Ionic Bonding, Structure & Metallic Properties – Full Study Notes

Metals form giant metallic lattices: regularly arranged cations immersed in a sea of delocalised electrons.
Metallic bond = strong electrostatic attraction between $ext{M}^{n+}$ ions and free electrons.
Key consequences of the structure
- High density – ions closely packed.
- Crystalline solids – long-range ordering of ions.
- High melting/boiling points – much energy required to overcome the strong attractions.
- Electrical conductivity – delocalised electrons flow toward the positive terminal.
- Malleability / ductility – layers of ions can slide without breaking the metallic bond.
- Strength – overall lattice cohesion from non-directional bonding.

Groups = vertical columns; elements share the same number of valence electrons → similar reactivity.
Periods = horizontal rows; period number = number of occupied electron shells.
Atomic number (Z) dictates proton number → also electron number in the neutral atom.
2n² rule – maximum electrons per shell $n$ : $2n^2$ (e.g.
1st shell: $2 \times 1^2 = 2$ ; 2nd: $8$ ; 3rd: $18$ …).
Electron configuration trends
- Left side: few valence electrons → metallic character, tendency to lose e⁻.
- Right side (p-block): nearly full valence shells → non-metallic, tendency to gain e⁻.
Special groups
- Group 1 (alkali metals): 1 e⁻ in outer shell, very reactive, form $+1$ ions.
- Group 2 (alkaline earths): 2 e⁻ outer shell, form $+2$ ions.
- Transition metals: multiple oxidation states, coloured compounds, useful catalysts.
- Halogens (Group 17): 7 e⁻ outer shell, form $-1$ ions, diatomic molecules, highly reactive.
- Noble gases (Group 18): full valence shells, very unreactive.
Excited electrons: absorb energy → jump to higher energy levels; on relaxing emit photons (spectral lines; basis for flame tests).

Ion = atom/ group of atoms with net charge owing to unequal proton/electron numbers.
Cation (positive ion) – produced by electron loss; typical for metals.
Anion (negative ion) – produced by electron gain; typical for non-metals.

Atom	e⁻ Configuration	e⁻ Lost	Resulting Ion & Notation
Lithium	$2.1$	$1$	$[2]^+ \;(\text{Li}^+)$
Sodium	$2.8.1$	$1$	$[2.8]^+ \;(\text{Na}^+)$
Magnesium	$2.8.2$	$2$	$[2.8]^{2+} \;(\text{Mg}^{2+})$
Aluminium	$2.8.3$	$3$	$[2.8]^{3+} \;(\text{Al}^{3+})$

Atom	e⁻ Configuration	e⁻ Gained	Resulting Ion & Notation
Fluorine	$2.7$	$1$	$[2.8]^- \;(\text{F}^-)$
Chlorine	$2.8.7$	$1$	$[2.8.8]^- \;(\text{Cl}^-)$
Oxygen	$2.6$	$2$	$[2.8]^{2-} \;(\text{O}^{2-})$
Sulfur	$2.8.6$	$2$	$[2.8.8]^{2-} \;(\text{S}^{2-})$

Iso-electronic ions: different elements yielding the same e⁻ configuration after ionisation (e.g.
$\text{O}^{2-}$ and $\text{F}^-$ both ⇒ $[2.8]$ ).

Ionic compound = combination of cations + anions held by strong Coulombic forces.
Formed chiefly between metals (e⁻ donors) and non-metals (e⁻ acceptors).
Process (NaCl exemplar)
1. $\text{Na}$ loses 1 e⁻ → $\text{Na}^+$ .
2. $\text{Cl}$ gains 1 e⁻ → $\text{Cl}^-$ .
3. Opposite charges attract → ionic bond.
Electrostatic attraction magnitude ∝ product of charges; higher charges → stronger bonds.

General algorithm

Sodium oxide: $\text{Na}^+$ vs $\text{O}^{2-}$ → need 2 × Na per O ⇒ $\text{Na}_2\text{O}$ (2:1 ratio).
Magnesium chloride: $\text{Mg}^{2+}$ vs $\text{Cl}^-$ → need 2 × Cl per Mg ⇒ $\text{MgCl}_2$ (1:2).
Aluminium bromide: $\text{Al}^{3+}$ vs $\text{Br}^-$ → need 3 × Br ⇒ $\text{AlBr}_3$ (1:3).
Aluminium oxide: $\text{Al}^{3+}$ vs $\text{O}^{2-}$ → LCM of charges = 6 → 2 × Al, 3 × O ⇒ $\text{Al}2\text{O}3$ (2:3).

$\text{LiF},\; \text{CaF}2,\; \text{AlF}3, \dots$
$\text{Na}2\text{O},\; \text{Al}2\text{O}_3, \dots$
$\text{Mg}3\text{N}2, \text{K}_3\text{N}, \dots$
(Full grid provided in transcript covers Li, Ca, Na, Mg, Al, K with F, O, N, Br, S, Cl.)

Ionic lattice = vast 3-D repeating network; each ion surrounded by oppositely charged neighbours → maximised attraction, minimised repulsion.
Crystals: external cubic shapes reflect internal cubic lattice (e.g.
NaCl).

Compound	Charges	$T_m$ (°C)	$T_b$ (°C)	Explanation
NaCl	$+1/-1$	$801$	$1413$	Weaker attraction vs 2+/2-
MgO	$+2/-2$	$2852$	$3600$	Higher ionic charge → stronger bond → more heat needed

General rule: bigger charges & smaller ionic radii → higher melting/boiling points.

Solid: ions fixed; no conduction.
Molten / aqueous solution: lattice broken, ions mobile → carry current.
Solubility in water arises because polar $\text{H}_2\text{O}$ molecules pull ions from lattice.

When stress shifts a layer, like-charges may align → strong repulsion → lattice cracks.

Metallic: cations + sea of e⁻; conducts as solid; malleable.
Ionic: cations + anions; non-conductor solid, conductor molten/solution; brittle; high $T_m$ .
Covalent (introduced only): shared e⁻ pairs; discrete molecules or giant networks; properties depend on structure (low $Tm$ for simple molecules, very high for networks like $\text{SiO}2$ ).

Trace historical development of atomic theory.
Distinguish group vs period, relate Z to electrons.
Apply 2n² rule & determine configurations and valence.
Correlate valence electrons with metallic / non-metallic properties.
Describe electron excitation phenomena.
Catalogue properties of Group 1, 2, transition metals, halogens, noble gases.
Define & determine charges on ions from Groups 1, 2, 15, 16, 17.
Recognise polyatomic ions (e.g.
$\text{SO}4^{2-},\; \text{NO}3^-$ ) – future lessons.
Draw electron-transfer diagrams for ionic bonding & shared-pair diagrams for covalent molecules.
Contrast metallic compounds vs alloys (mixture of metals; enhanced hardness, corrosion resistance etc.).

Metals & ionic substances underpin manufacture of fuels, pharmaceuticals, structural materials.
Understanding conductivity and malleability informs electrical wiring & fabrication industries.
Proper handling of reactive Group 1 metals (stored under oil) – safety implication.

$2n^2$ electron capacity rule.
Charge annotations: $+$ or $-$ sign with magnitude (e.g.
$\text{Al}^{3+}$ , $\text{O}^{2-}$ ).
Example melting points: $Tm(\text{NaCl}) = 801\,^\circ!\text{C}$ ; $Tm(\text{MgO}) = 2852\,^\circ!\text{C}$ .
Formulas derived: $\text{Na}2\text{O}, \text{MgCl}2, \text{AlBr}3, \text{Al}2\text{O}_3$ .