Paper 1: Inorganic and Physical Chemistry

Give the full electron configuration of a chromium atom and explain why it is unusual.

Explain why the first ionisation energy decreases down Group 2.

Describe how ions are formed and detected in a time-of-flight mass spectrometer using electrospray ionisation.

0.300 g of an impure sample of calcium carbonate (Mr 100.1) reacted with 36.0 cm³ of 0.150 mol dm⁻³ HCl. Calculate the percentage purity of the sample.

Calculate the volume, in dm³, of 0.250 mol of gas at 298 K and 100 kPa.

Explain why sodium chloride has a high melting point but conducts electricity only when molten or dissolved.

Explain why tetrachloromethane is non-polar even though C–Cl bonds are polar.

Describe a dative covalent bond, using the ammonium ion as an example.

Use Hess's law to calculate ΔfH of propane from enthalpies of combustion (kJ mol⁻¹): C −394, H₂ −286, C₃H₈ −2220.

Define standard enthalpy of combustion.

Sketch (describe) Maxwell–Boltzmann distributions at two temperatures and explain how they show why rate increases with temperature.

Explain why the rate of reaction between marble chips and hydrochloric acid decreases as the reaction proceeds.

Write the Kc expression for 2SO₂ + O₂ ⇌ 2SO₃ and give its units.

Explain the effect of increasing pressure on the yield of SO₃ and on Kc.

Balance: Cr₂O₇²⁻ + Fe²⁺ + H⁺ → Cr³⁺ + Fe³⁺ + H₂O.

In the reaction between dichromate(VI) and iron(II) ions, identify the oxidising agent and the reducing agent.

Use the data to calculate the enthalpy of solution of MgCl₂ (data are illustrative, written for this paper): lattice enthalpy of dissociation +2526 kJ mol⁻¹; hydration enthalpy Mg²⁺ −1920, Cl⁻ −364 kJ mol⁻¹. Explain whether it dissolves exothermically.

Explain why a reaction with positive ΔH and positive ΔS becomes feasible at high temperature.

The rate equation for 2NO + O₂ → 2NO₂ is rate = k[NO]²[O₂]. Explain the effect on the rate of doubling [NO] and halving [O₂].

Explain how the rate-determining step relates to the rate equation, using a reaction whose rate equation is rate = k[(CH₃)₃CBr].

For H₂(g) + CO₂(g) ⇌ H₂O(g) + CO(g), equilibrium partial pressures are p(H₂) 30 kPa, p(CO₂) 30 kPa, p(H₂O) 20 kPa, p(CO) 20 kPa (data are illustrative, written for this paper). Calculate Kp and explain why it has no units.

Use E° values to predict whether Fe³⁺ will oxidise I⁻. (Fe³⁺/Fe²⁺ +0.77 V; I₂/I⁻ +0.54 V)

Explain why predictions using E° values may not happen in practice.

A weak acid HA has pH 2.70 at 0.100 mol dm⁻³. Calculate Ka.

Explain how an ethanoic acid / ethanoate buffer resists a change in pH when a small amount of alkali is added.

Calculate the pH after 25.0 cm³ of 0.200 mol dm⁻³ HCl is added to 25.0 cm³ of 0.100 mol dm⁻³ NaOH.

Explain the trend in melting points of the Period 3 elements sodium, magnesium and aluminium.

Describe the reactions of magnesium with steam and of calcium with cold water. Include equations.

Give one use of barium sulfate and explain why the toxic barium ion is safe to use.

Describe how silver nitrate and ammonia can be used to identify chloride, bromide and iodide ions.

Explain the trend in oxidising ability of the halogens.

Write an ionic equation for the reaction of chlorine water with potassium bromide solution and give the colour change.

Explain why sodium oxide has a high melting point and phosphorus(V) oxide a much lower one.

Explain the meaning of the term ligand substitution, using [Co(H₂O)₆]²⁺ and concentrated HCl as an example, and state the colour change.

Explain the chelate effect using EDTA⁴⁻ and [Cu(H₂O)₆]²⁺.

Describe what you would see when sodium carbonate solution is added to [Al(H₂O)₆]³⁺ and explain.

Methanol is made industrially: CO(g) + 2H₂(g) ⇌ CH₃OH(g), ΔH = −91 kJ mol⁻¹. Discuss the choice of temperature, pressure and catalyst, referring to yield, rate, feasibility and cost.