Editing Molecular orbital theory (section)

==Types of orbitals==
[[File:MO diagram dihydrogen.png|thumb|right|300px|MO diagram showing the formation of molecular orbitals of H<sub>2</sub> (centre) from atomic orbitals of two H atoms. The lower-energy MO is bonding with electron density concentrated between the two H nuclei. The higher-energy MO is anti-bonding with electron density concentrated behind each H nucleus.]]
Molecular orbital (MO) theory uses a [[linear combination of atomic orbitals]] (LCAO) to represent molecular orbitals resulting from bonds between atoms. These are often divided into three types, [[bonding molecular orbital|bonding]], [[antibonding molecular orbital|antibonding]], and [[non-bonding orbital|non-bonding]]. A bonding orbital concentrates electron density in the region ''between'' a given pair of atoms, so that its electron density will tend to attract each of the two nuclei toward the other and hold the two atoms together.<ref name="Tarr 2013">Miessler and Tarr (2013), ''Inorganic Chemistry'', 5th ed, 117-165, 475-534.</ref> An anti-bonding orbital concentrates electron density "behind" each nucleus (i.e. on the side of each atom which is farthest from the other atom), and so tends to pull each of the two nuclei away from the other and actually weaken the bond between the two nuclei. Electrons in non-bonding orbitals tend to be associated with atomic orbitals that do not interact positively or negatively with one another, and electrons in these orbitals neither contribute to nor detract from bond strength.<ref name="Tarr 2013"/>

Molecular orbitals are further divided according to the types of [[atomic orbital]]s they are formed from. Chemical substances will form bonding interactions if their orbitals become lower in energy when they interact with each other. Different bonding orbitals are distinguished that differ by [[electron configuration]] (electron cloud shape) and by [[energy level]]s.

The molecular orbitals of a molecule can be illustrated in [[molecular orbital diagram]]s.

Common bonding orbitals are [[Sigma bond|sigma (σ) orbitals]] which are symmetric about the bond axis and [[pi bond|pi (π) orbitals]] with a [[Node (physics)|nodal plane]] along the bond axis. Less common are [[delta bond|delta (δ) orbitals]] and [[Phi bond|phi (φ) orbitals]] with two and three nodal planes respectively along the bond axis. Antibonding orbitals are signified by the addition of an asterisk. For example, an antibonding pi orbital may be shown as π*.