Polarisability in Ionic Bonding

In a purely ionic model, ions are treated as hard spheres of fixed charge that simply attract and repel each other according to Coulomb's law. In reality, the electron cloud around an ion is not rigid. It can be pulled and distorted by the electric field of a neighbouring ion. This distortion is called polarisation, and polarisability is a measure of how easily an ion's electron cloud can be distorted in this way.

When polarisation happens, electron density is pulled away from the anion and shifted towards the cation. This buildup of shared electron density between the two nuclei is exactly what a covalent bond looks like, so a highly polarised ionic bond gains some genuine covalent character. Most real ionic compounds sit somewhere on a spectrum between "purely ionic" and "purely covalent", and polarisability is the key idea that explains where a compound sits on that spectrum.

It is important to keep two related but different ideas separate:

• Polarising power describes the ability of a cation to distort a neighbouring anion's electron cloud. Cations with a small ionic radius and a high charge have a high charge density, so they exert a strong distorting pull on nearby electrons.
• Polarisability describes how easily an anion's own electron cloud can be distorted. Anions with a large ionic radius have their outer electrons held loosely and far from the nucleus, so they are more easily distorted.

The greatest degree of covalent character therefore appears in compounds pairing a small, highly charged cation with a large, singly (or multiply) charged anion, for example \(Al^{3+}\) with \(I^-\).

Fajans' rules give a simple qualitative way to predict when a compound is likely to show significant covalent character rather than behaving as a "clean" ionic solid. Covalent character increases when:

• The cation is small and/or highly charged, giving it a high charge density and strong polarising power.
• The anion is large and/or highly charged, making its electron cloud loosely held and easily polarisable.
• The cation does not have a noble-gas electron configuration (for example, transition metal cations polarise anions more strongly than an s-block cation of similar size and charge, because their electron shielding is less effective).

These rules explain patterns such as why silver halides show progressively more covalent character down the group (\(AgF\) is largely ionic, while \(AgI\) has substantial covalent character), and why aluminium chloride behaves more like a covalent, molecular compound than an ionic lattice.

Because polarisability depends on ionic size and charge, it follows predictable trends across the periodic table:

• Down a group, ions get larger, so anions become more polarisable and cations become less strongly polarising (their charge is spread over a bigger sphere, lowering charge density).
• Across a period, for ions of the same type, increasing nuclear charge and decreasing ionic radius increase a cation's polarising power.
• Higher ionic charge increases both polarising power (for cations) and, to a lesser extent, polarisability (for anions), since more charge is concentrated in a similar volume.

These same size and charge factors also govern how strongly an ion interacts with water molecules, which is discussed further in the lesson on hydration enthalpy.

The simple ionic model, used to build a theoretical lattice enthalpy from electrostatic attraction alone, assumes ions are non-polarisable point charges. When an ion pair is significantly polarised, the bond has real covalent character, and the lattice is held together more strongly than the pure ionic model predicts. As a result, the experimental lattice enthalpy (obtained indirectly, often via a Born-Haber cycle) is more negative (more exothermic) than the theoretical value calculated from an ionic model.

This discrepancy is one of the key pieces of evidence chemists use to judge how much covalent character a supposedly "ionic" compound actually has. You can see how these theoretical and experimental values are compared in the lesson on calculations using Born-Haber cycles, and how lattice enthalpy itself is defined and calculated in Enthalpy: lattice energy, atomisation and electron affinity.

Question: Explain why silver iodide, \(AgI\), shows more covalent character than silver fluoride, \(AgF\).

Answer: Both compounds have the same cation, \(Ag^+\), so the polarising power of the cation is fixed. The difference lies in the anion. The iodide ion, \(I^-\), is much larger than the fluoride ion, \(F^-\), because it has more electron shells. Its outer electrons are held further from the nucleus and less tightly, so its electron cloud is more easily distorted (more polarisable). This greater polarisation means more electron density is pulled into the region between the \(Ag^+\) and \(I^-\) nuclei, giving the bond more covalent character than in \(AgF\), where the small, tightly held fluoride ion resists distortion.

• Identify which ion is being polarised (usually the anion) and which is doing the polarising (usually the cation).
• Compare ionic radius: smaller cation and larger anion both increase polarisation and covalent character.
• Compare ionic charge: higher charge on either ion increases the effect.
• Check for non-noble-gas electron configurations on the cation, which strengthen polarising power further.

Related lessons

• Uk A Level Chemistry