Comparing ionic and covalent bonds

Two mechanisms, two architectures. Ionic bonding transfers one or more electrons from an electropositive atom to an electronegative one, producing a cation and an anion that attract each other in every direction at once. There is no single "molecule" of $\text{NaCl}$ — only an endlessly repeating crystal lattice of alternating ions. Covalent bonding instead shares electron pairs between two atoms so each reaches a stable configuration; the result is a discrete molecule, internally strong but only weakly attracted to its neighbours.

Why the properties diverge. Melting an ionic solid means breaking Coulombic attractions running through the entire lattice — a huge energy cost, hence high melting points and hardness. Melting a molecular covalent solid only requires overcoming the weak intermolecular forces between molecules; the strong covalent bonds inside each molecule stay intact, hence low melting points. The same logic explains conductivity: ionic compounds conduct only when their ions are freed to move, i.e. molten or in solution, never as a rigid solid; molecular covalent compounds generally lack free charge carriers altogether.

A rough quantitative handle on lattice strength comes from treating the ions as point charges obeying Coulomb's law:

where: $q^+$ and $q^-$ are the magnitudes of the cation and anion charge, and $r_0$ is the sum of the ionic radii (the centre-to-centre distance). Higher charges and smaller ions pack more Coulombic attraction into the lattice, which is exactly why $\text{MgO}$ ($q^+q^- = 4$) melts far higher than $\text{NaCl}$ ($q^+q^- = 1$).

The same ratio, read differently, predicts covalent character. Define the ionic potential of the cation as $\phi = q^+ / r^+$ — charge packed into a small volume. A high $\phi$ means the cation is a strong polariser: it distorts, or "pulls on," the electron cloud of the neighbouring anion. Fajans' rules turn this into three practical predictors of increased covalent character in an otherwise ionic compound: a small, highly charged cation (high polarising power); a large, easily distorted anion (high polarisability); and a cation with a pseudo-noble-gas ($d^{10}$) electron configuration, which polarises more strongly than a true noble-gas cation of equal charge and size.

The physical driver behind every ionic-versus-covalent contrast is where the shared electron density actually sits. In a purely ionic picture the electron has moved completely onto the anion, so the bond is nothing but electrostatic attraction between two spherical, non-directional charge clouds — free to pack in whichever geometry maximises attraction and minimises repulsion, which is exactly a lattice. That is why ionic bonding gives no fixed "bond angle": each ion is simply surrounded by as many opposite-charge neighbours as will fit. Covalent bonding, by contrast, concentrates electron density between two specific nuclei, so the bond has a definite direction and a definite partner — which is why covalent compounds form discrete molecules with fixed bond angles rather than open-ended lattices. Fajans' rules matter because no real bond is a perfect point-charge transfer. When a small, highly charged cation such as $\text{Al}^{3+}$ sits next to a large anion such as $\text{Cl}^-$, its concentrated positive charge drags the anion's outer electron cloud toward itself. Push that distortion far enough and electron density starts accumulating between the two nuclei — the defining feature of a covalent bond — even though the compound formally consists of a metal and a non-metal. So covalent character is not an exception to ionic bonding; it is what happens whenever polarisation becomes large enough to matter, and the ionic potential $\phi = q^+/r^+$ is the single number that tells you how far along that spectrum a compound sits.

Rank $\text{NaCl}$, $\text{MgCl}_2$ and $\text{AlCl}_3$ in order of increasing covalent character, and explain the melting-point trend $1074\ \text{K} \to 987\ \text{K} \to 465\ \text{K}$ (sublimes) that goes with it.

Step 1. All three share the same anion, $\text{Cl}^-$, so by Fajans' rules the polarising power of the cation — governed by $\phi = q^+/r^+$ — controls the trend. The cation charges are $+1$, $+2$ and $+3$ respectively, and ionic radius shrinks as charge rises across a period ($\text{Na}^+ > \text{Mg}^{2+} > \text{Al}^{3+}$), so $\phi$ rises sharply from $\text{Na}^+$ to $\text{Mg}^{2+}$ to $\text{Al}^{3+}$.

Step 2. Higher $\phi$ means stronger distortion of the shared $\text{Cl}^-$ cloud, so covalent character increases in the order $\text{NaCl} < \text{MgCl}_2 < \text{AlCl}_3$. $\text{AlCl}_3$ is polarised so strongly that it behaves like a covalent compound in many respects — it sublimes rather than melts at a high temperature, dissolves in organic solvents, and exists as the dimer $\text{Al}_2\text{Cl}_6$ in the vapour phase.

Step 3. The melting-point fall follows directly: $\text{NaCl}$, essentially fully ionic, needs the most energy to break its extended Coulombic lattice ($1074\ \text{K}$). $\text{MgCl}_2$, more polarised, melts lower ($987\ \text{K}$) but is still a genuine ionic lattice. $\text{AlCl}_3$, so covalent that it consists of discrete molecular units held by weak intermolecular forces, needs comparatively little energy to separate its units and sublimes at only $465\ \text{K}$.

Answer: covalent character rises $\text{NaCl} < \text{MgCl}_2 < \text{AlCl}_3$, tracking the rising ionic potential of the cation, and melting point falls in the same order because the lattice progressively weakens into a molecular solid.

• Treating "ionic" and "covalent" as two watertight boxes. Fajans' rules exist precisely because real bonds sit on a continuum; a metal–non-metal pair can still be substantially covalent.
• Forgetting that anion size matters too. Polarising power is about the cation, but polarisability — how easily the electron cloud distorts — is an anion property; a large, diffuse anion like $\text{I}^-$ is polarised far more easily than a small, tightly-held anion like $\text{F}^-$.
• Confusing solid-state and molten/solution conductivity. Ionic compounds never conduct as rigid solids — the ions must be free to move, which happens only molten or dissolved.
• Assuming low melting point always means "purely covalent." Giant covalent networks such as diamond have very high melting points despite being covalent; the low-melting-point rule applies specifically to molecular covalent compounds, where only weak intermolecular forces need breaking.

Ask the tutor: Want to work through Fajans' rules for a different cation-anion pair, such as comparing MgCl2 and MgO for covalent character?