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In Chemistry, a covalent bond is a Chemical bond in which each dependant Atome S shares one or more doublets of electron S of its external layers. It is what produces mutual attraction between atoms.

The covalent bond generally implies, the equal share of only one pair of electrons, called flexible doublet; each atom providing an electron; the pair of electrons being delocalized between the two atoms. The division of two or three pairs of electrons is called connection doubles or connection triples .
The triple connections are relatively rare in nature; one finds it for example in the Carbon monoxide CO. Sometimes, an electron is delocalized on several atoms, like in the case of the aromatic cycles and other resounding structures, such as the Benzène. When the electrons are delocalized on many atoms, there is a metal Liaison.

The covalent bond most frequently occurs between similar atoms of electronegativities. The difference in energy level between the two atoms is not sufficient to produce “the flight” of electrons of an atom towards the other. However the distribution of the electrons in a covalent bond between different atoms, will not be exactly equal; indeed the point of gravity will be moved towards the most electronegative atom (electronegativity means the capacity which the pair of electrons has an atom to attract towards him); it is said that the covalent bond is polarized. The direction of polarization is given by partial loads (delta + for the least electronegative atom and delta - for most electronegative). The larger the difference in electronegativity is between the atoms, the more the partial loads are high; it is said that the connection is polarized and has a character " ionique". The covalent bonds are more common between Non-métaux, while the ionic Connection is more frequent when one or each of the two atoms is metals.

The covalent bond tends to being stronger than other types of connection, such as the ionic connection. Moreover, contrary to the ionic connections where the atoms are bound by Coulomb attraction not-directional, the covalent bonds are strongly directional. Consequently, the molecules bound by covalence tend to adopt forms characteristic having of the specific angles of connection.

Function

The atoms seek stability and the covalent bonds cure it. The electrons brought, make it possible to saturate the last layer with electrons and thus, stabilize the atom.

History of the covalent bond

The idea of the covalent bond goes back to Lewis, which in 1916 describes the division of pairs of electrons between atoms. It presented the “notation of Lewis” in whom the electrons of valence (of the external electron shell) are represented like points around the atomic symbols. The pairs of electrons located between the atoms represent the covalent bonds and the multiple pairs represent double linkings or triple.

While the idea of the divided pairs of electrons provides an effective qualitative image of the covalent bond, the quantum Mécanique is necessary to include/understand the nature of these connections and to be able to envisage the structures and the properties of the simple molecules. In 1927, Walter Heitler and Fritz London were credited with the first correct quantum explanation of the chemical bond, specifically that of the molecule of dihydrogene. Their work was based on the model of connection of valence, which supposes that a chemical bond is formed when there is a good covering between orbital atomic. These orbital is known to have specific angles the ones with the others. The model of bond of valence could thus envisage successfully the angles of connection observed in simple molecules.

Today the model of the connections of valence was mainly supplanted by the model of the molecular orbital . In this model, when the atoms approach, their orbital atomic interacts and forms then a whole of molecular orbits which extends on all the molecule. Half of these orbits tends to being the orbital flexible ones and other half anti-flexible. The electrons of orbital flexible cause the formation of a chemical bond, whereas those of orbital anti-flexible tend to prevent them. The formation of a chemical bond is not possible that when the electrons occupying of orbital flexible are more numerous than those occupying of orbital anti-flexible.

The order of connection

The difference between the number of pairs of electrons contained in orbital flexible and anti-flexible ones determines the order of the connection. For example, in a diatomic molecule, there is a simple connection which is formed if there is an excess of two electrons in the orbital flexible ones (H₂), a double linking if four electrons are in excess (O₂) and a triple connection if this excess is of six electrons (N₂).

The order of connection does not need to be a Integer and the bonds can be delocalized on more than two atoms. For example, in the Benzene, the order of bond between two adjacent carbon atoms is of 3/2. I.e. the electrons of orbital flexible are distributed uniformly on each of the 6 carbon atoms; this phenomenon is called resonance . Other example are the anions derived from dioxygene: the addition of electrons in orbital π* antiliantes makes pass the order of connection from 2 to 3/2 for the Superoxyde O₂^- and to 1 for the Peroxyde O₂^2- The lengths and energies of dissociation of the connections are inversely proportional to the order of connection; the higher the order of connection is, plus this connection is short and strong.

Measurements

By using quantum mechanics it is possible to calculate the electronic structure, the energy levels, the distances and the angles of connections, the dipole moments, and the spectrum of simple molecules with a great exactitude. Currently, the distances and the angles from connections can be calculated as precisely as they can be measured (some pm for the distances and some degree S for the angles).

For the small molecules, calculations concerning the energy levels are sufficiently precise, so that it can be used in order to determine to them Heat of formation and the energy of their kinetic barrier of activation.

See too

Simple: Covalent jump

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