Collision theory

The collision theory states that reacting particles must collide, and that only collisions with enough energy and the correct orientation actually lead to a reaction. Malay: teori perlanggaran · Chinese: 碰撞理论.

This model explains why reactions happen at all and why their rates change. For any two particles to react, they first have to meet, so a collision is essential. But not every collision produces a reaction. The colliding particles must have energy equal to or greater than the activation energy, and they must approach in the correct orientation. A collision that meets both conditions is called an effective collision. The rate of reaction depends on the frequency of these effective collisions, that is, how many happen each second.

Example. In the reaction between hydrogen and iodine, two molecules that drift into each other gently, or that meet end-on in the wrong way, simply bounce apart unchanged. Only when they collide hard enough and line up correctly do the bonds break and new bonds form. Heating the mixture makes the molecules move faster, so they collide more often and a larger fraction of collisions has enough energy, and the rate rises.

Confusion to avoid. Do not say that increasing a factor simply gives more collisions and stop there. A full answer must explain the effect through the frequency of effective collisions, mentioning energy or orientation as appropriate. Merely having more collisions is not enough if they are not effective.

The collision theory is the engine of this whole chapter: it is the single idea used to explain how temperature, concentration, surface area, pressure and a catalyst each change the rate, so SPM Chemistry rewards an answer phrased in its exact language.