AQA A-Level Chemistry Paper 3, June 2022: Question 17

📋 What this question tests

This question assesses your understanding of trends descending Group 1 (the alkali metals), integrating multiple physical chemistry topics:

Enthalpy of hydration: The relationship between ionic radius, charge density, and electrostatic attraction to polar water molecules.
Metallic bonding & physical properties: How atomic radius influences metallic bond strength, enthalpy of atomisation, and melting points.
Successive ionisation energies: Shielding, distance from nucleus, and electron removal from inner principal quantum energy levels.

The second ionisation energy of rubidium is lower than the second ionisation energy of lithium.

Mark Scheme: D (AO3) — 1 mark

The second ionisation energy represents the removal of an electron from a 1+ gaseous ion:

M⁺(g) → M²⁺(g) + e⁻

For Li⁺ (1s²), the electron is removed from the principal quantum level n = 1 .
For Rb⁺ (1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶), the electron is removed from n = 4 (a 4p orbital).
The electron being removed from Rb⁺ is significantly further from the nucleus and experiences greater shielding from inner shells.
Therefore, electrostatic attraction to the nucleus is weaker in Rb⁺, so less energy is required to remove the second electron from rubidium than from lithium.

Option A is incorrect: Hydration enthalpy (Δ_hydH) depends on charge density (charge ÷ ionic radius). Since Cs⁺ has a larger ionic radius than Rb⁺ while both carry a 1+ charge, Cs⁺ has lower charge density. It attracts water dipoles less strongly, meaning Cs⁺ has a less negative (less exothermic) enthalpy of hydration than Rb⁺.
Option B is incorrect: Enthalpy of atomisation reflects the strength of metallic bonding in the solid lattice:
M(s) → M(g)
As Group 1 is descended, ionic radii increase (K⁺ is larger than Na⁺). The delocalised electron cloud is further from the cation centres, leading to weaker electrostatic attraction in the metallic lattice. Thus, potassium has a lower enthalpy of atomisation than sodium.
Option C is incorrect: For the same reason as atomisation enthalpy, weaker metallic bonding down the group means melting points decrease down Group 1. Therefore, potassium has a lower melting point than sodium (Na m.p. ≈ 98 °C, K m.p. ≈ 63 °C).

Link related properties: Notice that Option B and Option C both depend directly on the strength of metallic bonding. If metallic bonding weakens down the group, both enthalpy of atomisation and melting point must decrease. This allows you to eliminate both B and C simultaneously.
Watch the sign of enthalpy: "More negative" means "more exothermic" (stronger interaction). Since Cs⁺ is larger than Rb⁺, its interactions are weaker, so its hydration enthalpy is less negative. That rules out A.
Write electronic configurations: When comparing 2nd ionisation energies, mentally visualise the ion: Li⁺ has electrons in shell 1; Rb⁺ has valence electrons in shell 4. Lower energy down the group is guaranteed by increased shielding and radius.

Confusing 1st and 2nd Ionisation Energies: Students often mistakenly think that because second ionisation energies for Group 1 elements are huge (breaking into a full shell), the normal trend does not apply. Down a group, all successive ionisation energies follow the general decreasing trend due to increased shielding and atomic/ionic radius.
"More negative" confusion: Forgetting that more exothermic means a larger magnitude with a negative sign. Cs⁺ has a smaller absolute value, so it is less negative.
Reversing metallic bonding trends: Assuming metals become harder or have higher melting points down the group (confusing Group 1 metallic trends with halogen intermolecular trends).

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