Topic 6 The Rate and extent of chemical change
Unlock the secrets behind the speed of chemical reactions in Topic 6 of the AQA GCSE Chemistry specification: The Rate and Extent of Chemical Change. Ever wondered why some reactions are explosive while others take years? In this essential topic, you'll master how to calculate reaction rates, delve into the crucial principles of collision theory, and explore the dynamic world of reversible reactions and equilibrium. By understanding these concepts, you'll gain the power to predict and control how chemical reactions happen.
The rate of a chemical reaction is the measure of how fast reactants are converted into products. Some reactions are very fast, like explosions, while others are extremely slow, like the rusting of iron.
We can calculate the mean (average) rate of a reaction by measuring how much of a reactant is used up or how much of a product is formed over a specific period.
The formulas are:
Mean rate of reaction = Amount of reactant used / Time taken
Mean rate of reaction = Amount of product formed / Time taken
The amount can be measured in mass (grams, g), volume (cubic centimetres, cm³), or moles (mol). This means the units for the rate of reaction will typically be g/s, cm³/s, or mol/s.
Exam Tip: Always check the units given in the question! If a question gives you the mass of a product in grams and the time in seconds, your answer for the rate must be in g/s.
To calculate the rate, we first need to measure the change. Here are three common methods:
Precipitation and Colour Change: This method involves timing how long it takes for a solution to become opaque (cloudy) or change colour. A precipitate is a solid that forms in a solution during a chemical reaction.
Example: In the reaction between sodium thiosulfate and hydrochloric acid, a yellow sulfur precipitate forms. You can time how long it takes for a cross drawn on paper under the conical flask to disappear.
Change in Mass: This method is used when a gas is produced in a reaction. The reaction flask is placed on a mass balance, and the loss in mass is recorded over time as the gas escapes.
Example: When calcium carbonate reacts with hydrochloric acid, carbon dioxide gas is released. The mass of the flask and its contents will decrease.
Volume of Gas Produced: If a gas is produced, you can collect it in a gas syringe and measure the volume produced at regular time intervals.
Example: When magnesium reacts with dilute acid, hydrogen gas is produced and can be collected and measured.
Exam Tip: If a question asks you to suggest a method for measuring the rate of a reaction, look at the state symbols in the chemical equation. If you see (g) for a gas on the product side, measuring the volume of gas produced or the change in mass are both excellent answers.
Graphs are often used to show how the amount of reactant or product changes over time. The steeper the line on the graph, the faster the rate of reaction. The reaction stops when the line becomes horizontal (flat), as no more product is being formed or reactant is being used up.
To find the rate of reaction at a specific point in time, you need to calculate the gradient (steepness) of the curve at that point.
Here’s how:
Draw a Tangent: Use a ruler to draw a straight line that just touches the curve at the time you are interested in.
Draw a Triangle: Make this tangent the longest side of a large right-angled triangle.
Calculate the Gradient: The rate of reaction is the gradient of the tangent.
Gradient = Change in Y / Change in X
Example Calculation:
Change in Y (Volume of product) = 48 cm³
Change in X (Time) = 60 s
Rate = 48 / 60 = 0.8 cm³/s
Exam Tip: When drawing the triangle to calculate the gradient, make it as large as possible. This will make your readings more accurate and reduce the chance of making a calculation error, securing you the marks!
The speed of a chemical reaction can be changed. Understanding how to control the rate of reaction is crucial for industrial processes, like manufacturing medicines or fertilisers, where producing the product quickly and efficiently saves money.
There are five key factors that can be changed to alter the rate of a chemical reaction:
Temperature
Concentration (of solutions) or Pressure (of gases)
Surface area (of solid reactants)
The presence of a catalyst
To understand how these factors work, we need to use collision theory. This theory states two things must happen for a reaction to occur:
Particles must collide with each other.
The colliding particles must have enough energy to react.
The minimum amount of energy that particles need to react is called the activation energy (). If particles collide with less energy than the activation energy, they will just bounce off each other and no reaction will happen. A successful collision that leads to a chemical reaction is often called an effective collision.
So, to speed up a reaction, we need to increase the number of effective collisions. We can do this in two ways:
Increase the frequency of collisions (make particles collide more often).
Increase the energy of the collisions (make each collision more likely to be successful).
Exam Tip: When asked to explain how a factor affects the rate of reaction, you MUST refer to collision theory. A perfect answer will mention both the frequency and/or the energy of the collisions and link this to the activation energy.
Increasing the temperature increases the rate of reaction.
How it works: When you heat particles, they gain more kinetic energy and move faster. This means they will collide more frequently. More importantly, the collisions will be more energetic, so a much higher proportion of the collisions will have energy equal to or greater than the activation energy, leading to more effective collisions.
Exam Tip: A common mistake is to only say that particles collide more often. The main reason increasing the temperature speeds up a reaction is that the collisions are more energetic. Make sure you mention both points for full marks!
Increasing the concentration of a reactant in a solution, or increasing the pressure of reacting gases, increases the rate of reaction.
How it works: Increasing the concentration or pressure means there are more reactant particles packed into the same volume. This makes collisions between reactant particles more frequent, which leads to more effective collisions over a given time.
Exam Tip: Be precise with your language. For reactions in solution, use the word concentration. For reactions involving gases, use the word pressure. The explanation for both is the same: more particles in a given volume leads to more frequent collisions.
Increasing the surface area of a solid reactant increases the rate of reaction. For example, a powder will react faster than a large lump of the same solid.
How it works: Breaking a solid into smaller pieces exposes more of its particles to the other reactants. This increases the surface area to volume ratio. With more particles exposed, the frequency of collisions between the reactant particles increases, speeding up the reaction.
Exam Tip: If a question involves a solid reacting (e.g., a metal in acid, or a carbonate in acid), a guaranteed way to increase the rate is to crush the solid into a powder. Remember to explain why this works using the terms ‘surface area’ and ‘frequency of collisions’.
A catalyst is a substance that speeds up a chemical reaction but is not used up in the reaction itself. Because it isn’t consumed, it can be used over and over again. Enzymes are biological catalysts that speed up reactions in living things.
How it works: Catalysts work by providing an alternative reaction pathway that has a lower activation energy (). With a lower energy barrier to overcome, more of the colliding particles will have sufficient energy to react, meaning the frequency of effective collisions increases.
Exam Tip: A reaction profile diagram showing two “humps” – a higher one for the uncatalysed reaction and a lower one for the catalysed reaction – is a very common exam question. You must be able to label the activation energy for both pathways and explain that the catalyst provides a different route with a lower activation energy.
Why do some reactions happen at the blink of an eye, while others take years? The answer lies in collision theory, which is the fundamental idea that explains the speed of all chemical reactions.
For a chemical reaction to happen, two key conditions must be met:
Particles must collide: The reactant particles (atoms, molecules, or ions) have to physically bump into each other.
Collisions must be energetic: The particles must collide with at least a certain minimum amount of energy.
A collision that meets both of these conditions and results in a reaction is called a successful collision or an effective collision. To increase the rate of a reaction, we simply need to increase the number of successful collisions that happen every second.
Exam Tip: When explaining reaction rates, always use the phrase “successful collisions”. For example, “This increases the frequency of successful collisions.” It shows the examiner you understand that not all collisions lead to a reaction.
Activation energy is the scientific term for the “minimum amount of energy” needed for a reaction to occur.
Definition: The activation energy () is the minimum energy that colliding particles must have in order to react.
Think of it as an energy barrier or a hurdle that the reactants must get over to become products. If the colliding particles don’t have enough energy to overcome this barrier, they just bounce off each other without reacting.
Exam Tip: You must be able to label the activation energy on a reaction profile diagram. It’s the “hump” that goes from the reactants’ energy level to the peak of the curve.
The five factors that change the rate of reaction all work by changing the number of successful collisions. They do this by either increasing the frequency of collisions or the energy of collisions.
Temperature: Increasing the temperature gives particles more kinetic energy. They move faster, leading to more frequent and, more importantly, more energetic collisions. A higher proportion of collisions will overcome the activation energy.
Concentration / Pressure: Increasing the concentration or pressure packs more particles into the same volume. This doesn’t make the particles more energetic, but it does make collisions more frequent.
Surface Area: Increasing the surface area of a solid (e.g., by crushing it into a powder) exposes more particles. This leads to more frequent collisions.
Exam Tip: In an exam, if you are asked to explain how temperature, concentration, or surface area affects the rate, your answer must be linked back to collision theory. State whether the factor increases the frequency or energy of collisions and why this leads to more successful collisions per second.
Catalysts are clever chemicals that increase the reaction rate in a unique way.
Function: A catalyst speeds up a reaction without being chemically changed or used up itself.
Mechanism: A catalyst works by providing an alternative reaction pathway that has a lower activation energy ().
By lowering the energy “hurdle,” a much larger proportion of the colliding particles now have enough energy to react successfully. This dramatically increases the rate of reaction.
Exam Tip: A common mistake is to say that catalysts give particles more energy. This is incorrect! Catalysts do not change the particles; they change the reaction pathway. The key phrase to remember is: “Catalysts provide an alternative pathway with a lower activation energy.”
Sometimes, a chemical reaction is too slow for practical use. In these cases, we can often add a special substance to speed it up. This substance is called a catalyst.
A catalyst is a substance that increases the rate of a chemical reaction but is not used up in the reaction itself.
This is a key definition! Because the catalyst isn’t consumed, a small amount can be used to process a huge amount of reactants, which is very useful in industrial chemistry.
Specificity: Different reactions require different catalysts. There is no “one size fits all” catalyst.
Biological Catalysts: In living organisms, proteins called enzymes act as biological catalysts. For example, enzymes in your digestive system break down food much faster than would otherwise be possible.
Exam Tip: When asked to define a catalyst, you must include two points for full marks: 1) it speeds up the rate of a reaction, and 2) it is not used up during the reaction.
Catalysts work by making it easier for a reaction to happen. They do this by changing the route the reaction takes.
Mechanism: A catalyst provides an alternative reaction pathway that has a lower activation energy ().
By lowering the activation energy “hurdle,” a much larger proportion of the colliding particles will have enough energy to react successfully. This increases the frequency of successful collisions, leading to a faster reaction rate.
The reaction profile below shows how a catalyst lowers the activation energy barrier.
Exam Tip: This is a crucial point that is often tested. A common mistake is to say catalysts “give particles more energy.” This is wrong. Catalysts don’t change the particles at all. They change the pathway of the reaction. Always use the phrase: “A catalyst provides an alternative pathway with a lower activation energy.”
Catalysts are essential in industry because they save time and money, and help protect the environment. Many industrial processes (like making ammonia for fertilisers) would not be economically viable without them.
The advantages include:
Lower Energy Costs: By lowering the activation energy, reactions can happen at lower temperatures and pressures. This saves a huge amount of money on energy bills.
Increased Efficiency: Products can be made much more quickly, increasing the output of a factory.
Environmental Benefits: Using less energy means burning fewer fossil fuels, which reduces the emission of pollutants like carbon dioxide (a greenhouse gas) and sulfur dioxide (which causes acid rain).
Exam Tip: For higher-level questions, you may be asked to explain the economic or environmental advantages of using a catalyst. Link your answer back to the fact that catalysts allow reactions to run at lower temperatures and pressures.
Most of the chemical reactions you’ve seen so far go in one direction only; the reactants turn into products, and the reaction stops. However, some reactions are reversible, meaning they can go in both directions.
In a reversible reaction, the products can react with each other to re-form the original reactants. This means the reaction can proceed forwards (reactants ⟶ products) and backwards (products ⟶ reactants) at the same time.
To show that a reaction is reversible, we use a special double arrow symbol: ⇌.
So, a general reversible reaction is written as: A + B ⇌ C + D
The direction of the reaction can be changed by altering the conditions, such as temperature or pressure.
Exam Tip: Make sure you use the correct symbol (⇌) in your exam answers for reversible reactions. Using a normal arrow (⟶) would be incorrect and could lose you marks.
All chemical reactions involve an energy change. They are either exothermic (release energy) or endothermic (take in energy). For reversible reactions, there’s a simple rule:
If a reaction is exothermic in one direction, it must be endothermic in the opposite direction.
The amount of energy transferred is exactly the same in both directions.
Example 1: Ammonium Chloride
Heating solid ammonium chloride causes it to break down into two gases: ammonia and hydrogen chloride. This is the forward reaction and it is endothermic (it takes in heat). ammonium chloride ⇌ ammonia + hydrogen chlorideNH₄Cl(s) ⇌ NH₃(g) + HCl(g)
If you let the hot gases cool down, they react together to form solid ammonium chloride again. This is the reverse reaction and it is exothermic (it releases heat).
Hydrated salts are crystals that have water molecules locked into their structure. Heating them removes this water, leaving an anhydrous (without water) salt. This change is often reversible.
Blue, hydrated copper(II) sulfate crystals can be heated to produce a white, anhydrous powder and water vapour. This forward reaction is endothermic. hydrated copper(II) sulfate ⇌ anhydrous copper(II) sulfate + water
If you add a few drops of water to the white anhydrous powder, the reverse reaction happens. It turns back into the blue hydrated crystals and the reaction is exothermic (it gets hot). This is a common test for the presence of water.
Exam Tip: Questions about the copper(II) sulfate reaction are very common. Remember: Blue to White, Endothermic (when heating) and White to Blue, Exothermic (when adding water). The colour change is the key observation.
If a reversible reaction takes place in a closed system, it will eventually reach equilibrium.
A closed system means that no substances can get in or out.
At equilibrium, the forward and reverse reactions are still happening, but they are occurring at exactly the same rate.
Because the rates are equal, the overall concentrations of the reactants and products remain constant.
It’s called a dynamic equilibrium because the reactions haven’t stopped – they are just perfectly balanced. Think of two people on a seesaw, perfectly balanced in the middle. They are both still there, but their net movement is zero.
Exam Tip: To get the marks for defining dynamic equilibrium, you must state that the rate of the forward reaction is equal to the rate of the reverse reaction and that this occurs in a closed system. Simply saying “the amounts are the same” is not enough.
When a reversible reaction occurs, there is always an energy change. Understanding this relationship is key to controlling the reaction.
The fundamental rule for energy in reversible reactions is:
If a reaction is exothermic in one direction, it must be endothermic in the opposite direction.
This means if the forward reaction gives out heat (gets hotter), the reverse reaction must take in the exact same amount of heat (gets colder).
Exam Tip: You must remember this rule. A common question might give you the energy change for the forward reaction and ask about the reverse reaction. It will always be the same number but with the opposite sign.
This is a classic reversible reaction that shows both an energy change and a colour change.
First, let’s define some key terms:
Hydrated: A salt that contains water molecules locked within its crystal structure (water of crystallisation).
Anhydrous: A salt that has had its water of crystallisation removed, usually by heating.
The reversible reaction is: hydrated copper(II) sulfate ⇌ anhydrous copper(II) sulfate + water CuSO₄·5H₂O(s) ⇌ CuSO₄(s) + 5H₂O(l)
Forward Reaction (Endothermic): If you heat the blue crystals of hydrated copper(II) sulfate, they turn into a whitepowder of anhydrous copper(II) sulfate. This reaction takes in heat from the surroundings, so it is endothermic.
Reverse Reaction (Exothermic): If you add a few drops of water to the white anhydrous powder, it turns back into the blue hydrated crystals. This reaction gives out heat, so it is exothermic.
Exam Tip: You must know the terms anhydrous and hydrated and the specific colours for the copper(II) sulfate test. Remember: Blue = Hydrated, White = Anhydrous. This reaction is the chemical test for water.
The decomposition of ammonium chloride is another good example.
The reversible reaction is: ammonium chloride ⇌ ammonia + hydrogen chloride NH₄Cl(s) ⇌ NH₃(g) + HCl(g)
Forward Reaction (Endothermic): Heating solid ammonium chloride breaks it down into ammonia gas and hydrogen chloride gas. This reaction is endothermic as it requires heat energy to proceed.
Reverse Reaction (Exothermic): As the hot ammonia and hydrogen chloride gases cool down, they react together to re-form solid ammonium chloride. This reaction is exothermic, releasing heat energy.
In a reversible reaction happening inside a sealed container, a special point is eventually reached where the reaction appears to have stopped. This state of balance is called equilibrium.
Equilibrium is reached in a reversible reaction when the rate of the forward reaction is exactly equal to the rate of the reverse reaction
Imagine a shop where customers are entering and leaving at the same rate. The total number of people inside the shop stays constant, even though individuals are constantly moving in and out. This is what happens at equilibrium.
Exam Tip: This definition is essential. To get full marks, you must state that the rates of the forward and reverse reactions are equal. Just saying the reactions are “balanced” is not specific enough.
A system at equilibrium has two key features:
The reaction is dynamic: The forward and reverse reactions have not stopped
Concentrations are constant: Because the forward and reverse reactions are happening at the same rate, the overall amounts (concentrations) of the reactants and products do not change
Exam Tip: Questions often test the meaning of “dynamic” in this context. It means the reactions are still ongoing. A common mistake is to think that everything stops at equilibrium, which is incorrect.
Equilibrium can only be achieved in a closed system
A closed system is one where no substances can escape or be added
If the system is open, products (especially gases) can escape, which means the reverse reaction can’t happen at the same rate, and equilibrium will never be reached.
Exam Tip: If a question asks why equilibrium is only reached in a closed system, explain that it prevents reactants or products from escaping, allowing the reverse reaction to occur at the same rate as the forward reaction.
Once a reversible reaction reaches equilibrium, the amounts of reactants and products stay constant. However, if we change the conditions (like temperature, pressure, or concentration), this balance can be disturbed. The system will then adjust itself to a new equilibrium position. This is all explained by Le Chatelier’s Principle.
Le Chatelier’s Principle states that:
If a change is made to the conditions of a system at equilibrium, the system will move to counteract the change.
Think of it as the system trying to “undo” whatever change you make to it.
If the system favours the forward reaction, we say the equilibrium shifts to the right. This will increase the amount of products.
If the system favours the reverse reaction, we say the equilibrium shifts to the left. This will increase the amount of reactants.
Exam Tip: You must be able to state Le Chatelier’s Principle clearly. The key idea is that the system opposes or counteracts the change you impose on it. Using these keywords will score you the marks.
When a reversible reaction is at equilibrium, the amounts of reactants and products are balanced. If you change the concentration of any substance, the system will adjust to restore this balance. The rules for this are explained by Le Chatelier’s Principle.
Le Chatelier’s Principle states: When a change is made to the conditions of a system at equilibrium, the system automatically moves to oppose the change.
When you change the concentration of a substance, the equilibrium “shifts” to either the right (favouring the forward reaction) or the left (favouring the reverse reaction) to counteract what you’ve just done.
Exam Tip: The key is that the system tries to “undo” the change. If you add something, it tries to use it up. If you remove something, it tries to make more of it.
If you INCREASE the concentration of a reactant:
The system will try to remove the extra reactant.
It favours the forward reaction to use up the reactant.
The equilibrium shifts to the right, and more product is formed.
If you DECREASE the concentration of a reactant:
The system will try to replace the reactant that has been removed.
It favours the reverse reaction to make more reactant.
The equilibrium shifts to the left, and less product is formed.
The same principle applies if you change the concentration of the products.
If you DECREASE the concentration of a product (by removing it):
The system will try to replace the product that has been removed.
It favours the forward reaction to make more product.
The equilibrium shifts to the right, increasing the overall yield of the product.
If you INCREASE the concentration of a product:
The system will try to remove the extra product.
It favours the reverse reaction to use up the product.
The equilibrium shifts to the left, and more reactants are formed.
Exam Tip: In industrial processes, products are often removed from the reaction vessel as they are made. This forces the equilibrium to constantly shift to the right to produce more product, resulting in a much higher yield. This is a common application question.
Consider this reversible reaction:
ICl(l) + Cl₂(g) ⇌ ICl₃(s) (dark brown) + (gas) ⇌ (yellow solid)
If we increase the concentration of Cl₂ (a reactant), the equilibrium will shift to the right to use it up. This will produce more yellow ICl₃ solid.
Temperature is a critical factor in controlling the position of equilibrium. Its effect is predicted by Le Chatelier’s Principle, but to use it, you must first know if the forward reaction is exothermic or endothermic.
Le Chatelier’s Principle states: When a change is made to the conditions of a system at equilibrium, the system automatically moves to oppose the change.
Remember, in any reversible reaction, one direction is exothermic (releases heat) and the opposite direction is endothermic (absorbs heat). The system uses this to counteract temperature changes.
The rules are:
If you INCREASE the temperature (add heat): The system will try to cool down by absorbing the extra heat. It favours the endothermic reaction. The equilibrium shifts in the endothermic direction.
If you DECREASE the temperature (remove heat): The system will try to warm up by releasing heat. It favours the exothermic reaction. The equilibrium shifts in the exothermic direction.
Exam Tip: A simple way to remember this: Heat it, it cools (endothermic). Cool it, it heats (exothermic). The system always does the opposite of what you do to it.
Consider the reaction between iodine monochloride and chlorine. The forward reaction is exothermic.
ICl(l) + Cl₂(g) ⇌ ICl₃(s) (dark brown) ⇌ (yellow)
If we increase the temperature:
The system will oppose the change by absorbing heat.
It will favour the endothermic reaction.
Since the forward reaction is exothermic, the reverse reaction must be endothermic.
Therefore, the equilibrium will shift to the left.
This produces more of the dark brown reactant (ICl), and the mixture will become browner.
The Haber process, used to make ammonia, is a vital industrial reaction. The forward reaction is exothermic.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
If we decrease the temperature:
The system will oppose the change by releasing heat.
It will favour the exothermic reaction.
The forward reaction is exothermic.
Therefore, the equilibrium will shift to the right.
This increases the yield of ammonia (NH₃).
Exam Tip: In industrial processes like the Haber Process, there is a trade-off between rate and yield. A low temperature gives a high yield of ammonia (good!), but it also makes the reaction very slow (bad!). A high temperature makes the reaction fast (good!), but gives a low yield (bad!). Therefore, a compromise temperature is used to get an acceptable yield at a reasonable rate.
Changing the pressure is another way to control the position of equilibrium in certain reactions. Like temperature and concentration, the outcome is predicted by Le Chatelier’s Principle.
Le Chatelier’s Principle states: When a change is made to the conditions of a system at equilibrium, the system automatically moves to oppose the change.
Crucially, pressure changes only affect reactions that involve gases. If there are no gases, or if the number of gas molecules is the same on both sides of the equation, pressure will have no effect.
The rules are:
If you INCREASE the pressure: The system will try to reduce the pressure. It does this by favouring the side of the reaction with the fewer molecules (moles) of gas. The equilibrium will shift to the side with the least gas moles.
If you DECREASE the pressure: The system will try to increase the pressure. It does this by favouring the side of the reaction with the more molecules (moles) of gas. The equilibrium will shift to the side with the most gas moles.
Exam Tip: The first step is always to look at the balanced symbol equation and count the number of moles of gas on the reactant side and the product side. Ignore any solids (s), liquids (l), or aqueous solutions (aq).
A classic example is the reversible reaction between brown nitrogen dioxide gas and colourless dinitrogen tetroxide gas.
2NO₂(g) ⇌ N₂O₄(g) (dark brown) ⇌ (colourless)
Count the gas moles:
Left side (reactants): 2 moles of gas.
Right side (products): 1 mole of gas.
Apply the principle:
If we increase the pressure, the equilibrium will shift to the side with fewer moles to reduce the pressure. It will shift to the right. The mixture will become paler / more colourless.
If we decrease the pressure, the equilibrium will shift to the side with more moles to increase the pressure. It will shift to the left. The mixture will become darker brown.
The Haber process is the industrial method for making ammonia. High pressure is essential for a good yield.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Count the gas moles:
Left side (reactants): 1 + 3 = 4 moles of gas.
Right side (products): 2 moles of gas.
Apply the principle:
To get the highest possible yield of ammonia, we need to shift the equilibrium as far to the right as possible.
According to Le Chatelier’s principle, an increase in pressure will favour the side with the fewer gas molecules.
Therefore, high pressures are used (around 200 atm) to shift the equilibrium to the right and increase the yield of ammonia.
Exam Tip: Questions about the Haber process are very common. You should be able to explain the compromise between rate, yield, and cost for both temperature and pressure. For pressure, very high pressures give a great yield but are expensive and dangerous to maintain, so a compromise pressure is chosen.