1. Equilibrium & LeChâtelier’s Principle
Chemistry Unit 12: Reaction Rates & Equilibrium
Lecture #2

2. Objectives Describe what chemical equilibrium IS, and what it ISN’T
Explain LeChâtelier’s Principle
Describe the effects different stresses have on a system in equilibrium

3. A state of balance due to the equal action of opposing processes
Equilibrium
A state of balance due to the equal action of opposing processes
Static Equilibrium:
A state of balance that is NOT moving
Dynamic Equilibrium:
A state of balance that IS in motion

4. Chemical Equilibrium
A state of dynamic equilibrium when the rate of the forward reaction is equal to the rate of the reverse reaction
Reactants are becoming products just as quickly as products are turning back into reactants

5. What Chemical Equilibrium IS…and what it IS NOT
When the rate of the forward reaction equals the rate of the reverse reaction
What it IS NOT…
The reaction does NOT stop. The forward and reverse are just happening at the same rate, so concentrations stop changing
Not (necessarily) when the amount of reactants equals the amount of products

6. Elevator Analogy Roof-Top Restaurant Lobby (Capacity 100 people)
(Capacity Unlimited)

7. Equilibrium Particulars
Only happens for reversible reactions
It is not reached immediately
The quantity of reactants and products does NOT have to be the same when equilibrium is reached
The rate of exchange (reactants to products and products to reactants) is the same once equilibrium is reached.

8. Le Châtelier’s Principle
When a stress is applied to a system in equilibrium, the equilibrium will shift to relieve the stress.
This means…either the forward or reverse reaction, is favored (will speed up) until a new equilibrium point is reached

9. Possible Stressors
Concentration
Temperature
Pressure

10. H2CO3(aq)  CO2(g) + H2O(l)
Concentration
Increasing the concentration of a substance causes more collisions on that side of the equation, and the equilibrium will shift to the other side
H2CO3(aq)  CO2(g) + H2O(l)
Increasing the concentration of the reactants…
Equilibrium shifts toward the products
The products are favored

11. Temperature (put heat into the equation first)
Adding heat causes more collisions on that side of the equation, so equilibrium shifts to the other side
2SO2(g) + O2(g)  2SO3(g) + heat
If the temperature is increased…
Equilibrium shifts toward the reactants
If the temperature is decreased…
Equilibrium shifts toward the products

12. Pressure (determine the number of moles on each side first)
Only affects gases
Increasing the pressure shifts the equilibrium to the side with fewer moles of gas
If the # of moles are equal on both sides…no shift in equilibrium.

13. H2CO3(aq)  CO2(g) + H2O(l)
Pressure
H2CO3(aq)  CO2(g) + H2O(l)
0 mol of (g) on left…1 mol of (g) on right
Increasing the pressure…
Decreasing the pressure (like leaving the cap off the bottle of soda)…

14. 2SO2(g) + O2(g)  2SO3(g) + heat
Pressure
2SO2(g) + O2(g)  2SO3(g) + heat
mol (g) on left… mol (g) on right
Increase in pressure shifts equilibrium to the
Decrease in pressure shifts equilibrium to the

15. PCl5(g) + heat  PCl3(g) + Cl2(g)
Practice Problem
Describe what effect each of the following will have on the equilibrium:
PCl5(g) + heat  PCl3(g) + Cl2(g)
Addition of Cl2
An increase in pressure
Removal of heat
Removal of PCl3(g) as it forms

16. Write Summary & Questions