Reaction rate is the change in concentration over time. Concentration, temperature, surface area, and catalysts all affect reaction speed by altering collision frequency and molecular energy.
Observe how changing one factor at a time affects reaction rate in experiments or simulations.
From your study of chemical kinetics you know that reactions proceed at measurable speeds — some explosively fast, others imperceptibly slow. Reaction rate formalizes this idea: it is the change in concentration of a reactant or product per unit time, typically expressed in mol·L⁻¹·s⁻¹. For a reaction A → B, the rate can be written as −Δ[A]/Δt (negative because reactant concentration decreases) or +Δ[B]/Δt. The key insight is that rate is not a fixed property of a reaction — it changes as conditions change, and understanding *which* conditions matter and *why* is the core of this topic.
Concentration is the most intuitive factor. If you double the number of reactant molecules in a given volume, collisions between them become more frequent, and the reaction speeds up. Think of it like a crowded dance floor: the more people packed into the room, the more often they bump into each other. This is why many reactions start fast (high concentration) and slow down as reactants are consumed. The precise mathematical relationship between concentration and rate — the rate law — is the subject of the next topic, but the qualitative principle is straightforward: more molecules per liter means more collisions per second.
Temperature affects rate through molecular energy, not just collision frequency. Raising the temperature does increase how often molecules collide, but the dominant effect is that a larger fraction of collisions now carry enough energy to overcome the activation energy barrier — the minimum energy required for bonds to break and reform. A useful rule of thumb is that many reactions roughly double in rate for every 10 °C increase. This is why refrigeration slows food spoilage (fewer molecules have the energy to drive decomposition reactions) and why a spark can ignite a fuel-air mixture (locally raising temperature past the activation threshold).
Surface area matters for reactions involving solids. A sugar cube dissolves slowly in water, but the same mass of powdered sugar dissolves almost instantly. The total amount of sugar is the same, but the powder exposes vastly more surface to the water, allowing many more collisions between sugar molecules and water molecules at any given moment. This factor is especially important in industrial chemistry, where catalysts are often ground into fine particles or spread across porous supports to maximize the reactive surface.
Catalysts increase reaction rate without being consumed. They work by providing an alternative reaction pathway with a lower activation energy. The reactants and products are unchanged — the catalyst simply makes it easier for molecules to reach the transition state. Enzymes in biological systems are a familiar example: they accelerate reactions by factors of millions, allowing life-sustaining chemistry to proceed at body temperature. Crucially, catalysts do not shift the position of equilibrium; they speed up both the forward and reverse reactions equally, so the system reaches equilibrium faster but at the same concentrations.