Name common acids and alkalis with Brønsted–Lowry theory
Recognise the syllabus's common acids and alkalis and describe proton donation and acceptance.
Section 7 · Equilibria
AS Level; Cambridge International Chemistry 9701, 2025–2027 syllabus. Equilibrium constants apply at a fixed temperature; include only gaseous or aqueous species in the relevant expressions. For acid–base work, distinguish strength from concentration. Written practice is self-checked
Before you begin
- Balance equations, interpret reversible arrows and use concentration in mol dm⁻³.
Work on paper. Try each question before revealing help, and explain why your method works.
Quick readiness check
What does a Brønsted–Lowry acid do?
What you will learn
- State the named acid and alkali formulas
- Define Brønsted–Lowry acids and bases
- Identify conjugate acid–base pairs in equations
Learn the limited syllabus list
Common acids: hydrochloric acid HCl, sulfuric acid H₂SO₄, nitric acid HNO₃ and ethanoic acid CH₃COOH. Common alkalis: sodium hydroxide NaOH, potassium hydroxide KOH and ammonia NH₃.
An alkali is a base that produces OH⁻ ions in aqueous solution. Some alkalis contain OH⁻ directly; ammonia reacts with water to produce NH₄⁺ and OH⁻.
Acids donate protons; bases accept them
A Brønsted–Lowry acid is a proton donor. A Brønsted–Lowry base is a proton acceptor. Proton transfer creates a conjugate acid–base pair: after donating H⁺, the acid becomes its conjugate base; after accepting H⁺, the base becomes its conjugate acid.
Water can act as either acid or base depending on its reaction partner. Track H⁺ transfer rather than relying only on whether the formula contains an H or OH group.
Pause and explain
Which is a common acid in the syllabus list?
Worked example
In HCl + H₂O → H₃O⁺ + Cl⁻, identify the acid, base and conjugate pair.
Show the worked solution
- HCl donates H⁺ to water.
- HCl is the acid and H₂O is the base.
- Cl⁻ is the conjugate base; H₃O⁺ is the conjugate acid.
Answer HCl is the Brønsted acid, H₂O is the base, and the conjugate pairs are HCl/Cl⁻ and H₂O/H₃O⁺.
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Write your answer and reasoning first. Use a hint only if you are stuck. The model solution is for self-checking; your written work is not automatically marked.
List the four common acids and three common alkalis specified in this syllabus topic.
Give me a hint
Separate the two lists and include formulae.
Compare my reasoning
- Acids: HCl, H₂SO₄, HNO₃, CH₃COOH.
- Alkalis: NaOH, KOH, NH₃.
- Attach the correct names to the formulas.
Acids: hydrochloric acid HCl, sulfuric acid H₂SO₄, nitric acid HNO₃ and ethanoic acid CH₃COOH. Alkalis: sodium hydroxide NaOH, potassium hydroxide KOH and ammonia NH₃.
Look for these in your work
- All four acids
- All three alkalis
- Names matched to formulas
For NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, identify the Brønsted acid and base.
Give me a hint
Follow the proton from donor to acceptor.
Compare my reasoning
- Water donates H⁺ and becomes OH⁻.
- NH₃ accepts H⁺ and becomes NH₄⁺.
- Thus water is acid and ammonia is base.
H₂O is the Brønsted acid; NH₃ is the Brønsted base. Their conjugates are OH⁻ and NH₄⁺.
Look for these in your work
- Proton donor identified
- Proton acceptor identified
- Conjugates matched
Explain how H₂SO₄ and HSO₄⁻ can be a conjugate acid–base pair.
Give me a hint
A conjugate pair differs by one proton.
Compare my reasoning
- H₂SO₄ can donate H⁺.
- After donation, the species is HSO₄⁻.
- The pair differs by one proton.
H₂SO₄ is the acid and HSO₄⁻ its conjugate base; they differ by one H⁺.
Look for these in your work
- Proton transfer stated
- Conjugate role correct
- One-proton difference identified
Common mistakes
- Defining a Brønsted acid as an OH⁻ donor
- Calling proton acceptance oxidation
- Forgetting ammonia acts as an alkali in water
- Mismatching a species with its conjugate partner
What can you explain now?
What does a Brønsted–Lowry acid do?
Compare with the explanation
Donates a proton
An acid is a proton donor.
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Make it stick
How do you identify a Brønsted acid, base and their conjugates from a reaction equation?
If you needed the solution, close it and solve the problem again from a blank page. A correct answer is stronger when you can explain the reason for each step.