1. Lecture 21

2. Water-Related Complexes Ferric iron, will form a Fe(H 2 O) 6 3+ aquo-complex. The positive charge of the central ion tends to repel hydrogens in the water molecules, so that water molecules in these aquo-complexes are more readily hydrolyzed than otherwise. Thus these aquo- complexes can act as weak acids. For example: Fe(H 2 O) 6 3+ ⇄ Fe(H 2 O) 5 (OH) 2+ + H+ ⇄ Fe(H 2 O) 4 (OH) 2 + + 2H + ⇄ Fe(H 2 O) 3 (OH) 3 + 3H + ⇄ Fe(H 2 O) 2 (OH) 4 - + 4H + Thus equilibrium between these depends strongly upon pH.

3. Hydroxo- and Oxo-Complexes Loss of hydrogens from the solvation shell results in hydroxo-complexes. The repulsion between the central metal ion and protons in water molecules of the solvation shell will increase with decreasing diameter of the central ion and with increasing charge of the central ion. Not surprisingly, the complexes also depend on the abundance of H + and OH – ions, i.e., pH. For highly charged species, the repulsion of the central ion is sufficiently strong that all hydrogens are repelled and it is surrounded only by oxygens. Such complexes, for example, and, are known as oxo-complexes. Oxo-complexes are generally more soluble than hydroxo-complexes, so 6+ ions of a metal can be more soluble than lesser charged species (e.g., U). Intermediate types in which the central ion is surrounded, or coordinated, by both oxygens and hydroxyls are also possible, for example MnO 3 (OH) and CrO 3 (OH) –, and are known as hydroxo-oxo complexes.

4. Summary of Water Complexes For most natural waters, metals in valence states I and II will be present as “free ions”, that is, aquo- complexes, valence III metals will be present as aquo- and hydroxo- complexes (depending on pH), and those with higher charge will be present as oxo- complexes.

5. Polynuclear Complexes Polynuclear hydroxo- and oxo-complexes, containing two or more metal ions, are also possible. The extent to which such polymeric species form increases with increasing metal ion concentration. Most highly-charged metal ions (3+ through 5+) are highly insoluble in aqueous solution. This is due in part to the readiness with which they form hydroxo- complexes, which can in turn be related to the dissociation of surrounding water molecules as a result of their high charge. When such ions are present at high concentration, formation of polymeric species such as those above quickly follows formation of the hydroxo- complex. At sufficient concentration, formation of these polymeric species leads to the formation of colloids and ultimately to precipitation. In this sense, these polymeric species can be viewed as intermediate products of precipitation reactions.

6. Cu Hydroxides Interestingly enough, however, the tendency of metal ions to hydrolyze decreases with concentration. The reason for this is the effect of the dissociation reaction on pH. For example, increasing the concentration of dissolved copper decreases the pH, which in turn tends to drive the hydrolysis reaction to the left. To understand this, consider the following reaction: Cu 2+ + H 2 O ⇄ CuOH + + H + for which the apparent equilibrium constant is K app = 10 –8. We can express the fraction of copper present as CuOH +, α CuOH+ as: For a fixed amount of dissolved Cu, we can also write a proton balance equation: and a mass balance equation. Combining these with the equilibrium constant expression, we can calculate both α and pH as a function of Cu T

7. Other Complexes When non-metals are present in solution, as they would inevitably be in natural waters, then other complexes are possible. We can divide the elements into four classes: o Ligand Formers: Non-metals, which form anions or anion groups. o “A-type” or “hard” metals They can be viewed as hard, charged spheres. They preferentially complex with fluorine and ligands having oxygen as the donor atoms (e.g., OH –,CO 3 2-,SO 4 2- ). o B-type, or “soft”, metals. Their electron sheaths are readily deformed by the electrical fields of other. They preferentially form complexes with bases having S, I, Br, Cl, or N (such as ammonia; not nitrate) as the donor atom. Bonding is primarily covalent and is comparatively strong. Thus Pb forms strong complexes with Cl – and S 2–. o First transition series metals. Their electron sheaths are not spherically symmetric, but they are not so readily polarizable as the B-type metals. On the whole, however, their complex-forming behavior is similar to that of the B-type metals.

8. Irving-Williams Series the Irving-Williams series is the sequence of complex stability among first transition metals: Mn 2+ Zn 2+. In the figure, all the sulfate complexes have approximately the same stability, a reflection of the predominantly electrostatic bonding between sulfate and metal. Pronounced differences are observed for organic ligands. The figure demonstrates an interesting feature of organic ligands: although the absolute value of stability complexes varies from ligand to ligand, the relative affinity of ligands having the same donor atom for these metals is always similar.

9. Complexation Computations Where only one metal is involved, the complexation calculations are straightforward, as in Example 6.7. Natural waters, however, contain many ions. The most abundant of these are Na +, K +, Mg 2+, Ca 2+, Cl –, SO 4 2-, HCO 3 –, CO 3 2-, and there are many possible complexes between them as well as with H + and OH –. To calculate the speciation state of such solutions, an iterative approach is required, such as Example 6.08. Most major ions are not complexed in most situations. free ionOH – HCO 3 - CO 3 2- SO 4 2- free ion 1x10 -06 9.12x10 -04 1.12x10 -06 1.65x10 -04 H+H+ 1x10 -08 —2.06x10 -05 9.12x10 -04 1.58x10 -10 Na + 3.03x10 -04 —1.57x10 -07 2.35x10 -08 5.64x10 -07 K+K+ 5.69x10 -05 ———8.40x10 -08 Mg 2+ 1.40x10 -04 5.03x10 -08 1.84x10 -06 1.45x10 -06 5.14x10 -06 Ca 2+ 2.80x10 -04 3.92x10 -09 4.64x10 -06 1.83x10 -06 9.16×10 -06 Na + 99.76%Cl – 100% K+K+ 99.85%SO 4 2- 91.7% Mg 2+ 94.29%HCO 3 – 99.3% Ca 2+ 94.71%CO 3 2- 25.3% Percent Free Ion in Stream Water

10. Dissolution and Precipitation

11. Carbonate Solubility Carbonate is the most common kind of chemical sediment and carbonate components (Ca, Mg, CO 3 ) are often the dominant species in natural waters. Using equilibrium constants, we can calculate calcite solubility as a function of P CO2 : Thus calcite solubility increases with 1/3 power of P CO2. One consequence is that calcite shells tend to dissolve in deep ocean water. A second is that calcite will dissolve out of soils when microbial activity is present.

13. Carbonate Solubility Another interesting feature is because of this non-linearity, mixing of two saturated waters can produce an undersaturated water.

14. Carbonate Solubility Open system solutions, those in equilibrium with CO 2 gas in the atmosphere or soil (in red), can dissolve more calcite than closed systems waters (in black). In a sense, this is because dissolving CO 2 in water lowers pH, resulting in greater dissolution.

15. Mg Solubility Several Mg-bearing minerals can precipitate from solution: o brucite, Mg(OH) 2 o dolomite CaMg(CO 3 ) 2 o magnesite MgCO 3 We can use equilibrium constant expressions, such as: to construct a predominance diagram showing which phase will precipitate under a given set of conditions. We construct these in the same way we constructed pe-pH diagrams, namely manipulate the log equilibrium constant expressions to get [Mg 2+ ] on one side of the equation and pH on the other. o We simplify things by calculating equilibrium with only 1 carbonate species at a time and ignoring the others. ΣCO 2 = 10 -2.5 M

16. Mg Solubility For magnesite: We can think of three limiting cases: where carbonic acid, bicarbonate ion, and carbonate ion predominate. In the latter case, [CO 3 2- ] ≈ ΣCO 2. When bicarbonate ion predominates, [HCO 3 - ] ≈ ΣCO 2, and the carbonate ion concentration is: log [CO 3 2– ] = pK 2 + log [HCO 3 – ] +pH = - 12.88+pH What about when H 2 CO 3 dominates? ΣCO 2 = 10 -2.5 M

17. Mg solubility as a function of CO 2 & pH [Mg 2+ ] = 10 -4 M

18. Mg solubility as a function of CO 2 & Ca/Mg Mg 2+ = 10 -4 M

19. Constructing stability diagrams This diagram shows the stability of ferrous iron minerals as a function of pH and sulfide for fixed total Fe and CO 2. Procedure: manipulate equilibrium constant expressions to obtain and expression for ΣS in terms of pH. For example: o FeCO 3 + H + ⇋ Fe 2+ + HCO 3 – o FeCO 3 + 2H 2 O ⇋ HCO 3 - H + + Fe(OH) 2 o FeS + 2H 2 O ⇋ Fe(OH) 2 + H + + HS - Trick: simplify by ignoring species present at low conc. (e.g., CO 3 2- at low pH). ΣCO 2 = 5x 10 -2 M ➄ ③ (Pyrrhotite) (Siderite) [Fe 2+ ] = 10 -6 M ②