1. Unit 13: Equilibrium 13.1 Equilibrium 13.2 Le Chatelier’s Principle
13.3 Equilibrium Expressions

2. 13.1 Equilibrium Learning Targets
Explain why chemical reactions are reversible.
Define chemical equilibrium.

3. 13.1 Equilibrium Irreversible
Chemical reaction that only runs in one direction
“Goes to completion”
Products can’t collide to react in reverse
Mg + 2 HCl  MgCl2 + H2

4. 13.1 Equilibrium Reversible
Chemical reaction that runs in both directions
Forward: Reactants to products
Reverse: Products to reactants
2 NO2 (g) N2O4 (g)
NO2 is a brown gas
N2O4 is a colorless gas

5. 13.1 Equilibrium
Some NO2 must be remade even as it is used up

6. 13.1 Equilibrium Why reversible?
Reactants need activation energy to go to products.
If products get reverse activation energy, they will go back to reactants.
If there is sufficient energy, the reaction will go in both directions.

7. 13.1 Equilibrium Equilibrium
Chemical reaction that is balanced and the forward and reverse reactions occur at same rate
1. Dynamic: forward and reverse reactions do not stop
2. Concentrations don’t change over time
Concentrations of reactants and products are not equal
3. Closed system: Nothing gets in or out
4. Temperature is held constant

9. 13.2 Le Chatelier’s Principle
Learning Targets
Define Le Chatelier’s principle.
Predict how equilibrium will change if pressure, concentration, or temperature are changed.
State the effect of a catalyst or inhibitor on equilibrium.

10. 13.2 Le Chatelier’s Principle
Lake Nyos
CO2 built up at the bottom of the lake
Gas was in equilibrium with water
Avalanche disrupted the equilibrium, caused the gas to come out of the water
Events can change equilibrium

11. 13.2 Le Chatelier’s Principle
Henri le Chatelier
French, 1850 to 1936
Professor and scientist

12. 13.2 Le Chatelier’s Principle
If a stress is added to a chemical reaction in equilibrium, it will change to relieve that stress and form a new equilibrium
New equilibrium: new amounts of products and reactants
Ways to disturb equilibrium (stresses):
Concentration
Pressure
Temperature

13. 13.2 Le Chatelier’s Principle
Concentration
Adding more of a reactant or product (increase concentration)
2 H2 (g) + O2 (g) H2O (g) + energy
Add H2: 2 H2 (g) + O2 (g) H2O (g) + energy
Add O2: 2 H2 (g) + O2 (g) H2O (g) + energy
Add H2O: 2 H2 (g) + O2 (g) H2O (g) + energy

14. 13.2 Le Chatelier’s Principle
Concentration
Remove a reactant or product (decrease concentration)
2 H2 (g) + O2 (g) H2O (g) + energy
Remove H2: 2 H2 (g) + O2 (g) H2O (g) + energy
Remove O2: 2 H2 (g) + O2 (g) H2O (g) + energy
Remove H2O: 2 H2 (g) + O2 (g) H2O (g) + energy

15. 13.2 Le Chatelier’s Principle
You try!
Co(H2O)6 +2(aq) + 4 Cl- (aq) CoCl4-2 (aq) + 6 H2O (l)
Add Cl-1: Co(H2O)6 +2(aq) + 4 Cl- (aq) CoCl4-2 (aq) + 6 H2O (l)
Add H2O: Co(H2O)6 +2(aq) + 4 Cl- (aq) CoCl4-2 (aq) + 6 H2O (l)
Remove Cl-1- : Co(H2O)6 +2(aq) + 4 Cl- (aq) CoCl4-2 (aq) + 6 H2O (l)

16. 13.2 Le Chatelier’s Principle
Pressure
Chemical reaction responds by relieving changes in pressure
*Only changes if gases are involved*
Increase pressure
Increases amounts of substances on side with fewest moles of gas
Add coefficients of gases
Decrease pressure
Increases amounts of substances on side with most moles of gas

17. 13.2 Le Chatelier’s Principle
Pressure
Increase pressure
2 H2 (g) + O2 (g) H2O (g)
3 gas molecules on reactants side (2+1)
2 gas molecules on products side (2)
2 H2 (g) + O2 (g) H2O (g)

18. 13.2 Le Chatelier’s Principle
Pressure
Decrease pressure
2 H2 (g) + O2 (g) H2O (g)
2 H2 (g) + O2 (g) H2O (g)

19. 13.2 Le Chatelier’s Principle
Pressure
H2 (g) + Cl2 (g) HCl(g)
2 gas molecules on reactants side (2)
2 gas molecules on products side (2)
Equal moles on both sides
Changing pressure does not change equilibrium

20. 13.2 Le Chatelier’s Principle
You try!
HCl (aq) + NaHCO3 (aq) H2O (l) + CO2 (g) + NaCl (aq)
Increase pressure
HCl (aq) + NaHCO3 (aq) H2O (l) + CO2 (g) + NaCl (aq)
Decrease pressure

21. 13.2 Le Chatelier’s Principle
Temperature
Heat appears in reaction equation
Adding heat favors endothermic (ΔH is positive)
Removing heat favors exothermic (ΔH is negative)

22. 13.2 Le Chatelier’s Principle
Temperature
2 H2 (g) + O2 (g) H2O (g) + heat
Heat can be a “reactant” or “product”
Exothermic: product
Endothermic: reactant
“Endo” as a reactant, “exo” as a product

23. 13.2 Le Chatelier’s Principle
Temperature
2 H2 (g) + O2 (g) H2O (g) + energy
Remove heat: 2 H2 (g) + O2 (g) H2O (g) + energy
Add heat: 2 H2 (g) + O2 (g) H2O (g) + energy

24. 13.2 Le Chatelier’s Principle
You try!
Ba(OH)2 + NH4SCN + energy Ba(SCN)2 + NH3 + H2O
Remove heat:
Ba(OH)2 + NH4SCN + energy Ba(SCN)2 + NH3 + H2O
Add heat:

25. 13.2 Le Chatelier’s Principle
Catalyst and Inhibitor
Speed up or slow down BOTH forward and reverse reactions
Do NOT change equilibrium

27. 13.3 Equilibrium Expressions
Learning Targets
Use a balanced chemical equation to write the equilibrium law for a reaction.

28. 13.3 Equilibrium Expressions
Equation that compares the forward reaction rate with the reverse reaction rate
Rate Laws:
Rate laws are used to relate reaction rate to concentration of reactants

29. 13.3 Equilibrium Expressions
To write an equilibrium expression:
Products go over reactants
Keq: equilibrium constant

30. 13.3 Equilibrium Expressions
To write an equilibrium expression:
Only count gases and aqueous reactants/products
Solids and liquids can’t easily react in a reversible way

31. 13.3 Equilibrium Expressions
Meaning of Keq
If Keq > 1, products are favored
If Keq < 1, reactants are favored
If Keq = 1, reactants and products are equal

32. 13.3 Equilibrium Expressions
Example 1:
__ CO (g) + __ H2O (g) __ CO2 (g) + __ H2 (g)

33. 13.3 Equilibrium Expressions
Example 2:
__ NO2 (g) __ N2O4 (g)

34. 13.3 Equilibrium Expressions
Example 3:
__ C (s) + __ H2O (g) __ CO2 (g) + __ H2 (g)