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Some reactions are over in a flash — an explosion, or magnesium fizzing in acid — while others, like iron rusting, take years. The rate of reaction measures how fast a reaction goes, and collision theory explains why changing the conditions can speed a reaction up or slow it down. Understanding this is one of the central goals of Topic C5 of OCR Gateway Science A, because it lets a chemist control how quickly a reaction happens. This lesson defines the rate of reaction, sets out collision theory, and explains each of the factors that affect the rate in terms of collisions.
By the end of this lesson you should be able to define the rate of reaction, state collision theory, and explain how temperature, concentration, pressure, surface area and a catalyst change the rate by affecting collisions.
The rate of reaction is a measure of how fast reactants are used up, or how fast products are formed, in a given time. A fast reaction uses up its reactants quickly; a slow reaction uses them up slowly.
For a reaction to happen at all, the particles of the reactants must come together and react. This simple requirement is the basis of collision theory, which explains everything about rates.
Exam Tip: Define rate as how quickly reactants are used up or products are formed. Any explanation of a change in rate should then be built on collision theory — examiners expect the reasoning, not just the fact that a change speeds things up.
The range of reaction rates in nature is enormous. Some reactions are effectively instantaneous — an explosion of a fuel–air mixture, or a precipitate forming the moment two solutions are mixed. Others are extremely slow — the rusting of iron takes years, and the chemical weathering of rock can take centuries. Most reactions you meet in the laboratory lie somewhere in between, fast enough to follow over a few minutes. The point of studying rates is to understand what makes a reaction fast or slow, so that a chemist can deliberately speed up a useful reaction (to make a product quickly) or slow down an unwanted one (such as food spoiling or metal corroding). Every one of those choices comes back to controlling how often the particles collide and how energetically.
Collision theory states that for particles to react, they must:
A collision that has at least the activation energy is a successful collision and leads to a reaction. A collision with less than the activation energy is unsuccessful — the particles simply bounce apart unchanged.
So the rate of a reaction depends on two things:
Anything that makes collisions more frequent, or makes a greater proportion successful, increases the rate.
The activation energy is the key to why not every collision leads to a reaction. Even in a fast reaction, the great majority of collisions are unsuccessful — the particles simply do not have enough energy to break the existing bonds, so they bounce apart unchanged. Only the small fraction of collisions that have at least the activation energy actually react. This is why a reaction with a high activation energy is slow at room temperature: very few collisions are energetic enough. Anything that increases the fraction of successful collisions — chiefly raising the temperature, or lowering the activation energy with a catalyst — therefore has a powerful effect on the rate.
Exam Tip: The two conditions for a successful collision are: particles must collide and collide with at least the activation energy. Build every rate explanation from these — "more frequent collisions" and/or "a greater proportion of successful collisions".
Five factors increase the rate of reaction, and each can be explained by collision theory.
| Factor (increase) | Effect on collisions | Effect on rate |
|---|---|---|
| Higher temperature | Particles move faster → collisions more frequent and more have the activation energy (more energetic) | Increases |
| Higher concentration (solutions) | More particles in the same volume → collisions more frequent | Increases |
| Higher pressure (gases) | Gas particles squeezed closer → collisions more frequent | Increases |
| Larger surface area (solids) | More particles exposed → collisions more frequent | Increases |
| Catalyst | Provides a pathway of lower activation energy → a greater proportion of collisions succeed | Increases |
Raising the temperature speeds a reaction up in two ways. The particles gain kinetic energy, so they move faster and collide more frequently. More importantly, a greater proportion of the collisions now have at least the activation energy, so more of them are successful. The energy effect is the larger of the two — which is why even a modest rise in temperature can dramatically increase the rate.
Increasing the concentration of a solution means there are more reactant particles in the same volume, so the particles collide more often — more frequent collisions means a faster rate. For gases, increasing the pressure squeezes the same number of particles into a smaller volume, which has exactly the same effect: the particles are closer together, so they collide more frequently and the rate increases.
Breaking a solid into smaller pieces (or grinding it to a powder) increases its surface area. Only particles on the surface of a solid can be hit by the other reactant, so a larger surface area means more particles are exposed and available to collide. This makes collisions more frequent and increases the rate. This is why powdered solids react faster than large lumps — and why fine dusts can react dangerously fast.
Exam Tip: Explain concentration, pressure and surface area all the same way: they put more particles in reach, so collisions become more frequent. Explain temperature differently: it makes collisions more frequent and (mainly) more energetic, so more reach the activation energy.
A catalyst speeds up a reaction without being used up. It works by providing a different reaction pathway with a lower activation energy. Because the activation energy is lower, a greater proportion of collisions now have enough energy to be successful, so the reaction goes faster. The catalyst is not used up — it is unchanged at the end and can be used again.
A reaction profile shows this clearly: the catalysed pathway has a lower hump, while the reactant and product energy levels (and so the overall energy change) stay the same.
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