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This document presents a critical evaluation of the equilibrium constants and reaction enthalpies for the complex formation reactions between aqueous Hg(II) and the common environmental inorganic ligands Cl–, OH–, CO32–, SO42–, and PO43–. The analysis used data from the IUPAC Stability Constants database, SC-Database, focusing particularly on values for 25 °C and perchlorate media. Specific ion interaction theory (SIT) was applied to reliable data available for the ionic strength range Ic < 3.0 mol dm–3. Recommended values of log10βp,q,r° and the associated reaction enthalpies, ∆rHm°, valid at Im = 0 mol kg–1 and 25 °C, were obtained by weighted linear regression using the SIT equations. Also reported are the equations and specific ion interaction coefficients required to calculate log10βp,q,r° values at higher ionic strengths and other temperatures. A similar analysis is reported for the reactions of H+ with CO32– and PO43–. Diagrams are presented to show the calculated distribution of Hg(II) amongst these inorganic ligands in model natural waters. Under typical environmental conditions, Hg(II) speciation is dominated by the formation of HgCl2(aq), Hg(OH)Cl(aq), and Hg(OH)2(aq).

Complex formation between Pb II and the common environmental inorganic ligands, Cl -, OH -, CO 3 2-, SO 4 2-, and PO 4 3-, can be significant in natural waters with low concentrations of organic matter. Numerical modeling of the speciation of Pb II amongst these inorganic ligands requires reliable values for the relevant stability (formation) constants. This paper provides a critical review of such constants and related thermodynamic data. It recommends values of log 10 β p,q,r°v alid at I m = 0 mol kg -1 and 25 °C (298.15 K), along with the equations and empirical coefficients required to calculate log 10 β p,q,r values at higher ionic strengths using the Brønsted-Guggenheim-Scatchard specific ion interaction theory (SIT). Some values for reaction enthalpies, Δ r H, are also reported.In weakly acidic fresh water systems (-log 10 {[H + ]/c°} < 6), the speciation of Pb II is similar to that of Cu II . In the absence of organic ligands, Pb II speciation is dominated by Pb 2+ (aq), with PbSO 4 (aq) as a minor species. In weakly alkaline solutions, 8.0 < -log 10 {[H + ]/c°} < 9.0, the speciation is dominated by the carbonato species PbCO 3 (aq) and Pb(CO 3 ) 2 2-. In weakly acidic saline systems (-log 10 {[H + ]/c°} < 6), the speciation is dominated by PbCl n (2-n)+ complexes, (n = 0-3), with Pb 2+ (aq) as a minor species. In this medium (and in seawater), the speciation contrasts with that of Cu II because of the higher stability of the Pb 2+ -chlorido-complexes. In seawater at -log 10 {[H + ]/c°} = 8.2, the calculated speciation is less well defined, although it is clearly dominated by the uncharged species PbCO 3 (aq) (41 % of [Pb] T ) with a significant contribution (16 %) from Pb(CO 3 )Cl -and minor contributions (5-10 %) from PbCl n (2-n)+ , (n = 0-3) and Pb(CO 3 ) 2 2-. The uncertainty in calculations of Pb II speciation in seawater arises from (a) the large uncertainty in the stability constant for the apparently dominant species PbCO 3 (aq), (b) the reliance on statistical predictions for stability constants of the ternary species Pb(CO 3 )Cl -and Pb(CO 3 )OH -, and (c) the uncertainty in the stability constant for PbCl 4 2-, the available value being considered

Complex formation between CuII and the common environmental ligands Cl-, OH-, CO32-, SO42-, and PO43- can have a significant effect on CuII speciation in natural waters with low concentrations of organic matter. Copper(II) complexes are labile, so the CuII distribution amongst these inorganic ligands can be estimated by numerical modeling if reliable values for the relevant stability (formation) constants are available. This paper provides a critical review of such constants and related thermodynamic data. It recommends values of log10βp,q,r° valid at Im = 0 mol kg-1 and 25 °C (298.15 K), along with the equations and specific ion interaction coefficients required to calculate log10βp,q,r values at higher ionic strengths. Some values for reaction enthalpies, ΔrHm, are also reported where available. In weakly acidic fresh water systems, in the absence of organic ligands, CuII speciation is dominated by the species Cu2+(aq), with CuSO4(aq) as a minor species. In seawater, it is dominated by CuCO3(aq), with Cu(OH)+, Cu2+(aq), CuCl+, Cu(CO3)OH-, Cu(OH)2(aq), and Cu(CO3)22- as minor species. In weakly acidic saline systems, it is dominated by Cu2+(aq) and CuCl+, with CuSO4(aq) and CuCl2(aq) as minor species.

Equilibrium constants are reported for the aluminum(III) complexes with four ligands that contain the 1,2-dihydroxyaryl group: trans-3-(3′,4′-dihydroxyphenyl)propenoic acid (caffeic acid), 3-(3′,4′-dihydroxycinnamoyl)-1,3,4,5-tetrahydroxycyclohexanecarboxylic acid (chlorogenic acid), 3,4-dihydroxybenzylamine (DHB), and 1,2-dihydroxyanthraquinone-3-sulfonic acid (DASA). The protonation and complexation reactions were studied by potentiometric and spectrophotometric titrations in aqueous solution at 25 °C (I ) 0.10 M KCl). The complexation reactions were studied for ligand/metal ratios in the range from 1.3 to 5.9. For caffeic and chlorogenic acids (H 3 L), the experimental data were consistent with the formation of six mononuclear Al-ligand species, each involving catecholate coordination: AlHL + , AlL, Al(OH)L -, AlL 2 3-, Al(OH)L 2 4-, and AlL 3 6-. For DHB (H 2 L), the data were consistent with the formation of six monomeric and one dimeric Al species: AlHL 2+ , Al(HL) 2 + , Al(HL) 3 0 , Al(HL) 2 L -, Al(HL)L 2 2-, AlL 3 3-, and (OH) 2 Al 2 (HL) 2 2+ , in which HL has a protonated amino group. On the basis of the DASA (H 2 L -) titration data, three monomeric and one dimeric Al-ligand species were proposed: AlL 0 , AlL 2 3-, Al(OH)L 2 4-, and Al 2 (OH) 2 L 4 8-. The relative stability of the 1:1 complexes, AlL, was DASA > chlorogenic acid > caffeic acid > DHB. This indicates the progressive electron-withdrawing effects of the ring substituents -SO 3 -< -CHdCH‚CO‚O‚(C 6 H 10 O 3 )‚COOH < -CHdCH‚COOH < -NH 3 + .