Some reactions take a fraction of a second; others take millions of years. Learn what controls reaction speed, why collision theory explains every factor, and how to interpret rate-of-reaction graphs in exams.
The rate of reaction measures how quickly reactants are converted into products. You can express it as the amount of reactant used up per unit time, or as the amount of product formed per unit time.
Rate = Amount of product formed ÷ Time taken
A fast reaction (like an explosion) is over in milliseconds. A slow reaction (like iron rusting) takes days or years. Understanding what controls this speed has enormous practical value — it is why food is refrigerated, why car engines use catalytic converters, and why tablets dissolve faster when crushed.
Collision theory explains all rate-of-reaction factors in one framework: for a reaction to occur, reactant particles must collide with each other. But not every collision leads to a reaction — the colliding particles must also have enough energy (called the activation energy) and must hit each other in the correct orientation.
So the rate of reaction depends on two things:
Every factor that increases the rate of reaction does so by increasing collision frequency, increasing collision energy, or both. Every factor that decreases the rate works the other way.
Increasing the concentration of a dissolved reactant increases the number of particles in a given volume. More particles in the same space means more frequent collisions — so the rate increases.
In everyday terms: a crowded room means you bump into people more often. A dilute solution is like a nearly empty room — collisions are rare.
Marble chips (calcium carbonate) react with hydrochloric acid: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. Using 2 mol/dm³ HCl produces carbon dioxide much faster than using 0.5 mol/dm³ HCl, because the higher concentration puts more HCl molecules in contact with the marble surface per second.
Increasing the temperature increases the rate because it does two things simultaneously: it makes particles move faster (more frequent collisions), and — more importantly — it gives more particles enough energy to exceed the activation energy (more successful collisions).
A rough rule of thumb: raising temperature by 10 °C approximately doubles the rate of many reactions. This is why reactions in a cold lab run noticeably slower than at room temperature.
Sodium thiosulfate reacting with hydrochloric acid produces a sulfur precipitate that makes the solution go cloudy. This reaction is used to measure rate by timing how long it takes for a cross drawn on paper beneath the flask to disappear. At 50 °C the cross disappears much faster than at 20 °C.
Increasing the surface area of a solid reactant increases the rate because more of the solid is exposed to the other reactant. Only particles on the surface can collide with other reactants — particles buried inside the solid cannot react until those on the outside have been used up.
Cutting a solid into smaller pieces, or using a powder instead of lumps, dramatically increases the surface area without changing the total amount of reactant.
Large marble chips react slowly with acid. Powdered calcium carbonate reacts much faster with the same volume and concentration of acid — even though the total mass is identical. Flour mills must control dust carefully because very fine flour particles suspended in air can ignite explosively (a real industrial hazard).
A catalyst is a substance that increases the rate of reaction without being used up in the reaction. Catalysts work by providing an alternative reaction pathway with a lower activation energy. More particles have enough energy to react via this easier path, so collisions are more frequently successful.
Crucially, a catalyst does not change how much product is made — it only changes how quickly the product is made. After the reaction is complete, the catalyst can be recovered unchanged.
For reactions involving gases, increasing the pressure pushes gas particles closer together, increasing collision frequency. The effect is equivalent to increasing the concentration of a dissolved reactant. Pressure only affects gaseous reactants — it has no effect on solids or liquids.
A typical rate graph plots product formed (or reactant remaining) against time. Key features to know:
Question: A student doubles the concentration of hydrochloric acid used to react with zinc. Explain why the rate increases, using collision theory.
Answer: Doubling the concentration puts more HCl molecules in the same volume of solution. This increases the frequency of collisions between HCl molecules and zinc atoms. More collisions per second means more successful collisions per second, so the rate increases.
Question: Hydrogen peroxide decomposes slowly at room temperature. Adding manganese dioxide speeds up the reaction considerably. What is manganese dioxide acting as? How does it work?
Answer: Manganese dioxide is a catalyst. It provides an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions between H₂O₂ molecules now have enough energy to be successful. The manganese dioxide is not used up — it can be filtered off unchanged at the end.
Question: Two experiments use the same mass of marble chips with hydrochloric acid. Experiment A uses large chips; Experiment B uses powder. Both produce the same total volume of CO₂. Describe how the two curves on a volume-vs-time graph would differ.
Answer: Experiment B (powder) would have a steeper initial gradient — the reaction starts faster because the powder has greater surface area, giving more frequent collisions. However, both curves level off at the same final volume because the total amount of marble (and therefore the amount of CO₂ that can form) is identical in both experiments.
Question: Explain, in terms of particle energy, why increasing temperature increases the rate of reaction more than you would expect from just the increase in collision frequency.
Answer: At higher temperature, particles move faster, so they collide more often (increased frequency). More importantly, a much larger proportion of particles now have energy that meets or exceeds the activation energy threshold. These two effects combine, which is why a 10 °C rise can double the rate — the energy effect is larger than the frequency effect alone.
Rates of reaction is one of the most reliably examined topics in O Level chemistry. Every question — whether it involves graphs, experiments, or explanations — can be answered by applying collision theory consistently. Practise stating the factor, the collision theory explanation, and the result on rate, and you will cover almost every mark available on this topic.