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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. Equilibrium Constants • The Law of Mass Action • This is an empirical law discovered in 1864 • Every reaction has a constant associated with it telling us where the equilibrium position is. jA + kB lC + mD • K = Equilibrium Constant = tells us where the equilibrium position is • K > 1 tells us the equilibrium lies to the right • K < 1 tells us the equilibrium lies to the left • If we know the concentrations, we can find K from its equation • K is written without units, even in cases where there are units left not cancelled. This is correct for nonideal behavior of molecules. • Sample Ex. 13.1 Write K for: 4NH3 + 7O2 4NO2 + 6H2O • Don’t include solvents, pure liquids or pure solids in the K equation K

  3. III. Kinetics A. Rate Laws a) Describe how fast a reaction occurs and how we can effect that speed • For simple organic reactions, we can directly write the rate law based on the stoichiometry of the reactants 1. NO2 + NO2 NO3 + NO rate = k[NO2]2 2. NO3 + CO NO2 + CO2rate = k[NO3][CO] • Other examples: A + A + B products rate = k[A]2[B] A + B + C products rate = k[A][B][C] A + B C k = a constant unique to each reaction [A], [B] = concentration of reactants (M) rate = k[A][B]

  4. IV. Bronsted-Lowery Model of Acids and Bases • Acid is an H+ donor • Base is an H+ acceptor • HCl + H2O H3O+ + Cl- 1) General Acid Equation HA + H2O H3O+ + A- • Conjugate base = what is left after H+ leaves acid • Conjugate acid = base + H+ • Conjugate acid-base pair are related by loss/gain of H+ d) Competition for H+ by A- and H2O; strongest base wins base hydronium ion acid conjugate acid conjugate base acid base

  5. 2) Ka = acid dissociation constant 3) Sample Exercise: Write simple Ionizations for: HCl, HC2H3O2, NH4+, C6H5NH3+, Al(H2O)63+ 4) Bronsted-Lowery theory allows for non-aqueous solutions NH3 + HCl NH4Cl

  6. Acid Strength • Acid strength describes the equilibrium position of the ionization reaction HA + H2O H3O+ + A- • Strong Acid = equilibrium lies far to the right (Large Ka) • Almost all HA has ionized to H+ and A- ([H+] = [HA]0) • A strong acid has a weak conjugate base • To ionize fully, the conjugate base must have low proton affinity • The conjugate base must be weaker that water • Weak Acid = equilibrium lies far to the left (Small Ka) • Almost all HA remains unionized ([H+] << [HA]0) • A weak acid has a strong conjugate base • The conjugate base is much stronger than water Strong acid Weak Acid

  7. Water as an Acid and Base • An amphoteric substance can behave as an acid or a base (water) • Autoionization of water (reaction with itself) H2O + H2O H3O+ + OH- • Ionization constant for water = KW = [H3O+][OH-] = [H+][OH-] • For any water solution at 25 oC, [OH-] x [H+] = KW = 1 x 10-14 • Neutral solutions (pure water) have [OH-] = [H+] = 1 x 10-7 • Acidic solutions: [H+] > [OH-] • Basic solutions: [OH-] > [H+] • Sample Ex. Calculate [OH-] or [H+] for the following: • [OH-] = 1 x 10-5 M • [OH-] = 1 x 10-7 M • [H+] = 10 M

  8. pH Scale • pH = -log[H+] (simplifies working with small numbers) • If [H+] = 1.0 x 10-7, pH = -log(1 x 10-7) = -(-7.00) = 7.00 3) pOH = -log[OH-] pKa = -logKa 4) pH changes by 1 unit for every power of 10 change in [H+] • pH = 3 [H+] = 10 times the [H+] at pH = 4 • pH decreases as [H+] increases (pH = 2 more acidic than pH = 3) • Meaning of pKa HA + H2O H3O+ + A- The lower the pKa, the stronger the acid

  9. 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 = Electrophile • Lewis Base = electron pair donor = Nucleophile • Some covalently bonded molecules can be considered Lewis Acid/Base pairs • Dissociation of a Lewis Acid/Base Pair (Mechanisms)

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