Physics:Quantum antibonding molecular orbital - Revision history

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	Antibonding molecular orbitals (MOs) are normally ''higher'' in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. 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 [[Physics:Quantum number\|quantum number]]s.<ref>{{cite web \| url=https://www.chemistry.mcmaster.ca/esam/Chapter_6/section_2.html \| title=The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding }}</ref> Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1''s'') 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 H<sub>2</sub>. The higher-energy orbital is the antibonding orbital, which is less stable and opposes bonding if it is occupied. In a molecule such as H<sub>2</sub>, the two electrons normally occupy the lower-energy bonding orbital, so that the molecule is more stable than the separate H atoms.		Antibonding molecular orbitals (MOs) are normally ''higher'' in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. 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 [[Physics:Quantum number\|quantum number]]s.<ref>{{cite web \| url=https://www.chemistry.mcmaster.ca/esam/Chapter_6/section_2.html \| title=The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding }}</ref> Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1''s'') 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 H<sub>2</sub>. The higher-energy orbital is the antibonding orbital, which is less stable and opposes bonding if it is occupied. In a molecule such as H<sub>2</sub>, the two electrons normally occupy the lower-energy bonding orbital, so that the molecule is more stable than the separate H atoms.

	[[File:He2 antibonding orbital.svg\|250px\|right\|thumb\|He<sub>2</sub> 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.<ref>{{cite journal \| doi=10.3390/molecules25112667 \| doi-access=free \| title=The Basics of Covalent Bonding in Terms of Energy and Dynamics \| year=2020 \| last1=Nordholm \| first1=Sture \| last2=Bacskay \| first2=George B. \| journal=Molecules \| volume=25 \| issue=11 \| page=2667 \| pmid=32521828 \| pmc=7321125 }}</ref> 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.		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.<ref>{{cite journal \| doi=10.3390/molecules25112667 \| doi-access=free \| title=The Basics of Covalent Bonding in Terms of Energy and Dynamics \| year=2020 \| last1=Nordholm \| first1=Sture \| last2=Bacskay \| first2=George B. \| journal=Molecules \| volume=25 \| issue=11 \| page=2667 \| pmid=32521828 \| pmc=7321125 }}</ref> 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.

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	* Orchin, M. Jaffe, H.H. (1967) ''The Importance of Antibonding Orbitals''. Houghton Mifflin. ISBN B0006BPT5O		* Orchin, M. Jaffe, H.H. (1967) ''The Importance of Antibonding Orbitals''. Houghton Mifflin. ISBN B0006BPT5O

	~~{{Chemical bonding theory}}~~

	~~[[Category:Chemical bonding]]~~

	{{Sourceattribution\|Antibonding molecular orbital}}		{{Sourceattribution\|Antibonding molecular orbital}}

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	[[Image:Dihydrogen-LUMO-phase-3D-balls.png\|thumb\|right\|150px\|H<sub>2</sub> 1sσ* antibonding molecular orbital]]		[[Image:Dihydrogen-LUMO-phase-3D-balls.png\|thumb\|right\|150px\|H<sub>2</sub> 1sσ* antibonding molecular orbital]]

	In ~~[[Chemistry:Theoretical chemistry\|~~theoretical chemistry]], an '''antibonding orbital''' is a type of ~~[[Chemistry:Molecular orbital\|~~molecular orbital]] that weakens the ~~[[Chemistry:Chemical bond\|~~chemical bond]] between two ~~[[Atom\|atom]]s~~ and helps to raise the ~~[[Physics:Energy level\|~~energy]] of the ~~[[Physics:Molecule\|~~molecule]] relative to the separated atoms. Such an orbital has one or more ~~[[Node\|node]]s~~ in the bonding region between the ~~[[Physics:Atomic nucleus\|~~nuclei]]. The ~~[[Physics:Electron~~ density~~\|density]]~~ of the ~~[[Physics:Electron\|electron]]s~~ 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.<ref>Atkins P. and de Paula J. ''Atkins Physical Chemistry''. 8th ed. (W.H. Freeman 2006), p.371 {{ISBN\|0-7167-8759-8}}</ref><ref>Miessler G.L. and Tarr D.A., ''Inorganic Chemistry'' 2nd ed. (Prentice-Hall 1999), p.111 {{ISBN\|0-13-841891-8}}</ref> This is in contrast to a ~~[[Chemistry:Bonding molecular orbital\|~~bonding molecular orbital]], which has a lower energy than that of the separate atoms, and is responsible for chemical bonds.		In theoretical chemistry, 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.<ref>Atkins P. and de Paula J. ''Atkins Physical Chemistry''. 8th ed. (W.H. Freeman 2006), p.371 {{ISBN\|0-7167-8759-8}}</ref><ref>Miessler G.L. and Tarr D.A., ''Inorganic Chemistry'' 2nd ed. (Prentice-Hall 1999), p.111 {{ISBN\|0-13-841891-8}}</ref> 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.
	</div>		</div>

	<div style="width:300px;">		<div style="width:300px;">
	[[File:~~File not found~~.png\|thumb\|280px\|~~No image available~~.]]		[[File:Butadiene-pi-MOs-Spartan-3D-balls.png\|thumb\|280px\|antibonding molecular orbital in the Quantum Collection.]]
	</div>		</div>

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	==Diatomic molecules==		==Diatomic molecules==
	Antibonding ~~[[Chemistry:Molecular orbital\|~~molecular ~~orbital]]s~~ (MOs) are normally ''higher'' in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. If two ~~[[Software:Hydrogen\|~~hydrogen]] atoms are initially far apart, they have identical ~~[[Physics:Atomic orbital\|~~atomic ~~orbital]]s~~. However, as the spacing between the two atoms becomes smaller, the electron ~~[[Wave function\|~~wave ~~function]]s~~ begin to overlap. The ~~[[Physics:Pauli exclusion principle\|~~Pauli exclusion principle]] prohibits any two electrons (e-) in a molecule from having the same set of [[Physics:Quantum number\|quantum number]]s.<ref>{{cite web \| url=https://www.chemistry.mcmaster.ca/esam/Chapter_6/section_2.html \| title=The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding }}</ref> Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1''s'') 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 H<sub>2</sub>. The higher-energy orbital is the antibonding orbital, which is less stable and opposes bonding if it is occupied. In a molecule such as H<sub>2</sub>, the two electrons normally occupy the lower-energy bonding orbital, so that the molecule is more stable than the separate H atoms.		Antibonding molecular orbitals (MOs) are normally ''higher'' in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. 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 [[Physics:Quantum number\|quantum number]]s.<ref>{{cite web \| url=https://www.chemistry.mcmaster.ca/esam/Chapter_6/section_2.html \| title=The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding }}</ref> Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1''s'') 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 H<sub>2</sub>. The higher-energy orbital is the antibonding orbital, which is less stable and opposes bonding if it is occupied. In a molecule such as H<sub>2</sub>, the two electrons normally occupy the lower-energy bonding orbital, so that the molecule is more stable than the separate H atoms.

	[[File:He2 antibonding orbital.svg\|250px\|right\|thumb\|He<sub>2</sub> electron configuration. The four electrons occupy one bonding orbital at lower energy, and one antibonding orbital at higher energy than the atomic orbitals.]]		[[File:He2 antibonding orbital.svg\|250px\|right\|thumb\|He<sub>2</sub> 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 ~~[[Physics:Electron density\|~~electron density]] between the two nuclei than there would be if there were no bonding interaction at all.<ref>{{cite journal \| doi=10.3390/molecules25112667 \| doi-access=free \| title=The Basics of Covalent Bonding in Terms of Energy and Dynamics \| year=2020 \| last1=Nordholm \| first1=Sture \| last2=Bacskay \| first2=George B. \| journal=Molecules \| volume=25 \| issue=11 \| page=2667 \| pmid=32521828 \| pmc=7321125 }}</ref> 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.		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.<ref>{{cite journal \| doi=10.3390/molecules25112667 \| doi-access=free \| title=The Basics of Covalent Bonding in Terms of Energy and Dynamics \| year=2020 \| last1=Nordholm \| first1=Sture \| last2=Bacskay \| first2=George B. \| journal=Molecules \| volume=25 \| issue=11 \| page=2667 \| pmid=32521828 \| pmc=7321125 }}</ref> 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.


	==Polyatomic molecules==		==Polyatomic molecules==
	[[File:Butadiene-pi-MOs-Spartan-3D-balls.png\|thumb\|right\|200px\|Butadiene pi molecular orbitals. The two colors show opposite signs of the wavefunction.]]		[[File:Butadiene-pi-MOs-Spartan-3D-balls.png\|thumb\|right\|200px\|Butadiene pi molecular orbitals. The two colors show opposite signs of the wavefunction.]]
	In molecules with several atoms, some orbitals may be ~~[[Physics:Delocalized electron\|~~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''.		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, ~~[[Chemistry:Butadiene\|~~butadiene]] has pi orbitals which are delocalized over all four carbon atoms. There are two bonding pi orbitals which are occupied in the ~~[[Physics:Ground state\|~~ground state]]: π<sub>1</sub> is bonding between all carbons, while π<sub>2</sub> is bonding between C<sub>1</sub> and C<sub>2</sub> and between C<sub>3</sub> and C<sub>4</sub>, and antibonding between C<sub>2</sub> and C<sub>3</sub>. There are also antibonding pi orbitals with two and three antibonding interactions as shown in the diagram; these are vacant in the ~~[[Physics:Ground state\|~~ground state]], but may be occupied in ~~[[Physics:Excited state\|~~excited ~~state]]s~~.		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: π<sub>1</sub> is bonding between all carbons, while π<sub>2</sub> is bonding between C<sub>1</sub> and C<sub>2</sub> and between C<sub>3</sub> and C<sub>4</sub>, and antibonding between C<sub>2</sub> and C<sub>3</sub>. 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 ~~[[Chemistry:Benzene\|~~benzene]] with six carbon atoms has three bonding pi orbitals and three antibonding pi orbitals. Since each ~~[[Chemistry:Carbon\|~~carbon]] atom contributes one electron to the ~~[[Chemistry:Pi bond\|~~π-system]] of benzene, there are six pi electrons which fill the three lowest-energy pi molecular orbitals (the bonding pi orbitals).		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 ~~[[Chemistry:Chemical reaction\|~~chemical ~~reaction]]s~~ in terms of molecular orbital theory. [[Biography:Roald Hoffmann\|Roald Hoffmann]] and [[Biography:Kenichi Fukui\|Kenichi Fukui]] shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions.<ref>{{cite web \| title = The Nobel Prize in Chemistry 1981 \| publisher = Nobelprize.org \| url = http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| access-date = 15 March 2022 \| archive-url = https://web.archive.org/web/20081221120050/http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| archive-date = 21 December 2008 \| url-status = live }}</ref>		Antibonding orbitals are also important for explaining chemical reactions in terms of molecular orbital theory. [[Biography:Roald Hoffmann\|Roald Hoffmann]] and [[Biography:Kenichi Fukui\|Kenichi Fukui]] shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions.<ref>{{cite web \| title = The Nobel Prize in Chemistry 1981 \| publisher = Nobelprize.org \| url = http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| access-date = 15 March 2022 \| archive-url = https://web.archive.org/web/20081221120050/http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| archive-date = 21 December 2008 \| url-status = live }}</ref>

	== See also ==		== See also ==

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	Antibonding orbitals are also important for explaining [[Chemistry:Chemical reaction\|chemical reaction]]s in terms of molecular orbital theory. [[Biography:Roald Hoffmann\|Roald Hoffmann]] and [[Biography:Kenichi Fukui\|Kenichi Fukui]] shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions.<ref>{{cite web \| title = The Nobel Prize in Chemistry 1981 \| publisher = Nobelprize.org \| url = http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| access-date = 15 March 2022 \| archive-url = https://web.archive.org/web/20081221120050/http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| archive-date = 21 December 2008 \| url-status = live }}</ref>		Antibonding orbitals are also important for explaining [[Chemistry:Chemical reaction\|chemical reaction]]s in terms of molecular orbital theory. [[Biography:Roald Hoffmann\|Roald Hoffmann]] and [[Biography:Kenichi Fukui\|Kenichi Fukui]] shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions.<ref>{{cite web \| title = The Nobel Prize in Chemistry 1981 \| publisher = Nobelprize.org \| url = http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| access-date = 15 March 2022 \| archive-url = https://web.archive.org/web/20081221120050/http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html \| archive-date = 21 December 2008 \| url-status = live }}</ref>

	==See also==		== See also ==
	*[[Chemistry:~~Bonding molecular orbital~~\|~~Bonding molecular orbital]]~~		{{#invoke:PhysicsQC\|tocHeadingAndList\|Physics:Quantum basics/See also/Matter}}
	*[[Physics:~~Valence and conduction bands\|Valence and conduction bands]]~~
	*[[Chemistry:Valence bond theory\|Valence bond theory]]
	*[[Chemistry:Molecular orbital theory\|Molecular orbital theory]]
	*[[Chemistry:Conjugated system\|Conjugated system]]

	==References==		==References==

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 {{Short description|Molecular orbital which weakens chemical bonding}}
 [[Image:Dihydrogen-LUMO-phase-3D-balls.png|thumb|right|150px|H<sub>2</sub> 1sσ* antibonding molecular orbital]]
 In [[Chemistry:Theoretical chemistry|theoretical chemistry]], an '''antibonding orbital''' is a type of [[Chemistry:Molecular orbital|molecular orbital]] that weakens the [[Chemistry:Chemical bond|chemical bond]] between two [[Atom|atom]]s and helps to raise the [[Physics:Energy level|energy]] of the [[Physics:Molecule|molecule]] relative to the separated atoms. Such an orbital has one or more [[Node|node]]s in the bonding region between the [[Physics:Atomic nucleus|nuclei]]. The [[Physics:Electron density|density]] of the [[Physics:Electron|electron]]s 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.<ref>Atkins P. and de Paula J. ''Atkins Physical Chemistry''. 8th ed. (W.H. Freeman 2006), p.371 {{ISBN|0-7167-8759-8}}</ref><ref>Miessler G.L. and Tarr D.A., ''Inorganic Chemistry'' 2nd ed. (Prentice-Hall 1999), p.111 {{ISBN|0-13-841891-8}}</ref>  This is in contrast to a [[Chemistry:Bonding molecular orbital|bonding molecular orbital]], which has a lower energy than that of the separate atoms, and is responsible for chemical bonds.
 ==Diatomic molecules==

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imported>WikiHarold: WikiHarold moved page Physics:Antibonding molecular orbital to Physics:Quantum antibonding molecular orbital without leaving a redirect

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WikiHarold moved page Physics:Antibonding molecular orbital to Physics:Quantum antibonding molecular orbital without leaving a redirect

New page

{{Short description|Molecular orbital which weakens chemical bonding}}
[[Image:Dihydrogen-LUMO-phase-3D-balls.png|thumb|right|150px|H2 1sσ* antibonding molecular orbital]]

In [[Chemistry:Theoretical chemistry|theoretical chemistry]], an '''antibonding orbital''' is a type of [[Chemistry:Molecular orbital|molecular orbital]] that weakens the [[Chemistry:Chemical bond|chemical bond]] between two [[Atom|atom]]s and helps to raise the [[Physics:Energy level|energy]] of the [[Physics:Molecule|molecule]] relative to the separated atoms. Such an orbital has one or more [[Node|node]]s in the bonding region between the [[Physics:Atomic nucleus|nuclei]]. The [[Physics:Electron density|density]] of the [[Physics:Electron|electron]]s 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.<ref>Atkins P. and de Paula J. ''Atkins Physical Chemistry''. 8th ed. (W.H. Freeman 2006), p.371 {{ISBN|0-7167-8759-8}}</ref><ref>Miessler G.L. and Tarr D.A., ''Inorganic Chemistry'' 2nd ed. (Prentice-Hall 1999), p.111 {{ISBN|0-13-841891-8}}</ref> This is in contrast to a [[Chemistry:Bonding molecular orbital|bonding molecular orbital]], which has a lower energy than that of the separate atoms, and is responsible for chemical bonds.

==Diatomic molecules==
Antibonding [[Chemistry:Molecular orbital|molecular orbital]]s (MOs) are normally ''higher'' in energy than bonding molecular orbitals. Bonding and antibonding orbitals form when atoms combine into molecules. If two [[Software:Hydrogen|hydrogen]] atoms are initially far apart, they have identical [[Physics:Atomic orbital|atomic orbital]]s. However, as the spacing between the two atoms becomes smaller, the electron [[Wave function|wave function]]s begin to overlap. The [[Physics:Pauli exclusion principle|Pauli exclusion principle]] prohibits any two electrons (e-) in a molecule from having the same set of [[Physics:Quantum number|quantum number]]s.<ref>{{cite web | url=https://www.chemistry.mcmaster.ca/esam/Chapter_6/section_2.html | title=The Chemical Bond - the Effect of the Pauli Principle on Chemical Binding }}</ref> Therefore each original atomic orbital of the isolated atoms (for example, the ground state energy level, 1''s'') 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.

[[File:He2 antibonding orbital.svg|250px|right|thumb|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 [[Physics:Electron density|electron density]] between the two nuclei than there would be if there were no bonding interaction at all.<ref>{{cite journal | doi=10.3390/molecules25112667 | doi-access=free | title=The Basics of Covalent Bonding in Terms of Energy and Dynamics | year=2020 | last1=Nordholm | first1=Sture | last2=Bacskay | first2=George B. | journal=Molecules | volume=25 | issue=11 | page=2667 | pmid=32521828 | pmc=7321125 }}</ref> 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.

==Polyatomic molecules==
[[File:Butadiene-pi-MOs-Spartan-3D-balls.png|thumb|right|200px|Butadiene pi molecular orbitals. The two colors show opposite signs of the wavefunction.]]
In molecules with several atoms, some orbitals may be [[Physics:Delocalized electron|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, [[Chemistry:Butadiene|butadiene]] has pi orbitals which are delocalized over all four carbon atoms. There are two bonding pi orbitals which are occupied in the [[Physics:Ground state|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 [[Physics:Ground state|ground state]], but may be occupied in [[Physics:Excited state|excited state]]s.

Similarly [[Chemistry:Benzene|benzene]] with six carbon atoms has three bonding pi orbitals and three antibonding pi orbitals. Since each [[Chemistry:Carbon|carbon]] atom contributes one electron to the [[Chemistry:Pi bond|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 [[Chemistry:Chemical reaction|chemical reaction]]s in terms of molecular orbital theory. [[Biography:Roald Hoffmann|Roald Hoffmann]] and [[Biography:Kenichi Fukui|Kenichi Fukui]] shared the 1981 Nobel Prize in Chemistry for their work and further development of qualitative molecular orbital explanations for chemical reactions.<ref>{{cite web | title = The Nobel Prize in Chemistry 1981 | publisher = Nobelprize.org | url = http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html | access-date = 15 March 2022 | archive-url = https://web.archive.org/web/20081221120050/http://nobelprize.org/nobel_prizes/chemistry/laureates/1981/index.html | archive-date = 21 December 2008 | url-status = live }}</ref>

==See also==
*[[Chemistry:Bonding molecular orbital|Bonding molecular orbital]]
*[[Physics:Valence and conduction bands|Valence and conduction bands]]
*[[Chemistry:Valence bond theory|Valence bond theory]]
*[[Chemistry:Molecular orbital theory|Molecular orbital theory]]
*[[Chemistry:Conjugated system|Conjugated system]]

==References==
{{Reflist}}

==Further reading==

* Orchin, M. Jaffe, H.H. (1967) ''The Importance of Antibonding Orbitals''. Houghton Mifflin. ISBN B0006BPT5O

{{Chemical bonding theory}}

[[Category:Chemical bonding]]

{{Sourceattribution|Antibonding molecular orbital}}