AP Chemistry Unit 8 Review: Acids and Bases

Every Unit 8 topic, what's actually tested, and what the exam explicitly excludes -- the second-largest unit on the exam, verified against the current CED.

Unit 8 applies the equilibrium principles from Unit 7 to the reaction type that dominates real chemistry: proton transfer. At 11–15% of the multiple-choice section across 11 topics and 14–16 class periods, it's the second-heaviest unit after Unit 3. Here's everything in it, topic by topic, verified against the current CED -- including four exclusion statements that cut real content most study guides still teach as if it were fair game.

Unit 8 at a Glance

Unit 8 (Acids and Bases) builds directly on Unit 7's equilibrium framework. See the full AP Chemistry units breakdown for how it fits alongside the other 8 units, and the Course and Exam Description for the full framework.

Introduction to Acids and Bases

pH = −log[H₃O⁺] and pOH = −log[OH⁻]. Water autoionizes with K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C, so at 25°C, pK_w = 14 = pH + pOH. Because K_w is temperature-dependent, neutral water's pH shifts away from 7.0 at temperatures other than 25°C -- neutral doesn't always mean pH 7.

Conversions between pH, pOH, [H₃O⁺], and [OH⁻], the pH of common substances, and a worked example of neutral water at 50°C are in the AP Chemistry pH scale guide.

The acid-base definitions behind all of this are covered with worked examples in Bronsted-Lowry acids, bases, and conjugate pairs.

pH and pOH of Strong Acids and Bases

The six strong acids (HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃) ionize completely, so [H₃O⁺] equals the acid's initial concentration directly. Strong bases (group I and II hydroxides) dissociate completely too -- but a group II hydroxide like Ca(OH)₂ releases two OH⁻ per formula unit, so [OH⁻] is double the initial concentration.

Worked example: What's the pH of 0.025 M HCl? Since HCl is a strong acid, [H₃O⁺] = 0.025 M directly. pH = −log(0.025) = 1.60.

For the full strong acid and base lists, plus Ba(OH)₂, concentration-from-pH, and acid-base mixing problems, see AP Chemistry strong acids and bases.

Weak Acid and Base Equilibria

A weak acid only partially ionizes, governed by K_a = [H₃O⁺][A⁻] / [HA]. A weak base only partially ionizes, governed by K_b = [OH⁻][HB⁺] / [B]. For any conjugate acid-base pair, K_w = K_a × K_b, so pK_w = pK_a + pK_b.

Worked example: Find the pH of 0.10 M HF (K_a = 6.8 × 10⁻⁴). Set up the equilibrium expression: 6.8 × 10⁻⁴ = x² / (0.10 − x). Since K_a is small, approximate 0.10 − x ≈ 0.10: x² = 6.8 × 10⁻⁵, so x = [H₃O⁺] = 8.2 × 10⁻³ M. Check the approximation: 8.2 × 10⁻³ M is 8.2% of 0.10 M, above the 5% limit, so solve the quadratic instead: x = 7.9 × 10⁻³ M and pH = 2.10.

The K_a and pK_a scale, a pK_a table of common weak acids, and more weak acid pH problems, including when the 5% check fails, are in AP Chemistry weak acids, Ka, and pKa.

Acid-Base Reactions and Buffers

What happens when you mix an acid and a base depends on which species (if any) is left over:

Strong + strong reacts completely to form water; pH comes from whichever reagent is in excess.
Weak acid + strong base, weak acid in excess forms a buffer -- use the Henderson-Hasselbalch equation.
Weak acid + strong base, exactly equimolar leaves only the conjugate base, which hydrolyzes water to give a slightly basic pH.
Weak base + strong acid, exactly equimolar leaves only the conjugate acid, giving a slightly acidic pH.

The same conjugate-ion reasoning predicts the pH of any salt solution, with worked pH calculations, in AP Chemistry salt hydrolysis.

Acid-Base Titrations

A titration curve plots pH against volume of titrant added. At the equivalence point, moles of titrant equal moles of analyte. For a weak acid or base titration, the half-equivalence point is special: since [HA] = [A⁻] there, pH = pK_a exactly -- the fastest way to read a K_a straight off a graph.

At the equivalence point of a strong acid/strong base titration, pH is neutral (7 at 25°C). For a weak acid titrated with a strong base, the equivalence point is basic (pH > 7) because the leftover conjugate base hydrolyzes water. For polyprotic acids, the titration curve shows multiple equivalence points, each with its own half-equivalence pK_a.

Not tested: computing the exact concentration of every species at each point along a polyprotic acid's titration curve is explicitly excluded. You still need qualitative reasoning about which species dominates at each stage, and full quantitative work for monoprotic titrations remains in scope.

For the full quantitative method, calculating pH before, at, and after the equivalence point, see the AP Chemistry titration curve pH calculations guide.

Molecular Structure of Acids and Bases

Strong acids have very weak conjugate bases, stabilized by electronegativity, inductive effects, or resonance. Carboxylic acids are the classic weak-acid structural class. Strong bases (group I/II hydroxides) have very weak conjugate acids. Common weak bases include nitrogenous bases like ammonia and carboxylate ions. In general, more electronegative atoms near the acidic proton stabilize the conjugate base and increase acid strength.

pH and pK_a

Comparing solution pH to an acid's pK_a tells you which form dominates: if pH < pK_a, the protonated (acid) form dominates; if pH > pK_a, the deprotonated (base) form dominates. Acid-base indicators exploit this -- they change color between their protonated and deprotonated forms, so a good indicator for a titration has a pK_a close to the pH at the equivalence point.

Properties of Buffers

A buffer contains a large concentration of both members of a conjugate acid-base pair. Added acid reacts with the conjugate base component; added base reacts with the conjugate acid component -- that's the entire mechanism behind pH stabilization.

For examples of which mixtures count as buffers, the reactions with added acid or base, and how to choose a buffer by pK_a, see AP Chemistry buffers.

Henderson-Hasselbalch Equation

pH = pK_a + log([A⁻]/[HA]). Small additions of acid or base barely shift the [A⁻]/[HA] ratio, which is exactly why a buffer resists large pH swings.

Worked example: A buffer contains 0.30 M acetic acid (pK_a = 4.74) and 0.45 M sodium acetate. pH = 4.74 + log(0.45/0.30) = 4.74 + log(1.5) = 4.74 + 0.18 = 4.92.

Not tested (two exclusions in this topic): deriving the Henderson-Hasselbalch equation is explicitly excluded -- you apply the formula, you don't prove it. Computing the numerical pH change after adding acid or base to a buffer is also excluded -- you explain qualitatively why the change is small, without calculating the new pH.

For more practice, including solving for a target ratio and finding the pH of a buffer formed by a reaction, see the full Henderson-Hasselbalch equation guide.

Buffer Capacity

Raising both buffer components' concentrations at a fixed ratio keeps pH the same but increases the buffer's capacity to absorb added acid or base. A buffer with more conjugate acid than base has greater capacity against added base; one with more conjugate base than acid has greater capacity against added acid.

pH and Solubility

A salt's solubility is pH-sensitive whenever one of its ions is a weak acid, a weak base, or the hydroxide ion -- reasoned qualitatively through Le Chatelier's principle.

Not tested: calculating solubility as a numerical function of pH is explicitly excluded. This topic is qualitative-only on the AP Exam.

Common Mistakes in Unit 8

Assuming every weak-acid titration hits pH 7 at the equivalence point. Only strong acid/strong base equivalence points are neutral -- weak acid equivalence points are basic, weak base equivalence points are acidic.
Forgetting the factor of 2 for group II hydroxides. Ca(OH)₂ and Ba(OH)₂ release two OH⁻ ions per formula unit -- halving that factor is a common pOH error.
Trying to calculate an exact new pH after adding acid to a buffer. That numerical computation is explicitly excluded -- only qualitative reasoning about buffer stability is required.
Applying Henderson-Hasselbalch when no buffer exists. The equation only applies when both the conjugate acid and conjugate base are present in significant amounts -- not at the equivalence point, and not for a pure weak acid solution.
Treating pH-solubility connections as a calculation problem. The CED keeps this topic qualitative; spend your study time on the Le Chatelier reasoning, not a formula.

How Unit 8 Connects to the Rest of the Course

Every calculation here is a direct application of the equilibrium-constant reasoning from Unit 7 -- K_a, K_b, and K_sp all share the same algebraic structure.
The proton-transfer redox parallel resurfaces in Unit 9's electrochemistry, where half-reactions play the same "two conjugate species" role that acids and bases play here.
Buffer and titration reasoning is a favorite AP free-response context precisely because it forces you to combine stoichiometry, equilibrium, and graph-reading in one question.

Frequently Asked Questions

What topics are in AP Chemistry Unit 8?

Introduction to Acids and Bases, pH and pOH of Strong Acids and Bases, Weak Acid and Base Equilibria, Acid-Base Reactions and Buffers, Acid-Base Titrations, Molecular Structure of Acids and Bases, pH and pKa, Properties of Buffers, the Henderson-Hasselbalch Equation, Buffer Capacity, and pH and Solubility -- 11 topics in total.

How much is Unit 8 worth on the AP Chemistry exam?

Eleven to fifteen percent of the multiple-choice section -- the second-largest unit after Unit 3, across roughly 14 to 16 class periods.

Does AP Chemistry require calculating the pH change after adding acid or base to a buffer?

No. The CED explicitly excludes computing the numerical pH change after an acid or base addition -- you only need to explain qualitatively why a buffered solution resists a large pH change.

Do I need to derive the Henderson-Hasselbalch equation?

No. Derivation of the Henderson-Hasselbalch equation is explicitly excluded. You need to apply pH = pKa + log([A-]/[HA]), not prove where it comes from.

Does AP Chemistry test solubility as a function of pH with calculations?

No. Computations of solubility as a function of pH are explicitly excluded -- pH's effect on a salt's solubility is tested only qualitatively, using Le Chatelier's principle.

Sourced from College Board's official AP Chemistry Course and Exam Description, Effective Fall 2024. This page describes the document's real content and current exclusion statements; it is not a copy of it and is not affiliated with or endorsed by College Board.