Processing the Data — practice questions

A 20.0 mL aliquot of a solution of sodium carbonate required 18.60 mL of 0.1004 mol L⁻¹ hydrochloric acid to reach the endpoint. The equation is 2HCl + Na₂CO₃ → 2NaCl + H₂O + CO₂. Calculate the concentration of the sodium carbonate solution.

A vitamin C tablet is dissolved and made up to 250.0 mL. A 25.0 mL aliquot required 14.85 mL of 0.01012 mol L⁻¹ iodine solution. The reaction is C₆H₈O₆ + I₂ → C₆H₆O₆ + 2HI. Calculate the mass of vitamin C in the tablet and compare it with a label claim of 500 mg. (M(C₆H₈O₆) = 176.14 g mol⁻¹)

A student calculating the concentration of ethanoic acid in vinegar obtains 8.25 mol L⁻¹. The label states 4% acidity. Evaluate this result, identify the most likely error, and justify your reasoning.

State the three quantities of reagent that must be calculated in a back titration, and the relationship between them.

1.000 g of impure calcium carbonate was added to 40.0 mL of 0.500 mol L⁻¹ HCl. The excess acid required 22.50 mL of 0.200 mol L⁻¹ NaOH. Calculate the percentage purity of the calcium carbonate. (M = 100.09 g mol⁻¹)

A student performing a back titration on an antacid tablet does not warm the mixture and titrates it after only two minutes. Some undissolved solid is still visible. Analyse the effect this has on the calculated mass of calcium carbonate, and justify your reasoning.

A calculation uses a titre of 18.65 mL, a standard solution of 0.1004 mol L⁻¹ and a 25.0 mL pipetted aliquot. State how many significant figures the final answer should have and explain why.

Explain why rounding intermediate values in a titration calculation is poor practice, and illustrate with a numerical example.

A student determines a household bleach to contain sodium hypochlorite at 0.732 mol L⁻¹, against a label claim of "4.2% available chlorine". Evaluate this outcome in relation to the consumer product. (M(NaOCl) = 74.44 g mol⁻¹; % available chlorine is conventionally expressed as the equivalent mass of Cl₂, M = 70.90 g mol⁻¹)