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Aqueous Environmental Geochemistry by aR8wG9

VIEWS: 5 PAGES: 12

									                                Fall 2007




Aqueous Environmental
    Geochemistry

          Brooklyn College
   The City University of New York
Natural Systems
    Open and Close Natural Systems

•   Natural systems can be large or small
•   Different phases
•   In relative sense
•   Reaction rates vs. Flux rates
         Gibbs Phase Rule
• F=C–P+2
  What do C and P mean?

• Example:
    H2O (ice) = H2O (liquid) = H2O (vapor)

• F = C’-R – P + 2:
     What are C’ and R? Give one example.
                Ehthalpy (H)
•   Heat content at constant pressure
•   Heat of formation from elements
•   ΔHf0 for (most stable form) element is zero
•   ΔHr0 : Heat exchange
       ΔHr0 : H>0, endothermic
       ΔHr0 H<0, exothermic
• How T & P change affect the equilibrium of
  a reaction?
                Entropy (S)
• Degree of randomness or disorder
                      T   Cp
       ST  S 0            dT
                     0    T
• Cp: heat capacity. Energy needed to raise the
  temperature of 1 mole by 1 K
• At 0 K, most substances S=0
• S increases as temperature increases
      Gibbs Free Energy (G)
• At equilibrium condition:
           ΔGr0 = - RT ln Keq
           R: gas constant, 1.9872 cal/mol K
           T: 25 ºC, or 298.15 K
• Can be transformed into:
           ΔGr0 = -0.59248 ln Keq
                = -1.3642 log Keq
     Direction of Reaction

        ΔGr = ΔGr0 + RT ln Q
• Q: reaction quotient
• Sign of ΔGr indicate the direction of reaction
         Chemical Potential (µ)
• Similar to Gibbs Free Energy
• µi = µi0 + RT ln(αi)
  - Chemical potential is for individual component

  - αi is activity, which is different than concentration!

• At equilibrium, is the same for the same
  component in different phases!
                         Solutions
• In solutions
               (αi) = λi · Ni
  -λ   is the rational activity coefficient
  - for ideal solutions, λ ≈ 1 for major components (Raoult’s Law);
  - for ideal solutions, λ ≠ 1 for minor components or dilute solutions
  (Henry’s Law)
  - See Fig. 1.3 for differences
     Solid Solutions (binary mixing)
• Only equations to remember:
     ln λA = α0 NB2
     ln λB = α0 NA2

 Can be treated as exchange reaction:

               KAX   AmA    BNB 
         Kex                    
               KEX   BmB    ANA 
              Keq Dependence on T
•   Empirical constants
•   Determined in experiments
•   Table 1.1 (P39-45)
•   Equation:

                                        F
    log Keq  A  BT   D log T  ET 
                       C
                       T
                                      2
                                            GT 1/ 2

                                        T2
• Calculate ΔGr0 , ΔHr0 ΔCpr0 ΔSr0
• ΔHr0 controls the dependence.

								
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