The Concentration Reaction Quotient

• All concentrations must be expressed in Molarity.

• Product concentrations are in the numerator (multiplied together).

• Reactant concentrations are in the denominator (multiplied together).

• Each concentration is raised to the power of its coefficient from the balanced equation.

• Solids are always omitted from the expressions for Q and K

• The numeric value of Qc for a given reaction can vary prior to equilibrium.

• The value of Qc depends on the concentration of products and reactants present at that particular moment.

• Qc can be calculated at any point in a reaction.

• We will often calculate Qc at the start of the reaction using initial concentrations.

Equilibrium Constant, K

• Equilibrium constant (K): the value of Q when the reaction is at equilibrium

• Don’t confuse this with the kinetic rate constant (k)

• If K is very small, the mixture contains mostly reactants at equilibrium.

• If K is very large, the mixture contains mostly products at equilibrium.

• The value of K gives no indication as to whether the reaction is fast or slow.

• The value of the equilibrium constant is independent of the starting amounts of the reactants and products.

The Direction of the Reaction

• A system that is not at equilibrium will proceed in the direction that establishes equilibrium.

• By comparing Q to K, it is possible to determine which direction the system will proceed to achieve equilibrium.

• When Q < K*:* reaction must shift FORWARD

• When Q > K*:* reaction must shift BACKWARD

• When Q = K*:* reaction is at equilibrium, and will maintain constant concentration

Homogenous Equilibrium

• Homogenous equilibrium: one in which all of the reactants and products are present in the same phase.

• Most commonly are either liquid or gaseous phases.

• Reaction quotients include concentration or pressure terms only for gaseous and solute species.

• For gas-phase solutions, the equilibrium constant may be expressed in terms of either the molar concentrations (Kc) or partial pressures (Kp) of the reactants and products.

Heterogenous Equilibrium

• Heterogenous equilibria: contain reactants and products that are in two or more different phases.

• Pure solids and pure liquids do not appear in the K expression.

• The position of equilibrium is independent of the amount of solid or liquid present, as long as at least some is present in the reaction mixture.

Le  Châtelier’s Principle

• Le Châtelier’s Principle: when a chemical system at equilibrium is disturbed, it returns to equilibrium by counteracting the disturbance.

  • At equilibrium Q = K.

  • The disturbance causes a change in Q.

  • The reaction will shift to re-establish Q = K.

• In the case of a temperature change, the disturbance changes the value of K.

  • The direction of that change depends on whether the reaction is exothermic or endothermic.

• If a chemical equilibrium is disturbed by adding a reactant or product, the system will proceed in the direction that consumes part of the added species.

• If a chemical equilibrium is disturbed by removing a reactant or product, the system will proceed in the direction that restores part of the removed species.

• The system responds in the way that restores equilibrium and therefore allows Q = K again.

• If what is added or removed is a SOLID or liquid, the reaction does not shift at all

  • However, while the amount of solid does not affect the equilibrium, any shift in equilibrium DOES change the amount of solid.

  • This is because pure liquids and solids do not appear in the equilibrium expression.

Effect of Temperature

• An increase in temperature will change K.  It will increase K for an endothermic, and decrease K for an exothermic

• It will shift the reaction so as to favor whichever direction is endothermic.

• A decrease in temperature will also change K.  It will Increase K for an exothermic,  and decrease K for an endo.

• It will shift the reaction so as to favor the exothermic direction.

Catalysts

• A catalyst speeds up the rate of a reaction.

• For reversible reactions, catalysts increase the rates of the forward and reverse reactions.

• Result: A catalyst causes the system to reach equilibrium more quickly.

• But a catalyst does not affect the equilibrium concentrations or value of the equilibrium constant.