Why Does Rate Change?
The rate of a chemical reaction is not fixed — it can be made faster or slower by changing the conditions. Understanding which factors affect rate (and why) is one of the most practical parts of chemistry. Industries use this knowledge to control how fast reactions happen in manufacturing, food production, and medicine.
There are four main factors that affect reaction rate:
- Temperature
- Concentration (or pressure for gases)
- Surface area (particle size)
- Catalysts
Each of these can be explained using collision theory: reactions happen when particles collide with enough energy and in the correct orientation. Any change that increases the number of successful collisions per second increases the rate.
Factor 1: Temperature
Higher temperature = faster reaction.
When you heat a substance, its particles gain kinetic energy and move faster. Faster-moving particles collide more often and — crucially — with more energy. This means a greater proportion of collisions have enough energy to overcome the activation energy barrier.
A useful rule of thumb in chemistry: raising the temperature by 10 °C roughly doubles the reaction rate for many common reactions.
Real world: Cooking food
Cooking is simply speeding up chemical reactions that change raw food into cooked food. Boiling water at 100 °C cooks an egg in about 10 minutes. Pressure cookers raise the boiling point of water to about 120 °C, cooking the same egg in around 4 minutes — roughly twice as fast, consistent with the 10 °C doubling rule.
Temperature affects both the frequency of collisions (particles move faster, so they meet more often) and the energy of those collisions. Both effects together make temperature the most powerful factor affecting reaction rate.
Factor 2: Concentration (and Pressure)
Higher concentration = faster reaction.
Concentration tells us how many particles of a substance are dissolved in a given volume. A higher concentration means more particles in the same space — so particles collide more frequently.
For gases, increasing the pressure has the same effect: squeezing the same number of gas particles into a smaller volume makes collisions more frequent.
| Concentration | Number of particles per litre | Collision frequency |
|---|---|---|
| Low | Few | Low |
| High | Many | High |
Real world: Hydrogen peroxide and wounds
When you pour hydrogen peroxide (H₂O₂) on a wound, it bubbles — oxygen gas is released as it decomposes. The 3% solution sold in pharmacies reacts slowly and safely. A 30% solution (ten times more concentrated) would react violently and dangerously. More H₂O₂ molecules per millilitre means far more successful collisions per second.
Factor 3: Surface Area
Smaller particles = larger surface area = faster reaction.
Reactions between a solid and a liquid (or gas) can only happen at the surface of the solid — the interior cannot react until the outer layers have been used up. If the solid is cut into smaller pieces, more surface is exposed to the other reactant, so more collisions can happen simultaneously.
Think about dissolving sugar: a single large sugar cube dissolves much more slowly in water than the same mass of granulated sugar, because granulated sugar has far more surface exposed to the water.
| Particle Size | Surface Area | Reaction Rate |
|---|---|---|
| Large lumps | Small | Slow |
| Small granules | Medium | Medium |
| Fine powder | Very large | Fast |
Finely divided powders of normally safe materials — flour, sugar, coal dust — can form explosive clouds when mixed with air. Dust explosions in grain silos and coal mines have caused many industrial accidents. The enormous surface area of fine powder allows combustion to happen almost simultaneously throughout the entire cloud.
Factor 4: Catalysts
A catalyst speeds up a reaction without being used up itself.
A catalyst works by providing an alternative reaction pathway that has a lower activation energy. This means more collisions have enough energy to be successful — so the rate increases — even though the temperature has not changed.
Because the catalyst is not permanently changed or consumed, it can catalyse many reaction cycles. A tiny amount of catalyst can therefore speed up a very large amount of reaction.
Real world: Catalytic converters in cars
Car exhaust contains harmful gases such as carbon monoxide (CO) and nitrogen oxides (NOₓ). A catalytic converter contains platinum and palladium metals, which act as catalysts. They speed up the conversion of these toxic gases into harmless nitrogen (N₂), carbon dioxide (CO₂), and water vapour — without being consumed. Without catalytic converters, city air pollution would be far worse.
Comparing All Four Factors
| Factor | What Changes | Why Rate Increases |
|---|---|---|
| Temperature | Particles move faster | More frequent and more energetic collisions |
| Concentration / Pressure | More particles per volume | More frequent collisions |
| Surface area | More surface exposed | More surface for collisions to occur |
| Catalyst | Lower activation energy | Greater proportion of collisions are successful |
Catalysts do not change the overall energy released or absorbed by a reaction, and they do not change the products. They only change the rate at which equilibrium is reached.
Putting It All Together
These factors are routinely combined in industry. For example, the Haber process (making ammonia from nitrogen and hydrogen) uses:
- High temperature (about 450 °C) to increase rate
- High pressure (about 200 atmospheres) to increase concentration of gas
- An iron catalyst to lower activation energy
Each factor contributes to making the industrial reaction economically viable.
Practice Problems
A student tests two marble chip samples with the same mass of hydrochloric acid. Sample A uses large marble lumps; Sample B uses powdered marble. Both use acid of the same concentration and temperature. Which reacts faster and why?
Show answer
Explain how a catalyst increases reaction rate without changing the temperature.