S3.1Structure 3 — Classification of matter

The periodic table: classification of elements — IB Chemistry SL Revision Notes & Topical Questions

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S3.1 The periodic table: classification of elements — IB CHEMISTRY SL notes: Relate position in the periodic table to electron configuration, explain periodic trends, describe group 1, group 17 and oxide behaviour, and assign oxidation states.Arrangement — elements are ordered by…

Special cases
Naming groups
Reading a configuration
Across a period
Down a group
Electron affinity

Part 1 of 10Position and electron configuration

Position and electron configuration

Arrangement

Arrangement elements are ordered by increasing atomic number (1 to 118), in periods (rows) and groups (columns).

Period

Period the period number equals the number of occupied shells.
- (the highest principal quantum number)

Group

Group for s- and p-block elements the group number (1, 2, 13–18) indicates the valence electrons, adding 10 for groups 13–18 in IUPAC numbering.

Block

Block the last sublevel filled gives the block: s, p (groups 13–18), d (groups 3–12) and f (lanthanoids and actinoids).
- (groups 1–2 plus He)

Types

Types metals are on the left and centre, non-metals on the right and metalloids lie on the diagonal boundary.
- (B, Si, Ge, As, Sb, Te)

Special cases

Special cases hydrogen has its own position and helium sits in group 18 because it is unreactive.

Naming groups

Naming groups group 1 alkali metals, group 17 halogens, group 18 noble gases and groups 3–12 transition elements.

Reading a configuration

[Ne] 3s² 3p³ is period 3, group 15, p-block, so it is phosphorus.

Part 2 of 10Periodic trends and their explanation

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
Ionic radius	Cations decrease; anions are larger and decrease	Increases
First ionisation energy	Generally increases	Decreases
Electronegativity	Increases	Decreases
Electron affinity	Generally more exothermic	Generally less exothermic
Metallic character	Decreases	Increases

Identify the change in nuclear charge.

Identify the change in number of shells and distance.

Consider shielding by inner electrons.

Decide the net attraction on the valence electrons.

Relate it to radius, ionisation energy or reactivity.

Periodic trends and their explanation

Across a period

Across a period nuclear charge rises while shielding stays roughly constant, so valence electrons are pulled closer.

Down a group

Down a group extra shells increase distance and shielding, which outweigh the extra nuclear charge.

Ions

Ions cations are smaller than their atoms because a shell is lost, and anions are larger because of extra electron repulsion.

Part 3 of 10Electron affinity and electronegativity

1Electron affinity

Electron affinity the energy change when one mole of gaseous atoms gains one mole of electrons to form gaseous 1− ions, in kJ mol⁻¹.

2First value

First value usually exothermic, e.g. F(g) + e⁻ → F⁻(g), but endothermic for group 2 atoms and N.

3Second value

Second value endothermic, e.g. O⁻(g) + e⁻ → O²⁻(g), because the electron is repelled by the negative ion.

4Electronegativity

Electronegativity the ability of an atom to attract the shared pair of electrons in a covalent bond; it has no units and is not directly measured.

5Comparison

Comparison electron affinity is a measurable enthalpy change for forming anions, and both increase across a period and decrease down a group.

Part 4 of 10Group 1 metals with water and halogens

Equation
- 2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g), giving an alkaline solution.

Metal	Observations with water
Lithium	Floats and fizzes steadily, forming colourless LiOH solution
Sodium	Reacts faster, melts into a ball that moves on the surface
Potassium	Reacts violently, ignites the hydrogen with a lilac flame

Group 1 metals with water and halogens

Reactivity trend

Reactivity trend increases down the group because the single outer electron is further from the nucleus, more shielded and more easily lost.
- (lower ionisation energy)

Physical properties

Physical properties soft, shiny when cut, good conductors with low melting points that fall down the group and low densities.

Storage

Storage kept under oil because they react with air and moisture.

With halogens

With halogens form ionic halides, e.g. 2Na(s) + Cl₂(g) → 2NaCl(s), most vigorously for metals low in group 1 and halogens high in group 17,
- such as Cs with F₂.

Part 5 of 10Group 17 and halide ions

Group 17 and halide ions

Structure

Structure halogens are diatomic with seven outer electrons and form halide ions by gaining one electron.

States and colours

F₂ yellow gas, Cl₂ pale green gas, Br₂ red-brown liquid and I₂ grey-black solid, darkening down the group.

Solutions in water

Solutions in water chlorine green-blue, bromine orange and iodine dark brown.

Melting and boiling points

Melting and boiling points rise down the group because larger molecules have more electrons and stronger London forces.

Reactivity

Reactivity decreases down the group because the incoming electron is further away, more shielded and less strongly attracted.

Displacement

Displacement a more reactive halogen displaces a less reactive halide, e.g. Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq), giving an orange solution.

Reactivity order

Cl₂ > Br₂ > I₂, so Br₂ displaces iodide to give a brown solution.

Halide tests

Halide tests acidified AgNO₃(aq) gives white AgCl, cream AgBr and yellow AgI.

Part 6 of 10Oxides across period 3

Oxide	Bonding and structure	Nature	With water
Na₂O	Giant ionic	Basic	NaOH, strongly alkaline
MgO	Giant ionic	Basic	Mg(OH)₂, weakly alkaline
Al₂O₃	Ionic with covalent character	Amphoteric	No reaction; reacts with acids and bases
SiO₂	Giant covalent	Acidic	No reaction
P₄O₁₀, SO₃	Simple molecular	Acidic	H₃PO₄, H₂SO₄

Oxides across period 3

Trend

Trend oxides change from basic through amphoteric to acidic across the period.

Cause

Cause the electronegativity difference with oxygen falls across the period, so bonding changes from ionic to covalent.

Basic oxide equations

Na₂O + H₂O → 2NaOH and MgO + H₂O → Mg(OH)₂, and MgO + 2HCl → MgCl₂ + H₂O.

Acidic oxide equations

P₄O₁₀ + 6H₂O → 4H₃PO₄ and SO₃ + H₂O → H₂SO₄.

Down a group

Down a group oxides become more ionic.

Part 7 of 10Acid deposition and ocean acidification

Acid rain

Acid rain non-metal oxides dissolve in atmospheric water, SO₂ + H₂O → H₂SO₃ and 2NO₂ + H₂O → HNO₃ + HNO₂, lowering rainwater pH.

Ocean acidification

CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq), a weak acid that partly dissociates: H₂CO₃ ⇌ H⁺ + HCO₃⁻.

Effect

Effect the extra H⁺ lowers seawater pH and harms organisms that build calcium carbonate shells.

Part 8 of 10Oxidation and reduction

Definition	Oxidation	Reduction
Oxygen	Gain of oxygen	Loss of oxygen
Hydrogen	Loss of hydrogen	Gain of hydrogen
Electrons	Loss of electrons	Gain of electrons
Oxidation number	Increase	Decrease

Memory aid

OIL RIGOxidation Is Loss, Reduction Is Gain of electrons.

Redox reaction

Redox reaction any reaction in which an element changes its oxidation number.

Part 9 of 10Assigning oxidation states

Give uncombined elements 0 and monatomic ions their charge.

Give group 1 +1, group 2 +2 and fluorine −1.

Give hydrogen +1, except −1 in metal hydrides.

Give oxygen −2, except −1 in peroxides and +2 in OF₂.

Make the sum equal zero for a compound or the charge for an ion.

Solve for the unknown and write Roman numerals in names, e.g. manganate(VII).

Electronegativity

Electronegativity the more electronegative element takes the negative oxidation state,
- so OF₂ has F = −1 and O = +2,
- whereas Cl₂O has O = −2 and Cl = +1.

Fractional states

Fractional states an average over several atoms can be fractional, e.g. +8/3 for Fe in Fe₃O₄, but a single atom always has a whole-number state.

Examples

Mn is +7 in KMnO₄, Cr is +6 in Cr₂O₇²⁻ and S is +6 in SO₄²⁻.

Writing

Writing the sign comes before the number, as +2.

Part 10 of 10More from the specification

1Effective nuclear charge (Zeff)

The net positive nuclear attraction experienced by an outermost electron after accounting for inner shell electron shielding.

2Atomic radius periodicity

Atomic radius decreases across a period as added protons increase Z_eff, pulling the outermost electron shell closer to the nucleus.

3First ionisation energy trends

General increase across a period reflecting greater electrostatic pull on the outermost electron.

4Oxidation states and overall charge

In polyatomic ions, the algebraic sum of individual oxidation states of all atoms equals the overall charge of the ion.

Key terms

Term	Meaning
Period	A horizontal row of the periodic table; the number equals the occupied main energy levels.
Group	A vertical column of elements with the same number of valence electrons.
Block	A region of the table named after the sublevel being filled (s, p, d or f).
Electron affinity	The energy change when one mole of gaseous atoms gains electrons to form 1− ions.
Shielding	The reduction in nuclear attraction caused by inner-shell electrons.
Oxidation state	The charge an atom would have if all its bonds were ionic.
Amphoteric	Able to react as both an acid and a base, e.g. Al₂O₃.

Avoid these mistakes

Saying atomic radius increases across a period — nuclear charge increases and the radius decreases.

Assuming all halogens become more reactive down the group — reactivity decreases down group 17 but increases down group 1.

Forgetting oxygen is −1 in peroxides — it is not always −2.

Confusing electronegativity with electron affinity — electronegativity concerns bonded electrons and electron affinity is a measured enthalpy change.

Last updated: 14 October 2026