ZnO solubility and Zn2+ complexation by chloride and sulfate in acidic solutions to 290°C with in-situ pH measurement

The solubility of zincite in mildly to strongly acidic aqueous solutions, according to the reaction ZnO + 2H(+) double left right arrow Zn2+ + H2O, has been measured at ionic strengths of 0.03-1.0 (stoichiometric molal basis) from 50 to 290 degrees C at saturation vapor pressure in sodium triffuoromethanesulfonate solutions (NaTriflate, a noncomplexing 1:1 electrolyte). The hydrogen-electrode concentration cells employed in this study permit continuous and highly accurate pH measurement at elevated temperatures, and periodic sampling to determine the dissolved metal content of the experimental solution. The solubility of zincite is shown to be reversible at 200 degrees C by addition of acidic and basic titrants, at constant ionic strength. The equilibrium constant is precisely described (+/-0.05 log units) by the function log K = -4.0168 + 4527.66/T. One additional adjustable parameter, together with an extended Debye-Huckel function, is sufficient to model the ionic strength dependence of the reaction. The solubility product at infinite dilution obtained from this study is in quantitative agreement with the thermodynamic model of Ziemniak (1992). This experimental approach is demonstrated to be advantageous in studying the complexation of Zn2+ with Cl- and SO42-, by titrations involving the appropriate anion into NaTriflate solutions pre-equilibrated with zincite at constant temperature and ionic strength. Formation constants in 0.1 molal NaTriflate for the reaction Zn2+ + yL(z-) double left right arrow Zn(L)(y)(2-yz) are reported for ZnCl+, ZnCl2 degrees and ZnSO4 degrees at 200 degrees C (log Q = 1.7 +/- 0.1, 3.0 +/- 0.1, and 2.6 +/- 0.1, respectively). Estimates of the equilibrium constants for the chloride species at infinite dilution and 200 degrees C are log K = 2.5 +/- 0.1 (ZnCl+), and 4.2 +/- 0.1 (ZnCl2 degrees). This value for the dichlorozinc complex agrees quantitatively with values reported by Bourcier and Barnes (1987) and Ruaya and Seward (1986). However, the latter authors give a value for the monochlorozinc complex (log K = 4.01 +/- 0.02) that is markedly different from our result and that of Bourcier and Barnes (1987) (log K = 3.1 +/- 0.3). Copyrights (C) 1998 Elsevier Science Ltd.