Chemical reaction rate is simply a measure of how quickly reactants are consumed and products are formed. Observing that a chemical process speeds up when heat is applied is a common phenomenon, from cooking food faster to industrial manufacturing. This increase in speed is often not linear; a small rise in temperature can lead to a dramatic acceleration of the entire reaction. Understanding this relationship requires looking closely at the movement of molecules and the energy transfers happening at the atomic level.
Collision Theory: The Foundation of Reactions
For a chemical reaction to occur, the reacting particles—whether they are atoms, ions, or molecules—must physically come into contact with one another. This foundational principle is known as collision theory, which posits that molecular collisions are the prerequisite for any chemical transformation. Reactant molecules exist in constant, random motion, continuously bouncing off the walls of their container and each other.
The rate at which a reaction proceeds is directly proportional to how often these particles collide. If the frequency of collisions increases, the potential for a reaction to take place also increases. However, simply colliding is not enough to guarantee that the reactant particles will reorganize themselves into new product molecules.
Temperature and Kinetic Energy
The temperature of a substance is a direct measure of the average kinetic energy of its constituent particles. Kinetic energy is the energy associated with motion, meaning that as heat is added to a system, the average speed of the molecules within that system increases. This addition of thermal energy causes molecules to move more vigorously.
Since molecules are moving more quickly, they naturally traverse a given volume faster and encounter other molecules more often. An increase in temperature therefore leads to a measurable increase in the total frequency of collisions between reactant particles. While this higher collision frequency contributes to a faster reaction rate, calculations show that this modest increase alone cannot account for the dramatic speed-up observed when a substance is heated. The primary reason for acceleration must lie in a separate, more impactful factor related to energy.
Activation Energy: The Required Threshold
Not every molecular collision results in a chemical reaction because the reactants must overcome an energy barrier known as the activation energy, symbolized as $E_a$. This barrier represents the minimum amount of energy required to distort the existing bonds within the reactant molecules and arrange them into a high-energy, unstable intermediate state called the transition state. Without reaching the transition state, the existing chemical bonds will not break, and new ones cannot form.
One way to visualize this is imagining an object that needs to be pushed over a hill. The height of the hill represents the activation energy, and the object needs sufficient kinetic energy to reach the peak. Only collisions that possess energy equal to or greater than this specific threshold will be successful in forming products.
Furthermore, even an energetic collision may fail if the reactant molecules are not aligned correctly in a specific geometric orientation. Complex molecules must collide at the precise site where the bond-forming or bond-breaking action is supposed to occur. Achieving both the correct energy and the specific orientation simultaneously is a requirement for a successful collision.
Explaining the Exponential Increase
The dramatic acceleration of reaction rates with a modest temperature increase is best explained by the distribution of molecular kinetic energies within a system. In any collection of molecules, the energy is not uniform; some move slowly, and others move very fast, following a statistical pattern described by the Maxwell-Boltzmann distribution. At lower temperatures, only a small fraction of the molecules possess kinetic energy exceeding the required activation energy ($E_a$).
When the temperature is raised, the entire energy distribution curve shifts toward higher average energies. This shift causes a disproportionately large increase in the area under the curve that lies past the activation energy threshold. Even a $10^{\circ}\text{C}$ rise in temperature can double or triple the reaction rate because the number of molecules capable of overcoming the energy barrier increases exponentially.
This large increase in molecules with sufficient energy is the primary reason for the rapid speed-up, dwarfing the effect of simply increasing the total number of collisions. The reaction rate is governed far more by how many of those collisions are “effective”—meaning they possess both the correct orientation and the necessary activation energy to successfully transition into products.