- Online Text
- 1. Introduction
- 2. Microstates, Macrostates, and Entropy
- 3. The Entropy of Energy Quanta
- 4. Entropy and States of Matter
- 5. Spontaneity and Gibbs Free Energy
- 6. Coupling Reactions
- 7. Equilibrium
- 8. The Equilibrium Constant Expression
- 9. Le Chatelier's Principle
- 10. Temperature and Equilibrium
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 9: Le Chatelier's Principle
When a chemical system reaches equilibrium, it is in its most stable state. If something disturbs the equilibrium—for example, the concentration of a chemical is changed—the system will return to equilibrium just as a ball pushed up a hill will roll back down. As it returns to equilibrium, it counteracts the original disturbance. For example, if the concentration of a reactant increases, the reaction will lower its concentration. If a product is removed, it will be restored. If heat is added, the reaction will cool back down, etc. In 1888, French chemist Henri Louis Le Chatelier (1850–1936) described this phenomenon that bears his name:
Every change of one of the factors of an equilibrium occasions a rearrangement of the system in such a direction that the factor in question experiences a change in a sense opposite to the original change.
When a reaction at equilibrium is disturbed, it will shift to the right or the left to return to equilibrium. Shifting to the right means some reactants are consumed and more products are made. Shifting to the left means some products are consumed and more reactants are made. There are a number of ways in which an equilibrium can be disturbed:
Adding or Removing a Product or Reactant
If more A is added, the reaction will shift to the right to restore the equilibrium ratio. Adding B will shift the reaction left. Removing A shifts the reaction left; removing B shifts it right.
Figure 9-12. Equilibrium Analogy
Two containers of water are connected by a tube. If water is added or removed from one side, water will flow from one side to the other to restore equilibrium.
© Science Media Group.
To remember how an equilibrium shifts when chemicals are added or removed, consider the analogy in Figure 9-12. Two containers of water represent the reactants and products of a chemical reaction. A tube allows water to flow from one to the other. The system starts out at equilibrium, and the water levels are even. If extra water is added to the container on the left, the water flows from left to right until the levels equalize and equilibrium is reached again.
Likewise, if water is removed from the reactant (left) side, the flow from right to left will restore the equilibrium.
In chemical syntheses, the goal is to maximize the amount of product. In this case, equilibrium is not a desirable state; once equilibrium is reached, the product ceases to accumulate. The Haber-Bosch process for synthesizing ammonia from atmosphere nitrogen (discussed in Unit 7 and illustrated in the Control the Haber-Bosch Ammonia Plant interactive) faces this problem. The reaction reaches equilibrium:
N2(g) + 3H2(g) 2NH3(g)
To prevent the reaction from stopping, the ammonia is continually removed from the reaction apparatus. Keeping the amount of NH3 low forces the reaction to shift to the right until the reactants are used up.
Change in Pressure
Pressure only affects reactions involving gases; it does not affect reactions where no gases are present. When pressure on the system increases, the reaction will shift to counteract the increase in pressure; it does this by shifting toward the side of the reaction with fewer moles of gas. For example:
A(g) B(g) + 2C(g)
In this reaction, there is one mole of gas on the left and three moles of gas on the right. If the pressure increases, the reaction will try to decrease the pressure by decreasing the amount of gas—it will shift to the left. Conversely, if the pressure decreases, the reaction will try to increase the pressure by increasing the amount of gas—it will shift to the right.
The shifting due to pressure is also used to maximize yield in the Haber-Bosch process. There are four moles of gas on the reactant side of the equation and two moles on the product side. To shift the reaction toward the product side, the reaction mixture is kept under high pressure.
In Unit 12, we will explore how catalysts speed up reactions. At equilibrium, both the forward and reverse reactions are happening at the same rate. Adding a catalyst speeds up both reactions, and therefore the equilibrium balance is maintained, but a reaction can reach equilibrium faster. Catalysts do not cause shifts in equilibrium.