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Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases. Review of Simple Kinetics and Thermodynamics Definitions Thermodynamics = changes in energy during a process or reaction. Determines extent of completion of the reaction or process

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chapter 2 lecture 1 kinetics thermo acids bases
Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases
  • Review of Simple Kinetics and Thermodynamics
    • Definitions
      • Thermodynamics = changes in energy during a process or reaction. Determines extent of completion of the reaction or process
      • Kinetics = rate of a process or reaction. Determines how fast the reaction or process occurs.
    • Equilibria
      • Equilibrium = state of a system in which the concentrations of reactants and products are no longer changing.
      • Equilibrium Constant
        • If K is large, reaction goes forward
        • If K is small, reaction goes in reverse
slide2
Relating Gibb’s Free Energy Change to Equilibrium Constants
    • DG0 = Gibb’s Free Energy Change = describes the overall energy change as a reaction reaches equilibrium
    • DG0 = -RTlnK

R = 1.986 cal/deg mol

T = temperature in Kelvins (oC + 273)

    • When K = 10, DG0 = -1.36 kcal/mol (at T = 298K)
    • When K = 0.1, DG0 = +1.36 kcal/mol
    • When K = 1, DG0 = 0
  • Relating DG0 to Enthalpy and Entropy
    • DG0 = DH0 - TDS0
    • DH0 = Enthalpy = Broken Bond Strengths – Formed Bond Strengths
      • -DH0 = Exothermic reaction (gives off energy)
      • + DH0 = Endothermic reaction (requires energy input)
    • DS0 = Entropy = Amount of order in the system
      • - DS0 = less disorder (fewer molecules in the system)
      • + DS0 = more disorder (more molecules in the system)
slide3
Reaction Rates
    • Activation Energy determines reaction rates
      • Small Ea = fast reaction
      • Large Ea = slow reaction
    • Rate Constants
slide4
The Arrhenius Equation

A = maximum rate constant possible = different for each reaction

High T -Ea/RT becomes small e0 = 1 k = A

  • Review of Acids and Bases
    • Bronsted Acids/Bases
      • Acid = H+ donor
      • Base = H+ acceptor
      • Ionization of Water: H2O + H2O H3O+ + OH-
      • pH = -log[H3O+]
      • pKa = -logKa = pH at which HA is half-dissociated
      • pKa + pKb = 14

If you know Ka, Kb, pKa, pKb, you can find all others

Ka

slide5
Predicting Acid/Base Strength
    • Size of A-: HI > HBr > HCl > HF
      • F- is small, more concentrated charge, holds on to H+
      • I- is large, less concentrated charge, gives up H+
    • Electronegativity of A-: HF > H2O > NH3 > CH4
    • Resonance Forms of A-
  • Lewis Acids and Bases
    • Lewis Acid = electron pair acceptor
    • Lewis Base = electron pair donor
    • Some covalently bonded molecules can be considered Lewis Acid/Base pairs
    • Dissociation of a Lewis Acid/Base Pair (Mechanisms)