Covers Lessons 13–19: Buffers, Titration (standard solutions, technique, calculations), Indicators, Titration Curves, Titration & Indicator Mastery, Back Titration, and Conductometric Titration.
Lesson Summaries, Quick Review
A buffer resists pH change upon addition of small amounts of acid or base. Composed of a weak acid and its conjugate base (or weak base and conjugate acid) in comparable concentrations. Mechanism: added H⁺ is consumed by A⁻; added OH⁻ is consumed by HA. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Blood pH (7.35–7.45) is maintained by the carbonate buffer system.
A standard solution has an accurately known concentration. Primary standards (e.g. Na₂CO₃, oxalic acid) are used to standardise solutions. Titration technique: wash burette with titrant, rinse conical flask with distilled water (not analyte), add indicator, titrate to persistent colour change. Calculate using n = cV; for acid–base: n(acid) × stoichiometric ratio = n(base).
Indicators are weak acids (HIn) where the acid form and base form (In⁻) have different colours. They change colour at their pKa. Select an indicator whose colour change range overlaps the equivalence point. Strong acid + strong base: any indicator works. Weak acid + strong base: alkaline indicator (phenolphthalein). Strong acid + weak base: acidic indicator (methyl orange).
Four curve types: SA+SB (pH jump 4–10, steep), WA+SB (starts higher, equivalence ~pH 8–9), SA+WB (equivalence ~pH 4–6), WA+WB (gradual, no sharp jump). The shape reveals: initial pH, buffer region gradient, half-equivalence point (pH = pKa for WA+SB), equivalence point pH, and the steepness of the jump.
Consolidation: selecting correct indicators for given titration types, calculating unknown concentrations from titration data, interpreting and sketching titration curves, identifying errors in titration technique. Band 6 responses explain indicator selection in terms of indicator pKa and equivalence point pH, not just memorised rules.
Back titration: add excess known reagent, titrate the unreacted excess. Used when the analyte dissolves slowly (e.g. CaCO₃ in antacid tablets) or decomposes. Conductometric titration: monitor conductance instead of colour, no indicator needed. Conductance drops to minimum at equivalence point then rises; useful for coloured or turbid solutions.
Industrial acid/base analysis uses automated titrators, pH probes, and potentiometric endpoints. Digital titration eliminates subjectivity in endpoint detection. Applications: quality control in food/pharmaceutical manufacture, waste water pH management, soil analysis. Data logging allows derivative plots (dpH/dV) to precisely locate the equivalence point.
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Multiple Choice, 20 Questions (1 mark each)
A buffer solution contains CH₃COOH and CH₃COONa. When HCl is added to this buffer, the primary reaction that resists pH change is:
25.00 mL of NaOH solution is titrated against 0.100 mol L⁻¹ HCl. The average titre is 22.50 mL. What is the concentration of the NaOH solution?
Which indicator is MOST appropriate for the titration of a weak acid with a strong base?
In a titration curve for a weak acid titrated with strong base, the pH at the half-equivalence point equals:
A student dissolves an antacid tablet (containing CaCO₃) in 50.00 mL of 1.00 mol L⁻¹ HCl (excess). The unreacted HCl is then titrated with 0.500 mol L⁻¹ NaOH; the titre is 28.00 mL. How many moles of HCl reacted with the CaCO₃?
A buffer contains equal concentrations of a weak acid and its conjugate base. Which statement is correct?
Which buffer would have the greatest capacity to resist pH change, assuming the ratio [A⁻]/[HA] is close to 1 in each case?
Which property is essential for a substance to be used as a primary standard?
Before filling a burette with NaOH solution, the burette should be rinsed with:
An indicator changes colour because:
Which titration type typically has an equivalence point below pH 7?
At the half-equivalence point in a weak acid-strong base titration, the solution contains:
In a phenolphthalein titration, the endpoint is identified when the solution:
The main purpose of obtaining concordant titres is to:
Which sequence best describes a back titration?
Why is conductometric titration useful for coloured solutions?
During the conductometric titration of HCl with NaOH, conductance typically:
A potentiometric endpoint in acid-base analysis is detected by measuring:
Why are derivative plots (dpH/dV against volume) useful in automated titration?
Which is the best reason industrial laboratories often use automated titrators?
Short Answer, 3 Questions
5 marksA buffer is prepared by mixing 0.15 mol of CH₃COOH and 0.10 mol of CH₃COONa in 1.00 L of solution (Ka = 1.8 × 10⁻⁵). (a) Calculate the pH of this buffer using the Henderson-Hasselbalch equation. (b) Predict and explain what happens to the pH when 0.01 mol of NaOH is added.
4 marksA student titrates 20.00 mL of ethanoic acid with 0.100 mol L⁻¹ NaOH. The average concordant titre is 18.50 mL. (a) Calculate the concentration of the ethanoic acid. (b) State and justify which indicator the student should use.
4 marksSketch and annotate a titration curve for the titration of 25 mL of 0.1 mol L⁻¹ CH₃COOH with 0.1 mol L⁻¹ NaOH. Label: initial pH, half-equivalence point, equivalence point pH, and suitable indicator. Explain why phenolphthalein is preferred over methyl orange for this titration.
The conjugate base (CH₃COO⁻) consumes the added H⁺: CH₃COO⁻ + H⁺ → CH₃COOH. This reaction removes the added acid without significantly changing the ratio [A⁻]/[HA], so pH changes very little.
n(HCl) = 0.02250 L × 0.100 mol L⁻¹ = 2.25 × 10⁻³ mol. Since HCl + NaOH → NaCl + H₂O (1:1), n(NaOH) = 2.25 × 10⁻³ mol. c(NaOH) = 2.25×10⁻³ / 0.02500 = 0.0900 mol L⁻¹.
Weak acid + strong base → equivalence point at pH ~8–9 (basic, because the conjugate base hydrolyses). Phenolphthalein changes colour at pH 8.3–10.0, which matches this equivalence point. Methyl orange changes at pH 3.1–4.4 and would indicate endpoint far too early.
At the half-equivalence point, exactly half the weak acid has been converted to its conjugate base, so [HA] = [A⁻]. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]) = pKa + log(1) = pKa + 0 = pKa. This is used to determine the pKa of an unknown weak acid experimentally.
Total n(HCl) = 0.0500 L × 1.00 mol L⁻¹ = 0.0500 mol. n(NaOH) used = 0.0280 L × 0.500 mol L⁻¹ = 0.0140 mol = unreacted HCl. n(HCl) reacted with CaCO₃ = 0.0500 − 0.0140 = 0.0360 mol.
From Henderson-Hasselbalch, pH = pKa + log([A⁻]/[HA]). If [A⁻] = [HA], then log(1) = 0, so pH = pKa. This is why the half-equivalence point is so useful when analysing weak-acid titration curves.
Buffer capacity depends mainly on the total concentration of the weak acid/conjugate base pair. All else equal, the most concentrated buffer resists added acid or base the best, so the 1.00 mol L⁻¹ pair has the highest capacity.
A primary standard must be available in very high purity, stable during storage and weighing, and not significantly hygroscopic or reactive with air. That allows the measured mass to represent a known amount of substance accurately.
The burette should be rinsed with the solution it will contain, here NaOH. If rinsed only with water, residual water dilutes the titrant and introduces a systematic concentration error.
Indicators are weak acids or bases whose two forms have different structures and therefore different colours. The relative amounts of those forms shift with pH, producing the visible colour change.
In a strong acid-weak base titration, the salt formed contains the conjugate acid of the weak base, which hydrolyses and makes the equivalence point acidic. That is why the equivalence pH lies below 7.
At the half-equivalence point, half of the original weak acid has been converted into its conjugate base. The solution therefore contains equal amounts of HA and A⁻, giving pH = pKa.
With phenolphthalein, the correct endpoint is the first faint pink colour that persists after swirling. A deep pink colour means excess titrant has been added and the endpoint has been overshot.
Concordant titres show that repeated measurements agree closely, improving precision and confidence in the average titre. They do not guarantee perfect accuracy, but they reduce random error.
Back titration involves adding a known excess of one reagent to the analyte, then titrating the amount left over. The difference between the amount added and the amount remaining is used to find how much reacted with the analyte.
Conductometric titration monitors conductivity, so no visual colour change is required. That makes it useful for coloured, opaque, or otherwise hard-to-observe solutions where ordinary indicators are unreliable.
In HCl titrated with NaOH, highly mobile H⁺ ions are progressively replaced by the less mobile Na⁺ ions, so conductance falls. After equivalence, excess OH⁻ accumulates and conductance rises again, giving a minimum near equivalence.
Potentiometric methods use a probe to measure pH or electrode potential as titrant is added. The endpoint is identified from the sharp change in the measured signal, often with software assistance.
A derivative plot makes the steepest part of the titration curve stand out as a peak. That helps locate the equivalence point more precisely than estimating it visually from the original curve.
Automated titrators improve repeatability, reduce operator subjectivity, and can detect endpoints using probes and digital analysis. This is especially valuable in industrial quality-control settings where consistency matters.
(a) pH = pKa + log([A⁻]/[HA]) = −log(1.8×10⁻⁵) + log(0.10/0.15). (1 mark) pKa = 4.74. log(0.10/0.15) = log(0.667) = −0.176. pH = 4.74 − 0.18 = 4.56. (1 mark)
(b) NaOH reacts with CH₃COOH: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O. New [HA] = 0.14 mol, [A⁻] = 0.11 mol. (1 mark) pH = 4.74 + log(0.11/0.14) = 4.74 − 0.104 = 4.64. (1 mark) pH increases only slightly (~0.08 units) because the buffer capacity resists change by consuming the added OH⁻. (1 mark)
(a) n(NaOH) = 0.01850 × 0.100 = 1.85 × 10⁻³ mol. Since 1:1 ratio, n(CH₃COOH) = 1.85 × 10⁻³ mol. c(CH₃COOH) = 1.85×10⁻³ / 0.02000 = 0.0925 mol L⁻¹. (2 marks)
(b) Phenolphthalein (pH 8.3–10.0). Weak acid + strong base gives an equivalence point at pH ~8–9 (basic due to hydrolysis of CH₃COO⁻). Phenolphthalein's colour change overlaps this region; methyl orange changes at pH 3–4, well below the equivalence point pH. (2 marks)
Curve description (1 mark each for 2 key features):
• Starts at pH ~2.9 (weak acid, partial dissociation). Rises gradually through a buffer region. Half-equivalence at 12.5 mL NaOH: pH = pKa ≈ 4.74. Equivalence at 25 mL NaOH: pH ≈ 8.7 (basic). Curve flattens after equivalence.
Indicator justification (2 marks): The equivalence point is at ~pH 8.7. Phenolphthalein (pH 8.3–10.0) changes colour within this range → accurate endpoint. Methyl orange changes at pH 3.1–4.4, which falls in the buffer region, far before the equivalence point → endpoint detected too early → systematic error in concentration calculation.
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