Inorganic chemistry, abent bond, also known as abanana bond, is a type ofcovalentchemical bond with a geometry somewhat reminiscent of abanana. The term itself is a general representation ofelectron density or configuration resembling a similar "bent" structure within small ring molecules, such ascyclopropane (C3H6) or as a representation of double or triple bonds within a compound that is an alternative to thesigma andpi bond model.
Bent bonds[1][2][3][4] are a special type ofchemical bonding in which the ordinaryhybridization state of two atoms making up a chemical bond are modified with increased or decreaseds-orbital character in order to accommodate a particularmolecular geometry. Bent bonds are found in strainedorganic compounds such ascyclopropane,oxirane andaziridine.
In these compounds, it is not possible for the carbon atoms to assume the 109.5°bond angles with standard sp3 hybridization. Increasing the p-character to sp5 (i.e.1⁄6 s-density and5⁄6 p-density)[5] makes it possible to reduce the bond angles to 60°. At the same time, the carbon-to-hydrogen bonds gain more s-character, which shortens them. In cyclopropane, the maximum electron density between two carbon atoms does not correspond to the internuclear axis, hence the namebent bond. In cyclopropane, theinterorbital angle is 104°. This bending can be observed experimentally byX-ray diffraction of certain cyclopropane derivatives: thedeformation density is outside the line of centers between the two carbon atoms. The carbon–carbonbond lengths are shorter than in a regular alkane bond: 151 pm versus 153 pm.[6]
Cyclobutane is a larger ring, but still has bent bonds. In this molecule, the carbon bond angles are 90° for the planar conformation and 88° for the puckered one. Unlike in cyclopropane, the C–C bond lengths actually increase rather than decrease; this is mainly due to 1,3-nonbondedsteric repulsion. In terms of reactivity, cyclobutane is relatively inert and behaves like ordinaryalkanes.
An alternative model utilizes semi-localizedWalsh orbitals in which cyclopropane is described as a carbon sp2 sigma bonding and in-plane pi bonding system. Critics of the Walsh orbital theory argue that this model does not represent theground state of cyclopropane as it cannot be transformed into the localized or fully delocalized descriptions via a unitary transformation.[3]
Two different explanations for the nature of double and triplecovalent bonds inorganic molecules were proposed in the 1930s.Linus Pauling proposed that the double bond results from two equivalent tetrahedral orbitals from each atom,[7] which later came to be calledbanana bonds ortau bonds.[8]Erich Hückel proposed a representation of the double bond as a combination of asigma bond plus api bond.[9][10][11] The Hückel representation is the better-known one, and it is the one found in most textbooks since the late-20th century.
Both models represent the same total electron density, with the orbitals related by aunitary transformation. We can construct the two equivalent bent bond orbitalsh andh' by taking linear combinationsh =c1σ +c2π andh' =c1σ –c2π for an appropriate choice of coefficientsc1 andc2. In a 1996 review,Kenneth B. Wiberg concluded that "although a conclusive statement cannot be made on the basis of the currently available information, it seems likely that we can continue to consider the σ/π and bent-bond descriptions of ethylene to be equivalent."[3]Ian Fleming goes further in a 2010 textbook, noting that "the overall distribution of electrons [...] is exactly the same" in the two models.[12]
The bent bond theory can also explain other phenomena in organic molecules. Influoromethane (CH3F), for instance, the experimental F–C–H bond angle is 109°, which is greater than the calculated value. This is because according toBent's rule, the C–F bond gains p-orbital character leading to high s-character in the C–H bonds, and H–C–H bond angles approaching those of sp2 orbitals – e.g. 120° – leaving less for the F–C–H bond angle. The difference is again explained in terms of bent bonds.[3]
Bent bonds also come into play in thegauche effect, explaining the preference for gauche conformations in certain substituted alkanes and thealkenecis effect associated with some unusually stable alkenecis isomers.[3]
