1. Unit 9: Chemical Equilibrium
Collision theory
Rates of reactions
Catalysts
Reversible reactions
Chemical equilibrium
Le Chatelier’s Principle
Concentration
Temperature
Volume

2. A. Collision Theory
Reaction rate depends on the collisions between reacting particles.
Successful collisions occur if the particles...
collide with each other
have the correct orientation
have enough kinetic energy to break bonds

3. Activation energy: minimum energy required for a reaction to occur
Exothermic
Endothermic
Activation energy
Time
Energy
Time
Energy
Energy of reaction

4. A. Collision Theory Activation Energy depends on reactants
low Ea = fast rxn rate Ea

5. 16.2: Rates of Reactions 1. SURFACE AREA
Chemical kinetics: the study of the rate (the speed) of a reaction
Rate of a chemical reaction depends on:
1. SURFACE AREA
2. CONCENTRATION of reactants
3. TEMPERATURE (T) of reactants
4. Presence/absence of a CATALYST

6. SURFACE AREA Surface Area high SA = fast rxn rate
more opportunities for collisions
Increase surface area by…
using smaller particles
dissolving in water

7. Effect of Concentration on Rate
KMT (Kinetic-Molecular theory) states that increasing concentration of reactants results in more collisions.
More collisions result in more reactions, increasing the rate of the reaction.

8. Effect of Temperature on Rate
Increasing T increases particle speed.
Faster reactants means more collisions have the activation energy, which increases the rate of the reaction.

9. Effect of Catalysts on Rate
A catalyst:
A chemical that influences a reaction, but is not consumed in the reaction. (It can be recovered unchanged at the end of the reaction.)
Lowers the activation energy of the reaction.
Activation energy
Time
Energy
Activation energy with catalyst

10. 16.1: Reversible Reactions
* Thus far, we have considered only one-way reactions: A + B → C + D
Some reactions are reversible:
They go forward (“to the right”) : A + B → C + D
and backwards (“to the left”) : A + B ← C + D
Written with a two-way arrow:
A + B ↔ C + D
Examples:
Boiling & condensing
Freezing & melting
Recharging a “rechargeable battery”

11. Examples of irreversible reactions:
Striking a match / burning paper
Dropping an egg
Cooking (destroys proteins)

12. 16.3: Chemical Equilibrium
For a reversible reaction, when the forward rate equals the backward rate, a chemical equilibrium has been established.
Both the forward and backward reactions continue, but there is a balance of products “un-reacting” and reactants reacting.
A + B ↔ C + D A B + A B + A B + C D + C D + C D +

13. * Le Chatelier’s Principle is about reducing stress – a stress applied to a chemical equilibrium
Relax! Reduce stress brought on by chemical equilibrium with me, Henri Le Chatelier!
(1850 – 1936)

14. 16.4: Le Chatelier’s Principle
When a stress is applied to a system (i.e. reactants and products) at equilibrium, the system responds to relieve the stress.
The system shifts in the direction of the reaction that is favored by the stress.
A stress is a change in:
Concentration
Temperature
Volume

15. 16.5: Stress: Change Concentration
Ex: Co(H2O) Cl1- ↔ CoCl H2O
(pink) (blue)
Stress Result
Add Cl1- Forward rxn favored
Shifts forward to reduce extra Cl1-
More CoCl42- will form
Add H2O Backward rxn favored
Shifts backward to reduce extra H2O
More Co(H2O)62+ will form

16. 16.7: Stress: Change Temperature
Ex: heat + Co(H2O) Cl1- ↔ CoCl H2O
(pink) (blue)
This reaction is endothermic. For Le Chatelier’s principle, consider “heat” as a chemical.
Stress Result
Increase T Forward rxn favored; shifts forward to reduce extra heat
More CoCl42- will form
Decrease T Backward rxn favored; shifts backward to replace “lost” heat
More Co(H2O)62+ will form

17. 16.6: Stress: Change Volume
Ex: 1 N2 (g) + 3 H2(g) ↔ 2 NH3(g)
(1 + 3 = 4 moles of gas) ↔ (2 moles of gas)
Stress Result
Decrease V Forward rxn favored; shifts forward to side with fewer moles of gas (reduces # of molecules packed into this smaller volume)
Increase V Backward rxn favored; shifts backward to side with more moles of gas (to fill the larger volume with more molecules)

18. 16.7: Catalysts & Equilibrium
MnO2
Ex: 2 H2O2 (aq) ↔ 2 H2O (l) + O2 (g)
Since a catalyst increases the forward and backward rates equally, it will not shift the equilibrium.

19. NaCl (s) ↔ Na+ (aq) + Cl- (aq)
Ex: saturated salt solution
NaCl (s) ↔ Na+ (aq) + Cl- (aq)
Dissolving (forward rate) decreases…
Reaction Rate
Time
Equilibrium is established:
Forward rate = Backward rate
Crystallization (backward rate) increases…