5.1 Elementary reactions and molecularity

Elementary Reactions and Molecularity

Every complex reaction you encounter can be broken down into simpler, single-step events called elementary reactions. These are the individual molecular-level steps that, taken together, make up a full reaction mechanism. Understanding them is essential because the rate law for an elementary reaction can be written directly from its chemical equation, something you cannot do for an overall reaction.

Elementary Reactions in Kinetics

An elementary reaction describes a single molecular event: reactants collide (or a single molecule rearranges) and products form in one step, with no intermediates along the way.

Key properties of elementary reactions:

• They represent actual molecular events, not just balanced equations
• Reactants convert directly into products in a single step
• Their rate laws come directly from the stoichiometry of that step (this is unique to elementary reactions)
• They serve as building blocks for multi-step mechanisms
• In a multi-step mechanism, the slowest elementary step (the rate-determining step) controls the overall reaction rate

Types of Elementary Reactions

The classification is based on molecularity, which is the number of reactant molecules (or atoms/ions) that participate in a single elementary step.

• Unimolecular reactions involve one reactant molecule undergoing a change on its own, such as decomposition or isomerization.
  • Example: $A \rightarrow B + C$
  • Rate law: $\text{Rate} = k[A]$ (first-order)
  • A real example is the decomposition of $N_2O_5$ into $NO_2$ and $NO_3$.
• Bimolecular reactions involve two reactant molecules colliding and reacting. These are the most common type of elementary reaction.
  • The two molecules can be the same species ($A + A \rightarrow \text{products}$) or different species ($A + B \rightarrow C + D$).
  • Rate law: $\text{Rate} = k[A][B]$ (second-order overall)
  • If both molecules are the same species: $\text{Rate} = k[A]^2$
• Termolecular reactions involve three molecules colliding simultaneously. These are extremely rare because the probability of three molecules colliding at the same instant with the right orientation and sufficient energy is very low.
  • Example: $A + B + C \rightarrow D + E$
  • Rate law: $\text{Rate} = k[A][B][C]$ (third-order overall)

Reactions involving four or more molecules in a single step are essentially never observed.

Molecularity and Rate Laws

For elementary reactions, the rate law exponents equal the stoichiometric coefficients of the reactants in that step. This is the defining feature that makes elementary reactions so useful in kinetics.

The overall order of an elementary reaction is simply the sum of the exponents, which equals the molecularity.

Identifying Reaction Molecularity

To determine molecularity, count the number of individual reactant particles in the elementary step:

1. One reactant particle → unimolecular (e.g., $O_3 \rightarrow O_2 + O$)
2. Two reactant particles → bimolecular (e.g., $NO + O_3 \rightarrow NO_2 + O_2$)
3. Three reactant particles → termolecular (e.g., $2NO + O_2 \rightarrow 2NO_2$)

Watch out for stoichiometric coefficients. The reaction $2A \rightarrow B + C$ involves two molecules of A colliding, so it is bimolecular, not unimolecular. The coefficient "2" tells you two separate A molecules participate. The rate law would be $\text{Rate} = k[A]^2$.