Covalent bond

Early concepts in covalent bonding arose from this type of image of a molecule of methane. Here, four hydrogen atoms are bound to a central carbon atom by four covalent bonds. Each covalent bond is formed by the sharing of two electrons—one from the carbon atom and one from a hydrogen atom.

Covalent bonding is a description of chemical bonding that is characterized by the sharing of one or more electrons between two atoms. In general, bonds are defined by a mutual attraction that holds the resultant molecule together. Often, bonding occurs in such a way that the outer electron shells of the participating atoms become filled. In contrast to electrostatic interactions labeled as "ionic bonds," the strength of covalent bond depends on the angular relation between atoms in polyatomic molecules. Covalent bonding is most important between atoms with similar electronegativities. Covalent bonding is often delocalized. Covalent bonding is a broad concept and includes many kinds of interactions, including σ-bonding, π-bonding, metal-metal bonding, agostic interactions, and three-center two-electron bonds.

History

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.

Bond polarity

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 or equal to 2.1 but greater or equal to 0.5. 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).

File:CovalentBond.png

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

Bond order

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, the 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 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 are found in the transition metals. Molybdenum and rhenium 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 have been found to exist in certain dichromium compounds.
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.

Most bonding of course, is not localized, so the following classification, while powerful and pervasive, is of limited validity. Three center bond do not conform readily to the above conventions.

Coordinate covalent bonds

A special case is called a coordinate covalent bond, also known as a dative covalent bond, which occurs when one atom gives both of the electrons in the bond with the other ion. The classic example is borane-ammonia.

777777777 A coordinate covalent bond (also known as dative bond) is a description of covalent bonding in many kinds of compounds. The distinction from ordinary covalent bonding is artificial, but the terminology is popular in textbooks, especially those describing coordination compounds. Once the bonds have been formed using this, its strength and description is no different from that of other polar covalent bonds.

Coordinate covalent bonds are invoked when a Lewis base (an electron donor or giver) donates a pair of electrons to a Lewis acid (an electron acceptor) to give a so-called adduct. The process of forming a dative bond is called coordination. The electron donor acquires a positive formal charge, while the electron acceptor acquires a negative formal charge.

Examples

Classically, any compound that contains a lone pair of electrons is capable of forming a coordinate bond. The bonding in diverse chemical compounds can be described as coordinate covalent bonding.

• carbon monoxide (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the carbon atom and the oxygen atom. This highly unusual description illustrates the flexibility of this bonding description. Thus in CO, carbon is the electron acceptor and oxygen is the electron donor.
• ammonium ion (NH₄⁺), can be viewed as consisting of four coordinate covalent bonds between the protons (the H⁺ ion) and the nitrogen trianion "N^3-".
• beryllium dichloride (BeCl₂) is described as electron deficient in the sense that the triatomic species (which does exist in the gas phase) features Be centers with four valence electrons. When treated with excess chloride, the Be²⁺ ion binds four chloride ions to form tetrachloroberyllate anion, BeCl₄^2-, wherein all ions achieve the octet configuration of electrons.

Coordination compounds

Coordinate bonding is popularly used to describe coordination complexes, especially involving metal ions. In such complexes, several Lewis bases "donate" their "free" pairs of electrons to an otherwise naked metal cation, which acts as a Lewis acid and "accepts" the electrons. Coordinate bonds form and the resulting compound is called a coordination complex, and the electron donors are called ligands. A more useful description of bonding in coordination compounds is provided by Ligand Field Theory, which embraces molecular orbitals as a description of bonding in such polyatomic compounds.

Many chemical compounds can serve as ligands, often these contain oxygen, sulfur, nitrogen, and halide ions. The most common ligand is water (H₂O), which form coordination complexes with metal metal ions, e.g. [Cu(H₂O)₆]²⁺. Ammonia (NH₃) is also a common ligand. Anions are common ligands, especially fluoride (F^-), chloride (Cl^-), and cyanide (CN^-).

A coordinate bond is sometimes represented by an arrow pointing from the donor of the electron pair to the acceptor of the electron pair.

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.

See also

• Chemical bond
• Ionic bond
• Linear combination of atomic orbitals
• Metallic bonding
• Hybridisation
• Hydrogen bond
• Noncovalent bonding
• Disulfide bond
• Sigma bond
• Pi bond
• Delta bond

References

ISBN links support NWE through referral fees

• March, J. “Advanced Organic Chemistry” $th Ed. J. Wiley and Sons, 1992: New York. ISBN 0-471-60180-2.
• G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.
• http://www.chemguide.co.uk/atoms/bonding/covalent.html
• http://www.chem.ox.ac.uk/vrchemistry/electronsandbonds/intro1.htm
• http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/bond.html#c5

External links

• Covalent Bonds and Molecular Structure
• Bonding tutorials
• Chemguide

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