Why Table Salt Is Ionic but Aluminium Chloride Behaves Like a Covalent Compound

Atoms bond because the bonded arrangement has lower energy than the separate atoms. Some atoms transfer electrons and form ions held by electrostatic attraction, others share electrons, and many real compounds lie somewhere in between.

This lesson covers ionic bonding and lattice energy, Lewis structures, and Fajans' rules for covalent character.

An ionic bond forms when a metal atom with low ionisation enthalpy transfers electrons to a non-metal atom with a strongly negative electron gain enthalpy, and the large energy released as the ions pack into a lattice makes the process favourable; the Born-Haber cycle calculates that lattice energy using Hess's law.

Conditions for ionic bonding:

- Low ionisation enthalpy of the metal
- Strongly negative electron gain enthalpy of the non-metal
- High lattice enthalpy of the resulting solid

Lattice enthalpy. The energy needed to separate one mole of a solid ionic compound completely into gaseous ions.

Born-Haber cycle for sodium chloride:

- Sublimation of sodium, $\mathrm{Na}(\mathrm{s})\to \mathrm{Na}(\mathrm{g})$: $+108$ kJ mol${}^{-1}$
- Ionisation of sodium, $\mathrm{Na}(\mathrm{g})\to \mathrm{N}{\mathrm{a}}^{+}(\mathrm{g})+{\mathrm{e}}^{-}$: $+496$ kJ mol${}^{-1}$
- Half the bond dissociation of chlorine, $\frac{1}{2}\mathrm{C}{\mathrm{l}}_2(\mathrm{g})\to \mathrm{Cl}(\mathrm{g})$: $+121$ kJ mol${}^{-1}$
- Electron gain by chlorine, $\mathrm{Cl}(\mathrm{g})+{\mathrm{e}}^{-}\to \mathrm{C}{\mathrm{l}}^{-}(\mathrm{g})$: $-349$ kJ mol${}^{-1}$
- Enthalpy of formation of sodium chloride: $-411$ kJ mol${}^{-1}$

Worked example. By Hess's law, the enthalpy of formation equals the sum of all the steps, including forming the lattice from gaseous ions, which releases energy U:

$$-411=108+496+121-349-U$$

$$U=108+496+121-349+411=787{\text{ kJ mol}}^{-1}$$

An everyday example. Sea salt crystals drying in the salt pans of Gujarat are held together by this strong lattice energy, which is why salt melts only at a very high temperature.

The substance. Ionising sodium on its own costs energy — the ionic compound forms only because the lattice energy released more than pays for it.

A Lewis structure shows valence electrons as dots or lines, with shared pairs between bonded atoms and lone pairs on individual atoms, arranged so that most atoms reach an octet, or a duplet for hydrogen.

Steps:

- Add up the valence electrons of all atoms, adding one for each negative charge and subtracting one for each positive charge
- Choose the central atom, usually the least electronegative one and never hydrogen
- Join each outer atom to the centre with a single bond
- Place the remaining electrons as lone pairs, completing outer atoms' octets first
- If the central atom still lacks an octet, turn lone pairs into double or triple bonds

**Worked example 1 — ammonia, $\mathrm{N}{\mathrm{H}}_3$.** Valence electrons $=5+3(1)=8$. Three N-H single bonds use 6 electrons, leaving one lone pair on nitrogen.

**Worked example 2 — carbon dioxide, $\mathrm{C}{\mathrm{O}}_2$.** Valence electrons $=4+2(6)=16$. Two C=O double bonds use 8 electrons, and each oxygen carries two lone pairs, so every atom has an octet.

**Worked example 3 — nitrate ion, $\mathrm{N}{\mathrm{O}}_3^{-}$.** Valence electrons $=5+3(6)+1=24$. Nitrogen forms one double bond and two single bonds to oxygen, and the double bond can be drawn to any of the three oxygens — an example of resonance.

Exceptions to the octet rule:

- Incomplete octet — boron in $\mathrm{B}{\mathrm{F}}_3$ has only 6 electrons
- Expanded octet — phosphorus in $\mathrm{PC}{\mathrm{l}}_5$ and sulphur in $\mathrm{S}{\mathrm{F}}_6$ have more than 8
- Odd-electron molecules — nitric oxide, NO, has an unpaired electron

An everyday example. The bubbles in a glass of soda water are carbon dioxide molecules held together by two double bonds, just as the Lewis structure shows.

The substance. Lewis structures show bonding, not shape — predicting molecular geometry needs VSEPR theory.

Fajans' rules state that an ionic bond gains covalent character when a small, highly charged cation distorts — polarises — the electron cloud of a large anion, and when the cation has a pseudo-noble-gas configuration rather than a noble-gas configuration.

Covalent character increases when:

- The cation is small — a small cation has a high charge density and polarises the anion strongly
- The anion is large — its outer electrons are far from its nucleus and easily distorted
- The charge on either ion is high — $\mathrm{A}{\mathrm{l}}^{3+}$ polarises far more strongly than $\mathrm{N}{\mathrm{a}}^{+}$
- The cation has a pseudo-noble-gas configuration — such as $\mathrm{C}{\mathrm{u}}^{+}$ or $\mathrm{A}{\mathrm{g}}^{+}$, whose d electrons shield the nucleus poorly

Worked examples:

- Lithium halides — covalent character increases from LiF to LiCl to LiBr to LiI as the anion grows larger
- Period 3 chlorides — covalent character increases from NaCl to $\mathrm{MgC}{\mathrm{l}}_2$ to $\mathrm{AlC}{\mathrm{l}}_3$ as cation size falls and charge rises; aluminium chloride is largely covalent
- NaCl versus CuCl — both cations carry a +1 charge, but $\mathrm{C}{\mathrm{u}}^{+}$ has a pseudo-noble-gas configuration, so copper(I) chloride is more covalent

Effects of covalent character:

- Lower melting points — aluminium chloride sublimes at a far lower temperature than sodium chloride melts
- Lower solubility in water and greater solubility in organic solvents

An everyday example. Silver chloride forming as an insoluble white precipitate in a school test for chloride ions reflects the extra covalent character that $\mathrm{A}{\mathrm{g}}^{+}$ brings compared with $\mathrm{N}{\mathrm{a}}^{+}$.

The substance. No bond is purely ionic — even in sodium chloride the chloride ion is slightly polarised, so every ionic compound has some covalent character.