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

These notes condense Structure 3.1, The periodic table: classification of elements, following the IB Chemistry guide for the Diploma Programme (February 2023 edition, first examined in 2025). They cover understandings 3.1.1 to 3.1.10. Sections 3.1.1 to 3.1.6 apply at SL and HL. Sections 3.1.7 to 3.1.10 come from the additional higher level part of the guide and are marked HL only. The notes apply to the course from 2025 onwards.

For full explanations and worked examples, read the Structure 3.1 study guide first. Then try the practice questions. To spot gaps in other units, the free diagnostics help. Keep the course hub and the checklist you can print to hand.

3.1.1 Layout of the table

Block	Groups	Sublevel being filled
s	1, 2 (and He)	s
p	13 to 18	p
d	3 to 12	d
f	Lanthanoids, actinoids	f

• Metals: left and centre. Non-metals: upper right, plus hydrogen.
• Metalloids lie along the stepped boundary: B, Si, Ge, As, Sb, Te.
• The guide notes that the data booklet includes a periodic table.

3.1.2 Position and configuration

• Period = outer occupied main level (n).
• Group = shared number of valence electrons. Groups 1 and 2: same as group number. Groups 13 to 18: group number minus 10.
• Names to know: alkali metals (group 1), halogens (group 17), noble gases (group 18), transition elements (d-block).

Method in steps: position to configuration

1. Read the period: this gives the outer n.
2. Read the block: this tells you which sublevel the last electron enters.
3. Count across the block to place the last electron.
4. Fill in order 1s 2s 2p 3s 3p 4s 3d 4p, and check the total equals Z.

Reminder: calcium (period 4, group 2, s-block) is [Ar] 4s². Krypton (period 4, group 18) is [Ar] 3d¹⁰ 4s² 4p⁶. Exceptions up to Z = 36: Cr [Ar] 3d⁵ 4s¹ and Cu [Ar] 3d¹⁰ 4s¹.

3.1.3 Periodic trends

Two explanations do nearly all the work:

• Across: more protons, same main level, similar shielding → outer electrons pulled harder.
• Down: extra main level, more shielding → outer electrons further away and held less tightly.

Property	Across a period	Down a group
Atomic radius	↓	↑
Ionic radius	Falls along cations, jumps up at the first anion, then falls	↑
First ionization energy	↑ (with dips)	↓
Electron affinity	Generally more exothermic	Generally less exothermic
Electronegativity	↑	↓

Definitions

• First ionization energy: the energy, per mole, to take the outermost electron from each atom of a gaseous element in its ground state, giving gaseous 1+ ions.
• Electron affinity: the energy change, per mole, as each gaseous atom picks up one electron to become a gaseous 1− ion.
• Electronegativity: an atom’s ability to attract a bonding pair of electrons in a covalent bond.

Ions

• Cation smaller than its atom: one fewer occupied main level.
• Anion larger than its atom: more electron repulsion, same nuclear charge.
• Isoelectronic ions: more protons → smaller. Example: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ (all 10 electrons).

3.1.4 Group 1 and group 17

Group 1 + water: 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g). Alkaline solution; fizzing. Reactivity increases down the group as the outer electron is further away, more shielded and lost more easily.

Group 17 + halide ions: a halogen oxidizes halide ions of halogens below it.

Halogen added	Cl⁻(aq)	Br⁻(aq)	I⁻(aq)
Cl₂(aq)	–	Orange (Br₂ forms)	Brown (I₂ forms)
Br₂(aq)	No reaction	–	Brown (I₂ forms)
I₂(aq)	No reaction	No reaction	–

Reactivity decreases down the group: a bigger atom attracts an incoming electron less strongly.

Small reminder: Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq).

3.1.5 Oxides

Continuum across a period: basic → amphoteric (Al₂O₃) → acidic.

Oxide	With water
Li₂O	Li₂O(s) + H₂O(l) → 2LiOH(aq)
MgO	MgO(s) + H₂O(l) → Mg(OH)₂(s) (slightly soluble, weakly alkaline)
BaO	BaO(s) + H₂O(l) → Ba(OH)₂(aq)
CO₂	CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq)
SO₂	SO₂(g) + H₂O(l) → H₂SO₃(aq)
SO₃	SO₃(g) + H₂O(l) → H₂SO₄(aq)

• Acid rain: SO₂ and nitrogen oxides from burning fuels dissolve in rain to form acids.
• Ocean acidification: extra CO₂ dissolves → H₂CO₃ ⇌ H⁺ + HCO₃⁻ → pH falls; H⁺ removes CO₃²⁻ as HCO₃⁻, so less carbonate is available for shells.

3.1.6 Oxidation states

Method in steps

1. Give elements 0.
2. Fix the “certain” atoms: group 1 +1, group 2 +2, F −1, H +1, O −2.
3. Override for exceptions: H −1 in metal hydrides; O −1 in peroxides.
4. Set the sum equal to the overall charge and solve.

Small reminders:

I in HIO₃:    (+1) + x + 3(−2) = 0    -> x = +5
N in HNO₂:    (+1) + x + 2(−2) = 0    -> x = +3

• Write the sign before the number: +3, not 3+ (3+ is an ion charge).
• An element’s oxidation state is zero because bonding between identical atoms involves no electron transfer.
• Oxyanion names: NO₃⁻ nitrate(V)/nitrate, NO₂⁻ nitrate(III)/nitrite, SO₄²⁻ sulfate(VI)/sulfate, SO₃²⁻ sulfate(IV)/sulfite.

HL only: 3.1.7 Dips in first ionization energy

• Group 2 → 13 (Be→B, Mg→Al): the electron removed is in a p sublevel, higher in energy than the s sublevel, so it is easier to remove.
• Group 15 → 16 (N→O, P→S): the electron removed is paired; repulsion raises its energy, so it is easier to remove.
• Base the answer on the energy of the electron removed. Do not rely on “half-filled stability”.

HL only: 3.1.8 Transition element properties

Variable oxidation states · high melting points · magnetic properties · catalytic properties · coloured compounds · complex ions with ligands.

• Cause: an incomplete d-sublevel.
• Zn excluded: 3d¹⁰ in the atom and in Zn²⁺.
• Scandium: debated (Sc³⁺ has no d electrons).
• Ligand: a molecule or ion with a lone pair that it donates to the central ion.

HL only: 3.1.9 Variable oxidation states

• 4s and 3d are close in energy → successive ionization energies rise steadily, without the large jump a group 2 element shows after two electrons.
• Ions: 4s out first.

Ion	Configuration
Ti²⁺	[Ar] 3d²
Mn³⁺	[Ar] 3d⁴
Co³⁺	[Ar] 3d⁶
Cu²⁺	[Ar] 3d⁹

HL only: 3.1.10 Colour

1. Ligands split the d orbitals into two sets.
2. An electron absorbs visible light of energy equal to the gap and is promoted.
3. The colour seen is the complementary colour, opposite on the colour wheel.
4. Use c = λf to convert. Put λ in metres.

The guide says the data booklet supplies both the colour wheel and c = λf. Approximate pairs: red ↔ green, orange ↔ blue, yellow ↔ violet.

Must-know distinctions

• Group (column, valence electrons) vs period (row, outer main level).
• Electron affinity (energy change when gaseous atoms gain electrons, kJ mol⁻¹) vs electronegativity (pull on a shared pair; no units).
• Halogen (the element, Cl₂) vs halide (the ion, Cl⁻).
• Oxidation state (+2, sign first) vs ion charge (2+, number first).
• (HL) Colour absorbed vs colour seen.

Quick self-test

1. State the period, group and block of the element with configuration [Ar] 3d¹⁰ 4s² 4p⁴, and name it.
2. How many valence electrons does the element in period 3, group 13 have?
3. Which is larger, a chlorine atom or a chloride ion? Give a reason.
4. Write an equation for lithium reacting with water.
5. Aqueous chlorine is added to potassium bromide solution. State the colour change and write the ionic equation.
6. Write an equation for sulfur trioxide reacting with water.
7. Deduce the oxidation state of phosphorus in P₂O₇⁴⁻ and of nitrogen in N₂F₄.
8. Give the systematic name of NaClO.
9. (HL) Why is the first ionization energy of aluminium lower than that of magnesium?
10. (HL) Write the configurations of Ni²⁺ and Cr³⁺.
11. (HL) A complex absorbs at 440 nm. State the colour seen and calculate the frequency absorbed (c = 3.00 × 10⁸ m s⁻¹).
12. (HL) A complex absorbs light of frequency 6.00 × 10¹⁴ s⁻¹. Find the wavelength and the colour seen.

Answers

1. Period 4, group 16, p-block: selenium (Z = 34).
2. Three (13 − 10).
3. The chloride ion: one extra electron adds repulsion with the same 17 protons.
4. 2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)
5. Pale green/colourless to orange; Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq).
6. SO₃(g) + H₂O(l) → H₂SO₄(aq)
7. P: 2x + 7(−2) = −4, so +5. N: 2x + 4(−1) = 0, so +2.
8. Sodium chlorate(I).
9. Aluminium’s electron is removed from 3p, higher in energy than magnesium’s 3s electron.
10. Ni²⁺ [Ar] 3d⁸; Cr³⁺ [Ar] 3d³.
11. 440 nm is violet, so yellow is seen. f = 3.00 × 10⁸ / (440 × 10⁻⁹) = 6.82 × 10¹⁴ s⁻¹.
12. λ = 3.00 × 10⁸ / 6.00 × 10¹⁴ = 5.00 × 10⁻⁷ m = 500 nm, green absorbed, so red is seen.

Where marks are usually lost

• Saying a group 15 element has 15 valence electrons (it has 5).
• Explaining a trend with “more protons” alone, without the same-level or extra-level shielding point.
• Describing the halogen colour you add instead of the colour of the halogen that forms.
• Missing state symbols, or not balancing group 1 + water equations as 2 : 2 → 2 : 1.
• Showing the CO₂ and water reaction as going to completion, or leaving out the carbonate step in ocean acidification.
• Forgetting that metal hydrides put hydrogen at −1 and peroxides put oxygen at −1.
• Writing an oxidation state as “2+” rather than “+2”.
• (HL) Using “half-filled stability” for the group 15 to 16 dip.
• (HL) Taking 3d electrons out before 4s.
• (HL) Leaving λ in nanometres in c = λf.

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.