Collision Theory: How Chemical Reactions Occur
In order for a chemical reaction to take place, the reactants must collide. The collision between the molecules in a chemical reaction provides the kinetic energy needed to break the necessary bonds so that new bonds can be formed.
Sometimes, even if there is a collision, not enough kinetic energy is available to be transferred — the molecules aren’t moving fast enough. You can help the situation somewhat by heating the mixture of reactants. The temperature is a measure of the average kinetic energy of the molecules; raising the temperature increases the kinetic energy available to break bonds during collisions.
The molecules must also collide in the right orientation, or hit at the right spot, in order for the reaction to occur. Here’s an example: Suppose you have an equation showing molecule A-B reacting with C to form C-A and B, like this:
A-B + C→C-A + B
The way this equation is written, the reaction requires that reactant C collide with A-B on the A end of the molecule. If it hits the B end, nothing will happen. The A end of this hypothetical molecule is called the reactive site, the place on the molecule that the collision must take place in order for the reaction to occur.
If C collides at the A end of the molecule, then there’s a chance that enough energy can be transferred to break the A-B bond. After the A-B bond is broken, the C-A bond can be formed. The equation for this reaction process can be shown in this way:
C~A~B→C-A + B
So in order for this reaction to occur, there must be a collision between C and A-B at the reactive site. The collision between C and A-B has to transfer enough energy to break the A-B bond, allowing the C-A bond to form.
Energy is required to break a bond between atoms.
This example is a simple one. Many reactions are one-step, but many others require several steps in going from reactants to final products. In the process, several compounds may be formed that react with each other to give the final products. These compounds are called intermediates.
An exothermic example of chemical reactions
Imagine that the hypothetical reaction A-B + C→C-A + B is exothermic — a reaction in which heat is given off (released) when going from reactants to products. The reactants start off at a higher energy state than the products, so energy is released in going from reactants to products.
In the diagram below, the activation energy for the reaction (the energy that you have to put in to get the reaction going) is shown as:
The energy diagram shows the collision of C and A-B with the breaking of the A-B bond and the forming of the C-A bond at the top of an activation energy hill. This grouping of reactants at the top of the activation energy hill is sometimes called the transition state of the reaction. The difference in the energy level of the reactants and the energy level of the products is the amount of energy (heat) that is released in the reaction.
An endothermic example of chemical reactions
Suppose that the hypothetical reaction A-B + C→C-A + B is endothermic — a reaction in which heat is absorbed in going from reactants to products — so the reactants are at a lower energy state than the products. The following energy diagram shows this reaction.
This diagram also shows that an activation energy is associated with the reaction. In going from reactants to products, you have to put in more energy initially to get the reaction started, and then you get that energy back out as the reaction proceeds.
Notice that the transition state appears at the top of the activation energy hill — just like in the exothermic-reaction energy diagram. The difference is that, in going from reactants to products, energy (heat) must be absorbed in the endothermic example.