Lecture 17 - Open Systems (Solid/Liquid Equilibrium)

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Lecture 17 - Open Systems (Solid/Liquid
Equilibrium)

• Exam 3 (Dec 5 _at_ 130pm)
• Precipitation Dissolution
• Solubility product (Ks0)
• Estimating solubility
• From Ks0 - ignoring complex formation
• Effects of complexation
• pC-pH diagram (estimate max/min solubility)

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Example - Ferric iron solubility

• Fe(OH)3 Fe3 3OH- log Kso -38.8

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Ferric Iron Solubility
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Metal Solubility
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Metal Solubility - Effect of Complexes

• Metal cations exist in solution as species other
  than the aquo ion
• These complexes impact metal solubility
  substantially and must be accounted for
• Example - Ferric Iron

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Fe(III) Hydrolysis
(A) Fe3 OH- FeOH2 logK1 11.81 (B)
Fe3 2OH- Fe(OH)2 logb2 23.4 (C)
Fe3 4OH- Fe(OH)4- logb4
34.4 OR (A) Fe3 H2O FeOH2
H logK1 -2.19 (B) Fe3 2H2O
Fe(OH)2 2H logb2 -4.60 (C) Fe3
4H2O Fe(OH)4- 4H logb4 -21.6
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System for Fe(III) solubility hydroxo complexes

• System Open
• Components Fe3 H2O
• Species Fe3, FeOH2, Fe(OH)2, Fe(OH)3(s),
  Fe(OH)4-, H2O, OH-, H
• Equilibria
• H2O H OH- Kw HOH- 10-14
• (plus 4 reactions for Fe-hydoxo
  Fe-solubility)
• Mass Balance
• Don't need!
• Charge Balance
• Don't need!

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Combine solubility and complex formation
equilibria

• Fe3/Fe(OH)3(s)
• Fe(OH)3 3H Fe3 3H2O log Ks 3.20
• FeOH2/Fe(OH)3(s)
• Fe(OH)3 2H FeOH2 2H2O logKs1 1.01
• Fe(OH)2/Fe(OH)3(s)
• Fe(OH)3 H Fe(OH)2 H2O logKs2 -1.40
• Fe(OH)4-/Fe(OH)3(s)
• Fe(OH)3 H2O Fe(OH)4- H logKs4 -18.40

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Fe(III) - solubility
Fe(OH)3(s)
FeOH2
Fe(OH)4-
Fe3
Fe(OH)2
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Ferric Iron Solubility Predominance
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So What?

• Complexes increase mineral solubility
• Diagram gives pH of min or max solubility
• What are these for ferric iron?
• Many other ligands form solids (e.g., CO32-)
• For systems dominated by a particular solid we
  can estimate total metal concentration and
  dominant species.

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Carbonate System - Mixed Open System

• System Open
• Components Ca2, CO2(g) H2O
• Species H2CO3, HCO3-,CO32-, CO2(g), H2O, OH-,
  H, CaCO3, Ca2
• Equilibria
• H2O H OH- Kw HOH- 10-14
• H2CO3 H HCO3- KA1 HHCO3-/H2CO3
  10-6.35
• HCO3- H CO32- KA2 HCO32-/HCO3-
  10-10.33
• CO2(g) CO2(l) KH CO2(l)/PCO2 10-1.5
  (mol/L-atm)
• CaCO3 Ca2 CO32- Kso 10-8.48
• Mass Balance CT H2CO3 HCO3- CO32-
• CaT Ca2
• Charge Balance 2CO32- HCO3- OH-
  H 2Ca
• Proton Balance Invalid because CO2 is proton
  active!!

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What do we know?

• Our carbonate species in open systems
• Total Ca based on solid/liquid

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Final Equation

• Substitute in charge balance equation
• Assume water pH near neutral, thus H, OH-
  CO32- are negligible.
• Solve for H to get pH 8.24

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Importance of mineral dissolution/precipitation

• Natural Waters
• Dictate major cation concentrations (e.g., Ca2,
  Mg2)
• Dictate carbonate system (particularly in g.w.)
• Treatment systems
• Hardness removal
• Iron removal by aeration
• Phosphate removal
• Polluted waters
• Acid mine drainage results from pyrite dissolution

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Importance of mineral dissolution/precipitation

• Typically non-equilibrium processes (due to slow
  kinetics), but we can use equilibrium approaches
  because
• Often one solid phase controls solution
  composition
• The optimum pH for mineral precipitation dictated
  by equilibrium solubility constant

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Solubility

• Described like all other reactions
• AaBb(s) aA bB Ks
• Ks termed solubility product
• For uncomplexed ligand we get Ks
• CaCO3(s) Ca2 CO32- Kso
• For complexed ligand we use Ks
• CaCO3(s) H Ca2 HCO3- Ks

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Aluminum Hydroxide