Lecture 3: Mineral Solubility

Solubility Controls Biomineralization

• Organisms produce hard parts by exceeding the solubility of the mineral component

• Increased CO2 in the oceans increases carbonate mineral solubility, making biomineralizationmore difficult

De Yoreo and Dove (2004) Science 306, 1301

Solubility Determines Deep-Sea Sediment Types

from: Marine Chemistry by Schulz and Zabel

Solubility Controls Contaminant Fate

McKinley et al. (2007) Vadose Zone J. Stubbs et al. (2009) Geochim. Cosmochim. Acta

Solubility Affects “Soil” Development on Mars

Layered ferric sulfate, calcium sulfate, and iron oxide, Columbia Hills, Mars

Subsurface perchlorate salts, a sign of soil water transport,

Northern Plains, Mars

Solubility Thermodynamics

• Equilibrium constant for mineral solubility is called the solubility product, Ksp– Convention is to write as a dissolution reaction

• Example: GypsumCaSO4·2H2O(s) = Ca2+ + SO4

2- + 2H2O Ksp = [Ca2+][SO4

2-] log Ksp = -4.58– [CaSO4·2H2O] and [H2O] assumed to equal 1

• We’ll address when these assumption do not hold later today and in the next lecture

Evaluating the Saturation State of Natural Waters

• For minerals we use the saturation index (SI) to evaluate saturation state

SI = log (Q/K)– SI = 0 when solution is saturated

• Mineral in equilibrium with solution– SI < 0 when solution is undersaturated

• Mineral, if present, should dissolve– SI > 0 when solution is supersaturated

• Mineral should precipitate

• Sometimes see SI = Q/K, or Ω = Q/K– “Ion Activity Product (IAP)” = Q– We don’t use these!!!

Barite Q/K in the Central Pacific OCean

“The Dolomite Problem”

Ordered Dolomite: Q/K = 1950, SI = 3.29Disordered Dolomite: Q/K = 550, SI = 2.74

Ocean Acidification and Carbonate Mineral Saturation State

Hoegh-Guldberg et al. (2007) Science

Silica Saturation States of Natural Waters

Type of Water Range (ppm) SIQuartz SIAm. Silica

Rivers/Lakes 5-25 -0.1 to 0.6 -1.4 to -0.7Seawater 0.01-7 -2.8 to 0.05 -4.1 to -1.2Soil Porewater 1-117 -0.8 to 1.3 -2.1 to 0.0Groundwater 5-85 -0.1 to 1.1 -1.4 to -0.1Oil Field Brine 5-60 -0.1 to 1.0 -1.4 to -0.3Hot Springs* 100-600 1.2 to 2.0 -0.1 to 0.7

* SI values calculated for 25°C. Hot spring temperatures range as high as 100°C at the surface, and SIquartz ~ 0 at spring water temperature for many hot springs. When water cools, silica precipitates, forming sinter.

SiO2(s) + 2H2O = H4SiO4(aq)Ksp = [H4SiO4]

log Ksp,Quartz = -4.0 log Ksp,Am. Silica = -2.7

Mineral Solubility and Stability• Thermodynamics predicts that the lowest energy

state should occur– This state is said to be thermodynamically stable

• A solution that is supersaturated is not stable– The saturated mineral phase should precipitate

• However, sometimes mineral precipitation is kinetically-inhibited– Metastable phases often form instead– Solutions must be supersaturated with respect to a

metastable phase for it to precipitate• Metastable phases are always more soluble

than stable phases!!!!

Silica Saturation States of Natural Waters

Type of Water Range (ppm) SIQuartz SIAm. Silica

Rivers/Lakes 5-25 -0.1 to 0.6 -1.4 to -0.7Seawater 0.01-7 -2.8 to 0.05 -4.1 to -1.2Soil Porewater 1-117 -0.8 to 1.3 -2.1 to 0.0Groundwater 5-85 -0.1 to 1.1 -1.4 to -0.1Oil Field Brine 5-60 -0.1 to 1.0 -1.4 to -0.3Hot Springs* 100-600 1.2 to 2.0 -0.1 to 0.7

* SI values calculated for 25°C. Hot spring temperatures range as high as 100°C at the surface, and SIquartz ~ 0 at spring water temperature for many hot springs. When water cools, silica precipitates, forming sinter.

SiO2(s) + 2H2O = H4SiO4(aq)Ksp = [H4SiO4]

log Ksp,Quartz = -4.0 log Ksp,Am. Silica = -2.7

Stable Versus Metastable• Carbonates

– STABLE: Calcite [CaCO3], Dolomite [CaMg(CO3)2]– METASTABLE: Aragonite [CaCO3], Mg-calcite*

• Sulfides– STABLE: Pyrite [FeS2]– METASTABLE: Mackinawite [Nanocrystalline FeS]

• Iron oxides– STABLE: Hematite [Fe2O3], Magnetite [Fe3O4]– METASTABLE: Ferrihydrite [Fe(OH)3]– Goethite [FeOOH] is metastable with respect to

hematite at 25° but stable below ~15°C

Stable Versus Metastable• Clays and Zeolites

– Talc, Muscovite stable– Kaolinite*, Illite, Smectite, Clinoptilolite metastable

• Clay stability difficult to assess– Rarely occur as simple phases with definitive

compositions– Difficult to measure thermodynamic properties

• Metastable phases may interconvert:– Smectite to Illite in marine sediments

Solid Solutions

• A solid solution occurs when an element substitutes into a mineral, and the substituting element can occur on its own in an isostructural phase

• Terminology needed:– Isostructures: Two minerals of different

composition but same structure– Polymorphs: Two minerals of the same

composition but different structures

Isostructures that Form Solid Solutions

AragoniteCaCO3

StrontianiteSrCO3

CalciteCaCO3

OtaviteCdCO3

Corundumα-Al2O3

Hematiteα-Fe2O3

Diasporeα-AlOOH

Goethiteα-FeOOH

Solubility of Ideal Solid Solutions

• Consider the CaCO3-CdCO3 solid solutionCaCO3 = Ca2+ + CO3

2- log Ksp = -8.48CdCO3 = Cd2+ + CO3

2- log Ksp = -12.1• For a normal solid, we set concentration of

mineral equal to 1 in the equilibrium equation• For a solid solution, these are set to the mole

fraction:

[ ] Cd

SS

XCaCd

CdCdCOmolmol

mol==

+

3

See Section 4.5 of Textbook for more details and examples

Effect of T on Solubility• The van’t Hoff equation demonstrates how

T affects an equilibrium constant:log K2 – log K1 = (ΔHr

o/2.303R)(1/T1 – 1/T2)– If ΔHr

o < 0, mineral solubility decreases with increasing T

– If ΔHro > 0, mineral solubility increases with

increasing T• Examples: ΔHr

o(Calcite) = -10.6 kJ/mol; ΔHr

o(Quartz) = 25.1The example here refers to mineral solubility, but the van’t Hoff equation applies

to any equilibrium constant. (Hint: This applies to gas solubility on PS1)

Effect of T on Solubility

Common Ion Effect

• Solutes in real systems often have more than one origin

• This leads to the common ion effect– Predicted concentration of a solute in complex

systems differs from in simple systems– This makes the solid less soluble than would

be expected for single-phase system• Example: Effect of Na2SO4 on Gypsum

solubility

Common Ion Effect in Groundwater: Fluorite Solubility