Position of equilibrium
Aims of this page
After studying this page, you should be able to:
- recall that imposing a change on a chemical reaction can change the position of equilibrium
- describe and explain how a change in conditions affect the position of equilibrium.
What is position of equilibrium?
| Position of equilbrium | Equilibrium mixture |
|---|---|
| Lies to the left | Reactant concentrations higher |
| Lies to the right | Product concentrations higher |
You cannot tell the position of equilibrium by looking at the balanced symbol equation for the reaction. You need to know the concentrations of reactants and products at equilibrium.
The position of equilibrium can be changed by changing the reaction conditions.
Le Chatelier’s Principle provides a way to predict what happens:
- A change in a condition affecting a dynamic equilibrium causes the position of equilibrium to move to counteract that change.
This lets you know what will happen, but it doesn’t fully explain why it will happen. However, the detailed explanations are more complicated than you need for your GCSE (and probably for your A Level too).
Changing a concentration
The position of equilibrium will move in the direction away from an added substance (and towards any substance that is removed).
The reaction between iron(III) ions and thiocyanate ions is easy to do in the school lab:
| Fe3+(aq) | + | SCN–(aq) | ⇌ | FeSCN2+(aq) |
| iron(III) ion | thiocyanate ion | thiocyanatoiron(III) ion | ||
| pale yellow | colourless | blood red |
| Action | Change | Result |
|---|---|---|
| Add more Fe3+ | Fe3+ concentration ⇧ | Turns red |
| Add more SCN– | SCN– concentration ⇧ | Turns red |
| Add NH4Cl powder | Fe3+ concentration ⇩ | Turns yellow |
Explanation
In terms of Le Chatelier’s principle:
- adding Fe3+ or SCN– causes the position of equilibrium to move to the right to use up these ions
- removing Fe3+ causes the position of equilibrium to move to the left to produce more of these ions.
Changing the temperature
The position of equilibrium will move in the direction of:
- the exothermic reaction if the temperature is decreased
- the endothermic reaction if the temperature is increased.
It may help if you only learn one of these ideas (exothermic or endothermic) because the other is just the other way round.
This reaction is easy to do in the school lab:
\(\ce{$\underset{\large\textsf{blue}}{\ce{Cu^2+(aq)}}$ + $\underset{\large\textsf{colourless}}{\ce{4Cl-(aq)}}$ <=>[\large\textsf{endothermic}][\large\textsf{exothermic}] $\underset{\large\textsf{green}}{\ce{[CuCl4]^2-(aq)}}$}\)
The forward reaction is endothermic (and so the reverse reaction is exothermic). You can change the temperature of the reaction mixture by heating the test tube in boiling water or by cooling it in iced water.
| Action | Change | Result |
|---|---|---|
| Put in boiling water | Temperature ⇧ | Turns green |
| Put in iced water | Temperature ⇩ | Turns blue |
Explanation
In terms of Le Chatelier’s principle:
- increasing the temperature causes the position of equilibrium to move to the right (in the direction of the endothermic reaction)
- this causes more [CuCl4]2– to form, so the mixture changes from blue to green.
Changing the pressure
If the pressure is increased, the position of equilibrium moves in the direction of the fewest molecules of reacting gas.
You need to look at the balanced chemical equation to work out which side has the fewest molecules of gas. For example, look at this reaction:
| 2NO2(g) | ⇌ | N2O4(g) |
| nitrogen dioxide | dinitrogen tetroxide | |
| red-brown | colourless |
There are 2 molecules of gas on the left-hand side of the equation, and only 1 molecule of gas on the right-hand side. This means that:
- increasing the pressure causes the position of equilibrium to move to the right (in the direction of the fewest molecules of gas)
- this causes the reaction mixture to become paler.
This video shows what happens.
The syringe contains nitrogen dioxide. It is sealed at the end so that the pressure increases when the plunger is pushed down. The bottle on the right contains nitrogen dioxide at atmospheric pressure.
Explanation
There are four main events in the video:
- The plunger is rapidly pushed in, which reduces the volume and increases the pressure
- The colour darkens because the concentration of NO2 increases
- The colour lightens slightly as the equilibrium shifts and some NO2 is converted to N2O4
- The plunger is released, the colour lightens noticeably, and then darkens slightly as the equilibrium shifts again.
Adding a catalyst
A catalyst is a substance that increases the rate of a reaction, without being used up or changed chemically by the end of the reaction.
A catalyst reduces the time taken to reach equilibrium, but it does not alter the position of equilibrium. Think of someone on a treadmill. They adjust their running speed to match the speed of the belt to stay in the same position.
Explanation
A catalyst increases the rate of the forward reaction and reverse reaction by the same ratio.
Summary
| Change | Position of equilibrium |
|---|---|
| Add more reactant | Moves to the right |
| Add more product | Moves to the left |
| Increase the temperature | Moves in the direction of the endothermic reaction |
| Increase the pressure | Moves in the direction of the fewest molecules of reacting gas |
| Add a catalyst | No change |
