Electron Configuration and Chemical Behaviour

Main-group elements gain or lose electrons to reach the nearest noble gas configuration (octet). The number of electrons gained or lost determines the ion charge.

Octet rule: main-group atoms lose or gain electrons to match the nearest noble gas (full s²p⁶ outer shell). Cations: Group 1 → 1+, Group 2 → 2+, Group 13 → 3+. Anions: Group 17 → 1−, Group 16 → 2−, Group 15 → 3−. Isoelectronic species (e.g. N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺, all 10 e⁻) have the same electron count but different nuclear charges; more protons → smaller ion. When ranking by size, compare the ions: Ne is the neutral atom of the set and its radius is defined differently (van der Waals), so it is not directly comparable with the ionic radii. (And note the octet rule is a model: noble gases are generally unreactive under ordinary conditions, but a few do form compounds, e.g. XeF₂ and XeF₄.)

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As you learned in L15, group number corresponds to valence electron count. All elements in the same group have the same number of valence electrons in the same type of subshell, and it is the valence electrons (not the core electrons) that determine chemical behaviour.

Same reaction type, same products, just increasingly vigorous because atomic radius increases down the group, making the valence electron progressively easier to lose (as established in L17). The type of reaction is determined by the valence configuration (1 electron to lose); the vigour is determined by the position in the group.

We just saw how electron configuration determines ion charge and the octet rule. That raises a question: why do all elements in the same group behave similarly in reactions? This card answers it → elements in the same group have identical valence electron configurations, so they undergo the same type of reaction; their position down the group only affects the vigour.

All Group 1 metals have valence configuration ns¹ → same number and type of valence electrons → same kind of chemistry (react with water to give H₂ + metal hydroxide: 2M + 2H₂O → 2MOH + H₂). Reaction type depends on valence configuration; reaction vigour depends on position down the group (K more vigorous than Li because larger atom, lower IE).

Add the highlighted group chemistry rule to your notes before the check below.

Transition metals (d-block) can form multiple oxidation states because their 3d and 4s electrons have similar energies, varying numbers of both can be lost in different reactions.

Transition metal compounds are often coloured because partially filled d orbitals allow electronic transitions (electrons jumping between d orbitals) in the visible light range, absorbing specific wavelengths. Compounds with full (Zn²⁺: 3d¹⁰) or empty (Sc³⁺: 3d⁰) d subshells are colourless. The exact colour is not fixed by the oxidation state alone, it also depends on the ligands attached, the geometry, the concentration and the pH, so colour is supporting evidence rather than a unique label for an oxidation state.

We just saw how valence electron configuration explains group behaviour. That raises a question: why can transition metals like iron form more than one stable ion (Fe²⁺ and Fe³⁺), unlike sodium which only forms Na⁺? This card answers it → 3d and 4s energies are similar, so successive electron removals are feasible without breaking into a stable noble-gas core.

Transition metals form multiple oxidation states because 3d and 4s electrons have similar energies, so variable numbers can be removed. 4s electrons are lost first, then 3d: Fe = [Ar]3d⁶4s² → Fe²⁺ = [Ar]3d⁶ → Fe³⁺ = [Ar]3d⁵. Transition-metal compounds are usually coloured due to d–d electron transitions (partially filled d orbitals absorb visible light); empty (Sc³⁺, 3d⁰) or full (Zn²⁺, 3d¹⁰) d subshells give colourless compounds.

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We just saw why transition metals form multiple oxidation states and why their compounds are coloured. That raises a question: once you know the ions and formula a compound contains, what are the actual rules for naming it correctly? This card answers it → the IUPAC naming conventions for inorganic substances, a syllabus dot point in its own right (investigate the nomenclature of inorganic substances using IUPAC naming conventions, CH11-8 dot point 3).

Chemists use a single, internationally agreed naming system, IUPAC nomenclature, so that any chemist anywhere can write or read a formula unambiguously. There are four main patterns you need for Module 1.

IUPAC naming rules: binary ionic = metal name + non-metal stem + "-ide" (NaCl = sodium chloride). Variable-charge cations get a Roman numeral worked out from the anion's charge (FeCl₃ = iron(III) chloride). Polyatomic ions are named directly, not broken into elements (Na₂SO₄ = sodium sulfate). Covalent binary compounds use Greek prefixes mono-/di-/tri-/tetra- (CO₂ = carbon dioxide; "mono-" dropped on the first element, e.g. CO = carbon monoxide, not monocarbon monoxide).

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We just saw the rules for naming compounds from their formula. That raises a question: how do you show the electron arrangement that a formula implies, the electrons actually being transferred or shared? This card answers it → Lewis dot (electron dot) diagrams, part of the syllabus dot point on using nomenclature, valency and chemical formulae (including Lewis dot diagrams) (CH11-8 dot point 16).

A Lewis dot diagram shows only the valence electrons of an atom or ion, drawn as dots around the element's symbol. The inner, core electrons are not shown, only the valence shell determines chemistry, bonding and reactivity.

Ionic species, electron transfer

For an ionic compound, draw the metal losing its valence electron dots completely (becoming a bare cation with no dots, in square brackets with its charge), and the non-metal gaining those dots to complete its octet (becoming an anion, in square brackets with its charge, showing a full outer shell of 8 dots).

Covalent species, shared pairs

For a covalent molecule, atoms share pairs of electrons (bonding pairs, shown as a shared dot pair or a line between atoms) so that each atom reaches a full outer shell. Electrons not involved in bonding remain as lone pairs on the atom.

Lewis dot diagrams show only valence electrons as dots around the element symbol. Ionic: the metal transfers its dots completely to the non-metal (Na → [Na]⁺ + [:Cl:]⁻). Covalent: atoms share bonding pairs to reach a full outer shell, remaining electrons stay as lone pairs (H₂O: O has 2 bonding pairs + 2 lone pairs; each H has 1 shared pair). Hydrogen's stable configuration is 2 electrons (a duet), not 8.

Pause, sketch the highlighted NaCl and H₂O Lewis diagrams into your book before moving on.

We just saw how to draw Lewis dot diagrams for ionic and covalent species. That raises a question: how do you write precise exam answers on ion charges, isoelectronic series, nomenclature and Lewis dot diagrams? This card answers it → for each question type, cite the electron configuration or valence electron count as evidence and apply the specific rule (octet, isoelectronic ranking, 4s-first removal, Stock notation, electron transfer/sharing).

For exam answers on ion charge: state the octet rule, name the nearest noble gas, then give the electron change. For isoelectronic series: same e⁻, more protons → smaller; list smallest to largest by proton number. For nomenclature: work out the Roman numeral from the anion's charge. For Lewis dot diagrams: count valence electrons first, then show transfer (ionic) or sharing (covalent) to reach a full outer shell.

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