Ionic Bonding and Lattices

Spec Mapping — OCR H432 Module 2.2.2 — Bonding and structure, covering ionic bonding as the electrostatic attraction between oppositely-charged ions, prediction of ion charges from periodic-table position, dot-and-cross diagrams for simple ionic compounds, the giant ionic lattice structure of NaCl, the physical properties of ionic compounds (m.p., conductivity, brittleness, solubility), and an introduction to the link between charge, ionic radius, and lattice strength (refer to the official OCR H432 specification document for exact wording).

Ionic bonding is one of the three core bonding models (alongside covalent and metallic) that you must be able to deploy, compare, and use to predict physical properties. An ionic compound is held together by the electrostatic force between ions of opposite charge, arranged in a regular three-dimensional giant ionic lattice. This lesson covers how ionic bonds form (electron transfer from metal to non-metal to reach noble-gas configurations), how to predict ion charges from periodic-table position, how to draw dot-and-cross diagrams, the structure of the NaCl lattice (Bragg's 1913 X-ray determination — the very first crystal structure ever solved), and the link between lattice strength, ionic charge, ionic radius and observable physical properties (melting point, electrical conductivity, brittleness, solubility). The qualitative ideas here are extended quantitatively in Module 5.2.1 (Born-Haber cycles, lattice enthalpy) in Year 13.

Key Definition: an ionic bond is the electrostatic attraction between oppositely-charged ions formed by complete electron transfer (typically from a metal to a non-metal). The bond is non-directional — each ion attracts all neighbouring ions of opposite charge.

---

How an ionic bond forms

A metal atom donates one or more electrons to a non-metal atom, both achieving stable noble-gas electron configurations:

$\text{Na (1s}^2\text{2s}^2\text{2p}^6\text{3s}^1\text{)} \rightarrow \text{Na}^+\text{(1s}^2\text{2s}^2\text{2p}^6\text{)} + \text{e}^-$Na (1s22s22p63s1)→Na+(1s22s22p6)+e−

$\text{Cl (1s}^2\text{2s}^2\text{2p}^6\text{3s}^2\text{3p}^5\text{)} + \text{e}^- \rightarrow \text{Cl}^-\text{(1s}^2\text{2s}^2\text{2p}^6\text{3s}^2\text{3p}^6\text{)}$Cl (1s22s22p63s23p5)+e−→Cl−(1s22s22p63s23p6)

Both ions now have the closed-shell electron configuration of a noble gas (Ne for Na⁺, Ar for Cl⁻). The energy released when the resulting Na⁺(g) and Cl⁻(g) ions assemble into a giant ionic lattice (the lattice enthalpy, ≈ −787 kJ mol⁻¹ for NaCl) is what makes ionic bonding favourable overall.

Ionic bonding occurs between elements with a large difference in electronegativity — typically a Group 1, 2 or 13 metal reacting with a Group 15, 16 or 17 non-metal. Pauling's empirical rule of thumb places the ionic/polar-covalent boundary at an electronegativity difference of about 1.7; NaCl with $\Delta\chi = 2.1$Δχ=2.1 is essentially ionic; HCl with $\Delta\chi = 0.9$Δχ=0.9 is polar covalent.

---

Predicting ion charges from periodic-table position

Group	Charge of monatomic ion	Examples
1	+1	Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺
2	+2	Be²⁺, Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺
13	+3	Al³⁺ (B is too small to form B³⁺ ionically)
14	± 4 (rare for ions)	Sn²⁺, Sn⁴⁺, Pb²⁺
15	−3	N³⁻, P³⁻
16	−2	O²⁻, S²⁻
17	−1	F⁻, Cl⁻, Br⁻, I⁻

Common polyatomic (compound) ions you must know:

Ion	Formula	Charge	Source acid
Ammonium	NH₄⁺	+1	NH₃ + H⁺
Hydroxide	OH⁻	−1	H₂O
Nitrate	NO₃⁻	−1	HNO₃
Hydrogencarbonate	HCO₃⁻	−1	H₂CO₃
Ethanoate	CH₃COO⁻	−1	CH₃COOH
Carbonate	CO₃²⁻	−2	H₂CO₃
Sulfate	SO₄²⁻	−2	H₂SO₄
Sulfite	SO₃²⁻	−2	H₂SO₃
Phosphate	PO₄³⁻	−3	H₃PO₄
Manganate(VII)	MnO₄⁻	−1	(KMnO₄)
Dichromate(VI)	Cr₂O₇²⁻	−2	(K₂Cr₂O₇)

The compound's overall formula must be charge-neutral — total positive = total negative.

---

Dot-and-cross diagrams for ionic compounds

A dot-and-cross diagram uses dots for one element's electrons and crosses for the other, with each ion enclosed in square brackets and the charge written outside the bracket (upper right). Only the outermost (valence) electrons are shown.

NaCl [ Na ]⁺ . . [ :Cl: × ]⁻ 1 e⁻ transferred

MgO [ Mg ]²⁺ [ :Ö: ××]²⁻ 2 e⁻ transferred

CaCl₂ [ Ca ]²⁺ 2 × [ :Cl: × ]⁻ 2 e⁻ transferred (one to each Cl)

Na₂O 2 × [ Na ]⁺ [ :Ö: ××]²⁻ 2 e⁻ transferred (one from each Na)

Convention • dots = original outer electrons of one element • crosses = original outer electrons of the other • brackets + charge for each ion separately • show outer shell only (typically 8 e⁻ per anion)

Rules

1. Each ion is in its own bracket, with the charge written immediately outside the upper-right corner.
2. Only outermost (valence) electrons are shown; the inner-shell electrons (core electrons) are not drawn.
3. Use dots for one element's original electrons, crosses for the other's — this lets the reader trace where each electron came from.
4. The cation usually shows zero electrons in its bracket (its outer shell has been emptied), with a positive charge. The anion shows all eight outer electrons (typically four dot/cross pairs).
5. Total electrons transferred must balance: 2 e⁻ from Ca for 2 Cl atoms (one each); 1 e⁻ from each of 2 Na atoms to 1 O.

---

Giant ionic lattice — the NaCl structure

In the solid state, ionic compounds form a giant ionic lattice — a regularly repeating 3D array of cations and anions held together by electrostatic forces extending throughout the entire crystal. There are no discrete "NaCl molecules"; the smallest stoichiometric unit (the formula unit) is just Na⁺Cl⁻ as a counting convention.

In sodium chloride specifically:

• Each Na⁺ is surrounded by 6 Cl⁻ in octahedral coordination.
• Each Cl⁻ is surrounded by 6 Na⁺ in octahedral coordination.
• Coordination ratio (cation : anion) = 6 : 6.
• The lattice is face-centred cubic (fcc) — Na⁺ and Cl⁻ together occupy the corners and face-centres of a cubic unit cell.

Key features: • fcc unit cell • 6:6 coordination • Cl⁻ larger than Na⁺ • Bragg (1913): first X-ray solved • Repeating in 3D no discrete molecules

The structure was solved by William Lawrence Bragg (and his father William Henry Bragg) in 1913 using X-ray diffraction — the first crystal structure ever determined experimentally, and the work that earned the Braggs the 1915 Nobel Prize in Physics (W. L. Bragg, aged 25, remains the youngest Nobel laureate in physics).

Not every ionic compound adopts the NaCl structure. CsCl, with a larger cation, adopts an 8:8 coordination (body-centred cubic). Fluorite (CaF₂) adopts a 8:4 coordination. The structure adopted is determined by the radius ratio $r_+/r_-$r+​/r−​ — small cations prefer lower coordination numbers, large cations prefer higher.

---

Coulomb's law and the strength of an ionic bond

The electrostatic force between two point charges is given by Coulomb's law:

$F = \frac{Q_1 Q_2}{4\pi \varepsilon_0 r^2}$F=4πε0​r2Q1​Q2​​

and the potential energy by:

$E = -\frac{Q_1 Q_2}{4\pi \varepsilon_0 r}$E=−4πε0​rQ1​Q2​​

For ionic compounds we replace point charges with ionic charges and $r$r with the centre-to-centre distance. Two qualitative consequences:

1. Higher charges → stronger bond. Doubling both charges multiplies the force/energy by 4. So MgO (+2/−2) is bound about four times more strongly per ion pair than NaCl (+1/−1).
2. Smaller ions → stronger bond. Halving the separation distance doubles the force and quadruples it back at the energy level. So NaF (with smaller F⁻ than Cl⁻) has a stronger lattice than NaCl.

Combined: the lattice strength scales roughly as $Q_+ \cdot Q_- / (r_+ + r_-)$Q+​⋅Q−​/(r+​+r−​), and the melting point of an ionic compound tracks lattice strength closely.

Compound	Ionic charges	Approx. r₊ + r₋ (pm)	m.p. (°C)	Lattice enthalpy (kJ mol⁻¹)
NaCl	+1, −1	283	801	−787
NaF	+1, −1	235	993	−929
KCl	+1, −1	314	770	−711
MgO	+2, −2	212	2852	−3791
CaO	+2, −2	240	2613	−3401
Al₂O₃	+3, −2	194	2072	−15916

Notice the dramatic 4× jump from NaCl to MgO (both monatomic, both NaCl-structure) — the doubled charges dominate. The Al₂O₃ value is per mole of Al₂O₃ formula unit (five ions, with the +3 driving a vastly more negative number).

---

Physical properties of ionic compounds

Property	Observation	Explanation
Melting / boiling point	High (typically > 700 °C)	Many strong electrostatic attractions must be overcome to separate ions
Hardness	Hard	Strong lattice resists deformation
Brittleness	Brittle	Shifted layer brings like charges into alignment → strong repulsion → cleavage
Solid conductivity	Non-conducting	Ions are fixed in the lattice; no mobile charge carriers
Molten / aqueous conductivity	Conducts	Ions are now free to migrate under an applied field
Solubility in water	Often soluble	Polar water molecules solvate ions; hydration enthalpy compensates for breaking lattice
Solubility in non-polar solvents	Insoluble	Non-polar solvents cannot stabilise the separated ions

Why ionic solids are brittle

Apply a mechanical stress to an ionic crystal: one layer of ions shifts by exactly one ionic radius relative to the next. Suddenly Na⁺ ions sit above Na⁺ ions and Cl⁻ above Cl⁻ — like charges repel, the layer splits along the fault, and the crystal cleaves cleanly along that plane. This is why salt crystals always break along flat planes (their cleavage planes), not along curved fractures like metals or glass.

Aligned (before stress) Cations attract anions above and below → stable lattice

Mis-aligned (after stress) Slip: now Na⁺ above Na⁺, Cl⁻ above Cl⁻ Like-charge repulsion → cleavage along dashed line

Why ionic solids dissolve in water

Water molecules are polar — the O bears a $\delta^-$δ− partial charge and each H a $\delta^+$δ+. When an ionic solid contacts water:

1. The $\delta^+$δ+ ends of water molecules orient toward each anion in the lattice surface.
2. The $\delta^-$δ− end of water molecules orients toward each cation.
3. The ion-dipole interactions release energy (hydration enthalpy) as each ion is solvated.
4. If the total hydration enthalpy exceeds the lattice enthalpy, dissolution is energetically favourable.