Reaction rate tells how quickly reactants are converted into products, which matters in everything from cooking and rusting to engines, medicines, and industrial chemical production. A faster reaction has more successful particle collisions per second, while a slower reaction has fewer. Chemists can control reaction rate by changing conditions such as concentration, temperature, surface area, pressure, and catalysts.
Understanding these factors helps explain both everyday changes and carefully designed laboratory reactions.
Collision theory says that particles must collide with enough energy and the correct orientation for a reaction to occur. Increasing the number of collisions, increasing the fraction of collisions with enough energy, or lowering the activation energy can all speed up a reaction. Temperature mainly changes particle energy, concentration and pressure mainly change collision frequency, surface area changes how much solid is exposed, and catalysts provide an alternate pathway.
These ideas let chemists predict and control reaction speed without changing the overall chemical equation.
Key Facts
- Reaction rate = change in concentration / change in time, often written as rate = Δ[product] / Δt or rate = -Δ[reactant] / Δt.
- Collision theory: reactions occur when particles collide with enough energy and the correct orientation.
- Higher concentration usually increases rate because there are more reactant particles per unit volume and more frequent collisions.
- Higher temperature increases rate because particles move faster and a larger fraction have energy greater than or equal to the activation energy.
- For gases, higher pressure usually increases rate because particles are squeezed into a smaller volume, increasing collision frequency.
- A catalyst increases rate by lowering activation energy, often shown as Ea,catalyzed < Ea,uncatalyzed, and is not consumed overall.
Vocabulary
- Reaction rate
- Reaction rate is the speed at which reactants are used up or products are formed during a chemical reaction.
- Collision theory
- Collision theory explains reaction rate by stating that particles must collide with enough energy and proper orientation to react.
- Activation energy
- Activation energy is the minimum energy particles must have during a collision for a reaction to occur.
- Catalyst
- A catalyst is a substance that speeds up a reaction by providing a lower-energy pathway without being used up overall.
- Surface area
- Surface area is the exposed area of a solid reactant where collisions with other particles can happen.
Common Mistakes to Avoid
- Saying higher concentration makes each particle move faster is wrong because concentration mainly increases the number of particles per volume, not their speed.
- Assuming every collision causes a reaction is wrong because collisions must have enough energy and the correct orientation to be successful.
- Thinking a catalyst changes the amount of product formed is wrong because a catalyst changes how fast equilibrium is reached, not the balanced equation or the final yield for a reversible reaction at equilibrium.
- Ignoring surface area for solids is wrong because a large chunk may have the same mass as powder, but the powder exposes more particles for collisions.
Practice Questions
- 1 A reaction produces 0.80 mol/L of product in 40 s. What is the average reaction rate in mol/L/s?
- 2 A gas reaction is run in a 2.0 L container, then the same amount of gas is compressed to 1.0 L at the same temperature. If the reaction rate is proportional to gas concentration, by what factor does the rate change?
- 3 A student reacts hydrochloric acid with a solid magnesium ribbon, then repeats the reaction using the same mass of magnesium powder. Explain which reaction is faster and why using collision theory.