Double bond

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A double bond in chemistry is a chemical bond between two chemical elements involving four bonding electrons instead of the usual two. The most common double bond, that is between two carbon atoms, can be found in alkenes. Many types of double bonds exist between two different elements. For example, in a carbonyl group with 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 skeletal formula the double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, the equals sign is used for this.[1][2] Double bonds were first introduced in chemical notation by prominent Russian chemist Alexander Butlerov.

Double bonds involving carbon are stronger than single bonds and are also shorter. The bond order is two. Double bonds are also electron-rich, which makes them reactive.

Chemical compounds with double bonds
Ethene structural.svg
Leuckart-Wallach-Reaktion Aceton.svg
Dimethylsulfoxid.svg
Trans-diazene-2D.png
Ethylene
Carbon-carbon
double bond
Acetone
Carbon-oxygen
double bond
Dimethyl sulfoxide
Sulfur-oxygen
double bond
Diazene
Nitrogen-nitrogen
double bond

Double bonds in alkenes[edit]

Geometry of ethylene

The type of bonding can be explained in terms of orbital hybridization. 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.

Double bond presentation Double bond presentation
2 sp2 orbitals (total of 3 such orbitals) approach to form a sp2-sp2 sigma bond Two p-orbitals overlap to form a pi-bond in a plane parallel to the sigma plane

With 133 pm, the C=C bond length is shorter than the C−C length in ethane with 154 pm. The double bond is also stronger, 636 (KJ/mol) versus 368 kJ/mole 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 a bent bond.[3]

Types of double bonds between atoms[edit]

C O N S
C alkene carbonyl group imine thioketone, thial
O dioxygen nitroso compound sulfoxide, sulfone, sulfinic acid, sulfonic acid
N azo compound
S disulfur

Variations[edit]

In molecules, with alternating double bonds and single bonds, p-orbital overlap can exist over multiple atoms in a chain, giving rise to a conjugated system. Conjugation can be found in systems such as dienes and enones. In cyclic molecules, conjugation can lead to aromaticity. In cumulenes, two double bonds are adjacent.

Double bonds are common for period 2 elements carbon, nitrogen, and oxygen, and less common with elements of higher periods. Metals, too, can engage in multiple bonding in a metal ligand multiple bond.

Group 14 alkene homologs[edit]

Double bonded compounds, alkene homologs, R2E=ER2 are now known for all of the heavier group 14 elements. Unlike the alkenes these compounds are not planar but adopt twisted and/or trans bent structures. These effects become more pronounced for the heavier elements. The distannene (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 each tin atom, and readily dissociates in solution to form (Me3Si)2CHSn: (stannanediyl, a carbene analog). The bonding comprises of 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, in disilenes 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] The Carter-Goddard-Malrieu-Trinquier model can be used to predict the nature of the bonding.[4]

References[edit]

  1. ^ March, Jerry (1985), Advanced Organic Chemistry: Reactions, Mechanisms, and Structure (3rd ed.), New York: Wiley, ISBN 0-471-85472-7 .
  2. ^ Organic Chemistry 2nd Ed. John McMurry.
  3. ^ Advanced Organic Chemistry Carey, Francis A., Sundberg, Richard J. 5th ed. 2007.
  4. ^ a b c Power, Philip P. (1999). "π-Bonding and the Lone Pair Effect in Multiple Bonds between Heavier Main Group Elements". Chemical Reviews 99 (12): 3463–3504. doi:10.1021/cr9408989. 
  5. ^ a b Wang, Yuzhong; Robinson, Gregory H. (2009). "Unique homonuclear multiple bonding in main group compounds". Chemical Communications (Royal Society of Chemistry) (35): 5201–5213. doi:10.1039/B908048A. Retrieved 21 January 2015. (subscription required (help)). 
  • Pyykkö, Pekka; Riedel, Sebastian; Patzschke, Michael (2005). "Triple-Bond Covalent Radii". Chemistry - A European Journal 11 (12): 3511–20. doi:10.1002/chem.200401299. PMID 15832398.