In chemistry, adouble bond is acovalent bond between twoatoms involving fourbonding electrons as opposed to two in asingle bond. Double bonds occur most commonly between two carbon atoms, for example inalkenes. Many double bonds exist between two different elements: for example, in acarbonyl group between a carbon atom and an oxygen atom. Other common double bonds are found inazo compounds (N=N),imines (C=N), andsulfoxides (S=O). In askeletal formula, a double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, theequals sign is used for this.[1][2] Double bonds were introduced in chemical notation byRussian chemistAlexander Butlerov.[citation needed]
Double bonds involving carbon are stronger and shorter thansingle bonds. Thebond order is two. Double bonds are also electron-rich, which makes them potentially more reactive in the presence of a strong electron acceptor (as inaddition reactions of the halogens).
The type of bonding can be explained in terms oforbital hybridisation. Inethylene each carbon atom has threesp2 orbitals and onep-orbital. The three sp2 orbitals lie in a plane with ~120° angles. The p-orbital is perpendicular to this plane. When the carbon atoms approach each other, two of the sp2 orbitals overlap to form asigma bond. At the same time, the two p-orbitals approach (again in the same plane) and together they form api bond. For maximum overlap, the p-orbitals have to remain parallel, and, therefore, rotation around the central bond is not possible. This property gives rise tocis-trans isomerism. Double bonds are shorter than single bonds because p-orbital overlap is maximized.
With 133 pm, the ethyleneC=Cbond length is shorter than the C−C length inethane with 154 pm. The double bond is also stronger, 636 kJ mol−1 versus 368 kJ mol−1 but not twice as much as the pi-bond is weaker than the sigma bond due to less effective pi-overlap.
In an alternative representation, the double bond results from two overlapping sp3 orbitals as in abent bond.[3]
In molecules with alternating double bonds and single bonds, p-orbital overlap can exist over multiple atoms in a chain, giving rise to aconjugated system. Conjugation can be found in systems such asdienes andenones. Incyclic molecules, conjugation can lead toaromaticity. Incumulenes, two double bonds are adjacent and do not overlap with each other.
Double bonds are common forperiod 2 elementscarbon,nitrogen, andoxygen, and less common with elements of higher periods, a distinction called thedouble bond rule. Metals, too, can engage in multiple bonding in ametal–ligand multiple bond.
Double bonded compounds,alkene homologues, R2E=ER2 are now known for all of the heaviergroup 14 elements. Unlike the alkenes these compounds are not planar but adopt twisted and/ortrans bent structures. These effects become more pronounced for the heavier elements. Thedistannene (Me3Si)2CHSn=SnCH(SiMe3)2 has a tin-tin bond length just a little shorter than a single bond, a trans bent structure with pyramidal coordination at eachtin atom, and readily dissociates in solution to form (Me3Si)2CHSn: (stannanediyl, a carbene analog). The bonding comprises two weak donor acceptor bonds, the lone pair on each tin atom overlapping with the empty p orbital on the other.[4][5] In contrast, indisilenes each silicon atom has planar coordination but the substituents are twisted so that the molecule as a whole is not planar. In diplumbenes the Pb=Pb bond length can be longer than that of many corresponding single bonds[5] Plumbenes and stannenes generally dissociate in solution into monomers with bond enthalpies that are just a fraction of the corresponding single bonds. Some double bonds plumbenes and stannenes are similar in strength to hydrogen bonds.[4] TheCarter–Goddard–Malrieu–Trinquier model can be used to predict the nature of the bonding.[4]
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