Ionic bonding

Evidence for the Existence of Ions

Direct experimental evidence supports the existence of ions as discrete charged species held together in ionic compounds by electrostatic forces.

The first line of evidence comes from physical properties. Ionic compounds possess high melting and boiling points because large amounts of energy are required to overcome the strong electrostatic attractions between oppositely charged ions in the lattice.

Ionic compounds conduct electricity only when molten or dissolved in water, never when solid. This is because the ions become free to move and carry charge in the liquid or aqueous state, but are locked in fixed positions in the solid lattice.

A second line of evidence comes from electron density maps, produced by X-ray diffraction. These maps show contour lines representing electron density around each ion. The contours form closed regions around each ion with a clear gap between them, demonstrating that electrons belong to discrete ions rather than being shared.

A third line of evidence is the migration of ions during electrolysis. When a coloured ionic compound such as copper(II) chromate(VI) is placed on moist filter paper between electrodes, the blue Cu²⁺ ions migrate towards the cathode and the yellow CrO₄²⁻ ions migrate towards the anode. This visible separation confirms that ions carry opposite charges.

Formation of Ions by Electron Transfer

Ions form when atoms lose or gain electrons to achieve a stable noble gas electronic configuration.

Metal atoms lose electrons from their outer shell to form positively charged cations. For example, sodium loses one electron to form Na⁺, and magnesium loses two electrons to form Mg²⁺.

Non-metal atoms gain electrons into their outer shell to form negatively charged anions. For example, chlorine gains one electron to form Cl⁻, and oxygen gains two electrons to form O²⁻.

The number of electrons lost or gained equals the magnitude of the ionic charge. After transfer, both species achieve the electronic configuration of the nearest noble gas, which is energetically more stable than the parent atom.

The overall process is summarised by half-equations such as Na → Na⁺ + e⁻ and Cl + e⁻ → Cl⁻. The total number of electrons lost by metals equals the total number gained by non-metals, ensuring that ionic compounds are electrically neutral overall.

Dot-and-Cross Diagrams of Ions

A dot-and-cross diagram for an ionic compound shows the transfer of outer-shell electrons from metal to non-metal, with each ion drawn separately inside square brackets with its charge.

In the convention, dots represent electrons originally from one atom and crosses represent electrons originally from the other. After the transfer, the cation is shown with its outer shell empty (or with the previous shell now becoming the outer shell), while the anion is shown with a complete outer octet.

Each ion is enclosed in square brackets with the charge shown as a superscript outside the bracket, for example [Na]⁺ or [Cl]²⁻ as appropriate. Only outer-shell electrons are usually drawn for clarity.

Giant Ionic Lattices

Ionic compounds exist as giant ionic lattices in which positive and negative ions are arranged in a regular, repeating three-dimensional pattern.

Each cation is surrounded by a fixed number of nearest-neighbour anions, and each anion is surrounded by the same number of nearest-neighbour cations. In sodium chloride, for example, each Na⁺ is surrounded by six Cl⁻ ions and each Cl⁻ is surrounded by six Na⁺ ions.

The lattice extends in all directions over the whole crystal, which is why it is described as “giant.” The structure is held together by the net electrostatic attractions between all the oppositely charged ions throughout the lattice.

Ionic Bonding as Electrostatic Attraction

Ionic bonding is the strong net electrostatic attraction between oppositely charged ions in a giant ionic lattice.

The word “net” is essential. Within the lattice, every ion experiences attractions to all the oppositely charged ions around it and repulsions from all the like-charged ions. The bond strength reflects the overall balance, in which attractions dominate.

Ionic bonding is non-directional, meaning the attraction acts equally in all directions around each ion. This is why ionic compounds form regular three-dimensional lattices rather than discrete molecules.

MisconceptionIonic bonding is not simply “the transfer of electrons” between atoms. Electron transfer forms the ions; the bond itself is the electrostatic attraction between the ions afterwards. Confusing the two costs a definition mark.

Exam TipDefine ionic bonding as the strong electrostatic attraction between oppositely charged ions, not as electron transfer.

Factors Affecting Ionic Bond Strength

The strength of ionic bonding depends on ionic charge and ionic radius, because both control the magnitude of the electrostatic attraction between ions.

Larger ionic charges produce stronger attractions. For example, the attraction between Mg²⁺ and O²⁻ is much stronger than that between Na⁺ and Cl⁻ because the charges are doubled on both ions.

Smaller ionic radii produce stronger attractions. When ions are smaller, their centres of charge can sit closer together in the lattice, which increases the strength of the electrostatic force between them.

Stronger ionic bonding means more energy is required to break apart the lattice. This is reflected in higher melting points, higher boiling points, and more exothermic lattice energies. For example, MgO has a much higher melting point than NaCl because Mg²⁺ and O²⁻ carry larger charges and have smaller radii than Na⁺ and Cl⁻.

Trends in Ionic Radii

Ionic radius changes systematically down a group and across a set of isoelectronic ions, reflecting changes in the number of electron shells and the nuclear charge experienced by the outer electrons.

Down a group, ionic radius increases. Each successive ion has one more occupied electron shell than the one above, so the outer electrons are further from the nucleus. For example, in Group 1, the radii increase Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺. The same pattern is seen in Group 7: F⁻ < Cl⁻ < Br⁻ < I⁻.

A set of isoelectronic ions all have the same number of electrons but different numbers of protons. The ions N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, and Al³⁺ all have ten electrons (the configuration of neon).

Across this isoelectronic series, ionic radius decreases as nuclear charge increases. With the same number of electrons but more protons, the nuclear charge per electron increases, so the electron cloud is pulled in more tightly. The order of decreasing radius is N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺.

Examiner InsightWhen asked to explain the radius decrease across isoelectronic ions, state explicitly that the number of electrons stays the same while the number of protons increases. Vague answers about “more protons attracting electrons” without identifying the constant electron count lose marks.

Exam TipName the isoelectronic feature first, then explain the change in nuclear charge.

Polarisation of Ions

Polarisation is the distortion of the electron cloud of an anion caused by a nearby cation, which pulls electron density towards itself.

In a perfectly ionic bond, the anion would be a spherical electron cloud unaffected by its neighbours. In reality, every cation exerts some attractive pull on the electron cloud of nearby anions. The resulting distortion concentrates electron density in the region between the two ions, introducing a degree of covalent character into the bond.

The polarising power of a cation is its ability to distort the electron cloud of a neighbouring anion. Polarising power increases with greater positive charge and smaller ionic radius. A small, highly charged cation such as Al³⁺ is far more polarising than a large, singly charged cation such as Cs⁺.

The polarisability of an anion is the ease with which its electron cloud can be distorted by a nearby cation. Polarisability increases with greater negative charge and larger ionic radius. A large, highly charged anion such as I⁻ or S²⁻ has loosely held outer electrons that are easily distorted, while a small anion such as F⁻ holds its electrons tightly and is only weakly polarisable.

The combined effect determines how much covalent character an ionic compound shows. For example, AlCl₃ has significant covalent character because Al³⁺ is highly polarising and Cl⁻ is moderately polarisable, while NaCl behaves as an almost purely ionic compound because Na⁺ has weak polarising power.

MnemonicSmall Highly-charged Cations Polarise Strongly.