Environ. Sci. Technol. 2000, 34, 2369-2370
Response to Comment on “Thallium Speciation in the Great Lakes” SIR: We thank the Editor for the opportunity to response to perceptive comments by Cheam (1) regarding our recent paper (2). His comments basically center around the following three issues: first, that monovalent thallium species are favored thermodynamically in natural waters (3-12); second, that the equilibrium distribution of Tl species may have been altered in the presence of complexing agents, particulates, and colloids; and third, that the proportion of trivalent species may have been overestimated as a result of sample acidification. We agree with Cheam that changes in species distribution can occur during sample collection, preservation, storage, and analysis (13). We disagree, however, that any of these factors would have biased the data reported in our paper. Cheam (1, 14) and many others before him (3-12) claim that monovalent species should be the dominant forms of Tl in natural waters because of the high standard reduction potential (+1.28 V) for the reaction Tl3+(aq) + 2e- f Tl+(aq). Our study in essence debunked this belief in showing that the Great Lakes are far from being at equilibrium with respect to Tl chemistry. Our field data are consistent with the results of our laboratory-based experiments, which show that the hydrolysis behavior of Tl in dilute aqueous systems deviates significantly from what one would predict using available thermodynamic data (15). In fact, Cheam himself makes a strong case that Tl behavior cannot be described by equilibrium thermodynamics considering the reactions with ligands in natural waters that can stabilize the Tl3+ species. He rightly noted that the distribution of Tl species can be influenced by complexing agents, particulates, and colloids in water, and it is our contention that these factors are responsible for the unexpected Tl3+/Tl+ ratios that we have reported. There is no evidence to support the claim that Tl+ is oxidized to Tl3+ by dilute (∼0.2 M) nitric acid. Such oxidative reaction between nitrate and thallic ions can be conceptualized as
Tl+ + NO3- + 2H+ f Tl3+ + NO2- + H2O From available standard free energies, log K for the reaction is estimated to be -13.2. Assuming that 5% of the nitric acid is reduced (i.e., [NO2-]/[NO3-] ) 0.05), the ratio of [Tl3+]/[Tl+] is estimated to be 6.32 × 10-16 for a NO3concentration of 0.2 M and a pH of 1.5. Assuming that most of the thallium (∼1.0 × 10-10 M) in the samples was monovalent, acidification of the samples to pH 1.5 would result in a Tl3+ concentration of 6.32 × 10-26, or an insignificant amount of the initial concentration. If one assumes the formation of an ion pair between Tl3+ and NO3under same condition, the ratio of [TlNO32+]/[Tl3+] is estimated to be 1.6 × 107, and the resulting [TlNO32+] /[Tl+] ratio is estimated to be 10-9. In other words, formation of such an ion pair would not enhance the conversion of Tl+ to Tl3+ to any significant degree. In the very unlikely scenario where the nitrate is reduced to ammonia, the reaction can be considered to be
4Tl+ + NO3- + 10H+ f 4Tl3+ + NH4+ + 3H2O Log K for this reaction is estimated to be -54. Assuming 10.1021/es0020052 CCC: $19.00 Published on Web 04/29/2000
2000 American Chemical Society
reduction of 5% of the NO3- as a result of acidification, the [Tl3+]/Tl+] ratio is estimated to be 1.3 × 10-17 for a NO3- concentration of 0.2 M and a pH of 1.5. Assuming initial Tl+ concentration in water samples of 1.0 × 10-10 M, the Tl3+ concentration is estimated to be less that 1.3 × 10-5% of the original Tl+ concentration. These calculations suggest that the nitrate ion is not an effective oxidant for Tl+ and vitiate Cheam’s contention that “at 10-10 M thallium and pH 1.5 or even at pH 2 there is a substantial amount of Tl3+ formed” has no theoretical basis whatsoever. There is some experiments evidence to support our inference that Tl3+ is formed naturally in the Great Lakes and is not an artifact of sample acidification. Batley and Florence (18) used several methods to demonstrate the predominance of Tl3+ species in their seawater samples. They used anionexchange resins to trap any Tl(III) species present without adding bromine water to oxidize the Tl(I) species in the sample (18, 20). They found that the recovery of Tl without addition of bromine water was about 77% of the total Tl concentration in the original seawater sample. Batley and Florence (18) also reported that the 204Tl(I) spiked in river water samples was not oxidized to Tl(III) when the samples were acidified to pH ∼1.0 with HCl; they found that 50% of the 204Tl(I) spike in seawater was oxidized to Tl(III) in 5 days however. In our own experiments, we found that the average recovery in nine samples spiked with 40 µg/L Tl+ and acidified with HNO3 was 99 ( 2.6%, suggesting that most of the Tl+ passing through the resin was not oxidized (19). In the Chelex-100 recovery experiments, we found average recovery of Tl(I) to be 92% (2, 19). Cheam attributed the 8% deficit in extraction efficiency to the oxidation of Tl(I) during the separation step. We disagree with this highly capricious interpretation of our results. It is rare to get 100% recovery in spike experiments with natural waters, and the discrepancy can be due to any number of factors in sample analysis. One cannot, for instance, rule out the possibility that the discrepancy was due to incomplete oxidation of Tl+ with bromine. In fact, we were more concerned with the reduction of Tl(III) in the sample than with the oxidation of Tl(I) during the recovery and selectivity tests.
Literature Cited (1) Cheam, V. Environ. Sci. Technol. 2000, 34, 2367-2368. (2) Lin, T.-S.; Nriagu, J. Environ. Sci. Technol. 1999, 33, 3394. (3) Kaplan, D. I.; Mattigod, S. V. In Thallium in the Environment; Nriagu, J. O., Ed.; John Wiley & Sons: New York, 1998; Chapter 2. (4) Flegal, A. R.; Patterson, C. C. Mar. Chem. 1985, 15, 327. (5) Bodek, I.; Lyman, W. J.; Reehl, W. F.; Rosenblatt, D. H. Environmental Inorganic Chemistry; Pergamon: New York, 1988. (6) Vink, B. W. Chem. Geol. 1993, 19, 119. (7) Smith, I. C.; Carson, B. L. Trace Metals in the Environment Chemistry V(I); Ann Arbor Science: Ann Arbor, MI, 1977. (8) Schoer, J. In The Handbook of Environmental Chemistry, Vol. 3, Part C; Hutzinger, O., Ed.; Springer-Verlag: New York, 1984. (9) Lee, A. G. The Chemistry of Thallium; Elsevier Publishing Co.: New York, 1971. (10) Shaw, D. M. Geochim. Cosmochim. Acta 1952, 2, 118. (11) Landford, A. Effect of trace metals on stream ecology. Presented at the 1969 Cooling Tower Institute Annual Meeting, Anaheim, CA, January 20, 1969. VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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(12) Wedepohl, K. H. Handbook of Geochemistry, Element 81, Vol 2, No. 3; Springer-Verlag: Berlin, 1972. (13) Pickering, W. F. In Chemical Speciation in the Environment; Ure, A. M., Davidson, C. M., Eds.; Blackie Academic & Professional: London, 1995. (14) Cheam, V.; Lechner, J.; Desrosiers, R.; Sekerka, I.; Lawson, G.; Mudroch, A. J. Great Lakes Res. 1995, 21 (3), 384. (15) Lin, T.-S.; Nriagu, J. J. Air Waste Manage. Assoc. 1998, 48, 151. (16) Baes, C. E.; Mesmer, R. E. The Hydrolysis of Cations; John Wiley & Sons: New York, 1976. (17) Nordstrom, D. K. Water, Air, Soil Pollut. 1996, 90, 257. (18) Batley, G. E.; Florence, T. M. J. Electroanal. Chem. Interfacial Electrochem. 1975, 61, 205. (19) Lin, T.-S.; Nriagu, J. Anal. Chim. Acta 1999, 395, 301. (20) Matthews, A. D.; Riley, J. P. Anal. Chim. Acta 1969, 48, 25.
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ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 11, 2000
Tser-Sheng Lin* Department of Environmental Engineering and Health Yuanpei Technical College 306 Yuan-Pei Street Hsin Chu City, 300 Taiwan
Jerome Nriagu Department of Environmental Health Sciences School of Public Health University of Michigan Ann Arbor, Michigan 48109 ES0020052
