Chemical Kinetics: Mechanisms | General Chemistry 2

Activation energy E_a (in J.mol^-1):

The minimum amount of energy that must be provided to the reactants to initiate a chemical reaction. This energy corresponds to the minimum energy that the colliding molecules must have for the collision to be effective

The Arrhenius equation is an equation that gives the dependence of the rate constant k of a reaction on temperature:
 

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

Elementary reaction:

A reaction that occurs in a single collision of the reactant molecules (single step). The reaction order of an elementary reaction is equal to the stoichiometric coefficients of the reactants ⇒ the rate law of an elementary reaction can be written directly using the stoichiometric coefficient of each reactant as exponent in the rate law
 

The formation of NO₂ can be divided in 2 elementary steps (1) and (2):
Step (1): 2 NO → N₂O₂
Step (2): N₂O₂ + O₂ → 2 NO₂
Overall reaction (1) + (2): 2 NO + O₂ → 2 NO₂

The rate law of elementary reactions can be written directly using the stoichiometric coefficients:
rate law of (1) = k₁ [NO]²
rate law of (2) = k₂ [N₂O₂] [O₂]

Reversible reaction:

A reaction in which products can react to form reactants. This type of reaction proceeds in both the forward and reverse directions. At equilibrium, the rate of the forward reaction = the rate of the reverse reaction
 

2 NO $\rightleftharpoons$ N₂O₄ is a reversible reaction
At equilibrium: k₁ [NO]² = k_-1 [N₂O₄]

Intermediate vs. catalyst:

Intermediate: a species that is produced in one step and consumed in a subsequent step
Catalyst: a species that is first consumed and later regenerated
Neither intermediates nor catalysts appear in the overall balanced equation
 

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

N₂O₂ is an intermediate: it is produced in step (1) and consumed in the subsequent step (2) 

Rate-determining step:

The step in a reaction that is much slower than all the other steps. It 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

Catalyst:

A substance that increases the rate of a chemical reaction without being consumed during the reaction.  A catalyst speeds up the reaction by lowering the activation energy barrier to the reaction:

Catalysis:

A process by which a catalyst increases the reaction rate. Catalysis may be:

• Homogeneous: catalyst and reactants exist in the same phase
• Heterogeneous: the catalyst and the reactants exist in different phases. It is by far the most important type of catalysis in industrial chemistry

Enzymes:

Biological catalysts. They increase the rate of biochemical reactions and are also highly specific