Chemical energetics (A Level: lattice energy, Born–Haber, hydration) Banker

1. (a) Define the term lattice energy. [2]
2. (b) Write an equation, with state symbols, for the lattice energy of magnesium chloride. [1]
3. (c) State whether lattice energy is exothermic or endothermic and explain why. [1]

Data (kJ/mol): ΔHf⦵ NaCl(s) −411; ΔHat⦵ Na(s) +107; first ionisation energy of Na +496; ΔHat⦵ of chlorine (½Cl₂ → Cl) +122; first electron affinity of Cl −349.

1. (a) Use the data to calculate the lattice energy of sodium chloride. [2]
2. (b) Name the enthalpy change ½Cl₂(g) → Cl(g). [1]

1. (a) Define the standard enthalpy change of atomisation of an element. [2]
2. (b) Write equations, with state symbols, for the standard enthalpy change of atomisation of (i) sodium and (ii) chlorine. [2]
3. (c) Explain why the enthalpy change of atomisation of chlorine is half the Cl–Cl bond energy. [1]

Data (kJ/mol): ΔHat⦵ Li +159; 1st IE Li +520; ΔHat⦵ F (½F₂ → F) +79; 1st EA F −328; lattice energy LiF −1046.

1. (a) Use the data to calculate the standard enthalpy change of formation of lithium fluoride. [2]
2. (b) Write the equation for the standard enthalpy change of formation of lithium fluoride. [1]

1. (a) List, in order, the enthalpy changes in a Born–Haber cycle that lead from K(s) and ½Cl₂(g) to K⁺(g) and Cl⁻(g), giving an equation for each. [4]
2. (b) State which of the steps in (a) are exothermic. [1]
3. (c) Given ΔHf⦵ KCl = −437, ΔHat K = +89, IE K = +419, ΔHat Cl = +122 and EA Cl = −349 (all kJ/mol), calculate the lattice energy of KCl. [1]

Data (kJ/mol): ΔHf⦵ MgO(s) −602; ΔHat⦵ Mg +148; 1st IE Mg +738; 2nd IE Mg +1451; ΔHat⦵ O (½O₂ → O) +249; 1st EA O −141; 2nd EA O +798.

1. (a) Calculate the lattice energy of magnesium oxide. [3]
2. (b) Suggest why magnesium oxide is used to line the inside of furnaces. [1]

Calcium chloride is used to melt ice on roads. Data (kJ/mol): ΔHf⦵ CaCl₂(s) −796; ΔHat⦵ Ca +178; 1st IE Ca +590; 2nd IE Ca +1145; ΔHat⦵ Cl (½Cl₂ → Cl) +122; 1st EA Cl −349.

1. (a) Calculate the lattice energy of calcium chloride. [3]
2. (b) State the two ionisation energy equations included in the cycle. [1]

Data (kJ/mol): ΔHf⦵ KBr(s) −394; ΔHat⦵ K +89; 1st IE K +419; ΔHat⦵ Br (½Br₂(l) → Br(g)) +112; 1st EA Br −325.

1. (a) Calculate the lattice energy of potassium bromide. [2]
2. (b) The Br–Br bond energy is 193 kJ/mol. Explain why the enthalpy change of atomisation of bromine (+112 kJ/mol) is greater than ½ × 193. [2]

Lattice energies (kJ/mol): NaF −918; NaCl −787; NaBr −747.

1. (a) Explain the trend in lattice energy from NaF to NaBr. [2]
2. (b) Explain why the lattice energy of MgO (−3845 kJ/mol) is much more exothermic than that of NaF. [2]

Data (kJ/mol): ΔHf⦵ SrO(s) −592; 1st IE Sr +550; 2nd IE Sr +1064; ΔHat⦵ O +249; 1st EA O −141; 2nd EA O +798; lattice energy SrO −3276.

1. (a) Use the data to calculate the enthalpy change of atomisation of strontium. [3]
2. (b) Suggest why the lattice energy of SrO is less exothermic than that of CaO (−3454 kJ/mol). [1]

Data (kJ/mol): ΔHf⦵ MgF₂(s) −1124; ΔHat⦵ Mg +148; 1st IE Mg +738; 2nd IE Mg +1451; ΔHat⦵ F +79; 1st EA F −328.

1. (a) Calculate the lattice energy of magnesium fluoride. [2]
2. (b) The lattice energy of MgO is −3845 kJ/mol. Predict which of MgO and MgF₂ has the higher melting point and explain your answer. [2]

Data (kJ/mol): ΔHf⦵ CaO(s) −635; ΔHat⦵ Ca +178; 1st IE Ca +590; 2nd IE Ca +1145; ΔHat⦵ O +249; 1st EA O −141; lattice energy CaO −3454.

1. (a) Use the data to calculate the second electron affinity of oxygen. [3]
2. (b) Explain why the second electron affinity of oxygen is endothermic. [2]

Data (kJ/mol): ΔHat⦵ Mg +148; 1st IE Mg +738; 2nd IE Mg +1451; ΔHat⦵ Cl +122; 1st EA Cl −349.

1. (a) Calculate ΔHf⦵ for the hypothetical compound MgCl(s), taking its lattice energy as −753 kJ/mol. [2]
2. (b) Calculate ΔHf⦵ for MgCl₂(s), whose lattice energy is −2526 kJ/mol. [2]
3. (c) Use your answers to explain why magnesium forms MgCl₂ and not MgCl, even though forming Mg²⁺ needs much more energy than forming Mg⁺. [2]
4. (d) Suggest why MgCl₃ does not form. [1]

Aluminium oxide is mined as bauxite. Data (kJ/mol): ΔHf⦵ Al₂O₃(s) −1676; ΔHat⦵ Al +326; IE of Al: 1st +578, 2nd +1817, 3rd +2745; ΔHat⦵ O +249; 1st EA O −141; 2nd EA O +798.

1. (a) Write the equation for the lattice energy of aluminium oxide. [1]
2. (b) Calculate the lattice energy of aluminium oxide. [3]

Data (kJ/mol): ΔHf⦵ CsCl(s) −443; ΔHat⦵ Cs +76; 1st IE Cs +376; ΔHat⦵ Cl +122; lattice energy CsCl −668.

1. (a) Use the data to calculate the first electron affinity of chlorine. [3]
2. (b) Suggest why the first ionisation energy of caesium is low, and explain why the lattice energy of CsCl (−668 kJ/mol) is less exothermic than that of NaCl (−787 kJ/mol). [2]

1. (a) Define the standard enthalpy change of hydration of an ion. [2]
2. (b) Define the standard enthalpy change of solution. [1]
3. (c) Write equations, with state symbols, for (i) the enthalpy change of hydration of K⁺ and (ii) the enthalpy change of solution of NaCl. [2]

Data (kJ/mol): lattice energy NaCl −787; ΔHhyd⦵ Na⁺ −419; ΔHhyd⦵ Cl⁻ −364.

1. (a) Describe, in words, an energy cycle linking the lattice energy, the hydration enthalpies and the enthalpy change of solution of NaCl. [1]
2. (b) Calculate the enthalpy change of solution of NaCl. [2]

1. (a) Explain why the hydration enthalpy of F⁻ (−506 kJ/mol) is more exothermic than that of Cl⁻ (−364 kJ/mol). [2]
2. (b) State the type of interaction between an ion and the water molecules around it. [1]

Hydration enthalpies (kJ/mol): Li⁺ −519; Na⁺ −419; K⁺ −322; Mg²⁺ −1920.

1. (a) Explain the trend in hydration enthalpy from Li⁺ to K⁺. [2]
2. (b) Explain why the hydration enthalpy of Mg²⁺ is much more exothermic than that of Na⁺. [2]
3. (c) Describe how water molecules are arranged around a cation and around an anion. [1]

Data (kJ/mol): enthalpy change of solution of KCl +17.2; lattice energy KCl −703; ΔHhyd⦵ K⁺ −322.

1. (a) Calculate the hydration enthalpy of the chloride ion. Give your answer to 3 significant figures. [3]
2. (b) Predict whether the temperature rises or falls when KCl dissolves in water. [1]

A student investigates instant cold packs used by sports teams. 8.00 g of ammonium nitrate is added to 50.0 g of water in a polystyrene cup. The temperature falls from 21.5 °C to 10.3 °C.

1. (a) Calculate the heat energy absorbed from the water. [c = 4.18 J g⁻¹ K⁻¹] [1]
2. (b) Calculate the enthalpy change of solution of ammonium nitrate, in kJ/mol, including its sign. [Mr: NH₄NO₃, 80.0] [2]
3. (c) State two assumptions made in your calculation. [2]
4. (d) Suggest one improvement to the method that would make the result more accurate. [1]

Data (kJ/mol): lattice energy MgCl₂ −2526; ΔHhyd⦵ Mg²⁺ −1920; ΔHhyd⦵ Cl⁻ −364.

1. (a) Calculate the enthalpy change of solution of magnesium chloride. [2]
2. (b) Explain why the chloride hydration enthalpy must be multiplied by 2 in this calculation. [1]

Data (kJ/mol): lattice energy CaCl₂ −2255; ΔHhyd⦵ Ca²⁺ −1650; ΔHhyd⦵ Cl⁻ −364.

1. (a) Calculate the enthalpy change of solution of calcium chloride. [2]
2. (b) 5.55 g of anhydrous CaCl₂ is dissolved in 100 g of water. Using your answer to (a), calculate the expected temperature rise. [Mr: CaCl₂, 111.1; c = 4.18 J g⁻¹ K⁻¹; assume the solution has mass 100 g] [2]

Data (kJ/mol): enthalpy change of solution of LiF +4.9; ΔHhyd⦵ Li⁺ −519; ΔHhyd⦵ F⁻ −506.

1. (a) Calculate the lattice energy of lithium fluoride from these data. Give your answer to 3 significant figures. [3]
2. (b) Lithium fluoride is only slightly soluble in water, but lithium chloride is very soluble. Suggest an explanation in terms of the enthalpy changes involved. [2]

Data (kJ/mol): lattice energy MgSO₄ −2833, BaSO₄ −2469; ΔHhyd⦵ Mg²⁺ −1920, Ba²⁺ −1360, SO₄²⁻ −1059.

1. (a) Calculate the enthalpy change of solution of MgSO₄ and of BaSO₄. [2]
2. (b) Use the data to explain why the solubility of the Group 2 sulfates decreases down the group. [3]

A student dissolves different masses of potassium nitrate, KNO₃, in 50.0 g portions of water and records the temperature fall.mass of KNO₃ / g: 2.02, 4.04, 6.06, 8.08temperature fall / °C: 3.2, 6.5, 8.1, 13.0

1. (a) Identify the anomalous result and suggest one reason for it. [2]
2. (b) Use the result for 8.08 g to calculate the enthalpy change of solution of potassium nitrate. [Mr: KNO₃, 101.1; c = 4.18 J g⁻¹ K⁻¹] [3]
3. (c) The data-book value is +34.9 kJ/mol. Suggest why the experimental value is lower and how this error could be reduced. [2]

Data (kJ/mol): ΔHf⦵ NaCl(s) −411; ΔHat⦵ Na +107; 1st IE Na +496; ΔHat⦵ Cl +122; 1st EA Cl −349; enthalpy change of solution of NaCl +3.9; ΔHhyd⦵ Cl⁻ −364.

1. (a) Calculate the lattice energy of NaCl. [1]
2. (b) Hence calculate the hydration enthalpy of Na⁺. Give your answer to 3 significant figures. [2]
3. (c) Calculate the enthalpy change for Na(s) + ½Cl₂(g) → Na⁺(aq) + Cl⁻(aq). [1]

1. (a) Define the first electron affinity of an element. [2]
2. (b) Write the equation for the first electron affinity of oxygen. [1]
3. (c) Explain why the first electron affinity of oxygen is exothermic. [1]

First electron affinities (kJ/mol): F −328; Cl −349; Br −325; I −295.

1. (a) Explain the general trend from chlorine to iodine. [2]
2. (b) Suggest why fluorine's first electron affinity is less exothermic than chlorine's. [1]
3. (c) The second electron affinity of sulfur is +640 kJ/mol. Calculate the enthalpy change for S(g) + 2e⁻ → S²⁻(g) given that the first electron affinity of sulfur is −200 kJ/mol. [1]

Data (kJ/mol): 1st IE Mg +738; 2nd IE Mg +1451; 1st EA O −141; 2nd EA O +798.

1. (a) Calculate the enthalpy change for O(g) + 2e⁻ → O²⁻(g). [1]
2. (b) Calculate the total enthalpy change for Mg(g) + O(g) → Mg²⁺(g) + O²⁻(g). [1]
3. (c) Forming the gaseous ions is very endothermic, yet magnesium burns vigorously in oxygen to form MgO. Explain why. [2]