Lecture 3: Mineral Solubility
Transcript of Lecture 3: Mineral Solubility
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Lecture 3: Mineral Solubility
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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
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Solubility Determines Deep-Sea Sediment Types
from: Marine Chemistry by Schulz and Zabel
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Solubility Controls Contaminant Fate
McKinley et al. (2007) Vadose Zone J. Stubbs et al. (2009) Geochim. Cosmochim. Acta
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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
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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
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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
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“The Dolomite Problem”
Ordered Dolomite: Q/K = 1950, SI = 3.29Disordered Dolomite: Q/K = 550, SI = 2.74
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Ocean Acidification and Carbonate Mineral Saturation State
Hoegh-Guldberg et al. (2007) Science
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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
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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!!!!
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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
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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
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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
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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
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Isostructures that Form Solid Solutions
AragoniteCaCO3
StrontianiteSrCO3
CalciteCaCO3
OtaviteCdCO3
Corundumα-Al2O3
Hematiteα-Fe2O3
Diasporeα-AlOOH
Goethiteα-FeOOH
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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
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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)
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Effect of T on Solubility
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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
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Common Ion Effect in Groundwater: Fluorite Solubility