Stability constants for metal arsenates
Janice S. Lee A B C and Jerome O. Nriagu AA Department of Environmental Health Sciences, School of Public Health, University of Michigan, Ann Arbor, MI 48109, USA.
B Present address: National Health and Environmental Effects Laboratory, US Environmental Protection Agency, Research Triangle Park, NC 27711, USA.
C Corresponding author. Email: Lee.JaniceS@epa.gov
Environmental Chemistry 4(2) 123-133 https://doi.org/10.1071/EN06070
Submitted: 15 November 2006 Accepted: 15 March 2007 Published: 17 April 2007
Environmental context. The behaviour of arsenic compounds is controlled by their solubility, which in turn controls both the forms and mobility of arsenic in the environment. Current knowledge on arsenic chemistry may be distorted because information available on the solubility of arsenic compounds usually does not include possible formation of metal–arsenate complexes. Our results show the formation of stable metal–arsenate complexes that have not been considered before, and this new data can be used to further examine the effect of these complexes on controlling the fate and transport of arsenic in the environment.
Abstract. The formation of solid metal arsenates could conceivably reduce the concentrations of arsenate and metal ions in natural and contaminated aqueous ecosystems, and possibly in human body fluids. In this study, solid metal arsenates were dissolved isothermally in solutions with different molar concentrations of arsenic acid. The saturated solutions were analysed and the results processed to derive the solubility products (Ksp) for solid phases and association constants (K) for metal arsenate ion-pairs. Ion chromatography was used to confirm the presence of ion-pairs, some of which had never before been considered. Association constants were determined for the following ion-pairs: FeHAsO4+ (log K = 4.88), CoHAsO40 (log K = 1.50), ZnHAsO40 (log K = 3.28), SrH2AsO4+ (log K = 1.72), and Ag2H2AsO4+ (log K = 4.50). The following metal ions apparently do not form stable complexes with HAsO42–: Cd2+, Cr3+, Cu2+, Mg2+, Mn2+, Ni2+, Pb2+, and Sn2+. Standard state solubility products (Ksp°) were redetermined for the following compounds: Ag3AsO4, Cd3(AsO4)2, Co3(AsO4)2, CrAsO4, Cu3(AsO4)2, FeAsO4, Mg3(AsO4)2, MnHAsO4, NiHAsO4, PbHAsO4, Sn3(AsO4)2, Sr3(AsO4)2, Zn3(AsO4)2·Zn3(AsO4)2·8H2O (koettigite), Cu2Al7(AsO4)4(OH)13·12H2O (ceruleite), and Pb2CuAsO4CrO4OH (fornacite). Our results show the formation of ion-pairs for some metal arsenates and indicate that previous studies have overestimated the solubilities of many arsenates.
Acknowledgements
We thank Ali Bazzi and the Department of Natural Sciences of the University of Michigan at Dearborne for assistance and permission to use the FAAS. We also appreciate the efforts of Chris Palenik in the Geological Sciences Department at the University of Michigan for his assistance with SEM and XRD analysis. We thank AWWRF and the Horace Rackham School of Graduate Studies of the University of Michigan for financial support.
0.1 N solution (total volume 10 mL) of cadmium acetate (Cd(C2H3O2)2·2H2O) from Mallinckrodt Baker and 0.1 N solution (total volume 10 mL) of sodium arsenate (Na2HAsO4) were mixed in stoichiometric amounts at 40–50°C. The precipitate was left standing in the original solution for several days, and then washed with water and dried. A 0.5 N solution (total volume 50 mL) of sodium arsenate was added to a 1.0 N solution (total volume 25 mL) of chromium potassium sulfate (CrK(SO4)2·12H2O), from J.T. Baker, in stoichiometric amounts at a temperature of 50–60°C. The precipitate was left standing in the original solution for a few days before being washed and subsequently dried at 50°C. A 0.1 N solution (total volume 150 mL) of sodium arsenate was mixed with a 0.1 N solution (total volume 150 mL) of copper nitrate (Cu(NO3)2·3H2O) in stoichiometric amounts with stirring. The mixture contained an excess of 3% nitrate; copper nitrate was obtained from Aldrich. The temperature of the solutions was 50–60°C. The precipitate was kept in the original solution overnight before being washed and dried. Heated 0.1 N solution (total volume 300 mL) of ferric sulfate (Fe2(SO4)3·nH2O) from J.T. Baker and 0.1 N solution (total volume 300 mL) of sodium arsenate were mixed together in stoichiometric amounts at 50–60°C. The resulting precipitate was kept overnight in the original solution, washed, and dried at 60°C. 0.1 N solution (total volume 150 mL) of lead acetate (Pb(C2H3O2)2 ·3H2O) obtained from MCB Manufacturing Chemists and 0.1 N solution (total volume 150 mL) of sodium arsenate were mixed together in stoichiometric proportions at 60°C. The lead acetate solution included 20 g L–1 of free acetic acid. The precipitate was filtered, washed, and dried at 50°C. 0.1 N solution (total volume 10 mL) of magnesium sulfate (MgSO4·7H2O) from Mallinckrodt Baker and 0.1 N solution (total volume 10 mL) of sodium arsenate were mixed in stoichiometric amounts at 50°C. The precipitate was left overnight in the original solution, and then washed and dried. Stoichiometric proportions of 0.1 N solution (total volume 300 mL) of manganese sulfate (MnSO4·H2O) obtained from Mallinckrodt Baker and 0.1 N solution (total volume 300 mL) of sodium arsenate were mixed together. The solution was heated and the precipitate was kept in the original solution for a few days. The precipitate was washed and then dried at 70°C. 0.5 N solution (total volume 140 mL) of nickel nitrate (Ni(NO3)2·6H2O from Mallinckrodt Baker and 0.5 N solution (total volume 140 mL) of sodium arsenate were mixed in stoichiometric proportions at 50–60°C. The precipitate was left standing in the original solution for a few days and then washed and dried at 40°C. 0.1 N (total volume 150 mL) solution of silver nitrate (AgNO3) obtained from Fisher Scientific and 0.1 N solution (total volume 150 mL) of sodium arsenate were mixed in stoichiometric amounts at 50–60°C. The precipitate was left standing in the original solution for a few days before washing and drying. 0.05 N solution (total volume 500 mL) of strontium chloride (SrCl2) from General Chemical Co. (New York, NY) and 0.05 N solution (total volume 600 mL) of sodium arsenate were mixed in stoichiometric amounts at 50–60°C. The precipitate was left standing in the original solution for a few days before washing and drying. 0.05 N solution (total volume 400 mL) of stannous chloride (SnCl2·2H2O) from Laboratory Chem, Inc. (Pittsburgh, PA) and 0.05 N solution (total volume 600 mL) of sodium arsenate were mixed in stoichiometric amounts at 50–60°C. The precipitate was left standing in the original solution for a few days and then washed and dried. The ion chromatographic system used a Waters IC-PAC Anion Column (4.6 × 150 mm) and a Wescan Cation Column (4.6 × 50 mm) to detect negatively and positively charged species respectively. Chromatographic identification of anions was performed under the following experimental conditions: mobile phase 0.85 mM NaHCO3/0.9 mM Na2CO3, flow rate 1.5 mL min–1, column temperature 25°C, and injection volume 50 µL. The following experimental conditions were employed in cation analysis: mobile phase 2 mM tartaric acid/1 mM oxalic acid, flow rate 1 mL min–1, column temperature 25°C, and injection volume 50 µL. Conductivity detector was used for both anion and cation analysis. The sensitivity, baseline stability, and resolving power were improved in anion analysis by using an ERIS autosuppressor.
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Appendix I: Recipes for metal arsenate synthesis (all syntheses yielded greater than 2.0 g of metal arsenate)
Cadmium arsenate (Cd3(AsO4)2)
Chromium arsenate (CrAsO4)
Copper arsenate (Cu3(AsO4)2)
Ferric arsenate (FeAsO4)
Lead arsenate (HPbAsO4)
Magnesium arsenate (Mg3(AsO4)2)
Manganese arsenate (MnHAsO4)
Nickel arsenate (NiHAsO4)
Silver arsenate (Ag3AsO4)
Strontium arsenate (Sr3(AsO4)2)
Tin arsenate (Sn3(AsO4)2)
Appendix II: Ion chromatography conditions
