Acids and Bases Study Guide

Acids and bases is one of the biggest topics in any chemistry course, and it connects to almost everything else: equilibrium, energy changes, titrations, organic chemistry and biology. This study guide pulls it all together in one place. Use it to check what you know, find the gaps, and jump to the detailed article for anything that needs more work.

The guide is split into three levels. Foundation material is needed by everyone; intermediate is typical of the final years of secondary school; advanced covers the harder content found in A-level, AP, IB Higher Level and first-year university courses.

Level 1: foundation

Checklist

• I can describe the properties of acids and alkalis (taste, indicators, reactions, pH).
• I know that acids produce H⁺ ions and alkalis produce OH⁻ ions in water.
• I can use the pH scale: below 7 acidic, 7 neutral, above 7 alkaline.
• I know the colours of litmus, universal indicator and phenolphthalein in acid and alkali.
• I can write word and symbol equations for the reactions of acids with metals, bases and carbonates.
• I can name the salt produced from any common acid.
• I know that an alkali is a soluble base.

Key facts

• Acid + metal → salt + hydrogen
• Acid + base → salt + water
• Acid + carbonate → salt + water + carbon dioxide
• Neutralisation (ionic): H⁺ + OH⁻ → H₂O
• Hydrochloric acid → chlorides; sulfuric acid → sulfates; nitric acid → nitrates.

Tests to remember

Test	Result
Hydrogen	lighted splint gives a squeaky pop
Carbon dioxide	turns limewater milky
Acid	turns blue litmus red
Alkali	turns red litmus blue

Go deeper

What is an acid? · What is a base? · Acids vs bases · Bases vs alkalis · Naming salts · Litmus paper · Universal indicator

Level 2: intermediate

Checklist

• I can explain the difference between strong and weak acids in terms of ionisation.
• I can explain the difference between concentrated and dilute.
• I can calculate pH from [H⁺] and [H⁺] from pH.
• I know that a tenfold change in [H⁺] changes pH by one unit.
• I can use Brønsted–Lowry definitions and identify conjugate acid–base pairs.
• I can describe how to carry out a titration accurately.
• I can do titration calculations, including 1 : 2 ratios.
• I can explain how to prepare soluble and insoluble salts.

Key equations

Quantity	Equation
pH	pH = −log[H⁺]
[H⁺] from pH	[H⁺] = 10⁻ᵖᴴ
Moles in solution	n = c × V (V in dm³)
Dilution	c₁V₁ = c₂V₂

Key ideas

• Strong acid: completely ionised (HCl, HNO₃, H₂SO₄). Weak acid: partially ionised (ethanoic, citric, carbonic).
• Strength ≠ concentration. A concentrated weak acid can have a lower pH than a dilute strong one.
• Brønsted–Lowry acid = proton donor; base = proton acceptor. Conjugate pairs differ by one H⁺.
• In titrations, rinse the burette and pipette with their own solutions; average only concordant titres (within 0.10 cm³).

Go deeper

Strong vs weak acids · Concentrated vs strong · How to calculate pH · Brønsted–Lowry acids and bases · Conjugate pairs · Acid–base titration · Titration calculations · The strong acids

Level 3: advanced

Checklist

• I can write and use expressions for Ka, Kb and Kw.
• I can calculate the pH of strong bases, weak acids and weak bases.
• I can use pKa and pKb, and know pKa + pKb = 14 at 25 °C for a conjugate pair.
• I can explain why neutral pH changes with temperature.
• I can sketch and interpret titration curves for all four strong/weak combinations.
• I can choose an appropriate indicator.
• I can explain how buffers work and calculate buffer pH, including after additions.
• I can predict whether a salt solution is acidic, neutral or basic.
• I can use the Lewis definition of acids and bases.

Key equations

Quantity	Equation
Ionic product of water	Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C
pH + pOH	= 14.00 at 25 °C
Acid dissociation constant	Ka = [H⁺][A⁻] ÷ [HA]
pKa	pKa = −log Ka
Weak acid pH (approx.)	[H⁺] ≈ √(Ka × c)
Weak base (approx.)	[OH⁻] ≈ √(Kb × c)
Conjugate pair	Ka × Kb = Kw
Henderson–Hasselbalch	pH = pKa + log([A⁻] ÷ [HA])

Key ideas

• Half-equivalence in a weak acid titration: pH = pKa.
• Equivalence pH: 7 (strong/strong), above 7 (weak acid/strong base), below 7 (strong acid/weak base).
• Buffer range ≈ pKa ± 1; capacity is highest when [HA] = [A⁻].
• Neutral means [H⁺] = [OH⁻], not necessarily pH 7.
• Salt hydrolysis: anions of weak acids are basic; cations of weak bases and small, highly charged metal ions are acidic.

Go deeper

pOH and pH · Kw · Ka and pKa · Kb and pKb · pH of weak acids · pH of weak bases · Ka × Kb = Kw · Titration curves · Choosing an indicator · Henderson–Hasselbalch · Buffer calculations · Buffer capacity · Salt hydrolysis · Lewis acids and bases · Polyprotic acids

Common exam question types

1. “Describe” observations: fizzing, colour changes, solids disappearing, temperature rise.
2. “Explain” differences in strength, often comparing two acids at equal concentration using pH, conductivity or reaction rate.
3. Calculations: pH, titrations, Ka from pH, buffer pH.
4. Practical: titration method, sources of error, choosing an indicator.
5. Graphs: interpreting titration curves, reading pKa from half-equivalence.
6. Definitions: Brønsted–Lowry, strong/weak, buffer, salt.

Exam technique

• Definitions must be precise. “A strong acid is completely ionised in aqueous solution” earns the mark; “a strong acid is very acidic” doesn’t.
• Always show the mole ratio in titration calculations, even when it’s 1 : 1.
• Give pH to two decimal places unless told otherwise.
• State symbols are often required in ionic equations.
• Check direction: acids below 7, alkalis above 7, and buffers change pH by only small amounts.
• For “explain” questions, link the observation to particles: “more H⁺ ions per dm³, so more frequent collisions with magnesium atoms, so faster fizzing”.

A one-week revision plan

Day	Focus	Activity
1	Definitions and properties	Make flashcards for every definition in the checklists
2	Reactions and salts	Write 20 word and symbol equations from memory
3	pH calculations	Work through pH practice problems
4	Titrations	Do titration calculations and learn the method
5	Ka, Kb and weak acids (advanced)	Redo worked examples without looking
6	Buffers and curves (advanced)	Sketch all four titration curves from memory
7	Mixed practice	Try acid–base reactions practice under timed conditions

Self-test: ten quick questions

1. What ion do all acids release in water?
2. What is the pH of 0.001 mol/dm³ HCl?
3. Name the salt from potassium hydroxide and sulfuric acid.
4. Is ammonia a strong or weak base?
5. What is the conjugate base of HSO₄⁻?
6. Which indicator suits a weak acid–strong base titration?
7. What’s the pH at the half-equivalence point of a titration of an acid with pKa 4.2?
8. Is sodium ethanoate solution acidic, neutral or basic?
9. What is Kw at 25 °C?
10. What is a buffer?

Answers: (1) H⁺ (H₃O⁺). (2) 3.00. (3) Potassium sulfate. (4) Weak. (5) SO₄²⁻. (6) Phenolphthalein. (7) 4.2. (8) Basic. (9) 1.0 × 10⁻¹⁴ mol² dm⁻⁶. (10) A solution that resists changes in pH when small amounts of acid or base are added.

Key takeaways

• Master the foundation reactions and naming first; everything else builds on them.
• Strength vs concentration and the log nature of pH are the two ideas most often tested.
• At higher levels, Ka, buffers and titration curves are all the same equilibrium idea in different forms.
• Use the pH calculator to check your practice answers.