of a significantly larger charging current in intermodular ac polarography favors the use of second harmonic ac polarography. The magnitude of the charging current is the main factor determining the sensitivity of a polarographic technique. The limit of detection of the intermodular is therefore expected to be higher than the second harmonic. In the present work, it was found that the second harmonic method is about an order of magnitude more sensitive. Careful tuning of phase angles to minimum background current may slightly modify this conclusion, but this would not be an analytically expedient procedure. The instrumentation for the intermodular developed in this work is also more complicated than the simple multiplier approach developed (I1) for phase-selective second harmonic ac polarography. The need for an additional oscillator and phase-selective detector make the intermodular technique less attractive from the cost point of view. Finally, second harmonic polarography may be used over a wide amplitude and frequency range without any adjustment of the circuitry. On the other hand, the intermodular method has a restrictive frequency range and any changes to either frequency or amplitude require balancing of the two oscillators for the new conditions. In conclusion, for most analytical or electrode-kinetic applications no advantage is seen in using phase-selective intermodular over second harmonic ac polarography, and the latter technique is strongly recommended.
LITERATURE CITED R. Neeb, Naturwissenschaffen, 49, 447 (1962). J. Paynter, Doctoral Thesis, Columbia University, New York, 1964. W. H. Reinmuth, Anal. Chem., 36, 211R (1964). T. G. McCord, E. R. Brown, and D. E. Smith, Anal. Chem., 38, 1615 (1966). (5)R. D. Jee, Fresenius’ 2.Anal. Chem., 264, 143 (1973). (6) H. H. Bauer and P. J. Elving, Anal. Chem., 30, 341 (1958). (7) D. E. Smith in “Electroanalytical Chemistry”, A. J. Bard, Ed., M. Dekker, New York, 1966, Chap. 1 and references clted therein. (8) D. E. Smith, Crit. Rev. Anal. Chem., 2, 247 (1971) and references cited therein. (9) A. L. Woodson and D. E. Smith, Anal. Chem., 42, 242 (1970). (IO) H. Blutstein and A. M. Bond, Anal. Chem., 46, 1531 (1974). (1 1) H. Blutstein, A. M. Bond, and A. Norris, Anal. Chem., 46, 1754 (1974). (12) H. Blutstein and A. M. Bond, Anal. Chem., 46, 1934 (1974). (13) H. Blutstein and A. M. Bond, J. Electroanal. Chem., 56, 177 (1974). (14) Instruction Manual for PAR Electrochemistry System, Model 170, Section 11, Detector and Signal Processing Board, p IX-3, Princeton Applied Research Corp., Princeton, N.J., 1972. (15) J. R. Delmastro and D. E. Smith, Anal. Chem., 38, 109 (1966). (16) J. R. Delmastro and D. E. Smith, J. Electroanal. Chem., a, 192 (1965). (17) T. G. McCord and D. E. Smith, Anal. Chem., 41, 131 (1969). (18) T. G. McCord and D. E. Smith, Anal. Cheni., 42, 126 (1970). (19) I. Ruzic and D. E. Smith, Anal. Chem., 47, 530 (1975). (20) A. M. Bond, J. Electrochem. Soc., 118, 1588 (1971) and references clted therein. (21) A. Zatka, J. Nectroanal. Chem., 27, 164 (1970). (22) L. A. Matheson and N. Nichols, Trans. Am. Electrochem. Soc., 73, 193 (1938). (23) C. I. Mooring, Polarogr. Ber., 6, 63 (1958). (1) (2) (3) (4)
RECEIVEDfor review April 5,1976. Accepted August 6,1976. The authors express their appreciation to the Australian Research Grants Committee for financial support.
Intermetallic Compound Formation between Copper and Zinc in Mercury and Its Effects on Anodic Stripping Voltammetry Mark S. Shuman* and George P. Woodward, Jr. Department of Environmental Sciences and Engineering, School of Public Health, University of North Carolina, Chapel Hill, N.C. 275 14
Several Cu-Zn intermetallic compounds form during anodic stripping voltammetry (ASV) analysis of solutions containing both copper and zinc. There are three soluble compounds with copper to zinc ratlos of 1:1, 1:2, and 1:3. In addition, an insoluble compound also forms at high amalgam concentrations and has a copper to zinc ratio of 1:3. The formation of these compounds decreases the ASV zinc current and increases the copper current since the soluble compounds are electroactive and are oxidized at a potential very close to copper stripping potential. This interferencein the determination of copper and zinc by ASV is most serious with thin film electrodes where the small mercury volume leads to very high amalgam concentrations. Stepwlse instability constants for CuZn, CuZnp and CuZn3, respectively,were found to be K, = 1.9 X K2 = 7.6 X K3 = 2.1. The solubility product for the insoluble compound, CuZn3(s), was estimated to be Ksp= 3.1 X 10-5.
The formation of intermetallic compounds can cause error in the analysis of metals by anodic stripping voltammetry. For example, Cu-Zn compounds can interfere in the determination of Cu or Zn in environmental samples where, of the trace metals readily analyzed by ASV, these two metals are in greatest abundance. The formation of Cu-Zn compounds can be written
aCu
+ b Zn = Cu, Znb
Russell et al. ( I ) studied these by dissolving copper and zinc in a Hg pool and reported formation of several compounds with varying stoichiometry, a and b. Kemula et al. (2) were the first to recognize that these compounds interfered with the ASV analysis of Cu and Zn. The observed effect was a decrease in sensitivity of ASV to Zn when Cu was present. More recently Stromberg and co-workers ( 3 ) studied the formation of Cu-Zn compounds by ASV. Their method was to measure the diminished Zn stripping current when Cu was varied from 0.1 to 2 times the concentration of Zn. They concluded that the stoichiometry was 1:l and by assuming a priori that the compound in mercury was insoluble, calculated a solubility product of 5 X lo+. Kozlovsky and Zebreva ( 4 ) in reviewing the literature, observed that the solubility product values obtained from potentiometric data ranged from 2.8 X 10-5 to 7.1 X Further work by Stromberg and co-workers (5)and by Mesyats et al. (6) in which the stripping current of both copper and zinc was followed as a function of pre-electrolysis time indicated that a 1:l Cu-Zn compound was formed and was sparingly soluble with a solubility product of 1 X lov6,although their data did not fit their theory based on limited solubility. Rudolph (7) has also studied the electrochemical formation of these compounds and assumed a priori a limited solubility.
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
1979
He found a 1:l stoichiometry and a solubility product of 6.6 x 10-6. The effects of this intermetallic compound formation on the practical application of ASV in the analysis of Cu and Zn has been generally disregarded. Several papers, however, are worth mentioning. Bradford (8) used a thin mercury film electrode (TFE) and noted suppression of the zinc stripping peak by the presence of copper in artificial seawater. He as: sumed that the stripping peak height was proportional to free zinc in the amalgam and also that the area of the Zn stripping peak reflected both free zinc and zinc reacted with copper in the amalgam (the free and complexed zinc was assumed to be separated by an anodic tailing of the zinc stripping process). Using these assumptions and holding zinc always in excess, he concluded that the intermetallic compound had a 1:l stoichiometry and a formation constant of 1 X lo5. Seitz (9) in contrast noted that the area of Zn peak decreased with additions of Cu to seawater. Shuman and Woodward (10) found that not only did the zinc stripping peak diminish in height and area in the presence of copper, but that the copper stripping peak was enhanced in the presence of zinc. Analysis of the same solutions with atomic absorption confirmed this interference. The zinc peak height and area was diminished to about 15% when copper and zinc were present at nearly equal molar concentrations M), and the copper peak area was larger than for a solution containing copper alone. The data also suggested a stoichiometry different than 1:l. Recently, Copeland et al. (11)eliminated the formation of Cu-Zn compounds by adding Ga which appears to form intermetallic compounds itself but not to cause interference. Crosmun et al. (12)noted a decrease in zinc stripping current with added copper and assumed a 1:l stoichiometry. The data presented in their paper could also be interpreted as indicating formation of compounds with Zn to Cu ratios greater than unity. Several approaches to estimating the stoichiometry of intermetallic compounds have been suggested by previous workers using either ASV or controlled potential electrolysis. Stromberg et al. ( 3 , 5 )derived solubility product expressions to use with ASV data and assumed that stripping currents reflected the concentration of unassociated and soluble Cu or Zn in the amalgam. To solve for the stoichiometric ratio of the Cu-Zn compound, either the solution concentration of the two metals or the pre-electrolysis time was varied in a region where the solubility product was exceeded and then stripping currents were measured. This method, of course, ignored soluble intermetallic compounds that could form and be oxidized during the stripping step. Hovsepian and Shain ( 1 3 ) used another approach to study Co-Zn compounds. The compounds they studied are definitely soluble in mercury and their formation was followed by the decrease of Zn current in the presence of Co. Simultaneous equilibrium expressions were solved with data obtained by ASV at two pairs of metal concentrations and gave both the stoichiometry and the formation constant. Rodgers and Meites (14) used controlled potential electrolysis to study Ni-Zn compounds. They assumed that when zinc was in excess, all the nickel was consumed by intermetallic compound formation, and that the stoichiometric ratio of zinc to nickel could then be determined by the calculated ratio of bound zinc to the total analytical concentration of nickel present in the amalgam, where unbound zinc was estimated from the zinc stripping current. They found that this ratio increased with increasing excess zinc which suggested the formation of more than one Ni-Zn compound. It was the intention of the work described here to gain an understanding of the stoichiometry and stability of Cu-Zn intermetallic compounds in order to predict interferences and to generate equations that would be useful in correcting for
i~eo
this interference. A number of separate methods were used. First, solubility expressions were derived for the formation of soluble Cu-Zn compounds with stoichiometry a = 1and b = 1to 3 and the formation of insoluble species a = 1,b = 1to 3. These were tested against experimental data for the best fit. Second, the ratio of free metal in the amalgam to total metal in the amalgam was followed by measuring Cu and Zn stripping currents as a function of amalgam concentration varied by increasing pre-electrolysis time. Third, this same ratio for zinc was followed as a function of the analytical concentration ratio of Cu:Zn from 0-0.8. Finally, the mechanism of the formation of these intermetallic compounds was examined by comparing cyclic voltammetry theory (15, 16) with current-voltage curves obtained by reducing Zn a t a hanging drop electrode composed of Cu amalgam.
EXPERIMENTAL The electronic potentiostat and three-electrode cell used in this work were identical to those described previously (10). The anodic stripping voltammetry was carried out with a thin film of mercury plated on a glassy carbon electrode (Chemtrix Corp.) or with a microburet hanging mercury-drop electrode (Metrohm). The plating bath was approximately 0.02 g of HgO dissolved in 100 ml of 1 M HC104 and had a mercury pool anode. A 30-second plate gave a thin film volume of 1.44 X cm3. The volume of the mercury drop was 6.23 X cm3. All stripping voltammetry experiments were done with a 43 mV/s linear scan rate from -1.250 V to +0.150 V vs. SCE. Pre-electrolysis was carried out at -1.250 V vs. SCE for 1-10 min with the thin film and for 1 min with the hanging drop. The supporting electrolyte for all thin film experiments was 0.05 M acetate buffer a t pH 4.5 and for all hanging drop experiments was 0.1 M KBr. The cyclic experiments were made from -0.525 V to -1.125 V vs. SCE with a scan rate of 40 mV/s and after a 60-s deposition of copper a t -0.525 V. Copper and zinc solutions were prepared from the metals.
RESULTS AND DISCUSSION Stoichiometry. Cyclic Voltammetry. Cyclic voltammetry was performed on a series of solutions that contained a constant concentration of zinc and increasing concentrations of copper. One solution contained 10 ppm Zn and 0-60 ppm Cu in order that Cu would be in excess in the amalgam, and two other solutions contained 200 and 300 ppm Zn with Cu concentrations 0-60 ppm in order that Zn would always be in excess. The initial potential of these cyclic voltammograms was selected anodic of zinc reduction and cathodic of copper reduction so that the zinc when reduced was reduced at a copper amalgam electrode. Analysis of the cathodic peak of zinc, its height and peak width, its anodic peak height, and ratio of cathodic to anodic peak heights was made with regard to electrode reactions orders other than simple first-order reactions of the type bZn+2e-
Cu amalgam
CuZnb
The results for a 200-ppm Zn solution are shown in Figure 1. This figure shows that the anodic peak height of zinc decreases sharply as the copper concentration is increased from zero to 6.3 X lob4M (40 ppm) and that the cathodic peak also decreases slightly. Figure 2 presents the cathodic E , - E+ for increasing copper concentrations where E , is the potential at the peak and Ep/2is the potential at half peak height. This method of measuring peak height is that suggested by Nicholson and Shain (15). In the same copper range as Figure 1, the peak width is seen to broaden from 35 to 50.5 mV. Zinc is in large excess in these voltammograms and, if it is assumed that reaction with copper in the amalgam is nearly complete, it is obvious that the decrease in the anodic zinc current with added copper reflects formation of an intermetallic compound with a zinc to copper ratio of 2:l. This would suggest an electrode reaction mechanism second-order in zinc. For a twoelectron, second-order reduction, the cathodic E , - E,/z is
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
50 -
40 #I
4
9 30-
a
i
3
0
LT
u
0
g
Figure 3. Qz,/@,
a
eo+
N
vs. R; Dashed line is slope of -1
( A ) 0.5 ppm zinc. (0)1.0 pprn zinc. ( 0 )2.0 ppm zinc
A
IO r
I
O
h
03
04 R
A
J
00
20 COPPER
40
60
80
100
CONCENTRATION x IO4 M
Figure 1. Zinc peak currents for cyclic voltammograms of 200 ppm zinc solution as a function of copper concentration ( 0 )Anodic peak. ( 0 )Cathodic peak
6ol 50
00
-
06
0 7
0 8
R
[ A ) 0.5 ppm zinc. (0)1.0 ppm zinc. ( 0 )2.0 pprn zinc
1-
10
+ Q& VS.
05
a
'b 2 30-
-
02
Figure 4. Qzn t QcU/Q;,
40-
20
01
I
40.7 mV compared to 28.5 mV for first-order, and the cathodic current function is 0.353 compared to 0.443 for first-order (16). Therefore Figures 1and 2 indicate that as copper concentration increases and a greater fraction of the total reduced zinc reacts with copper in the amalgam, the cathodic peak current and peak width take on greater second-order character. Between 0 and 40 ppm copper, the cathodic peak height decreases from 27.5 to 22.5 MAor a decrease of 18%,nearly the 20% decrease expected for a shift from first- to second-order. Over this same copper range, a 15-mV increase in peak width is observed compared to a 12-mV increase expected if the reaction were to change from first- to second-order. Similar observations were made for 300 ppm Zn except, in its case, the decrease in anodic peak height with added copper suggests a zinc to copper ratio for the intermetallic compound of 3:l. As with 200 ppm Zn, a decrease in cathodic peak height and broadening is observed. The peak broadens from 35 to 61 mV from 0 to 40 ppm Cu, a larger potential shift than in the previous case. The correlations with 300 ppm Zn are less clean cut than at 200 ppm, possibly because of the larger excess zinc and perhaps because the electrode reaction is a mixture of secondand third-order reactions. Cyclic voltammograms where copper is in excess of zinc, 10 ppm Zn, indicated an intermetallic compound with zinc to copper ratio of 1:l.Although the cathodic peak broadens and
decreases in height with increased copper concentration, the relative change in these two parameters is smaller than when zinc is in excess. Thus the mechanism, although not exclusively first-order in this instance, appears to have greater first-order character than when zinc is in excess. Cu-Zn Compounds in a Thin Mercury Film Electrode. Anodic stripping experiments were carried out using a TFE and three concentrations of zinc, 0.5,1.0, and 2.0 ppm. Copper was varied from 0.1 to 0.4 ppm to give a molar ratio, R , of copper to zinc from approximately 0.05 to 0.8. Since it was found that the accumulation coefficients for Cu and Zn were practically identical (IO),this also is the molar ratio of the metal within the amalgam. Figure 3 shows the integrated stripping current of zinc expressed as a ratio of the zinc current with copper present, Qzn, to the zinc current at the same concentration with copper absent, &in,, plotted as a function of R. If a 1:l stoichiometry prevailed and if the intermetallic compound had a very small dissociation quotient or solubility product, all data points should be on the line with slope of -1 and intercept of y = 1.0, x = 0. The data points, however, have a slope steeper than -1 indicating stoichiometry greater than 1:l and show considerable dissociation or solubility a t R = 0.8. The question of whether the compounds are soluble or insoluble can also be addressed using these data. In Figure 4, the ratio of total copper and zinc stripping currents (both in the same solution) to the sum of the stripping currents of copper and of zinc (each in separate solutions but a t the same concentration) is plotted vs. R. Although these data suffer from lack of precision for the 0.5-ppm zinc solutions where replicate stripping peaks had a relative average deviation of 5%, they suggest that the solubility product is exceeded at least for the 1.0- and 2.0-ppm zinc solutions (where the relative average deviation was about 2%) and that a larger percentage of the total metal is precipitated at the higher zinc concentration. It appears that no more than about 16%of the total metal is electroinactive and assumed to be precipitated, indicating that the major portion of the intermetallic compounds formed in the experiments is soluble.
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
1981
I:
0 0
30/
U s 50
0
s40W
O
a O30-
z
0 I
0
4 5 6 7 ACCUMULATION TIME IN MIN.
2
3
8
9
0
20 -
0 IO
0
Figure 5. Q,/@ vs. accumulation time, 1 ppm zinc and 0.4 ppm cop-
I
per
00
I
20
0
0
02-
0
0
P
If instead of varying the solution concentrations of the metals, the time of accumulation is varied from 1to 10 min in order to increase the amalgam concentration of the metals, further information on the stoichiometry is obtained. The data in this case were examined at each accumulation time by taking the ratio Q t / Q f ,the stripping current of one metal in the presence of the other to the stripping current for that metal in the absence of the other metal and also at that accumulation time. For excess zinc, 1 ppm Zn and 0.4 Cu, this ratio for copper approaches 3.0 at 10 min indicating a stoichiometry of CuZnz (Figure 5 ) , whereas for nearly equal molar concentrations of copper and zinc, 1ppm each, the same ratio is about 1.8 and also increasing with accumulation time indicating a stoichiometry of CuZn (Figure 6). Therefore as might be expected, the stoichiometry is dependent on the relative concentrations of the two metals in the amalgam.
100 CONCENTRATION x 104M
60
COPPER
(0) Copper stripping current. (0) Zinc stripping current
0
40
80
12.0
140
Figure 7. Zinc stripping current vs. copper concentration for hanging drop electrode and 10 ppm zinc
From these experiments, a complex picture develops; within the mercury, several soluble Cu-Zn compounds form which are in equilibrium with one or more insoluble Cu-Zn compounds. Instability Constants for Intermetallic Compounds. Hanging Drop Electrode. Anodic stripping voltammetry experiments were carried out with the hanging drop electrode in solutions of 10 ppm zinc and 0-80 ppm Cu. In Figure 7, the zinc stripping current for a 1-min accumulation is plotted as a function of copper concentration. The accumulation coefficient for zinc, that is, the ratio of zinc amalgam concentration to zinc solution concentration after 1-min accumulation ( 17) was 19.7 obtained by using a method suggested by Shain and Lewinson (18). The copper accumulation coefficient was identical within experimental error. The dissociation constant was calculated from these data for copper concentrations equal to or in excess of zinc where the compound was considered soluble with a stoichiometry of 1:l and where the zinc stripping current was taken as a measure of unreacted zinc. The calculated constant was 1.9 f 0.2 X If instead, the 1:lcompound was assumed a priori to be insoluble, a solubility results from these data. product of 3.7 f 1.1 X Thin Film Electrode. When a thin film electrode is employed, the solubility of some Cu-Zn compounds appears to be exceeded. Whenever this occurs, mass balance and solubility expressions can be formulated for the species assumed to be present. If soluble CuZn, CuZnz, and CuZns species are present and an insoluble species, CuZns(s) forms, then solubility = [CUI [CuZn] [CuZnz] [CuZns], or,
+
+
+
Table I. Data from Thin Film Electrode Experimentsa
Cgu, P P ~ C$, ppm 0.0 0.1 0.2 0.3 0.4 0.0 0.1 0.2 0.3 0.4 0.0 0.2 0.4 a
0.5 0.5 0.5 0.5 0.5
1.0 1.0 1.0 1.0 1.0 2.0 2.0 2.0
Q&, PC
Qcu,kC
QZn, PC
0.0 3.84 7.52 9.80 13.7 0.0 3.84 7.52 9.80 13.7 0.0 7.52 13.7
0.0 7.06 12.2 18.4 22.5 0.0 9.80 15.8 20.2 25.4 0.0 15.3 30.9
18.1 13.6 10.2 7.64 6.27 40.7 33.4 30.3 24.5 21.4 85.8 71.2 56.3
Qzn/Q8n
1.0 0.75 0.56 0.42 0.35 1.0 0.82 0.74 0.60 0.53 1.0 0.83 0.66
Values in this table are the mean of four to seven replicate stripping curves.
1982
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13, NOVEMBER 1976
AQcu,zn, CLC
... 1.72 3.30 1.94 2.74
...
1.94 2.28 5.84 7.76
...
6.26 20.4
[Znl,M 0.66 0.49 0.37 0.27 0.22 1.5 1.2 1.09 0.88 0.77 3.1 2.6 2.0
Solubility, M
*.. 0.12 0.24 0.33 0.47
...
0.12 0.25 0.30 0.42
...
0.21 0.32
Solubility = K,,
(m+ 1
1 K1[Zn12 K1Kz[ZnI ~
l
+
where bracketed concentrations refer to amalgam concentrations and K,, = [Cu][Zn13,K3 = ([Zn][CuZnz])/[CuZns], K2 = ([Zn][CuZn])/[CuZn2], K1 = ([Zn][Cu])/[CuZn]. The analytical concentration of copper species in the amalgam is
+ 2CuZns(s)
(4)
where Qg, is the total coulombs of electricity obtained from the stripping curve of copper without zinc present; V is the volume of the mercury electrode; and F is the faraday. Also, the difference between the total electricity obtained for a solution containing both copper and zinc and that obtained for two solutions, one with zinc and one with copper, AQcu,zn = 8CuZn3(s). Therefore, the solubility can be expressed by the experimentally accessible quantity, QOcu AQCu,Zn solubility = - -12VF 8 VF The data and solubilities obtained for the T F E experiments are in Table I. To obtain instability constants, solubility was plotted vs. 1/[ZnI2 for 0.5 ppm Zn solutions in which zinc was in least excess and where CuZn appeared to predominate. Its slope, according to Equation 3, is K,,IK1. Using K1 of 1.9 X obtained from hanging drop electrode data, K,, was calculated to be 3.1 X Similarly, solubility was plotted against 1/ [Zn] for 1and 2 ppm Zn where zinc was in greatest excess and
CuZn appeared to form. Its slope of K,,/KlKz gave Kz = 7.6 X and its intercept gave K 3 = 2.1. Several other combinations of species and stoichiometry were attempted and several sets of solubility expressions like Equation 3 were formulated. These gave much larger variabilities in calculated solubility products and instability constants than the assumed stoichiometry.
LITERATURE CITED (1) A. S. Russell, P. V. F. Cazalet, and N. M. Irving, J. Chem. Soc., 652
(1932). (2)W. Kemula and 2 . Kiblik, Nature (London), 182, 1228 (1956). (3)A. G. Stromberg and V. E. Gorodovykh, Zh. Neorg. Khim., 6, 2355 (1963). (4) M. Kozlovsky and A. Zebreva, in "Progress in Polarography", Vol. Ill, P. Zuman and I. M. Kolthoff, Ed., Interscience, New York, 1972,p 157. (5) A. G. Stromberg, M. S. Zakharov, and N. A. Mesyats, Elektrokhimiya, 3, 1440 (1967). (6)N. A. Mesyats, A. G. Stromberg, and M. S. Zakharov, Elektrokhimiya, 4, 987 (1966). (7)R. G. Rudolph, Ph.D. thesis, University of Nebraska, Lincoln, Neb., 1969. ( 6 ) W. L. Bradford, Chesapeake Bay Institute, Report No. 76,Johns Hopkins University, Baltimore, Md.. 1972. (9) W. R. Seitz, Ph.D. Thesis, MiT, Cambridge, Mass., 1970. (IO)M. S.Shuman and G. P. Woodward, in "Trace Substances in Environmental Health VI, Proceedings of the University of Missouri Sixth Conference on Trace Substances in Environmental Health", D. D. Hemphill, Ed., 1972, p 269. (11) T.R. Copeland, R. A. Osteryoung, and R. K. Skogerboe, Anal. Chem., 46,
2093 (1974). (12)S. T. Crosmun, J. A. Dean, and J. R. Stokely, Anal. Chim. Acta, 75, 421 11975). - -, (13)B. K. Hovespian and I. Shain, J. Electroanal. Chem., 14, 1 (1967). (14)R. S.Rodgers and L. Meites, Electroanal. Chem., 38, 359 (1972). (15) R. S.Nicholson and I. Shain, Anal. Chem., 36, 706 (1964). (16)M. S.Shuman, Anal. Chem., 41, 142 (1969). (17)E. Barendrecht, in "Electroanalytical Chemistry", Vol. 4,A. J. Bard, Ed., Marcel Dekker, New York, 1966,p 53. (18)I. Shain and J. Lewinson, Anal. Chem., 33, 167 (1961). \
RECEIVEDfor review June 10,1976. Accepted August 4,1976. This research was supported by the Oceanographic Section, National Science Foundation, NSF Grant DES 73-06451 A02.
Coulometric-Spect ropolarimetric Titrations of Metal Ions with Chiral Ligands Richard A. Gibbs' and Robert J. Palma, Sr." Department of Chemistry, Midwestern State University, Wichita Falls, Texas 76308
The determination of metal ions In aqueous solution based on the reaction between mercury(l1) complexes of ( R ) - ( - ) -propylenedlamlnetetraacetic acid and (R,R)-(-)-trans1,2-cyclohexanediamlnetetraaceticacid with representative metal Ions Is described. Two reaction sequences account for the shapes of the tltration curves obtained. A mercury(l1) complex is amperostatically reduced and the chiral chelon is quantitatively liberated. Alternately, the stability constant of the mercury complex Is lowered with auxlllary complexing agents, so that the metal Ions can exchange and liberate mercury(II) Ions for subsequent reduction. Calcium, magnesium, lead, cadmium, and zinc have been determined In the range of 10 pequlv to 250 pequlv wlth a mean error of 0.29% and a mean relative average deviation of 0.51%. Present address, Department of Chemistry, Purdue University, West Lafayette, Ind. 47907.
Spectropolarimetric titrimetry was first introduced by Kirschner et al. (1, 2) and has enjoyed considerable recent success. This was due to the introduction of several relatively inexpensive photoelectric polarimeters and subsequent improvement in their design ( 3 , 4 )and utility. The applications of this technique have been significantly broadened by Pearson et al., who utilized the two chiral and stereospecific aminopolycarboxylic acids (R)-(-)-1,2-propylenediaminetetraacetic acid, [R(-)PDTA], and (R,R)-(-)-trans-1,2cyclohexanediaminetetraacetic acid, [R,R(-)CDTA], in the volumetric titrimetric analysis of over 39 metals (5-17). In other studies, molybdenum(V1) was determined with R (-) PDTA in the presence of equimolar tungsten(V1) (18). The stability constants of most metal complexes with PDTA or CDTA are considerably higher than those of EDTA complexes (19). The titrant and complexes served as self-indicators, thus permitting the maximum quantitative pH range of the metal complexes to be utilized in these analyses. The end point was
ANALYTICAL CHEMISTRY, VOL. 48, NO. 13. NOVEMBER 1976
1983
