Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases
I.
Review of Simple Kinetics and Thermodynamics
A.
Definitions
1) Thermodynamics = changes in energy during a process or reaction.
Determines extent of completion of the reaction or process
2)
B.
Kinetics = rate of a process or reaction. Determines how fast the reaction
or process occurs.
Equilibria
1) Equilibrium = state of a system in which the concentrations of reactants
and products are no longer changing.
2)
Equilibrium Constant
a) If K is large, reaction goes forward
b) If K is small, reaction goes in reverse
A  B  C  D
K
K
[C][D]
[A][B]
3)
4)
Relating Gibb’s Free Energy Change to Equilibrium Constants
a) DG0 = Gibb’s Free Energy Change = describes the overall energy
change as a reaction reaches equilibrium
b)
DG0 = -RTlnK
R = 1.986 cal/deg mol
T = temperature in Kelvins (oC + 273)
c)
d)
e)
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
a) DG0 = DH0 - TDS0
b) DH0 = Enthalpy = Broken Bond Strengths – Formed Bond Strengths
1) -DH0 = Exothermic reaction (gives off energy)
2)  DH0 = Endothermic reaction (requires energy input)
c) DS0 = Entropy = Amount of order in the system
1) - DS0 = less disorder (fewer molecules in the system)
2) + DS0 = more disorder (more molecules in the system)
C.
Reaction Rates
1) Activation Energy determines reaction rates
a) Small Ea = fast reaction
b) Large Ea = slow reaction
2)
Rate Constants
AB 
 C  D
k
rate  k[A][B] L-1s -1
3)
The Arrhenius Equation
k  Ae
- E a / RT
A = maximum rate constant possible = different for each reaction
High T
-Ea/RT becomes small
e0 = 1
k=A
II.
Review of Acids and Bases
A.
Bronsted Acids/Bases
1) Acid = H+ donor
2) Base = H+ acceptor
Ka
H3O+ + OH-
3)
Ionization of Water: H2O + H2O
4)
pH = -log[H3O+]
5)
pKa = -logKa = pH at which HA is half-dissociated
6)
pKa + pKb = 14
If you know Ka, Kb, pKa, pKb, you can find all others
B.
C.
Predicting Acid/Base Strength
1) 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-
Lewis Acids and Bases
1) Lewis Acid = electron pair acceptor
2) Lewis Base = electron pair donor
3) Some covalently bonded molecules can be considered Lewis Acid/Base
pairs
NH3
4)
C CH3
CH3
H3N
BCl3
BCl3
Dissociation of a Lewis Acid/Base Pair (Mechanisms)
CH3
CH3
Cl
+
-
Cl
+
+
C CH3
CH3
CH3
CH3
OH2
H2O
+
C CH3
CH3
+
H
+
HO C CH3
CH3