The pH Scale & Neutralisation
What makes a solution acidic or alkaline comes down to two ions:
- Acids release hydrogen ions, H+, in aqueous solution — the more H+, the more acidic.
- Alkalis release hydroxide ions, OH−, in aqueous solution — the more OH−, the more alkaline.
The pH scale runs from 0 to 14 and measures how acidic or alkaline a solution is. pH 7 is neutral; below 7 is acidic; above 7 is alkaline. You can measure pH with universal indicator (which changes colour) or, more precisely, with a pH probe.
The pH scale. Universal indicator turns red in strong acids, green at neutral, and purple in strong alkalis.
When an acid neutralises an alkali, the hydrogen ions from the acid react with the hydroxide ions from the alkali to make water. This is true for every acid–alkali neutralisation:
H+(aq) + OH−(aq) → H2O(l)
The leftover metal and non-metal ions stay in solution as the dissolved salt — which you can recover by crystallisation.
🧪 Exam-style questions
A solution turns universal indicator purple. Which row is correct?
Titrations Triple Only
A titration finds the exact volumes of acid and alkali that react together. This topic is Triple (Chemistry only). The method is for all triple students; the calculations are Higher Tier only.
- Use a pipette to measure a fixed volume of alkali into a conical flask, and add a few drops of a suitable indicator. Stand the flask on a white tile so the colour change shows up clearly.
- Fill a burette with the acid and record the starting reading (from the bottom of the meniscus — see below).
- Do a rough (trial) titration first: run the acid in fairly quickly, swirling, until the indicator just changes colour. This gives an approximate idea of how much acid you will need.
- Repeat carefully: run the acid in quickly at first, then dropwise as you approach the end-point, until the indicator just changes colour. Record the final reading.
- Repeat until you get concordant results (within 0.10 cm3) and take a mean (ignoring the rough titre).
Use a single-colour-change indicator such as phenolphthalein (pink→colourless) or methyl orange (yellow→red) — not universal indicator, whose gradual rainbow of colours gives no sharp end-point.
A pipette measures one fixed volume (e.g. 25.0 cm3) — it delivers the same amount every time, which is why it is used for the alkali. A burette measures variable volumes — you read off exactly how much liquid has run out, which is why it is used for the acid whose volume you are trying to find (the titre).
The required practical in two stages: first pipette a fixed volume of alkali into the flask. Add a few drops of indicator. Then run acid in from the burette drop by drop until the colour just changes.
Reading the burette
The surface of the liquid in a burette (or pipette) curves into a meniscus. Always take the reading from the bottom of the meniscus, with your eye level with it to avoid a parallax error. Burette readings are taken to the nearest 0.05 cm3.
The meniscus curves up at the glass walls and dips in the middle. Line your eye up with the lowest point and read from there — reading the higher edges gives a smaller, wrong value.
Titration calculations Higher
Turning your mean titre into a concentration is a calculation. In short: find the moles of the solution you know (moles = concentration × volume in dm3), use the balanced equation to get the moles of the other solution, then divide by its volume to find its concentration in mol/dm3. To express that concentration in g/dm3, multiply by the Mr — Higher Tier titration questions can ask for either unit.
🧪 Exam-style questions
A student titrates a sodium hydroxide solution against dilute hydrochloric acid. This is part of the method the student used.
2. Add a few drops of indicator to the flask.
3. Fill a burette with the dilute hydrochloric acid.
Describe how the student would complete the titration.
Show answer
- Add the hydrochloric acid to the flask until there is a (permanent) colour change. 1 mark
- Measure / record the volume of acid added — i.e. take the final (and initial) burette reading. 1 mark
- Plus any one good-technique point: swirl the flask · stand it on a white tile · add the acid dropwise (slowly) near the end-point · repeat and calculate a mean. 1 mark
The examiner ignores any mention of which colours the indicator turns — just that there is a colour change at the end-point.
Give two reasons why a burette is used for the hydrochloric acid.
Show answer
Any two from:
- It can add the acid in small increments (drop by drop / slowly). 1 mark
- It can measure variable volumes (it has a scale). 1 mark
- It is more accurate than a measuring cylinder. 1 mark
Any two of these reasons score the 2 marks — both are about reading off the exact titre precisely, which a pipette (one fixed volume) or measuring cylinder cannot do.
25.0 cm3 of 0.100 mol/dm3 sodium hydroxide solution is exactly neutralised by 20.0 cm3 of dilute hydrochloric acid. The equation is NaOH + HCl → NaCl + H2O. Calculate the concentration of the hydrochloric acid in mol/dm3.
Show answer
- Moles of NaOH = concentration × volume (in dm3) = 0.100 × (25.0 ÷ 1000) = 0.00250 mol. 1 mark
- From the equation the ratio NaOH : HCl is 1 : 1, so moles of HCl = 0.00250 mol. 1 mark
- Concentration of HCl = moles ÷ volume (in dm3) = 0.00250 ÷ (20.0 ÷ 1000) = 0.125 mol/dm3. 1 mark
Allow the correct answer following an arithmetic slip earlier (error carried forward). The full method — including expressing the answer in g/dm3 — is in C3: Quantitative Chemistry.
Strong & Weak Acids Higher
This whole section is Higher Tier only. The key is to separate two ideas that sound similar but mean different things: strong vs weak, and concentrated vs dilute.
A strong acid is completely ionised in aqueous solution — every molecule splits up to release its H+ ions. Examples: hydrochloric, nitric and sulfuric acids.
A weak acid is only partially ionised — just a small fraction of the molecules release their H+ ions at any moment. Examples: ethanoic, citric and carbonic acids.
Strong / weak is about the degree of ionisation — what fraction of molecules split up. Concentrated / dilute is about the amount of acid dissolved in a given volume (the number of moles per dm3). You can have a dilute strong acid or a concentrated weak acid — they are independent ideas.
For solutions of the same concentration, a stronger acid releases more H+ ions, so it has a lower pH.
The pH scale is logarithmic: as the pH decreases by one unit, the hydrogen-ion concentration increases by a factor of 10.
So an acid at pH 2 has 10× the H+ concentration of one at pH 3, and 100× that of one at pH 4.
Because the scale is logarithmic, each whole pH unit is a factor of 10, not an equal step. For a change of two pH units the H+ concentration changes by 10 × 10 = 100× — so a fall of 2 pH units means you multiply H+ by 100, not divide by 100 (and never by 2). Decide the direction first: lower pH = more H+.
🧪 Strong & weak acids — ionisation, concentration & pH
Both acids start at the same concentration. Notice the strong acid is far more acidic. Then drag its slider to dilute it — and see how far you have to go before it only matches the weak acid’s pH.
● H+ hydrogen ion · ● A− acid anion (Cl− or CH3COO−) · ● H–A un-ionised acid molecule
≈100% ionised · picture schematic
· picture schematic
pH scale — compare the two acids
🧪 Exam-style questions
Which statement describes a weak acid?
Acid A has pH 1 and acid B has pH 4. How does the H+ concentration of A compare with B?