Groups in the periodic table
Group 1 — alkali metals (Li, Na, K, Rb, Cs, Fr)
All Group 1 metals have 1 outer electron which they readily lose, forming M⁺ ions and reducing agents.
Physical properties (going down Group 1):
- Decreasing melting point / boiling point (weaker metallic bonds — outer electron further from nucleus)
- Decreasing hardness
- Increasing density
Reaction with water (increasing reactivity down the group): 2Na + 2H₂O → 2NaOH + H₂↑ 2K + 2H₂O → 2KOH + H₂↑ (burns lilac with a squeaky pop) Products: metal hydroxide (alkaline solution) + hydrogen gas.
Why reactivity increases down Group 1: outer electron is in a higher energy shell, further from the nucleus, with more shielding from inner electrons → weaker attraction → lost more easily.
Reaction with oxygen: 4Na + O₂ → 2Na₂O (tarnishes quickly — stored in oil to prevent oxidation). Reaction with chlorine: 2Na + Cl₂ → 2NaCl (white solid, ionic).
Group 7 — halogens (F, Cl, Br, I, At)
All Group 7 elements have 7 outer electrons and gain 1 electron to form X⁻ ions.
Physical states at room temperature: F₂ (yellow gas), Cl₂ (yellow-green gas), Br₂ (orange/brown liquid), I₂ (grey-black solid with purple vapour).
Reactivity decreases down the group: each successive halogen has its outer electron in a higher shell, further from the nucleus, making it harder to attract another electron.
Displacement reactions: a more reactive halogen displaces a less reactive halide from its aqueous solution: Cl₂ + 2KBr → 2KCl + Br₂ (chlorine displaces bromine — chlorine is more reactive) Cl₂ + 2KI → 2KCl + I₂ (iodine solution turns brown/violet)
Reaction with alkali metals (form ionic compounds — salts): 2Na + Cl₂ → 2NaCl; 2Fe + 3Cl₂ → 2FeCl₃.
Reaction with hydrogen (diatomic hydrogen): H₂ + Cl₂ → 2HCl (hydrogen chloride — acidic gas) Reactivity: F₂ reacts explosively at room temperature; I₂ requires UV light.
Group 0 — noble gases (He, Ne, Ar, Kr, Xe)
Full outer shells (He: 2; all others: 8) → extremely unreactive (inert). Monoatomic: exist as single atoms (He, Ne, Ar, etc.). Uses: Ar in light bulbs/welding (inert atmosphere); He in balloons (less dense than air, non-flammable); Ne in neon signs; Kr/Xe in specialist lamps.
Density and boiling point increase going down the group (more electrons → stronger London dispersion forces).
Transition metals
Located in the d-block (Groups 3–12). Key properties distinguishing them from Group 1/2:
- High melting points and densities (stronger metallic bonding)
- Hard and strong — useful as structural metals
- Variable oxidation states: Fe²⁺ (iron(II)) and Fe³⁺ (iron(III)); Cu⁺ and Cu²⁺
- Coloured compounds: FeSO₄ = green, Fe₂(SO₄)₃ = yellow, CuSO₄ = blue
- Catalytic activity: Fe (Haber process), Pt/Pd (catalytic converters), MnO₂ (H₂O₂ decomposition), Ni (hydrogenation of oils)
⚠Common mistakes
- Halogen reactivity decreases down Group 7 (opposite to Group 1). Don't mix up the trends.
- Noble gases do not form simple ions: they have full shells and do not lose or gain electrons in standard conditions.
- Transition metals vs Group 1: Group 1 have low densities, soft, low m.p.; transition metals are dense, hard, high m.p.
- Displacement test colour change: Cl₂ + KI → iodine produced (brown/violet in hexane), not blue — blue is the iodine/starch test.
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