When a reaction reaches equilibrium, it uses up reactants to make products at the same rate that it uses up products to re-make reactants. In order for these rates to remain equal, the reaction must take place in a closed system--that is, a system where no material or energy is being added or removed.
If the system is open to outside stresses, then the equilibrium can be disrupted. A stress is an outside influence that changes something about the reaction. Le Chatelier’s Principle states that if a stress is added to a system (i.e., a reaction) at equilibrium, then the system will shift the equilibrium to relieve the added stress.
Le Chatelier’s Principle in Simpler Terms
If something is added, the reaction will shift away to consume the excess that was added. If something is taken away, then the equilibrium will shift towards the void to replace what was taken.
Although these stresses can cause a shift in the reaction’s equilibrium, the value of the equilibrium constant (K) remains constant except in the instance of a temperature change.
Read each tab to learn more about the different ways that a reaction can be stressed and to see how Le Chatelier’s principle predicts how these stresses will affect the equilibrium of the reaction.
A reaction can be stressed by adding more of one of the substances (i.e., increasing its concentration). The equilibrium will shift away from the addition so that it can use up the excess chemical and get back to the equilibrium ratio.
Conversely, a reaction can also be stressed by removing some of a substance (i.e., decreasing its concentration). In this case, the equilibrium will shift towards the removal so that it can make up the deficit of the missing chemical and get back to the equilibrium ratio.
Click each stress in this chart to see how adding or removing different substances would affect the equilibrium for this reaction:
N2 (g) + 3 H2 (g) ⇌ 2 NH3 (g)
| Stress | Equilibrium Shift | Reasoning |
|---|---|---|
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→ (towards products) |
More reactants were added, so the reaction will favor the forward direction, using up the excess reactants and making more products until it can get back to the equilibrium ratio. |
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→ (towards products) |
More reactants were added, so the reaction will favor the forward direction, using up the excess reactants and making more products until it can get back to the equilibrium ratio. |
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← (towards reactants) |
More products were added, so the reaction will favor the reverse direction, using up the excess products and making more reactants until it can get back to the equilibrium ratio.
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← (towards reactants) |
Reactants were removed, so the reaction will favor the reverse direction, making more reactants to make up for the deficit until it can get back to the equilibrium ratio.
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← (towards reactants) |
Reactants were removed, so the reaction will favor the reverse direction, making more reactants to make up for the deficit until it can get back to the equilibrium ratio. |
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→ (towards products) |
Products were removed, so the reaction will favor the forward direction, making more products to make up for the deficit until it can get back to the equilibrium ratio. |
Recall that reactions can be described as endothermic or exothermic. An endothermic reaction is one that absorbs heat energy, while an exothermic reaction is one that releases heat energy. A reversible reaction is endothermic in one direction and exothermic in the opposite direction.
Raising the temperature requires adding heat, so the reaction will shift in the endothermic direction so that it can use up the excess heat and get back to the equilibrium ratio. Conversely, lowering the temperature requires removing heat, so the reaction will shift in the exothermic direction to produce more heat to make up for the deficit and get back to the equilibrium ratio.
Because an exothermic reaction produces heat, it can be helpful to think of the heat energy as a product of that reaction. Conversely, because an endothermic reaction absorbs heat, think of the heat energy as a reactant necessary for that reaction to proceed.
Click each stress in this chart to see how increasing or decreasing the temperature affects the equilibrium for this exothermic reaction:
N2 (g) + 3 H2 (g) ⇌ 2 NH3 (g) + 22 kcal
| Stress | Equilibrium Shift | Reasoning |
|---|---|---|
|
← (towards reactants) |
Heat is a “product” so when more heat energy is added, the reaction will favor the reverse (endothermic) direction, using up the excess heat until it can get back to the equilibrium ratio. |
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→ (towards products) |
Heat is a “product” so when heat energy is removed, the reaction will favor the forward (exothermic) direction, producing more heat until it can get back to the equilibrium ratio. |
Click each stress in this chart to see how increasing or decreasing the temperature affects the equilibrium for this endothermic reaction:
12.6 kcal + H2 (g) + I2 (g) ⇌ 2HI (g)
| Stress | Equilibrium Shift | Reasoning |
|---|---|---|
|
→ (towards products) |
Heat is a “reactant” so when more heat energy is added, the reaction will favor the forward (endothermic) direction, using up the excess heat until it can get back to the equilibrium ratio. |
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← (towards reactants) |
Heat is a “reactant” so when heat energy is removed, the reaction will favor the reverse (exothermic) direction, producing more heat until it can get back to the equilibrium ratio. |
Recall that pressure and volume have an inverse relationship: when the volume of a gas decreases (as when squeezing a balloon), the pressure of the gas increases. The opposite relationship is also true (when the volume increases, the pressure will decrease).
These pressure changes can cause an equilibrium shift for reactions involving gases. If the pressure is increased, then the equilibrium will shift in the direction that decreases the overall pressure. That is, it will shift in the direction that will produce fewer gas molecules. Conversely, if the pressure of the system is decreased, then the equilibrium will shift in the direction that increases the overall pressure, which is the direction that will produce more gas molecules.
It is important to note that only the number of gas particles matters when determining equilibrium shifts due to pressure changes. Aqueous, solid, or liquid substances will not be affected by pressure.
Click each stress in this chart to see how increasing or decreasing the volume/pressure affects the equilibrium for this reaction involving gases:
N2 (g) + 3 H2 (g) ⇌ 2 NH3 (g)
Notice there are 4 reactant gas molecules (1 N2 + 3 H2), while there are only 2 product gas molecules (2 NH3). This will be important for determining how pressure changes affect equilibrium for this reaction.
| Stress | Equilibrium Shift | Reasoning |
|---|---|---|
|
→ (towards products) |
There are fewer gas molecules on the product side, so when the pressure is increased, the reaction will favor the forward direction to produce fewer gas particles in order to lower the overall pressure and get back to the equilibrium ratio. |
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← (towards reactants) |
There are more gas molecules on the reactant side, so when the pressure is decreased, the reaction will favor the reverse direction to produce more gas particles in order to make up for the pressure deficit and get back to the equilibrium ratio. |
Recall that adding a catalyst will increase the rate (speed) of the reaction. For a reaction at equilibrium, this means that both the forward and reverse directions will increase equally. Therefore, the reaction will remain at equilibrium and no shift will occur.
