The Periodic Table

From Na to Cl the nuclear charge increases. Added electrons enter the third shell, so the shielding of these outer electrons changes relatively little. Stronger effective attraction generally draws the outer shell closer: atomic radius decreases, while first ionisation energy and electronegativity generally increase.

Explain the two first-ionisation-energy exceptions using orbitals. Al loses a higher-energy 3p electron rather than Mg's 3s electron. S loses a paired 3p electron, with extra electron-electron repulsion, whereas P has three singly occupied 3p orbitals. The resulting dips do not mean the general nuclear-charge trend is wrong.

Down Cl to I, the outer shell is farther from the nucleus and is more shielded. These effects outweigh the increased nuclear charge: atomic radius increases, first ionisation energy decreases and electronegativity decreases. The number of outer electrons stays seven.

For ionic radii, compare species with similar electron arrangements. Na⁺, Mg²⁺ and Al³⁺ each have ten electrons: increasing nuclear charge contracts this isoelectronic series. P^3-, S^2- and Cl^- each have eighteen; their radii also decrease in that order. Crossing from a cation series to an anion series introduces a large jump, so ionic radius does not form one smooth decreasing line across Period 3. Down Cl^-, Br^-, I^-, more occupied shells give larger ions.

The molecular comparison uses white phosphorus, P₄. S₈ has a larger, more easily polarised electron cloud than P₄, so sulfur has stronger intermolecular attractions and a higher melting point. Cl₂ is much smaller and melts much lower. Other allotropes can have different structures; an element name alone does not specify an allotrope.

Na, Mg and Al conduct electricity through mobile delocalised electrons. Conductivity generally increases from Na to Al: each atom supplies one, two and three valence electrons respectively, giving a greater density of mobile charge carriers. The ions remain in the solid lattice. Silicon's covalent network has far fewer mobile carriers at room temperature, so it is a semiconductor with much lower conductivity than the metals. P₄, S₈ and Cl₂ lack mobile charged particles and are electrical insulators in their pure forms.

Down the group, the molecules have more electrons and more polarisable electron clouds. Instantaneous dipole-induced dipole attractions become stronger, so more energy is needed to separate molecules and volatility decreases. This is an intermolecular explanation; do not use the X-X covalent bond strength to explain boiling.

In these highest oxides, oxygen has oxidation number -2. The element's highest oxidation number rises across the period: Na +1, Mg +2, Al +3, Si +4, P +5 and S +6. More outer electrons can participate in bonding across this sequence. With oxygen fixed, the electronegativity difference generally decreases, and bonding changes from predominantly ionic to covalent.

Na₂O and MgO are basic: their oxide ions accept protons. For example, MgO(s) + 2H⁺(aq) → Mg²⁺(aq) + H₂O(l). Their corresponding hydroxides neutralise acids: NaOH(aq) + H⁺(aq) → Na⁺(aq) + H₂O(l), and Mg(OH)₂(s) + 2H⁺(aq) → Mg²⁺(aq) + 2H₂O(l). Magnesium hydroxide is only sparingly soluble, so an excess solid does not make an arbitrarily concentrated solution.

Al₂O₃ and Al(OH)₃ are amphoteric: each reacts with acid and with strong alkali. The specified alkali is sodium hydroxide; heat the oxide with concentrated aqueous NaOH, whereas freshly precipitated Al(OH)₃ dissolves in excess aqueous NaOH. With acid, Al₂O₃ + 6H⁺ → 2Al³⁺ + 3H₂O; Al(OH)₃ + 3H⁺ → Al³⁺ + 3H₂O. In aqueous hydroxide, soluble aluminate species form.

SiO₂, P₄O₁₀ and SO₃ are acidic oxides and react with alkali. Insoluble SiO₂ does not need to react with water to be acidic: SiO₂ + 2NaOH → Na₂SiO₃ + H₂O with hot concentrated NaOH (or fused NaOH). With sufficient NaOH, P₄O₁₀ + 12NaOH → 4Na₃PO₄ + 6H₂O; SO₃ + 2NaOH → Na₂SO₄ + H₂O.

Across the series the highest element oxidation number rises with the available outer electrons. Increasing electronegativity of the Period 3 element generally reduces the difference from chlorine and favours covalent bonding. AlCl₃ is the stated exception to the simple electronegativity classification; the small, highly charged aluminium centre strongly polarises chloride electron density. Its molecular dimer Al₂Cl₆ is discussed in Bonding.

For NaCl, separating ions into water is principally dissolution. A hydrated Mg²⁺ or Al³⁺ ion attracts electron density from its water ligands, weakening O-H bonds and allowing proton release. The effect is stronger for Al³⁺, with greater charge density. It is the hydrated cation, not Cl^- simply splitting into HCl, that explains this acidity.

For example, [Al(H₂O)₆]³⁺ + H₂O ⇌ [Al(H₂O)₅(OH)]²⁺ + H₃O⁺. The equation makes proton release and charge conservation explicit. By contrast, SiCl₄ and PCl₅ undergo extensive hydrolysis of their covalent chlorides.

A reducing agent supplies electrons and is itself oxidised. Group 1 atoms have one outer electron: M → M⁺ + e^-. From lithium through sodium, potassium, rubidium and caesium, the outer electron is farther from the nucleus and more shielded. It is generally lost more easily, so the familiar chemical reactivity increases down the group.

A common reaction is 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g). The metal is oxidised and water is reduced. For this H1 comparison, explain the trend through ease of electron loss; do not import an H2 electrode-potential ranking, which includes additional energetic factors.

An oxidising agent accepts electrons and is itself reduced. X₂ + 2e^- → 2X^-. From chlorine to bromine to iodine, larger size and greater shielding weaken attraction for an incoming electron. Oxidising ability decreases in the order Cl₂ > Br₂ > I₂.

Thermal stability of the hydrogen halides is a different comparison. HCl is more thermally stable than HBr, which is more stable than HI. Down the group, the halogen atom becomes larger, the H-X bond longer and overlap poorer, so bond energy decreases. Less energy is needed to decompose HI: 2HI(g) ⇌ H₂(g) + I₂(g). This concerns covalent H-X bond breaking, not the intermolecular forces used to explain halogen volatility.

When inferring bonding from an oxide or chloride, combine melting behaviour, conductivity in different states and chemical reactions with water. High melting point does not uniquely establish ionic bonding: SiO₂ is a covalent network. An acidic chloride solution also does not uniquely establish molecular hydrolysis: hydrated metal ions can release protons.