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Covalent bond

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Covalent bonding is a form of chemical bonding that is characterized by the sharing of pairs of electrons between atoms. In short, attraction-to-repulsion stability that forms between atoms when they share electrons is known as covalent bonding. The term "covalence", in regards to bonding, was first used in 1919 by Irving Langmuir in a Journal of American Chemical Society article entitled "Isomorphism, Isosterism, and Covalence" wherein he states:[1]

I propose that the number of electrons which any given atom shares with the adjacent atoms be called the covalence of that atom.

Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-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. Because covalent bonding entails sharing of electrons, is 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.

Schemes depicting covalent (left) and polar covalent (right) bonding in a diatomic molecule. The arrows represent electrons provided by the participating atoms.


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Resonance

Many bonding situations can be described with 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 may 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.

Current theory

Today the valence bond model has been supplanted 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.

References

  1. ^ Langmuir, I. (1919). "Isomorphism, Isosterism and Covalence." J. Am. Chem. Soc.; 1919; 41(10); 1543-1559.
  2. ^ March, J. “Advanced Organic Chemistry” $th Ed. J. Wiley and Sons, 1992: 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).

See also

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