TLDR

• Understand what acids and bases are using Arrhenius and Brønsted–Lowry definitions.

• Calculate pH, pOH, Ka, and Kb correctly.

• Use ICE tables for weak acids and bases.

• Master buffers, titrations, and conjugate pairs.

• Learn how equilibrium explains acid and base strength.

Why This Unit Matters

Acids and bases are everywhere on the AP Chemistry exam. They appear in multiple choice, FRQs, titration curves, equilibrium problems, and lab-based questions.

This unit tests whether you can reason with equilibrium, not just compute pH. Students who understand why pH changes and how Ka and Kb control dissociation consistently score higher than those who rely on memorization.

1. What Is an Acid or a Base

There are two core definitions used on the AP exam.

Arrhenius definition:

• Acid produces H⁺ in water.

• Base produces OH⁻ in water.

Brønsted–Lowry definition:

• Acid donates a proton.

• Base accepts a proton.

Every acid–base reaction forms a conjugate acid–base pair.
A strong acid always has a weak conjugate base, and vice versa.

2. Strong vs Weak Acids and Bases

Strong acids and bases fully dissociate in water. Weak ones establish equilibrium.

Strong acids to memorize:
HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄

Strong bases:
Group 1 hydroxides, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

Weak acids and bases partially dissociate, which is why Ka and Kb matter.

3. pH, pOH, and Kw

The pH scale measures acidity based on hydrogen ion concentration.

Key relationships:

• pH = −log[H⁺]

• pOH = −log[OH⁻]

• Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴

• pH + pOH = 14

Interpretation:

• pH < 7 acidic

• pH = 7 neutral

• pH > 7 basic

4. Acid and Base Strength (Ka and Kb)

Weak acids and bases are described using equilibrium constants.

• Ka = ([H⁺][A⁻]) / [HA]

• Kb = ([BH⁺][OH⁻]) / [B]

Larger Ka or Kb means greater dissociation and stronger acid or base.

For conjugate pairs:
Ka × Kb = Kw

This relationship frequently appears in AP free-response questions.

5. pH of Strong Acids and Bases

For strong acids, assume complete dissociation:

• [H⁺] = initial acid concentration.

For strong bases:

• [OH⁻] = base concentration × number of OH⁻ ions.

Always convert to pH or pOH using logarithms.

6. pH of Weak Acids and Bases

Weak acids and bases require ICE tables.

Example setup for a weak acid:
HA + H₂O ⇌ H⁺ + A⁻

Steps:

1. Write initial concentrations.

2. Let change be x.

3. Solve using Ka or Kb.

If Ka or Kb is very small, x is often negligible. This assumption must be justified.

7. Buffers

Buffers resist changes in pH when small amounts of acid or base are added.

A buffer contains:

• A weak acid and its conjugate base, or

• A weak base and its conjugate acid.

Key equation:
pH = pKa + log([A⁻]/[HA])

Buffers work best when the concentrations of acid and conjugate base are similar.

8. Titrations

Titrations are used to determine unknown concentrations.

Important points:

• Equivalence point: moles acid = moles base.

• Endpoint: indicator color change.

Common cases:

• Strong acid–strong base → pH = 7 at equivalence.

• Weak acid–strong base → pH > 7 at equivalence.

• Strong acid–weak base → pH < 7 at equivalence.

Use mole ratios unless the reaction is strictly 1:1.

9. Polyprotic Acids

Polyprotic acids donate more than one proton.

Examples:
H₂SO₄, H₂CO₃, H₃PO₄

Each proton has its own Ka value.
The first dissociation dominates pH, so later Ka values are usually ignored in pH calculations.

Common Pitfalls

• Treating weak acids as fully dissociated.

• Forgetting to multiply OH⁻ concentration for strong bases.

• Mixing up pH and pOH.

• Using Henderson–Hasselbalch when no buffer is present.

• Assuming strong conjugate pairs exist together.

Tutor Tip

On AP free-response questions, never say an acid is strong or weak without justification.

Strong answer example:
“The larger Ka value indicates greater dissociation, so the acid is stronger.”

Using equilibrium language consistently earns full reasoning credit.

Ready to Master Acids and Bases?

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