Reaction Rates and Equilibrium Ch. 19

Rates of Reaction 19-1

Collision Theory The rate or speed of a reaction varies tremendously depending on the reaction Rate = measures the speed of any change over time. Collision Theory = atoms/ions/molecules react to form products when they collide, if they have enough kinetic energy. –If they lack enough kinetic energy, will just bounce off each other when they collide

Activation Energy Activation Energy (E a ) = minimum amount of energy that particles must have in order to react.

Energy Diagrams – Exothermic Reaction Ex: 2H 2 + O 2  2H 2 O + 286kJ

Energy Diagrams – Endothermic Reaction Ex: 2H 2 O + 286kJ  2H 2 + O 2

DRAW! R P P R +ΔH -ΔH ENDOTHERMICEXOTHERMIC REACTION  ENERGY 

Factors Affecting Reaction Rates 1.Temperature Reaction Rate -Inc. temp., inc. kinetic E, inc. # of collisions, inc. rate 2.ConcentrationReaction Rate -Inc. # of particles, more collisions, inc. rate 3.Surface AreaReaction Rate -More surface area, inc. reactant exposure to collisions, inc. rate

4. Catalyst = substance that increases rate of reaction without being used up in rxn. –Catalyst’s lower the activation energy of a rxn. –Catalyst written above the “yield arrow” 2H 2 (g) + O 2 (g)  2H 2 O(l) 5. Inhibitor = substance that interferes with action of catalyst. Pt

Reaction without a catalyst Reaction with a catalyst Effect of a Catalyst

-H-H DRAW!!!! ( Or has an Inhibitor)

Reversible Reactions and Equilibrium 19-2

Reversible Reactions Reversible Reactions = reactions which occur simultaneously in both directions. 2SO 2 (g) + O 2 (g)  2SO 3 (g) Chemical Equilibrium = forward and reverse reactions take place at the same rate.

Factors affecting Equilibrium: Le Chatelier’s Principle Le Chatelier’s Principle = if a stress is applied to equilibrium, the system will change to relieve the stress. 1.Concentration: changing the amount of any reactant or product at equilibrium disturbs EQ. H 2 CO 3 (aq)  CO 2 (aq) + H 2 O(l) -If add CO 2 (product), rxn direction shifts LEFT -If remove CO 2 (product), rxn direction shifts RIGHT -If add H 2 CO 3 (reactant), rxn direction shifts RIGHT -If remove H 2 CO 3 (reactant), rxn direction shifts LEFT

2.Temperature = increasing temperature causes equilibrium to shift in direction that absorbs heat. Exothermic: 2SO 2 (g) + O 2 (g)  2SO 3 (g) + heat -If add heat (product), rxn shifts LEFT -If remove heat (product), rxn shifts RIGHT Endothermic: N 2 (g) + O 2 (g) + heat  2NO(g) -If add heat (reactant), rxn shifts RIGHT -If remove heat (reactant), rxn shifts LEFT

3.Pressure = change in pressure affects only equilibrium rxn’s with an unequal # of moles of gaseous reactants AND products. N 2 (g) + 3H 2 (g)  2NH 3 (g) 4 moles reactant to 2 moles product -If increase pressure, shift RIGHT (move to less moles side) -If decrease pressure, shift LEFT (move to more moles side)

Another pressure example: 2Ag 2 O(g)  4Ag(g) + O 2 (g) 2 moles reactant to 5 moles product -If increase pressure, shift LEFT (move to less moles side) -If decrease pressure, shift RIGHT (move to more moles side)

Equilibrium Constants Equilibrium Constant (K eq ) = ratio of product concentration to reactant concentrations at equilibrium. –aA + bB  cC + dD –K eq = [C] c x [D] d [A] a x [B] b –Put product concentrations over reactant concentrations. –Coefficients become exponents of concentrations. –Concentrations are in mol/L or M –K eq will change if temperature changes –If K eq < 1, then reactants are favored (more reactants than products) –If K eq > 1, then products are favored (more products than reactants)

Equilibrium example: –Using the following equation; 4.00 moles/L of Cl 2, 1.2 mol/L BrCl, and 4.00 moles/L of Br 2. What is the K eq ? 2BrCl(g)  Cl 2 (g) + Br 2 (g) K eq = [Cl 2 ] x [Br 2 ] [BrCl] 2 K eq = 4.00mol/L x 4.00mol/L (1.2mol/L) 2 K eq = 11.1 K eq > 1 Products are favored!!