Strong and weak acids (HT only) — AQA GCSE Chemistry
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Strong and weak acids (HT only) explained
Strong and weak acids (HT only) covers 5 specification points, including A strong acid is completely ionised in aqueous… and A weak acid is only partially ionised in….
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A strong acid is completely ionised in aqueous solution. Examples of strong acids are hydrochloric, nitric and sulfuric acids.
In aqueous solution, a strong acid is fully dissociated into ions, so essentially every acid molecule releases a hydrogen ion, H⁺, to a water molecule. For hydrochloric acid the equation is HCl(aq) → H⁺(aq) + Cl⁻(aq), and the single arrow shows complete ionisation. Nitric acid ionises to H⁺ and NO₃⁻, and sulfuric acid ionises to release two H⁺ ions per molecule, forming SO₄²⁻. Because ionisation is complete, the concentration of H⁺ ions equals the concentration of the acid for monoprotic acids, giving a low pH. This contrasts with weak acids, which ionise only partially and are shown with an equilibrium arrow.
A weak acid is only partially ionised in aqueous solution. Examples of weak acids are ethanoic, citric and carbonic acids.
In aqueous solution, a weak acid releases hydrogen ions only partially: most of its molecules remain intact and only a small fraction ionise. This is shown by an equilibrium, for example CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq), where the position lies well to the left. Because few H⁺ ions form, a weak acid of a given concentration has a higher pH and reacts more slowly than a strong acid of the same concentration. Ethanoic acid (in vinegar), citric acid (in citrus fruit) and carbonic acid (formed when carbon dioxide dissolves in water) are common examples. The double arrow is essential: it shows that ionisation is reversible and incomplete, unlike the single arrow used for a strong acid such as hydrochloric acid.
For a given concentration of aqueous solutions, the stronger an acid, the lower the pH.
Acid strength depends on the degree of ionisation, not on how much acid is dissolved. If two aqueous solutions have the same concentration, the stronger acid ionises more completely, so it releases a greater concentration of hydrogen ions, H⁺. Since pH is a measure of H⁺ concentration, a higher H⁺ concentration gives a lower pH value. For example, 0.1 mol/dm³ hydrochloric acid (a strong acid) has a lower pH than 0.1 mol/dm³ ethanoic acid (a weak acid). The comparison is only valid at equal concentration: a very dilute strong acid can have a higher pH than a concentrated weak acid. Always state that concentration is fixed before comparing strength.
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: each decrease of one pH unit represents a tenfold increase in hydrogen ion concentration, [H⁺]. For example, a solution of pH 3 has ten times the H⁺ concentration of a solution of pH 4, and one hundred times that of a solution of pH 5. Conversely, increasing pH by one unit divides the H⁺ concentration by 10. This relationship explains why small pH changes matter greatly in reactions and in living systems. To use it, compare pH values, find the difference, and raise 10 to that power: a fall of 2 pH units means 10² = 100 times the H⁺ concentration. The rule applies to whole-number pH values at this level.
Students should be able to: 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 describe neutrality and relative acidity in terms of the effect on hydrogen ion concentration and the numerical value of pH (whole numbers only). AQA GCSE Chemistry 8462. GCSE exams Jun
This skill combines two pairs of terms. Dilute and concentrated describe how much acid is dissolved per unit volume: a dilute solution has a small amount of substance in a given volume, a concentrated solution has a large amount. Weak and strong describe the degree of ionisation: a strong acid ionises completely, a weak acid only partially. The pairs are independent, so a weak acid can be concentrated and a strong acid can be dilute. Neutrality means [H⁺] equals [OH⁻], giving pH 7 for aqueous solutions at the usual laboratory temperature. Relative acidity depends on hydrogen ion concentration: the greater the [H⁺], the lower the numerical pH. At this level, pH values are whole numbers, so you compare and order solutions using values such as pH 1, 4 and 7.
Your focus
- Define a strong acid in terms of complete ionisation.
- Write ionisation equations for hydrochloric, nitric and sulfuric acids.
- Explain how complete ionisation affects the concentration of H⁺ ions and the pH.
Show all 15 objectives
- State that a weak acid is only partially ionised in aqueous solution.
- Represent the ionisation of a weak acid using a reversible equation.
- Name ethanoic, citric and carbonic acids as examples of weak acids.
- Explain why a stronger acid has a lower pH at a given concentration.
- Relate pH to hydrogen ion concentration.
- Distinguish between acid strength and acid concentration when comparing pH values.
- State that a one-unit fall in pH increases H⁺ concentration tenfold.
- Calculate the factor change in H⁺ concentration between two whole-number pH values.
- Interpret pH changes in terms of powers of 10.
- Use and explain dilute, concentrated, weak and strong correctly in relation to acids.
- Describe neutrality and relative acidity using hydrogen ion concentration and whole-number pH values.
- Distinguish clearly between concentration and strength when comparing acid solutions.
Strong and weak acids (HT only) exam tips
Marking Points
- States that a strong acid is completely ionised in aqueous solution.
- Writes ionisation equations using a single arrow to show complete dissociation.
- Names hydrochloric, nitric and sulfuric acids as examples of strong acids.
- Identifies the ions formed, including H⁺ and the appropriate negative ion.
- Links complete ionisation to a high concentration of H⁺ ions and a low pH.
- A weak acid is only partially ionised in aqueous solution, so only a small proportion of its molecules release hydrogen ions.
- The ionisation of a weak acid is represented as an equilibrium using a reversible arrow, for example CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq).
- Because ionisation is incomplete, the concentration of H⁺ ions is lower than for a strong acid of the same concentration, giving a higher pH.
- Ethanoic acid, citric acid and carbonic acid are named examples of weak acids.
- The contrast with strong acids is the degree of ionisation: strong acids ionise completely, weak acids do not.
- A weak acid can still be concentrated; weak refers to degree of ionisation, not to amount of substance dissolved.
- At the same concentration, a stronger acid ionises more completely and so produces a higher concentration of H⁺ ions.
- A higher concentration of H⁺ ions corresponds to a lower pH value.
- Therefore, for equal concentrations, the stronger acid has the lower pH.
- The comparison must specify a given concentration, because changing concentration also changes pH.
- Strong acids such as hydrochloric acid ionise completely, whereas weak acids such as ethanoic acid ionise only partially.
- pH is a measure of hydrogen ion concentration, so it links directly to the degree of ionisation.
- A decrease of one pH unit corresponds to a tenfold increase in hydrogen ion concentration.
- A decrease of two pH units corresponds to a hundredfold increase, because 10² = 100.
- An increase of one pH unit corresponds to dividing the hydrogen ion concentration by 10.
- The relationship is logarithmic, so equal pH steps represent equal multiplicative changes in [H⁺].
- To compare two solutions, subtract the pH values and use the difference as the power of 10.
- The rule links the numerical pH scale directly to the concentration of H⁺ ions.
- Dilute and concentrated refer to the amount of substance dissolved per unit volume of solution.
- Weak and strong refer to the degree of ionisation of the acid in aqueous solution.
- The two descriptions are independent: a weak acid can be concentrated, and a strong acid can be dilute.
- Neutrality corresponds to equal concentrations of H⁺ and OH⁻ ions, giving pH 7 in aqueous solution at the usual laboratory temperature.
- Relative acidity increases as hydrogen ion concentration increases, which corresponds to a lower numerical pH.
- At this level, pH values are compared as whole numbers, for example pH 2 is more acidic than pH 5.
Examiner Tips
- 💡Use the words completely ionised or fully dissociated when defining a strong acid.
- 💡Write state symbols in ionisation equations to show the ions are in aqueous solution.
- 💡Compare strong and weak acids by referring to the extent of ionisation, not to pH alone.
- 💡When asked to explain why a weak acid has a higher pH than a strong acid of the same concentration, link partial ionisation to a lower H⁺ concentration.
- 💡Quote at least one named weak acid from the specification, such as ethanoic acid, to support your answer.
- 💡If you write an equation for a weak acid, use the reversible arrow and include state symbols to show it is in aqueous solution.
- 💡Begin comparison answers with the phrase 'for the same concentration' to show you understand the condition.
- 💡Link the lower pH to a higher H⁺ concentration, not simply to the label 'strong acid'.
- 💡Use a named pair such as hydrochloric acid and ethanoic acid to make the comparison concrete.
- 💡Write the pH difference as a power of 10 before calculating, for example 10² = 100.
- 💡Check the direction: lower pH means more H⁺ ions, so the factor should be greater than 1.
- 💡Show your working clearly when comparing two pH values so the examiner can follow your reasoning.
- 💡Define each term separately and then state clearly that concentration and strength are independent.
- 💡When describing neutrality, refer to equal H⁺ and OH⁻ concentrations and give pH 7 for aqueous solutions at the usual laboratory temperature.
- 💡Use whole-number pH values when ordering solutions by acidity, and justify the order using H⁺ concentration.
Common Mistakes
- Describing strong acids as concentrated; correct by explaining that strength refers to complete ionisation, while concentration refers to amount of acid per unit volume.
- Using an equilibrium arrow for a strong acid; correct by using a single arrow to show complete ionisation.
- Forgetting that sulfuric acid releases two H⁺ ions per molecule; correct by writing the equation with two H⁺ ions and one SO₄²⁻ ion.
- Confusing weak with dilute: a weak acid is partially ionised, whereas a dilute acid simply has a small amount of acid per unit volume; a weak acid can be concentrated.
- Writing a single arrow for the ionisation of a weak acid; the correct notation uses a reversible arrow to show incomplete ionisation.
- Believing that a weak acid produces no ions at all; it does produce H⁺ ions, but only a small fraction of its molecules ionise.
- Comparing a strong acid and a weak acid without fixing the concentration; the statement only holds for a given concentration.
- Thinking that a stronger acid always has a lower pH regardless of concentration; a dilute strong acid may have a higher pH than a concentrated weak acid.
- Confusing acid strength with concentration; strength is about degree of ionisation, concentration is about amount of substance per unit volume.
- Treating the change as additive, for example saying a fall of 2 pH units doubles the H⁺ concentration; the correct change is a factor of 10² = 100.
- Reversing the direction, for example saying a pH decrease lowers the H⁺ concentration; a lower pH means a higher H⁺ concentration.
- Applying the factor of 10 to the pH value itself rather than to the hydrogen ion concentration.
- Using dilute and weak as if they mean the same thing; dilute concerns amount of substance, weak concerns degree of ionisation.
- Assuming all acids have pH values below 7 and all alkalis above 7 without linking this to H⁺ and OH⁻ concentrations; neutrality needs equal concentrations of both ions.
- Treating pH as a linear measure of acidity; a change of one pH unit represents a tenfold change in H⁺ concentration.