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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

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Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases

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  1. 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

  2. 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)

  3. Reaction Rates • Activation Energy determines reaction rates • Small Ea = fast reaction • Large Ea = slow reaction • Rate Constants

  4. 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

  5. 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)

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