AP Chemistry · AP Chemistry CED Unit 7 · 14 min read
1. Dynamic Equilibrium★★☆☆☆⏱ 4 min
The key distinction between dynamic equilibrium and static equilibrium is that reactions do not stop at equilibrium. For example, when liquid bromine is sealed in a flask, the brown color of bromine vapor stops changing intensity once equilibrium is reached: the rate of evaporation of liquid bromine equals the rate of condensation of gaseous bromine, so vapor concentration remains constant even as both processes continue.
Exam tip: On AP MCQ, any answer that claims 'at equilibrium, all reaction stops' or 'concentrations of reactants equal concentrations of products' is always wrong. Remember: equal rates, not equal concentrations, and reactions never stop at equilibrium.
2. The Equilibrium Constant and Law of Mass Action★★★☆☆⏱ 5 min
The law of mass action defines the form of the equilibrium constant expression for any balanced reversible reaction. For the general balanced reaction:
The equilibrium constant $K$ is defined as the ratio of product concentrations (raised to their stoichiometric coefficients) over reactant concentrations (raised to their stoichiometric coefficients):
$K_c$ uses molar concentrations (mol/L) for all species. For gas-phase reactions, we can also write $K_p$, which uses partial pressures of gases (typically in atm) instead of concentrations, with the same ratio of products to reactants. A critical rule for writing any equilibrium expression is: **pure solids, pure liquids, and solvents in dilute solutions do not appear in the expression**. Their concentrations are constant at constant temperature, so they are incorporated into the value of $K$ and do not need to be included explicitly. The relationship between $K_p$ and $K_c$ is:
Where $\Delta n = \text{moles of gaseous products} - \text{moles of gaseous reactants}$, $R = 0.0821 \ \text{L·atm/mol·K}$, and $T$ is absolute temperature in Kelvin.
Exam tip: Always check for pure solids and liquids before writing a $K$ or $Q$ expression on FRQ. If you include them, your expression will be marked wrong, which is an easy point to avoid losing.
3. The Reaction Quotient and Predicting Reaction Direction★★★☆☆⏱ 5 min
The reaction quotient ($Q$) is a value calculated from non-equilibrium concentrations (or partial pressures) of reactants and products at any point in a reaction before equilibrium is reached. $Q$ has the exact same form as the equilibrium constant $K$: the only difference is that $K$ uses equilibrium concentrations, while $Q$ uses current non-equilibrium concentrations. By comparing $Q$ to $K$, we can predict which direction the reaction will proceed to reach equilibrium:
If $Q < K$: The product terms (numerator) are too small, so the reaction proceeds forward (toward products) to increase $Q$ until $Q=K$.
If $Q > K$: The product terms (numerator) are too large, so the reaction proceeds reverse (toward reactants) to decrease $Q$ until $Q=K$.
If $Q = K$: The system is already at equilibrium, so no net change occurs.
Common Pitfalls
Why: Students are used to including all species in stoichiometry and rate law problems, so they forget the rule that constant concentrations do not appear.
Why: Students confuse equal rates of forward and reverse reaction with equal concentrations of reactants and products.
Why: Students mix up the direction of the reaction or mis-memorize the ratio order.
Why: Students rush through calculations and skip checking exponents.
Why: Students reverse the direction rule because they misremember which value corresponds to which shift.