The Periodic Table

Order: increasing proton number. Hydrogen ($Z = 1$) at top-left; oganesson ($Z = 118$) at bottom-right.

Row = period. Each new row starts a new electron shell.

• Period 1: H, He.
• Period 2: Li to Ne.
• Period 3: Na to Ar.

Column = group. Same group → same number of outer-shell electrons → similar chemistry.

• Group 1: alkali metals (1 outer e⁻).
• Group 2: alkaline earth metals (2 outer e⁻).
• Group 17: halogens (7 outer e⁻).
• Group 18: noble gases (8 outer e⁻ — full shell).

Modern numbering uses 1-18. Older/short-form numbering used I-VIII (e.g. Group I = Group 1; Group VII = Group 17). Cambridge accepts both but prefers numbers 1-18.

Worked qualitative. Calcium ($Z = 20$, configuration $2, 8, 8, 2$). Period 4 (4 shells), Group 2 (2 outer electrons).

Why arrange this way? Mendeleev (1869) observed PERIODIC repetition of properties. Modern understanding: same outer-shell configuration = same chemistry → naturally falls into groups.

Metals (left side and middle). Most elements. Properties:

• Solid (except mercury) at room temp.
• Shiny when polished.
• Good conductors of heat and electricity.
• Malleable and ductile.
• Form positive ions (lose electrons).
• Form basic oxides.
• High m.p. and density (mostly).

Non-metals (top right). Limited number. Properties:

• Often gas or low-m.p. solid (some liquid, e.g. Br₂).
• Dull (when solid).
• Poor conductors (insulators or semiconductors).
• Brittle (when solid).
• Form negative ions (gain electrons), or share electrons covalently.
• Form acidic oxides.

Metalloids (the staircase). Boron, silicon, germanium, arsenic, antimony, tellurium. Properties between metals and non-metals.

Worked qualitative. Copper is a typical metal: silvery-orange shine, conducts electricity, ductile (drawn into wires), forms Cu²⁺ ion, CuO is a basic oxide.

Transition metals sit in the middle of the table (Groups 3-12 in modern numbering). Examples: titanium, vanadium, chromium, manganese, iron, cobalt, nickel, copper, zinc.

Distinctive properties.

1. Variable oxidation states. Iron can be Fe²⁺ or Fe³⁺. Copper can be Cu⁺ or Cu²⁺. Cambridge usually quotes the charge in the question, e.g. iron(II) chloride = FeCl₂.

2. Coloured compounds. Cu²⁺ blue, Fe²⁺ green, Fe³⁺ orange/brown, Mn²⁺ pale pink, Cr³⁺ green. Useful for identification.

3. Catalytic activity. Iron in Haber, V₂O₅ in Contact, nickel in hydrogenation, platinum in catalytic converters.

4. High density and m.p. Most transition metals are denser and have higher m.p. than Group 1/2 metals.

5. Form complex ions. Beyond IGCSE — but a hint: the colour comes from electronic transitions in d-orbitals.

Comparison: Group 1 vs transition metals.

Property	Group 1 (alkali metals)	Transition metals
Density	Low	High
Hardness	Soft (cut with knife)	Hard
m.p.	Low (Cs melts in hand)	High
Reactivity	Very reactive	Less reactive
Compounds	Usually colourless / white	Often coloured
Oxidation states	Always $+1$	Variable
Catalytic?	Rarely	Often

Across a period (left to right).

• Atomic radius: DECREASES (more protons pulling same number of shells closer).
• Metallic character: DECREASES (going from metal to non-metal).
• Reactivity of metals: DECREASES (Na > Mg > Al across period 3).
• Reactivity of non-metals: INCREASES (P < S < Cl across period 3).

Down a group.

• Atomic radius: INCREASES (more shells).
• Metallic character: INCREASES.
• Group 1 reactivity: INCREASES (Li < Na < K < Rb < Cs).
• Group 17 reactivity: DECREASES (F > Cl > Br > I).

Why opposite trends for metals vs non-metals?

• Metals lose electrons: easier when outer e⁻ is far from nucleus (down the group) or when the atom is bigger.
• Non-metals gain electrons: easier when the nucleus pulls strongly on incoming e⁻ (smaller atom, top of group).

(For full detail, see Periodic Trends and Group Properties study notes.)