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Rate of reaction is affected by concentration, pressure, surface area, temperature and the presence of a catalyst. Free notes and 9 questions answered.
What affects how fast reactions happen? is part of Rate of reaction in GCSE AQA Higher Chemistry.
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Rate of reaction is affected by concentration, pressure, surface area, temperature and the presence of a catalyst. Free notes and 9 questions answered.
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Use this lesson to stay inside Rate of reaction while connecting the detail back to the wider Topic 6: The rate and extent of chemical change topic.
Read the core explanations from What affects how fast reactions happen? before testing yourself from memory.
Drop a piece of magnesium in acid and it fizzes violently in seconds. Leave an iron nail outside and it takes months to rust. Same idea - a metal reacting - but completely different speeds. What's going on?
The answer lies in understanding what makes particles react. Think of chemical reactions like a busy dance floor - particles are constantly moving and bumping into each other. But not every collision leads to a reaction. Only the right collisions, with enough energy and the correct orientation, actually break bonds and form new ones.
Here's the key insight: if we can increase the number of successful collisions per second, we speed up the reaction. If we decrease them, the reaction slows down. Five main factors control this collision rate, and understanding them gives you complete control over reaction speed.
Picture a crowded nightclub. The more people (concentration), the more collisions. Make the dance floor smaller (pressure for gases), more collisions again. Now imagine everyone moves faster (temperature) - even more collisions. These everyday analogies help us understand the five factors that control reaction rates.
1. Concentration of reactants in solution: More particles in the same volume means more frequent collisions. Double the concentration, roughly double the collision rate. Think of it like rush hour traffic - more cars, more chance of accidents.
2. Pressure of reacting gases: Squashing gas particles into a smaller space has the same effect as increasing concentration. Higher pressure = particles closer together = more collisions per second.
3. Surface area of solid reactants: Reactions happen at surfaces where particles meet. Powder has millions more surface particles than a lump. More surface = more collision sites = faster reaction.
Worked example: Surface Area Calculation
A 2cm × 2cm × 2cm cube is cut into 1cm × 1cm × 1cm cubes. How does the total surface area change?
Answer: Surface area doubles from 24 cm² to 48 cm²
4. Temperature: Heat gives particles more kinetic energy. They move faster, collide more often, and more collisions have enough energy to overcome the activation energy barrier. A 10°C rise typically doubles the reaction rate.
5. Presence of catalysts: These clever substances provide an alternative reaction pathway with lower activation energy. More collisions now have enough energy to react, even though the catalyst itself remains unchanged.
Let's zoom in to the molecular level and see exactly what happens when we change these factors. Imagine you're watching particles through a super-powerful microscope as they dance around, collide, and occasionally react.
In a dilute solution, particles are spread out like people in a nearly empty shopping centre. Collisions are rare. Add more reactant (increase concentration) and suddenly it's like Black Friday - particles are everywhere, bumping into each other constantly.
For gases, imagine particles as ping-pong balls bouncing in a box. Increase the pressure (make the box smaller) and the balls hit each other more often. The collision frequency goes up dramatically.
Surface area is like the difference between trying to dissolve a whole sugar cube versus granulated sugar in your tea. The cube has sugar particles locked inside - only the surface ones can react with water. Granulated sugar exposes millions more particles to the water, so it dissolves much faster.
Temperature is the energy controller. Cold particles move sluggishly - even when they collide, they often just bounce off without reacting. Hot particles zip around with high kinetic energy. When they collide, they hit hard enough to break bonds and form new ones.
Catalysts are the master key of chemistry. They don't push particles together or give them more energy. Instead, they provide a completely different route for the reaction - like finding a mountain pass instead of climbing straight over the peak.
Think of activation energy as a high wall that particles must jump over to react. Most particles don't have enough energy to make this jump. A catalyst builds a ramp or tunnel through the wall - suddenly, many more particles can get across and react.
The brilliant thing about catalysts is they're not consumed in the reaction. They help the reaction happen, then regenerate ready to help again. It's like having a friend who helps you climb over a fence, then climbs back to help the next person.
Examples are everywhere: enzymes in your body catalyse thousands of reactions, platinum in car exhausts breaks down harmful gases, and iron speeds up ammonia production in the Haber process. Without catalysts, many reactions would be too slow to be useful.
Rate of reaction questions follow predictable patterns. Master these techniques and you'll spot the marks coming from miles away. The key is linking particle behaviour to observable changes.
Common exam mistakes to avoid: Don't say catalysts 'speed up particles' - they lower energy barriers. Don't confuse concentration with amount - it's particles per unit volume that matters. Don't forget that pressure only affects gases, not solids or liquids.
Worked example: Perfect Rate Explanation Structure
Explain why increasing temperature increases the rate of reaction between magnesium and hydrochloric acid. [4 marks]
Answer: Temperature → more kinetic energy → faster particles → more frequent collisions → more collisions exceed activation energy → faster rate
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Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
The five factors that affect the rate of reaction are the concentration of reactants in solution, the pressure of reacting gases, the surface area of solid reactants, temperature, and the presence of a catalyst. Each one changes the number of successful collisions per second.
Rate of reaction is how fast reactants are used up or products are formed in a chemical reaction, measured in units such as g/s or mol/s.
Increasing temperature increases the rate because particles gain more kinetic energy, move faster and collide more frequently, and more of the collisions have energy greater than the activation energy. So more successful collisions occur per second.
Concentration is the number of particles per unit volume, not the total amount of reactant. More particles in the same volume means more frequent collisions, so doubling the concentration roughly doubles the collision rate.
Some reactions happen faster than others because they have more successful collisions per second. Only collisions with enough energy and the correct orientation break bonds and form new ones; magnesium in acid fizzes in seconds, while an iron nail takes months to rust.
A catalyst increases the rate of reaction by providing an alternative pathway with lower activation energy. More collisions now have enough energy to react, even though the catalyst itself is not used up and does not give particles more energy.
A 10°C rise typically doubles the reaction rate. Heating gives particles more kinetic energy, so they move faster, collide more often, and more collisions have enough energy to overcome the activation energy barrier.
Pressure only affects reactions involving gases, not solids or liquids. Higher pressure squashes gas particles closer together, which has the same effect as increasing concentration, so there are more collisions per second.
Cutting a solid into smaller pieces increases its total surface area. A 2cm cube has a surface area of 24 cm², but cut into eight 1cm cubes the total is 8 × 6 = 48 cm², so the surface area doubles and the reaction is faster.