approach gives the concentration ratio [EDB]/[Br-] = 0.96. A second experimental approach is to measure the concentration of bromide released in the RP experiment by standard addition of bromide. This would probably be the favored procedure if this technique were being used to determine EDB. Typical standard addition curves are shown in Figure 4, and the resulting values are [EDB]/[Br-] = 1.20 and 1.15 for [EDB] = 0.10 and 0.22 mM, respectively. Either approach relates a standard solution of bromide to a standard solution of EDB and allows determination of unknown solutions of EDB by measuring the R P current a t +0.17 V and using bromide standards for calibration. This has three advantages: first, bromide standards are more easily prepared, maintained, and used than are EDB standards; second, the wave for EDB is more drawn out and more easily shifted by changes in chemical conditions than that for bromide; and third, interferences are less likely at ca. +0.2 V than at the very negative potential required to reduce EDB. Coulometry at Controlled Potential. Since the wave of EDB is diffusion controlled in the limiting region, we per-
ANALYTICAL CHEMISTRY, VOL. 58, NO. 9, AUGUST 1986
1967
Table 111. Evaluation of n o for EDB from RP Polarographic Data"
.
.
[EDB],mM
LRP/Ldc
Pld
0.1 0.2 0.3 0.4
1.94 1.83 1.88 1.89 1.89 1.88
0.97 0.92 0.94 0.95 0.95 0.94
0.5 0.6
"0.1 M TEAP in H20, Ei = -1.8 V. 0.017 5 t , 50.057 a.
*From eq
2
using td = 1a,
value of n. The number of electrons involved during the reduction of EDB was determined accurately by using reverse pulse polarography (25). Consider our system
[ ~ r - 1 ,~o-*M
S
Figure 4. RP currents at +0.3 V for ED6 with standard additions of bromide, (0)0.1 mM ED9 and (0)0.22 mM EDB. Other conditions are given in Figure 1.
1 2 3 4
5 6 7 8
electrolysis potential, V -1.5 -1.5 -1.5 -1.7 -1.7 -1.7 -1.7 -1.7
-iRp/iDC QbhLtb
0.20 0.28 0.30 0.20
0.40 0.48 0.46 0.40
C
QEDB,
C
16.30 15.00 16.90 18.24 17.00 16.90 16.90 18.00
+ Hg
-
-
+ pBr'/2HgzBrz + nReP
(1)
where no is the value we wish to determine and nR is known to equal unity. The value of no is related to the limiting currents and the time parameters of the experiment by
Table 11. Evaluation of n by Coulometry at Constant Potential" sample
Br-
+ nee-
=. (pnR/no)((3td/7tP)'/'
- [3td/(3td
n 2.08 1.91 2.15 2.34 2.15 2.13 2.13 2.28 2.15 f 0.12
"10 mM EDB in 0.1 M TEAP, V = 8 mL. *Blank charge was obtained in the same solution but without EDB after the same time of electrolysis.
formed controlled potential electrolysis at -1.9 V. Exhaustive electrolysis was done in a sample of 8 mL of 0.1 M LiC10, in water without and after addition of 0.8 pmol of EDB. After the charge consumed in electrolysis of the supporting electrolyte alone was subtracted, the number of coulombs passed was 16.70, which corresponds to n = 2.1. The same procedure was repeated in 0.1 M TEAP and gave n equal to 2.1 as shown in Table 11. Coulometry at controlled potential was also done while monitoring the unreduced EDB by sampled dc polarography. It was performed in 0.1 M TEAP in water and for comparison in acetonitrile. The plot of id vs. 8 is linear and extrapolation from the last point at Q = 1.5 C to i d = 0 gives the charge 2.2. The theoretical value for n = 2 is 1.9 C. The experiments performed in acetonitrile gave similar results; the charge corresponds to n = 2.2. Long time electrolysis in water solutions generates hydroxide as shown by a pink color on addition of phenolphthalein. This effect could be expected in view of possible reaction mechanisms discussed below. The main product of electrolysis is gaseous. It appeared to be unsaturated hydrocarbon as detected by passing the gas produced in the cell through a solution of bromine or permanganate. The products of electrolysis were collected in a cold trap and analyzed by gas chromatography. The major product was identified as ethene. Ethyl bromide was also detected, but ethanol was not found. Determination of n for the Reduction of EDB by Reverse Pulse Polarography. As is frequently the case, the coulometric study at controlled potential only suggested the
+ 4t,)]'/') (2)
where t d is the drop time and tp is the pulse width. As the current is averaged over a period t,, the effective values of t d and t, are the nominal values less t,/2. It is important to realize that the ratio iRp/iX is independent of (1) the electron transfer kinetics of the reactions of species 0 and R, (2) the bulk concentration of 0, and (3) the properties of the electrode-solution interface. The experimental data for p/no are shown for various concentrations of EDB in Table 111. The mean value of p/n, is 0.94 f 0.01. The calculation of no requires that we know p , the stoichiometric coefficient of bromide in eq 1. It may be evaluated from eq 2 and -iNPB,
= -iDCB,(3td/7tp)'/'
(3)
The ratio of slopes of calibration plots is given by
( ~ ~ R P ~ J ~ C E/ D ( ~B~) N P ~ , / ~=C B J dDEDB/DB,)"'~1
- [ 7 t p / ( 3 t d + 4tp)1'/21 (4)
In Figure 3 for EDB the slope is 7.12 pA/mM and in Figure 2 for NaBr the slope is 6.83 pA/mM. The diffusion coefficients are 1.00 X and 1.50 X cm2/s for EDB and bromide, respectively. The p value obtained is 1.90 which gives no = 2.02. This confirms the less accurate number obtained in the coulometric study and shows that over a wide range of time scale the overall reaction has the same stoichiometry. A t this point it should be mentioned that hydroxide also gives an anodic wave in the potential region of that for bromide (26). The RP polarogram in the supporting electrolyte alone with initial potential at -1.8 V does not show any anodic wave before the dissolution of mercury (cf. curve a, Figure 1). This point was explored further using LiC104 as the supporting electrolyte. The cathodic wave for EDB occurs at more negative potential in LiClO, than in TEAP. But R P polarography with initial potential in the limiting region for reduction of EDB gave an anodic wave due to bromide oxidation identical in all respects with that in TEAP, and no wave was observed in the supporting electrolyte alone. Bulk electrolysis of a solution of EDB at a large pool electrode gave hydroxyl ions, as the solution turned pink after phenolphthalein was added. To rule out the possibility that the
1968
ANALYTICAL CHEMISTRY, VOL. 58, NO. 9. AUGUST 1986
in comparison with the rate of displacement of the second halogen. Nevertheless, evidence of ethyl bromide formation according to the reaction
Table IV. Evaluation of n o for EDB from RP Polarographic Data a -iwlb W A 8.63 6.13 4.88 4.13 3.38
ide,b
CLA
2.05 2.05 2.05 2.05 2.05
t,, s
Plno
PC
0.0168 0.0240 0.0325 0.0405 0.0490
0.89 0.83 0.83 0.83 0.78
1.78 1.66 1.66 1.66 1.56
BrCH2CH2- + H20
"0.564 mM EDB in 1.28 mM HCIOl + 0.1 M LiCIOI, Ei = -1.8 *See Figure 2. Calculated from eq 2 with nR = 1 and no = 2.
measurement of the height of the bromide wave is affected seriously by hydroxide, experiments were repeated in acid solution. In these experiments the hydrogen ion concentration must not be too large or discharge of H+ will interfere with reduction of EDB. Results are presented in Table IV. The values of p obtained are lower than those obtained in neutral solution, and there seems to be a trend to larger values at longer times. This is reasonable in view of the possible reaction mechanisms discussed below. Kinetics of Cathodic Reduction of EDB. As described above, the reduction of EDB is not diffusion controlled on the rising portion of the wave. Furthermore, the reduction waves of EDB in the N P and R P modes obviously cannot be superimposed. The half-wave potential for reduction of EDB in the N P mode depends on pulse width. All these characteristics indicate that the reduction of EDB is irreversible. At sufficiently positive values of E - E I j z the , rate-determining step is electron transfer (27). Normal pulse experimenta were carried out on the rising portion of the wave and analyzed according to the procedure of Oldham and Parry (28) using the equations -
Eli2
(0.0592/2ana) log ( ~ ~ ( 1 .+7 5x2)/(1 - x ) ! (5)
=E
+ (0.0592/ana) log (2.31k(t,/D)1'2)
(6)
where k is the heterogeneous rate constant at potential E , and x is the normalized current, @)/id. A similar treatment can be carried out for the dc mode, for which
E = El,? - (0.0592/ana) log { 2 ~ (-3~ ) / 5 ( 1- x)] (7) Eli2
=E
+ (0.0592/ana) log (1.35/~(td/D)~'~) (8)
where t d is the drop time. Plots of eq 5 and eq 7 are linear over about one unit change in the logarithm and plots of eq 6 and eq 8 are linear over about 2.5 units change in log k. The value of an, was found to be 0.24 and the value of k was found to be 3 X cm/s at -1.00 V. Mechanism of Reduction of EDB. In aqueous solutions it appears that ethene is formed as the result of two-electron irreversible reduction of EDB. The process may be interpreted as a concerted elimination according to Fry (29) and Casanova and Rogers (30) BrCH2CHzBr+ 2e-
-
[Br-C-C-BrI2-
-2Br-
H,C=CH2 or as a process with discrete carbanion formation (31) BrCH2CH2Br + 2e-
y
-BrCH2CHz-
H,CCH2Br
+ OH-
was obtained during long time scale electrolysis. In slightly acid solution the value of p is significantly lower than two, which suggests involvement of the overall reaction
v, t d = 1 s, t , = 7.5 ms.
E = El,z
-
-Br-
H2C=CH2 In the reverse pulse experiments the stoichiometric coefficient obtained for bromides is slightly lower than two, which suggests the intervention of carbanion intermediates. The extraction of protons from water appears to be a slow process
BrCH2CH2Br +
H++ 2e-
-Br-
H3CCH2Br
and confirms the participation of carbanion in the reduction process of EDB. It should be added that at potentials more negative than -2.0 V, hydroxide ions are formed in the vicinity of the electrode during electrolysis in neutral solutions of supporting electrolyte alone. The presence of OH- ions prevents the formation of ethyl bromide. This explains the failure to yield products of proton capture during the electrolysis at very negative potentials (30) but does not imply the absence of carbanion. Analytical Application. A typical calibration curve in the dc mode fot EDB has standard deviation about the curve of0.020 pA and slope 3.79 pA/mM, which yields the estimate ( n = 5 ) of 15 pM EDB for the detection limit. This corresponds to about 2 pg of EDB in 1 mL of solution. The detection limit in the N P mode is about 10 times less. Alternatively, one may consider determination of EDB by using the anodic wave in the presence of bromide. Data similar to those ofFigure 4 give detection limits of about 5 pM. It would be possible to determine both inorganic bromide and EDB in the same sample by measuring inorganic bromide using the limiting current for bromide in the NP mode with initial potential at -0.2 V and then measuring EDB either by direct means, using its limiting current, or by difference using the bromide wave. The development of specific analytical methods for foodstuffs requires detailed considerations of sample type and treatment and is beyond the scope of the present work.
ACKNOWLEDGMENT We are grateful to Joe Magee for performing the GC experiments and to Chenniah Nanjundiah and Eric Mittlefehldt for their helpful collaboration, and we wish to thank Hector Fernandez for valuable suggestions. Registry No. EDB, 106-93-4; H20, 7732-18-5. LITERATURE CITED (1) Borman, S. A. Anal. Chem. 1084, 56, 573A. (2) Morris, S. C.; Rippon, L. E.; Haiamek, R. J . Chromatogr. 1982, 246, 136. (3) Iwata, Y . ; Duesch. M. E.; Gunther, F. A. J . A@. Fow'Chem. 1983, 3 1 , 171. (4) Keough, T.; Strife, R. J.; Rodriquez, P. A.; Sanders, R. A. J . Chromatogr. 1084, 372, 450. (5) Jonsson, A.; Berg, S. J . ChrMBtogr. 1980, 790,97. (6) Dumas, T.; Bond, E. J. J . Assoc. Off. Anal. Chem. 1982, 65, 1379. (7) Coilins, M.; Barker, M. J. Int. Lab. 1983, 13, 106. (8) Cairns, T.; Siegmund, E. G.; Doose, G. M.; Hundiey, H. K.; Barry, T.; Petzlnger, G. Anal. Chem. 1084, 56, 2138. (9) Carnahan, J. W.;Caruso. J. A. Anal. Chim. Acta 1982, 136, 261. (IO) Clower, M. Laboratory Information Bulletin; U S . Food and Drug Admtnlstratkm. U S . Government Prlnting Office: Washington, DC, 1980; 2338C. (11) von Stackelberg, M.; Stracke, W. 2. Electrochem. 1940, 53, 118. (12) Meites, L.; Zurnan, P. CRC Handbook Series in Electrochemistry; CRC Press: Cleveland, OH, 1978. (13) Feoktlstov, L. 0. In Ofpnic Ebctrochemstry; Baker, M. M., Lund. H., Eds.; Marcel Dekker: New York, 1983. (14) Abet, R. H.; Christie, J. H.; Jackson, L. L.; Osteryoung, J. G.; Osteryoung, R. A. Chem. Instrum. 1078, 7, 123. (15) Jackson, L. L.; Yarnitzky, Ch.; Osteryoung, R. A.; Osteryoung, J. G. Chem. B & d . Environ. Insbum. 1000, 10, 175. (18) Brumleve. T. R.; O'Dea. J. J.; Osteryoung, R. A.; Osteryoung, J. G. Anal. Chem. 1981, 53, 702. (17) Kemula. W.;Kublik, Z. Anal. Chim. Acta 1058, 18, 104. (18) Broadbeult, A. D. J . Chem. Educ. 1967, 4 4 , 145. (19) Flanagan, J. B.; Takahast, K.; Anson, F. C. J . Electroanel. Chem. 1977, 257. 257.
1969
Anal. Chem. 1986, 58, 1969-1972 Seidell, A. Solubilities of Inorganic and Organlc Compounds; D. van Nostrand: New York, 1940. Fry, A. J. I n Synthetic Organic Electrochemistry; Harper and Row: New York, 1972; p 171. Revenda, J. J. Collect. Czech. Chem. Commun. 1934, 6 , 453. Kolthoff, J. M.; Miller, C. W. J. Am. Chem. SOC. 1941, 63, 2732. biegler, T. J. €kcfroanal. Chem. 1963, 6 , 365. Osteryoung, J.; Kirowa-Eisner, E. Anal. Chem. 1980, 52, 62. Klrowa-Eisner, E.; Osteryoung, J. Anal. Chem. 1978, 5 0 , 1062. Bard, A. J.; Faulkner, L. R. Electrochemical Methods, Fundamentals and Applications; Wiley: New York, 1980; pp 168, 434. Oldham, K. B.; Parry, E. P. Anal. Chem. 1968, 4 0 , 65. Fry. A. J. Fortschr. Chem. Forsch. 1972, 3 4 , 1. Casanova, J.; Rogers, H. R. J. Org. Chem. 1974, 3 9 , 2409.
(31) Mann, C. K.; Barnes, K. K. €kctrochemica/ Reacflons in Nonaqueous Systems Marcel Dekker: New York, 1970; p 212.
RECEIVED for review November 12, 1985.
Resubmitted April
7, 1986. Accepted April 7,1986. This work was supported in part by the National Science Foundation under Grant Nos. CHE 7917543 and 8305748. R.T. gratefully acknowledges support from the Brazilian government foundation FAPESP (Fundacao de Amparo a Pesquisa do Estado de Sao Paulo). J.O. thanks the Guggenheim Foundation for their support.
Simultaneous Polarographic Determination of Parathion and Paraoxon. Catalytic Hydrolysis of Parathion by Palladium(I I) J. HernPndez M6ndez,* R. Carabias Martinez, and J. Sdnchez Martin Department of Analytical Chemistry, Faculty of Chemistry, University of Salamanca, Salamanca, Spain
The present work describes an electroanalyticalstudy of the poiarographlc behavior (DPP) of the pesticides paratwon and paraoxon In the presence of Pd( I I). Thls metalilc Ion shows affinity for the thiophosphate group and catalyzes the hydrolysis of paratMon but not paraoxon. A method Is proposed for the simultaneous determination of both pestlcldes based on the fact that parathion can be determlned by measuring the p-nitrophenol formed after the addnlon of Pd(II), whereas paraoxon can be measured directly by Its reductton peak. I n the determination of parathion, acceptable errors were found as long as the parathlon/paraoxon ratio Is greater than 1/45. I n the determlnatlon of paraoxon, satlsfactory results were obtained for paraoxon/parathion ratlos greater than 1/70.
The determination of mixtures of parathion and paraoxon is of great interest from an analytical point of view in part because the control of these pesticides in environmental analyses is important and also because parathion is metabolized in vivo to yield paraoxon and p-nitrophenol, among other species ( 1 ) . parathion
paraoxon
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C 2 H 5 0 \ ! - 0 0 N 0 2
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C2H50'
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The difficulty inherent to the simultaneous electroanalytical determination of the two pesticides lies in the fact that both of them possess the same electroactive group (R-NOJ and the remaining parts of both molecules are not very different either in size or in polarity. However, p-nitrophenol is reduced a t more negative potentials, such that the simultaneous determination of parathion or paraoxon with p-nitrophenol is possible (2, 3). The simultaneous electroanalytical determination of parathion and paraoxon has been studied elsewhere (4, 5), although the methods proposed do not seem to have been very satisfactory. Smyth et al. (5) have proposed an indirect method based on the different hydrolysis rates in alkaline media; measurements are carried out after 25 min, determining 0003-2700/S6/0358-1969$01.50/0
parathion directly and paraoxon by difference. The method is only applicable if the paraoxon/parathion ratio is lower than 3. In the present work two polarographic procedures (DPP) are proposed for the simultaneous determination of parathion and paraoxon; both of them are based on the different hydrolysis rates exhibited by these pesticides in the presence of Pd(II), a cation which shows a selective affinity for the thiophosphate group (6, 7)and which catalyzes the hydrolysis of parathion but not paraoxon. Other metallic ions (Cu2+and Hg2+) of known catalytic activity in the hydrolysis of organophosphorus pesticides (8-10) were also assayed though the best results were obtained with Pd(I1).
EXPERIMENTAL SECTION Reagents. Solutions of parathion and paraoxon in 50% MeOH/H20 (v/v) prepared from 99% pure commercial products (Riedel-De-AG, Seelze-Hannover). Solutions of p-nitrophenol and of potassium 0,O-dimethyl thiophosphate were prepared in 50% MeOH/H20 (v/v). Aqueous solutions of Pd(I1) were prepared from PdC12. Buffer H9P04,HAC, HB02, and aqueous solutions of NaOH were also used. All reagents were of analytical grade. Apparatus. A Metrohm E-506 polarograph, P-Selecta thermostat, and Crison 501 pH meter were used. The electrodes used were Metrohm EA-1019/1 mercury-drop electrode, an auxiliary Metrohm EA-285 platinum electrode, and a KCl saturated calomel electrode. F rocedure. Fifty milliliters of a solution containing parathion and/or paraoxon in 30% MeOH/H20 medium (v/v), 0.12 M buffer, and variable amounts of NaOH are placed in a cell thermostated at 25 "C. The oxygen is eliminated by bubbling N2 through the solution for 15 min, and the corresponding polarogram (DPP) is recorded between 0.0 and -2.0 V with a pulse amplitude of -50 mV and a scan rate of 4 mV s-l. Following this, variable amounts of Pd(I1) are added, and the concentration of p-nitrophenol in the solution is followed polarographically or amperometrically with time. The pH was measured with a glass electrode calibrated with aqueous buffer solutions. The pH value measured is corrected to obtain true values of proton activity in hydroalcoholic medium (pH* = pH - 0.05). RESULTS AND DISCUSSION Reaction of Parathion and Paraoxon with Pd(I1). In 50% (v/v) MeOH/H20-0.1 M HAc-O.1 M NaAc medium, 0 1986 American Chemical Soclety
