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Concept
Double bond
Summary
In chemistry, a double bond is a covalent bond between two atoms involving four bonding electrons as opposed to two in a single bond. Double bonds occur most commonly between two carbon atoms, for example in alkenes. Many double bonds exist between two different elements: for example, in a carbonyl group between a carbon atom and an oxygen atom. Other common double bonds are found in azo compounds (N=N), imines (C=N), and sulfoxides (S=O). In a skeletal formula, a double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, the equals sign is used for this. Double bonds were first introduced in chemical notation by Russian chemist Alexander Butlerov.
Double bonds involving carbon are stronger and shorter than single bonds. The bond order is two. Double bonds are also electron-rich, which makes them potentially more reactive in the presence of a strong electron acceptor (as in addition reactions of the halogens).
File:Ethene structural.svg|[[Ethylene]] {{nowrap|Carbon-carbon}} double bond
File:Leuckart-Wallach-Reaktion Aceton.svg|[[Acetone]] {{nowrap|Carbon-oxygen}} double bond
File:Dimethylsulfoxid.svg|[[Dimethyl sulfoxide]] {{nowrap|Sulfur-oxygen}} double bond
File:Trans-diazene-2D.svg|[[Diazene]] {{nowrap|Nitrogen-nitrogen}} double bond
The type of bonding can be explained in terms of orbital hybridisation. In ethylene each carbon atom has three sp2 orbitals and one p-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 a sigma bond. At the same time, the two p-orbitals approach (again in the same plane) and together they form a pi 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 to cis-trans isomerism. Double bonds are shorter than single bonds because p-orbital overlap is maximized.
Image:Doppelbindung1.
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Orbital hybridisation
In chemistry, orbital hybridisation (or hybridization) is the concept of mixing atomic orbitals to form new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. For example, in a carbon atom which forms four single bonds the valence-shell s orbital combines with three valence-shell p orbitals to form four equivalent sp3 mixtures in a tetrahedral arrangement around the carbon to bond to four different atoms.
Allenes
In organic chemistry, allenes are organic compounds in which one carbon atom has double bonds with each of its two adjacent carbon atoms (, where R is H or some organyl group). Allenes are classified as cumulated dienes. The parent compound of this class is propadiene (), which is itself also called allene. An group of the structure is called allenyl, where R is H or some alkyl group. Compounds with an allene-type structure but with more than three carbon atoms are members of a larger class of compounds called cumulenes with bonding.
Pi bond
In chemistry, pi bonds (π bonds) are covalent chemical bonds, in each of which two lobes of an orbital on one atom overlap with two lobes of an orbital on another atom, and in which this overlap occurs laterally. Each of these atomic orbitals has an electron density of zero at a shared nodal plane that passes through the two bonded nuclei. This plane also is a nodal plane for the molecular orbital of the pi bond. Pi bonds can form in double and triple bonds but do not form in single bonds in most cases.