Chemical Equilibria

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.

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