Chemical Kinetics Notes

Chemical kinetics, also known as reaction kinetics, is the study of the rates of chemical processes.
It examines how quickly a reaction proceeds.
It investigates the influence of experimental conditions on reaction speed.
It yields information about the reaction's mechanism and transition states.
It enables the construction of mathematical models to describe reaction characteristics.

The rate of a reaction is the change in concentration of a product or reactant per unit time.
$rate = \frac{\Delta[product]}{time}$
$rate = -\frac{\Delta[reactant]}{time}$

Chemical kinetics can be analyzed from two primary viewpoints:
- Collision Theory
- Transition State Theory

Based on kinetic theory.
Assumes particles must collide with the correct orientation and sufficient kinetic energy for reactants to convert into products.
For the reaction A + B → products, particles (reactants) with sufficient kinetic energy collide and form products in a one-step process.
Not all collisions are successful because not all particles possess sufficient kinetic energy (activation energy).
Activation energy is the minimum energy required for a reaction to occur.

This theory posits that as reactants approach each other, a transitory activated complex (transition state) is formed at a potential energy maximum.
At the activated complex, original bonds weaken, and new bonds are partially formed.
Collisions are not necessary for a successful reaction: A + B → [T.S.] → products
In transition state theory, activation energy is the energy difference between reactants and the potential energy maximum.

As reactant concentration increases, the frequency of molecular collisions increases, raising the probability of successful collisions and, thus, the reaction rate.

Molecules at higher temperatures have greater average kinetic energy.
While collision frequency increases at higher temperatures, its contribution to the increase in reaction rate is small.
The proportion of reactant molecules with sufficient energy to react (energy greater than activation energy: E > E_a) is significantly greater, leading to more successful collisions and a faster reaction rate.

A catalyst is a substance that accelerates the rate of a chemical reaction but remains chemically unchanged afterward.
Catalysts increase the reaction rate by providing a different reaction mechanism with a lower activation energy, enabling more particles to collide successfully.

Increasing the pressure in a gaseous reaction increases the frequency of collisions between reactant particles.
This leads to a corresponding increase in the frequency of successful collisions and, thus, an increase in the rate of reaction.

Reactants in the solid state react faster when subdivided into smaller particles.
This results in a greater surface area per unit volume, increasing contact with other reactant particles, and thus accelerating the reaction.