Subject: Chemistry | Level: GCSE | Exam Board: WJEC
Master the speed of chemical reactions and the delicate balance of dynamic equilibrium. This topic is essential for understanding how industrial processes are optimised to save time and money, and it's a guaranteed high-mark area in your GCSE Chemistry exams.
Revision Notes & Key Concepts
Revision Podcast Transcript
Welcome to your GCSE Chemistry revision podcast. I'm your tutor, and today we're diving into one of the most important and frequently examined topics in your chemistry course: Rate of Chemical Change and Dynamic Equilibrium. Whether you're sitting AQA, Edexcel, or OCR, this topic comes up every single year, and by the end of this episode, you'll have everything you need to tackle it with confidence. So grab your revision notes, get comfortable, and let's get started. Let's begin with the big picture. Chemical reactions happen at different speeds. Some, like an explosion, happen in a fraction of a second. Others, like iron rusting, take years. The rate of reaction tells us how quickly reactants are converted into products. Understanding what controls this rate — and being able to explain it using particle theory — is absolutely central to your exam success. And then there's the second half of this topic: reversible reactions and dynamic equilibrium. This is where chemistry gets really elegant, because we're looking at reactions that can go both ways simultaneously, and the beautiful balance that results. Now let's get into the core concepts. First up: how do we actually measure the rate of a reaction? Examiners love this, and you need to know three key methods. The first method is gas collection. If your reaction produces a gas, you can collect it in a gas syringe or over water in an inverted measuring cylinder. You measure the volume of gas produced at regular time intervals. The rate is faster at the start — when you see the gas collecting quickly — and slows down as reactants are used up. A classic example is marble chips reacting with hydrochloric acid, producing carbon dioxide gas. The second method is measuring loss of mass. If a gas is produced and escapes from the reaction vessel, the total mass of the flask and its contents decreases over time. You place the flask on a balance and record the mass every thirty seconds. The gradient of the mass-versus-time graph tells you the rate. The steeper the gradient, the faster the rate. This is a really clean method because you can see the rate changing in real time. The third method is the precipitation method, also called the disappearing cross experiment. You use a reaction that produces a cloudy precipitate — typically sodium thiosulfate reacting with hydrochloric acid, which produces sulfur. You draw a cross on paper, place the flask on top, and time how long it takes for the cross to disappear from view. The shorter the time, the faster the rate. This is a classic required practical, so make sure you know it inside out. Now, here's the critical thing examiners want you to do: calculate the rate from a graph. The rate at any point is the gradient of the tangent to the curve. At the start of the reaction, the gradient is steepest — that's where the rate is highest. As the reaction proceeds, the gradient decreases, and eventually the line becomes flat when the reaction is complete. A really common mistake is to describe the rate as "the steepness of the line" without actually calculating it. If the question says calculate, you must draw a tangent, find the rise over run, and include units — typically centimetres cubed per second, or grams per second. Right, let's move on to the particle model and collision theory. This is the heart of the topic, and you need to be able to apply it to every single factor that affects rate. The key idea is this: for a reaction to occur, particles must collide with sufficient energy. That minimum energy is called the activation energy. So the rate of reaction depends on two things: the frequency of collisions — how often particles collide — and the energy of those collisions — whether they have enough energy to overcome the activation energy barrier. Let's go through each factor in turn. Temperature: when you increase the temperature, particles move faster. This means they collide more frequently, and crucially, a greater proportion of particles have energy greater than or equal to the activation energy. Both effects increase the rate. In your exam answer, you must mention both collision frequency and the proportion of particles exceeding the activation energy to get full marks. Concentration: when you increase the concentration of a solution, there are more particles per unit volume. This increases the collision frequency, so the rate increases. Note: concentration does not change the energy of individual particles — it only affects how often they meet. Pressure: for reactions involving gases, increasing the pressure has the same effect as increasing concentration. The gas particles are pushed closer together, so there are more particles per unit volume, increasing collision frequency. Surface area: when a solid reactant is broken into smaller pieces — or ground into a powder — the total surface area exposed to the other reactant increases. More surface area means more particles are available to collide at any one time, so the collision frequency increases. Think about it this way: if you have a large marble chip, only the particles on the outside surface can react. But if you crush it into powder, suddenly thousands more particles are exposed. This is why powders react faster than lumps. Catalysts: a catalyst is a substance that increases the rate of a reaction without being used up. It does this by providing an alternative reaction pathway with a lower activation energy. Because the activation energy is lower, a greater proportion of particles have enough energy to react, so the rate increases. A critical point for your exam: a catalyst does not change the collision frequency. It only lowers the activation energy. Also, a catalyst is not consumed in the reaction — it is regenerated. This is why you only need a small amount. Now let's turn to the second major area: reversible reactions and dynamic equilibrium. Some chemical reactions are reversible — the products can react together to reform the original reactants. We write these with a double-headed arrow. A classic example is the reaction between nitrogen and hydrogen to form ammonia: nitrogen plus three hydrogen gives two ammonia, and this can also go in reverse. When a reversible reaction takes place in a closed system — meaning nothing can enter or leave — something fascinating happens. At first, the forward reaction dominates because there are lots of reactants and few products. But as products build up, the reverse reaction starts to happen more and more. Eventually, the rate of the forward reaction equals the rate of the reverse reaction. This is dynamic equilibrium. The word dynamic is crucial. It does not mean the reactions have stopped. Both the forward and reverse reactions are still happening simultaneously — it's just that they're happening at the same rate, so the concentrations of reactants and products remain constant. This is one of the most common misconceptions in this topic, and examiners specifically test it. Now, what happens when you change the conditions at equilibrium? The system responds to oppose the change — this is Le Chatelier's Principle, though at GCSE you don't need to name it. You just need to predict and explain the shift. Temperature: if you increase the temperature, the equilibrium shifts in the direction of the endothermic reaction — the reaction that absorbs heat. This is because the system is trying to reduce the temperature by absorbing the extra heat energy. If you decrease the temperature, the equilibrium shifts in the direction of the exothermic reaction. Concentration: if you increase the concentration of a reactant, the equilibrium shifts to the right — towards the products — to use up the extra reactant. If you remove a product, the equilibrium also shifts to the right to replace it. Pressure: if you increase the pressure, the equilibrium shifts towards the side with fewer moles of gas. This reduces the pressure by reducing the number of gas molecules. If both sides have the same number of moles of gas, changing pressure has no effect on the position of equilibrium. Now let's talk exam technique and common mistakes. This is where marks are won and lost. The number one mistake is describing dynamic equilibrium as a state where the reaction has stopped. It has not stopped. Both reactions are still occurring. Always write: "the forward and reverse reactions occur at the same rate." The second major mistake is forgetting to mention collision frequency or activation energy when explaining rate changes. If a question asks you to explain why increasing temperature increases the rate, you must say: particles move faster, collision frequency increases, and a greater proportion of particles have energy greater than or equal to the activation energy. Missing any of these points costs you marks. The third mistake is confusing catalysts with changing reaction conditions. A catalyst lowers the activation energy — it does not change the temperature, concentration, or pressure. It also does not change the position of equilibrium in a reversible reaction — it just helps the system reach equilibrium faster. The fourth mistake is misreading rate graphs. Remember: the rate is the gradient of the graph, not the total amount produced. A steeper gradient means a faster rate. When the line becomes horizontal, the reaction has stopped — but this does not mean equilibrium has been reached unless the reaction is reversible in a closed system. Now for your quick-fire recall quiz. I'll ask the questions, and you try to answer before I give the answer. Ready? Question one: What is the activation energy? The minimum energy required for a collision to result in a reaction. Question two: Name three methods for measuring the rate of a reaction. Gas collection, loss of mass, and precipitation — the disappearing cross. Question three: What does a catalyst do? It provides an alternative reaction pathway with a lower activation energy, without being used up. Question four: What is dynamic equilibrium? A state in a closed system where the forward and reverse reactions occur at the same rate, so concentrations remain constant. Question five: If you increase the pressure on a gaseous equilibrium system, which way does it shift? Towards the side with fewer moles of gas. Question six: What two things must you mention when explaining why increased concentration increases the rate? More particles per unit volume, and increased collision frequency. How did you do? If you got all six, you're in great shape. If you missed any, go back and re-read those sections. Let's wrap up with a quick summary of the key points to take away. One: rate of reaction measures how quickly reactants are converted to products. Two: rate can be measured by gas collection, loss of mass, or precipitation. Three: rate is calculated from the gradient of a graph — steeper gradient means faster rate. Four: increasing temperature, concentration, pressure, or surface area all increase collision frequency. Temperature also increases the proportion of particles with energy above the activation energy. Five: catalysts lower the activation energy by providing an alternative pathway — they do not increase collision frequency. Six: dynamic equilibrium occurs in a closed system when forward and reverse rates are equal — both reactions are still happening. Seven: changing temperature, concentration, or pressure shifts the equilibrium position to oppose the change. That's everything for today's episode. You've covered collision theory, the five factors affecting rate, methods of measurement, reversible reactions, and dynamic equilibrium. These are all high-value topics that appear regularly in GCSE Chemistry exams. Make sure you can explain each factor using the particle model, and practice drawing and interpreting rate graphs. Good luck with your revision — you've got this!
Key Terms & Definitions
- Rate of Reaction
- The change in concentration of a reactant or product per unit time.
- Activation Energy
- The minimum amount of energy that particles must have to react when they collide.
- Catalyst
- A substance that increases the rate of a chemical reaction without being changed or used up in the reaction.
- Reversible Reaction
- A reaction in which the products can react to reform the original reactants.
- Dynamic Equilibrium
- The point in a reversible reaction where the forward and reverse reactions occur at exactly the same rate in a closed system.
- Closed System
- An apparatus in which no substances can enter or leave during a reaction.
Worked Examples
Worked Example
Question: A student investigated the rate of reaction between marble chips and hydrochloric acid. Explain, in terms of particles, how and why the rate of reaction changes during the reaction. (4 marks)
Solution: Step 1: State what happens to the rate. The rate of reaction decreases as the reaction proceeds. Step 2: Explain why in terms of concentration. This is because the concentration of hydrochloric acid decreases as the acid particles are used up. Step 3: Explain the effect on collisions. Therefore, there are fewer acid particles per unit volume. Step 4: Link to collision frequency. This results in a lower frequency of successful collisions between the acid particles and the marble chips.
Worked Example
Question: The equation for the production of ammonia is: N₂(g) + 3H₂(g) ⇌ 2NH₃(g). The forward reaction is exothermic. Predict and explain the effect of increasing the pressure on the yield of ammonia. (3 marks)
Solution: Step 1: Predict the effect on yield. Increasing the pressure will increase the yield of ammonia. Step 2: Explain in terms of moles of gas. This is because there are 4 moles of gas on the reactant side (1 N₂ + 3 H₂) and only 2 moles of gas on the product side (2 NH₃). Step 3: Link to Le Chatelier's principle. The equilibrium shifts to the right (towards the products) to reduce the pressure by producing fewer molecules.
Worked Example
Question: Explain how a catalyst increases the rate of a reaction. (2 marks)
Solution: Step 1: State the function of a catalyst. A catalyst provides an alternative reaction pathway. Step 2: Explain the effect on energy. This alternative pathway has a lower activation energy.
Practice Questions
Question: A student investigates the reaction between magnesium and dilute hydrochloric acid. Suggest two ways the student could increase the rate of this reaction. (2 marks)
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Question: Explain, in terms of particles and collisions, the effect of increasing the temperature on the rate of a reaction. (3 marks)
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Question: Hydrogen reacts with iodine to form hydrogen iodide in a reversible reaction: H₂(g) + I₂(g) ⇌ 2HI(g). The forward reaction is exothermic. State and explain what happens to the yield of hydrogen iodide if the temperature is increased. (3 marks)
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Question: For the same reaction: H₂(g) + I₂(g) ⇌ 2HI(g), state the effect of increasing the pressure on the position of equilibrium. Give a reason for your answer. (2 marks)
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Question: A student measures the rate of a reaction by collecting the gas produced. How could the student calculate the rate of reaction at exactly 30 seconds from their volume-time graph? (2 marks)
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