Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases

I. Review of Simple Kinetics and ThermodynamicsA. Definitions

1) Thermodynamics = changes in energy during a process or reaction. Determines extent of completion of the reaction or process

2) Kinetics = rate of a process or reaction. Determines how fast the reaction or process occurs.

B. Equilibria1) Equilibrium = state of a system in which the concentrations of reactants

and products are no longer changing.

2) Equilibrium Constanta) If K is large, reaction goes forwardb) If K is small, reaction goes in reverse

D C B A K [A][B][C][D]K

II. Equilibrium ConstantsA. The Law of Mass Action

1. This is an empirical law discovered in 18642. Every reaction has a constant associated with it telling us where the

equilibrium position is.

jA + kB lC + mD

3. K = Equilibrium Constant = tells us where the equilibrium position isa) K > 1 tells us the equilibrium lies to the rightb) K < 1 tells us the equilibrium lies to the left

4. If we know the concentrations, we can find K from its equation

5. K is written without units, even in cases where there are units left not cancelled. This is correct for nonideal behavior of molecules.

6. Sample Ex. 13.1 Write K for: 4NH3 + 7O2 4NO2 + 6H2O

7. Don’t include solvents, pure liquids or pure solids in the K equation

kj

ml

[B][A][D][C]K K

III. KineticsA. Rate Laws

a) Describe how fast a reaction occurs and how we can effect that speed

1. 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 + CO2 rate = k[NO3][CO]

c) Other examples:A + A + B products rate = k[A]2[B]A + B + C productsrate = k[A][B][C]

rate = k[A][B]

A + B C k = a constant unique to each reaction[A], [B] = concentration of reactants (M)

IV. Bronsted-Lowery Model of Acids and Basesa) Acid is an H+ donorb) Base is an H+ acceptorc) HCl + H2O H3O+ + Cl-

1) General Acid Equation

HA + H2O H3O+ + A-

a) Conjugate base = what is left after H+ leaves acidb) Conjugate acid = base + H+ c) Conjugate acid-base pair are related by loss/gain of H+

d) Competition for H+ by A- and H2O; strongest base wins

acid base hydronium ion

H O

HH Cl H O

H

H Cl+ +

acid baseconjugateacid

conjugatebase

2) Ka = acid dissociation constant

3) Sample Exercise: Write simple Ionizations for:HCl, HC2H3O2, NH4

+, C6H5NH3+, Al(H2O)6

3+

4) Bronsted-Lowery theory allows for non-aqueous solutionsNH3 + HCl NH4Cl

[HA]]][A[H

[HA]]][AO[HK 3

a

H N

H

H

H Cl H N

H

H

H

Cl+ +

B. Acid Strength1) Acid strength describes the equilibrium position of the ionization reaction

HA + H2O H3O+ + A-

2) Strong Acid = equilibrium lies far to the right (Large Ka)

a) Almost all HA has ionized to H+ and A- ([H+] = [HA]0)

b) A strong acid has a weak conjugate basei. To ionize fully, the conjugate base must have low proton affinityii. The conjugate base must be weaker that water

3) Weak Acid = equilibrium lies far to the left (Small Ka)

a) Almost all HA remains unionized ([H+] << [HA]0)

b) A weak acid has a strong conjugate basec) The conjugate base is much stronger than water

Strong acid Weak Acid

C. Water as an Acid and Base1) An amphoteric substance can behave as an acid or a base (water)2) Autoionization of water (reaction with itself)

H2O + H2O H3O+ + OH-

3) Ionization constant for water = KW = [H3O+][OH-] = [H+][OH-]

a) For any water solution at 25 oC, [OH-] x [H+] = KW = 1 x 10-14

b) Neutral solutions (pure water) have [OH-] = [H+] = 1 x 10-7 c) Acidic solutions: [H+] > [OH-] d) Basic solutions: [OH-] > [H+]

e) Sample Ex. Calculate [OH-] or [H+] for the following:i. [OH-] = 1 x 10-5 Mii. [OH-] = 1 x 10-7 Miii. [H+] = 10 M

D. pH Scale1) pH = -log[H+] (simplifies working with small numbers)

2) 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+]

a) pH = 3 [H+] = 10 times the [H+] at pH = 4a) pH decreases as [H+] increases (pH = 2 more acidic than pH = 3)

E. Meaning of pKa

HA + H2O H3O+ + A-

The lower the pKa, the stronger the acid

5 )10 x -log(1 -logKpK 10 x 1HA

AHK 5-aa

5a

B. Predicting Acid/Base Strength1) Size of A-: HI > HBr > HCl > HF

a) F- is small, more concentrated charge, holds on to H+

b) I- is large, less concentrated charge, gives up H+

2) Electronegativity of A-: HF > H2O > NH3 > CH4

3) Resonance Forms of A-

C. Lewis Acids and Bases1) Lewis Acid = electron pair acceptor = Electrophile2) Lewis Base = electron pair donor = Nucleophile3) Some covalently bonded molecules can be considered Lewis Acid/Base

pairs

4) Dissociation of a Lewis Acid/Base Pair (Mechanisms)

NH3 BCl3+ H3N BCl3

C CH3

CH3

Cl

CH3

C CH3

CH3

CH3

+Cl- +

OH2C CH3

CH3

H2O+

CH3

H++ C CH3

CH3

HO

CH3