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

The periodic table is arranged in periods (rows) and groups (columns, numbered 1–18). The period number equals the outer main energy level occupied by electrons: sodium, [Ne]3s¹, is in period 3. Elements in the same group have the same number of valence electrons, which is why they react similarly. The table also divides into four blocks named after the sublevel being filled: s-block (groups 1–2), p-block (groups 13–18), d-block (groups 3–12) and f-block (lanthanoids and actinoids). Metals lie to the left, non-metals to the right, and the metalloids (such as B, Si, Ge, As, Sb, Te) sit along the staircase boundary between them.

Read the configuration directly from the table: move along the periods, filling s in groups 1–2, d in groups 3–12 (one level behind: 3d in period 4) and p in groups 13–18. So group 15, period 4 means [Ar]3d¹⁰4s²4p³ (arsenic), and the third d-block element in period 4 is [Ar]3d³4s² (vanadium). Work backwards too: 1s²2s²2p⁶3s²3p⁴ has outer level n = 3 (period 3) and six outer electrons (group 16), so it is sulfur. Remember the exceptions Cr = [Ar]3d⁵4s¹ and Cu = [Ar]3d¹⁰4s¹.

Periodicity is the repeating pattern of properties across periods and down groups. Every trend is explained by three factors: nuclear charge, distance of the outer electrons from the nucleus and shielding by inner electrons.

Across a period: nuclear charge rises, shielding stays similar (same main level), so outer electrons are held more tightly. Atomic radius decreases; first ionisation energy, electron affinity (energy released on gaining an electron, generally more exothermic) and electronegativity all increase.

Down a group: outer electrons are in higher main levels, further away and more shielded, which outweighs the larger nuclear charge. Atomic radius increases; ionisation energy and electronegativity decrease; electron affinity generally becomes less exothermic.

Cations are smaller than their atoms: losing the outer electrons usually empties a whole main level (Na 186 pm → Na⁺ 102 pm). Anions are larger than their atoms: extra electrons increase repulsion with the same nuclear charge (Cl 99 pm → Cl⁻ 181 pm). In an isoelectronic series (same number of electrons), radius falls as nuclear charge rises: P³⁻ > S²⁻ > Cl⁻ and Na⁺ > Mg²⁺ > Al³⁺.

Group 1 metals react with water to form an alkaline hydroxide and hydrogen: 2Na + 2H₂O → 2NaOH + H₂. Reactivity increases down the group (Li fizzes, Na melts, K ignites) because the single outer electron is further from the nucleus and more shielded, so it is lost more easily: metallic character increases down the group.

Group 17 halogens gain an electron to form halide ions. A more reactive halogen displaces a less reactive one from its halide: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂ (orange/yellow solution); Br₂ + 2I⁻ → 2Br⁻ + I₂ (brown solution). Reactivity decreases down the group because the incoming electron is further from the nucleus and more shielded: non-metallic character decreases down the group.

Metallic and non-metallic properties form a continuum, shown by the oxides of period 3: basic metal oxides (Na₂O, MgO), amphoteric Al₂O₃ (reacts with acids and bases), then acidic non-metal oxides (SiO₂, P₄O₁₀, SO₂, SO₃).

Equations with water:

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

An oxidation state shows the number of electrons an atom has transferred in forming bonds: it is the charge the atom would have if the compound were fully ionic. Rules: elements = 0; the sum equals the charge on the species; O is usually −2 and H usually +1; the more electronegative atom takes the negative value. Examples: S is +4 in H₂SO₃, +6 in H₂SO₄ and SO₃, −2 in H₂S; Cl is +5 in ClO₃⁻. Name compounds with Roman numerals where needed: iron(III) oxide, chlorate(V).