Kinetics I

Full Notes

Collision Theory and Factors Affecting Rate

The rate of a reaction measures how quickly reactants are converted into products.

It depends on how frequently reacting particles collide and how many of those collisions are successful.

According to collision theory, for a reaction to occur:

• Reactant particles must collide.
• Collisions must have sufficient energy (≥ activation energy, E_a).

(Particles must also be correctly oriented - this is covered in more detail later in the course, see here)

Activation Energy (E_a)

Activation Energy (E_a) is the minimum energy that colliding particles must have to cause a reaction. Only collisions with energy ≥ E_a result in bond-breaking and forming.

If collision energy is lower than activation energy, the reaction does not occur.

Factors that increase rate

• Concentration (solutions): More particles in the same volume means more frequent collisions.
• Pressure (gases): More particles per unit volume means more frequent collisions.
• Temperature: Increases kinetic energy means more collisions and more occur with the required activation energy.
• Surface area (solids): More area exposed means more collisions.

Calculating Rate of Reaction

There are two main methods to calculate rate:

From amounts data:

• Rate = change in concentration or mass / time

Units depend on the quantity being measured (e.g. g s^-1, mol dm⁻³s^-1)

From a concentration-time graph:

• Initial rate: Draw tangent at t = 0, find gradient.
• Rate at time t: Draw tangent at that point, calculate gradient.

Maxwell-Boltzmann Distribution

The Maxwell-Boltzmann distribution is a graph that shows the spread of molecular energies in a gas.

It helps explain why:

• Not all molecules have the same energy.
• Only a small fraction of molecules have enough energy to react.
• Increasing temperature increases reaction rate.

Features of the Maxwell-Boltzmann Distribution Curve

• Starts at the origin (0,0): No particles have zero energy.
• Peaks at the most probable energy: The energy that most particles have.
• Has a long tail to the right: A few molecules have very high energy (never crosses the x axis again).
• Area under the curve = Total number of molecules.
• Only molecules with energy ≥ activation energy (E_a) can react.

Increasing temperature:

If temperature is increased, the curve shifts to the right and becomes more spread out.

• Total area stays the same (same number of particles).
• More particles exceed activation energy, meaning an increased rate.
• Even a small increase in temperature can greatly increase rate.

Catalysts

A catalyst is a substance that speeds up a reaction without being chemically changed or used up.

Catalysts work by providing an alternative reaction pathway with a lower activation energy (E_a).

This increases the frequency of successful collisions, leading to a faster reaction.

Maxwell-Boltzmann with a catalyst:

If a catalyst is used, the Maxwell-Boltzmann curve shape does not change, but the activation energy marker shifts left.

This represents a greater proportion of particles having the required activation energy and explains the increase in rate of reaction when a catalyst is used.

Reaction Profile Diagrams

Reaction profile diagrams show how the energies of reactants change during a reaction:

In a reaction profile diagram a catalyst causes a lower peak (activation energy) however the enthalpy change (ΔH) is the same with or without the catalyst.

Heterogeneous Catalysts

Heterogenous catalysts have been covered in more detail along with examples here.

Heterogeneous catalysts are in a different phase to reactants.

The reaction happens on the surface of the catalyst.

Heterogenous catalysts are used in many industrial processes.

Economic Benefits of Catalysts

Catalysts are cost-effective because they:

• Allow reactions to occur at lower temperatures and pressures, saving energy
• Increase reaction rate, giving higher productivity

Summary

• Reactions occur when reactant particles collide with enough energy (activation energy, E_a) and correct orientation.
• Rates increase with higher concentration, pressure, temperature and surface area.
• Only molecules with energy greater than or equal to the activation energy can react.
• Higher temperatures means more molecules can collide with the required activation energy so rate increases.
• Catalysts provide an alternative pathway with lower activation energy and are not used up.
• Heterogeneous catalysts are in a different phase to reactants and reactions occur at the surface of the catalyst. They are widely used in industry.