Collision theory explains how chemical reactions occur. The molecules or atoms colliding must have a minimum amount of energy and a specific orientation for a chemical reaction to take place.
Collision theory describes how chemical reactions occur through molecular collisions and why reaction rates vary between reactions. It can be used to explain why certain factors affect reaction rate. According to collision theory, collisions lead to chemical change. However, not all collisions result in a chemical change. The collisions must occur with a specific orientation and must occur with sufficient energy. The majority of collisions do not result in a chemical change. The following describes the reaction between nitric oxide (NO) and ozone (O3).
According to collision theory, this chemical reaction occurs through collisions between nitric oxide (NO) and ozone (O3) molecules. For a reaction to occur, the nitrogen atom of a nitric oxide (NO) molecule must strike one of the oxygen atoms in the ozone (O3) molecule. Furthermore, the angle of the two molecules also plays a role. A reaction will occur when the two molecules strike each other at a specific range of angles. A strike on the central oxygen on the ozone or a strike between two oxygen atoms does not result in a chemical reaction. A strike between correct atoms but at an unfavorable angle will also not result in a chemical reaction.
Effective and Ineffective Collisions
Atoms in a molecule have bonds. Bonds release energy as they form: Atoms in a molecule have less energy than those same atoms if the molecular bond is broken. A molecule is a state of low potential energy of the atoms.
Recall that kinetic energy, the energy of motion, is related to mass and speed. If two molecules collide at sufficient speeds, the kinetic energy of the collision can break some existing bonds. This can lead to formation of new bonds. Thus, according to collision theory, another factor that affects whether or not a reaction will occur during a collision is the kinetic energy, or speed, of the colliding molecules. If the molecules do not have enough energy, they will bounce off each other without a chemical reaction occurring.
During a collision, the minimum energy needed for a chemical reaction to initiate is called the activation energy. Activation energy is an increase in potential energy relative to the reactants and represents a conversion from kinetic energy to potential energy. Activation energy depends on the nature of the chemical reaction. Some reactions have higher activation energy compared to others.
This increase in energy can be demonstrated using the chemical reaction between nitric oxide (NO) and ozone (O3) as an example.
The total potential energy of the reactants is at one level when they are outside a collision. During a collision, the kinetic energy is converted to potential energy. The collision between reactants, such as NO and O3, must have more energy than the activation energy of the reaction in order for a reaction to occur. If there is sufficient energy in the collision, then the two reactants form an activated complex. An activated complex, also called a transition state, is an intermediate configuration of atoms during a chemical reaction with high potential energy. An activated complex is an inherently unstable structure. It will break down to the products of the reaction, with formation of new bonds. For this reaction, the total potential energy of the products is at a lower level than that of the reactants, and energy is released. In some reactions the reverse is true. Energy is absorbed, and products have more energy than the reactants.