Chapter 7 - Reversible Reactions and Chemical Equilibrium

Equilibrium: products and reactants are occurring simultaneously at the same rate. water in a closed jar has the number of evaporating molecules and number of condensing molecules equal
Chemical processes that reach equilibrium: * Homogeneous equilibrium: the equilibrium reactions in which the reactants and products are in the same phase * Heterogeneous equilibrium: the equilibrium of reactions in which the reactants and products are in different phases

Favorable change: change that has a natural tendency to happen under certain conditionsEnthalpy: Exothermic: when products have less enthalpy than reactants, it releases energyEndothermic: when reaction absorbs energy Temperature: changes with the direction
Entropy S: a tendency towards randomness or disorder in a system
The second law of thermodynamics: total entropy of the universe is constantly increasing must add together changes in the entropy of the system, ∆Ssys, and changes in the entropy of the surroundings, ∆Ssur
Free energy: available energy; useful work obtained from reaction * Also called gibbs free energy G
∆G = ∆H − T∆S * Change in free energy * Change in enthalpy * Change in entropy kelvin temperature
∆G * Negative: forward reaction is favorable * Zero: reaction at equilibrium * Positive: reaction favorable in reverse direction only

Law of chemical equilibrium: At equilibrium, there is a constant ratio between the concentrations of the products and reactants in any change
equilibrium constant keq or kc: forward rate constant divided by the reverse rate constant * kf / kr = keq * kc: uses concentrated values when concentration is at equilibrium
ICE table: table is used to record the initial, change, and equilibrium values of the reacting species,
K > 1, products are favoured. The equilibrium lies far to the right. Reactions where K is greater than 1010 are usually regarded as going to completion
K ≈ 1, there are approximately equal concentrations of reactants and products at equilibrium
K < 1, reactants are favoured. The equilibrium lies far to the left. Reactions in which K is smaller than 10−10 are usually regarded as not taking place at all

Reaction quotient, Qc: expression that is identical to equilibrium constant expression, but concentration not necessarily at equilibrium
Qc = ( [R]c[S]d ) / ( [P]a[Q]b ) * If Qc equal to kc then it is Qc close to equilibrium
Le Châtelier’s principle: a dynamic equilibrium tends to respond so as to relieve the effect of any change in the conditions that affect the equilibrium * Common ion effect: common ion involves adding ion to a solution in which theion is already present, the equilibrium shifts away from the added ion * applies Le Châtelier’s principle to ions in an aqueous solution * Endothermic change (∆H > 0): * An increase in temperature shifts the equilibrium to the right, forming more products. * Kc increases. * A decrease in temperature shifts the equilibrium to the left, forming more reactants. * Kc decreases. * Exothermic change (∆H < 0): * An increase in temperature shifts the equilibrium to the left, forming more reactants. * Kc decreases. * A decrease in temperature shifts the equilibrium to the right, forming more products. * Kc increases. * Reducing the volume of an equilibrium mixture of gas (constant temperature), causes sift a shift in equilibrium in direction with few gas molecules * Catalyst does not affect the position of equilibrium, only affects the time