2104 (23) (24) (25) (26) (27) (28) (29) (30)
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978 R. A. Johnsonand J. D. O'Rourke, J. Am. Chem. Soc., 76, 2124(1954). A. E. Nielson, J . Colloid. Sci., IO, 576 (1955). J. D. O'Rourke and R. A. Johnson, Anal. Chem., 27, 1699 (1955). S. A. Goidstein and J. P. Walters, Spectrochim. Acta, Part 6 ,31, 201 (1976). S. A. Goldstein and J. P. Walters, Spectrochim. Acta, Part B , 31, 295 (1976). S. G. Salmon and J. A. Holcombe, Anal. Chem., 50, 1714 (1978). L. de Gabn and G. F. Samaey, Spectrochim. Acta. Part 6,25, 245 (1970). J. E. Chester, R. M. Dagnall, and M. R . G. Taylor, Anal. Chim. Acta, 51, 95 (1970).
(31) R. D. Cadle, "Particle Size Theory and Industrial Applications", Reinhold Publishing Corp., New York, 1965. (32) G. M. Hidy and J. R . Brock, "The Dynamics of Aerocolloidal Systems", Pergamon Press, New York, 1970. (33) W. E. Ranz and V. B. Wong, Ind. Eng. Chem., 4 4 , 1371 (1952). (34) R. K. Skogerboe and K. W. Olson, Appl. Spectrosc., 32, 181 (1978). (35) C. B. Boss and G. M. Hieftje, Anal. Chem., 49, 2112 (1977). (36) D. J. Halls, Anal. Chim. Acta, 88, 69 (1977).
RECEIVED for review June 1, 1978. Accepted October 2, 1978.
Reversibility of Copper in Dilute Aqueous Carbonate and Its Significance to Anodic Stripping Voltammetry of Copper in Natural Waters Mark S. Shuman" and Larry C. Michael Department of Environmental Sciences and Engineering, School of Public Health, University of North Carolina, Chapel Hi//, North Carolina
275 14
Cyclic voltammetry and anodic stripping voltammetry of copper in dilute carbonate solutions approximating natural fresh water indicated carbonate alkalinity and pH affected copper reversibility, possibly through variations in buffer capacity or rate of C 0 2 hydration. Recommendation was made that shifts in anodic stripping peak potential or changes in stripping current magnitude with gross changes in solution composition of pH and alkalinity not be used to indicate changes in trace metal speciation.
Knowledge of the chemical state of trace metals in solution is important to understanding trace metal transport, reactivity, and toxicity in aquatic and marine environments. Ion selective electrodes and anodic stripping voltammetry (ASV) are the most popular techniques for studying speciation because of their selectivity and sensitivity. There are several approaches to the application of ASV. In one approach, shifts in stripping peak potentials are interpreted in t h e same way half-wave potentials are interpreted in classical polarography, as a n indication of the formation of reducible complexes in solution (1-4). There is no theoretical basis for this analogy with polarography, but assumptions about electrochemical reversibility are usually made to make the analogy plausible. In a second approach, peak current variations with pH, organic chelate content or carbonate alkalinity are observed and related to t h e formation or dissolution of nonreducible complexes, solid phases, or colloidal species (1.5-9). In yet another approach, ASV is used to measure reducible metal during a complexometric titration. Serial additions of metal are made t o a solution containing ligands which form nonreducible complexes (1,10-12). T h e end point k r e l a t e d to t h e solution concentration of these ligands and the current magnitude is related t o t h e conditional formation constant of t h e complex (13). T h e first two approaches rely on strict electrochemical reversibility, or adherence to a given degree of irreversibility, throughout gross solution changes of pH, alkalinity, dissolved organic and dissolved solid content, a requirement that is probably unrealistic. T h e p H and alkalinity of natural water samples vary widely depending on sample location, season, dissolved organic carbon content, temperature, and other factors, and these variations can be expected to affect one or more steps in the electrochemical 0003-2700/78/0350-2104$01.OO/O
reduction of metals during ASV analysis. Therefore, it is important if these approaches are to be used that solution variation effects on reversibility be investigated. Although equilibrium models of trace metals in natural waters are available t h a t consider wide ranges of solution conditions (14-16), there is considerable disagreement among models for Cu in fresh water (3,17,18) and none can be used with certainity to predict principal Cu species a t near neutral p H or to predict electrochemical reversibility. No systematic electrochemical study of Cu reduction in dilute neutral carbonate solutions has been reported although Meites (19) studied the polarographic reduction of copper in 0.25 M carbonate, p H 9.5 to 11.0, and found a doublet wave which he attributed to a stepwise reduction through the cuprous state. Above p H 11.0, a blue precipitate of cupric hydroxide appeared in his solutions. Ernst et al. ( 3 ) performed differential pulse polarography and differential pulse anodic stripping on solutions of varying carbonate alkalinity and pH and reported Cu reduction approximated reversible behavior throughout these changes; however, they reported formation constants for CuC03 based on measurements of peak potential shifts an order of magnitude lower than previously reported values. Low ligand numbers and low formation constants were also obtained for a suggested C U ( C O ~ ) ~species. 'In addition to carbonate, their solutions contained 0.1 M K N 0 3 . T h e principal objective of the work reported here was to investigate reversibility of Cu reduction in dilute aqueous carbonate solutions that approximated conditions in natural fresh water (except elevated Cu concentrations were used) and to relate ASV of Cu in these solutions to reversibility. Cyclic voltammetry was used and carbonate alkalinity, pH, and scan rate were varied over a limited range. A secondary objective was to evaluate the method of using voltammetric peak potentials to identify Cu species undergoing reduction.
EXPERIMENTAL All reagents were reagent grade and used without further purification. Standard copper(I1) solution was prepared by dissolving 1.00000 g of 99.999% pure copper wire (J. T. Baker and lo-* N alkalinity were Chemical Co.). Solutions of W4, prepared from 0.1 N sodium carbonate (J.T. Baker Chemical Co.). Potassium perchlorate (Fisher Chemical Co.) and potassium nitrate (Matheson, Coleman and Bell) were used to reduce cell resistance in some experiments. The pH was maintained by purging the cell solutions with a mixture of seaford grade N2 C 1978 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
Table I. Cyclic Voltammetry of Copper in Aqueous Carbonate Solutions cathodic alkalinity 10-4N
pH
6.0 7.0
10-3N
6.0 7.5
volts/s 0.040 0.100 0.160 0.040 0.100 0.160 0.040 0.100 0.160 0.040
u,
0.100
10-2N
7.0
0.160 0.040 0.100
7.5
0.160 0.040 0.100
0.160
i,, nA 123 175 199 90 180 222 120 189 225 109 206 232 181
386 469 123 237 251
E,, volts vs. SCE t 0.008
0.005 - 0.019 - 0.015 - 0.033 - 0.039 + 0.006 -
- 0.012 - 0.019 - 0.055 - 0.072
0.085 - 0.036 -- 0.053 - - 0.059 - 0.062 - 0.076 - 0.083 -
E, ?E:
37 45 51 44 43 47 36 46 49 36 38 36 29 36 36 40 35 34
(National Welders Supply Co.) and 5% C02/95% N2 (Matheson Gas Products) with flow rates regulated with Lab-Crest Century Series 100 flowmeters (Fisher and Porter). All solutions were prepared with ultra-high purity deionized water supplied by I.B.M. Laboratories, Research Triangle Park, K.C. The electrochemical cell, constructed of Vycor (Corning Glass Works), contained a hanging mercury drop working electrode (HMDE), a saturated calomel reference electrode (SCE),and a platinum wire counter electrode. The reference electrode was electrically connected to the cell solution by a saturated potassium nitrate liquid junction and ultimately with a Luggin capillary. Mercury drops were delivered for each individual analysis from a micrometer syringe working electrode (Metrohm). Electrode volume was 9.3 X cm3. A micro-combination pH electrode (Sargent-Welch Co.) and a Fisher-Accument Model 520 pH/Ion meter provided continuous pH monitoring. The electronic apparatus consisted of a Wenking potentiostat (G.Bank Electronik), an Exact Model 505A function generator, a Hewlett-Packard 3440A digital voltmeter and a Hewlett-Packard i004B x-y recorder. Copper standard was delivered from Eppendorf pipets (Brinkmann Instruments, Inc.) to solutions of and lo-' N alkalinity at pH 6.0 and 7.0 and ASV performed after each addition. Accumulation of copper was made at 4.700 V vs. SCE for 5 min while the solution was stirred at 600 rpm with a Sargent synchronous rotator. Scan rate for stripping was 40 mV/s. Cyclic voltammograms were performed on Cu solutions of 7.7 X M and at scan rates of 40, 100, and 160 mV/s.
anodic
-
ip/v'l 616 566 497 450 580 556 600 598 562 545 652 580 905 1220 1172 615 749 628
2105
i,, nA 121 161 185 67 123 140 130 190 190 80 118
114 162 208 249 128 175 208
E, E;$,
E,, volts vs. SCE
0.067 0.081 0.087 0.052 0.069 0.083 0.058 0.071 0.081 0.010 0.023 0.028 0.013 0.018 0.029 -0.014 -0.008
36 44 47 35 46 57 37 38 37 47 52 46 31 35 52 33 40 50
-0.001
AEp,
mV 59 87 107 66 102 122 51
83 100 66 95 113 50 70 89 48 68 84
i,li,
0.98 0.92 0.93 0.76 0.68 0.63 1.06 1.00 0.84 0.74 0.57 0.49 0.89 0.54 0.53 1.04 0.74 0.83
__ _. i.
Figure 1. Cyclic voltammograms for pH 6 and line is 40 m V / s , solid line IS 160 m V / s
N alkalinity Dotted
RESULTS Cyclic voltammetry was performed on solutions of IO-', and lo-' N carbonate alkalinity; pH 6.0, 7.0, and 7.5; and a t voltage scan rates, u, of 40, 100, and 160 mV/s. Some of these voltammograms appear in Figures 1-3 and current-potential data are summarized in Table I. Voltammograms for Figures 1-3 were chosen to illustrate three features of the reduction mechanism; nonreversibility, a preceding chemical reaction, and adsorption. A trend among data with pH. alkalinity, or scan rate in any one figure will not necessarily be found throughout the entire series of voltammograms. General trends throughout the series are obscured by complexity of the reduction mechanism, especially by adsorption a t high alkalinity, pH, and scan rate. Two voltammograms of a solution with alkalinity, N, and p H 6.0 a t two different scan rates, 40 and 160 mV/s are shown in Figure 1. T h e current axis in this figure and in Figures 2 and 3 is normalized by dividing by u112. As the scan rate increased, both cathodic and anodic peak currents decreased in magnitude (the anodic peak current was measured
-~ irr
Figure 2. Cyclic voltammograms for 100 r V / s and Dotted line IS pH 6 and solid line is pH 7.5
N alkalinity.
from cathodic current decay) and the cathodic and anodic waves broadened as judged by E , - E p I 2 .In addition, the
2106
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
Figure 3. Cyclic voltammograms for 160 mV/s and pH 7 5 Dotted line is N alkalinity and solid line is 10.’ N alkalinity
Y V
k’
f-
-
~
~
..
.-=1
. L
-
C~
.
.
~
~ ~ ~-
~ _
12 ~
~
~
~
Figure 4. Copper stripping current vs. copper concentration at lo-’ N alkalinity. pH 6 (0,W) and p H 7 ( 0 ) cathodic peak shifted to higher negative potentials, and 1E,, the difference between anodic and cathodic peak potentials, increased with increasing scan rate. When p H was increased N alkalinity at a constant scan rate of frurii 6.0 to 7.5 at 100 mV/s, the entire voltammogram shifted to higher cathodic potentials, the cathodic portion became narrower and symmetrical, the anodic portion reduced in height and broadened and the ratio i J i Cdecreased (Figure 2 ) . When alkalinity was raised from 10 N to N a t p H 7.5 and a single scan rate of‘ 160 mV/s, the cathodic peak shifted slightly positive, the ilrlodic height increased, and 1l3,decreased (Figure 3). Anodic stripping voltammetry was carried out on a series of solutions to mimic a complexonietric titration or standard addition procedure, and peak stripping currents vs. Cu concentration are plotted in Figure 4. The curves for N alkalinity, p H 6.0 and p H 7.0 were both linear up to about M added copper and the lower p H had approximately 25% greater slope. Identical experiments a t N showed siniilar results except slopes were smaller than those a t N. Above M Cu data points were curvilinear indicating progressive loss of sensitivity at higher copper concentrations. M copper and With the addition of approximately 2 x a t p H 7.0, a light blue precipitate was observed. Loss of sensitivity at high concentrations may be due to formation of nunreducible colloidal or solid Cu as suggested by Ernst et al. ( 3 ) .
DISCUSSION Cyciic voltamniograms suggest Cu reduction is nonreversible and becomes less reversible with increasing pH and decreasing alkalinity. There are three criteria for reversibility: E , - E,/z =. 56/n mV, i J i L = 1.0, and S, = 5 8 / n mV, independent of s c m rate (20). None of these criteria are met and the di-
agnostics all change with scan rate, alkalinity, and pH. If the reduction is considered merely from the formalism of a “quasi-reversible’’reaction, a heterogeneous rate constant on the order of cm/s is calculated from AEp using Table I of reference 2 1 and Do = 1.02 X cm2/s. There is also evidence for adsorption. In Figure 2, the cathodic portion of the cyclic voltammogram shows increased symmetry and narrowing, and a marked decrease in i a / i c suggesting , weak adsorption of electroactive Cu(I1) on the electrode surface prior to electron transfer (22). Reference 22 characterizes the effects of weak adsorption with two theoretical parameters, one of which Po, the adsorption parameter, increases with the free energy of adsorption and surface concentration of reactant. Increasing scan rate also increases Po. Increasing Po causes narrowing of the cathodic wave and decrease in iJic. If the criteria of wave narrowing and peak ratios are used, adsorption is judged to occur with solutions N alkalinity, pH 7.5, and lo-* N alkalinity, pH 7.0 and 7.5. Evidently, pH and alkalinity, not just scan rate, affect Po. Either an adsorbable reactant is formed in these solutions and not in solutions of lower p H and alkalinity, or the free energy of adsorption or surface concentration of adsorbed species is increased. As discussed below, overall reversibility also changes among these same solutions, making precise judgments about adsorption from these criteria difficult. Nonreversibility comes from iR distortion of the voltammogram ( 2 4 2 4 , a slow electron transfer reaction (211, or a slow solution reaction preceding the electron transfer reaction (20). Uncompensated iR drop was dismissed on the basis of several observations, but primarily because no change in the cyclic voltammograms was observed when NaC10, was added to increase ionic strength and because similar solutions measured previously in t h e cell had less than 200 Q uncompensated resistance. Two criteria for a reaction without adsorption preceding a reversible or a totally irreversible electron transfer are decreasing normalized cathodic current and broadening cathodic wave with increasing scan rate (20). The copper reduction mechanism in carbonate is not that simple but involves adsorption and a slow but not totally irreversible electron transfer; however, the two solutions at p H 6 fit these criteria and so does p H 7, N alkalinity for 100 and 160 mV/s, although the small ia/Lcsuggests strong influence of adsorption in this solution. At higher alkalinity and pH, adsorption obscures these diagnostics. No theory is available for this complex mechanism; therefore, l E p suggested by Nicholson for measuring reversibility of quasireversible reactions (21) was used to measure overall reversibility. Reversibility of Cu reduction increased with alkalinity and this could be due t o its influence on a chemical reaction
CUCOS
kf = Cu?+ + CO32-
(1)
kb
or k
Cu(OH)+
Cu2+
+ OH
kb
preceding the electron transfer reaction
where k f and k b are homogeneous rate constants and k, and a have their usual meaning. First of all, alkalinity can influence these reactions through buffer capacity. Reaction 1 creates a local increase in carbonate or hydroxide ion near the electrode during Cu(I1) reduction and failure to buffer sufficiently can lead to distortion of current-voltage curves (25). In the case of a conjugate base such as carbonate, total buffer concentration 10-20 times in excess of the metal is recommended, provided the acid-base equilibrium is rapid (26). For
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
.
> E -801
A
A
I
A
A
-1201
I40
2107
A
1 I
I
1
-4 5
-3 5
-2 5
-3 0
P Figure 5. LEpof cyclic voltammograms plotted as a function of buffer capacity, 8. 40 mV/s (O), 100 m V / s (W), 160 m V / s (A)
-2 0 l o g Pco*
log
alkalinities used here which imitate the range of values normally encountered in natural waters, these requirements of excess carbonate are met a t N and lo-* N alkalinity, are barely met a t N alkalinity and p H 6.0, and are not N alkalinity and p H 7.0 where the buffer capacity met at is about 40% lower. Copper reduction was least reversible in this solution. Alkalinity may influence reversibility in another way. In the p H range used in these experiments, the rate of neutralization is slow and limited by the hydration of aqueous C 0 2 (27,28). This slow hydration can affect dissociation of CuC03 prior to Cu(I1) reduction. The rate law for hydration of C02(aq) by the reaction COz + H20 c H2C03is given by the pseudo-first-order expression
(3) where kcOais 0.03 s-l a t 25 "C (28). In terms of the experiments in this study where p H was maintained with C 0 2 partial pressure, an increase in alkalinity at any given p H produces an increase in the rate of CO,(aq) hydration. It is expected then, that the rate of the dissociation step in Reaction 1 will increase with increasing alkalinity a t any given p H and decrease with increasing p H a t a given alkalinity, merely from this kinetic standpoint alone, aside from considerations of buffer capacity. T h e degree of reversibility as indicated by S, was plotted vs. the buffer capacity and Pco2 in Figures 5 and 6. These figures show increasing reversibility with both increasing Pco, and increasing buffer capacity for all scan rates. T h e cathodic peak potential a t the slowest scan rate (to minimize the effect of adsorption) was plotted vs. calculated carbonate concentration in Figure 7 . A general negative shift was observed with increasing carbonate and pH. The data are scattered and if peak potentials are compared with a -30-mV shift per decade increase in carbonate concentration, formation of CuC03 might be inferred; however, a t constant p H this shift is only 2-20 mV/decade suggesting simultaneous formation of Cu(OH)+. If the same data are plotted vs. p H and compared to a 30 mV/decade slope, formation of Cu(OH)+is implied. Again, data are scattered. It is very difficult to determine for any one solution and scan rate the step in the overall mechanism which has greatest influence on the location of the peak potential; therefore, for this and other nonreversible reduction mechanisms, plotting voltammetric peak potential shifts is not a reliable method for species identification. Shifts in stripping peak potential were also negative with increasing carbonate concentration, greater shifts observed a t p H 7.0 than p H 6.0, and were also dependent on copper
-I
0
Figure 6. A€p of cyclic voltammograms plotted as a function of C 0 2 partial pressure. 40 mV/s ( O ) ,100 mV/s (U), 160 mV/s (A)
I I
I
\ \ A
I
\
," -02r
\
\ -04
i
\
t
1
A
\ \
I
I
\
-80
-70
-60
log
-50
[co;--l
Figure 7. Copper peak potential, volts vs. SCE, of cyclic voltammograms at 40 mV/s. pH 6 ( O ) ,pH 7 (A),pH 7.5 (W). Dotted line is -30 mV/decade
concentration, shifting positive with increasing copper concentrations. These shifts were not interpretable in any simple way. Cyclic voltammetry indicates ASV pre-electrolysis is rate controlled and the proposed mechanism suggests ASV potentials are a function of complex adsorption, amalgam concentration, and the rate of copper complex formation. The overall reduction mechanism of copper from dilute carbonate solutions includes prior dissociation, slow electron transfer, and reactant adsorption. Under conditions of low alkalinity and high pH, the rate of reduction may be affected by buffer capacity or by the rate of C 0 2 hydration. Further research is necessary to establish definitively the mechanism, but these speculations are consistant with variations of the cyclic voltammograms with solution parameters. This mechanism could also explain the observed decrease in ASV sensitivity a t low copper concentrations with increasing p H and decreasing alkalinity since copper accumulation during the ASV pre-electrolysis step would be rate controlled, decreasing with increasing p H and decreasing alkalinity. Of course, ASV sensitivity may also be affected by adsorption or by changes in the diffusion coefficient of the reducible species. The diffusion coefficient would need to change 50-60% to account for the observed change in sensitivity. Adsorption did not appear to block reduction.
C0NCLUSI 0NS Further work is indicated to establish the mechanism conclusively, but copper reduction from dilute carbonate solutions appears influenced by dissociation of Cu(OH)+ or CuCO, prior to electron transfer. The reduction also involves adsorption and a slow electron transfer. Variations of car-
2108
ANALYTICAL CHEMISTRY, VOL. 50, NO. 14, DECEMBER 1978
bonate alkalinity and p H of the solution affect the reversibility of the reduction, possibly through variations in buffer capacity and rate of C 0 2 hydration. Sensitivity of ASV in these solutions may be related to this mechanism. I n light of these results, it is recommended that shifts in stripping peak potentials or changes in stripping current magnitude with gross changes in solution composition of p H and alkalinity not be used to indicate changes in metal speciation. Reversibility of the overall electrode reaction can affect these measurements in a manner that is complex and uninterpretable unless the reaction mechanism is known. ASV theory of complex mechanisms needs further development and mechanisms of metal reduction in carbonate solutions need further investigation before shifts in potential can be used with certainty. Anodic stripping used with an amperometric titration for studying metal complex formation is preferred. Any change of reversibility that affects ASV sensitivity is directly observable in this procedure and reflected in the data subsequent to t h e titration end point (11).
LITERATURE CITED (1) W. R. Matson, Ph.D. Thesis, Massachusetts Institute of Technology, Cambridge, Mass., 1968. (2) W. L. Bradford, Limnol. Oceanogr., 18, 757 (1973). (3) R. Ernst, H. E. Allen, and K. H. Mancy, Water Res., 9, 969 (1975). (4) T. A. O'Shea and K. H. Mancy, Anal. Chem., 48, 1603 (1976). (5) H. E. Allen, Ph.D. Thesis, University of Michigan, Ann Arbor. Mich., 1974. (6) A. Zirino and M. L. Healy, fnviron. Sci. Technol., 6, 243 (1972). (7) A. Piro, M. Bernhard, M. Branica, and M. Verzi, in "Radioactive Contamination of the Marine Environment", IAEA, Vienna, 1973. (8) R. W. Baier, J . fnviron. Qual., 6, 205 (1977).
(9) Y. K. Chau, R. Gachter, and K. Lum-Shue-Chan, J . Fish. Res. Board Can.. 31, 1515 (1974). (10) Y. K. Chan and P. T. S. Wong, in "Workshop on Toxicity to Biota of Metal Forms in Natural Water". International Joint Commission, Windsor, Ont.. 1976, p 187. (11) M. S. Shuman and G. P. Woodward, Jr., Environ. Scl. Technol.. 11, 809 (1977). (12) J. C. Duinker and C. J. M. Kramer, Mar. Chem.. 5, 207 (1977). (13) M. S. Shuman and G. P. Woodward, Jr., Anal. Chem.. 45, 2032 (1973). (14) F. Morel, R. E. McDuff, and J. J. Morgan in "Trace Metals and MetalOrganics Interactions in Natural Waters", Ann Arbor Science Publishers, Ann Arbor, Mich., 1974, p 157. (15) A . Zirino and S. Yamamoto, Llmnol. Oceanogr., 16, 779 (1972). (16) J. D. Hem and W. H.Durum, J. Am. Water Works Assoc., 65, 562 (1973). (17) D. T. Long and E. E. Angino, Geochim. Cosmochim. Acta, 41, 1183 ( 1977). (18) J. Vuceta, P h D Thesis, California Institute of Technology, Pasadena, Calif., 1976. (19) L. Meites, J . Am. Chem. Soc., 72, 184 (1950). (20) R . S. Nicholson and I.Shain, Anal. Chem., 36, 706 (1964). (21) R. S. Nicholson, Anal. Chem.. 37, 1351 (1965). (22) R. H. Wopschail and I.Shain, Anal. Chem., 39, 1514 (1967). (23) P. Dehhey, "New InstrumentalMethods in Electrochemisby", Interscience, New York, 1954. (24) R. S. Nicholson, Anal. Chem., 37, 667 (1965). (25) A. A. Vicek, in "Progress in Inorganic Chemistry", F. A. Cotten, Ed., Interscience, New York, 1963, p 211. (26) L. Meites, "Polarographic Techniques", Wiley-Interscience, New York, 1963. (27) W. Stumm and J. J. Morgan, "Aquatic Chemistry", Interscience, New York, 1970. (28) D. M. Kern, J , Chem. fduc., 37, 14 (1960).
RECEIVEDfor review June 1,1976. Resubmitted May 11,1978. Accepted October 5 , 1978. Work supported by funds from the Oceanographic Section, National Science Foundation, NSF Grant OCE73-21045
Measurement of Organomercury Species in Biological Samples by Liquid Chromatography with Differential Pulse Electrochemical Detection W. A. MacCrehan" and R. A. Dursl Center for Analytical Chemistry, National Bureau of Standards, Washington, D.C. 20234
A new measurement approach for organomercury cations is described employing liquid chromatography with electrochemical detection. Special considerations in constructing apparatus to optimize reductive electrochemical measurement are outlined. The added selectively of the differential pulse mode of detection is demonstrated. A charge-neutralization reversephase separation of methyl-, ethyl-, and phenylmercury cations has been developed. The limit of detection for methylmercury is 2 ng/g or 40 pg (190 fmol/20 1 L sample). Methylmercury is determined by the technique in two research materials-tuna fish and shark meat.
Many metal and metalloid elements can be converted by biological processes into organometals (I, 2). Also the alkyl and aryl compounds of Hg, Sn, P b , and As are finding increasing use in industrial manufacturing and pest control (2). In order to investigate fully the role of toxic metals in environmental and biological systems, it is necessary to measure the exact chemical form of the element. T h e measurement of organometals in complex "real world'' samples presents two difficult analytical requirements: high selectivity and very low detection limits. Perhaps the most
frequently used approach for the analysis of the cationic organometals is derivatization with a halogen and subsequent gas chromatography (GC) with electron capture detection (3, 4 ) . Detection limits with this approach are quite good, 2 ng/g for CH,HgBr, for example (3). However, the relatively poor selectively of this detection approach requires rather extensive sample "cleanup" before analysis ( 4 ) . Recently, the selectivity problem has been overcome by the use of element-specific detection such as microwave cavity emission detection ( 5 , 6). However, the gas chromatographic separation requires thermally stable, strong complexes of the cationic organometals to be made by derivatization before analysis. An alternate approach for the measurement of organometal species is to use liquid chromatographic (LC) separation coupled with selective detection. Graphite furnace atomic absorption (GFAA) has been used as a detection system with an automatic, periodic sampling and atomization of the chromatographic eluent. With the autosampled mode, the eluent can be sampled, desolvated, and atomized only every 50 s ( 7 ) . This sampling period is much too long to preserve the integrity of a high efficiency separation. T h e detection limits for the GFAA approach for LC detection are relatively poor (7). This reflects the difficulties in sampling, the poor atomization efficiency of volatile organometals, and the
This article not subject to U.S. Copyright. Published 1978 by the American Chemical Society
