General Chemistry: Thermodynamics & Kinetics
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Thermodynamics
ΔG = ΔH − TΔS
| ΔH | ΔS | Spontaneous? | |
|---|---|---|---|
| − | + | Always (all T) | Combustion |
| + | − | Never | Reverse combustion |
| − | − | Low T only | Freezing |
| + | + | High T only | Melting, boiling |
Key Point: ΔG < 0 = spontaneous. ΔG = 0 = equilibrium. ΔG > 0 = nonspontaneous.
ΔG° = −RT ln K | ΔG° = −nFE°
Kinetics
Rate Laws
| Order | Rate Law | Units of k | Half-life | Plot for straight line |
|---|---|---|---|---|
| Zero | rate = k | M/s | t½ = [A]₀/2k | [A] vs t |
| First | rate = k[A] | 1/s | t½ = 0.693/k | ln[A] vs t |
| Second | rate = k[A]² | 1/(M·s) | t½ = 1/(k[A]₀) | 1/[A] vs t |
Key Point: Rate law exponents are determined experimentally, NOT from balanced equation coefficients.
Arrhenius equation: k = Ae−Ea/RT
A catalyst lowers Ea without changing ΔG or equilibrium position. It speeds up both forward and reverse reactions equally.
Chemical Equilibrium
K = [products]coefficients / [reactants]coefficients
Key Point: Q < K → forward. Q > K → reverse. Q = K → equilibrium.
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