Editing Ionic bonding (section)

==Formation==
Ionic bonding can result from a [[redox]] reaction when atoms of an element (usually [[metal]]), whose [[ionization energy]] is low, give some of their electrons to achieve a stable electron configuration. In doing so, cations are formed. An atom of another element (usually nonmetal) with greater [[electron affinity]] accepts one or more electrons to attain a stable [[electron configuration]], and after accepting electrons an 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 [[electron configuration|stable electron configurations]] for [[d-block]] and [[f-block]] elements. The electrostatic attraction between the anions and cations leads to the formation of a solid with a [[crystallographic lattice]] in which the ions are stacked in an alternating fashion. In such a lattice, it is usually not possible to distinguish discrete molecular units, so that the compounds formed are not molecular. However, the ions themselves can be complex and form molecular ions like the acetate anion or the ammonium cation.
[[Image:Ionic Bonding LiF.svg|thumb|right|250px|Representation of ionic bonding between [[lithium]] and [[fluorine]] to form [[lithium fluoride]]. Lithium has a low ionization energy and readily gives up its lone [[valence electron]] to a fluorine atom, which has a positive electron affinity and accepts the electron that was donated by the lithium atom. The end-result is that lithium is [[isoelectronicity|isoelectronic]] with [[helium]] and fluorine is isoelectronic with [[neon]]. Electrostatic interaction occurs between the two resulting ions, but typically aggregation is not limited to two of them. Instead, aggregation into a whole lattice held together by ionic bonding is the result.]] 
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<sup>+</sup>), and the chlorine atoms each gain an electron to form anions (Cl<sup>−</sup>). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).
: Na + Cl → Na<sup>+</sup> + Cl<sup>−</sup> → NaCl 

However, to maintain charge neutrality, strict ratios between anions and cations are observed so that ionic compounds, in general, obey the rules of stoichiometry despite not being molecular compounds. For compounds that are transitional to the alloys and possess mixed ionic and metallic bonding, this may not be the case anymore. Many sulfides, e.g., do form non-stoichiometric compounds.

Many ionic compounds are referred to as '''salts''' as they can also be formed by the neutralization reaction of an Arrhenius base like NaOH with an Arrhenius acid like HCl

: NaOH + HCl →  NaCl + H<sub>2</sub>O

The salt NaCl is then said to consist of the acid rest Cl<sup>−</sup> and the base rest Na<sup>+</sup>.

The removal of electrons to form 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's accepting the cation's valence electrons and the subsequent attraction of the ions to each other releases (lattice) 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. In general, the reaction is exothermic, but, e.g., the formation of mercuric oxide (HgO) is endothermic. The charge of the resulting ions is a major factor in the strength of ionic bonding, e.g. a salt C<sup>+</sup>A<sup>−</sup> is held together by electrostatic forces roughly four times weaker than C<sup>2+</sup>A<sup>2−</sup> according to [[Coulomb's law]], where C and A represent a generic cation and anion respectively. The sizes of the ions and the particular packing of the lattice are ignored in this rather simplistic argument.