Paper 3: General and Practical Principles in Chemistry

A student titrated 25.0 cm³ of sodium hydroxide with 0.100 mol dm⁻³ hydrochloric acid. Her titres were 22.45, 22.10, 22.15 and 22.20 cm³. Calculate the mean titre using concordant results and the concentration of the sodium hydroxide.

The burette reading uncertainty is ±0.05 cm³. Calculate the percentage uncertainty in a titre of 22.15 cm³ and suggest one way to reduce it.

Explain why methyl orange or phenolphthalein would both be suitable for this titration.

A spirit burner containing propan-1-ol heated 100 g of water from 20.5 °C to 45.0 °C. The burner lost 0.62 g. Calculate the enthalpy of combustion of propan-1-ol (Mᵣ 60.0) and explain why the value is much less exothermic than the data-book value (−2021 kJ mol⁻¹).

Suggest three improvements to the apparatus.

Describe how to prepare a pure sample of 1-bromobutane from butan-1-ol, sodium bromide and concentrated sulfuric acid. Include purification steps.

Give three hazards in this preparation and how to control them.

Describe how a colorimeter could be used to determine the order of reaction with respect to iodine in the acid-catalysed iodination of propanone.

The rate equation is rate = k[CH₃COCH₃][H⁺]. Explain what this tells you about the role of iodine in the mechanism.

Methanol is made from CO and H₂: CO(g) + 2H₂(g) ⇌ CH₃OH(g), ΔH = −91 kJ mol⁻¹. Discuss the choice of conditions (about 250 °C, 50–100 atm, copper-based catalyst) in terms of yield, rate, cost and safety.

At equilibrium the partial pressures are p(CO) = 20 atm, p(H₂) = 40 atm, p(CH₃OH) = 15 atm. Calculate Kp with units.

An iron tablet containing iron(II) sulfate was dissolved in dilute sulfuric acid and titrated with 0.0100 mol dm⁻³ KMnO₄. The titre was 21.50 cm³. Calculate the mass of iron in the tablet.

Explain why no indicator is needed in this titration.

For N₂O₄(g) → 2NO₂(g), ΔH = +58 kJ mol⁻¹ and ΔS(system) = +176 J K⁻¹ mol⁻¹. Calculate the temperature above which the reaction is feasible and explain why a gas-phase dissociation has a positive entropy change.

N₂O₄ is colourless and NO₂ is brown. Predict and explain what you would see when a sealed tube of the equilibrium mixture is placed in hot water.

An organic compound contains 54.5% C, 9.1% H and 36.4% O by mass. Its mass spectrum shows M⁺ at m/z 88. It reacts with sodium carbonate to produce CO₂. Its ¹H NMR spectrum shows a doublet (6H), a multiplet (1H) and a singlet (1H) at δ 11.5. Identify the compound.

Draw or name two structural isomers of this compound that are esters.

Discuss how the type of bonding affects the physical properties of sodium chloride, magnesium, silicon dioxide and iodine.

Describe how to measure the standard electrode potential of the Cu²⁺/Cu half-cell, including the conditions.

What volume of 0.050 mol dm⁻³ H₂SO₄ is needed to neutralise 20.0 cm³ of 0.10 mol dm⁻³ NaOH?

Which is the oxidising agent in 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂?

Which reagent turns orange to green with a primary alcohol?

Which IR absorption confirms a carboxylic acid?

The units of a second-order rate constant are:

Which of these has the lowest pH at the same concentration?

Which ion has the electronic configuration [Ar]3d⁶?

Which compound reacts with both HCl and NaOH?

A student used thin-layer chromatography to check the purity of aspirin she had made. (a) Describe how to carry out the TLC and calculate Rf values. (b) The product showed two spots: Rf 0.62 (matching pure aspirin) and Rf 0.35 (matching salicylic acid). Explain what this shows and how she could purify the product. (c) Give one other method to check purity, with the expected result for a pure sample.

Describe how to determine the acid dissociation constant, Ka, of a weak acid using a pH meter and sodium hydroxide solution.

2.46 g of hydrated magnesium sulfate, MgSO₄·xH₂O, was heated in a crucible until it reached constant mass. 1.20 g of anhydrous MgSO₄ remained. Calculate x, and explain why the sample is heated to constant mass. (Mᵣ MgSO₄ = 120.4)

The enthalpy change for 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(l) + CO₂(g) cannot be measured directly. A student measured: Na₂CO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g) ΔH₁ = −35.6 kJ mol⁻¹ NaHCO₃(s) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g) ΔH₂ = +27.8 kJ mol⁻¹ Use a Hess cycle to calculate the enthalpy change for the decomposition, and explain why it cannot be measured directly.

Describe a test for each of the following ions and give the positive result: NH₄⁺, SO₄²⁻, CO₃²⁻, Br⁻.