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An ionic bond is a type of chemical bond formed through an electrostatic attraction between two oppositely charged ions. Ionic bonds are formed between a cation, which is usually a metal, and an anion, which is usually a nonmetal. Pure ionic bonding cannot exist: all ionic compounds have some degree of covalent bonding. Thus, an ionic bond is considered a bond where the ionic character is greater than the covalent character. The larger the difference in electronegativity between the two atoms involved in the bond, the more ionic (polar) the bond is. Bonds with partially ionic and partially covalent character are called polar covalent bonds. Ionic bonding is a form of noncovalent bonding.
A formation of an ionic bond proceeds when the atom of an element (usually metal), whose ionization energy is low, releases some of its electron(s) to achieve a stable electron configuration and after releasing some of its electron(s) the atom becomes a cation. The atom of another element (usually non metal), whose electron affinity is positive, then accepts the electron(s), again to attain a stable electron configuration and after accepting electron(s) the atom becomes an anion. Typically, the stable electron configuration is one of the noble gases for elements in the s-block and the p-block, and particular stable electron configurations for d-block and f-block elements. The electrostatic attraction between these two entities forms the ionic bond. It is formed to attain stable octate structure.
For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na+), and the chlorine atoms each gain an electron to form anions (Cl−). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).
- Na + Cl → Na+ + Cl− → NaCl
The removal of electrons from the cation is endothermic, raising the system's overall energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the action of the anion accepting the cation's valence electrons and the subsequent attraction of the ions to each other releases energy and thus lowers the overall energy of the system.
Ionic bonding will occur only if the overall energy change for the reaction is favorable – when the reaction is exothermic. The larger the resulting energy change, the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the reaction is most favorable between a metal and a non-metal.
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 MgO.
Bond strength 
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 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 inter-ionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The Born-Landé equation gives a reasonable fit to the lattice energy of e.g. sodium chloride where the calculated value is −756 kJ/mol which compares to −787 kJ/mol using the Born-Haber cycle.
Polarization effects 
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, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important 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 which refers to displacement of ions in the lattice due to the application of an electric field.
Comparison with covalent bonds 
In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons to attain stable electron configurations. In covalent bonding, the molecular geometry around each atom is determined by Valence shell electron pair repulsion VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.
Purely ionic bonds cannot exist, as the proximity of the entities involved in the bond allows some degree of sharing electron density between them. Therefore, all ionic bonds have some covalent character.
Thus, an ionic bond is considered a bond where the ionic character is greater than the covalent character. The larger the difference in electronegativity between the two atoms involved in the bond, the more ionic (polar) the bond is. Bonds with partially ionic and partially covalent character are called polar covalent bonds.
For example, Na–Cl and Mg–O bonds have a few percent covalency, while Si–O bonds are usually ~50% ionic and ~50% covalent.
Electrical conductivity 
Ionic compounds, if molten or dissolved, can conduct electricity because the ions in these conditions are free to move and carry electrons between the anode and the cathode. In the solid form, however, they cannot conduct because the electrons are held together too tightly for them to move. However, some ionic compounds can conduct electricity when solid. This is due to migration of the ions themselves under the influence of an electric field. These compounds are known as fast ion conductors.
See also 
- David Arthur Johnson, Metals and Chemical Change, Open University, Royal Society of Chemistry, 2002, ISBN 0-85404-665-8