By Michael Judge
Updated Mar 24, 2022
In chemistry, the speed at which a reaction proceeds is critical—especially for industrial processes. A reaction that is thermodynamically favorable but sluggish, such as the conversion of diamond to graphite, can be practically useless. Conversely, an overly rapid reaction may pose safety risks. Understanding and controlling the factors that influence reaction rates allows chemists to design safer, more efficient processes.
Raising temperature generally accelerates reactions. The underlying reason is the activation energy barrier that must be overcome for molecules to react. Higher thermal energy increases the kinetic energy of molecules, so more collisions achieve the critical activation energy. A useful rule of thumb is that, for many reactions, the rate roughly doubles for every 10 °C rise in temperature (Arrhenius behavior).
For reactions in the same phase—e.g., two solutes in water—higher concentrations raise the probability of productive collisions, thereby speeding the reaction. The magnitude of the effect depends on the reaction’s order with respect to each reactant. In the gas phase, increasing pressure similarly boosts collision frequency, often accelerating the reaction in proportion to the pressure rise.
The surrounding medium can significantly alter reaction rates. Solvents that stabilize charged or polar transition states, such as water or highly polar organic solvents, can lower the activation energy and speed up reactions involving ionic intermediates. Conversely, a nonpolar solvent might slow down a reaction that requires a polar transition state.
Catalysts lower the activation energy of a reaction by providing an alternative pathway. This can involve adsorption of reactants onto a catalytic surface, formation of intermediate complexes, or provision of an organized environment that favors the transition state. Because more molecules possess the lower energy barrier at a given temperature, the overall rate increases without the catalyst being consumed.
When one reactant is a solid, the exposed surface area limits the reaction to the interface with the other phase. Increasing the surface area—by crushing a solid into powder, for example—provides more active sites and thereby accelerates the reaction. Classic examples include faster rusting of finely divided iron compared to a solid block.
By strategically manipulating these factors, chemists can fine‑tune reaction rates to meet safety, efficiency, and economic goals.
