lecture 3: mineral solubility
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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
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