You are viewing a free preview of this lesson.
Subscribe to unlock all 10 lessons in this course and every other course on LearningBro.
For a chemical reaction to occur, reactant particles must collide. But not just any collision will do — the particles must collide with sufficient energy and in the correct orientation. This is the foundation of collision theory, and it explains why reactions happen at the rates they do.
Collision theory states that a reaction occurs when two conditions are met simultaneously:
The particles must collide with energy equal to or greater than the activation energy (Ea). The activation energy is the minimum energy required to break the bonds in the reactants and initiate the reaction. Collisions with less energy than Ea simply result in the particles bouncing apart unchanged.
The particles must collide with the correct orientation. Even if particles have sufficient energy, the reactive parts of the molecules must be aligned appropriately. For example, in the reaction between an OH⁻ ion and a halogenoalkane, the hydroxide must approach the carbon atom bonded to the halogen — not some other part of the molecule.
Only collisions that satisfy both conditions are called successful collisions (or effective collisions). The rate of reaction depends on the frequency of these successful collisions per unit time.
The activation energy can be visualised on an enthalpy profile diagram. For an exothermic reaction, the products are at a lower energy level than the reactants, but there is an energy "hump" that the reactants must overcome first. This hump represents Ea.
For an endothermic reaction, the products are at a higher energy level, and the activation energy is measured from the reactant energy level to the top of the energy barrier.
The higher the activation energy, the fewer particles have enough energy to react at a given temperature, and the slower the reaction.
At any given temperature, the particles in a gas or solution have a range of kinetic energies. Some are moving slowly, some are moving very fast, and most have intermediate energies. The Maxwell-Boltzmann distribution shows the spread of these energies.
Key features of the distribution curve:
This last point is critical: only particles in the shaded region beyond Ea can undergo successful collisions.
When temperature increases:
Even a modest temperature rise (say 10 °C) can dramatically increase the proportion of particles with energy ≥ Ea. This is the primary reason why increasing temperature increases the rate of reaction — there are far more particles capable of overcoming the activation energy barrier.
A common misconception is that increasing temperature simply makes particles move faster. While this is true, the key effect is the change in the shape of the distribution — the proportion of high-energy particles increases disproportionately.
Increasing the concentration of reactants in solution (or increasing the pressure of gaseous reactants) means there are more particles per unit volume. This increases the frequency of collisions, which increases the frequency of successful collisions, and therefore increases the rate.
Note: increasing concentration does not change the proportion of particles with energy ≥ Ea. It only changes how often they collide.
As discussed above, increasing temperature increases the proportion of particles with energy ≥ Ea. It also slightly increases collision frequency (particles move faster), but the dominant effect is the increased proportion above Ea.
For reactions involving solids, breaking the solid into smaller pieces increases the surface area exposed to the other reactant. More surface area means more collisions can occur per unit time. This is why powdered reactants react faster than lumps.
A catalyst provides an alternative reaction pathway with a lower activation energy. On a Maxwell-Boltzmann diagram, this shifts the Ea marker to the left. With a lower Ea, a much larger proportion of particles now have sufficient energy to react — so the rate increases without changing the temperature.
Importantly, a catalyst:
While collision theory provides a qualitative understanding of reaction rates, the quantitative relationship between concentration and rate is described by the rate equation, which you will study in detail in Lesson 3. For now, understand that collision theory is the conceptual foundation upon which the mathematical treatment of kinetics is built.
The rate of a reaction depends on the frequency of successful collisions — those with energy ≥ Ea and correct orientation. Temperature affects the energy distribution of particles; concentration affects collision frequency; surface area affects the number of collisions at the solid surface; and catalysts lower the energy barrier. The Maxwell-Boltzmann distribution is the key tool for understanding why temperature and catalysts have such a dramatic effect on rate.