Chapter 2 lecture 1 kinetics thermo acids bases
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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 l.jpg
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


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  • 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


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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]


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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


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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


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  • 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


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  • 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


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  • 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


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  • 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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