Why the elements line up the way they do — and how their arrangement lets you predict size, reactivity and the trends across a period.
Theme · StructureStructure 3.1
This is the Standard Level lesson, covering the core of S3.1. Taking Chemistry at Higher Level? The HL version adds the discontinuities in ionization energy and the transition elements (variable oxidation states, coloured complexes).
👆 Switch the trend in the explorer · read the period-3 pattern
1. How the table is built S3.1.1
Elements are arranged by increasing atomic number. A period (row) is the highest occupied energy level; a group (column) collects elements with the same number of valence electrons — which is why a group behaves alike. The table splits into blocks named after the sub-level being filled: s, p, d and f.
The blocks of the periodic table
The block tells you which sub-level is being filled — e.g. group-1 and 2 metals are the s-block; the halogens and noble gases are in the p-block.
2. Periodicity — the trends S3.1.3
Because properties depend on electron arrangement, they repeat periodically. Pick a trend and read it across period 3 (Na → Ar):
Across a period (left → right): atomic radius decreases; ionisation energy and electronegativity increase — because the nuclear charge rises while electrons enter the same shell, pulling everything in tighter.
Down a group the trends reverse for size (atoms get bigger — more shells) and for ionisation energy/electronegativity (they decrease — the outer electrons are further from the nucleus and shielded).
3. Metallic character & oxides S3.1.4
Across period 3 the elements change from metals (Na, Mg, Al) through a metalloid (Si) to non-metals (P, S, Cl, Ar). Metallic character decreases across a period and increases down a group. This shows up in their oxides — they shift from basic to amphoteric to acidic:
Na₂O, MgO basic
Al₂O₃ amphoteric
SiO₂ · P₄O₁₀ · SO₃ · Cl₂O₇ acidic
Metal oxides (left) react with acids and are basic; non-metal oxides (right) react with water to form acids — the cause of acid rain (SO₃, NO₂) and ocean acidification (CO₂).
Oxidation states (S3.1.5) are a bookkeeping number showing how many electrons an atom has gained or lost. Useful rules: elements = 0; group 1 = +1, group 2 = +2; oxygen is usually −2 (but −1 in peroxides); hydrogen is +1 (but −1 in metal hydrides); the sum across a neutral compound = 0.
Going to HL? The HL lesson adds the small dips in the ionisation-energy trend (evidence for sub-levels) and the transition elements: variable oxidation states, catalysis, magnetism and the coloured complexes formed by d-orbital splitting. → Open the HL version
Common mistakes examiners see
What happens to atomic radius across a period?✗ It increases (more electrons = bigger).✓ It decreases — the electrons go into the same shell but the nuclear charge rises, pulling them in more tightly.
Why does ionisation energy increase across a period?✗ Because atoms get bigger.✓ Greater nuclear charge and a smaller radius (electrons in the same shell) hold the outer electron more tightly, so more energy is needed to remove it.
What determines an element's group?✗ Its mass number.✓ Its number of valence (outer-shell) electrons — which is why a group reacts similarly.
Are non-metal oxides basic or acidic?✗ Basic.✓ Acidic — e.g. SO₃ and CO₂ dissolve in water to form acids. Metal oxides (Na₂O, MgO) are basic.