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A covalent bond is a form of chemical bonding that is characterized by the sharing of pairs of electrons between atoms, and other covalent bonds. In short, the attraction-to-repulsion stability that forms between atoms when they share electrons is known as covalent bonding.[1]

Covalent bonding includes many kinds of interaction, including σ-bonding, π-bonding, metal to metal bonding, agostic interactions, and three-center two-electron bonds.[2][3] The term covalent bond dates from 1939.[4] 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 H2, the hydrogen atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities. Thus, covalent bonding does not necessarily require the two atoms be of the same elements, only that they be of comparable electronegativity. Although covalent bonding entails sharing of electrons, it is not necessarily delocalized. Furthermore, in contrast to electrostatic interactions ("ionic bonds") the strength of covalent bond depends on the angular relation between atoms in polyatomic molecules.

Contents

History

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.

The term "covalence" in regard to bonding was first used in 1919 by Irving Langmuir in a Journal of the American Chemical Society article entitled "The Arrangement of Electrons in Atoms and Molecules". Langmuir wrote that "we shall denote by the term covalence the number of pairs of electrons which a given atom shares with its neighbors."[5]

The idea of covalent bonding can be traced several years before 1919 to Gilbert N. Lewis, who in 1916 described the sharing of electron pairs between atoms.[6] He introduced the so called 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 of representation, not shown here, has bond-forming electron pairs represented as solid lines. 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.[7] 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.

Resonance

Many bonding situations can be described with more than one valid Lewis dot structure (for example, ozone, O3). In a Lewis diagram of O3, the center atom will have a single bond with one atom and a double bond with the other. The Lewis 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 may alternate between single and double bonds according to their Lewis diagram. 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.

Lewis dot structures for molecules with resonance are shown by creating the dot structure for every possible form, placing brackets around each structure, and connecting the boxes with double-headed arrows.

Current theory

Today the valence bond model has been augmented 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.

Polarity of covalent bonds

Covalent bonds are affected by the electronegativity of the connected atoms. Two atoms with equal electronegativity will make non-polar covalent bonds such as H-H. An unequal relationship creates a polar covalent bond such as with Cl-H.

See also

References

  1. ^ Campbell, Neil A.; Brad Williamson; Robin J. Heyden (2006). Biology: Exploring Life. Boston, Massachusetts: Pearson Prentice Hall. ISBN 0-13-250882-6. http://www.phschool.com/el_marketing.html. 
  2. ^ March, J. “Advanced Organic Chemistry” 4th Ed. J. Wiley and Sons, 1991: New York. ISBN 0-471-60180-2.
  3. ^ G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.
  4. ^ Merriam-Webster - Collegiate Dictionary (2000).
  5. ^ Langmuir, Irving (1919-06-01). "The Arrangement of Electrons in Atoms and Molecules". Journal of the American Chemical Society 41 (6): 868-934. doi:10.1021/ja02227a002. http://dx.doi.org/10.1021/ja02227a002. Retrieved 2009-11-10. 
  6. ^ Lewis, Gilbert N. (1916-04-01). "THE ATOM AND THE MOLECULE.". Journal of the American Chemical Society 38 (4): 762-785. doi:10.1021/ja02261a002. http://dx.doi.org/10.1021/ja02261a002. Retrieved 2009-11-10. 
  7. ^ W. Heitler and F. London, Zeitschrift für Physik, vol. 44, p. 455 (1927). English translation in H. Hettema, Quantum Chemistry, Classic Scientific Papers, World Scientific, Singapore (2000).

Notes

External links

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