S3.1 The periodic table: Classification of elementsIB Chemistry HL: Revision notes

The table has periods (outer main energy level), groups (same number of valence electrons) and s, p, d and f blocks. Metals are on the left, non-metals on the right, and metalloids along the boundary. Deduce configurations up to Z = 36 from position, and vice versa.

Across a period, nuclear charge rises with similar shielding: atomic radius decreases while ionisation energy, electron affinity and electronegativity increase. Down a group, extra levels and shielding outweigh the rising nuclear charge, so the trends reverse. Cations are smaller and anions larger than their atoms; in isoelectronic series, radius falls as nuclear charge rises. Group 1 reactivity with water increases down the group; group 17 reactivity (displacement of halides) decreases.

Period 3 oxides show a continuum: basic (Na₂O, MgO) → amphoteric (Al₂O₃) → acidic (SiO₂, P₄O₁₀, SO₃). Large electronegativity differences give ionic oxides containing O²⁻, which form OH⁻ in water; small differences give covalent oxides that form acids.

• Na₂O + H₂O → 2NaOH; MgO + H₂O → Mg(OH)₂
• CO₂ + H₂O ⇌ H₂CO₃; SO₂ + H₂O → H₂SO₃; SO₃ + H₂O → H₂SO₄

The oxidation state is the charge an atom would have if the compound were fully ionic. In a complex ion, subtract the ligand charges from the overall charge: in [CoCl₄]²⁻, Co is +2.

First ionisation energy rises across a period but with two dips, which are evidence for sublevels. Explain them by the energy of the electron removed, not by the "special stability" of full or half-full sublevels.

• Group 2 → 13 (Be → B, Mg → Al): the electron is removed from a p sublevel, which is higher in energy than the s sublevel (partly shielded by s electrons), so less energy is needed.
• Group 15 → 16 (N → O, P → S): the electron is removed from a doubly occupied p orbital; repulsion from its partner raises its energy, so it is easier to remove.

The pattern 2, 3, 3 across each period matches one s orbital and three p orbitals.

A transition element has an incomplete d-sublevel in the atom or one of its common ions (Sc to Cu in period 4; Zn²⁺ is 3d¹⁰, so zinc is not a transition element). Characteristic properties:

• variable oxidation states (Mn: +2, +4, +7)
• high melting points (strong metallic bonding involving d electrons)
• magnetic properties: unpaired d electrons make species paramagnetic
• catalytic properties (Fe in the Haber process, MnO₂ with H₂O₂)
• coloured compounds
• complex ions, in which ligands donate lone pairs to the central ion by coordination bonds.

Variable oxidation states arise because the successive ionisation energies of the 4s and 3d electrons are close in value, with no large jump, so the energy cost of losing extra electrons can be recovered from bonding.

When first-row transition elements form ions, the 4s electrons are removed before the 3d electrons. Examples: Fe [Ar]3d⁶4s² → Fe²⁺ [Ar]3d⁶ → Fe³⁺ [Ar]3d⁵; Co²⁺ [Ar]3d⁷; Cu [Ar]3d¹⁰4s¹ → Cu²⁺ [Ar]3d⁹; Zn²⁺ [Ar]3d¹⁰; Mn²⁺ [Ar]3d⁵ (five unpaired electrons, strongly paramagnetic).

In a complex, the ligands split the d-sublevel into two sets of orbitals of slightly different energy. When an electron is promoted from the lower to the higher set it absorbs a photon of visible light; the energy gap fixes the frequency absorbed. The colour observed is the complementary colour, opposite the absorbed colour on the colour wheel in the data booklet (red–green, orange–blue, yellow–violet). A complex absorbing at 610 nm (orange) looks blue; f = c/λ = 4.92 × 10¹⁴ Hz. Ions with empty (3d⁰) or full (3d¹⁰) d-sublevels (Sc³⁺, Zn²⁺) cannot undergo d–d promotion, so they are colourless.