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Encyclopedia > Double bond
Covalently bonded hydrogen and carbon (Methane)

Covalent bonding is a form of chemical bonding characterized by the sharing of one or more pairs of electrons between atoms, in order to produce a mutual attraction, which holds the resultant molecule together. Atoms tend to share electrons in such a way that their outer electron shells are filled. Such bonds are always stronger than the intermolecular hydrogen bond and similar in strength to or stronger than the ionic bond.


Covalent bonding most frequently occurs between atoms with similar (high) electronegativities, where to completely remove an electron from one atom requires too much energy. Covalent bonds are more common between non-metals, whereas ionic bonding is more common between a metal atom and a non-metal atom.


Covalent bonding tends to be stronger than other types of bonding, such as ionic bonding. Unlike ionic bonds, 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.

Contents

History of the covalent bond

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 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.


Bond order

Bond order is the scientific term used to describe the number of pairs of electrons shared between atoms forming a covalent bond. The most common type of covalent bond is the single bond, the sharing of only one pair of electrons between two individual atoms. All bonds with more than one shared pair are called multiple covalent bonds. The sharing of two pairs is called a double bond and the sharing of three pairs is called a triple bond. An example of a double bond is nitrous acid (between N and O), and an example of a triple bond is in hydrogen cyanide (between C and N).


Quadruple bonds, though rare, also exist. Both carbon and silicon can theoretically form these; however, the formed molecules are explosively unstable. The three shared orbitals in a triple bond can be imaged as left, right, and up. The fourth orbital must bend these three away, leading to instability: C2 molecules must be observed in a vacuum environment, and Si2 molecules are even more unstable. Stable quadruple bonds are observed as transition metal-metal bonds, usually between two transition metal atoms in organometallic compounds.


Sextuple bonds of order 6 have also been observed in transition metals in the gaseous phase and are even more rare.


A special case is called a dative covalent bond, also known as a coordinate covalent bond, which occurs when one atom gives both of the electrons in the bond.


Resonance

Some structures can have more than one valid Lewis Dot Structure; for example, Ozone, O3. The LDS tells us the center O will have a single bond with one O and a double bond with the other O. However, it does not tell us which O should receive the double bond, and, as a result, the first O and the second O have equal chance of having the double bond. These two structures are called resonance structures. In actuality, the true structure is a resonance hybrid between all of its resonance structures. In our example, 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 alternate between single and double bonds according to their LDS. In actuality, the electrons tend to be disambiguously spaced out evenly within the ring of atoms. Electrons in aromatic structures are often represented with a ring inside of the circle of atoms instead of lines or dots.


Current theory

Today the valence bond model has been supplemented with the molecular orbital model. In this model, as atoms are brought together, the atomic orbitals interact to form hybrid molecular 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 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.


 
 

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