Theories of acids and bases
Compare H1 Chemistry acid–base theories, distinguish strong and weak behaviour, calculate pH and explain indicators, buffers and ocean acidification.
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Before you begin
Acids and bases are everywhere, from the hydrochloric acid in your stomach to the buffers that keep your blood and the oceans at a steady pH. This topic starts with the Arrhenius and Brønsted–Lowry theories and conjugate acid–base pairs, explains the difference between strong and weak acids and bases, calculates pH for strong acids and bases, and then covers indicators, buffers and ocean acidification.
Be comfortable with: O-Level acids and bases, solution concentration, equilibrium constant Kc, and using logarithms on a calculator.
Learning goals
- Explain and apply the Arrhenius theory of acids and bases
- Explain and apply the Brønsted–Lowry theory of acids and bases, including conjugate acids and bases
- Explain how strong and weak acids and bases behave differently because of how fully they dissociate
- Explain pH, Ka, Kb and Kw
- Calculate [H⁺(aq)] and pH for strong acids and strong bases
- Choose a suitable indicator for an acid–base titration from data and the strengths of the acid and base
- Explain how a buffer solution controls pH
- Describe and explain uses of buffer solutions, including how CO₃²⁻/HCO₃⁻ keeps ocean pH steady and how rising atmospheric CO₂ has acidified the oceans
Syllabus statements covered
- show understanding of, and apply the Arrhenius theory of acids and bases
- show understanding of, and apply the Brønsted-Lowry theory of acids and bases, including the concept of conjugate acids and conjugate bases
- explain qualitatively the differences in behaviour between strong and weak acids and bases in terms of the extent of dissociation
- explain the terms pH; Ka; Kb; Kw [the relationship Kw = KaKb is not required]
- calculate [H+(aq)] and pH values for strong acids and strong bases
- explain the choice of suitable indicators for acid-base titrations, given appropriate data, in terms of the strengths of the acids and bases
- explain how buffer solutions control pH
- describe and explain the uses of buffer solutions including the role of CO32−/HCO3− in maintaining pH in oceans, and how the rapid increase in atmospheric carbon dioxide gas contributed to ocean acidification
Lessons
Work through them in order.
- Arrhenius and Brønsted–Lowry theoriesExplain aqueous ion formation and identify proton donors and acceptors from the changes in an equation.
- Conjugate acid–base pairsIdentify conjugate pairs and prove amphiprotic behaviour by tracking one proton and one charge unit.
- Strong and weak acids and basesSeparate acid strength from concentration, and interpret pH, Ka, Kb and Kw.
- pH of strong acids and basesCalculate strong-acid and strong-base pH after establishing ion stoichiometry and concentration.
- Choosing acid–base indicatorsSelect indicators from acid/base strengths and supplied transition ranges.
- Buffer action and capacityExplain how conjugate pairs remove added acid or base and why buffer capacity is finite.
- Buffer uses and ocean acidificationApply buffer action to practical uses and explain how carbon dioxide changes ocean pH and carbonate availability.
Practise and check
Or choose
Topic reference
Quick reference
| Question | First move | What to show |
|---|---|---|
| Brønsted–Lowry acid or base? | follow the proton through the equation | the donor is the acid; the acceptor is the base |
| Conjugate pair | the two species must differ by one H⁺ | the charge changes by one unit |
| Strong or weak? | complete or partial ionisation | not the same as concentrated or dilute |
| pH of a strong acid | [H⁺] from the concentration | pH = - log ₁₀[H⁺] |
| pH of a strong base | [OH⁻] from the formula | [H⁺] = Kw/[OH⁻], then pH |
| pH after mixing an acid and a base | moles of H⁺ and OH⁻ | excess, divided by the total volume |
| Choose an indicator | find the sharp pH change near equivalence | the whole transition range lies inside it |
| How a buffer works | which component reacts with the added H⁺ or OH⁻? | an equation, and a small change in pH |
Core knowledge to remember
- Arrhenius: an acid produces H⁺(aq) in water; a base produces OH⁻(aq) in water.
- Brønsted–Lowry: an acid is a proton donor; a base is a proton acceptor.
- A strong acid or base ionises completely in water; a weak one ionises only partially.
- pH = - log ₁₀[H⁺].
- Kₐ and Kb are the equilibrium constants for the ionisation of a weak acid and a weak base; the larger the value, the stronger the acid or base.
- Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ mol² dm⁻⁶ at 25 °C.
- A buffer is a solution that resists changes in pH when small amounts of acid or base are added. It contains a weak acid and its conjugate base.
- Ocean acidification: dissolved carbon dioxide lowers the pH of seawater and uses up carbonate ions.
Common mistakes
- Saying a base must contain hydroxide ions.
- Pairing species as conjugates when they differ by more than one proton.
- Calling a dilute acid weak, or a concentrated acid strong.
- Taking - log ₁₀[OH⁻] as the pH. It is the pOH.
- Choosing an indicator by its colour instead of its transition range.
- Saying a buffer keeps the pH constant, or has unlimited capacity.
- Saying ocean acidification means seawater has become acidic. Its pH has fallen but it is still alkaline.