The rate of a chemical reaction is defined by how quickly reactants are converted into products. Controlling this rate is a fundamental concept in chemistry, impacting everything from industrial manufacturing to metabolic pathways within a living cell. Manipulating the speed of a reaction relies on adjusting the conditions under which reactant molecules meet and interact. For any reaction to occur, reactant particles must collide with both sufficient energy and the correct physical orientation.
How Catalysts Accelerate Reactions
Catalysts accelerate chemical reactions by changing the reaction’s underlying mechanism. This acceleration is achieved by lowering the activation energy ($E_a$), which is the minimum energy barrier molecules must overcome to break existing bonds and form new ones. A catalyst offers an entirely new pathway for the reaction, one that requires significantly less energy to complete than the uncatalyzed route.
The catalyst is not consumed during the process, meaning it can facilitate the transformation of a large amount of reactants without being permanently changed. Catalysts participate by forming temporary intermediate species with the reactants, which then quickly break down to form the final products and regenerate the original catalyst molecule. Enzymes, such as catalase, are biological catalysts that perform this role with high selectivity. Industrially, catalysts like platinum and rhodium are used in catalytic converters to speed up the conversion of harmful vehicle emissions into less toxic gases.
Raising the Temperature
Increasing the temperature provides reactant molecules with greater kinetic energy, which accelerates the reaction rate. Since temperature measures the average kinetic energy of particles, a higher temperature causes molecules to move faster. This increased motion leads to a greater frequency of collisions per unit of time.
The more significant effect is the increased force of these collisions. For a reaction to succeed, molecules must collide with energy equal to or greater than the activation energy ($E_a$). Raising the temperature shifts the distribution of molecular energies so that a much larger fraction of molecules possess this minimum energy threshold. For many reactions, a temperature increase of just $10^\circ$ Celsius can approximately double the reaction rate.
Increasing Reactant Density
Chemical reactions proceed faster when the concentration or density of the reactants is increased, which is explained by collision theory. When more reactant molecules are packed into the same volume, the probability of collision increases. This applies to liquid solutions, where increasing concentration means more solute particles are present, and to gaseous reactions, where increasing the pressure forces particles closer together.
The relationship between density and reaction speed is tied to the frequency of molecular collisions. A higher density of reactants leads to a higher total number of collisions per second. Assuming the collision energy and orientation are sufficient, more total collisions translate into more effective collisions, increasing the reaction rate.
Maximizing Contact Points
For reactions involving reactants in different physical states, known as heterogeneous reactions, the total surface area available for interaction controls the reaction rate. When a solid reactant is involved, only the molecules on its exposed surface can contact the other reactant, such as a liquid or a gas. The interior of the solid is shielded from the reaction until the outer layer is consumed.
Breaking a large solid mass into smaller pieces, such as grinding it into a fine powder, increases the exposed surface area relative to its total volume. This maximizes the contact points between the solid and the other reactant, enabling a greater number of molecules to collide simultaneously. For example, fine sawdust burns much faster and more violently than a large log because the powdered wood offers a greater surface area for reaction with oxygen.