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WikiHarold moved page Chemistry:Bond order to Physics:Quantum bond order without leaving a redirect

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WikiHarold moved page Chemistry:Bond order to Physics:Quantum bond order without leaving a redirect

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{{Short description|Difference between the number of bonds and anti-bonds in a molecule}}
In [[Chemistry:Chemistry|chemistry]], '''bond order''' is a formal measure of the multiplicity of a [[Chemistry:Covalent bond|covalent bond]] between two atoms. As introduced by [[Biography:Linus Pauling|Linus Pauling]], bond order is defined as the difference between the numbers of electron pairs in [[Chemistry:Bonding molecular orbital|bonding]] and [[Chemistry:Antibonding molecular orbital|antibonding molecular orbital]]s.

Bond order gives a rough indication of the [[Physics:Bond-dissociation energy|stability of a bond]]. [[Chemistry:Isoelectronicity|Isoelectronic]] species have the same bond order.<ref>{{Cite book|title=Modern's abc Chemistry|last=Dr. S.P. Jauhar}}</ref>

==Examples==

The bond order itself is the number of [[Physics:Electron pair|electron pair]]s (covalent bonds) between two [[Physics:Atom|atom]]s.<ref>{{GoldBookRef|title=Bond number|file=B00705}}</ref> For example, in diatomic [[Chemistry:Nitrogen|nitrogen]] N≡N, the bond order between the two nitrogen atoms is 3 ([[Chemistry:Triple bond|triple bond]]). In [[Chemistry:Acetylene|acetylene]] H–C≡C–H, the bond order between the two [[Chemistry:Carbon|carbon]] atoms is also 3, and the C–H bond order is 1 ([[Chemistry:Single bond|single bond]]). In [[Chemistry:Carbon monoxide|carbon monoxide]], {{chem2|-C\tO+}}, the bond order between carbon and oxygen is 3. In [[Chemistry:Thiazyl trifluoride|thiazyl trifluoride]] {{chem2|N\tSF3}}, the bond order between [[Chemistry:Sulfur|sulfur]] and nitrogen is 3, and between sulfur and [[Chemistry:Fluorine|fluorine]] is 1. In diatomic [[Chemistry:Oxygen|oxygen]] O=O the bond order is 2 ([[Chemistry:Double bond|double bond]]). In [[Chemistry:Ethylene|ethylene]] {{chem2|H2C\dCH2}} the bond order between the two carbon atoms is also 2. The bond order between carbon and oxygen in [[Chemistry:Carbon dioxide|carbon dioxide]] O=C=O is also 2. In [[Chemistry:Phosgene|phosgene]] {{chem2|O\dCCl2}}, the bond order between carbon and oxygen is 2, and between carbon and [[Chemistry:Chlorine|chlorine]] is 1.

In some molecules, bond orders can be 4 ([[Chemistry:Quadruple bond|quadruple bond]]), 5 ([[Chemistry:Quintuple bond|quintuple bond]]) or even 6 ([[Chemistry:Sextuple bond|sextuple bond]]). For example, [[Chemistry:Potassium octachlorodimolybdate|potassium octachlorodimolybdate]] salt ({{chem2|K4[Mo2Cl8]}}) contains the {{chem2|[Cl4Mo\qMoCl4](4–)}} anion, in which the two [[Chemistry:Molybdenum|Mo]] atoms are linked to each other by a bond with order of 4. Each Mo atom is linked to four [[Chemistry:Chloride|{{chem2|Cl−}}]] ligands by a bond with order of 1. The compound ([[Chemistry:Terphenyl|terphenyl]])–[[Chemistry:Chromium|CrCr]]–(terphenyl) contains two [[Chemistry:Chromium|chromium]] atoms linked to each other by a bond with order of 5, and each chromium atom is linked to one terphenyl ligand by a single bond. A bond of order 6 is detected in ditungsten molecules [[Chemistry:Tungsten|{{chem2|W2}}]], which exist only in a gaseous phase.

===Non-integer bond orders===

In molecules which have [[Chemistry:Resonance|resonance]] or nonclassical bonding, bond order may not be an [[Integer|integer]]. In [[Chemistry:Benzene|benzene]], the delocalized [[Chemistry:Molecular orbital|molecular orbital]]s contain 6 pi electrons over six carbons, essentially yielding half a [[Chemistry:Pi bond|pi bond]] together with the [[Chemistry:Sigma bond|sigma bond]] for each pair of carbon atoms, giving a calculated bond order of 1.5 (one and a half bond). Furthermore, bond orders of 1.1 (eleven tenths bond), 4/3 (or 1.333333..., four thirds bond) or 0.5 (half bond), for example, can occur in some molecules and essentially refer to bond strength relative to bonds with order 1. In the [[Chemistry:Nitrate|nitrate]] anion ({{chem2|NO3−}}), the bond order for each bond between nitrogen and oxygen is 4/3 (or 1.333333...). Bonding in [[Chemistry:Dihydrogen cation|dihydrogen cation]] {{chem2|H2+}} can be described as a covalent one-electron bond, thus the bonding between the two hydrogen atoms has bond order of 0.5.<ref>{{Cite book|author=Clark R. Landis|author2=Frank Weinhold |title=Valency and bonding: a natural bond orbital donor-acceptor perspective |publisher=Cambridge University Press |location=Cambridge, UK |year=2005 |pages=91–92 |isbn=978-0-521-83128-4 }}</ref>

==Bond order in molecular orbital theory==

In [[Chemistry:Molecular orbital theory|molecular orbital theory]], bond order is defined as half the difference between the number of bonding electrons and the number of antibonding electrons as per the equation below.<ref>{{cite book | author1 = Jonathan Clayden | last2 = Greeves | first2 = Nick | author3 = Stuart Warren | title = Organic Chemistry | edition = 2nd | publisher = Oxford University Press | date = 2012 | isbn = 978-0-19-927029-3 | page = 91}}</ref><ref>{{cite book | title = Inorganic Chemistry | last1 = Housecroft | first1 = C. E. | last2 = Sharpe | first2 = A. G. | year = 2012 | publisher = Prentice Hall | edition = 4th | isbn = 978-0-273-74275-3 | pages = 35–37}}</ref> This often but not always yields similar results for bonds near their equilibrium lengths, but it does not work for stretched bonds.<ref name = Manz2017>{{cite journal | doi = 10.1039/c7ra07400j | journal = RSC Adv. | title = Introducing DDEC6 atomic population analysis: part 3. Comprehensive method to compute bond orders | author1 = T. A. Manz | year = 2017 | volume = 7 | issue = 72 | pages = 45552–45581| bibcode = 2017RSCAd...745552M |doi-access = free}}</ref> Bond order is also an index of bond strength and is also used extensively in [[Chemistry:Valence bond theory|valence bond theory]].

:''bond order'' = {{sfrac|''number of bonding electrons'' - ''number of antibonding electrons''|2}}

Generally, the higher the bond order, the stronger the bond. Bond orders of one-half may be stable, as shown by the stability of {{chem2|H2+}} (bond length 106 pm, bond energy 269 kJ/mol) and {{chem2|He2+}} (bond length 108 pm, bond energy 251 kJ/mol).<ref>Bruce Averill and Patricia Eldredge, ''Chemistry: Principles, Patterns, and Applications'' (Pearson/Prentice Hall, 2007), 409.</ref>

Hückel molecular orbital theory offers another approach for defining bond orders based on molecular orbital coefficients, for planar molecules with delocalized π bonding. The theory divides bonding into a sigma framework and a pi system. The π-bond order between atoms ''r'' and ''s'' derived from Hückel theory was defined by [[Biography:Charles Coulson|Charles Coulson]] by using the orbital coefficients of the Hückel MOs:<ref>{{cite book |last1=Levine |first1=Ira N. |title=Quantum Chemistry |date=1991 |publisher=Prentice-Hall |isbn=0-205-12770-3 |page=567 |edition=4th}}</ref><ref>{{cite journal |last1=Coulson |first1=Charles Alfred |title=The electronic structure of some polyenes and aromatic molecules. VII. Bonds of fractional order by the molecular orbital method |journal=Proceedings of the Royal Society A |date=7 February 1939 |volume=169 |issue=938 |pages=413–428 |doi=10.1098/rspa.1939.0006 |bibcode=1939RSPSA.169..413C |doi-access=free }}</ref>{{clarification needed|Should we read the writer's mind and guess what "MOs" stands for??? "Molecular Orbitals"??? Using jargon and abbreviations without clear explanation what that abbreviations stand for is annoying for readers who are not experts in this field. Even experts must guess what the "MOs" stands for.|date=August 2022}}

:<math>p_{rs} = \sum_i n_ic_{ri}c_{si}</math>,

Here the sum extends over π molecular orbitals only, and ''ni'' is the number of electrons occupying orbital ''i'' with coefficients ''cri'' and ''csi'' on atoms ''r'' and ''s'' respectively. Assuming a bond order contribution of 1 from the sigma component this gives a total bond order (σ + π) of 5/3 = 1.67 for benzene, rather than the commonly cited bond order of 1.5, showing some degree of ambiguity in how the concept of bond order is defined.

For more elaborate forms of molecular orbital theory involving larger [[Chemistry:Basis set|basis sets]], still other definitions have been proposed.<ref>{{cite journal |last1=Sannigrahi |first1=A. B. |last2=Kar |first2=Tapas |title=Molecular orbital theory of bond order and valency |journal=Journal of Chemical Education |date=August 1988 |volume=65 |issue=8 |pages=674–676 |doi=10.1021/ed065p674 |bibcode=1988JChEd..65..674S |url=https://pubs.acs.org/doi/abs/10.1021/ed065p674 |access-date=5 December 2020}}</ref> A standard [[Physics:Quantum mechanical|quantum mechanical]] definition for bond order has been debated for a long time.<ref>IUPAC Gold Book [http://goldbook.iupac.org/BT07005.html ''bond order'']</ref> A comprehensive method to compute bond orders from quantum chemistry calculations was published in 2017.<ref name = Manz2017 />

==Other definitions==
The bond order concept is used in [[Physics:Molecular dynamics|molecular dynamics]] and [[Chemistry:Bond order potential|bond order potential]]s. The magnitude of the bond order is associated with the [[Bond length|bond length]]. According to Linus Pauling in 1947, the bond order between atoms ''i'' and ''j'' is experimentally described as

:<math>s_{ij} = \exp{\left[\frac{d_{1} - d_{ij}}{b}\right]}</math>

where ''d''1 is the single bond length, ''dij'' is the bond length experimentally measured, and ''b'' is a constant, depending on the atoms. Pauling suggested a value of 0.353 [[Angstrom|Å]] for ''b'', for carbon-carbon bonds in the original equation:<ref>{{cite journal | last = Pauling | first = Linus | title = Atomic Radii and Interatomic Distances in Metals | journal = Journal of the American Chemical Society | volume = 69 | issue =3 | pages = 542–553 | date = March 1, 1947 | doi=10.1021/ja01195a024}}</ref>

:<math>d_{1} - d_{ij} = 0.353~\text{ln}(s_{ij})</math>

The value of the constant ''b'' depends on the atoms. This definition of bond order is somewhat ''ad hoc'' and only easy to apply for diatomic molecules.

==References==
{{Reflist}}

{{Chemical bonding theory}}

[[Category:Chemical bond properties|Order]]

{{Sourceattribution|Bond order}}

Physics:Quantum bond order - Revision history

imported>WikiHarold: WikiHarold moved page Chemistry:Bond order to Physics:Quantum bond order without leaving a redirect

imported>WikiHarold: WikiHarold moved page Chemistry:Bond order to Physics:Quantum bond order without leaving a redirect