KCSE 2025 Chemistry P2 Q1 — Periodicity Across Period 3 (Radius, Ionisation, Oxides)
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The Question
“This question is about the elements of Period 3. (a) Explain why, across the period from magnesium to chlorine, (i) the atomic radius decreases and (ii) the first ionisation energy increases. (b) In terms of structure and bonding, explain why (i) aluminium has a higher melting point than magnesium and (ii) phosphorus has a higher melting point than chlorine. (c) Write the formulae of the oxides of Mg, Al, Si and P and state whether each is acidic, basic or amphoteric. (d) Sulphur forms the chloride S₂Cl₂. State (i) the type of bond in the molecule and (ii) the oxidation number of sulphur. (e) Write a balanced equation for the fuming of silicon(IV) chloride, SiCl₄, in moist air. (f) (i) Define a metalloid and (ii) give one use of silicon.”
(a)(i) Why the atomic radius decreases
Going from magnesium to chlorine, each successive atom gains one more proton in the nucleus, so the nuclear charge rises from 12 up to 17. The extra electrons are added to the same outer shell (the third energy level), which does not increase the shielding much. With a stronger nuclear pull on electrons that are essentially the same distance out, the outer shell is drawn in closer, so the atomic radius decreases across the period.
(a)(ii) Why the first ionisation energy increases
Ionisation energy is the energy needed to remove the outermost electron. It follows directly from the radius trend: as the atoms get smaller and the nuclear charge gets larger, the outer electron is held more tightly to the nucleus. A more tightly held electron takes more energy to pull away, so the first ionisation energy increases across the period.
(b)(i) Why aluminium melts higher than magnesium
Both are metals with giant metallic structures — a lattice of positive ions in a sea of delocalised electrons. The difference is how many electrons each atom donates: magnesium releases 2 outer electrons per atom, aluminium releases 3. More delocalised electrons and a higher ionic charge mean a stronger electrostatic attraction between the ions and the electron sea, so aluminium has stronger metallic bonds and a higher melting point.
(b)(ii) Why phosphorus melts higher than chlorine
Phosphorus and chlorine are simple molecular substances, so melting them only breaks the weak van der Waals forces between molecules, not the strong covalent bonds inside them. Phosphorus exists as large P₄ molecules while chlorine is a small Cl₂ molecule. The larger, heavier P₄ molecule has stronger van der Waals forces between molecules than Cl₂, so phosphorus melts at a higher temperature.
(c) Formulae and acid–base character of the oxides
Across the period the oxides change from basic (metals on the left) through amphoteric to acidic (non-metals on the right). Magnesium oxide is basic, aluminium oxide is amphoteric (it reacts with both acids and bases), and silicon and phosphorus oxides are acidic.
(d)(i) Type of bond in S₂Cl₂
S₂Cl₂ is made only of non-metal atoms (sulphur and chlorine), which share electrons rather than transfer them. The bonding in the molecule is therefore covalent.
(d)(ii) Oxidation number of sulphur
Let the oxidation number of sulphur be x. Chlorine in a chloride is −1, and there are two chlorine atoms; the molecule is neutral, so the oxidation numbers must sum to zero. Solving for x gives the oxidation number of each sulphur atom.
(e) Balanced equation for SiCl₄ fuming in moist air
Silicon(IV) chloride fumes because it reacts with water vapour in the air. It hydrolyses to solid silicon(IV) oxide and hydrogen chloride gas (the white fumes). Check the balance: one Si, four Cl, four H and two O on each side.
(f) Metalloid and a use of silicon
A metalloid is an element that shows both metallic and non-metallic properties. Silicon is the classic example, and its most important use is as a semiconductor — in computer chips, transistors and solar cells.
Final Result
Across Period 3 the atomic radius decreases and the first ionisation energy increases (rising nuclear charge on the same shell). Al melts higher than Mg (3 vs 2 delocalised electrons); P melts higher than Cl (larger P₄ molecule, stronger van der Waals forces). Oxides: MgO basic, Al₂O₃ amphoteric, SiO₂ and P₄O₁₀ acidic. In S₂Cl₂ the bond is covalent and sulphur is +1. SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g). A metalloid has both metallic and non-metallic properties; silicon is used in semiconductors.
Why this method works
Almost the whole question flows from one idea: across a period the nuclear charge rises while electrons fill the same shell, so the atoms pull their outer electrons in harder. That single fact explains the shrinking radius and the rising ionisation energy directly, and it underpins the shift in the oxides from basic metals to acidic non-metals. The melting-point parts test a different skill — recognising the type of structure first. Metals (Mg, Al) are compared by counting delocalised electrons, whereas simple molecules (P₄, Cl₂) are compared by the strength of the weak intermolecular forces, which grows with molecular size. Keeping 'what structure is this?' as the first question is what stops students from wrongly comparing covalent bond strengths when only van der Waals forces are melting.
The oxidation-number check is self-verifying: +1 for each sulphur and −1 for each chlorine gives 2(+1) + 2(−1) = 0, matching the neutral molecule.