Chemical Kinetics: Mechanisms | General Chemistry 2

Mechanisms in chemical kinetics are studied in this chapter: elementary reactions, intermediates, activation energy, Arrhenius equation, rate-determining step, reversible reactions, catalysis


Many reactions involve more than one step. They can be divided in elementary reactions

Elementary reactions:
- occur in a single step
- have reaction orders equal to the stoichiometric coefficients for each reactant

Intermediate: species produced in one step and consumed in a subsequent step
⇒ does not appear in the overall equation


Step (1): 2 NO → N2O2
Step (2): N2O2 + O2 → 2 NO2
Overall reaction (1) + (2): 2 NO + O2 → 2 NO2

2 NO + O2 → 2 NO2 can be divided in 2 elementary steps (1) and (2) and N2O2 is an intermediate.

rate law of (1) = k1 [NO]2
rate law of (2) = k2 [N2O2] [O2]

Activation Energy

Activation energy Ea (in J.mol-1):
energy that must be provided to compounds to result in a chemical reaction 

Arrhenius Equation

k = Ae-Ea/RT

ln k = - EaRT + ln A

k = reaction-rate constant
Ea = activation energy (in J.mol-1)
R = molar gas constant = 8.314 J.mol-1.K-1
T = temperature (in K)
A = constant which depends on the reaction


Rate-Determining Step

Rate-determining step: step which is much slower than any of the other steps
⇒ controls the overall reaction rate


Step (1): 2 NO2 → NO + NO3 (slow)
Step (2): NO3 + CO → NO2 + CO2 (fast)
Overall: NO2 + CO → NO + CO2

NO2 + CO → NO + CO2 is a two-step reaction and (1) is the rate-determining step
Law rate of the overall reaction = law rate of (1) = k [NO2]2

Reversible Reactions

Reversible reactions: proceed in both the forward direction and the reverse direction
At equilibrium: rate of forward reaction = rate of reverse reaction


2 NO N2O4 is a reversible reaction
At equilibrium: k1 [NO]2 = k-1 [N2O4]


Catalyst: substance that increases the reaction rate but is not consumed in the reaction
Homogeneous catalyst: same phase at the reaction mixture
Heterogeneous catalyst: different phase at the reaction mixture

A catalyst lowers the activation energy barrier to the reaction and thereby increases the reaction rate: