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Anal. Chem. 1982, 5 4 , 2536-2539
Solvent Extraction of Dithiocarbamate Complexes and Back-Extraction with Mercury(I I) for Determination of Trace Metals in Seawater by Atomic Absorption Spectrometry J. M. Lo Institute of Nuclear Science, Natlonai Tslng Hue University, Hsinchu, Taiwan, Republlc of China
J. C. Yu, F. I. Hutchison,' and C. M. Wai" Department of Chemistry, University of Idaho, Moscow, Idaho 83843
This two-step preconcentratlon method whlch Involves the extractlon of metal-dlthiocarbamatecomplexes into chioroform at pH 4.5 followed by back-extractlon with a dilute Hg(I1) solution proves to be an efficient and fast way of preconcenlratinga number of trace metals such as Cd, Co, Cu, Fe, Mn, NI, Pb, and Zn In seawater by graphlte furnace AAS. The back-extractlon procedure Is based on the fact that the extractlon constant of Hg( II)dlthlocarbamate Is much greater than those of the other metals under conslderatlon. Because of its low atomlratlontemperature, the presence of a low concentratlon of Hg(I1) in the back-extracted solution does not cause any noticeable Interferences.
The dithiocarbamate extraction method has been one of the most widely used techniques of preconcentration for trace metal analysis by atomic absorption spectrometry (1-10). This extraction method can be generally classified into two major categories. The first one involves the extraction of metaldithiocarbamate complexes into oxygenated organic solvents such as methyl isobutyl ketone (MIBK) and then analyzing the solvents directly (2-5). The other one is to extract the metal complexes into oxygenated or chlorinated organic solvents such as chloroform, MIBK, etc. followed by a nitric acid back-extraction, and then analyzing the trace elements in the acid solution. The latter category has been the subject of a number of recent reports (6-10). There are several drawbacks associated with the acid back-extraction of metal dithiocarbamates. The kinetics is generally slow and the efficiency of acid extraction is poor for certain metals such as Co, Cu, and Fe (8). We have recently developed a new back-extraction method using a dilute mercury(I1) solution instead of nitric acid. This back-extraction method is based on the fact that the extraction constant of the mercury(I1) dithiocarbamate complex is much greater than most of the common trace metals of environmental importance. The substitution of mercury(I1) for other metals in the form of dithiocarbamate complex is extremely fast and the efficiency of recovery is nearly 100% for a number of metals tested including Co, Cu, and Fe. In addition, the back-extracted solution contains a low concentration of mercury(I1) which is virtually interference free in GFAAS due to its high volatility. This paper describes the application of this proposed method to trace metal analysis in seawater which is always a challenging system because of its complex matrix and low metal concentrations. EXPERIMENTAL SECTION Reagents. Ammonium pyrrolidinedithiocarbamate (APDC) and sodium diethyldithiocarbamate (NaDDC) were purchased 'Present address: Idaho Bureau of Mines and Geology, University
of Idaho, Moscow, ID 83843.
from the Fisher Scientific Co. and were Fisher reagent grade and certified ACS grade, respectively. Chloroform and mercury(I1) oxide used in the extraction were Baker analyzed reagents from the J. T. Baker Chemical Co. Nitric acid used in the experiments was of Ultrex grade from Baker Chemical Co. Demineralized water was obtained by treatment of distilled water through an ion exchange column (Barnstead Ultrapure Water Purification Cartridge) and a 0.2-pm filter assembly (PallCorp, Ultipor DFA). All of the metal standard solutions (including Cd, Co, Cu, Fe, Mn, Ni, Pb, and Zn) were prepared in accordance with the EPAs instructions (11)and stored in polyethylene bottles. The mercury(I1) back-extraction solution was prepared by dissolving HgO in 2.4% HNOBand diluted to a final concentration of lo00 pg/mL and a pH about 1.6. Synthetic seawater was prepared by mixing 25.40 g of NaC1,5.083 g of MgCl,, 1.104 g of CaCl,, 0.722 g of KC1, 0.026 g of HSBOS,and 0.203 g of SrC1, in a volumetric flask and made up to 1L with demineralized water (12). Natural seawater (surface) was collected from the coastal waters near the Seattle area using a 15-L polyethylene bottle and filtered through a 0.45-pm Millipore membrane filter before use. The salinity of this seawater was about 27%0. All containers were first washed with 2% Liqui-Nox detergent (Alconox, Inc.) solution and then soaked in 10% NHOBfor at least 24 h. After soaking, they were rinsed with demineralized water several times. The cleaned labware were stored in a class 100 clean hood equipped with a vertical laminar flow HEPA filter (CCI). Extraction Procedures. A mixture of APDC and NaDDC solution (1% each) was prepared by dissolving 1 g of each compound in 100 mL of pH 3 water (HNOJ. The solution was always prepared fresh prior to use and was purified by shaking with 20 mL of chloroform for 20 min. This solution is referred to as the purified extraction agent. Normally, 400-mL aliquots of sample solution were placed into 500-mL ground-glass stopped Erlenmeyer flasks for extraction. The sample solutions included unspiked and spiked demineralized water, synthetic seawater, and natural seawater. The concentrations of the spiked metals varied from 0.25 to 100 pg/L. To each sample 4 mL of the purified extraction agent was added. The pH of the solution was then adjusted to 4.0-4.5 using nitric acid. After pH adjustment, exactly 20 mL of chloroform was pipetted into each flask. The mixture was shaken vigorously for at least 20 min to ensure all metal ions in the aqueous phase had been extracted into the organic phase. After a wait of about 5 min to allow phase separation, 16 mL of the organic phase was pipetted out from the flask and placed into a 20-mL Beckman polyvial with a fast turn cap. To backextract metals, 4 mL of the lo00 fig/mL mercury(I1) solution was added to the vial. Usually, 3 min of shaking is sufficient to complete the extraction. The aqueous phase was then transferred to another vial for AA analysis. The concentrations of metal impurities in the nitric acid and mercury(I1) solution used in this experiment are shown in Table I. The nitric acid was found to contain some Zn and Fe. The procedure blank (i.e., 400 mL of demineralized water that went through the same procedure as the seawater sample) was estimated to have 0.003 pg/L Cd, 0.03 fig/L Cu, 0.05 pg/L Fe, 0.01 pg/L Mn, 0.03 pg/L Ni, 0.04 pg/L Pb, and 0.21 pg/L Zn. The blank appeared to be satisfactorily low and reproducible with relative standard deviations of less than 20% for the metals listed.
0003-2700/82/0354-2538$01.25/00 1982 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 54, NO. 14, DECEMBER 1982
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Table I. GFAAS of the Seawater Samples by the Proposed Method Cd
co
demineralized water 0.000 pH 1.6 HNO, 0.000 pH 1.6 HNO, t 1000 pg/mL Hg2+ 0.000
0.000 0.000
0.003 0.002
20 pg/L M m + in pH 1.6 HNO, 20pg/LMm+inpH 1.6 HNO, 1000 pg/mL Hg2+
0.370 0.461
0.120 0.120
0.170 0.173
0.005 0.442 (95%) 0.127 0.273 X 2 (gn%)
0.000
0.000
metal absorbancea Fe Mn
cu Solution 0.000
Ni
Pb
Zn
0.003 0.017 0.020
0.000 0.000 0.003
0.000 0.000 0.000
0.000 0.000 0.000
0.007 0.067 0.068
0.189 0.190
0.398 0.413
0.076 0.074
0.162 0.190
0.357 0.460
0.035 0.184 (88%) 0.320X2.5 0.394 X 2.5 (109%)
0.016 0.390 (91%) 0.292X2 0.456X 2
0.009 0.080 (96%) 0.075 0.143 (92%)
0.028 0.202 (92%) 0.080 0.252 (91%)
0.384 0.390 X 2 (101%) 0.511X21 0.577 X 2 1 (>loo%)
Standard
+
Sample demineralized water demineralized water spiked with 0.25 pg/L Mm+ natural seawater natural seawater spiked with 0.25 pg/L M m +
0.024 0.113 0.201 (94%) (102%) 0.000 0.236X6 0.115 0.265 X 6 (96%) (100%)
(80%)
All absorbances were measured at peak height and those with multiple sign mean the solutions were diluted due to high concentration. Twenty-five microliters of solutions for all metals was injected into the graphite furnace except those for Cd All samples were and Zn which were 1 0 p L. Numbers in parentheses denote the percent recovery of the spiked metals. started with 400 mL and run through the proposed procedure, The final solutions for GFAAS were all in the mercury(I1)HNO, solution at pH 1.6. The preconcentration factor was 80. a
Atomic Absorption Spectrometry. IEach set of samples consisting of blanks, original samples, and spiked samples was analyzed either by GFAAS or by flame AAS depending on the metal concentrations in solution. In some cases a proper dilution with the mercury(I1) solution was necessary for GFAAS because the metal concentrationn in sample solution were too high. In general, samples spiked with more than 5 pg/L of various metals were analyzed by flame AIS using an Instrumentation Laboratory 353 spectrophotometer. A Perkin-Elmer Model 603 spectrophotometer equipped with a HGA 2100 graphite furnace was used for GFAAS. Metals were analyzed by using Perkin-Elmer pyrolytically coated graphite tubes and operating conditions prescribed by the EPAs inadmctions ( 1 1 ) .
RESULTES AND DISCUBSION Extraction of metal ions by dithiocarbamate is known to depend on several factors such as type and amount of dithiocarbamate, organic solvent, chemical form of metal ion, pH of solution, and shaking time (1-10). We have reinvestigated the extraction ]process in order to obtain optimum conditions which can be combined with our proposed mercury(I1) back-extractioin procedure for AAS. A mixture of APDC and NaDDC rather than APDC OY NaDDC alone was used as the extraction agent in this procedure because of the availability of a broader working pH range (3)and the synergistic effect of the miixture for quantitative extraction of manganese (13).Chloroform was chosen a13 the organic solvent because of its low solubility in water and its high extraction efficiency for most of the metal dithiocarbamates (14). Oxygenated solvents like MIBK, which has a significant solubility in water, are not suitable for the extraction of low levels of metals in a large volume of seawater (7,8).Unlike oxygenated solvents, chloroform alno will not extract C1- and Br- which may cause interference iin AAS (8). The effects of pH on the extraction of various metals according to our results are shown in Figure 1. These results were obtained by using the synthetic seawater spiked with 20 pg/L each of the following metals: Cd, Co, Cu, Fe, Mn, Ni, Pb, and Zn. The pH values were measured after the addition of the APDC and NaDDC mixture because it altered the pH of the initial solution. According to Figure 1, the optimum pH range for quantitative extraction of all of the eight metals investigated is around 4 to 6. Shaking time for the extraction of metal dithiocarbamates by chloroform was studied with the same synthetic seawater at an adjusted pH
cu Nl
Pb
co
Cd
PH
Flgure 1. pH dependence of the extraction of eight trace metals in
seawater by a mixture of diethyldithiocarbamateand pyrrolidinedithiocarbamate into chloroform. 100-
a0
.
8 6 c .-U .-
60.
L Y
111
c Y
20 Mn
0
I
I
4
I
I
a
I
I
12
I
I
16
I
I
20
S h a k i n g t i m e (min)
Extraction times for the spiked metal dithiocarbamates from seawater at pH 4.5 Into chloroform. Flgure 2.
of 4.5. The results are shown in Figure 2. In general, the required shaking time increases with increasing metal concentration in solution. It is obvious from Figure 2 that 20 min of shaking is sufficient to quantitatively extract all of the eight
ANALYTICAL CHEMISTRY, VOL. 54, NO. 14, DECEMBER 1982
2538
100
-
E
.-e0 U
i ,E:
90-
I
A
v
U
; ;
0..
m u
-
a I
80-
I 70
Pb
Y
cod 0.5
1 .o
I
I
1.5
2.0
S h a k i n g t i m e (inin)
Figure 3. Recovery of the metals from the organic phase (CHCI,) into a 1000 ppm Hg(I1) solution at pH 1.6 (HNO,) vs. extractlon time.
metals at concentrations of 20 pg/L or less. For solutions with higher metal concentrations, the required shaking time should be somewhat longer. The mercury(I1) back-extraction process can be expressed by the following equation:
2 Hg2+ + --M(DTC)m(org)
m
2 * -M"+ m
+ Hg(DTC)2,org,
(1)
where DTC denotes dithiocarbamate and subscript org represents organic phase. Thermodynamically, eq 1 should be favored to the right because of the large equilibrium constant = KI-Ig(DTC)2/KM(DTC),2"
(2)
where KHg(DTC)2and KM(DTC), are the extraction constants of Hg(DTC)2and M(DTC),, respectively. The extraction constant is defined as (3) According to previous reports, KHg(DTC)2 is in the oder of loa which is much greater than the K values of the eight metals included in this study (15-25). For example, the extraction constant of Cu is only in the order of lo2', the highest among the eight metals investigated. Basically, this mercury(I1) back-extraction procedure is different from the nitric acid back-extraction which involves the breakdown of metal dithiocarbamates by the acid. In the latter case, the final acid solution will contain various organic species which may cause interference in GFAAS. Kinetically, eq 1 proceeds very rapidly according to our experiments. The percentage of recovery of each metal during mercury(I1) back-extraction as a function of shaking time is shown in Figure 3. Under our experimental conditions, all of the metals studied with initial concentrations of 100 pg/L in the seawater could be back-extracted into the mercury(I1) solution with >95% recovery within 2 min. Those started with 5 pg/L in seawater required even shorter time, less than 0.5 min, for near total recovery. In comparison with the acid back-extraction procedure reported in the literature, this mercury(I1) back-extraction is an extremely efficient process to transfer metals from the organic phase to aqueous solution. The acid back-extraction procedure usually requires shaking for 1h for most metals and even overnight for some metals as reported by Magnusson and Westerlund (8). The same authors also indicated that certain metals like Co, Cu, and Fe cannot be quantitatively transferred from the organic phase to aqueous solution by acid back-extraction (8). It should be pointed out that the mercury(I1) solution used in the back-extraction should be prepared from either HgO
or Hg(NO& but not from HgCl,. When a mercury(I1) solution prepared from HgCll was used in our experiment, the efficiency of back-extraction was found to decrease for Co and Ni. One plausible explanation is that mercury(I1) can form stable chloride complexes, e.g., HgCld2-which inhibit the replacement reaction as shown in eq 1for both Co and Ni (26). The actual mercury(I1) solution used in this study was prepared from HgO with a concentration of 1000 pg/mL and at a pH of 1.6. This mercury(I1) solution has been applied satisfactorily to water samples containing up to 100 pg/L each of the eight metals investigated because of the suprastoichiometric amount of mercury relative to the metals being replaced. During the extraction of metal-dithiocarbamate complexes from seawater, it is possible that trace amounts of C1- end up in the organic phase. This trace amount of Cl- may complex with mercury(I1) in the back-extraction process and thus reduce the mercury(I1) ions available for the replacement of metals in the organic phase. Keeping an excess amount of mercury(I1) in the system also serves the purpose of compensating for this potential loss. Since the atomization temperature of mercury is relatively low (-360 "C) compared with the other metals, the presence of this excess amount of Hg2+ should not interfere with the analysis of the other metals by GFAAS. This is one of the advantages of using mercury(I1) as a back-extracting agent. Generally speaking, a nonvolatile metal present in large amount might affect the atomization of other trace metals in graphite furnace (27-29). The choice of pH 1.6 for the back-extraction solution is based on the fact that mercury(I1) exists in aqueous solution essentially as pure Hg2+at this pH (30). In addition, metal ions back-extracted into the aqueous phase at this pH will not be adsorbed by the container walls during short-term storage (11). In practice, dilute nitric acid is the preferred matrix for GFAAS. If flame AAS is used for the analysis, the proposed preconcentration method should be applicable to any water samples as long as the concentration of the metal of interest is high enough to be detected in the back-extracted solution. Assuming a preconcentration factor of 100, this method combined with flame AAS should be able to detect the specified metals in water samples with concentrations as low as several micrograms per liter. However, most of the metals in natural seawater are usually present at the submicrogram per liter level. In this case, GFAAS which is 2 to 3 orders of magnitude more sensitive than flame AAS is needed to detect these metals in the back-extracted solution. Table I summarizes the results of our GFAAS analysis of the spiked and unspiked natural seawater, the standards, and the blank solutions. All data given in Table I are expressed in absorbance at peak height for each element. Two sets of standards were analyzed for this experiment; both contained Cd, Co, Cu, Fe, Mn, Ni, Pb, and Zn at 20 pg/L. One set of the standards was prepared by using HNOBat pH 1.6 only and the other set had an additional 1000 pg/mL of Hg2+in it. The latter one is similar to the back-extraction solution used in the experiment. Results obtained from these two sets of standards were about the same (Table I) indicating that the presence of 1000 pg/mL of Hg2+in solution did not have any significant effects on the analysis of these eight trace metals by GFAAS. To evaluate the percentage of recovery by this method, the natural seawater was spiked with 0.25 pg/L of each of the eight metals. With a preconcentration factor of 80 according to the procedure described in this experiment, the final solution after back-extraction from the spiked samples should contain 20 pg/L of each metal being recovered. As shown in Table I the percentage of recovery of each metal calculated from our data is quite satisfactory (>go% for most of the cases). The con-
ANALYTICAL CHEMISTRY, VOL. 54, NO. 14, DECEMBER 1982
centration and standard deviation of each metal in the natural seawater sample were calculated to be 0.07 f 0.01 pg/L Cd, 2.04 f 0.12 pg/L Cu, 1.13 f 0.12 pg/L Fe, 0.35 f 0.01 pg/L Mn, 0.22 f 0.04 pg/L Ni, 0.07 f 0.02 pg/L Pb, and 6.60 f 0.50 pg/L Zn. Cobalt wan not detected in the natural seawater; probably its concentration was too low to be determined by this procedure. The concentrations of the seven metals listed above are comparable to those reported for other coastal seawaters (4,7). However, the Cu and Zn values in our sample are higher than those reported in seawater collected from central Puget Sound (111). The seawater sample studied in this work may not be representative of the body of water from which it was taken. The main purpose of this work is to evaluate the kinetics and efficiency of the new preconcentration method. By use of the two-step preconcentration procedure described in this paper, Le., extraction into CHC13 and backextraction by mercwy(Il), the enrichment factor for the metals investigated is 80. If necessary, higher enrichment factors can be easily achieved by using a larger volume of seawater to start with and/or a smaller volume of the mercury(I1) back-extracting solution in the second step. Besides preconcentrating trace metals, another major advantage of this procedure is that the matrix species of seawater such as Ca, S, Al, Na, C1, K, Sr, Mg, Ba, etc. which clo not complex with dithiocarbamate can be simultaneously removed. Many authors have reported that the interferences of the seawater matrix species are serious in determination of trace metals by GFAAS (32-34). In principle, the proposeld method can be applied to other transition metals with their extraction constants lower than that of mercury. Only noble metals such as Au, Pd, and Pt whose extraction constamts are known to be higher than that of mercury are excluded from its application (35-37).
ACKNOWLEDGMENT The authors thank F. M. Chen for technical assistance in the early development of this work.
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RECEIVED for review April 12,1982. Resubmitted August 23, 1982. Accepted September 7,1982. J. M. Lo is indebted to the National Tsing Hua University and the Taiwan Power Company for a fellowship which allowed him to carry out this research at the University of Idaho.
