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

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

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

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

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

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

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.

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.

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

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”

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

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 +

* 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)

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

16.5: Stress: Change Concentration Ex: Co(H2O)62+ + 4 Cl1- ↔ CoCl42- + 6 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

16.7: Stress: Change Temperature Ex: heat + Co(H2O)62+ + 4 Cl1- ↔ CoCl42- + 6 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

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)

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.

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…