Covalent bonding is an intramolecular form of chemical bonding characterized by the sharing of one or more pairs of electrons between two components, producing a mutual attraction that 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.
In contrast to the ionic and metalic bond, the covalent bond is directional. I.e. the bond angles have a great impact on the strength of the bond. Because of the directional character of the bond, covalently bound materials are more difficult to deform than metals. The cause of the directionallity is the form of the S, P, and SP-hybrid orbitals.
Covalent bonding most frequently occurs between atoms with similar electronegativities. For this reason, non-metals tend to engage in covalent bonding more readily since metals have access to metallic bonding, where the easily-removed electrons are more free to roam about. For non-metals, liberating an electron is more difficult, so sharing is the only option when confronted with another species of similar electronegativity.
However, covalent bonding involving metals is particularly important, especially in industrial catalysis and process chemistry. Many polymerization techniques require catalysis involving metal-organic covalent bonds. In their more useful applications, metals often engage in more exotic covalent bonding, such as those between a metal and the σ bond of molecular hydrogen, or between a metal and the π bond of an alkane or alkene.
Covalently bonded hydrogen and carbon in a molecule of methane. One way of representing covalent bonding in a molecule is with a dot and cross diagram.
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 alternative form, in which bond-forming electron pairs are represented as solid lines, is shown alongside.
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.
There are two types of covalent bonds: Polar covalent bonds, and non-polar (or pure) covalent bonds. The most widely-accepted definition of polar covalence is the occurrence of the atoms involved of an electronegativity difference less than 1.67 (though some texts read 1.7), but greater than zero. A pure covalent bond is a bond that occurs when the atoms involved have an electronegativity difference of zero (though some texts read less than 0.2).
Pure covalent bonds (which are usually non-soluble, are electrically non-conductive, and tend to exist as individual molecules), and ionic bonds (which are soluble, are electrically conductive when molten or in solution, and, in general, tend to exist in a crystalline form) are on two opposite ends of the spectrum and have different properties. Polar covalent bonds fall in the middle and have properties of both.
Bond order is a term that describes the number of pairs of electrons shared between atoms forming a covalent bond.
1) The most common type of covalent bond is the single bond, sharing 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 covalent bonds.
2) 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.
3) 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.
4) Quadruple bonds, though rare, exist. Both carbon and silicon can theoretically form these; however, the formed molecules are explosively unstable. Stable quadruple bonds are observed as transition metal-metal bonds, usually between two transition metal atoms in organometallic compounds. Molybdenum and Ruthenium are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in Di-tungsten tetra(hpp).
5) Quintuple bonds are found to exist in certain chromium dimers.
6) Sextuple bonds, of order 6, have also been observed in transition metals in the gaseous phase at very low temperatures and are extremely rare.
Other more exotic bonds, such as three center bonds are known and defy the conventions of bond order. It is also important to note that bond order is an integer value only in the elementary sense and is often fractional in more advanced contexts.
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.
Some structures can have more than one valid Lewis Dot Structure (for example, ozone, O3). In an LDS diagram of O3, 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 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 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.
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 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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