AS Chemistry: Periodicity Across Period 3 — Practice Questions

These are original questions written for Marlbridge, for revision and practice on this content. They are not reproduced past-paper questions, and they do not replicate the exam’s exact structure, question count or mark tariffs — examination boards hold copyright in their own papers. Use these alongside the official past papers available free from your board.

Related: Periodicity revision notes

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Section A

1. Describe and explain the trend in atomic radius across period 3. [3]

2. Describe the trend in electronegativity across period 3 and explain it. [3]

3. State the structure and bonding of Na, Si and S₈. [3]

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Section B

4. The melting points across period 3 rise to a maximum at silicon and then fall sharply.

(a) Explain why melting point increases from sodium to aluminium. [3]

(b) Explain why silicon has the highest melting point. [2]

(c) Explain the sharp drop from silicon to phosphorus. [3]

(d) Explain why sulfur (S₈) has a higher melting point than phosphorus (P₄). [2]

5. Sodium oxide and sulfur dioxide are both added to water.

(a) Write an equation for each reaction. [2]

(b) Predict the approximate pH of each resulting solution. [2]

(c) Explain the trend in acid–base character of the period 3 oxides. [3]

6. Explain what happens when aluminium oxide is added to (a) hydrochloric acid, (b) sodium hydroxide, and state the term used to describe such behaviour. [3]

7. Sodium chloride and silicon tetrachloride are each added to water separately.

(a) Describe what is observed in each case. [2]

(b) Explain the difference in terms of bonding. [3]

8. First ionisation energy generally increases across Period 3, but there are dips at aluminium and at sulfur.

(a) Explain the general increasing trend across the period. [2]

(b) Explain why aluminium has a lower first ionisation energy than magnesium. [2]

(c) Explain why sulfur has a lower first ionisation energy than phosphorus. [2]

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Answers

1. Atomic radius decreases [1]. Nuclear charge increases while shielding stays approximately constant [1], so the outer electrons are pulled in more strongly [1].

2. Electronegativity increases [1]. The nuclear charge increases and the atomic radius decreases with similar shielding [1], so the nucleus attracts bonding electrons more strongly [1].

3. Na — giant metallic [1]. Si — giant covalent [1]. S₈ — simple molecular [1].

4. (a) All are giant metallic structures [1]. Across Na → Mg → Al the number of delocalised electrons per atom increases (1, 2, 3) and the ionic radius decreases [1], so the metallic bonding is stronger and more energy is needed to break it [1].

(b) Silicon is a giant covalent structure [1] in which many strong covalent bonds must be broken [1].

(c) Phosphorus is simple molecular (P₄) [1], so only weak instantaneous dipole–induced dipole forces between molecules must be overcome [1], not covalent bonds — far less energy is required [1].

(d) S₈ has more electrons than P₄ [1], so the instantaneous dipole–induced dipole forces between molecules are stronger [1].

5. (a) Na₂O + H₂O → 2NaOH [1]; SO₂ + H₂O → H₂SO₃ [1].

(b) Sodium hydroxide solution: pH about 13–14 [1]. Sulfurous acid: pH about 2–3 [1].

(c) Oxides change from basic (Na₂O, MgO) through amphoteric (Al₂O₃) to acidic (SiO₂, P₄O₁₀, SO₂) [1]. This reflects the change from metallic to non-metallic character across the period [1], and hence from ionic to covalent bonding in the oxides [1].

6. (a) It reacts, behaving as a base: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O [1]. (b) It also reacts, behaving as an acid, forming the aluminate ion [1]. It is described as amphoteric [1].

7. (a) Sodium chloride dissolves with no visible reaction, giving a neutral (pH 7) solution [1]. Silicon tetrachloride undergoes violent hydrolysis, giving off misty white fumes of hydrogen chloride and leaving an acidic (pH ≈ 2) solution [1]. (b) Sodium chloride is ionic, so it simply dissociates into its existing Na⁺ and Cl⁻ ions on dissolving [1]. Silicon tetrachloride is covalent, so instead of dissolving as intact molecules it reacts with water, breaking Si–Cl bonds and forming Si–O bonds as well as releasing HCl [1] — the transition from dissolving to hydrolysing tracks the change from ionic to covalent bonding across the period [1].

8. (a) Across the period, nuclear charge increases while electrons enter the same outer shell with roughly constant shielding [1], so successive outer electrons are held more strongly, raising the energy needed to remove one [1]. (b) Aluminium’s outer electron is in a 3p subshell, which is at higher energy and further from the nucleus (partially shielded by the 3s² electrons) than magnesium’s outer 3s electron [1], so despite aluminium’s greater nuclear charge, its outer electron is easier to remove [1]. (c) Sulfur’s fourth 3p electron must pair up in an already-occupied 3p orbital, unlike phosphorus’s three unpaired 3p electrons (one in each orbital) [1]; the extra electron–electron repulsion in that paired orbital makes sulfur’s outer electron easier to remove, despite sulfur’s greater nuclear charge [1].

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Where marks are usually lost

• Saying the melting point drop at phosphorus is due to weaker covalent bonds.
• Comparing “molecule size” rather than number of electrons for induced dipole strength.
• Forgetting that Al₂O₃ is amphoteric.
• Not linking the oxide trend to the metal/non-metal change.
• Describing SiCl₄’s reaction with water as “dissolving” — it is a chemical reaction (hydrolysis), not a physical change of state.
• Explaining the Al and S ionisation-energy dips with vague phrases like “it’s an exception” instead of naming the specific electron configuration reason (subshell energy for Al, electron pairing for S).

Questions 7 and 8 draw on the chlorides-with-water table and the first-ionisation-energy trend in the Periodicity revision notes, material the earlier questions on this page don’t reach.