Titration Calculations, Standard Solutions, Concordant Titres & Errors

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

We just saw what a primary standard is and why NaOH cannot be one. That raises a question: Once you have your standard solution prepared, what specific technique is required for each piece of titration equipment, and why does each step exist? This card answers it → rinse burette with titrant, pipette with analyte, flask with distilled water only; read at bottom of meniscus at eye level; endpoint ≠ equivalence point.

Every specific technique requirement in a titration exists for a quantitative reason, and being able to explain why each step is performed, not just describe what is done, is what distinguishes a Band 6 practical report from a Band 3 description.

Burette technique: The burette is a precision glass tube graduated to ±0.05 mL (0.01 mL graduations; read to nearest 0.05 mL). Before use: rinse twice with small volumes of the titrant solution, residual water would dilute the titrant and reduce its effective concentration. Fill above the zero mark, then open the stopcock to fill the tip and remove air bubbles, air bubbles in the tip produce an erroneously large delivered volume when released. Read the burette at the bottom of the meniscus, with the eye level with the liquid surface to eliminate parallax error. Record readings to ±0.05 mL (two decimal places: e.g. 16.45 mL). Titre = final reading − initial reading.

Pipette technique: A volumetric pipette delivers a single, fixed, accurately known volume (e.g. 10.00 mL or 25.00 mL). Before use: rinse twice with the analyte solution (not distilled water). Fill by suction using a pipette filler (never by mouth), drain to the calibration mark, touch the tip against the inside of the receiving flask to drain the last drop, do not blow out the remaining small volume (the calibrated volume assumes the last drop remains in the tip).

Conical flask technique: The conical flask holds the analyte. Use a conical flask rather than a beaker because: the tapered neck allows vigorous swirling without spilling; swirling mixes the reactants at the endpoint more effectively than stirring; the narrow opening reduces evaporation. The flask can be rinsed with distilled water during the titration, this does not change the moles of analyte. A white tile placed under the flask improves visibility of the indicator colour change.

Endpoint vs equivalence point: The equivalence point is the theoretical point at which moles of acid exactly equal moles of base (stoichiometric completion). The endpoint is the observed colour change of the indicator. A correct indicator choice makes these coincide; a wrong indicator produces a systematic error (endpoint ≠ equivalence point).

Burette: rinse with titrant; read bottom of meniscus at eye level; record to ±0.05 mL; remove air bubbles from tip. Pipette: rinse with analyte; do not blow out last drop. Conical flask: rinse with distilled water ONLY, never analyte; white tile underneath. Endpoint (indicator colour change) ≠ equivalence point (stoichiometric completion), coincide only with correct indicator choice.

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Syllabus: conduct practical investigations to analyse the concentration of an unknown acid or base by titration

We just saw the technique for each piece of volumetric glassware. That raises a question: how do those techniques combine into a result you can defend to three significant figures? This card answers it → standardise, titrate to concordance, then report with an uncertainty.

Inquiry question. What is the concentration of the supplied unknown sodium hydroxide solution, and what is the uncertainty in that value?

Equipment. 50 mL burette and stand, 20 or 25 mL volumetric pipette and filler, 250 mL conical flasks, 250 mL volumetric flask, funnel, white tile, wash bottle of distilled water, primary standard (anhydrous sodium carbonate or potassium hydrogen phthalate), analytical balance, approximately 0.1 mol L⁻¹ HCl, the unknown NaOH, methyl orange and phenolphthalein.

Analysis and evaluation.

• Write the balanced equation and state the mole ratio you used. Show the full calculation from mean titre to c(NaOH).
• Combine the uncertainties: two burette readings at ±0.05 mL each, the pipette tolerance, and the volumetric flask tolerance. Quote c(NaOH) with an absolute uncertainty and a sensible number of significant figures.
• Which single piece of glassware contributes most to your uncertainty? Support the answer with the percentage uncertainty of each.
• Justify your indicator choice from the equivalence-point pH for this titration, not from habit.
• Suppose the conical flask had been rinsed with the NaOH before the aliquot was added. Would your calculated c(NaOH) be too high or too low? Explain in terms of moles, not volumes.

Conclusion. Report c(NaOH) with its uncertainty and state the evidence for its reliability, which is the concordance of your titres, not the number of times you titrated.

Sequence: weigh primary standard by difference → make to the mark → standardise the HCl → condition burette and pipette with their own solution, flask with water only → rough titre → accurate titres to three concordant within 0.10 mL → average the concordant only → n = cV and the mole ratio → report with combined uncertainty.

Pause, write the highlighted sequence into your book.

We just saw the four-step concentration calculation applicable to any acid-base titration. That raises a question: Which specific household substance titrations appear in the HSC prescribed investigation, and what are the indicator choices and calculation routes for each? This card answers it → vinegar: CH₃COOH + NaOH (1:1), phenolphthalein; antacid: NaHCO₃ + HCl (1:1), methyl orange.

The HSC prescribed investigation requires chemical analysis of a common household substance, both vinegar (acetic acid concentration) and antacid tablets (sodium hydrogen carbonate or calcium carbonate content) are standard examples that connect titration technique directly to consumer chemistry.

Vinegar analysis (% acetic acid by mass): Vinegar is a dilute solution of acetic acid (CH₃COOH) in water. Commercial vinegar is labelled "5% acidity", approximately 5 g of acetic acid per 100 mL. To verify this: pipette 10.00 mL of vinegar into a conical flask; add 2–3 drops of phenolphthalein (appropriate because CH₃COOH + NaOH is weak acid + strong base, equivalence point pH ≈ 8.7, within phenolphthalein's range 8.3–10.0); titrate with standard NaOH until the first permanent faint pink colour persists for 30 seconds. Apply four-step method: write equation (CH₃COOH + NaOH → CH₃COONa + H₂O; 1:1); calculate n(NaOH) = c × V; n(CH₃COOH) = n(NaOH); mass(CH₃COOH) = n × 60.06 g/mol; % = (mass/mass of vinegar sample) × 100%.

Antacid analysis (NaHCO₃ or CaCO₃ content): Dissolve a known mass of antacid tablet in distilled water; titrate with standard HCl using methyl orange as indicator (NaHCO₃ + HCl reaction gives EP at acidic pH, within methyl orange's range 3.1–4.4). Equation: NaHCO₃ + HCl → NaCl + H₂O + CO₂ (1:1). Calculate n(HCl), then n(NaHCO₃) = n(HCl), then mass(NaHCO₃) = n × 84.01 g/mol, then % = (mass/mass of tablet) × 100%.

Vinegar: CH₃COOH + NaOH (1:1); phenolphthalein (EP pH ≈ 8.7; range 8.3–10.0); % = (n × 60.06 / mass of sample) × 100%. Antacid: NaHCO₃ + HCl (1:1); methyl orange (EP pH acidic; range 3.1–4.4); % = (n × 84.01 / mass of tablet) × 100%. For liquids: mass = V × density.

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