Physical chemistry

1. What is lattice enthalpy (of formation)?

1. Always positive
2. The enthalpy change forming one mole of ionic compound from gaseous ions, always exothermic
3. The energy to dissolve an ionic solid
4. The bond energy of a covalent bond

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B: The enthalpy change forming one mole of ionic compound from gaseous ions, always exothermic. Ions coming together to form a lattice always releases energy.

2. Why is lattice enthalpy more exothermic for smaller, more highly charged ions?

1. Weaker electrostatic attraction between ions
2. Stronger electrostatic attraction between ions
3. They form covalent bonds instead
4. There is no relationship

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B: Stronger electrostatic attraction between ions. Smaller, more highly charged ions attract each other more strongly.

3. What does entropy measure?

1. Temperature
2. The number of ways particles/energy can be arranged (disorder)
3. The mass of a substance
4. Enthalpy change

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B: The number of ways particles/energy can be arranged (disorder). Higher entropy means greater disorder.

4. In which state does a substance have the highest entropy?

1. Solid
2. Liquid
3. Gas
4. They are all equal

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C: Gas. Gas particles have the most freedom of arrangement and movement.

5. A reaction has positive ΔH and positive ΔS. Under what condition might it become feasible?

1. Low temperature
2. High temperature, so TΔS becomes larger than ΔH
3. It's never feasible
4. Only at 0 K

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B: High temperature, so TΔS becomes larger than ΔH. At high enough T, the TΔS term outweighs the positive ΔH, making ΔG negative.

6. What is the sign of the entropy change when a solid dissolves to form a solution?

1. Always negative
2. Usually positive, as disorder increases
3. Always zero
4. Cannot be determined

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B: Usually positive, as disorder increases. Particles become more randomly arranged in solution.

7. Which statement defines the enthalpy change of hydration of an ion?

1. ΔH to break one mole of ionic solid into gaseous ions
2. ΔH when one mole of solid dissolves in water
3. ΔH when one mole of water forms from its elements
4. ΔH when one mole of gaseous ions dissolves in water

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D: ΔH when one mole of gaseous ions dissolves in water. Hydration starts from gaseous ions and ends with aqueous ions, so it is always exothermic. Starting from the solid describes enthalpy of solution instead.

8. Which Born–Haber cycle step for NaCl is endothermic?

1. Na(g) → Na⁺(g) + e⁻
2. Na(s) + ½Cl₂(g) → NaCl(s)
3. Na⁺(g) + Cl⁻(g) → NaCl(s)
4. Cl(g) + e⁻ → Cl⁻(g)

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A: Na(g) → Na⁺(g) + e⁻. Ionisation needs energy to remove an electron. The electron affinity of Cl, lattice formation and ΔHf of NaCl are all exothermic.

9. CaCO₃ decomposition: ΔH = +178 kJ mol⁻¹, ΔS = +161 J K⁻¹ mol⁻¹. What is the minimum temperature for feasibility?

1. 1380 K
2. 1106 K
3. 833 K
4. 1.11 K

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B: 1106 K. T = ΔH/ΔS = 178 000 ÷ 161 = 1106 K. Forgetting to convert kJ to J gives 1.11 K.

10. Which compound has the most exothermic lattice enthalpy?

1. MgO
2. CaCl₂
3. KBr
4. NaCl

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A: MgO. Mg²⁺ and O²⁻ are both small and doubly charged, so the ionic attraction is strongest and the lattice enthalpy most exothermic.