A-Level · A-Level Chemistry
A-Level Chemistry titration
What is titration?
Titration is the quantitative technique that determines the concentration of an unknown solution by reacting it with a standard solution of known concentration until the end point is reached. A-Level Chemistry students must master the method (pipette, burette, conical flask, indicator), the mole calculation (n = C × V; mole ratio from balanced equation) and the sources of error.
Worked example
25.0 cm³ of 0.100 mol/dm³ sodium hydroxide is titrated with hydrochloric acid. 22.5 cm³ of acid is required to reach the end point. Find the concentration of the acid.
- Write the balanced equation: NaOH + HCl → NaCl + H₂O. Mole ratio is 1:1.
- Calculate moles of NaOH: n = C × V = 0.100 × (25.0/1000) = 0.00250 mol.
- Use the mole ratio: moles HCl = 0.00250 mol (1:1).
- Calculate concentration of HCl: C = n / V = 0.00250 / (22.5/1000) = 0.111 mol/dm³ (3 s.f.).
Apparatus and method
Standard A-Level apparatus:
- Pipette (calibrated 25 cm³ usually) for the analyte.
- Burette (calibrated 50 cm³) for the titrant.
- Conical flask (white tile underneath for end-point visibility).
- Indicator: phenolphthalein for strong base + weak acid; methyl orange for weak base + strong acid.
- Repeat titrations until two concordant results (within 0.10 cm³) are obtained.
End-point indicators
Indicator choice depends on the pH at equivalence. Phenolphthalein (colourless ↔ pink) changes around pH 8–10 — strong base + weak acid. Methyl orange (red ↔ yellow) changes around pH 3.1–4.4 — weak base + strong acid. Universal indicator is too gradual to use in titration.
Required-practical relevance
Acid–base titration is a required practical on every UK A-Level board (AQA 7405, Edexcel 9CH0, OCR H432). Written paper questions ask about apparatus precision, end-point detection, sources of error and concordant-result thresholds.
Common mistakes
- Not rinsing the burette with titrant before filling — leftover water dilutes the solution.
- Reading the burette at the top of the meniscus instead of the bottom.
- Using too much indicator — affects the end-point colour and equivalence point.
- Forgetting to convert cm³ to dm³ before applying n = C × V.
- Including non-concordant titres in the mean calculation.
Frequently asked
- Why do we repeat titrations?
- The first titration is a rough trial; subsequent titrations should produce concordant results (within 0.10 cm³ of each other). Concordant titres are averaged for the mean titre.
- What's the difference between equivalence point and end point?
- The equivalence point is the exact stoichiometric point where moles of acid = moles of base. The end point is when the indicator changes colour. A well-chosen indicator makes these nearly identical.
- Why can't universal indicator be used in titrations?
- Universal indicator changes colour gradually over a wide pH range, making it impossible to identify a sharp end point. A titration needs an indicator that changes colour sharply within ~0.5 pH units.
A-Level Chemistry glossary terms
- TitrationTitration is a quantitative analysis technique that determines the concentration of an unknown solution by reacting it with a standard solution of known concentration until a clear end point. For acid–base titrations, an indicator (phenolphthalein, methyl orange) changes colour at the end point; for redox titrations, a self-indicating species like potassium manganate(VII) does the job. Required practical on every UK board.
- Mole calculationsMole calculations convert between mass, moles and particles in chemistry. Core equations: n = m / Mr (moles from mass + relative formula mass), n = C × V (for solutions, with C in mol/dm³ and V in dm³), n = V / 24 (for gases at room temperature and pressure). Higher-tier GCSE Chemistry asks for mole ratios from balanced equations — read the equation as 'per mole of A, you need x moles of B'.
- Le Chatelier's principleLe Chatelier's principle states that when a dynamic equilibrium is disturbed, the system shifts to partially counteract the change. Increase the concentration of a reactant → equilibrium shifts toward products. Increase temperature → equilibrium shifts in the endothermic direction. Increase pressure → equilibrium shifts toward the side with fewer gas molecules. A-Level Chemistry uses it to predict yield changes in industrial processes (Haber, Contact).
- Required practicalsRequired practicals are specified science experiments that students must carry out to pass the practical-skills component of GCSE and A-Level Sciences. Exam boards (AQA, Edexcel, OCR) publish a list of around 8 GCSE and 12 A-Level required practicals per science. They are not directly marked, but written exams ask questions about apparatus, technique, hazards and analysis using the required practicals as context.
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