Collision theory explains why chemical reactions occur and why some are fast while others are slow. According to this theory, reactions occur when particles such as atoms, ions, or molecules collide with one another. However, not every collision produces one. Only collisions with adequate energy and proper orientation form products. These are termed successful or effective collisions.

Collision theory helps us understand and predict reaction rates. In other words, it explains why certain reactions occur rapidly while others proceed slowly.

Postulates of Collision Theory

1. Collisions Must Be Frequent

The reaction rate depends on the frequency of collisions.

Reaction Rate ∝ Number of Collisions per Unit Time

Conditions that increase collisions generally increase the reaction rate. These include:

• Higher concentration
• Greater surface area
• Increased pressure (for gases)
• The presence of a catalyst

2. Particles Need the Correct Orientation

When particles collide, they must approach in the proper direction so that the reacting parts of the molecules can come together. If the alignment is incorrect, no bond formation will occur, even if a collision takes place.

3. Particles Need Enough Energy

Colliding particles must possess enough kinetic energy to break old bonds and begin forming new ones. This minimum required energy is termed as activation energy (E_a)

Particles in a reaction mixture do not all have the same energy. Their energies spread over a range, described by the Maxwell-Boltzmann distribution. Only a certain fraction of particles have energy ≥ E_a, and these are the ones that successfully react. Increasing temperature increases this fraction, and therefore, the reaction rate rises.

Svante Arrhenius established a relationship between the reaction rate and the temperature, known as the Arrhenius equation:

k = A exp(-E_a / RT)

It shows that even a slight increase in temperature can substantially increase the number of molecules with enough energy to overcome E_a, making the reaction significantly faster.

Successful vs. Unsuccessful Collisions

Unsuccessful Collisions

Most collisions do not form products because:

• Energy is lower than the activation energy, or
• Particles collide with the wrong orientation

In such cases, particles bounce off each other without reacting.

Successful Collisions

A collision is successful when:

• Kinetic energy ≥ activation energy
• Orientation is correct

Under these conditions, particles form a short-lived, high-energy species called the activated complex or transition state. From here, bonds rearrange and result in products.

Example

Why does orientation matter

Consider the reaction between hydrogen chloride (HCl) and ethene (C₂H₄).

Ethene has an electron-rich C=C double bond that acts as a reactive site. For a reaction to occur:

• HCl must approach so that hydrogen faces the double bond
• Hydrogen bonds to one carbon
• Chloride attaches to the other

If HCl approaches incorrectly (with chlorine facing the double bond), the orientation is unfavorable, and no reaction occurs, even if the energy is sufficient.

Therefore, collision theory explains reaction rates by showing that:

• More collisions → faster reactions
• Correct orientation → colliding species can react
• Enough energy → reactions actually occur

Only collisions that satisfy all three conditions form products.