Environ. Sci. Technol. 1991, 25,1716-1722
Copper( I ) in Fogwater: Determination and Interactions with Sulfite Hanbin
Marla de Lurdes S. GonGaives,$ Max Reutlinger,t Laura Sigg, *,t and Werner Stummt
Institute for Water Resources and Water Pollution Control (EAWAG), Swiss Federal Institute of Technology, CH-8600 Dubendorf, Switzerland, and Centro de Quimica Estrutural, Complexo I , Instituto Superior Tecnico, 1096 Lisboa Codex, Portugal
H The copper(I)/(II) redox system was examined in fogwater with respect to the occurrence of Cu(I), the role of sulfite as a reductant of Cu(I1) and as a complexing ligand, and the speciation of Cu(1) and Cu(I1). Copper(1) was measured in fogwater by the bathocuproine method, which was evaluated for the conditions typically encountered in atmospheric water droplets. Concentrations of Cu(1) in the range 0.1-1 pM were found, which represented between 4 and >90% of the total copper in these samples. In experiments using concentration ranges of copper and S(IV) close to that in fogwater, the reduction of copper(I1) to copper(1) by sulfite was shown to be pH-dependent and to occur rapidly at pH >6. Calculations of the equilibrium complexation of Cu(1) and Cu(I1) under fogwater conditions show that complexes of Cu(1) with sulfite predominate, while for Cu(I1) oxalato complexes are important. Sulfite plays an important role as a ligand for Cu(1) in fogwater; Cu(1) may be produced by various reduction reactions, e.g., by organic compounds, and appears to be oxidized only slowly in the presence of S(1V).
Introduction Atmospheric water droplets (fog- and cloudwater) represent a reactive medium in which various chemical reactions may take place. Although liquid water represents only a very small part of the atmosphere, the processes in the aqueous phase may affect significantly the composition of the atmosphere (1-3). Transition metals such as Fe, Mn, and Cu are often present in significant concentrations in atmospheric water droplets (4-6). These elements are involved in redox and radical chain reactions (7-10). Especially the oxidation of S(1V) to S(V1) has been shown to be catalyzed by Fe and Mn (9,10). Cu(I)/(II) is involved in reactions with OH', Hz02,and H02' (1, 7,8). All these reactions are strongly dependent on the speciation of the metal ions. Redox couples such as Fe(II)/Fe(III), Cu(I)/Cu(II), and Mn(II)/Mn(III,IV) are thus of primary interest; in order to evaluate the role of these elements in different reactions, it is desirable to obtain information about their actual speciation (complexation and oxidation states). Recent work in this laboratory has shown the occurrence of substantial concentrations of Fe(I1) in fogwater (6), which were attributed to reduction of Fe(II1) by sulfite and under light conditions by radicals and by photochemical reactions (11). The copper(I)/(II) system has been studied under seawater conditions by Moffett and Zika (12-15) with respect to the possible oxidation and reduction reactions. They showed the occurrence of copper(1) in surface seawater (15). The oxidation reactions of Cu(1) in seawater medium were examined by Sharma and Miller0 (16,17). Formation of copper(I1)-sulfur(1V) complexes and the redox chemistry of these complexes, by which Cu(1) and sulfate are formed, have been reported by Conklin and Hoffmann (18, 19). Although numerous measurements of total or dissolved Cu in fog- and rainwater may be found in the literature 'Institute for Water Resources and Water Pollution Control. 3 Centro de Quimica Estrutural. 1716
Environ. Sci. Technol., Vol. 25, No. 10, 1991
( 4 , 5 ,20), little information about the speciation of Cu or about the occurrence of the oxidation state Cu(1) may be found. In this work, we consider different complementary aspects of the Cu(I)/(II) system in fogwater. The analytical problems involved in the determination of Cu(1) by the spectrophotometric method with bathocuproine (14, 21) are carefully evaluated for fogwater conditions, and concentrations of Cu(1) and their ratios to total copper are determined analytically in fogwater samples. In order to obtain insights about the predominant chemical reactions, thermodynamic and kinetic considerations are needed. The speciation of Cu(1) and Cu(I1) is thus evaluated for the composition of fogwater. Among the possible reduction reactions of Cu(I1) to Cu(I), the reduction by sulfite seems of special interest, since fogwater often contains significant concentrations of sulfite. We attempt thus to evaluate the role of aqueous S(1V) in the Cu chemistry, where it may act both as a ligand and as a reductant.
Experimental Section Fogwater Sampling. Fogwaters were collected at a sampling site located in Dubendorf, a surburb of Zurich (Switzerland), during radiation fogs, between October 1989 and January 1990. Radiation fogs usually form during nighttime and dissipate several hours after sunrise. Collection of complete fog events, from the time of formation to their dissipation, was attempted. Individual samples were usually collected for 1-2 h (Table I). The sampling device was a screen collector (22) and the techniques used were the same as described in ref 4. Liquid water contents, which were evaluated from the volume collected, ranged from 0.02 to 0.19 mL/m3. Temperature was in the range 0-5 "C. The samples were protected from light during collection by aluminum foil around the bottles. Immediately after collection, the samples were filtered through 0.45-pm membrane filters (Millipore) and divided into several aliquots for analytical determinations, and the spectrophotometric determinations were carried out without delay. Determination of Cu(1). Cu(1) was determined both in kinetic experiments with sulfite and in fogwater samples by use of the spectrophotometric method with bathocuproinedisulfonic acid (21). Absorbances of all samples were measured at 484 nm with a Kontron Uvikon 860 spectrophotometer against a blank in a 5- or 10-cm cell. Stock solutions of bathocuproinedisulfonic acid sodium salt were freshly prepared each week (1g/L, 1.77 X M). The concentration of bathocuproine in the measuring solution was 3.5 x M. Copper(1) chloride was dissolved in N,-deoxygenated 0.005 M HC1 and 1 M NaCl solution, which was continuously purged with suprapure nitrogen to eliminate oxygen; traces of oxygen were probably still present in these solutions. The cuprous stock solution M) was prepared within 4 h of use. Cu(1) standard series from the stock were freshly prepared and buffered with M citrate prior to each experiment. Linear calibration ranged over 5 X 10-8-2 X 10" M. The M. The standard detection limit for Cu(1) was 5 X deviation for five replicate calibration curves in the range
0013-936X/91/0925-1716$02.50/0
0 1991 American Chemical
Society
Table I.Concentrations of Cu(1) and of Other Components in Fogwater ratio of date
sampling timea
18 Oct 89
2:13-3:45 (D) 3:45-5:OO (D) 5:OO-6:15 (D) 6:15-7:30 (D) 7:30-8:45 (L) 8:45-1O:OO (L) 6:25-7:30 (D) 7:30-8:30 (L) 8:30-9:30 (L) 230-3:30 (D) 3:30-4:30 (D) 4:30-5:45 (D) 5:45-7:25 (D) 6:50-830 (L) 6:59-1O:OO (L) 23:45-1:45 (D) 1:45-3:45 (D) 3:45-5:45 (D) 5:45-7:45 (D) 7:45-9:45 (L) 9:45-11:25 (L) 2:40-3:40 (D) 3:40-5:45 (D) 5:45-7:45 (D)
27 Oct 89 28 Oct 89
10 Nov 21 Nov 2 1 Nov 22 Nov
89 89 89 89
22 Jan 90
a
D and L stand for
sample no.
Cu(I), lo-' M
total Cu
1 2 3 4 5 6
1.07 1.24 1.07 1.50 1.57 48.80 0.28 5.00 8.01 6.72 7.30 4.50 4.83 5.79 10.10 14.80 10.50 2.88 10.60 5.48 0.21 6.44 3.57 3.37
1
2 3 1 2
3 4 1 1 1 2
3 4 5 6 1
2 3
10-7 M
CU(1) to tot. c u
PH
X:[S(IV)l, 10-4 M
2.80 2.12 1.27 3.73 4.20 135.00 3.02 6.92 15.70 11.80 8.34 5.19 4.60 8.66 13.30 17.10 10.90 6.67 14.10 5.43 5.21 9.93 4.99 5.67
0.38 0.57 0.84 0.40 0.37 0.36 0.09 0.72 0.51 0.57 0.88 0.87 (1.00) 0.67 0.76 0.87 0.97 0.43 0.75 1.00 0.04 0.65 0.72 0.59
4.93 4.96 5.06 5.16 5.70 3.21 5.46 5.12 5.36 4.89 3.03 3.03 3.03 5.84 5.56 5.84 6.06 5.35 3.24 4.59 6.34 2.54 3.31 4.07
0.75 0.66 0.87 1.34 3.06 3.21 0.75 1.72 2.62 1.09 0.34 0.44 0.22 2.31 8.45 15.85 7.02 2.87 1.90 2.06 4.30 0.34 0.41 0.62
formaldehyde,
DOC,
10-5 M
mg/L
3.6 1.6 1.0 1.7 9.6 4.0
9.9 9.0 11.4 11.0
20.8 42.8
69.6 76.0 63.0 64.0 6.5 20.4 40.3 12.9 6.3 7.3 7.4 8.3
55.5 97.8 44.0 20.8 22.8 24.0 31.5
[S(VI)I, lo4 M 0.50 0.44 0.63 0.75 1.63 3.36 0.81 1.31 2.69 1.75 1.25 1.13 1.06 2.56 12.60 11-10 5.75 2.87 2.59 2.53 4.06 0.81 1.06 1.22
samdine conditions in dark or light. remectivelv.
5 X 1OW8-1.5X lo4 M was 3%. We also determined that the same values of absorbances were obtained either with copper(1) or with copper(I1) solutions in the presence of an excess of hydroxylamine hydrochloride as a reducing agent. The interferences of iron(I1) and sulfite in the copper(1) M SOS2-and 2 determination were examined in 1 X X lo4 M Fe(I1) solutions. No detectable interference could be measured. Interference by Fe(I1) and other common anions (such as nitrate or sulfate) have been ruled out (14). The bathocuproine method had to be optimized for fogwater conditions and potential interferences (mostly the ongoing reduction of Cu(I1) after addition of the reagents) had to be checked. The reduction of Cu(I1) by sulfite is slow under the conditions used for the analytical determination (pH (6, in presence of oxygen). In the presence of both sulfite and bathocuproine, however, solutions of Cu(I1) give an increasing absorbance with time. Moffett et al. (14) suggested the use of EDTA, NTA, or ethylenediamine in order to complex copper(I1) and to avoid further reduction reactions during the determination of Cu(1). In the presence of sulfite, the addition of these ligands did not, improve the results and they were not used subsequently. Different contact times of copper(I), copper(II), sulfite, bathocuproine, and the complexing ligands for copper(I1) were checked. The best agreement between measured and added Cu(1) concentrations was obtained by measuring the absorbance immediately after bathocuproine had been added, in the absence of additional ligands (Figure 1). Filtered fogwaters generally exhibit a yellow color, which interferes with the absorbance measurement at the 484-nm wavelength. The background absorbance was found to be significant in comparison to the total absorbance. Thus, the blank absorbance was subtracted from the total absorbance after addition of bathocuproine. The following procedure has been adopted for the determination of copper(1) in fogwater. The sample is filtered and quantitatively transferred to a 10-cm spectrophotometric cell; the blank absorbance of the sample is
1 I.o/ -
0.8
5 A
v
3
0
0.6
U
?!
a cn m
0.4
E 0.2
0.0
0.0
0.2
0.4
0.6
added Cu(l) pM Figure 1. Spectrophotometric measurements of Cu(1) by different procedures in a series of mixed Cu(1) and Cu(I1) solutions. Cu(1) is 55 % of total Cu at each concentration; pH of test solutions, 5-6; total M; bathocuproine concentration, 3.5 X M; sulfite, 1 X ethylenediamine (en), 2 X M (where present). Key: (-) theoretical 1:l curve for added Cu(1). (+) Total Cu (Cu(1) Cu(I1)) measured in the presence of an excess of hydroxylamine hydrochloride. (e)absorbance measured immediately after mixing sulfite and bathocuproine; bathocuproine added directly to the spectrophotometric cell containing Cu(I), Cu(II), and sulfite. The points fit well to the curve, which indicates that there was no obvious interference of Cu(I1) in the presence of sulfite for this fast measurement of Cu(1). (0)30 min after mixing sulfite and bathocuproine. The results were overestimated in comparison to the added Cu(1) and approached the total copper (copper(1) plus copper(I1)) concentration. (0)30 min after mixing sulfite, bathocuproine, and en. The results were underestimated or overestimated, depending on the added concentration of Cu(1). (B) immediately after mixing sulfite, bathocuproine, and en. The results were underestimated.
+
measured for 3 min with the kinetic mode of the instrument (one record each 6 s). Bathocuproine (3.5 X M) Environ. Sci. Technol., Voi. 25, No. 10, 1991
1717
is then added to the cell, and the variation of the total absorbance is again measured for 3 min. If the absorbance values are stable, average values are determined, otherwise the absorbance is extrapolated to time zero. For most samples the absorbances obtained for fogwater were stable within an error of 2% for a 3-min kinetic measurement. Only for a few samples did the absorbance values increase with time, indicating that the reduction of copper(I1) was still occurring. Other Analytical Methods. pH values were measured on a NBS scale with buffers for the kinetic experiments. In fogwaters, pH was calibrated by Gran plots on a concentration scale and measured at the temperature of fogwater or at 5 "C. Total Cu in fogwater was determined by graphite furnace AAS after appropriate dilution of the acidified Samples. Sulfite concentrations in the kinetic experiments were checked iodometrically. Sulfite in fogwater was determined by ion chromatography together with chloride, sulfate, and nitrate as described previously ( 4 ) . Dissolved organic carbon (DOC) was measured on samples filtered through washed 0.45-pm filters (Millipore) on a Dohrmann DC 80 instrument. Formaldehyde was determined by fluorometry according to ref 23. Kinetic Experiments. The kinetic studies of reduction of Cu(I1) by sulfite ions were carried out in a stirred 500-mL glass reaction vessel at 20 "C and I = 0.10 M (KN03) continuously purged with suprapure nitrogen and in the absence of light. Stock solutions of Cu(1) M) were prepared as described above. Na2S03was added as a solid directly to the reaction vessel and dissolved immediately. Cu(I1) was added to the sulfite solution, so that losses to the glass walls should not be significant. Progress of the reaction was monitored by measuring Cu(1) spectrophotometrically with bathocuproinedisulfonic acid (21) in aliquots. Cu(1) was determined immediately after the aliquots were removed from the reaction solution; in order to eliminate the influence of the continuing reaction, the absorbance values were read each minute for 6 min and extrapolated to time zero. Cu(I1) was found not to interfere in the concentration range used, since its addition did not affect the absorbance value determined. The stability of Cu(1) solutions in the presence of sulfite was monitored by the following experiment: Solutions of copper(1) (1 X lo4 M) with an excess of sulfite (2 X M) were prepared at different pH values and kept at room temperature in the dark; the concentration of copper(1) after different storage times was measured by the bathocuproine method. Similar experiments were performed in the presence of 0.01 and 1M chloride and in the absence of complexing ligands.
Results and Discussion Cu(1) in Fogwater Samples. A summary of the results obtained for Cu(1) and total copper in fog samples is presented in Table I. The concentrations of sulfite (total S(1V) in solution), DOC, and formaldehyde and the pH are also shown, since these parameters are relevant to the ratios Cu(I)/Cu total. The total Cu concentrations, as well as the pH range and composition of these fogwaters, are similar to previously obtained results at the same site ( 4 , 6). The ratio of Cu(1) to total dissolved copper is in most samples surprisingly high and varies between 4% and more than 90%. A large number of these samples were collected at night in the dark. No obvious correlation between Cu(1) and parameters such as pH, DOC, and sulfite is found in these samples, S(1V) is present in all samples at concen1718
Environ. Sci. Technol., Voi. 25, No. 10, 1991
I
f
I,
1 .o
0.8 r
0
0.6
"'1e 0.4
'-..a
,L]
:....~,.".-".-I._._.__.
0.0
0
10
20
30
40
Time (h) Figure 2. Variation of concentration of measured Cu(1) with time of storage in the presence of different ligands. The initial concentration of Cu(1) was 1 pM; pH in the solution was adjusted with 0.1 M HNO, or NaOH. The solid black symbols represent the resuits measured at pH 3, 5, or 7 in the absence of ligands, where the concentration of Cu(1) decreases rapidly due to oxidation of Cu(1). Key: (0)in the M sulfite at pH 3. (0)in the presence of 2 X presence of 2 X M sulfite at M sulfite at pH 5. (8)in the presence of 2 X pH 7. (A)in the presence of 1 M chloride at pH 3. (X) in the presence of 1 X lo-' M chloride at pH 7. I n the presence of sulfite at pH 3-7, 90% of the initially added Cu(1) is still found after 22 h. A high concentration of chloride ions also stabilizes Cu(1) to some extent, though less efficiently than sulfite.
trations of 10-4-10-3 M; formaldehyde concentrations are always lower than the total S(1V) concentrations. Sulfate (S(V1)) is present at concentration levels similar to total S(1V). DOC is present in large excess of Cu. Fog samples with pH =3 are at this sampling site most often related to inputs of hydrochloric acid; these samples contain (1-3) X M chloride. Role of Sulfite in Preserving Cu(1) from Oxidation. The concentration of copper(1) after storage of some of these fogwater samples was followed. In samples stored for 20 days (in the dark a t 4 "C), only slight changes in Cu(1) concentrations were found (either increase or decrease). The concentration of sulfite in these samples also appeared to be stable upon storage and was in large excess with respect to total Cu. In order to simulate this effect, the stability of Cu(1) in solutions containing sulfite was examined. Sulfite forms more stable complexes with Cu(1) than with Cu(I1) (18, 25). The variation of Cu(1) concentrations with time in solutions containing sulfite is shown in Figure 2. In the presence of sulfite, for pH between 3 and 7, only small changes in copper(1) (3, CuS0,- absolutely predominates; the concentration of free cuprous ion is some orders of magnitude lower than that of CuSO,-. (b) total ligands in the system included oxalate, sulfite, and chloride; [C,O,] = 5 X M, [SO,] = 4 X M, [CI] = 1 X lo-, M, [Cu(II)] = 1 X lo-' M. For pH >4, CuC20, is predominant; however, the ratio of free cupric ion to total Cu(I1) is much higher than that of free cuprous ion to total Cu(1).
where [Cu(I)] is the total concentration of Cu(1); [Cu(II)] is the total concentration of Cu(I1); al = [Cu+]([Cu(I)])-l; a2 = [CU~+]([CU(II)])-~; Pl'so and are the stability and constants of Cu(I)SO< and du(I)(SO,)~-;&IcI, &Ic1, 0 2 are ~ ~the stability constants of CuCl, CuCl,-, and CuC1,2-; Klc o,, Pxf14,and KHcfl4are the stability constants of the Cu(I1)-oxalato complexes CuC204,C U ( C ~ O ~and )~~-, CuHC204+;Plsq is the stability constant of Cu(II)S03;and Plcl is the stability constant of Cu(II)Cl+. From expression 4,it is clear that the ratio [Cu+]/[Cu2+] is low for small values of a1/a2. a1 depends on the available ligands for Cu(1); if strong ligands are present for Cu(I), a1and the concentration of free Cu+ are very small. On the other hand, if no strong ligands are available for Cu(II), a2is larger than a1and the ratio a1/a2becomes small. Assuming the same ligand concentrations as in Figure 4 (with all ligand concentrations in excess of the total Cu) and a given ratio of [Cu(I)]/[Cu(II)] (=l), the calculated ratio [Cu+]/[Cu2+]varies between about 0.3 and 1X between pH 3 and 7. At higher pH the ratio [Cu+]/ [Cu2+]decreases strongly due to the complexation of Cu(1) by sulfite. But the ratios of [Cu+]/[Cu2+]are still much higher than thermodynamically predicted in the presence of oxygen. The thermodynamic effect of the complexation of Cu(1) by sulfite is thus not sufficient to prevent its oxidation. For the data given in Table I, we can thus expect Cu(1) to be mostly present as sulfite complexes in these fogwaters; in some of the samples with low pH (=3) and high chloride concentrations, the chloride complexes are predominant. The extent of complexation of Cu(I1) is uncertain due to the uncertainties about the available organic ligands. For the oxidation reactions of Cu(1) by oxygen (16) and by H202(17), it has been shown that chloro complexes of copper(1) are oxidized more slowly than free Cu+ ions (Table 11). Wehrli (28) has established that the oxygen1720
Envlron. Sci. Technol., Vol. 25, No. 10, 1991
Table 11. Redox Reactions of Cu(I)/Cu(II) and Rate Constants reactions
rate constants"
-----
Oxidation of Cut Cut + O2 products CuClO + O2 products CuSOc + O2 products Cut + H202 Cu2++ OH- + OH' CU+ + o35 cu2+ + o2+ OHCu+ Fe3+ Fez+ + Cu2+ Cut + R02 Cu2++ R02Cu+ + H02' Cu2++ H202+ OHCut + 02- 3Cu2++ H,02 + 20H-
+
----
Reduction of Cu2+ Cu2++ H202 products [CU"SO~CU"] [ c u ~ s v o ~ c u " ] ~ ' Cu2++ RCHO Cut + RCO' + Ht Cu2++ R02' Cut + ROH + Ht + O2 Cu2++ H02' Cut + Ht + O2 CU2' + 02- c u t + 0 2
2.0 x 104 7.9 x 102 1.6 X lo2 1 x 102 1.7 x 105 1.0x 107 3.0 x 105 1.0x 109 1.0x 109 6.3 X lozc 5.7 x 10-3d 1.0 x 10-5
ref 16 16 b
15 8 8 8
27 27 15 19
4.0 X lo6
8 8
1.0x 107 8.0 X IOs
27 27
Units are M-l s? unless otherwise indicated. Estimated value on the basis of the Marcus relationship after Wehrli (28). Overall second-order rate constant for Cu(I1) reduction by H202in organic-free seawater. dFor this first-order reaction, the unit of the rate constant is SKI.
ation rate of Cu(1) species can be correlated in a Marcus-type relationship to -AG of the reaction Cu(1)X + 0
2
-
Cu(I1) + x
+ 02.-
It seems plausible that sulfite would act in a way similar to chloride and that the sulfite complexes would also be oxidized more slowly than free Cu+. By use of the relationship given by Wehrli (28), it can be estimated that the rate constant for the oxidation of the Cu(I)-sulfite complex by oxygen would be =2 orders of magnitude smaller than the rate constant for the oxidation of the free Cu+ ion. The
effect of sulfite in preserving Cu(1) from oxidation is thus probably a kinetic one. Different factors may contribute to the reduction of Cu(I1) to Cu(1). Possible oxidation and reduction reactions of copper(I1) and copper(1) are listed in Table 11. The reduction of copper(I1) by sulfite as indicated above may be efficient if the pH is sufficiently high (>6). Half-times for the reduction of Cu(I1) under conditions of fogwater may be estimated by using our experimental rate constants, which are of the order of kexp= 30 M-l s-l. For pH >6 and S(1V) concentration of 1 X lo4 M, the calculated half-times are about 2-20 min; for pH = 5 the calculated half-time is several hours. The reaction with sulfite could thus have been significant in some of our samples with pH 5.5-6.3; the high ratios of Cu(I)/Cu(total) found at low pH cannot be explained by this reaction. Radical reactions with H02'/02'- are likely to be important. The rate constant for the reduction of copper(I1) by hydrogen peroxide appears to be small (13),but this reduction rate has also been shown recently (29) to be increased in the presence of chloride and bromide. Reduction of Cu(I1) to Cu(1) by different types of organic compounds is possible, including the reduction by aldehydes and sugars, which is used as an analytical method (Fehling reaction), and by hydroquinone and phenolic compounds (30). As shown in Table I, both high DOC and high concentrations of formaldehyde are observed in fogwater. Only a small part of the organic compounds occurring in atmospheric water samples (fog- and rainwater) has been identified (31, 32);they include, for examples, phenols and nitrated phenols in concentrations up to the micromolar range (33,34). Numerous organic compounds able to act as reductants of Cu(I1) are thus present in fogwater. The total organic carbon is in very large excess of copper (Table I); only a small part of the DOC may be active for the reduction of Cu(II), so that no clear correlations of Cu(1) with DOC can be expected. Photoreduction of Cu(I1) organic complexes yielding Cu(1) has been shown for example for bis(2,g-dimethyl1,lO-phenanthroline) (35) and for different ligands under seawater conditions (15). Moffett and Zika (15) detected copper(1) in surface seawater, in agreement with a photochemical mechanism, copper(1) being in this case stabilized by chloride ions. Photochemical processes may have played a role in some of our samples that have been exposed to light. The light effect is however superimposed on the different.chemical factors. No conclusive difference between samples collected in the dark and in the light can be derived from our data.
Conclusions While Cu(1) i s not thermodynamically stable in atmospheric water (in the presence of oxygen and of other oxidants), we report here on the presence of high Cu(1) concentrations in fogwater. A qualitative explanation of these observations is given on the basis of the limited knowledge of the reactions of Cu(1) and Cu(I1) under these conditions. Cu(1) may be formed by reduction with sulfite at high pH, by reduction with organic compounds, and by various radical and photochemical reactions in the presence of light. The Cu(1) formed can be complexed by sulfite present in excess; in this way the rapid reoxidation of Cu(1) may be prevented. Sulfite is in this regard probably more important in acting as a ligand for Cu(1) than as a reductant. Copper(1) may in this way accumulate in fogwater. The occurrence of significant concentrations of copper in the Cu(1) oxidation state needs thus to be considered in atmospheric models.
Acknowledgments We thank Ph. Behra for discussions, J. Zobrist for help in fogwater sampling, and D. Kistler, C. Mlider, A. Meier, and T. Ruttimann for analytical work. Registry No. Cu, 7440-50-8.
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Received for review August 7,1990. Revised manuscript received April 15, 1991. Accepted manuscript May 31, 1991.
Fluidized-Bed Solid-Phase Extraction: A Novel Approach to Time-Integrated Sampling of Trace Metals in Surface Watert Johannes W. Hofstraat" Tidal Waters Division, Ministry of Transport and Public Works,
P.O. Box 20907, NL-2500 EX The Hague, The Netherlands
John A. Tlelroolj and Hajo Compaan Department of Analytical Chemistry, Division of Technology for Society, Netherlands Organization of Applied Scientific Research TNO, P.O. Box 217, NL-2600 AE Delft, The Netherlands
Wlm H. Mulder Institute for Inland Water Management and Wastewater Treatment, Ministry of Transport and Public Works, P.O. Box 17, NL-8200 AA Lelystad, The Netherlands
A new technique for solid-phase preconcentration of contaminants dissolved in surface waters is presented. The method is based on application of extraction material that is not packed, as in conventional solid-phase extraction systems, but instead is present in a freely floating form, as in a fluidized-bed reactor. The feasibility of the fluidized-bed extraction approach is demonstrated for the determination of heavy metals in surface waters using 8hydroxyquinoline attached to solid supports as complexing agent. Recoveries, repeatibility, and sensitivity appear satisfactory for this application, even when no filtration of the sample is done. As fluidized-bed extraction is based on free-floating, unpacked, extraction material, the pressure drop over the column is minimal and filtration is not required. Hence the technique seems eminently suited for deployment as an in situ, long-term, sampling method. As such it will provide time-integrated contamination levels that are not biased by biological variability or filtration artifacts, disadvantages of the commonly used methods for monitoring of contaminants in surface water. Introduction
Determination of dissolved contaminants is an important requirement for the assessment of the quality of both salty and fresh surface waters. The dissolved contaminants form the major part of the contaminants that are available for uptake by many aquatic biota. Even though the dissolved concentrations of environmental pollutants are extremely low (in particular for apolar organic compounds that have very low aqueous solubility), they can be harmful as a result of bioaccumulation. Data on concentrations of dissolved contaminants also provide information that is vital for the assessment of the effectiveness of regulations aimed at the reduction of the input of pollutants into the aqueous environment. 'Presented at the XVIIth FACSS Meeting, Cleveland, OH, October 1990. *Author to whom correspondence should be addressed: AKZO Research Laboratories, CRL, P.O. Box 9300, NL-6800 SB Arnhem, The Netherlands. 1722
Environ. Sci. Technol., Voi. 25, No. 10, 1991
The very low concentration of dissolved trace materials in surface waters, for heavy metals in the ppb range (I), and for organic micropollutants even down to the sub-ppt region ( 1 , 2 )requires the application of lengthy and time consuming analytical procedures. Vital steps in these procedures are preconcentration of the components of interest and the removal of sample constituents that interfere with the determination. For the determination of heavy metals in seawater in general the-timeconsuming-method described by Danielsson et al. is applied (3). The procedure consists of extraction of the metals from the seawater into an organic phase by application of dithiocarbamate as complexing agent. In this way the very abundant alkali and alkaline earth elements, which are not complexed and interfere with the subsequent determination, are effectively removed. Subsequently, the organic phase is acidified and the heavy metal ions are extracted back into an aqueous phase. Apart from cleanup this procedure also results in a preconcentration of the complexed ions by a factor of -50. For determination of trace organics even more rigorous extraction procedures have to be applied ( 2 , 4 , 5 ) .Apart from being laborious, such concentration methods have a significant risk of loss of compounds and of possible contamination. Alternatively, use can be made of solid-phase extraction of the analytes from the aqueous solution. In this case, a column filled with a specific adsorbent that only retains the materials of interest is used to bring about preconcentration and cleanup of the sample prior to analysis. For preconcentration of organics mostly apolar adsorbents, such as resins or silica spheres coated with alkyl chains, are used (2, 5 , 6). For heavy metals complex-forming agents attached to a variety of solid supports have been applied (7,8). As the extraction procedures, solid-phase techniques are most suitable for use in the laboratory. Furthermore, large sample volumes have to be used to obtain sufficiently high concentrations of the analytes in the final extract. The large volumes, which can be in the order of 10 L for determination of organic compounds, are unfavorable from a logistic point of view. An elegant alternative approach is to use biota for preconcentration in situ. In particular, mussels have been
0013-936X/91/0925-1722$02.50/0
0 1991 American Chemical Society
