Reactant molecules collide with each other but do not react. Explain why this might happen.
What this question is asking for is an explanation for one particular scenario. For this scenario, we can imagine two molecules hurtling towards one another. The two molecules undergo some form of collision and simply bounce off one another. We can think about this process using an energy reaction profile diagram.
Here we have the reactants labelled 𝑅 and the products labelled 𝑃. In between, we have an energy profile that shows a little bump. The height of this little bump above the reactant energy is called the activation energy. The activation energy is the minimum energy required by the reacting molecules for the reaction to occur. Let’s imagine that the kinetic energy of each molecule is 𝑥. If the total, two 𝑥, is less than the activation energy, the reaction cannot occur.
Before we go any further, I just want to point out that activation energy is a bulk property, whereas the energy for a reaction to occur on the molecular scale may be slightly different depending on the orientation of the molecules and so forth. However, it is generally true that the collision energy must be at or above the activation energy for the reaction to occur.
Now how do we put that into a full sentence? Reactant molecules collide with each other but do not react. Why? Because the reactant molecules during the collision do not have the minimum amount of kinetic energy required to react.