Chapter 7.1: Recognizing 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

Chapter 7.2: Thermodynamics and Equilibrium

• Favorable change: change that has a natural tendency to happen under certain conditions

• Enthalpy: Exothermic: when products have less enthalpy than reactants, it releases energy

• Endothermic: 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

Chapter 7.3 The equilibrium constant

• 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

7.4 Predicting the Direction of a Reaction

• 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