Strong and weak acids (HT only) — AQA GCSE Combined Science
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Strong and weak acids (HT only) explained
In aqueous solution, a strong acid donates essentially all of its acidic hydrogen to water, meaning ionisation is complete.
Read the full explanation
For hydrochloric acid, HCl(aq) → H⁺(aq) + Cl⁻(aq); every dissolved HCl molecule produces one hydrogen ion and one chloride ion, leaving no intact HCl. Nitric acid behaves similarly: HNO₃(aq) → H⁺(aq) + NO₃⁻(aq). Sulfuric acid is diprotic, so one formula unit releases two hydrogen ions: H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq). Complete ionisation creates a high concentration of H⁺(aq), resulting in low pH values and vigorous reactions with metals and bases. The key concept is the completeness of ionisation rather than concentration; therefore, even a dilute strong acid remains fully ionised.
A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids.
A weak acid dissolves in water but only some of its molecules donate a hydrogen ion, so ionisation is partial and reversible. Ethanoic acid establishes an equilibrium: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq). At any moment, many intact CH₃COOH molecules remain alongside a relatively low concentration of H⁺(aq) ions. Citric acid and carbonic acid behave in the same partial way. Because the H⁺(aq) concentration is lower than in a strong acid of the same concentration, the pH is higher and reactions with metals or bases are slower and less vigorous. The equilibrium arrow (⇌) is essential: it shows that both forward ionisation and reverse recombination occur simultaneously.
For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
Acids release hydrogen ions, H⁺, in aqueous solution. A strong acid, such as hydrochloric acid, is completely ionised, so a 0.1 mol/dm³ solution contains 0.1 mol/dm³ of H⁺ ions. A weak acid, such as ethanoic acid, is only partially ionised, so the same total concentration yields far fewer H⁺ ions. Because pH depends inversely on the concentration of H⁺ ions, the strong acid has the lower pH. For example, 0.1 mol/dm³ hydrochloric acid has a pH of 1, while 0.1 mol/dm³ ethanoic acid has a pH of about 3. The key comparison must be made at equal concentration: a stronger acid means a greater extent of ionisation, more H⁺ ions, and consequently a lower pH value.
As the pH decreases by one unit, the hydrogen ion concentration of the solution increases by a factor of 10.
pH is a logarithmic scale, meaning each unit change represents a tenfold change in hydrogen ion concentration. If the pH decreases by one unit, the solution becomes more acidic and the concentration of H⁺ ions becomes 10 times greater. For example, a solution of pH 4 has 10 times the H⁺ concentration of a solution of pH 5, and 100 times that of a solution of pH 6. Conversely, if pH increases by one unit, H⁺ concentration falls by a factor of 10. This mathematical relationship is assessed at Higher Tier only. It allows you to compare solutions without calculating exact concentrations: a two-unit drop in pH means a 10² = 100-fold increase in H⁺ concentration.
use and explain the terms dilute and concentrated (in terms of amount of substance), and weak and strong (in terms of the degree of ionisation) in relation to acids
This Higher Tier statement requires students to use and explain four key terms about acids. Dilute and concentrated describe the amount of substance dissolved in a given volume: a dilute acid contains a small amount of acid per unit volume, while a concentrated acid contains a large amount. Weak and strong describe the degree of ionisation in aqueous solution: a strong acid ionises completely, whereas a weak acid ionises partially, forming an equilibrium mixture with un-ionised molecules. For example, hydrochloric acid is strong and ethanoic acid is weak, but either can be dilute or concentrated. Students must use these terms accurately and explain how the degree of ionisation affects H⁺ ion concentration and pH.
describe neutrality and relative acidity in terms of the effect on hydrogen ion concentration and the numerical value of pH (whole numbers only).
The pH scale is a numerical way to express the relative acidity or alkalinity of an aqueous solution based on hydrogen ion (H⁺) concentration. A neutral solution, such as pure water, has a pH of exactly 7. Acidic solutions have a pH below 7; the lower the pH number, the higher the H⁺ concentration and the more acidic the solution. For example, a solution with pH 3 has a higher H⁺ concentration than one with pH 5. Alkaline solutions have a pH above 7, meaning their H⁺ concentration is lower than that of neutral water. The specification requires whole-number pH values only, so you should compare pH 1, 2, 3 and so on, rather than decimals.
Your focus
- Define a strong acid as one that is completely ionised in aqueous solution.
- Write balanced ionisation equations for hydrochloric, nitric and sulfuric acids, including state symbols.
- Explain how complete ionisation leads to a high concentration of H⁺(aq) ions and typical strong-acid behaviour.
Show all 18 objectives
- Define a weak acid as one that is only partially ionised in aqueous solution.
- Write the reversible ionisation equation for ethanoic acid and identify citric and carbonic acids as weak acids.
- Explain how partial ionisation leads to a lower H⁺(aq) concentration and less vigorous acidic behaviour than a strong acid of the same concentration.
- Describe strong acids as completely ionised and weak acids as partially ionised in aqueous solution.
- Relate the extent of ionisation to the concentration of hydrogen ions, H⁺, at a given acid concentration.
- Predict and justify which of two acids at the same concentration has the lower pH.
- State that a one-unit decrease in pH corresponds to a tenfold increase in H⁺ concentration.
- Calculate the factor change in H⁺ concentration for a pH change of more than one unit.
- Interpret pH differences to compare the acidity of two solutions.
- Use the terms dilute and concentrated correctly in terms of the amount of substance dissolved in a given volume.
- Explain the terms weak and strong in terms of the degree of ionisation of acids in aqueous solution.
- Apply these terms to named acids and use them to explain differences in pH between strong and weak acids of the same concentration.
- State the pH value of a neutral aqueous solution.
- Compare acidic and alkaline solutions in terms of their pH values and hydrogen ion concentrations.
- Use whole-number pH values to explain how the hydrogen ion concentration changes as acidity changes.
Strong and weak acids (HT only) exam tips
Marking Points
- A strong acid is one that ionises completely when dissolved in water, leaving no undissociated acid molecules.
- Hydrochloric acid ionises completely: HCl(aq) → H⁺(aq) + Cl⁻(aq), yielding one H⁺ ion per molecule.
- Nitric acid ionises completely: HNO₃(aq) → H⁺(aq) + NO₃⁻(aq), producing one H⁺ ion per molecule.
- Sulfuric acid is diprotic and ionises completely: H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq), releasing two H⁺ ions per molecule.
- Complete ionisation produces a high concentration of H⁺(aq) ions, leading to a low pH and vigorous reactions with metals.
- A weak acid is only partially ionised in aqueous solution, meaning a significant proportion of intact acid molecules remains.
- Ethanoic acid ionises reversibly: CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq).
- Citric acid and carbonic acid are additional examples of weak acids that ionise only partially in water.
- Partial ionisation results in a lower concentration of H⁺(aq) ions than a strong acid of the same concentration, leading to a higher pH.
- The equilibrium symbol (⇌) must be used because ionisation is reversible, with ions recombining to reform the acid molecules.
- Acids produce hydrogen ions, H⁺, in aqueous solution.
- A strong acid is completely ionised, whereas a weak acid is only partially ionised.
- At the same total acid concentration, a greater degree of ionisation results in a higher concentration of H⁺ ions.
- As the concentration of H⁺ ions increases, the pH decreases, meaning the stronger acid has the lower pH.
- A valid comparison requires both acids to be at the same concentration, for example comparing 0.1 mol/dm³ hydrochloric acid with 0.1 mol/dm³ ethanoic acid.
- State that the pH scale is logarithmic, based on hydrogen ion concentration.
- Explain that a decrease of one pH unit corresponds to a tenfold increase in H⁺ concentration.
- Apply the relationship to a numerical example, such as pH 3 having 10 times the H⁺ concentration of pH 4.
- Extend the reasoning to changes of more than one unit, for example a two-unit decrease giving a 10² = 100-fold increase.
- Recognise that an increase in pH by one unit means H⁺ concentration decreases by a factor of 10.
- Defines dilute and concentrated in terms of the amount of substance (acid) dissolved in a given volume of solution.
- Defines strong and weak in terms of the degree of ionisation of the acid in aqueous solution.
- Explains that a strong acid ionises completely, producing a high concentration of H⁺ ions for a given concentration of acid.
- Explains that a weak acid ionises partially, so an equilibrium mixture forms and the concentration of H⁺ ions is lower than for a strong acid of the same concentration.
- Uses the terms correctly with named examples, such as dilute hydrochloric acid, concentrated ethanoic acid, strong hydrochloric acid and weak ethanoic acid.
- Explains how the degree of ionisation affects pH, so that a strong acid has a lower pH than a weak acid at the same concentration.
- Neutral solutions have a pH of 7.
- Acidic solutions have a pH below 7 and a higher hydrogen ion concentration than neutral solutions.
- Alkaline solutions have a pH above 7 and a lower hydrogen ion concentration than neutral solutions.
- The lower the pH number, the greater the hydrogen ion concentration and the more acidic the solution.
- Whole-number pH values are sufficient when comparing acidity or neutrality in this specification.
Examiner Tips
- 💡When asked to explain strong acidity, explicitly state that ionisation is complete and link this to a high concentration of H⁺(aq) ions.
- 💡Use a single arrow (→) in equations for strong acids and include state symbols (aq) to demonstrate aqueous ionisation clearly.
- 💡Use the word 'partially' and include the equilibrium arrow (⇌) whenever you describe or write the ionisation of a weak acid.
- 💡Link partial ionisation to a lower concentration of H⁺(aq) ions, a higher pH and slower reactions compared with a strong acid of the same concentration.
- 💡Always state that the two acids are being compared at the same concentration before drawing a conclusion about their pH.
- 💡If asked to explain the difference in pH, write a linked sequence: greater ionisation → higher concentration of H⁺ ions → lower pH.
- 💡Remember that calculating the exact factor change in hydrogen ion concentration is a Higher Tier only skill.
- 💡Write the factor as a power of 10, such as 10² for a two-unit change, to show the multiplicative relationship clearly.
- 💡Check the direction of the change carefully: decreasing pH means increasing H⁺ concentration.
- 💡Read the question carefully to see whether it asks about concentration or strength, and use the matching pair of terms rather than mixing them.
- 💡Support each explanation with a named acid, for example hydrochloric acid for a strong acid and ethanoic acid for a weak acid.
- 💡When explaining pH differences, refer to the concentration of H⁺ ions produced by complete or partial ionisation rather than to the total amount of acid.
- 💡When comparing two acidic solutions, state both the pH values and the relative H⁺ concentrations, for example 'pH 3 has a higher H⁺ concentration than pH 5'.
- 💡Use the words 'hydrogen ion concentration' rather than 'amount of acid' when explaining pH.
Common Mistakes
- Confusing strong with concentrated. Correction: strength describes the extent of ionisation, whereas concentration describes the amount of acid per unit volume.
- Writing ionisation equations with equilibrium arrows, such as HCl(aq) ⇌ H⁺(aq) + Cl⁻(aq). Correction: use a single arrow (→) to represent the complete ionisation of a strong acid.
- Forgetting that sulfuric acid is diprotic and writing H₂SO₄(aq) → H⁺(aq) + HSO₄⁻(aq). Correction: balance the equation to show both protons being released as H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq).
- Writing a single arrow for ethanoic acid, such as CH₃COOH(aq) → H⁺(aq) + CH₃COO⁻(aq). Correction: use the equilibrium arrow (⇌) to show partial, reversible ionisation.
- Thinking weak acids do not ionise at all. Correction: they do ionise, but only partially, so some H⁺(aq) ions are present and the solution is acidic.
- Assuming a weak acid must be dilute or harmless. Correction: strength depends on the extent of ionisation, while concentration and hazard are separate concepts.
- Thinking that a strong acid is simply more concentrated. Correction: strength describes the extent of ionisation, while concentration describes moles of acid per unit volume.
- Assuming a weak acid always has a high pH regardless of concentration. Correction: the comparison only holds for a given concentration; a concentrated weak acid can still have a fairly low pH.
- Stating that a strong acid has a higher pH because it is stronger. Correction: a higher concentration of H⁺ ions results in a lower pH value.
- Treating pH as a linear scale and stating H⁺ concentration increases by one unit; correction: the change is multiplicative, increasing by a factor of 10 for each pH unit drop.
- Reversing the direction and claiming H⁺ concentration falls as pH decreases; correction: a lower pH value indicates a higher concentration of H⁺ ions.
- For a two-unit pH drop, adding 10 + 10 to get a 20-fold increase; correction: multiply the factors, so 10 × 10 = 100-fold increase.
- Treating dilute and weak as synonyms, or concentrated and strong as synonyms. Correction: dilute and concentrated describe how much acid is dissolved per unit volume, while weak and strong describe how completely the acid ionises.
- Assuming that a weak acid produces no ions. Correction: a weak acid ionises partially, so it produces some H⁺ ions but also contains un-ionised molecules.
- Comparing the pH of a strong and a weak acid without controlling the concentration. Correction: pH comparisons are only valid when the concentrations of the two acids are the same, because pH depends on the H⁺ ion concentration.
- Thinking that a higher pH number means a more acidic solution. Correction: lower pH numbers correspond to higher H⁺ concentrations and greater acidity.
- Confusing the terms 'strong' and 'concentrated'. Correction: pH measures H⁺ concentration, not acid strength; a dilute strong acid can have a higher pH than a concentrated weak acid.
- Believing that neutral solutions contain no ions. Correction: neutral water contains H⁺ and OH⁻ ions in equal, very small concentrations.