Chemical Equilibria

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): 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.

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: 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 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: 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.

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.

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.