Mission 1: Why is helium in group 18 when its atoms have only two electrons?

Periods, groups and blocks

• A period is a row. Its number is the outer occupied energy level.
• A group is a column, numbered 1 to 18.
• Blocks are named after the sublevel being filled: s, p, d, f.
• s-block: groups 1–2. p-block: groups 13–18. d-block: groups 3–12. f-block: the two rows below.
• The periodic table is in the data booklet.

Metals, metalloids, non-metals

• Metals fill the left and centre. Non-metals sit at the upper right, plus hydrogen.
• Metalloids lie along the stepped line: B, Si, Ge, As, Sb, Te.
• Metalloids have in-between properties: silicon is shiny but only a semiconductor.
• Aluminium touches the line but is a metal: shiny, malleable, a good conductor.

Valence electrons and the group

• Valence electrons: the electrons in the outer occupied energy level.
• Groups 1–2: valence electrons = group number.
• Groups 13–18: valence electrons = group number − 10.
• Group 16 atoms have 6 valence electrons, not 16. Group 13 atoms have 3.
• Helium is the one exception. The mission check explains it.

From position to configuration

• Period gives the outer level, n. Block gives the sublevel being filled.
• Bromine: period 4, p-block, group 17 → 7 valence electrons, 4s² 4p⁵.
• Bromine: [Ar] 3d¹⁰ 4s² 4p⁵. The 3d¹⁰ is inner: not valence.
• Manganese: period 4, d-block, group 7 → [Ar] 3d⁵ 4s².
• In the d-block, group = 4s + 3d electrons: 2 + 5 = 7.

From configuration to position

• Highest n → period. Sublevel being filled → block.
• [Ne] 3s² 3p¹: period 3, p-block, 3 outer electrons → group 13 (aluminium).
• [Ar] 3d⁸ 4s²: period 4, d-block, 8 + 2 → group 10 (nickel).
• Count the whole outer level, not just the last sublevel: 3s² 3p¹ is three.

The named families

• Alkali metals: group 1 metals, ns¹. Hydrogen sits above them but is a non-metal.
• Halogens: group 17, ns² np⁵, diatomic: F₂, Cl₂, Br₂, I₂.
• Noble gases: group 18, full outer level, very unreactive.
• Transition elements: d-block metals with an incomplete d sublevel. Zinc is not one.

Mission 2: Why is a chlorine atom smaller than a sodium atom, even though it has six more electrons?

Three factors explain every trend

• Nuclear charge: more protons, a stronger pull on the outer electrons.
• Distance: outer electrons in a higher level are further away, so pulled less.
• Shielding: inner electrons repel outer ones, reducing the pull they feel.
• Across a period: nuclear charge rises, same level, similar shielding → stronger pull.
• Down a group: a new level adds distance and shielding → weaker pull, despite more protons.

Atomic radius

• Across a period, atomic radius decreases: same level, more protons, electrons pulled in.
• Down a group, it increases: each period adds an occupied level.
• More electrons does not mean a bigger atom.

Ionic radius

• Cations are smaller than their atoms: the outer level is lost.
• Al [Ne] 3s² 3p¹ → Al³⁺ [Ne]: radius 124 → 54 pm (data booklet).
• Anions are larger than their atoms: same nuclear charge, more electron repulsion.
• Cl → Cl⁻: radius 100 → 181 pm (data booklet).
• Isoelectronic ions share a configuration; more protons → smaller ion.
• S²⁻ > Cl⁻ > K⁺ > Ca²⁺: all 1s² 2s² 2p⁶ 3s² 3p⁶, with 16, 17, 19, 20 protons.

Ionization energy

• First ionization energy: minimum energy to remove 1 mol of electrons from 1 mol of gaseous atoms in their ground state.
• X(g) → X⁺(g) + e⁻. Values are in the data booklet.
• Across a period it generally increases: more protons, similar shielding.
• Down a group it decreases: the outer electron is further away and more shielded.
• Group 2 (data booklet): Be 900, Mg 738, Ca 590 kJ mol⁻¹.
• The rise across a period has small dips; these are explained at HL.

Electron affinity

• First electron affinity: enthalpy change when 1 mol of gaseous atoms each gain one electron to form 1 mol of gaseous 1− ions.
• X(g) + e⁻ → X⁻(g). Usually negative: the nucleus attracts the incoming electron.
• Across a period: more protons, same level → the electron is attracted more strongly. EA becomes more negative.
• Down a group: further away, more shielded → attracted less. EA becomes less negative.
• Many exceptions. Chlorine (−349) is more negative than fluorine (−328, data booklet): fluorine's small, crowded atom repels the newcomer.

Electronegativity

• Electronegativity: the ability of an atom to attract a shared pair of electrons in a covalent bond.
• Pauling scale, no units (data booklet). Fluorine is highest, 4.0.
• Increases across a period and decreases down a group: the same three factors.
• No values for He, Ne, Ar: they form no ordinary bonds.
• Not electron affinity: that is measured on isolated gaseous atoms.

All the trends on one map

• Across a period → radius decreases; ionization energy and electronegativity increase.
• Down a group → radius increases; ionization energy and electronegativity decrease.
• Electron affinity: broadly more negative across a period, less negative down a group. There are many exceptions.
• Noble gases have no negative electron affinity.
• Metallic character increases down a group and towards the left.

Mission 3: Why does sodium fizz on water while caesium explodes?

Group 1 metals with water

• Alkali metals react with water to form a hydroxide and hydrogen.
• 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g). In general: 2M + 2H₂O → 2MOH + H₂.
• OH⁻ ions make the solution alkaline.
• Sodium floats, melts into a ball, fizzes and moves about.
• The reaction gets more vigorous down the group.

Why reactivity rises down group 1

• Group 1 metals react by losing their one outer electron.
• Down the group, that electron is further from the nucleus and more shielded.
• Lower ionization energy: it is lost more easily.
• So metallic character increases down group 1.

Halogens displace halide ions

• Halogens react by gaining electrons: X₂ + 2e⁻ → 2X⁻.
• A halogen displaces a halide below it in the group from solution.
• Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq).
• Colours in water: Cl₂ pale green, Br₂ yellow-orange, I₂ brown.
• Iodine cannot displace chloride or bromide: no change.

Why reactivity falls down group 17

• Down the group, atoms are larger: an incoming electron is further from the nucleus.
• It is also more shielded, so it is attracted less.
• So non-metallic character decreases down group 17.
• Chlorine pulls electrons from bromide ions; iodine cannot pull them from chloride ions.

Mission 4: If the oceans are "acidifying", why is seawater still alkaline?

Metal oxides are basic

• Group 1 and group 2 metal oxides (except BeO) are basic: they react with acids to form a salt and water.
• MgO(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l).
• Soluble ones give alkaline solutions: Na₂O(s) + H₂O(l) → 2NaOH(aq).
• Group 2: MO + H₂O → M(OH)₂, e.g. BaO(s) + H₂O(l) → Ba(OH)₂(aq).

Non-metal oxides are acidic

• Non-metal oxides react with water to form acids.
• CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq), carbonic acid, a weak acid.
• SO₂(g) + H₂O(l) → H₂SO₃(aq), sulfurous acid.
• SO₃(g) + H₂O(l) → H₂SO₄(aq), sulfuric acid.
• Carbon monoxide is a neutral oxide: it does not react with water.

Across a period: basic → amphoteric → acidic

• Period 3: Na₂O, MgO basic → Al₂O₃ amphoteric → SiO₂, P₄O₁₀, SO₃ acidic.
• Amphoteric: reacts with both acids and bases.
• Oxide character follows metallic character, which falls across a period.

Acid rain

• Rain is naturally slightly acidic, about pH 5.6, from dissolved CO₂.
• Acid rain, below 5.6, comes from sulfur oxides and nitrogen oxides.
• SO₂: sulfur impurities in coal and oil burn. In air, 2SO₂ + O₂ → 2SO₃.
• NO: N₂ + O₂ → 2NO in hot engines; then 2NO + O₂ → 2NO₂.
• 2NO₂(g) + H₂O(l) → HNO₃(aq) + HNO₂(aq).
• Effects: dissolves limestone and marble, damages trees, acidifies lakes.

Ocean acidification

• The oceans absorb part of the extra CO₂ released by burning fuels.
• CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq).
• More H⁺, lower pH: average surface pH has fallen from about 8.2 to 8.1 since pre-industrial times.
• Still alkaline: "acidification" means the pH is falling, not that it is below 7.
• H⁺ also removes carbonate ions, which corals and shellfish need to build CaCO₃.

Mission 5: How can nitrogen be −3 in ammonia but +5 in nitric acid?

What an oxidation state is

• An oxidation state is the charge an atom would have if all its bonds were ionic.
• Shared electrons are given to the more electronegative atom.
• Write the sign first: +2, −1. (Ion charges are written 2+, 1−.)
• Atoms in an element are 0: identical atoms, so no electrons are transferred.
• "Oxidation number" means the same thing.

The rules

• In a compound the oxidation states sum to zero; in an ion, to the ion's charge.
• O is −2, except −1 in peroxides, such as H₂O₂.
• H is +1, except −1 in metal hydrides, such as CaH₂.
• Group 1 metals +1, group 2 metals +2, fluorine always −1.
• Find the unknown atom last.

Worked examples

• NH₃: x + 3(+1) = 0, so N is −3.
• HNO₃: +1 + x + 3(−2) = 0, so N is +5.
• SO₄²⁻: x + 4(−2) = −2, so S is +6.
• H₂O₂: 2(+1) + 2x = 0, so O is −1 (a peroxide).
• CaH₂: +2 + 2x = 0, so H is −1 (a metal hydride).

Oxidation states in names

• Roman numerals give the oxidation state: iron(III) chloride, FeCl₃.
• KMnO₄: +1 + x + 4(−2) = 0, so Mn is +7: potassium manganate(VII).
• In oxyanions, -ate has the higher oxidation state and -ite the lower.
• NO₃⁻ nitrate, NO₂⁻ nitrite, SO₄²⁻ sulfate, SO₃²⁻ sulfite.
• Systematic names such as sulfate(IV) for sulfite are also accepted.

Mission 6: If nuclear charge rises across period 3, why does aluminium lose an electron more easily than magnesium?

Two dips in period 3

• First ionization energy rises across period 3, with two dips: Mg → Al and P → S.
• Mg 738, Al 578; P 1012, S 1000 kJ mol⁻¹ (data booklet).
• If every electron in a level had the same energy, the rise would be smooth.
• The dips are evidence that each level is split into sublevels.

Explaining each dip

• Mg → Al: aluminium's electron is removed from 3p, higher in energy than magnesium's 3s.
• P → S: sulfur's electron comes from a doubly occupied 3p orbital. Repulsion raises its energy.
• P 3p: [↑ ] [↑ ] [↑ ]. S 3p: [↑↓] [↑ ] [↑ ].
• Explain with the energy of the electron removed, not "stable half-full sublevels".

Mission 7: Why can manganese be +2 to +7, while calcium is only ever +2?

What counts as a transition element

• Transition element: an incomplete d sublevel in the atom or in one of its common ions.
• Scandium, [Ar] 3d¹ 4s²: incomplete 3d in the atom, so it is one.
• Zinc, [Ar] 3d¹⁰ 4s², forms only Zn²⁺, [Ar] 3d¹⁰: full 3d in both, so not one.
• "d-block" is a position; "transition element" depends on the configuration.

Characteristic properties

• Variable oxidation states: iron +2 and +3; manganese +2 to +7.
• High melting points, compared with the s-block metals of the same period (K, Ca).
• Coloured compounds.
• Magnetic properties, from unpaired d electrons.
• Catalysts: iron in the Haber process; MnO₂ catalyses the decomposition of H₂O₂.
• Complex ions with ligands.

Complex ions

• A ligand donates a lone pair to a central metal ion.
• Each ligand forms a coordination bond (a dative covalent bond).
• Common ligands: H₂O, NH₃, Cl⁻, CN⁻.
• [Co(H₂O)₆]²⁺: six water ligands around Co²⁺.
• Charge of a complex = metal ion charge + ligand charges: [CoCl₄]²⁻ is +2 + 4(−1).

Ions of transition elements

• Transition elements lose their 4s electrons first, then 3d.
• Mn [Ar] 3d⁵ 4s² → Mn²⁺ [Ar] 3d⁵.
• V [Ar] 3d³ 4s² → V³⁺ [Ar] 3d².
• Co [Ar] 3d⁷ 4s² → Co²⁺ [Ar] 3d⁷.

Why oxidation states vary

• The 4s and 3d sublevels are close in energy.
• So successive ionization energies rise gradually, with no huge jump among them.
• The energy for a higher state can be recovered from bonding, so several states are stable.
• Mn: 717, 1509, 3248 kJ mol⁻¹. Ca: 590, 1145, 4912 kJ mol⁻¹.
• Calcium's third electron would come from the [Ar] core: far too costly.

Mission 8: Why is the same cobalt(II) ion pink in water but blue in concentrated hydrochloric acid?

Ligands split the d sublevel

• In a complex, ligands split the five d orbitals into groups of different energy.
• The energy gap matches light in the visible range.
• A d electron absorbs that light and is promoted to a higher d orbital.
• We see the light that is not absorbed.

The colour wheel

• The colour seen is complementary to the colour absorbed: opposite on the wheel.
• Absorbs orange → looks blue. Absorbs green → looks red. Absorbs violet → looks yellow.
• The colour wheel, with wavelengths, is in the data booklet.
• The gap, and so the colour, depends on the metal, its oxidation state and the ligands.

From wavelength to frequency

• c = λf, so f = c/λ. c = 3.00 × 10⁸ m s⁻¹ (data booklet).
• Convert first: 1 nm = 10⁻⁹ m.
• Worked example: a complex absorbs most strongly at 600 nm (orange), so it looks blue.
• f = 3.00 × 10⁸ ÷ (6.00 × 10⁻⁷) = 5.00 × 10¹⁴ s⁻¹.

When there is no colour

• Colour needs a d electron to promote, and a d orbital to promote it into.
• 3d⁰ ions, such as Sc³⁺, have no d electron to promote.
• 3d¹⁰ ions, such as Zn²⁺, have no space to promote into.
• So their compounds are colourless (white as solids).

The big ideas

• Period = outer level; block = sublevel being filled. Groups 13–18: valence electrons = group − 10.
• Nuclear charge, distance and shielding explain every trend.
• Across a period, radius falls; ionization energy and electronegativity rise. Down a group, the reverse.
• Reactivity rises down group 1 and falls down group 17. A halogen displaces halides below it.
• Metal oxides basic, Al₂O₃ amphoteric, non-metal oxides acidic: the source of acid rain and ocean acidification.
• Oxidation states sum to zero in a compound, and to the charge in an ion.
• AHLIonization-energy dips show sublevels. Transition elements: incomplete d sublevel, variable states, coloured complexes.