The Periodic Table

Arrangement of elements

The periodic table arranges elements in order of increasing atomic number (Z). Elements with similar chemical properties appear in the same vertical column, called a group. Horizontal rows are called periods.

• Period: elements in the same period have the same number of electron shells.
• Group: elements in the same group have the same number of outer-shell electrons, giving similar chemical properties. Group number = number of outer electrons (for groups 1–7).
• Noble gases (Group 0/18): full outer shells, very unreactive.

Periodic trends

Across a period (left to right):

• Atomic number increases.
• Metallic character decreases (metals on left, non-metals on right).
• Reactivity of metals decreases; reactivity of non-metals increases.

Down a group:

• Atomic radius increases (more electron shells).
• Reactivity of Group 1 metals increases (outer electron further from nucleus, less energy to remove).
• Reactivity of Group 7 non-metals decreases (harder to gain an electron into the larger outer shell).

Group 1 — Alkali metals (Li, Na, K, Rb, Cs)

• Soft metals with low melting points.
• React vigorously with water to produce a metal hydroxide + hydrogen gas. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
• React with oxygen to form metal oxides (tarnish rapidly — stored under oil).
• React with halogens to form ionic salts (e.g. 2Na + Cl₂ → 2NaCl).
• Reactivity increases down the group (Li < Na < K < Rb < Cs).

Group 7 — Halogens (F, Cl, Br, I, At)

• Diatomic non-metals (F₂, Cl₂, Br₂, I₂).
• React with metals to form ionic salts (metal halides).
• React with hydrogen to form hydrogen halides (HF, HCl, HBr, HI).
• Displacement reactions: a more reactive halogen displaces a less reactive one from its salt solution. Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
• Reactivity decreases down the group (F > Cl > Br > I): harder to gain an electron into a larger shell.
• Colour and state at room temperature: F₂ (pale yellow, gas), Cl₂ (yellow-green, gas), Br₂ (orange-brown, liquid), I₂ (grey/purple, solid).

Group 0 — Noble gases

• Full outer shells: He (2), Ne, Ar, Kr, Xe (all 8 outer electrons).
• Monatomic, colourless gases at room temperature.
• Very low reactivity (no need to gain or lose electrons).
• Uses: Ar in light bulbs (inert atmosphere); He in balloons; Ne in neon signs.

Transition metals (d-block)

Located between groups 2 and 3 in periods 4 and 5. Properties distinct from Group 1 and 2 metals:

• Hard, dense, high melting points (except Hg).
• Form coloured compounds: Fe²⁺ (green), Fe³⁺ (orange/brown), Cu²⁺ (blue).
• Form ions with variable oxidation states: Fe²⁺ and Fe³⁺; Cu⁺ and Cu²⁺.
• Act as catalysts: Fe in Haber process (N₂ + 3H₂ ⇌ 2NH₃); Mn in decomposition of H₂O₂.

CCEA exam tips

• Halogen displacement: use a colour change to identify if displacement has occurred (e.g. colourless KBr + Cl₂ → orange/brown Br₂ formed).
• Know the trend words: Group 1 reactivity increases down; Group 7 reactivity decreases down.
• Transition metals always appear with coloured ions and catalytic behaviour in CCEA questions.