The pH scale and neutralisation — AQA GCSE Combined Science
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The pH scale and neutralisation explained
Acids produce hydrogen ions, H⁺, when dissolved in water.
Read the full explanation
In aqueous solution, an acid releases H⁺ ions, which are responsible for acidic properties such as low pH, reaction with metals to form hydrogen, and neutralisation by bases. For example, hydrochloric acid in water forms H⁺ and Cl⁻ ions; the H⁺ ions make the solution acidic. The concentration of H⁺ ions determines pH: higher H⁺ concentration gives lower pH. Strong acids ionise completely, while weak acids ionise partially, but both produce H⁺ in aqueous solution. This idea explains why acids only show acidic behaviour when water is present.
Aqueous solutions of alkalis contain hydroxide ions (OH–).
Alkalis are soluble bases, and when dissolved in water they produce hydroxide ions, OH⁻. This ion is what makes a solution alkaline: the greater the concentration of OH⁻, the higher the pH. For example, sodium hydroxide dissolves to give Na⁺ and OH⁻ ions, so the solution turns universal indicator blue or purple. The word 'aqueous' is essential because it means dissolved in water; solid alkalis such as sodium hydroxide pellets do not contain free hydroxide ions in the same way. Understanding this ion explains why alkalis neutralise acids: OH⁻ ions react with H⁺ ions to form water, H⁺ + OH⁻ → H₂O. This links the pH scale to ionic theory and to neutralisation reactions.
The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.
The pH scale runs from 0 to 14 and tells you how acidic or alkaline a solution is. A pH below 7 is acidic, pH 7 is neutral, and pH above 7 is alkaline. Universal indicator is a mixture of dyes that gives a colour for each pH value, such as red for strong acids and purple for strong alkalis. A pH probe gives a numerical reading and is more precise, so it is better when you need an exact value. For example, a solution with pH 2 is acidic, while one with pH 12 is alkaline. The scale is logarithmic, so each change of one pH unit represents a tenfold change in hydrogen ion concentration, though at GCSE you mainly need to read and interpret the scale.
A solution with pH 7 is neutral. Aqueous solutions of acids have pH values of less than 7 and aqueous solutions of alkalis have pH values greater than 7.
The pH scale runs from 0 to 14 and measures how acidic or alkaline an aqueous solution is. A solution with pH 7 is neutral, for example pure water at room temperature. Aqueous solutions of acids have pH values below 7: a strong acid such as hydrochloric acid may have pH 1, while a weak acid such as ethanoic acid may have pH 3. Aqueous solutions of alkalis have pH values above 7: sodium hydroxide solution may have pH 13, while ammonia solution may have pH 10. The further the pH is from 7, the more acidic or alkaline the solution. Universal indicator or a pH probe can measure pH, and the colour or reading is compared with a scale.
In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.
When an acid reacts with an alkali, the reaction is called neutralisation. In aqueous solution, the acid provides hydrogen ions, H⁺, and the alkali provides hydroxide ions, OH⁻. These ions react together to produce water: H⁺ + OH⁻ → H₂O. For example, hydrochloric acid reacts with sodium hydroxide solution to form sodium chloride and water: HCl + NaOH → NaCl + H₂O. The hydrogen ions and hydroxide ions are the reacting species, while the sodium ions and chloride ions are spectator ions that remain in solution. Because the hydrogen ions and hydroxide ions are used up, the pH moves towards 7. The same ionic equation applies whenever an acid is neutralised by an alkali.
This reaction can be represented by the equation:
In neutralisation, an acid reacts with an alkali or base to form a salt and water. The general word equation is: acid + alkali → salt + water. For example, hydrochloric acid reacting with sodium hydroxide solution gives sodium chloride and water: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). The H⁺ ions from the acid combine with the OH⁻ ions from the alkali to form water: H⁺(aq) + OH⁻(aq) → H₂O(l). The remaining ions form the dissolved salt. When writing the equation, you must balance atoms and include state symbols where required. The equation shows the reaction quantitatively, so the number of each type of atom is the same on both sides. This links pH, neutralisation and ionic equations.
describe the use of universal indicator or a wide range indicator to measure the approximate pH of a solution
Universal indicator is a mixture of dyes that changes colour gradually across the pH scale, so it shows whether a solution is strongly acidic, weakly acidic, neutral, weakly alkaline or strongly alkaline. To use it, add a few drops to the solution in a clean test tube or spotting tile, or dip a clean glass rod into the solution and touch it onto universal indicator paper. Compare the resulting colour with the colour chart and read off the nearest pH value. The reading is approximate because colour matching is subjective and the indicator gives a broad colour change rather than a sharp end point. A wide range indicator works in the same way, covering a broad pH interval.
use the pH scale to identify acidic or alkaline solutions.
The pH scale runs from 0 to 14 and shows how acidic or alkaline a solution is. A pH below 7 means the solution is acidic; the lower the pH, the more acidic the solution. A pH of exactly 7 means the solution is neutral, as in pure water. A pH above 7 means the solution is alkaline; the higher the pH, the more alkaline the solution. For example, a solution of pH 2 is acidic, a solution of pH 7 is neutral, and a solution of pH 12 is alkaline. The pH scale is related to the concentration of hydrogen ions, H⁺: acidic solutions have a higher concentration of H⁺ than neutral or alkaline solutions, while alkaline solutions have a lower concentration of H⁺.
Your focus
- State that acids produce hydrogen ions, H⁺, in aqueous solutions.
- Explain how the concentration of H⁺ ions relates to pH.
- Distinguish between strong and weak acids in terms of ionisation in water.
Show all 24 objectives
- Describe that aqueous solutions of alkalis contain hydroxide ions, OH⁻.
- Relate the presence and concentration of OH⁻ ions to alkaline pH values.
- Use the reaction between H⁺ and OH⁻ ions to explain neutralisation.
- State the range of the pH scale and what it measures.
- Describe how universal indicator and a pH probe can be used to measure pH.
- Classify solutions as acidic, neutral or alkaline from their pH values.
- Identify a solution as acidic, neutral or alkaline from its pH value.
- Place common acidic, neutral and alkaline solutions on the pH scale.
- Describe a practical method for measuring the pH of a solution.
- Write the ionic equation for the reaction between hydrogen ions and hydroxide ions.
- Explain how neutralisation changes the pH of a solution.
- Identify spectator ions in a neutralisation reaction between an acid and an alkali.
- Write a word equation for the reaction between an acid and an alkali to produce a salt and water.
- Construct a balanced symbol equation, including state symbols, for a neutralisation reaction.
- Deduce the ionic equation for neutralisation and explain the role of H⁺ and OH⁻ ions.
- Describe how universal indicator or a wide range indicator is used to test a solution.
- Interpret the colour produced by universal indicator using a colour chart to estimate pH.
- Explain why the pH value obtained with universal indicator is approximate rather than exact.
- Use pH values to classify solutions as acidic, neutral or alkaline.
- Compare solutions using pH values to decide which is more acidic or more alkaline.
- Relate pH to the concentration of hydrogen ions, H⁺, in a solution.
The pH scale and neutralisation exam tips
Marking Points
- Acids produce hydrogen ions, H⁺, when dissolved in water.
- The H⁺ ion is responsible for acidic properties, including low pH and neutralisation.
- In aqueous solution, an acid ionises or dissociates to release H⁺ ions.
- The concentration of H⁺ ions determines the pH of the solution.
- Strong acids ionise completely in water, while weak acids ionise partially, but both produce H⁺ ions.
- States that alkalis are soluble bases that produce hydroxide ions, OH⁻, when dissolved in water.
- Identifies OH⁻ as the ion responsible for alkaline properties and for pH values above 7.
- Explains that the concentration of OH⁻ ions determines how strongly alkaline a solution is.
- Uses the example of sodium hydroxide dissolving to form Na⁺ and OH⁻ ions.
- Links hydroxide ions to neutralisation: H⁺ + OH⁻ → H₂O.
- Distinguishes aqueous alkali solutions from solid bases, which do not contain free hydroxide ions.
- States that the pH scale runs from 0 to 14 and measures acidity or alkalinity.
- Identifies pH values below 7 as acidic, pH 7 as neutral and pH values above 7 as alkaline.
- Describes universal indicator as a indicator that gives a colour related to pH.
- Describes a pH probe as an instrument that gives a numerical pH reading and is more precise.
- Compares the two methods, noting that a pH probe gives a more accurate value than colour matching.
- Interprets given pH values correctly as acidic, neutral or alkaline.
- State that pH 7 is neutral, using pure water as the standard example.
- State that acids in aqueous solution have pH values below 7, and link lower values to greater acidity.
- State that alkalis in aqueous solution have pH values above 7, and link higher values to greater alkalinity.
- Use the pH scale from 0 to 14 and place given examples correctly, such as pH 1 for a strong acid and pH 13 for a strong alkali.
- Describe how universal indicator or a pH probe is used to measure pH and how the result is interpreted.
- Identify the hydrogen ion, H⁺, as the ion supplied by the acid in aqueous solution.
- Identify the hydroxide ion, OH⁻, as the ion supplied by the alkali in aqueous solution.
- Write the ionic equation H⁺ + OH⁻ → H₂O and state that water is the product.
- Explain that the hydrogen ions and hydroxide ions are used up, so the solution becomes less acidic and the pH moves towards 7.
- Apply the idea to an example such as HCl + NaOH → NaCl + H₂O, recognising that sodium ions and chloride ions are spectator ions.
- Write the correct word equation: acid + alkali → salt + water, naming the salt from the acid and the metal in the alkali.
- Write the balanced symbol equation, for example HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l), ensuring all atoms balance.
- Include state symbols correctly: (aq) for dissolved acid, alkali and salt, and (l) for water.
- Identify the ionic equation for neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l), showing that spectator ions are unchanged.
- Explain that the reaction occurs because H⁺ ions from the acid react with OH⁻ ions from the alkali to form water.
- Universal indicator is a mixture of indicators that produces a different colour at different pH values across the pH scale.
- The indicator is added to the solution, or the solution is applied to indicator paper, and the colour produced is compared with a standard colour chart.
- The colour chart links each colour to an approximate pH value, allowing the solution to be classified as acidic, neutral or alkaline.
- The result is approximate because colour matching is subjective and the indicator changes colour gradually over a range of pH values.
- A wide range indicator covers a broad pH interval and is used in the same way as universal indicator to estimate pH.
- The solution should be tested in a clean container with a clean dropping pipette or glass rod to avoid contamination by residues from previous tests.
- The pH scale runs from 0 to 14 and is used to classify solutions as acidic, neutral or alkaline.
- A pH value below 7 indicates an acidic solution, and a lower pH value indicates a more acidic solution.
- A pH value of exactly 7 indicates a neutral solution, such as pure water.
- A pH value above 7 indicates an alkaline solution, and a higher pH value indicates a more alkaline solution.
- Acidic solutions have a higher concentration of hydrogen ions, H⁺, than neutral or alkaline solutions.
- The pH scale can be used to compare the relative acidity or alkalinity of two solutions from their pH values.
Examiner Tips
- 💡Define an acid as a source of H⁺ ions in aqueous solution, and link this to acidic properties.
- 💡Use the formula H⁺ correctly in equations and explanations, including state symbols where appropriate.
- 💡When comparing strong and weak acids, refer to the extent of ionisation and the concentration of H⁺ ions.
- 💡Underline the word 'aqueous' in the question and make sure your answer refers to ions in solution.
- 💡When asked why a solution is alkaline, name OH⁻ and link its concentration to pH rather than just saying 'it is an alkali'.
- 💡Use the ionic equation H⁺ + OH⁻ → H₂O to show how hydroxide ions cause neutralisation.
- 💡Learn the colour range of universal indicator, from red at low pH through green at pH 7 to purple at high pH.
- 💡When comparing methods, mention that a pH probe gives a numerical value and avoids subjective colour judgement.
- 💡Check whether a question asks for a description or a comparison, and include both methods if it asks you to compare.
- 💡Quote pH values with the correct inequality, for example pH < 7 for acids and pH > 7 for alkalis.
- 💡When asked to compare solutions, give the pH value and state whether it is acidic, neutral or alkaline.
- 💡If describing a test, name the indicator or probe and state the colour or reading that shows acid, neutral or alkali.
- 💡Write the ionic equation with correct charges and state symbols, for example H⁺(aq) + OH⁻(aq) → H₂O(l).
- 💡If asked to explain neutralisation, name both reacting ions and the product rather than only naming the salt.
- 💡Use the ionic equation to explain pH changes, linking the removal of H⁺ and OH⁻ to the solution becoming neutral.
- 💡Read the question carefully to identify the acid and alkali before writing the equation.
- 💡Check that the equation is balanced by counting each type of atom on both sides.
- 💡If asked for an ionic equation, cancel spectator ions and write only the reacting ions and the product.
- 💡State clearly that the colour is matched to a chart to obtain an approximate pH value, not just that the indicator changes colour.
- 💡Use the colour chart language of the specification, such as red for strong acid and purple for strong alkali, and link each colour to an approximate pH.
- 💡When asked to describe a method, include the practical steps: place the solution in a clean container, add a few drops of universal indicator, and compare the colour with the chart.
- 💡Quote the pH value and then state the classification, for example pH 3 is acidic and pH 10 is alkaline.
- 💡When comparing two acids, refer to the lower pH value as the more acidic solution rather than saying it has a higher acidity number.
- 💡Link pH to the concentration of H⁺ ions to show understanding beyond simple colour changes.
Common Mistakes
- Writing H⁺ as H+ or H plus: the charge must be shown as a superscript, H⁺. Correction: use the correct ion symbol H⁺.
- Thinking acids produce H⁺ ions in any solvent: the statement specifies aqueous solutions. Correction: state that water is required for H⁺ production.
- Confusing strong and concentrated: strong refers to complete ionisation, concentrated refers to many moles per unit volume. Correction: use the terms precisely.
- Writing the hydroxide ion as OH+ or OH- without the correct Unicode superscript sign; the correct ion is OH⁻.
- Confusing alkalis with bases generally; all alkalis are bases, but not all bases are alkalis because only soluble bases form OH⁻ in water.
- Thinking that the OH⁻ ion is present in solid sodium hydroxide in the same free form as in solution; the ions are released when the solid dissolves.
- Saying that pH 7 is always neutral for every liquid; pH 7 is neutral in aqueous solutions at standard conditions, but the statement applies to solutions.
- Confusing universal indicator with litmus; litmus only shows acidic or alkaline, while universal indicator gives a range of colours across the pH scale.
- Reading the pH scale backwards, for example calling pH 12 acidic; lower numbers are acidic and higher numbers are alkaline.
- Saying that all acids have pH 7 or that neutral means pH 0; correct this by stating that neutral is pH 7 and acids are below 7.
- Confusing strong and weak with concentrated and dilute; correct this by explaining that strong or weak describes the degree of ionisation, while concentrated or dilute describes the amount of substance dissolved.
- Believing that alkalis are always harmless or that all bases dissolve in water; correct this by stating that alkalis are soluble bases and can still be corrosive.
- Writing H⁺ + OH⁻ → H₂O₂; correct this by balancing the equation so that one hydrogen ion and one hydroxide ion produce one water molecule, H₂O.
- Thinking that neutralisation always produces a pH of exactly 7; correct this by stating that the pH moves towards 7 and reaches 7 only when the acid and alkali are in the correct reacting amounts.
- Treating the salt as the only product; correct this by naming both water and the salt, and by identifying water as the product of H⁺ and OH⁻.
- Writing the salt as sodium chlorate or sodium hydroxide instead of sodium chloride; correct by naming the salt from the acid's negative ion and the metal ion from the alkali.
- Omitting state symbols or using (s) for the salt when it is dissolved; correct by using (aq) for solutions and (l) for liquid water.
- Forgetting to balance the equation, for example writing H₂SO₄ + NaOH → Na₂SO₄ + H₂O without checking atoms; correct by counting atoms on each side and adjusting coefficients to H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.
- Treating the universal indicator reading as exact: the colour chart gives an approximate pH, so a reading such as pH 3 should be described as approximate rather than precise.
- Confusing universal indicator with a single indicator such as litmus: litmus only shows acidic or alkaline, whereas universal indicator gives a range of colours linked to pH values.
- Using a dirty test tube or pipette: residues from a previous solution can change the colour and give a false pH reading, so apparatus must be clean.
- Thinking that a higher pH number means a stronger acid: the opposite is true, because a lower pH means a more acidic solution.
- Treating pH 7 as acidic or alkaline: pH 7 is neutral, and only values below 7 are acidic while values above 7 are alkaline.
- Assuming the pH scale ends at 14 in all contexts: the scale commonly runs from 0 to 14, but values outside this range are possible for very concentrated solutions.