Ionic Bonding

Ionic bonding arises from the electrostatic attraction between ions of opposite charge - a cation and an anion that have been produced by the transfer of one or more electrons from a metal to a non-metal. ^[1] The driving force behind the transfer is the same as for all valence-related bonding: every element seeks the low-energy, filled-shell electron configuration of a noble gas. Metals achieve that configuration by shedding a few outer electrons; non-metals achieve it by gaining a few. The transferred electrons make both parties stable - but in the process each acquires a charge, and those opposite charges then lock the pair together electrostatically.

The mineral halite (NaCl) illustrates the mechanism clearly. Sodium’s electron configuration is a neon noble-gas core plus a single electron in the 3s subshell. Losing that one electron gives sodium the complete neon configuration and a net charge of +1. Chlorine has a completely filled K, L shell, a filled 3s, and five electrons in the 3p subshell - one short of a full shell. The electron sodium discards is exactly what chlorine needs to fill its 3p subshell and acquire the argon configuration. ^[1] The result: Na⁺ and Cl^-, two ions with stable electron configurations and opposite charges that attract each other strongly.

The Attractive and Repulsive Forces

The magnitude of the electrostatic attraction between two ions is governed by Coulomb’s law: the attractive force F_A is proportional to the product of the ion charges and inversely proportional to the square of the distance between ion centres. ^[1] The larger the charges and the shorter the distance, the stronger the attraction. This is why divalent ions (charge ±2) form stronger ionic bonds than monovalent ions at equivalent distances.

The attraction alone would pull the ions together until they collapsed into each other. What prevents this is the Born repulsion: when the electron shells of the anion and cation begin to overlap, repulsion grows rapidly because electrons of like charge cannot occupy the same space without a large energy cost. ^[1] The repulsive force increases much more steeply with decreasing distance than the attractive force does, so it overwhelms the attraction at short range. The equilibrium bond length is the distance at which these two forces exactly balance - where the net force is zero. For Na⁺-Cl^-, this equilibrium distance is about 2.8 Å. ^[1] At greater distances the net force is attractive, pulling the ions back together; at shorter distances the net force is repulsive, pushing them apart. The bond sits at the energy minimum between those two regimes.

Charge Balance and Ion Packing

Ionic compounds must be electrically neutral overall, so ions bond in ratios that ensure the total positive charge equals the total negative charge. ^[1] One Na⁺ balances one Cl^-, giving NaCl. In fluorite, a Ca²⁺ carries twice the charge of a single F^-, so two fluoride ions are needed for every calcium, giving CaF₂. The formula of an ionic mineral therefore encodes both the charge balance and the stoichiometry of the constituent ions.

In the solid state, ions behave as charged spheres and pack together in the most symmetrical arrangement that satisfies charge balance, with positive and negative charges alternating through the structure. ^[1] In halite, this produces a face-centred cubic arrangement in which each Na⁺ is surrounded by six Cl^- neighbours and each Cl^- is surrounded by six Na⁺ neighbours. ^[1] This mutual six-fold coordination is not arbitrary - it reflects the relative sizes of the two ions and the geometry that minimises repulsion while maximising attraction.

Physical Consequences: Brittleness and Cleavage

The alternating charge arrangement that makes ionic structures stable also makes them brittle. If a force attempts to slide one plane of the crystal past an adjacent plane, cations that were once aligned with anions will shift to be aligned with other cations - and like charges repel. ^[1] The crystal resists plastic deformation. Push hard enough and instead of bending it snaps. Because ionic structures are often highly ordered and regular, fracture tends to follow specific crystallographic planes - those where the bonds across the plane are weakest or fewest. This is the mechanical origin of cleavage. ^[1] Halite cleaves perfectly along three mutually perpendicular planes parallel to the faces of the cubic unit cell, which is exactly what you observe when you split a halite crystal.

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