Volumetric (titration) quantitative analysis: Practice Questions

Which of the following procedures is the most appropriate way to rinse a burette before starting an acid-base titration?

A student uses 0.0500 M standard \(Na_2CO_3(aq)\) to titrate 25.00 \(cm^3\) of a sulfuric acid solution \(H_2SO_4(aq)\). If the average titre of \(Na_2CO_3(aq)\) required is 18.40 \(cm^3\), what is the molarity of the sulfuric acid?

Reaction: \(Na_2CO_3(aq) + H_2SO_4(aq) \rightarrow Na_2SO_4(aq) + H_2O(l) + CO_2(g)\)

A student weighs out \(8.50 \text{ g}\) of an impure sample of iron(II) sulfate-7-water (\(FeSO_4 \cdot 7H_2O\)) and dissolves it in deionised water to make \(250.0 \text{ cm}^3\) of solution. A \(25.00 \text{ cm}^3\) aliquot of this solution is acidified and then titrated with \(0.0200 \text{ M}\) potassium permanganate (\(KMnO_4\)) solution. The average titre required to reach the endpoint is \(28.50 \text{ cm}^3\).

What is the percentage purity by mass of \(FeSO_4 \cdot 7H_2O\) in the sample?
(Relative atomic masses: \(H = 1.0\), \(O = 16.0\), \(S = 32.1\), \(Fe = 55.8\))

When performing a titration to determine the concentration of an unknown alkali, a student chooses phenolphthalein as the indicator. What is the observed color change at the end point if the alkali is in the conical flask and the acid is in the burette?

A student aims to determine the percentage of available chlorine in a commercial bleach sample. The student first pipettes \(25.00 \text{ cm}^3\) of the bleach and dilutes it to \(250.0 \text{ cm}^3\) in a volumetric flask. Then, \(25.00 \text{ cm}^3\) of the diluted solution is transferred to a conical flask, followed by the addition of an excess of acidified potassium iodide (\(KI\)) solution. The liberated iodine is then titrated against \(0.100 \text{ M}\) sodium thiosulfate (\(Na_2S_2O_3\)) solution using starch indicator. The average titre is \(18.40 \text{ cm}^3\).

What is the concentration of sodium hypochlorite (\(NaOCl\)) in the original bleach sample in \(\text{mol dm}^{-3}\)?
(The relevant equations are:
\(ClO^-(aq) + 2I^-(aq) + 2H^+(aq) \rightarrow Cl^-(aq) + I_2(aq) + H_2O(l)\)
\(I_2(aq) + 2S_2O_3^{2-}(aq) \rightarrow 2I^-(aq) + S_4O_6^{2-}(aq)\))

In the titrimetric analysis of chloride ions using Mohr's method, the solution's pH must be carefully maintained between 6.5 and 9.0. Explain the chemical consequence regarding the indicator if the pH is significantly lower than this range.

In an iodometric titration to determine the amount of copper(II) ions, the student accidentally adds excess starch indicator prematurely, before the titration is near its end point. Explain the chemical reason for the irreversible colour change that often results and how this error would affect the calculated amount of copper(II) ions (increase, decrease, or no change).

In determining the concentration of aqueous ammonia (a weak base) using excess standard hydrochloric acid (strong acid) followed by titration with standard sodium hydroxide (strong base) (back titration), explain why methyl orange is a suitable indicator for the final titration step.

A primary standard solution of anhydrous sodium carbonate (\(\text{Na}_2\text{CO}_3\), Molar mass = \(106.0 \text{ g mol}^{-1}\)) is prepared by dissolving \(1.325 \text{ g}\) of solid \(\text{Na}_2\text{CO}_3\) in distilled water and making the solution up to \(250.0 \text{ cm}^3\) in a volumetric flask. This solution is used to standardise a dilute sulfuric acid (\(\text{H}_2\text{SO}_4\)) solution.

\(25.0 \text{ cm}^3\) portions of the \(\text{Na}_2\text{CO}_3\) solution required an average titre of \(23.85 \text{ cm}^3\) of the \(\text{H}_2\text{SO}_4\) solution for complete reaction.

a) Write the balanced ionic equation for the reaction occurring during the titration.

b) Calculate the molar concentration of the standard \(\text{Na}_2\text{CO}_3\) solution.

c) Determine the concentration of the \(\text{H}_2\text{SO}_4\) solution in \(\text{mol dm}^{-3}\).

A student intended to determine the concentration of magnesium ions in a sample of deep-sea water. However, the sample was first contaminated with a small amount of hydrochloric acid. To analyze the sample, a \(25.0 \text{ cm}^3\) portion of the sea water was first treated with excess sodium hydroxide solution to precipitate all magnesium ions as \(\text{Mg(OH)}_2(s)\). The mixture was filtered, and the residue was washed and then reacted with \(50.0 \text{ cm}^3\) of \(0.100 \text{ M}\) standard sulfuric acid (excess).

The resulting solution, which contained the unreacted sulfuric acid, was then titrated against a \(0.0500 \text{ M}\) sodium carbonate solution. It was found that \(18.40 \text{ cm}^3\) of the sodium carbonate solution was required to reach the methyl orange endpoint.

a) Explain why the sea water was treated with excess hydroxide before titration rather than titrating the original sea water sample directly with a base to find the \(\text{Mg}^{2+}\) concentration. (2 marks)

b) Calculate the concentration of \(\text{Mg}^{2+}\) ions in the original sea water sample in \(\text{g dm}^{-3}\). (4 marks)

c) A student suggests that using a phenolphthalein indicator instead of methyl orange in the final titration would result in a higher calculated value for the \(\text{Mg}^{2+}\) concentration. Evaluate this statement with reference to the pH range of the indicators and the nature of the reaction. (3 marks)

(Relative atomic mass: \(\text{Mg} = 24.3\))