1. To write equilibrium expressions for both homogeneous and heterogeneous equilibria.

2. To convert K_c to K_p and vice versa and to obtain the appropriate form of the equilibrium constant for a give situation.

3. To determine the reaction quotient and use it to determine whether or not equilibrium conditions exist, and which direction a reaction will shift to attain equilibrium.

4. To calculate the equilibrium constant from experimental data on the percent of reaction, the concentration of the reactant that actually undergoes the process, or the equilibrium concentrations of each species in the system.

5. To use Le Chatelier’s principle to predict how a given equilibrium system will be affected by temperature changes, volume changes, changes in concentrations of the system’s components, or the addition of a catalyst.

1. When given the equilibrium constant for one balanced equation, know how to obtain the new equilibrium constant if the original equation is reversed, multiplied by a constant, or combined with other balanced equations.

2. Remember that when writing equilibrium expressions, the quantities are multiplied, not added as they are in the balanced equation.

3. Using LeChatelier’s principle to predict the effects of changing system conditions on a chemical equilibrium seems easy at first, but it can be tricky. Remember that pure solids, pure liquids, and solvents in dilute solutions don’t appear in the equilibrium expression. Therefore, adding more of these substances to ( or removing them from) an equilibrium system will have no effect on the equilibrium. Also notice that the shift in the equilibrium doesn’t completely eliminate the original stress. For example, if I₂ is added to the system 2HI_(g) Ö H_2(g) + I_2(g) the equilibrium does shift to the left, and the concentrations of hydrogen and iodine gases are less than they were after the extra iodine was added. It is important to realize that the final concentration of iodine gas is greater that it was before you added the extra iodine.

4. Remember that an equilibrium expression is not a true equality unless all of the concentration values used are equilibrium concentrations. As the problems become more complicated, keep in mind that the ultimate purpose of what we are doing is usually to obtain equilibrium values so that we can use them in the equilibrium expression.

5. Only concentrations or pressures can be used in the equilibrium expression. In some problems the number of moles and the volume of the container or the number of moles and volume of the solution are given rather than concentrations. Don’t forget to convert these values to concentration units. Whenever a problem statement provides the volume of the aqueous solution or the volume of a gas container, check to see if a conversion is necessary to obtain molar concentrations.

6. Make sure you can use the quadratic equation.