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Covalent bonding most frequently occurs between atoms with similar electronegativities, where neither atom can provide sufficient energy to completely remove an electron from the other atom. Covalent bonds are more common between non-metals, whereas ionic bonding is more common between two metal atoms or a metal and a non-metal atom.
Covalent bonding tends to be stronger than other types of bonding, such as ionic bonding. In addition unlike ionic bonding, where ions are held together by a non-directional coulombic attraction, covalent bonds are highly directional. As a result, covalently bonded molecules tend to form in a relatively small number of characteristic shapes, exhibiting specific bonding angles.
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 Notation or Electron Dot Notation[?] 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 alternate form in which bond-forming electron pairs are represented as solid lines is shown in blue.
While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, Quantum mechanics is needed to understand the the nature of these bonds and predict the structures and properties of simple molecules. Heitler and 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.
Today the valence bond model has largely been supplanted by the molecular orbital model. In this model, as atoms are brought together, the atomic orbitals interact so as to form a set of molecular orbitals, which extend over the entire molecule. Half of these orbitals tend to be bonding orbitals, while the other half are anti-bonding orbitals. Electrons in bonding orbitals result in the formation of a chemical bond, while those in anti-bonding orbitals prevent bonding. Electrons may also occupy non-bonding orbitals, which are neither bonding nor anti-bonding. The formation of a chemical bond is only possible when more electrons occupy bonding orbitals than anti-bonding orbitals.
The number of pairs of electrons in bonding orbitals in excess of those in anti-bonding orbitals determines the bond order. For example, in a diatomic molecule, there is a single bond if there is a balance of two electrons in bonding orbitals (as for H2), a double bond if there is a balance of four electrons in bonding orbitals (as for O2), and a triple bond if there is a balance of six electrons in bonding orbitals (as for N2). Bond order needn't be integral, and bonds can be delocalized among more than two atoms. In benzene for instance, the bond order is 1.5 amongst any two carbon atoms. This is called resonance. Bond length and bond dissociation energy are related inversely with bond order--the higher the order of the bond, the shorter and stronger the bond, for any given set of atoms.
Both carbon and silicon can theoretically form quadruple bonds. However, they are explosively unstable. The three shared orbitals in a triple bond can be imaged as being left, right, and out through you computer screen. The fourth orbital must bend these three away, and that makes quadruple bonds explosive; C2 molecules must be observed in a vacuum environment. Si2 molecules are even more unstable.
Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and frequency spectra of simple molecules with a high degree of accuracy. Currently, 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 the case of small molecules, energy calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.
Dative covalent bonding occurs when one atom gives both of the electrons in the bond.