Kinetic molecular theory vs collision theory

The kinetic molecular theory describes the behaviour of particles in an ideal gas. It states that:

• Gas particles move randomly and collide with each other elastically (no energy is lost).
• The average kinetic energy of gas particles is proportional to the absolute temperature.
• At absolute zero (0 K), particle motion stops completely.

In contrast, the collision theory expands on these ideas to explain how chemical reactions occur:

• It allows for inelastic collisions where particles can undergo chemical changes.
• It applies to reacting species in any state of matter, not just gases.

Requirements for successful collisions

According to collision theory, most collisions are unsuccessful because the particles' kinetic energy is insufficient to break chemical bonds. These particles simply bounce off each other without undergoing chemical changes.

However, some collisions redistribute kinetic energy unequally, allowing certain particles to gain enough energy to break or rearrange bonds, transforming reactants into products. These are called successful collisions.

For a collision to be successful, three conditions must be met:

1. Two or more particles must collide.
2. The colliding particles must have the correct mutual orientation.
3. The total kinetic energy of the particles must be sufficient to initiate the reaction.

For example, consider the reaction:

AB + CD ➔ AC + BD

For this reaction to occur, the bonds A–B and C–D must break, and new bonds A–C and B–D must form. This is only possible if atom A is close to atom C and atom B is close to atom D during the collision. Other orientations result in unsuccessful collisions.

Reactions involving symmetrical species like atoms or monatomic ions tend to proceed faster, as their collision outcomes are less dependent on orientation compared to reactions with large, complex molecules.

Factors affecting reaction rates

The rate of a chemical reaction depends on the frequency of successful collisions per unit volume. Any factor that changes the collision frequency or average kinetic energy will affect the reaction rate.

1. Surface area (of solids)

• When the exposed surface area of the solid is increased, more particles on the surface are available to collide and react. This leads to a higher frequency of effective collisions between the solid and other reactants.
• For example, crushing a solid into a powder provides more exposed surface.
• Therefore, increasing the surface area of a solid reactant results in an increased reaction rate.

2. Concentration (of solutions)

• If the concentration of reactants in solution is increased, the particles will on average be closer together.
• Particles that are closer together will collide more frequently, increasing the number of effective collisions.
• Therefore, increasing the concentration increases the reaction rate.

3. Pressure (of gases)

• Raising the pressure forces the gas particles closer together.
• Particles that are closer together will collide more frequently, increasing the number of effective collisions.
• Therefore, increasing the pressure increases the reaction rate.

4. Temperature

• Raising the temperature increases the kinetic energy and speed of the particles.
• As the particles move faster, they collide more frequently, resulting in an increased frequency of collisions.
• Additionally, the increased kinetic energy means more particles have the necessary energy to overcome the activation energy, resulting in a greater proportion of effective collisions.
• Therefore, increasing temperature results in an increased reaction rate.

5. Catalyst

• A catalyst provides an alternative pathway or mechanism for the reaction that has a lower activation energy.
• A lower activation energy means particles require less kinetic energy to react.
• More particles have the activation energy, leading to an increased number of effective collisions.
• Consequently, adding a catalyst increases the overall rate of reaction.