This is the Higher Level lesson. The core of S3.1 (structure, trends, oxides) is shared with SL; the HL-only content — the discontinuities in ionisation energy and the transition elements — carries a purple HL badge.
Core — SL & HL
HL AHL only
👆 trend explorer · IE dips · colour wheel
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.
4. Discontinuities in ionisation energy HL
Ionisation energy generally rises across a period — but not smoothly. In period 3 there are small dips at aluminium and sulfur, and they are direct evidence for sub-levels. Click the red bars:
First ionisation energy across period 3 · click the dips
Click a red bar (Al or S)These dips break the general upward trend — each has a specific electron-arrangement reason.
5. The transition elements HL
The d-block transition elements share a set of properties, all stemming from partly-filled d-orbitals whose 4s and 3d energies are very close:
Variable oxidation states4s and 3d are close in energy, so different numbers of electrons can be lost — e.g. Fe²⁺/Fe³⁺, Cu⁺/Cu²⁺, Mn up to +7.
CatalysisIron in the Haber process, V₂O₅ in the Contact process, nickel in hydrogenation — variable oxidation states let them provide reaction pathways.
MagnetismUnpaired d-electrons make many transition-metal species paramagnetic.
Coloured compoundsTheir complexes absorb visible light — the colour depends on the metal, its oxidation state, the ligand and the coordination number.
Why complexes are coloured — the colour wheel HL
When ligands surround a transition-metal ion the d-orbitals split into two energy levels. The gap matches the energy of visible light, so the complex absorbs one colour and we see the complementary colour (the one opposite it on the wheel). Click a colour to see what is observed if that colour is absorbed:
Complementary colour wheel · click the absorbed colour
Click a colourA complex that absorbs a colour appears as the colour opposite it on the wheel.
Example: aqueous Cu²⁺ absorbs orange–red light, so it appears blue.
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.
Why does the first ionisation energy dip at aluminium?✗ Aluminium is bigger than magnesium.✓ Al's outer electron is removed from a 3p sub-level, which is higher in energy than Mg's 3s — so it is easier to remove. Evidence that sub-levels exist. HL
If a complex absorbs red light, what colour does it appear?✗ Red.✓ Green — we see the complementary colour (opposite on the wheel), not the absorbed one. HL