Collision theory

Collision theory is a theory proposed by Max Trautz and William Lewis in 1916 and 1918, that qualitatively explains how chemical reactions occur and why reaction rates differ for different reactions. The collision theory states that when suitable particles of the reactant hit each other, only a certain fraction of the collisions cause any noticeable or significant chemical change; these successful changes are called successful collisions. The successful collisions must have enough energy, also known as activation energy, at the moment of impact to break the preexisting bonds and form all new bonds. This results in the products of the reaction. Increasing the concentration of the reactant particles or raising the temperature, thus bringing about more collisions and therefore many more successful collisions, increases the rate of reaction.

History

The collision theory was proposed independently by Max Trautz in 1916 and William Lewis in 1918. The theory is based on the idea that molecules must collide to react, and that only a fraction of collisions are successful in producing a reaction.

Principles

The collision theory is based on the following principles:

• Molecules must collide in order to react.
• Not all collisions cause a reaction; only those with sufficient energy, known as the activation energy, cause a reaction.
• The orientation of the colliding molecules also affects whether a reaction occurs.

Factors affecting reaction rates

According to the collision theory, the rate of a chemical reaction is proportional to the number of collisions between reactant molecules. Several factors can affect the rate of a reaction:

• Concentration: Increasing the concentration of reactant molecules increases the number of collisions and thus the reaction rate.
• Temperature: Raising the temperature increases the kinetic energy of the molecules, leading to more frequent and more energetic collisions, and thus a higher reaction rate.
• Catalysts: Catalysts can lower the activation energy required for a reaction, leading to a higher reaction rate.
• Surface area: Increasing the surface area of a solid reactant exposes more of its particles to attack, thus increasing the reaction rate.

See also

• Chemical kinetics
• Rate equation
• Transition state theory

References

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