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-{{redirect|Covalent}}
-'''Covalent bonding''' is a form of [[chemical bond]]ing that is characterized by the ''sharing'' of pairs of [[electron]]s between [[atom]]s. In short, attraction-to-repulsion stability that forms between atoms when they share electrons is known as covalent bonding. The term "covalence", in regards to bonding, was first used in 1919 by [[Irving Langmuir]] in a ''Journal of American Chemical Society'' article entitled "[[Isomorphism]], [[Isosterism]], and Covalence" wherein he states:<ref>Langmuir, I. (1919). "[http://pubs.acs.org/wls/journals/query/subscriberResults.html?op=searchJournals Isomorphism, Isosterism and Covalence.]" ''J. Am. Chem. Soc''.; 1919; 41(10); 1543-1559.</ref>
-{{cquote|I propose that the number of electrons which any given atom shares with the adjacent atoms be called the '''covalence''' of that atom.}}
-Covalent bonding includes many kinds of interactions, including [[σ]]-bonding, [[π]]-bonding, metal-metal bonding, [[agostic complex|agostic interaction]]s, and [[three-center two-electron bond]]s.<ref>March, J. “Advanced Organic Chemistry” $th Ed.  J. Wiley and Sons, 1992: New York.  ISBN 0-471-60180-2.</ref><ref>G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.</ref>  The term ''covalent bond'' dates from 1939.<ref>[[Merriam-Webster]] - Collegiate Dictionary (2000).</ref>  The prefix ''co-'' means ''jointly, associated in action, partnered to a lesser degree, '' etc.; thus a "co-valent bond", essentially, means that the atoms share "valence", such as is discussed in [[valence bond theory]].  In the molecule H<sub>2</sub>, the hydrogen atoms share the two electrons via covalent bonding.  Covalency is greatest between atoms of similar [[electronegativity|electronegativities]].  Thus, covalent bonding does not necessarily require the two atoms be of the same elements, only that they be of comparable electronegativity.   Because covalent bonding entails sharing of electrons, is necessarily [[Delocalized electron|delocalized]].  Furthermore, in contrast to electrostatic interactions ("[[ionic bond]]s") the strength of covalent bond depends on the angular relation between atoms in polyatomic molecules.
-[[Image:CovalentBond.png|right|thumb|300px|Schemes depicting covalent (left) and polar covalent (right) bonding in a diatomic molecule.  The arrows represent electrons provided by the participating atoms.]]
-== History ==
-The idea of covalent bonding can be traced to [[Gilbert N. Lewis]], who in 1916 described the sharing of electron pairs between atoms. He introduced the so called ''[[Lewis Structure|Lewis notation]]'' or ''electron dot notation or The Lewis Dot Structure'' in which valence electrons (those in the outer shell) are represented as dots around the atomic symbols.  Pairs of electrons located between atoms represent covalent bonds.  Multiple pairs represent multiple bonds, such as double and triple bonds.  Some examples of Electron Dot Notation are shown in the following figure.  An alternative form, in which bond-forming electron pairs are represented as solid lines, is shown alongside.
-[[Image:covalent.svg|right|thumb|160px|Early concepts in covalent bonding arose from this kind of image of the molecule of [[methane]]. Covalent bonding is implied in the [[Lewis structure]] that indicates sharing of electrons between atoms.]]
-While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, [[quantum mechanics]] is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. [[Walter Heitler]] and [[Fritz London]] are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of [[molecular hydrogen]], in 1927.  Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the [[atomic orbitals]] of participating atoms.  These atomic orbitals are known to have specific angular relationships between each other, and thus the valence bond model can successfully predict the bond angles observed in simple molecules.
-== Bond polarity ==
-A "pure" covalent bond occurs between atoms with identical electronegativity, although some texts suggest that the term should be used when the difference is less than 0.2.  According to a widely-accepted definition, polar covalence describes bonds between atoms whose [[electronegativity|electronegativities]] differ by less than 2.1 but greater than 0.5.  Polar covalency also describes the so-called a [[coordinate covalent bond]], also known as a dative covalent bond, which occurs when one atom "gives" both of the electrons in the bond.  The classic example is [[borane-ammonia]].
-== Bond order ==
-[[Bond order]] is a number that indicates the number of pairs of electrons shared between atoms forming a covalent bond.  The term is only applicable to diatomic molecules, but is used to describe bonds within polyatomic compounds as well.
-# The most common type of covalent bond is the '''single bond''', the sharing of only one pair of electrons between two atoms. It usually consists of one [[sigma bond]]. All bonds with more than one shared pair are called '''multiple bonds'''.
-# Sharing two pairs is called a '''double bond'''. An example is in [[ethylene]] (between the carbon atoms). It usually consists of one [[sigma bond]] and one [[pi bond]].
-# Sharing three pairs is called a '''triple bond'''. An example is in [[hydrogen cyanide]] (between C and N). It usually consists of one sigma bond and two pi-bonds.
-# [[Quadruple bond]]s are found in the transition metals.  [[Molybdenum]] and [[rhenium]] are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in [[Di-tungsten tetra(hpp)]].
-# [[Quintuple bond]]s have been found to exist in certain di[[chromium]] compounds.
-# The only known molecules with true sextuple bonds (order 6) are diatomic [[molybdenum|Mo]]<sub>2</sub> and [[tungsten|W]]<sub>2</sub>, in the gaseous phase at very low temperatures. Although diatomic [[chromium|Cr]]<sub>2</sub> and [[uranium|U]]<sub>2</sub> have formal structures with twelve-electron bonds, their effective bond orders (derived from [[quantum chemistry]] calculations) are less than 5. There is strong evidence to believe that no two elements in the periodic table can form a bond with greater order than 6.<ref>{{cite journal | last = Roos | first = Björn O. | coauthors = Antonio C. Borin, and Laura Gagliardi | year = 2007 | month = January | title = Reaching the Maximum Multiplicity of the Covalent Chemical Bond | journal = Angewandte Chemie International Edition | doi = 10.1002/anie.200603600 }}</ref>
-Most bonding of course, is not localized, so the following classification, while powerful and pervasive, is of limited validity.  [[Three center bond]] do not conform readily to the above conventions.
-== Resonance ==
-Many bonding situations can be described with more than one valid Lewis Dot Structure (for example, [[ozone]], O<sub>3</sub>).  In an LDS diagram of O<sub>3</sub>, the center atom will have a single bond with one atom and a double bond with the other.  The LDS diagram cannot tell us which atom has the double bond; the first and second adjoining atoms have equal chances of having the double bond.  These two possible structures are called [[chemical resonance|resonance structures]].  In reality, the structure of ozone is a '''resonance hybrid''' between its two possible resonance structures.  Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.
-A special resonance case is exhibited in [[aromatic]] rings of atoms (for example, [[benzene]]).  Aromatic rings are composed of atoms arranged in a circle (held together by covalent bonds) that may alternate between single and double bonds according to their LDS.  In actuality, the electrons tend to be disambiguously and evenly spaced within the ring.  Electron sharing in aromatic structures is often represented with a ring inside the circle of atoms.
-== Current theory ==
-Today the valence bond model has been supplanted by the [[molecular orbital]] model.  In this model, as atoms are brought together, the ''atomic'' orbitals interact to form ''molecular'' orbitals, which are linear sums and differences of the atomic orbitals.  These molecular orbitals are a cross between the original atomic orbitals and generally extend between the two bonding atoms.
-Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and electromagnetic spectra of simple molecules with a high degree of accuracy.  Bond distances and angles can be calculated as accurately as they can be measured (distances to a few pm and bond angles to a few degrees).  For small molecules, calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.
-== References ==
-<references/>
-* http://www.chemguide.co.uk/atoms/bonding/covalent.html
-* http://www.chem.ox.ac.uk/vrchemistry/electronsandbonds/intro1.htm
-* http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/bond.html#c5
-== External links ==
-* [http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_07.html Covalent Bonds and Molecular Structure]
-* [http://www.chemistrycoach.com/tutorials-1.htm#Bonding2 Bonding tutorials]
-{{Organic chemistry}}
-==See also==
-* [[Chemical bond]]
-* [[Metallic bonding]]
-* [[Ionic bond]]
-* [[Linear combination of atomic orbitals]]
-* [[Orbital hybridisation|Hybridisation]]
-* [[Hydrogen bond]]
-* [[Noncovalent bonding]]
-* [[Disulfide bond]]
-[[Category:Chemical bonding]]
-[[su:Beungkeut kovalén]]
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