oscillator gate (upper right hand corner of Figure 3) were added. Prior to the addition of oscillator gate 2, it was observed that upon reset of the transient recorder (Le., change back to fast recording rate), a short burst of pulses occurred, which was capable of changing the time relationship between the two noise spikes within the SAD 100 (i.e., changing the relative states of the two internal ring counters). The oscillator gate (gate 2) now ensures that no clock pulses can leave the timing circuit during the reset cycle. The cue index circuit, composed of the modulo-50 counter and flag circuitry, generates a puke once per 100 clock cycles of the SAD 100. Since the noise spikes from the SAD 100 are generated by a particular state of the ring counters internal to the device, their time relationship to the cue index pulse will remain constant while the system is operating. This time relationship is not reproducible over successive power up cycles of the system hbwever, so it is necessary to collect a baseline scan each time the system is powered up, and then use this baseline scan to correct the data scans which follow., Figure 7 shows the results of a baseline subtraction made on a data set consisting of 100 data points collected at 10 MHz from a constant input voltage (-3V). The bottom trace (a) represents 100 averages of a low baseline voltage. The middle trace (b) is the single shot of data, and the upper trace (c) is the result of the baseline correction. The signal-to-noise (S/N) ratio of the uncorrected data scan is 210 and the S/N ratio of the baseline corrected scan is 378. This corresponds to nearly 9 bits of resolution. The S/N ratios were obtained by dividing the average of the data set by its standard deviation. Transient Recording. The accuracy of the transient recorder in a real application is illustrated in Figure 8. The transient resulting from an RC decay wasrecorded a t 10 MHz.
The transient was produced by applying a 200-11s current pulse to the parallel combination of a 300-R, O.Ol% resistor and a O.Ol-wF, 0.5% capacitor and monitoring the voltage decay. A weighted least sfluares fit was performed on the linearized data and the slope of the fit compared to the RC time constant for the components mentionedJt is seen from the figure that the measured result is in good agreement with the theoretical value. The cost of the SAD 100 is $100. The total parts cost of the transient recording system as described (not including computer and peripherals) was under $1000. The performance of the system with respect to speed and accuracy is certainly competitive with commercial transient recorders costing much more. The combination of cost and performance makes this method of transient recording very appealing, and the potential applications for this technique should increase still further as MOS monolithic technology expands.
LITERATURE CITED (1) J.E. Stewart, Am. Lab.,No. 11, 91-101 (1973). (2) W. P. Cargile and B. J. Moore, Res. Develop., 24, 32-35 (1973). (3) Data Sheets for SAD 100, obtained (on request) from Reticon Gorp., 910 Benicia Ave., Sunnyvale, Calif. 94086. (4) G. Horlick, Anal. Chem., 48, 783A-787A (1976). (5) H. Malmstadt, and C. G. Enke, "Digital Electronics for Scientists", W. A. Benjamin, New York, 1969, p 226.
RECEIVEDfor review June 25, 1976. Accepted October 1, 1976. The authors are grateful to Michigan State University for the Quill Fellowship received by one of us (T.A.L.). This paper was presented at the second annual meeting of the Federation of Analytical Chemistry and Spectroscopy Societies.
Loss of Carbon- 14 and Mercury-203 Labeled Methylmercury from Various Solutions Kenneth R. Olson' Department of Environmental Physiology, Rutgers University-The State University of New Jersey, New Brunswick, N.J. 08903
Disappearance of methylmercury and the lablllty of the methylmercury bond during volatilization was studied by the Incubation of submicromolar concentrations of I4CH3HgCIand CH3203HgCIfor 96 h in solutions routinely encountered In environmental and laboratory situations. Media with high ionic strength, 0.5 M NaCi and artificial sea water lost more mercury than deionized water, tap water, water from fish holding tanps, basic and acldlc soiutlons, and phosphate buffered saline. Loss from the former solutions appears to be a continuous process while occurring In the first few hours in the latter. Total loss did not exceed 20 % In any experiment and was due primarily to mercury in the Inorganic form.
Loss of inorganic mercury(I1) from dilute aqueous salt solutions has been observed by numerous investigators ( I - 5 ) , and, although there is somewhat less evidence, it appears that monomethylmercury is also lost from aqueous solutions (6, 7). Present address, Indiana University School of Medicine, South Bend Center for Medical Education, Medical Science Building, University of Notre Dame, Notre Dame, Ind. 46556.
Burrows and Krenkel(7) aerated submicromolar methylmercury 203Hgsolutions for 20 h and demonstrated considerable variation in loss of isotope depending on the nature of the aqueous media employed. Under these conditions 91% of the radioactivity was lost when methylmercury was placed in deionized tap water, 5%when in Nashville tap water, and 30% if the aforementioned tap water was charcoal filtered. Water collected from the shallows of a reservoir lost around 15%of the isotope. The authors suggest that the loss of mercury from the above solutions is probably in the methylated form although studies have not been done to verify this point. The purpose of the present investigation is to pursue the work of Burrows and Krenkel with respect to 1) characterization of aqueous solutions from which methylmercury can be lost and 2) the lability of the methyl-mercury bond during mercury loss from various aqueous solvents.
EXPERIMENTAL Glass staining disbes 94 X 70 X 55 mm (L X W X H)were used for the aeration vessels. It was thought that use of conventional aquarium air stones for aeration might provide a relatively large surface area for adsorption of the methylmercury so to prevent this possibility all aeration was with plastic tubing heat-sealed on one end with a hemostat qnd perforated 20 times with a 22-gauge stainless steel needle. ANALYTICAL CHEMISTRY, VOL. 49, NO. 1, JANUARY 1977
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loss from distilled-deionized water (A) 203HgExperiment A; (A)14CExperiment A, (0) 203HgExperiment B; (0)14C Figure 1. Methylmercury
IO 6
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Time (hrs) Figure 2. Methylmercury loss from 1% H2S04.Symbols as in Figure 1
Experiment B
All perforations were within 15 mm of the sealed end. Several aeration vessels were used simultaneously and air flow to each was regulated with a gang valve and bypass system to ensure adequate agitation and aeration (a fine stream of air bubbles) yet prevent spattering. Wire taped to the outside of the vessel and to the aeration tubing provided support to the aerators and allowed the tubing to be positioned in the bottom center of the vessel. Distilled, deionized water with a conductivity equivalent less than 1kg/ml NaCl was used for all reagents. Three hundred ml of the following solutions were used for experiment A: 1) Deionized water. 2) New Brunswick tap water. 3) Fish water (water from tanks holding 5 ) HzS04,10.6 ml/l. (1%). 6) rainbow trout). 4) NaOH, 10 g/l. (1%). Phosphate buffered physiological saline pH 7.3 (NaC1, 7.37 gfl.; KC1, 0.31 g/l.; CaC12, 0.10 g/l.; MgS04, 0.14 g/l.; KHzP04, 0.46 g/l.; Na?HPOJ, 2.02 g/l.). 7) NaC1, 29 g/l. 8) Artificial sea water (marine salt, Nektonics Research and Development Ltd., Brooklyn, N.Y.). These solutions were chosen to give a wide range of pH and salinity and because many of them approximate those routinely encountered in environmental and laboratory situations. Experiment B was identical to A except for a small amount of ethanol (5 X MA.) which was added to the stock mercuric chloride solution to aid in solubilization. In most experiments ethanol did not have any consistent effect on the experimental results when compared to Experiment A but the data are plotted separately as slight differences were found in several experiments. Methylmercuric chloride as l4CH3HgCl (specific activity 4.17 mCi/mM) and as CHaZo3HgCl(specific activity 921 mCi/mM) was obtained from New England Nuclear. After dilution in either deionized water (Experiment A) or dilute ethanol (Experiment B), the two radiolabeled mercurials were mixed and aliquots placed in the aerated baths. The maximum concentration of methylmercury for any one experiment was 6.0 X loT7M/1. (Experiment A) or 1.9 X MA. (Experiment B). One-ml samples were removed from the baths after 3 min mixing (0 h), and then at 1,2,4, 8,24,48, and 96-h intervals. These samples were mixed with 10 ml of New England Nuclear Formula 947 liquid scintillation counting cocktail and 2 ml of deionized water to form a thick gel. At the conclusion of Experiment A, approximately 1 inch of the end of the aeration tubing from each bath was cut off and placed in a scintillation vial to determine the absorption of mercury by the tubing. Since the energy spectra of I4C and 203Hgare somewhat similar (8) and a gamma counter was not available, the two isotopes were separated on the basis of the relatively short half-life of z03Hg(46.6 days). The samples were counted 4 or more times over the course of 4-6 weeks and stored at 5 "C in the interim. Prior to each counting period, the samples were warmed to 40 "C in an oven and shaken to reform the gel (9). This procedure did not appreciably alter the external standard channels ratio nor change the efficiency of prepared 14Cand z0,3Hgstandards. The following formulas were derived to determine the activity in counts per minute (cprn) of 20314gand 14Cand to correct for decay of 20'iHg:
where N1 = total cpm for z03Hgand I4C a t first counting period. Nz = total cpm for z03Hg and 14C at second counting period. X = 0.693/46.6 = 0.01487. t l = number of days elapsed between time of 24
ANALYTICAL CHEMISTRY, VOL. 49, NO. 1, JANUARY 1977
assay of 203Hgby New England Nuclear and first counting date. t z = number of days between assay of zo3Hg and second counting date. Approximate efficiency was determined for a series of 203Hgand I4C quenched standards prepared from the stock radiolabeled mercurials and, in addition, I4C efficiencies were compared to commercial quenched standards. Routinely, the activity (cprn) for zo3Hg and 14Cwas calculated from the first and third counting periods and from the second and fourth periods and these values were then averaged. Differences between the two calculated activities were usually under 5%.
RESULTS AND DISCUSSION The most significant finding of the present study was the failure of deionized water to lose substantial quantities of methylmercury during the 96-h experimental period (Figure 1).The percent of 203Hglost from deionized water is at most 9% (Experiment A) attributable primarily to the relatively large loss at the onset of the experiment as no more than 2 to 3%of the isotope is lost thereafter. This is consistent with the results of Experiment B, which showed essentially no mercury loss, and the 14Cdata from both experiments A and B again, with little loss of the label. These results in part quantitatively contradict those of Burrows and Krenkel (6),who found up to 91% of the methylmercury was lost from deionized water in 20 h. This discrepancy is difficult to explain. The large surface of the airstones used by Burrows and Krenkel (6) seemed to be a likely candidate for absorption of the mercurial (hence the use of plastic tubing in the present study); however, 1)airstones were also used by Burrows in solutions which lost essentially no mercury ( 6 ) , 2) methylmercury was lost from unaerated samples ( 7 ) ,and 3) the airstones appeared not to accumulate substantial amounts of mercury (IO). The only other explanations for this discrepancy would be the inherent characteristics of the water or those of solutes found with the stock mercurials. The former could be altered by various organics leached from the ion exchange resin, although it is doubtful that they would reach sufficient concentrations to bind stoichiometrically with the mercurial. The stock mercurial solutions were very similar between this study and that of Burrows and Krenkel (6) as both used CH3203HgC1from the same supplier (New England Nuclear) and l4CH3HgC1 used in the present study was obtained in crystalline form and dissolved in deionized water (Experiment A) or deionized water containing dilute ethanol solution (Experiment B). Aerator tubing was not a factor in the present experiment as it contained approximately the same 203Hg(but proportionately less I4C) as 1ml of medium. The low I4C probably does not reflect preferential absorption of *03Hgbut merely absorption of methylmercury (14C and *03Hg)into the plastic where escape of the energetically weaker 14C beta particle would be reduced. Loss of mercury from tap water and aquarium (fish) water was quantitatively and chronologically similar to values reported by Burrows and Krenkel (6, 7 ) and confirm their
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findings. Less than 5% of either label was lost from New Brunswick tap water in Experiment A or B and, with the exception of an initial 10% decrease in 203Hglevels in Experiment B, neither mercurial was lost from fish water during the 96-h experimental period. Inorganic mercury appeared to be lost from the basic medium (1%NaOH) as 10 and 14% of the mercury-203 was lost from Experiments A and B, respectively, without any corresponding loss of carbon-14. The short time course for this loss (within 8 h) might reflect, in part, loss of inorganic contaminants, Quantitatively similar losses of mercury-203, (6 and 14% for Experiments A and B, respectively) from the acid media were also observed (Figure 2); however, in Experiment B there was a concomitant 14% loss of carbon-14 and the concentration of both mercurials in Experiment B decreased slowly during the entire experiment suggesting some CH3-Hg cleavage. Why different results were observed between Experiments A and B in an acid medium is unknown. The solution chemistry for methylmercuric salts a t various pH's can involve formation of (CH3Hg)zOH+, (CH3Hg)30+, CHsHgOH, CH3HgOH2+ and possibly others (1I ); thus, ethanol could increase volatilization of one of these species within a specific pH range. The similarity, however, between the 203Hg lost in both acidic and basic media suggests that the disappearance of mercury is due to loss of mercury in the inorganic form. This inorganic mercury could be an intrinsic contaminant or from CHS-Hg bond cleavage in the incubation vessel. If the methyl group is split from mercury and oxidized to COZ, it would then be lost from an acid medium and retained in the basic medium as bicarbonate, whereas if methane was formed it would be lost from both acid and basic environments. Phosphate Buffered Saline (PBS) lost *03Hg from both experiments A and B (9 and 6%, respectively), while only Experiment B exhibited any loss (11%)of 14C.The z03Hgloss was fairly rapid in the first 6-10 h of both experiments and insignificant thereafter. The largest and most sustained mercury-203 and carbon-14 losses were from the high salt environments, 0.5 M NaCl and sea water. In both media the loss was continuous over the 96-h period (Figures 3 and 4) and relatively more z03Hgwas lost than 14C (20 vs. 10%and 16 vs. 0% for NaCl Experiments A and B, respectively, and 14 vs. 4% and 15 vs. 7% for sea water, Experiments A and B, respectively). Continued loss of label from NaCl and sea water could be explained by either a salting out effect increasing the volatilization of mercury or by chemical interactions which either increase volatilization or, more likely, precipitate the mercury or increase the adsorption/absorption process. Rose (12),in Amend et al. (13)states that ethyl mercuric ion precipitates with chloride. It is therefore not impossible that such events could happen with methyl and/or inorganic mercury. Lack of stoichiometry between 203Hgand 14C losses suggests that the inorganic species
6 2
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Time (hrs) Figure 3. Methylmercury loss from 0.5 M NaCI. Symbols as in Figure
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Time (hrs) Figure 4. Methylmercury loss from artificial sea water. Symbols as in Figure 1
is the prominent ion lost. It would be interesting to determine whether any other anions or cations produce this effect and to ascertain the importance of halides in this phenomenon. The time course for mercury loss is interesting in that many solutions lost mercury over the first few hours while only a few showed a prolonged disappearance rate. This would suggest that the early loss could be due to either: 1)Precipitation or volatilization of a small amount of the stock mercury not in the methylated form, 2) adsorption and/or absorption of methylmercury onto the glass of the staining dishes or aerator tubing, or 3) combination and/or breakdown of methylmercury with impurities found in trace amounts in the solutions used, producing a more volatile reaction product. The first two hypotheses seem more plausible since Burrows and Krenkel found approximately 5% inorganic mercury in methylmercury-203 samples from New England Nuclear (7) and they subsequently reported 1-2% absorption of methylmercury onto glass beakers (6).
SUMMARY Loss of methylmercury from water containing various solutes appears to be dependent on the nature of these solutes and proceeds with two somewhat typical rate constants. These data do not support the hypothesis that most of the methylmercury escapes from micromolar aqueous solutions via volatilization or any other absorption or precipitation process. They do, however, demonstrate the complexities in the solution chemistry of methylmercury and strongly suggest that further studies are needed to resolve the basic issue of how this environmental contaminant behaves in its most common environment, water. ACKNOWLEDGMENT The author thanks Helen 'Mozolic, Eugene Borbely, and Marc Katchen for their most capable technical assistance. LITERATURE CITED J. M. Lo and C. M. Wai, Anal. Chem., 47 1869 (1975). C. Feldman, Anal. Chem., 46, 99 (1974). Y. K. Chau and H. Saitoh, fnviron. Sci. Technol., 4, 839 (1970). A. M. lgoshin and L. N. Bogusevich, Gidrokhim. Mater., 47, 150 (1968). M. E. Hinkle and R. W. Learned, U.S. Geol. Surv., Prof. Pap. 650-0, D251 (1969). (6) W. D. Burrows and P. A. Krenkel, Anal. Chem., 46, 1613 (1974). (7) W. D. Burrows and P. A. Krenkel, Environ. Sci. Technol., 7, 1127 (1973). (8) . . R. J. Cousins, R. A. Wvnveen, K. S. Sauibb. and M. P. Richards, Anal. Biochem., 65,412 (1975). (9) Y. Kobayashi, New England Nuclear Corp. Boston, Mass., Personal Communication, 1975. (10) W. D. Burrows, Aware Inc. Nashville Tenn., Personal Communication, (1) (2) (3) (4) (5)
1975.
(1 1) D. L. Rabenstein, C. A. Evans, M. C. Tourangeau, and M. T. Fairhurst, Anal. Chem., 47, 338 (1975). (12) E. Rose and A. Rose, "The Condensed Chemical Dictionary", Reinhold, New York, 6th ed., 1961, p 1257. (13) D. F. Amend, W. T. Yasutake, and R. Morgan, Trans. Am. Fish Soc., 88 (3), 419 (1969).
ANALYTICAL CHEMISTRY, VOL. 49, NO. 1, JANUARY 1977
25
RECEIVEDfor review August 19, 1976. Accepted October 1, 1976. *his Research was supported in part by a Rutgers University Research Council Grant. Paper of the Journal
Series, New Jersey Agricultural Experiment Station, Cook College, Rutgers University-The State University of New Jersey, New Brunswick, N.J.
Mechanism of Reduction of Cadmium by Aminoiminomethanesulfinic Acid in Alkaline Media Juiian E. McGIIi’ and Frederick LindstromI Department of Chemistry and Geology, Clernson University, Clemson, S.C. 2963 1
Amlndiminomethanesulflnic acid, formamidinesuiflnlc acid, or thlourea dioxide Is the compound made by the reaction of thiburea and hydrogen peroxide. This compound Is remarkable In that It readily reduces cadmium to the metal from ahueous solutions. Zinc metal does not precipitate. Dithionlte, S2042-, occurs in alkaline solutions of amin~imlnomethanesulflnlcacid. Solbtions of dithionite do not redbce cadmium to the metalllc state and finding carbon In the precipitatedcadmium Indicated that the species responsible for the strong reducing properties of the compound is a radlcal Ion formed by a homolytic cleavage of the carbon-sulfur bond. A scheme to explain the above observations is presented.
Aminoiminomethanesulfinic acid, known also as formamidinesulfinic acid and thiourea dioxide, was first prepared in 1910 by de Barry Barnett ( I ) by the oxidation of thiourea with aqueous solutions of hydrogen peroxide. The fact that alkaline solutions of this compound exhibit strong reducing properties, reducing cadmium and those elements less electropositive than cadmium to the metallic state, was reported in 1936 by Boeseken ( 2 ) .Since this discovery there have been numerous commercial applications of this reducing property, especially in the textile industry. In addition to the commercial use, this compound has been utilized to a limited extent in chemical analysis. In acid media where it is reported to act as a complexing ligand rather than as a strong reilucer ( 3 ) ,it has been used for the quantitative separation of platinum from solution ( 4 ) ,for the precipitation of palladium ( 5 ) ,for the determination of rhodium (6),and for the separation of rhodium and iridium (7).In neutral solution with boiling and in alkaline solution with heating, it has been used to reduce uranium(V1) to uranium(1V) (a), and also in alkaline media for the microdetermination of tin by precipitating tin(I1) sulfide (9), and the reduction of iron(II1) hydroxide (IO). The decomposition of aminoiminomethanesulfinicacid in alkaline media is apparently quite complex. Some workers have attributed the fact that alkaline solutions of this compound possess greater reducing properties than dithionite to the formation of sulfoxalate ion (2, 11, 12) by hydrolysis of aminoiminomethanesulfinic acid while others attribute these properties to the compound itself (13, 14). Golovnya and Bolotova (3)studied the reducing properties and the decomposition products of aminoiminomethanesulfinic acid in various media. They reported that the magnitude and rate of change of potential was greatly influenced by the Present address, College of Pharmacy, Medical University of South Carolina, 80 Barre Street, Charleston, S.C. 29401. 26
ANALYTICAL CHEMISTRY, VOL. 49, NO. 1, JANUARY 1977
media, concentration, and temperature. Potentiometric studies showed that this compound was not a strong reducing agent in acid media and in neutral solution and that the acid itself reacted with strong oxidizing agents, the number of equivalents of oxygen consumed depending on the medium and the oxidizing agent. In alkaline media the acid and its decomposition products, which were more readily oxidized than the acid, reacted with oxidizing agents in freshly prepared solutions. Alkaline solutions were reported to decompose rapidly. The strong reducing properties in alkaline media were attributed to the decomposition products, chiefly to sulfoxylic acid and its salts. The ease and low cost with which aminoiminomethanesulfinic acid can be prepared and purified make it a very attractive analytical reagent. Its greatest potential usefulness as an analytical reagent is apparently due to the strong reducing properties exhibited in alkaline solution. Cadmium and those elements less electropositive are reduced to the metallic state while zinc, for example, is not reduced. This suggests that a better understanding of the mechanism of reduction in alkaline media could possibly lead to conditions that would make this separation and similar ones possible, thereby greatly enhancing the usefulness of this compound as an analytical reagent.
EXPERIMENTAL Apparatus and Material. Ultraviolet spectra were recorded using a Perkin-Elmer Model 4000-A UV/Visible Spectrophotometer and a Cary Model 14 Recording Spectrophotometer. Infrared spectra were made with a Perkin-Elmer Model 137 Spectrophotometer and electron paramagnetic resonance (EPR) studies were conducted on a Varian Model E-3 EPR Spectrometer. A Leeds & Northrup Model 62200 Electro-Chemograph was used for polarographic studies and measurements of pH were made with a Leeds & Northrup Model 7405 pH Indicator which had been previously standardized using appropriate National Bureau of Standards reference buffers. Measurements of potential were made with a Leeds & Northrup Model 7401 pH Indicator equipped with a platinum indicator electrode. X-ray diffraction patterns were recorded with a Norelco Automatic Recording X-ray Diffractometer. Allied Chemical special high-purity 59-Grade Cadmium “D” Metal, 99.999% pure, was used for preparation of stock solutions of cadmium. Technical grade thiourea was purified by recrystallization from ethanol prior to use. Practical grade sodium dithionite was purified using the procedure of Jellinek (15) with modifications suggested by Christiansen and Norton (16)with the exception that a nitrogen atmosphere was used instead of carbon dioxide. The water used was distilled and deionized by passage through an ion-exchange column charged with Bio-Rex RG-501-X8 mixed-bed resin. All other chemicals were of reagent grade unless otherwise noted. Stock solutions of cadmium were prepared by dissolving the metal with the appropriate inorganic acid, diluting to volume with water, and standardizing by titration with EDTA at pH 10 using Calmagite as indicator. A cross check of the standardization procedure was made using the gravimetric cadmium sulfate procedure ( I 7). Aminoimi-
