Physics:Quantum valence bond theory - Revision history

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Remove imported red links from Quantum page

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	[[File:Sigma bond.svg\|thumb\|left\|σ bond between two atoms: localization of electron density]]		[[File:Sigma bond.svg\|thumb\|left\|σ bond between two atoms: localization of electron density]]
	[[Image:Pi-Bond.svg\|thumb\|Two p-orbitals forming a ~~{{pi}}~~-bond.]]		[[Image:Pi-Bond.svg\|thumb\|Two p-orbitals forming a π-bond.]]
	The overlapping atomic orbitals can differ. The two types of overlapping orbitals are sigma and pi. Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head, with the electron density most concentrated between nuclei. Pi bonds occur when two orbitals overlap when they are parallel.<ref>{{Cite book \|last=House \|first=J. E. \|title=Inorganic chemistry \|date=2013 \|publisher=Elsevier/Academic Press \|isbn=978-0-12-385110-9 \|edition=2nd \|location=Waltham, MA}}</ref> For example, a bond between two ''s''-orbital electrons is a sigma bond, because two spheres are always coaxial. In terms of bond order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond, and triple bonds contain one sigma bond and two pi bonds. However, the atomic orbitals for bonding may be hybrids. Hybridization is a model that describes how atomic orbitals combine to form new orbitals that better match the geometry of molecules. Atomic orbitals that are similar in energy combine to make hybrid orbitals. For example, the carbon in methane (CH<sub>4</sub>) undergoes sp<sup>3</sup> hybridization to form four equivalent orbitals, resulting in a tetrahedral shape. Different types of hybridization, such as sp, sp<sup>2</sup>, and sp<sup>3</sup>, correspond to specific molecular geometries (linear, trigonal planar, and tetrahedral), influencing the bond angles observed in molecules. Hybrid orbitals provide additional directionality to sigma bonds, accurately explaining molecular geometries.<ref>{{Cite journal \|last=Barnes \|first=Craig \|date=2003 \|title=Inorganic Chemistry (Housecroft, Catherine E.; Sharpe, Alan G.) \|url=http://dx.doi.org/10.1021/ed080p747 \|journal=Journal of Chemical Education \|volume=80 \|issue=7 \|pages=747 \|doi=10.1021/ed080p747 \|issn=0021-9584\|url-access=subscription }}</ref>		The overlapping atomic orbitals can differ. The two types of overlapping orbitals are sigma and pi. Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head, with the electron density most concentrated between nuclei. Pi bonds occur when two orbitals overlap when they are parallel.<ref>{{Cite book \|last=House \|first=J. E. \|title=Inorganic chemistry \|date=2013 \|publisher=Elsevier/Academic Press \|isbn=978-0-12-385110-9 \|edition=2nd \|location=Waltham, MA}}</ref> For example, a bond between two ''s''-orbital electrons is a sigma bond, because two spheres are always coaxial. In terms of bond order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond, and triple bonds contain one sigma bond and two pi bonds. However, the atomic orbitals for bonding may be hybrids. Hybridization is a model that describes how atomic orbitals combine to form new orbitals that better match the geometry of molecules. Atomic orbitals that are similar in energy combine to make hybrid orbitals. For example, the carbon in methane (CH<sub>4</sub>) undergoes sp<sup>3</sup> hybridization to form four equivalent orbitals, resulting in a tetrahedral shape. Different types of hybridization, such as sp, sp<sup>2</sup>, and sp<sup>3</sup>, correspond to specific molecular geometries (linear, trigonal planar, and tetrahedral), influencing the bond angles observed in molecules. Hybrid orbitals provide additional directionality to sigma bonds, accurately explaining molecular geometries.<ref>{{Cite journal \|last=Barnes \|first=Craig \|date=2003 \|title=Inorganic Chemistry (Housecroft, Catherine E.; Sharpe, Alan G.) \|url=http://dx.doi.org/10.1021/ed080p747 \|journal=Journal of Chemical Education \|volume=80 \|issue=7 \|pages=747 \|doi=10.1021/ed080p747 \|issn=0021-9584\|url-access=subscription }}</ref>

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	If many terms are considered in the wave functions, the two theories approach mathematical equivalence, However MO is a more popular approach than VB due to its easier implementation in the early days of computational chemistry.		If many terms are considered in the wave functions, the two theories approach mathematical equivalence, However MO is a more popular approach than VB due to its easier implementation in the early days of computational chemistry.

	Valence bond theory views aromatic properties of molecules as due to spin coupling of the ~~{{pi}}~~ orbitals.<ref>{{cite journal\|last1=Cooper\|first1=David L.\|last2=Gerratt\|first2=Joseph\|last3=Raimondi\|first3=Mario\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=323\|page=699\|year=1986\|doi=10.1038/323699a0\|issue=6090\|bibcode = 1986Natur.323..699C \|s2cid=24349360}}</ref><ref>{{cite journal\|last1=Pauling\|first1=Linus\|title=Electronic structure of the benzene molecule\|journal=Nature\|volume=325\|page=396\|year=1987\|doi=10.1038/325396d0\|issue=6103\|bibcode = 1987Natur.325..396P \|s2cid=4261220\|doi-access=free}}</ref><ref>{{cite journal\|last1=Messmer\|first1=Richard P.\|last2=Schultz\|first2=Peter A.\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=329\|page=492\|year=1987\|doi=10.1038/329492a0\|issue=6139\|bibcode = 1987Natur.329..492M \|s2cid=45218186\|doi-access=free}}</ref><ref>{{cite journal\|last1=Harcourt\|first1=Richard D.\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=329\|page=491\|year=1987\|doi=10.1038/329491b0\|issue=6139\|bibcode = 1987Natur.329..491H \|s2cid=4268597}}</ref> This is essentially still the old idea of resonance between Friedrich August Kekulé von Stradonitz and James Dewar structures. In contrast, molecular orbital theory views aromaticity as delocalization of the ~~{{pi}}~~-electrons.<ref>{{Cite journal \|last=Vemulapalli \|first=G. K. \|date=2008-10-01 \|title=Theories of the chemical bond and its true nature \|url=http://link.springer.com/10.1007/s10698-008-9049-2 \|journal=Foundations of Chemistry \|language=en \|volume=10 \|issue=3 \|pages=167–176 \|doi=10.1007/s10698-008-9049-2 \|issn=1386-4238\|url-access=subscription }}</ref> Valence bond treatments are restricted to relatively small molecules, largely due to the lack of orthogonality between valence bond orbitals and between valence bond structures. The molecular orbitals are always orthogonal.		Valence bond theory views aromatic properties of molecules as due to spin coupling of the π orbitals.<ref>{{cite journal\|last1=Cooper\|first1=David L.\|last2=Gerratt\|first2=Joseph\|last3=Raimondi\|first3=Mario\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=323\|page=699\|year=1986\|doi=10.1038/323699a0\|issue=6090\|bibcode = 1986Natur.323..699C \|s2cid=24349360}}</ref><ref>{{cite journal\|last1=Pauling\|first1=Linus\|title=Electronic structure of the benzene molecule\|journal=Nature\|volume=325\|page=396\|year=1987\|doi=10.1038/325396d0\|issue=6103\|bibcode = 1987Natur.325..396P \|s2cid=4261220\|doi-access=free}}</ref><ref>{{cite journal\|last1=Messmer\|first1=Richard P.\|last2=Schultz\|first2=Peter A.\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=329\|page=492\|year=1987\|doi=10.1038/329492a0\|issue=6139\|bibcode = 1987Natur.329..492M \|s2cid=45218186\|doi-access=free}}</ref><ref>{{cite journal\|last1=Harcourt\|first1=Richard D.\|title=The electronic structure of the benzene molecule\|journal=Nature\|volume=329\|page=491\|year=1987\|doi=10.1038/329491b0\|issue=6139\|bibcode = 1987Natur.329..491H \|s2cid=4268597}}</ref> This is essentially still the old idea of resonance between Friedrich August Kekulé von Stradonitz and James Dewar structures. In contrast, molecular orbital theory views aromaticity as delocalization of the π-electrons.<ref>{{Cite journal \|last=Vemulapalli \|first=G. K. \|date=2008-10-01 \|title=Theories of the chemical bond and its true nature \|url=http://link.springer.com/10.1007/s10698-008-9049-2 \|journal=Foundations of Chemistry \|language=en \|volume=10 \|issue=3 \|pages=167–176 \|doi=10.1007/s10698-008-9049-2 \|issn=1386-4238\|url-access=subscription }}</ref> Valence bond treatments are restricted to relatively small molecules, largely due to the lack of orthogonality between valence bond orbitals and between valence bond structures. The molecular orbitals are always orthogonal.

	Valence bond theory cannot explain electronic transitions and spectroscopic properties as effectively as MO theory. While VB employs hybridization to explain bonding, it can oversimplify complex bonding situations, limiting its applicability in more intricate molecular geometries such as transition metal compounds.<ref name=":12" />		Valence bond theory cannot explain electronic transitions and spectroscopic properties as effectively as MO theory. While VB employs hybridization to explain bonding, it can oversimplify complex bonding situations, limiting its applicability in more intricate molecular geometries such as transition metal compounds.<ref name=":12" />
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	==Computational approaches==		==Computational approaches==
	~~{{Main\|Chemistry:Modern valence bond theory}}~~
	Modern valence bond theory replaces the overlapping atomic orbitals by overlapping valence bond orbitals that are expanded over a large number of basis functions, either centered each on one atom to give a classical valence bond picture, or centered on all atoms in the molecule. The resulting energies are more competitive with energies from calculations where electron correlation is introduced based on a Hartree–Fock reference wavefunction. The most recent text is by Shaik and Hiberty.<ref>{{cite book\| last = Shaik\| first = Sason S.\|author2=Phillipe C. Hiberty\| title = A Chemist's Guide to Valence Bond Theory\| publisher = Wiley-Interscience		Modern valence bond theory replaces the overlapping atomic orbitals by overlapping valence bond orbitals that are expanded over a large number of basis functions, either centered each on one atom to give a classical valence bond picture, or centered on all atoms in the molecule. The resulting energies are more competitive with energies from calculations where electron correlation is introduced based on a Hartree–Fock reference wavefunction. The most recent text is by Shaik and Hiberty.<ref>{{cite book\| last = Shaik\| first = Sason S.\|author2=Phillipe C. Hiberty\| title = A Chemist's Guide to Valence Bond Theory\| publisher = Wiley-Interscience
	\| year = 2008\| location = New Jersey\| isbn = 978-0-470-03735-5}}</ref>		\| year = 2008\| location = New Jersey\| isbn = 978-0-470-03735-5}}</ref>
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	==References==		==References==
	{{Reflist}}		{{Reflist}}

	~~{{Chemical bonding theory\|state=expanded}}~~
	~~{{Linus Pauling}}~~

	~~[[Category:Quantum chemistry]]~~
	~~[[Category:Chemical bonding]]~~
	~~[[Category:General chemistry]]~~

	{{Sourceattribution\|Valence bond theory}}		{{Sourceattribution\|Valence bond theory}}

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	<div style="flex:1; line-height:1.45; color:#006b45; column-count:2; column-gap:32px; column-rule:1px solid #b8d8c8;">		<div style="flex:1; line-height:1.45; color:#006b45; column-count:2; column-gap:32px; column-rule:1px solid #b8d8c8;">
	~~{{Redirect\|VBT}}~~		Valence bond theory is a Book II topic in the Quantum Collection. It explains chemical bonding by combining atomic orbitals into localized electron-pair bonds between atoms. In quantum mechanics a valence bond description uses spin pairing, orbital overlap, resonance structures, and hybridization to describe molecular structure. It complements molecular orbital theory, which often uses delocalized orbitals. Valence bond theory is useful for understanding covalent bonds, bond angles, resonance, reaction mechanisms, and the connection between chemical diagrams and many-electron wavefunctions.
	~~{{Electronic structure methods}}~~
	~~In [[HandWiki:Chemistry\|chemistry]], '''valence~~ bond ~~(VB)~~ theory~~'''~~ is ~~one of the two basic theories, along with [[Chemistry:Molecular orbital theory\|molecular orbital (MO) theory]], that were developed to use~~ the ~~methods of [[Physics:~~Quantum ~~mechanics\|quantum mechanics]] to describe [[Chemistry:Chemical bond\|chemical bond]]ing~~. It ~~focuses on how the [[Physics:Atomic orbital\|~~atomic ~~orbital]]s of the dissociated~~ atoms ~~combine to give individual chemical bonds when a molecule is formed~~. In ~~contrast~~, ~~[[Chemistry:Molecular~~ orbital ~~theory\|~~molecular orbital theory~~]] has~~ orbitals ~~that cover~~ the whole molecule.<ref>{{cite book\| last1 = Murrell\| first1 = J. N.\| last2 = Kettle\| first2 = S. F. A.\| last3 = Tedder\| first3 = J. M.\| title = The Chemical Bond\| edition = 2nd\| publisher = John Wiley & Sons\| year = 1985\| isbn = 0-471-90759-6\| url-~~access = registration\| url = https://archive~~.~~org/details/chemicalbond0000murr_e8r6}}</ref>~~
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	== See also ==		== See also ==
	* [[Chemistry:~~Valence bond programs~~\|~~Valence bond programs]]~~		{{#invoke:PhysicsQC\|tocHeadingAndList\|Physics:Quantum basics/See also/Matter}}

	==References==		==References==

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 {{Short description|One of two foundational theories of quantum chemistry}}
 {{Redirect|VBT}}
 {{Electronic structure methods}}
 In [[HandWiki:Chemistry|chemistry]], '''valence bond (VB) theory''' is one of the two basic theories, along with [[Chemistry:Molecular orbital theory|molecular orbital (MO) theory]], that were developed to use the methods of [[Physics:Quantum mechanics|quantum mechanics]] to describe [[Chemistry:Chemical bond|chemical bond]]ing. It focuses on how the [[Physics:Atomic orbital|atomic orbital]]s of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, [[Chemistry:Molecular orbital theory|molecular orbital theory]] has orbitals that cover the whole molecule.<ref>{{cite book| last1 = Murrell| first1 = J. N.| last2 = Kettle| first2 = S. F. A.| last3 = Tedder| first3 = J. M.| title = The Chemical Bond| edition = 2nd| publisher = John Wiley & Sons| year = 1985| isbn = 0-471-90759-6| url-access = registration| url = https://archive.org/details/chemicalbond0000murr_e8r6}}</ref>
 ==History==

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imported>WikiHarold: WikiHarold moved page Physics:Valence bond theory to Physics:Quantum valence bond theory without leaving a redirect

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{{Short description|One of two foundational theories of quantum chemistry}}
{{Redirect|VBT}}
{{Electronic structure methods}}
In [[HandWiki:Chemistry|chemistry]], '''valence bond (VB) theory''' is one of the two basic theories, along with [[Chemistry:Molecular orbital theory|molecular orbital (MO) theory]], that were developed to use the methods of [[Physics:Quantum mechanics|quantum mechanics]] to describe [[Chemistry:Chemical bond|chemical bond]]ing. It focuses on how the [[Physics:Atomic orbital|atomic orbital]]s of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, [[Chemistry:Molecular orbital theory|molecular orbital theory]] has orbitals that cover the whole molecule.<ref>{{cite book| last1 = Murrell| first1 = J. N.| last2 = Kettle| first2 = S. F. A.| last3 = Tedder| first3 = J. M.| title = The Chemical Bond| edition = 2nd| publisher = John Wiley & Sons| year = 1985| isbn = 0-471-90759-6| url-access = registration| url = https://archive.org/details/chemicalbond0000murr_e8r6}}</ref>

==History==

In 1916, G. N. Lewis proposed that a chemical bond forms by the interaction of two shared bonding electrons, with the representation of molecules as [[Chemistry:Lewis structure|Lewis structure]]s. In 1916, Kossel put forth his theory of the ionic chemical bond ([[Octet rule|octet rule]]), also independently advanced in the same year by [[Biography:Gilbert N. Lewis|Gilbert N. Lewis]].<ref>[http://www.ucc.ie/academic/chem/dolchem/html/dict/000c1.html University College Cork], [http://www.origin-life.gr.jp/2904/2904174/2904174.html University City Tübingen] {{Webarchive|url=https://web.archive.org/web/20161128185116/http://www.origin-life.gr.jp/2904/2904174/2904174.html |date=2016-11-28 }}, and (Pauling, 1960, p. 5).</ref><ref>Walther Kossel, “Uber Molkulbildung als Frage der Atombau”, Ann. Phys., 1916, 49:229–362.</ref> [[Biography:Walther Kossel|Walther Kossel]] put forward a theory similar to that of Lewis theory, except that Kossel supposed complete transfers of electrons between atoms, a model of ionic bonding. Both Lewis and Kossel based their bonding models on [[Chemistry:Abegg's rule|Abegg's rule]] (1904) that the difference between the maximum positive and negative valences of an element is frequently eight.

In 1921 the chemist Charles Rugeley Bury suggested that eight and eighteen electrons in a shell form stable configurations. Bury proposed that the electron configurations in transitional elements depended upon the valence electrons in their outer shell.<ref>{{Cite journal|last=Bury|first=Charles R.|date=July 1921|title=Langmuir's Theory of the Arrangement of Electrons in Atoms and Molecules|journal=[[Organization:Journal of the American Chemical Society|Journal of the American Chemical Society]]|language=en|volume=43|issue=7|pages=1602–1609|doi=10.1021/ja01440a023|bibcode=1921JAChS..43.1602B |issn=0002-7863|url=https://zenodo.org/record/1428812 }}</ref>

Although there is no mathematical formula either in chemistry or quantum mechanics for the arrangement of electrons in the atom, the hydrogen atom can be described by the [[Schrödinger equation]] and the Matrix Mechanics equation both derived in 1925. However, for hydrogen alone, in 1927 the Heitler–London theory was formulated which for the first time enabled the calculation of bonding properties of the [[Software:Hydrogen|hydrogen]] molecule H2 based on quantum mechanical considerations. Specifically, [[Biography:Walter Heitler|Walter Heitler]] determined how to use [[Schrödinger equation|Schrödinger's wave equation]] (1926) to show how two hydrogen atom wavefunctions join together, with plus, minus, and exchange terms, to form a [[Chemistry:Covalent bond|covalent bond]]. He then called up his associate [[Biography:Fritz London|Fritz London]] and they worked out the details of the theory over the course of the night.<ref>[https://scarc.library.oregonstate.edu/coll/pauling/bond/people/heitler.html Walter Heitler] – Key participants in the development of Linus Pauling's ''The Nature of the Chemical Bond''.</ref> Later, [[Biography:Linus Pauling|Linus Pauling]] used the pair bonding ideas of Lewis together with Heitler–London theory to develop two other key concepts in VB theory: [[Resonance|resonance]] (1928) and orbital hybridization (1930). According to [[Biography:Charles Coulson|Charles Coulson]], author of the noted 1952 book ''Valence'', this period marks the start of "modern valence bond theory", as contrasted with older valence bond theories, which are essentially electronic theories of [[Chemistry:Valence|valence]] couched in pre-wave-mechanical terms.<ref>{{Cite book |last=Coulson |first=Charles Alfred |url=https://books.google.com/books?id=Hv3BAAAAIAAJ |title=Valence |date=1952 |publisher=Clarendon Press |language=en}}</ref>

Linus Pauling published in 1931 his landmark paper on valence bond theory: "On the Nature of the Chemical Bond". Building on this article, Pauling's 1939 textbook: ''On the Nature of the Chemical Bond'' would become what some have called the bible of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to adequately address the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960s and 1970s as molecular orbital theory grew in usefulness as it was implemented in large digital computer programs. Since the 1980s, the more difficult problems, of implementing valence bond theory into computer programs, have been solved largely, and valence bond theory has seen a resurgence.<ref>{{Cite journal |last1=Shaik |first1=Sason |last2=Danovich |first2=David |last3=Hiberty |first3=Philippe C. |date=2021-03-15 |title=Valence Bond Theory—Its Birth, Struggles with Molecular Orbital Theory, Its Present State and Future Prospects |journal=Molecules |language=en |volume=26 |issue=6 |pages=1624 |doi=10.3390/molecules26061624 |doi-access=free |pmid=33804038 |pmc=8001733 |issn=1420-3049}}</ref>

==Theory==
According to this theory a '''covalent''' '''bond''' is formed between two atoms by the overlap of ''half filled valence'' atomic orbitals of each atom containing one unpaired electron. Valence Bond theory describes chemical bonding better than Lewis Theory, which states that atoms share or transfer electrons so that they achieve the octet rule. It does not take into account orbital interactions or bond angles, and treats all covalent bonds equally.<ref>{{Cite book |last1=Chang |first1=Raymond |title=General chemistry: the essential concepts |last2=Overby |first2=Jason Scott |date=2011 |publisher=McGraw-Hill |isbn=978-0-07-337563-2 |edition=6 |location=New York, NY}}</ref> A valence bond structure resembles a [[Chemistry:Lewis structure|Lewis structure]], but when a molecule cannot be fully represented by a single Lewis structure, multiple valence bond structures are used. Each of these VB structures represents a specific Lewis structure. This combination of valence bond structures is the main point of [[Resonance|resonance]] theory. Valence bond theory considers that the overlapping atomic orbitals of the participating atoms form a [[Chemistry:Chemical bond|chemical bond]]. Because of the overlapping, it is most [[Probability|probable]] that electrons should be in the bond region. Valence bond theory views bonds as weakly coupled orbitals (small overlap). Valence bond theory is typically easier to employ in [[Physics:Ground state|ground state]] molecules. The [[Physics:Core electron|core orbitals and electrons]] remain essentially unchanged during the formation of bonds.

[[File:Sigma bond.svg|thumb|left|σ bond between two atoms: localization of [[electron density]]]]
[[Image:Pi-Bond.svg|thumb|Two p-orbitals forming a {{pi}}-bond.]]
The overlapping atomic orbitals can differ. The two types of overlapping orbitals are sigma and pi. [[Chemistry:Sigma bond|Sigma bond]]s occur when the orbitals of two shared electrons overlap head-to-head, with the electron density most concentrated between nuclei. [[Chemistry:Pi bond|Pi bond]]s occur when two orbitals overlap when they are parallel.<ref>{{Cite book |last=House |first=J. E. |title=Inorganic chemistry |date=2013 |publisher=Elsevier/Academic Press |isbn=978-0-12-385110-9 |edition=2nd |location=Waltham, MA}}</ref> For example, a bond between two ''s''-orbital electrons is a sigma bond, because two spheres are always coaxial. In terms of bond order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond, and triple bonds contain one sigma bond and two pi bonds. However, the atomic orbitals for bonding may be hybrids. Hybridization is a model that describes how atomic orbitals combine to form new orbitals that better match the geometry of molecules. Atomic orbitals that are similar in energy combine to make hybrid orbitals. For example, the carbon in methane (CH4) undergoes sp3 hybridization to form four equivalent orbitals, resulting in a tetrahedral shape. Different types of hybridization, such as sp, sp2, and sp3, correspond to specific molecular geometries (linear, trigonal planar, and tetrahedral), influencing the bond angles observed in molecules. Hybrid orbitals provide additional directionality to sigma bonds, accurately explaining molecular geometries.<ref>{{Cite journal |last=Barnes |first=Craig |date=2003 |title=Inorganic Chemistry (Housecroft, Catherine E.; Sharpe, Alan G.) |url=http://dx.doi.org/10.1021/ed080p747 |journal=Journal of Chemical Education |volume=80 |issue=7 |pages=747 |doi=10.1021/ed080p747 |issn=0021-9584|url-access=subscription }}</ref>

==Comparison with MO theory==
Valence bond theory complements [[Chemistry:Molecular orbital theory|molecular orbital theory]] (MO), which does not adhere to the valence bond idea that electron pairs are localized between two specific atoms in a molecule, but that they are distributed in sets of [[Chemistry:Molecular orbital|molecular orbital]]s which can extend over the entire molecule. Although both theories describe chemical bonding, MO generally offer a clearer and more reliable framework for predicting magnetic and ionization properties (and therefore optical and IR spectra). In particular, molecular orbitals can effectively account for paramagnetism arising from unpaired electrons, whereas VBT struggles.<ref name=":12" />

Simple VB theory includes only covalent structures (for neutral molecules), but can be refined by adding terms for ionic structures to the wave function. Similarly simple MO theory considers a single electron configuration, with no correlation between the movement of different electrons, but correlation can be included by adding other configurations.

If many terms are considered in the wave functions, the two theories approach mathematical equivalence, However MO is a more popular approach than VB due to its easier implementation in the early days of computational chemistry.

Valence bond theory views [[Chemistry:Aromaticity|aromatic]] properties of molecules as due to spin coupling of the [[Chemistry:Pi bond|{{pi}} orbitals]].<ref>{{cite journal|last1=Cooper|first1=David L.|last2=Gerratt|first2=Joseph|last3=Raimondi|first3=Mario|title=The electronic structure of the benzene molecule|journal=Nature|volume=323|page=699|year=1986|doi=10.1038/323699a0|issue=6090|bibcode = 1986Natur.323..699C |s2cid=24349360}}</ref><ref>{{cite journal|last1=Pauling|first1=Linus|title=Electronic structure of the benzene molecule|journal=Nature|volume=325|page=396|year=1987|doi=10.1038/325396d0|issue=6103|bibcode = 1987Natur.325..396P |s2cid=4261220|doi-access=free}}</ref><ref>{{cite journal|last1=Messmer|first1=Richard P.|last2=Schultz|first2=Peter A.|title=The electronic structure of the benzene molecule|journal=Nature|volume=329|page=492|year=1987|doi=10.1038/329492a0|issue=6139|bibcode = 1987Natur.329..492M |s2cid=45218186|doi-access=free}}</ref><ref>{{cite journal|last1=Harcourt|first1=Richard D.|title=The electronic structure of the benzene molecule|journal=Nature|volume=329|page=491|year=1987|doi=10.1038/329491b0|issue=6139|bibcode = 1987Natur.329..491H |s2cid=4268597}}</ref> This is essentially still the old idea of resonance between Friedrich August Kekulé von Stradonitz and James Dewar structures. In contrast, molecular orbital theory views aromaticity as delocalization of the {{pi}}-electrons.<ref>{{Cite journal |last=Vemulapalli |first=G. K. |date=2008-10-01 |title=Theories of the chemical bond and its true nature |url=http://link.springer.com/10.1007/s10698-008-9049-2 |journal=Foundations of Chemistry |language=en |volume=10 |issue=3 |pages=167–176 |doi=10.1007/s10698-008-9049-2 |issn=1386-4238|url-access=subscription }}</ref> Valence bond treatments are restricted to relatively small molecules, largely due to the lack of orthogonality between valence bond orbitals and between valence bond structures. The molecular orbitals are always orthogonal.

Valence bond theory cannot explain electronic transitions and spectroscopic properties as effectively as MO theory. While VB employs hybridization to explain bonding, it can oversimplify complex bonding situations, limiting its applicability in more intricate molecular geometries such as transition metal compounds.<ref name=":12" />

On the other hand, VB theory provides a much more intuitive picture of the reorganization of electronic charge that takes place when bonds are broken and formed during the course of a chemical reaction.

Valence bond theory also correctly predicts the dissociation of homonuclear diatomic molecules into separate atoms even in the simplest models, while similarly crude MO approaches predict dissociation into a mixture of atoms and ions. For example, the MO function for dihydrogen is an equal mixture of the covalent and ionic valence bond structures and so predicts incorrectly that the molecule would dissociate into an equal mixture of hydrogen atoms and hydrogen positive and negative ions.

==Computational approaches==
{{Main|Chemistry:Modern valence bond theory}}
Modern valence bond theory replaces the overlapping atomic orbitals by overlapping valence bond orbitals that are expanded over a large number of [[Chemistry:Basis set|basis functions]], either centered each on one atom to give a classical valence bond picture, or centered on all atoms in the molecule. The resulting energies are more competitive with energies from calculations where electron correlation is introduced based on a Hartree–Fock reference wavefunction. The most recent text is by Shaik and Hiberty.<ref>{{cite book| last = Shaik| first = Sason S.|author2=Phillipe C. Hiberty| title = A Chemist's Guide to Valence Bond Theory| publisher = Wiley-Interscience
| year = 2008| location = New Jersey| isbn = 978-0-470-03735-5}}</ref>

==Applications==
An important aspect of the valence bond theory is the condition of maximum overlap, which leads to the formation of the strongest possible bonds. This theory is used to explain the covalent bond formation in many molecules.
[[File:Ch4_hybridization.svg|thumb|185x185px|sp3 hybridization in methane forms four equivalent sigma bonds with tetrahedral geometry.]]
For example, in the case of the [[Chemistry:Fluorine|F2]] molecule, the F−F bond is formed by the overlap of ''p''''z'' orbitals of the two F atoms, each containing an unpaired electron. Since the nature of the overlapping orbitals are different in H2 and F2 molecules, the bond strength and bond lengths differ between H2 and F2 molecules.

In [[Chemistry:Methane|methane]] (CH4), the carbon atom undergoes sp3 hybridization, allowing it to form four equivalent sigma bonds with hydrogen atoms, resulting in a tetrahedral geometry. Hybridization also explains the equal C-H bond strengths.<ref name=":12">{{Cite book |last1=Pettrucci |first1=Ralph H. |title=General Chemistry: Principles and Modern Applications |last2=Herring |first2=F. Geoffrey |last3=Madura |first3=Jeffrey D. |last4=Bissonnette |first4=Carey |year=2017 |isbn=978-0-13-293128-1 |edition=11th |pages=|publisher=Pearson }}</ref>

In an [[Chemistry:Hydrofluoric acid|HF]] molecule the covalent bond is formed by the overlap of the 1''s'' orbital of H and the 2''p''''z'' orbital of F, each containing an unpaired electron. Mutual sharing of electrons between H and F results in a covalent bond in HF.

== See also ==
* [[Chemistry:Valence bond programs|Valence bond programs]]

==References==
{{Reflist}}

{{Chemical bonding theory|state=expanded}}
{{Linus Pauling}}

[[Category:Quantum chemistry]]
[[Category:Chemical bonding]]
[[Category:General chemistry]]

{{Sourceattribution|Valence bond theory}}