Antibonding molecular orbital

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H2 1ss* antibonding molecular orbital Dihydrogen-LUMO-phase-3D-balls.png
H2 1sσ* antibonding molecular orbital

In chemical bonding theory, an antibonding orbital is a type of molecular orbital that weakens the chemical bond between two atoms and helps to raise the energy of the molecule relative to the separated atoms. Such an orbital has one or more nodes in the bonding region between the nuclei. The density of the electrons in the orbital is concentrated outside the bonding region and acts to pull one nucleus away from the other and tends to cause mutual repulsion between the two atoms. [1] [2] This is in contrast to a bonding molecular orbital, which has a lower energy than that of the separate atoms, and is responsible for chemical bonds.


Diatomic molecules

Antibonding molecular orbitals (MOs) are normally higher in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. [3] If two hydrogen atoms are initially far apart, they have identical atomic orbitals. However, as the spacing between the two atoms becomes smaller, the electron wave functions begin to overlap. The Pauli exclusion principle prohibits any two electrons (e-) in a molecule from having the same set of quantum numbers. [4] Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1s) splits into two molecular orbitals belonging to the pair, one lower in energy than the original atomic level and one higher. The orbital which is in a lower energy state than the orbitals of the separate atoms is the bonding orbital, which is more stable and promotes the bonding of the two H atoms into H2. The higher-energy orbital is the antibonding orbital, which is less stable and opposes bonding if it is occupied. In a molecule such as H2, the two electrons normally occupy the lower-energy bonding orbital, so that the molecule is more stable than the separate H atoms.

He2 electron configuration. The four electrons occupy one bonding orbital at lower energy, and one antibonding orbital at higher energy than the atomic orbitals. He2 antibonding orbital.png
He2 electron configuration. The four electrons occupy one bonding orbital at lower energy, and one antibonding orbital at higher energy than the atomic orbitals.

A molecular orbital becomes antibonding when there is less electron density between the two nuclei than there would be if there were no bonding interaction at all. [5] When a molecular orbital changes sign (from positive to negative) at a nodal plane between two atoms, it is said to be antibonding with respect to those atoms. Antibonding orbitals are often labelled with an asterisk (*) on molecular orbital diagrams.

In homonuclear diatomic molecules, σ* (sigma star) antibonding orbitals have no nodal planes passing through the two nuclei, like sigma bonds, and π* (pi star) orbitals have one nodal plane passing through the two nuclei, like pi bonds. The Pauli exclusion principle dictates that no two electrons in an interacting system may have the same quantum state. If the bonding orbitals are filled, then any additional electrons will occupy antibonding orbitals. This occurs in the He2 molecule, in which both the 1sσ and 1sσ* orbitals are filled. [6] Since the antibonding orbital is more antibonding than the bonding orbital is bonding, the molecule has a higher energy than two separated helium atoms, and it is therefore unstable.

Polyatomic molecules

Butadiene pi molecular orbitals. The two colors show opposite signs of the wavefunction. Butadiene-pi-MOs-Spartan-3D-balls.png
Butadiene pi molecular orbitals. The two colors show opposite signs of the wavefunction.

In molecules with several atoms, some orbitals may be delocalized over more than two atoms. A particular molecular orbital may be bonding with respect to some adjacent pairs of atoms and antibonding with respect to other pairs. If the bonding interactions outnumber the antibonding interactions, the MO is said to be bonding, whereas, if the antibonding interactions outnumber the bonding interactions, the molecular orbital is said to be antibonding.

For example, butadiene has pi orbitals which are delocalized over all four carbon atoms. There are two bonding pi orbitals which are occupied in the ground state: π1 is bonding between all carbons, while π2 is bonding between C1 and C2 and between C3 and C4, and antibonding between C2 and C3. There are also antibonding pi orbitals with two and three antibonding interactions as shown in the diagram; these are vacant in the ground state, but may be occupied in excited states.

Similarly benzene with six carbon atoms has three bonding pi orbitals and three antibonding pi orbitals. Since each carbon atom contributes one electron to the π-system of benzene, there are six pi electrons which fill the three lowest-energy pi molecular orbitals (the bonding pi orbitals).

Antibonding orbitals are also important for explaining chemical reactions in terms of molecular orbital theory. Roald Hoffmann and Kenichi Fukui shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions. [7]

See also

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Molecular orbital Wave-like behavior of an electron in a molecule

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Pi bond Type of chemical bond

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A non-bonding orbital, also known as non-bonding molecular orbital (NBMO), is a molecular orbital whose occupation by electrons neither increases nor decreases the bond order between the involved atoms. Non-bonding orbitals are often designated by the letter n in molecular orbital diagrams and electron transition notations. Non-bonding orbitals are the equivalent in molecular orbital theory of the lone pairs in Lewis structures. The energy level of a non-bonding orbital is typically in between the lower energy of a valence shell bonding orbital and the higher energy of a corresponding antibonding orbital. As such, a non-bonding orbital with electrons would commonly be a HOMO.

An electronic effect influences the structure, reactivity, or properties of molecule but is neither a traditional bond nor a steric effect. In organic chemistry, the term stereoelectronic effect is also used to emphasize the relation between the electronic structure and the geometry (stereochemistry) of a molecule.

Chemical bonding of water

Water is a simple triatomic bent molecule with C2v molecular symmetry and bond angle of 104.5° between the central oxygen atom and the hydrogen atoms. Despite being one of the simplest triatomic molecules, its chemical bonding scheme is nonetheless complex as many of its bonding properties such as bond angle, ionization energy, and electronic state energy cannot be explained by one unified bonding model. Instead, several traditional and advanced bonding models such as simple Lewis and VSEPR structure, valence bond theory, molecular orbital theory, isovalent hybridization, and Bent's rule are discussed below to provide a comprehensive bonding model for H
, explaining and rationalizing the various electronic and physical properties and features manifested by its peculiar bonding arrangements.

The bonding orbital is used in molecular orbital (MO) theory to describe the attractive interactions between the atomic orbitals of two or more atoms in a molecule. In MO theory, electrons are portrayed to move in waves. When more than one of these waves come close together, the in-phase combination of these waves produces an interaction that leads to a species that is greatly stabilized. The result of the waves’ constructive interference causes the density of the electrons to be found within the binding region, creating a stable bond between the two species.


  1. Atkins P. and de Paula J. Atkins Physical Chemistry. 8th ed. (W.H. Freeman 2006), p.371 ISBN   0-7167-8759-8
  2. Miessler G.L. and Tarr D.A., Inorganic Chemistry 2nd ed. (Prentice-Hall 1999), p.111 ISBN   0-13-841891-8
  3. "Molecular Orbital - an overview | ScienceDirect Topics".
  4. "The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding".
  5. Nordholm, Sture; Bacskay, George B. (2020). "The Basics of Covalent Bonding in Terms of Energy and Dynamics". Molecules. 25 (11): 2667. doi: 10.3390/molecules25112667 . PMC   7321125 . PMID   32521828.
  6. "2.1. Combining atomic orbitals, sigma and pi bonding | Organic Chemistry 1: An open textbook".
  7. "The Nobel Prize in Chemistry 1981". Archived from the original on 21 December 2008. Retrieved 15 March 2022.

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