Lesson 6.3: Explaining Reaction Rates: Collision Theory and Factors Affecting Reaction Rate

Chemical Kinetics

Explaining Reaction Rates

Chemical kinetics is the study of the rates at which chemical reactions occur and the factors that influence these rates. Understanding reaction rates is essential for controlling chemical processes in both laboratory and industrial settings.

Collision Theory

Principles of Collision Theory

Collision theory states that chemical reactions can only occur if reactant atoms, molecules, or ions collide with each other. However, not all collisions result in a reaction. For a collision to be effective and produce products, two main criteria must be met:

• Proper Orientation: Reactant entities must collide in a specific orientation that allows bonds to break and new bonds to form.

• Sufficient Kinetic Energy: The colliding entities must possess enough kinetic energy to overcome the activation energy barrier.

The frequency and proportion of effective collisions determine the reaction rate. Increasing the frequency of effective collisions increases the reaction rate.

Orientation and Collision Geometry

Only certain orientations during collisions lead to successful reactions. For example, in the decomposition of nitrosyl bromide gas, BrNO(g):

$2\ \mathrm{BrNO}(g) \rightarrow 2\ \mathrm{NO}(g) + \mathrm{Br}_2(g)$

Collisions where bromine atoms make direct contact can result in a reaction, while other orientations do not lead to product formation.

Activation Energy and the Activated Complex

Definition and Role of Activation Energy

Activation energy (E_a) is the minimum energy that reactant molecules must possess for a reaction to be successful. It is required to overcome electrostatic repulsion and to weaken the bonds of the reactants. Activation energy can be visualized as a potential energy barrier that reactants must overcome to form products.

$E_a = \text{Energy of activated complex} - \text{Energy of reactants}$

The activated complex (or transition state) is an unstable arrangement of atoms at the maximum potential energy point during the reaction.

Potential Energy Diagrams

• Exothermic Reaction: Products have lower potential energy than reactants; energy is released ($\Delta H < 0$).

• Endothermic Reaction: Products have higher potential energy than reactants; energy is absorbed ($\Delta H > 0$).

Factors Affecting Reaction Rate

Temperature

Increasing the temperature increases the average kinetic energy of reactant entities. This results in more entities having kinetic energy equal to or greater than the activation energy, leading to more effective collisions and a higher reaction rate. The relationship between kinetic energy and the number of entities is illustrated by the Maxwell–Boltzmann distribution.

Chemical Nature of Reactants

• Bond type, strength, and number affect the activation energy required for a successful collision.

• Reactions involving the breaking of fewer or weaker bonds proceed faster.

• Reactions between ions are usually faster than those between molecules due to the need to break covalent bonds in molecules.

• Larger and more complex molecules or ions are less likely to collide in an effective orientation, slowing the reaction rate.

Concentration and Surface Area

• Increasing the concentration of reactants increases the probability of collisions, thus increasing the reaction rate.

• Increasing the surface area of a solid reactant (e.g., by breaking it into smaller pieces or powdering it) exposes more particles to collisions, increasing the reaction rate.

Catalysts

A catalyst increases the reaction rate by providing an alternative reaction pathway with a lower activation energy. Catalysts do not increase the number of collisions or the kinetic energy of reactants but allow more collisions to be effective at a given temperature. Catalysts do not affect the overall energy change ($\Delta H$) of the reaction.

Summary Table: Factors Affecting Reaction Rate

Factor	Effect on Reaction Rate	Explanation
Temperature	Increases	More entities have energy ≥ activation energy; more effective collisions
Concentration	Increases	More frequent collisions between reactants
Surface Area	Increases	More particles exposed for collisions
Catalyst	Increases	Lowers activation energy, increasing fraction of effective collisions
Chemical Nature	Varies	Bond strength, complexity, and type affect ease of reaction

Key Terms and Definitions

• Collision theory: Chemical reactions occur only if reactants collide with proper orientation and sufficient energy.

• Activation energy (Ea): Minimum energy required for an effective collision.

• Activated complex (transition state): Unstable arrangement of atoms at the maximum potential energy point.

• Effective collision: A collision that results in the formation of products.

• Catalyst: Substance that increases reaction rate by lowering activation energy without being consumed.

Summary Points

• Collision theory explains the requirements for chemical reactions: collision, proper orientation, and sufficient energy.

• Activation energy is the energy barrier that must be overcome for a reaction to proceed.

• Reaction rates are affected by temperature, concentration, surface area, catalysts, and the chemical nature of reactants.

• Catalysts provide alternative pathways with lower activation energy, increasing the rate of reaction.