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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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• 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

• 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

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]

• 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

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

• 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

• 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

• 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

• 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)