AP Chemistry Equilibrium — Complete Guide with Practice Problems (2026)

Chemical equilibrium is one of the highest-weighted topics on AP Chemistry, appearing in both multiple choice and FRQ sections every year. This guide covers every equilibrium concept tested on the AP exam, with worked examples and practice problems.

What Is Chemical Equilibrium?

A reaction reaches equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant — but the reaction has not stopped. It is a dynamic equilibrium.

$$\text{reactants} \rightleftharpoons \text{products}$$

Key point: Equilibrium does not mean equal concentrations. It means the rate of forward = rate of reverse.

The Equilibrium Constant (Keq)

For a generic reaction: $aA + bB \rightleftharpoons cC + dD$

$$K_{eq} = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$

Rules for writing K expressions:

• Use molar concentrations (mol/L) for aqueous species and gases
• Pure solids and pure liquids are NOT included in K expressions (their concentration is constant)
• Coefficients in the balanced equation become exponents in K

Example: Write Keq for $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

$$K_c = \frac{[\text{NH}_3]^2}{[\text{N}_2][\text{H}_2]^3}$$

Interpreting K Values

Kp vs Kc

Kc uses molar concentrations (mol/L). Kp uses partial pressures (atm).

$$K_p = K_c(RT)^{\Delta n}$$

Where $\Delta n$ = moles of gaseous products − moles of gaseous reactants, R = 0.08206 L·atm/mol·K, T = temperature in Kelvin.

Example: For $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

$\Delta n = 2 - (1 + 3) = -2$

$$K_p = K_c(RT)^{-2} = \frac{K_c}{(RT)^2}$$

The Reaction Quotient Q

Q has the same form as K, but uses current concentrations (not equilibrium concentrations).

AP exam tip: Q vs K problems are extremely common. Always state which direction the reaction shifts and why.

ICE Tables

ICE (Initial, Change, Equilibrium) tables are the standard method for calculating equilibrium concentrations.

Example: 0.500 mol of N₂ and 0.500 mol of H₂ are placed in a 1.00 L container. At equilibrium, [NH₃] = 0.060 M. Find Kc.

Reaction: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$

Given: $2x = 0.060 \Rightarrow x = 0.030$

$[\text{N}_2] = 0.500 - 0.030 = 0.470$ M $[\text{H}_2] = 0.500 - 0.090 = 0.410$ M $[\text{NH}_3] = 0.060$ M

$$K_c = \frac{(0.060)^2}{(0.470)(0.410)^3} = \frac{0.0036}{0.0324} = \boxed{0.111}$$

Le Chatelier's Principle

If a system at equilibrium is disturbed, it will shift to minimize the disturbance and re-establish equilibrium.

Concentration Changes

Pressure/Volume Changes (Gas-phase reactions only)

Example: $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$

Increasing pressure shifts toward the side with fewer moles of gas: products (2 moles vs 4 moles). Reaction shifts forward → more NH₃ produced.

Temperature Changes

Key distinction: Concentration and pressure changes shift equilibrium but do NOT change K. Temperature changes shift equilibrium AND change K.

Gases in Equilibrium — Partial Pressures

For gas-phase equilibria, you can use partial pressures directly in Kp expressions.

Dalton's Law: $P_{total} = P_A + P_B + P_C + \cdots$

Mole fraction: $\chi_A = \frac{n_A}{n_{total}}$, then $P_A = \chi_A \cdot P_{total}$

Example FRQ-style: A mixture of 0.40 atm SO₂, 0.20 atm O₂, and 0.80 atm SO₃ is placed in a container.

Reaction: $2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g)$ with $K_p = 1.7 \times 10^2$

$$Q_p = \frac{(P_{\text{SO}3})^2}{(P{\text{SO}2})^2(P{\text{O}_2})} = \frac{(0.80)^2}{(0.40)^2(0.20)} = \frac{0.64}{0.032} = 20$$

Since $Q_p (20) < K_p (170)$, the reaction shifts forward to produce more SO₃.

Solubility Equilibrium (Ksp)

When a slightly soluble ionic compound dissolves:

$$\text{AB}(s) \rightleftharpoons \text{A}^+(aq) + \text{B}^-(aq)$$ $$K_{sp} = [\text{A}^+][\text{B}^-]$$

The solid is not included in Ksp because it's a pure solid.

Common Ksp expressions:

Will a Precipitate Form?

Compare the ion product Q to Ksp:

• Q > Ksp → precipitate forms (solution is supersaturated)
• Q < Ksp → no precipitate (solution is unsaturated)
• Q = Ksp → at saturation

AP Chemistry Equilibrium Practice Problems

Problem 1: For the reaction $\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g)$, Kc = 54.3 at 430°C. If [H₂] = [I₂] = 0.10 M and [HI] = 0.50 M, which direction does the reaction shift?

$Q = \frac{(0.50)^2}{(0.10)(0.10)} = \frac{0.25}{0.01} = 25$. Since Q (25) < K (54.3), reaction shifts forward.

Problem 2: The Ksp of BaSO₄ is 1.1 × 10⁻¹⁰. What is the molar solubility of BaSO₄?

$\text{BaSO}_4 \rightleftharpoons \text{Ba}^{2+} + \text{SO}4^{2-}$ Let solubility = s. Then $K{sp} = s \cdot s = s^2$ $s = \sqrt{1.1 \times 10^{-10}} = 1.05 \times 10^{-5}$ mol/L

Problem 3: If the temperature is raised for an endothermic reaction at equilibrium, what happens to K and to the equilibrium position?

K increases (temperature change always changes K). The equilibrium shifts forward (toward products), consistent with the higher K.

How Equilibrium Is Tested on AP Chemistry FRQs

Equilibrium appears on nearly every AP Chemistry FRQ section. Common question types:

• Write the Kc or Kp expression for a given reaction
• Use an ICE table to find equilibrium concentrations
• Calculate Q and determine shift direction
• Apply Le Chatelier's principle to predict the effect of a change
• Calculate molar solubility from Ksp or vice versa
• Determine whether a precipitate forms given ion concentrations

AP reader expectation: Show your ICE table. Label rows clearly (I, C, E). Show algebraic setup before plugging in numbers.

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