# Dissolution and Precipitation

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Dissolution and Precipitation
Lecture 23

Carbonate Solubility Carbonate is the most common kind of chemical sediment and carbonate components (Ca, Mg, CO3) are often the dominant species in natural waters. Using equilibrium constants, we can calculate calcite solubility as a function of PCO2: Thus calcite solubility increases with 1/3 power of PCO2. 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.

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

Carbonate Solubility Open system solutions, those in equilibrium with CO2 gas in the atmosphere or soil, can dissolve more calcite than closed systems waters. In a sense, this is because dissolving CO2 in water lowers pH, resulting in greater dissolution.

Mg Solubility ΣCO2 = M Several Mg-bearing minerals can precipitate from solution: brucite, Mg(OH)2 dolomite CaMg(CO3)2 magnesite MgCO3 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 [Mg2+] on one side of the equation and pH on the other. We simplify things by calculating equilibrium with only 1 carbonate species at a time and ignoring the others.

Mg solubility as a function of CO2 & pH
[Mg2+] = 10-4 M

Mg solubility as a function of CO2 & Ca/Mg
Mg2+ = 10-4 M

Constructing stability diagrams
ΣCO2 = 5x 10-2 M This diagram shows the stability of ferrous iron minerals as a function of pH and sulfide for fixed total Fe and CO2. Procedure: manipulate equilibrium constant expressions to obtain and expression for ΣS in terms of pH. For example: FeCO3 + H+ ⇋ Fe2+ + HCO3– FeCO3 + 2H2O ⇋ HCO3- H+ + Fe(OH)2 FeS + 2H2O ⇋ Fe(OH)2 + H+ + HS- Trick: simplify by ignoring species present at low conc. (e.g., CO32- at low pH). [Fe2+] = 10-6 M (Pyrrhotite) (Siderite)

Solubility of SiO2 Silica forms silicic acid (H4SiO4) in solution, which can then dissociate through a series of reactions, e.g., H4SiO4 ⇋ H3SiO4– + H K1 H3SiO4– ⇋ H2SiO42- + H+ K2 Solubility can be expressed as: where [H4SiO4] is controlled by solubility of either quartz or amorphous silica. As a consequence, its solubility is a function of pH: high only at high pH.

Solubility of Hydroxides
The hydroxide is the least soluble salt of many metals. Therefore, it is the solubility of their hydroxides that controls the solubility of these metals in natural waters. Since these dissolution reactions involve OH–, they are pH-dependent, and the slope of the solubility curve depends on the valence of the metal (e.g., -3 for Fe3+, -2 for Fe2+, -1 for Ag+).

Solubility of Al(OH)3 Solubility of gibbsite:
Al(OH)3 + 3H+ ⇋ Al3+ + H2O However, Al forms hydroxide complexes, e.g.: Al3+ + H2O ⇋ Al(OH)2+ + H+ The total dissolved Al will be the sum of all Al species in solution: A consequence of this is that acid rain leads to Al poisoning.

Solubility of Ferric Iron

Silicate Solubility We’ve looked at the solubility of Si, Al, Fe and other cations in the isolation of simple laboratory-like systems. The real world is usually more complex. Silicate rocks predominate at the surface of the Earth, thus Si, Al, and other cations will generally all be present. In addition to gibbsite and SiO2, some of the more common weathering products of silicate rocks include: gibbsite: Al(OH)3 kaolinite: Al2Si2O5(OH)4 pyrophyllite: Al2Si4O10(OH)2 illite (muscovite): KAl3Si3O10(OH)2

Silicate Solubility Gibbsite will precipitate from Al-bearing solution only at lowest concentrations of SiO2. Occurrence generally restricted to highly weathered soils where all the SiO2 has washed out.

Silicate Solubility

Clay Minerals Review section 6.5 to become familiar with clay minerals, but we will not cover it in class.