IB DP Chemistry Structure 3.1: The periodic table: classification of elements

This page teaches sub-topic Structure 3.1, The periodic table: classification of elements, as set by the Diploma Programme Chemistry guide from the IB (February 2023 edition, first examined in 2025). It covers understandings Structure 3.1.1 to 3.1.10. Understandings 3.1.1 to 3.1.6 are for SL and HL; 3.1.7 to 3.1.10 are additional higher level (AHL), so HL only. It applies to the course from its first assessment in 2025 onwards.

Check what has stuck using the revision notes, then try the practice questions. For weak spots elsewhere, try the free diagnostics.

What this unit covers

Understanding	Skill	Level
Structure 3.1.1	Place metals, metalloids, non-metals; s, p, d, f blocks	SL and HL
Structure 3.1.2	Position to configuration and back, Z ≤ 36	SL and HL
Structure 3.1.3	Explain five periodic trends	SL and HL
Structure 3.1.4	Group 1 with water; group 17 with halide ions	SL and HL
Structure 3.1.5	Oxides with water; acid rain; ocean acidification	SL and HL
Structure 3.1.6	Deduce oxidation states; name oxyanions	SL and HL
Structure 3.1.7	Ionization energy dips as evidence for sublevels	HL only
Structure 3.1.8	Transition element properties	HL only
Structure 3.1.9	Variable oxidation states; ion configurations	HL only
Structure 3.1.10	Colour, the colour wheel and c = λf	HL only

Periods, groups and blocks (3.1.1)

Rows are periods; columns are groups, numbered 1 to 18. The four blocks are named after the sublevel the highest-energy electron enters:

• s-block: groups 1 and 2 (plus helium)
• p-block: groups 13 to 18
• d-block: groups 3 to 12
• f-block: the lanthanoids and actinoids, usually printed below the main table

Metals fill the left and centre. Non-metals sit upper right, plus hydrogen. Between them runs a staircase of metalloids, usually boron, silicon, germanium, arsenic, antimony and tellurium, with properties between the two. The guide notes that the data booklet includes a periodic table.

Position and electron configuration (3.1.2)

The period number is the outer occupied main energy level. Every element in a group has the same number of valence electrons. In groups 1 and 2 this equals the group number; in groups 13 to 18 take 10 away from the group number (helium, with 2, is the exception).

Know these names: alkali metals (group 1, not hydrogen), halogens (group 17), noble gases (group 18) and transition elements (in the d-block; see 3.1.8).

Worked example 1: from position to configuration

Germanium is in period 4, group 14, in the p-block.

Period 4 -> outer level is n = 4
Group 14 -> 14 - 10 = 4 valence electrons: 4s² 4p²
Filling order up to it: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p²
Condensed: [Ar] 3d¹⁰ 4s² 4p²   (check: 18 + 10 + 2 + 2 = 32 = Z)

Worked example 2: from configuration to position

An atom has the configuration [Ne] 3s² 3p³.

Highest level occupied: n = 3 -> period 3
Last electron enters 3p -> p-block
Valence electrons: 2 + 3 = 5 -> group 15
Z = 10 + 5 = 15 -> phosphorus

Learn the two exceptions up to Z = 36: chromium [Ar] 3d⁵ 4s¹ and copper [Ar] 3d¹⁰ 4s¹.

Periodicity: trends across and down (3.1.3)

Periodicity means trends in properties across a period and down a group. Two ideas explain them:

• Across a period, nuclear charge rises by one each step, but new electrons enter the same main level, so shielding barely changes. Outer electrons feel a stronger pull.
• Down a group, each element has one more occupied main level. Outer electrons are further away and more shielded, which outweighs the extra nuclear charge.

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
First ionization energy	Increases (with dips)	Decreases
Electron affinity	Generally more exothermic	Generally less exothermic
Electronegativity	Increases	Decreases

Electron affinity is the energy change, per mole, when isolated gaseous atoms each gain an electron and become gaseous 1− ions. Electronegativity is an atom’s ability to attract the shared pair in a covalent bond; fluorine’s is highest.

Ionic radius:

• A positive ion is smaller than its atom: it has lost its outer main level.
• A negative ion is larger than its atom: extra electrons add repulsion with no extra nuclear charge.
• In an isoelectronic series (same electron count), the ion with the highest nuclear charge is smallest.

Worked example 3: ranking ions

Rank S²⁻, Cl⁻, K⁺ and Ca²⁺ by radius. Each has 18 electrons; nuclear charges are 16, 17, 19 and 20. More protons pull the same electrons closer: S²⁻ > Cl⁻ > K⁺ > Ca²⁺.

Group 1 and group 17 (3.1.4)

Metallic character increases down group 1. Non-metallic character decreases down group 17.

Group 1 with water. Every alkali metal reacts to give a hydroxide and hydrogen:

2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)

The solution turns alkaline. Lithium fizzes steadily; sodium melts into a moving ball; potassium ignites the hydrogen with a lilac flame. Reactivity rises down the group: with each extra main level the outer electron sits further out and has more inner electrons between it and the nucleus, so it is lost more easily.

Group 17 with halide ions. A halogen takes electrons from the halide ion of any halogen below it, displacing it:

Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

In water, chlorine is pale green, bromine orange and iodine brown. Reactivity falls down the group: larger atoms attract an incoming electron less strongly, because it is further from the nucleus and more shielded. So iodine gives no reaction with bromide ions.

Oxides: from basic to acidic (3.1.5)

Properties form a continuum across a period: basic metal oxides, then amphoteric oxides such as Al₂O₃ (reacting with acids and bases), then acidic non-metal oxides.

Oxide	Equation with water	Product
Na₂O (group 1)	Na₂O(s) + H₂O(l) → 2NaOH(aq)	Alkaline solution
CaO (group 2)	CaO(s) + H₂O(l) → Ca(OH)₂(aq)	Alkaline solution
CO₂	CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)	Weakly acidic
SO₂	SO₂(g) + H₂O(l) → H₂SO₃(aq)	Acidic
SO₃	SO₃(g) + H₂O(l) → H₂SO₄(aq)	Strongly acidic

Acid rain. Gaseous non-metal oxides from burning fuels dissolve in rain. SO₂ gives sulfurous acid, or is oxidized to SO₃, giving sulfuric acid. Nitrogen dioxide: 2NO₂(g) + H₂O(l) → HNO₃(aq) + HNO₂(aq).

Ocean acidification. As atmospheric CO₂ rises, more dissolves in seawater. Carbonic acid releases hydrogen ions, H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq), so the pH falls. The extra H⁺ also reacts with carbonate ions, H⁺(aq) + CO₃²⁻(aq) → HCO₃⁻(aq), leaving less carbonate for shell-building organisms.

Worked example 4

Strontium oxide is a group 2 oxide, so basic: SrO(s) + H₂O(l) → Sr(OH)₂(aq). Potassium oxide is basic too: K₂O(s) + H₂O(l) → 2KOH(aq).

Oxidation states (3.1.6)

Imagine every bond in a compound were fully ionic: the charge an atom would then carry is its oxidation state. Write the sign first: +2, −1. “Oxidation number” is an accepted alternative term.

Rules, highest priority first:

1. An uncombined element is 0: its atoms bond only to identical atoms, so no electrons transfer.
2. Oxidation states sum to 0 in a compound, or to the charge in an ion.
3. Group 1 is +1 and group 2 is +2. Fluorine is always −1.
4. Hydrogen is +1, but −1 when bonded to a metal in a metal hydride (NaH, for example).
5. Oxygen is −2, except in peroxides such as H₂O₂, where it is −1.

Naming oxyanions. Roman numerals give the central atom’s oxidation state: ClO₃⁻ is chlorate(V). Generic names are also accepted: NO₃⁻ nitrate, NO₂⁻ nitrite, SO₄²⁻ sulfate, SO₃²⁻ sulfite.

Worked example 5

Fe in FeO₄²⁻:  x + 4(−2) = −2  -> x = +6   (ferrate(VI))
Al in LiAlH₄:  (+1) + x + 4(−1) = 0 -> x = +3   (H is −1, a metal hydride)

HL only: discontinuities in first ionization energy (3.1.7)

First ionization energy rises across period 3 but dips twice. The dips are evidence for sublevels. Explain them by the energy of the electron removed, not “special stability”.

• Magnesium to aluminium. Aluminium loses a 3p electron, which is higher in energy than magnesium’s 3s electron, so less energy is needed.
• Phosphorus to sulfur. In sulfur two 3p electrons share an orbital. Their repulsion raises the energy of one, so it is removed more easily than any of phosphorus’s unpaired 3p electrons.

If every electron in a main level had equal energy, the rise would be smooth.

HL only: transition elements (3.1.8)

An incomplete d-sublevel gives transition elements these properties:

• variable oxidation states
• high melting points (both 3d and 4s electrons contribute to metallic bonding)
• magnetic properties, from unpaired d electrons (the types of magnetism are not assessed)
• catalytic properties, for example iron in ammonia synthesis
• coloured compounds
• complex ions, in which ligands donate lone pairs to the central ion

Zinc is not a transition element: the atom and Zn²⁺ both have a full 3d sublevel. Scandium is debated, because Sc³⁺ has no d electrons.

HL only: variable oxidation states and ion configurations (3.1.9)

The 4s and 3d sublevels are close in energy, so successive ionization energies rise gradually, and several different numbers of electrons can be lost.

Worked example 6: iron and strontium

Successive ionization energies in kJ mol⁻¹ (use these values):

Fe: 762, 1561, 2957, 5290
Sr: 549, 1064, 4138

Iron’s increases (799, 1396, 2333) show no huge jump, so Fe²⁺ and Fe³⁺ both form. Strontium’s third value is almost four times its second, as that electron comes from an inner main level: only Sr²⁺ forms.

Ion configurations. Remove the 4s electrons first, then 3d.

Fe³⁺: Fe is [Ar] 3d⁶ 4s² -> remove 4s², then one 3d -> [Ar] 3d⁵
V³⁺:  V is [Ar] 3d³ 4s²  -> remove 4s², then one 3d -> [Ar] 3d²
Cu⁺:  Cu is [Ar] 3d¹⁰ 4s¹ -> remove 4s¹ -> [Ar] 3d¹⁰

HL only: why complexes are coloured (3.1.10)

Ligands split the d-sublevel into two sets of orbitals. A d electron absorbs a photon matching the gap and is promoted to the higher set, removing that frequency from white light. The colour seen is complementary to the colour absorbed, opposite it on the colour wheel. Splitting patterns are not required. The guide says the data booklet supplies both the colour wheel and c = λf.

A simple colour wheel (approximate ranges; use the one you are given):

Colour	Wavelength / nm	Complementary colour
Violet	400–450	Yellow
Blue	450–490	Orange
Green	490–560	Red
Yellow	560–590	Violet
Orange	590–630	Blue
Red	630–700	Green

Worked example 7

A complex ion absorbs most strongly at 505 nm. Deduce the colour seen and the frequency absorbed (c = 3.00 × 10⁸ m s⁻¹).

505 nm lies in the green range -> colour seen is red
f = c / λ = (3.00 × 10⁸) / (505 × 10⁻⁹) = 5.94 × 10¹⁴ s⁻¹

A d¹⁰ ion (Zn²⁺) or d⁰ ion (Sc³⁺) cannot undergo this promotion, so its compounds are usually colourless or white.

Common errors

• Counting valence electrons for groups 13 to 18 as the group number, not the group number minus 10.
• Writing CO₂ + H₂O → H₂CO₃ as a complete, one-way reaction; it is an equilibrium.
• Giving oxygen −2 in a peroxide, or hydrogen +1 in a metal hydride.
• (HL) Removing 3d electrons before 4s when writing ion configurations.
• (HL) Giving the absorbed colour as the colour seen.

Next steps

• Free diagnostics across the course
• Structure 3.1 revision notes and practice questions
• Course overview on the chemistry hub; tick off progress on the checklist

Official syllabus

Content here maps to the International Baccalaureate Organization’s Diploma Programme Chemistry guide, published in February 2023 for first assessment in 2025. Sub-topic covered: Structure 3.1, The periodic table: classification of elements. Understandings 3.1.1 to 3.1.6 are SL and HL; 3.1.7 to 3.1.10 are additional higher level.