Welcome to Chemical Equilibrium!
Have you ever tried walking up a "down" escalator at the exact same speed it is moving down? You are walking, the escalator is moving, but to someone watching from the side, you stay in the exact same spot! In chemistry, this special type of balance is known as dynamic equilibrium.
Don't worry if this chapter seems tricky at first. We will break down every single idea into simple, bite-sized steps with everyday analogies to help you master your CCEA GCSE Double Award Science exam.
1. Reversible Reactions
In most chemical reactions you have studied so far, reactants turn into products and cannot easily turn back (like baking a cake—you cannot unbake it!). These are called irreversible reactions.
However, some reactions can go both forwards and backwards. These are called reversible reactions.
• Forward reaction: Reactants turn into products (left to right).
• Reverse (or backward) reaction: Products break down or react to reform the original reactants (right to left).
In chemical equations, we show a reversible reaction using a special double arrow: \(\rightleftharpoons\)
General Equation:
\(A + B \rightleftharpoons C + D\)
Example: The synthesis of ammonia in the Haber Process:
\(N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\)
Open Systems vs Closed Systems
For a reversible reaction to reach a balance, it must take place in a closed system.
• Closed system: A container where no substances can enter or leave (like a sealed bottle of soda).
• Open system: A container where gases or liquids can escape into the surroundings (like an open beaker of boiling water).
Key Takeaway: Reversible reactions use the \(\rightleftharpoons\) symbol and can only reach equilibrium in a closed system.
2. What is Dynamic Equilibrium?
When you place reactants into a sealed container, here is what happens step-by-step:
1. At the start, only the forward reaction occurs because there are no products yet.
2. As products are formed, the reverse reaction begins and gradually gets faster.
3. Eventually, the rate of the forward reaction becomes equal to the rate of the reverse reaction.
When this happens, the reaction has reached dynamic equilibrium.
The Two Golden Rules of Dynamic Equilibrium
You must memorize these two key features for your exam:
1. The rate of the forward reaction equals the rate of the reverse reaction.
2. The amounts (concentrations) of reactants and products remain constant.
Why is it called "dynamic"?
The word dynamic means active or moving. Even though the amounts of chemicals stop changing, the reactions have not stopped! Reactants are still turning into products, and products are still turning into reactants at the exact same speed.
Analogy Time: Imagine you are filling a bucket with water from a tap at a rate of 1 litre per minute, while water is leaking out of a hole in the bottom at the exact same rate of 1 litre per minute. The water level in the bucket never changes, even though water is constantly flowing in and out.
Key Takeaway: At dynamic equilibrium, both reactions occur at equal speeds, so concentrations remain constant.
3. Changing Conditions: Le Chatelier's Principle
Once a system reaches equilibrium, it likes to stay balanced. But what happens if we change the temperature, pressure, or concentration?
A scientist named Henri Le Chatelier discovered a simple rule to predict what will happen:
Le Chatelier's Principle: If a change in conditions is applied to a system in dynamic equilibrium, the position of equilibrium will shift to oppose (counteract) that change.
Memory Trick: Think of equilibrium like a stubborn teenager. Whatever you do to it, it will try to do the exact opposite!
• If you heat it up \(\rightarrow\) It tries to cool things down.
• If you increase the pressure \(\rightarrow\) It tries to decrease the pressure.
• If you add more chemical \(\rightarrow\) It tries to use that chemical up.
What does "Position of Equilibrium" mean?
• Shift to the right: More products are made (yield of products increases).
• Shift to the left: More reactants are made (yield of products decreases).
4. Factors Affecting Equilibrium
A. Changing Temperature
Every reversible reaction is exothermic (gives out heat, \(\Delta H\) is negative) in one direction and endothermic (takes in heat, \(\Delta H\) is positive) in the opposite direction by the exact same amount.
Example:
\(N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\) (Forward reaction is exothermic: \(\Delta H = -92\text{ kJ/mol}\))
• Increasing the temperature: The equilibrium shifts in the endothermic direction (to absorb the extra heat).
• Decreasing the temperature: The equilibrium shifts in the exothermic direction (to release more heat and warm it back up).
B. Changing Pressure (Gases Only)
Pressure is caused by gas particles colliding with the walls of the container. More gas molecules mean higher pressure.
To predict the effect of pressure, count the total moles (molecules) of gas on each side of the balanced equation:
Example:
\(N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\)
• Left side: \(1 + 3 = 4\text{ moles of gas}\)
• Right side: \(2\text{ moles of gas}\)
• Increasing the pressure: The equilibrium shifts to the side with fewer moles of gas (to reduce the pressure). In this example, it shifts to the right.
• Decreasing the pressure: The equilibrium shifts to the side with more moles of gas (to increase the pressure). In this example, it shifts to the left.
• If both sides have equal moles of gas: Changing the pressure has no effect on the position of equilibrium.
C. Changing Concentration
• Adding more reactant: The equilibrium shifts to the right to use up the added reactant and make more product.
• Removing a product: The equilibrium shifts to the right to replace the product that was removed.
• Adding more product: The equilibrium shifts to the left to remove the extra product.
D. Adding a Catalyst
This is a favorite CCEA exam question!
A catalyst increases the rate of the forward reaction and the reverse reaction equally.
Because both rates increase by the exact same amount:
• A catalyst does NOT change the position of equilibrium.
• A catalyst does NOT increase the yield of products.
• A catalyst DOES allow the system to reach equilibrium much faster.
Key Takeaway: Temperature, pressure, and concentration can change the yield by shifting equilibrium. A catalyst only speeds up the time taken to reach balance.
5. Quick Summary & Common Exam Mistakes
Summary Table of Equilibrium Shifts
• Increase Temperature: Shifts in endothermic direction (\(+\Delta H\))
• Decrease Temperature: Shifts in exothermic direction (\(-\Delta H\))
• Increase Pressure: Shifts to the side with fewer gas molecules
• Decrease Pressure: Shifts to the side with more gas molecules
• Add Reactant / Remove Product: Shifts to the right (makes more product)
• Add Catalyst: No shift in position (reaches equilibrium faster)
Common Mistakes to Avoid
Mistake 1: Saying "concentrations are equal at equilibrium."
Correction: Concentrations are constant (they stay the same), but the amount of product does not have to equal the amount of reactant.
Mistake 2: Thinking a catalyst increases the yield.
Correction: Catalysts only save time; they never increase the final amount of product made.
Mistake 3: Counting solid or liquid moles when looking at pressure.
Correction: Pressure only affects gases! Check state symbols carefully \((g)\).
Exam Practice Tip
When answering an exam question about Le Chatelier's principle, always structure your answer in three steps:
1. State the change made (e.g., "Temperature is increased").
2. State how equilibrium opposes the change (e.g., "Equilibrium moves in the endothermic direction to absorb heat").
3. State the direction of the shift and the effect on yield (e.g., "Therefore, it shifts to the left, decreasing the yield of ammonia").