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Chemical bonding involving attraction between ions

Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding, along with covalent bonding and metallic bonding. Ions are atoms (or groups of atoms) with an electrostatic charge. Atoms that gain electrons make negatively charged ions (called anions). Atoms that lose electrons make positively charged ions (called cations). This transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be more complex, e. g. molecular ions like NH4+ or SO42−. In simpler words, an ionic bond results from the transfer of electrons from a metal to a non metal to obtain a full valence shell for both atoms. Clean ionic bonding — in which one atom or molecule completely transfers an electron to another — cannot exist: all ionic compounds have some degree of covalent bonding or electron sharing. Thus, the term "ionic bonding" is given when the ionic character is greater than the covalent character – that is, a bond in which there is a large difference in electronegativity between the two atoms, causing the bonding to be more polar (ionic) than in covalent bonding where electrons are shared more equally. Bonds with partially ionic and partially covalent characters are called polar covalent bonds. Ionic compounds conduct electricity when molten or in solution, typically not when solid. Ionic compounds generally have a high melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the cohesive forces and the higher the melting point. They also tend to be soluble in water; the stronger the cohesive forces, the lower the solubility.

Atoms that have an almost full or almost empty valence shell tend to be very reactive. Strongly electronegative atoms (such as halogens) often have only one or two empty electron states in their valence shell, and frequently bond with other atoms or gain electrons to form anions. Weakly electronegative atoms (such as alkali metals) have relatively few valence electrons, which can easily be lost to strongly electronegative atoms. As a result, weakly electronegative atoms tend to distort their electron cloud and form cations.

They are considered to be among the strongest of all types of chemical bonds. This often causes ionic compounds to be very stable. Ionic bonds have high bond energy. Bond energy is the mean amount of energy required to break the bond in the gaseous state. Most ionic compounds exist in the form of a crystal structure, in which the ions occupy the corners of the crystal. Such a structure is called a crystal lattice. Ionic compounds lose their crystal lattice structure and break up into ions when dissolved in water or any other polar solvent. This process is called solvation. The presence of these free ions makes aqueous ionic compound solutions good conductors of electricity. The same occurs when the compounds are heated above their melting point in a process known as melting.

Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as magnesium oxide. Pauling's rules provide guidelines for predicting and rationalizing the crystal structures of ionic crystals

For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born–Haber cycle. It can also be calculated (predicted) using the Born–Landé equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short range repulsive potential energy term. The electrostatic potential can be expressed in terms of the interionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The further away from the nucleus the weaker the shield. The Born–Landé equation gives a reasonable fit to the lattice energy of, e. g., sodium chloride, where the calculated (predicted) value is −756 kJ/mol, which compares to −787 kJ/mol using the Born–Haber cycle. In aqueous solution the binding strength can be described by the Bjerrum or Fuoss equation as function of the ion charges, rather independent of the nature of the ions such as polarizability or size. The strength of salt bridges is most often evaluated by measurements of equilibria between molecules containing cationic and anionic sites, most often in solution. Equilibrium constants in water indicate additive free energy contributions for each salt bridge. Another method for the identification of hydrogen bonds also in complicated molecules is crystallography, sometimes also NMR spectroscopy. The attractive forces defining the strength of ionic bonding can be modeled by Coulomb's Law. Ionic bond strengths are typically (cited ranges vary) between 170 and 1500 kJ/mol.

Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised in Fajans' rules. This polarization of the negative ion leads to a build up of extra charge density between the two nuclei, that is, to partial covalency. Larger negative ions are more easily polarized, but the effect is usually important only when positive ions with charges of 3+ (e. g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e. g., LiI is ionic but has some covalent bonding present). Note that this is not the ionic polarization effect that refers to displacement of ions in the lattice due to the application of an electric field.