"Covalent" redirects here. For other uses, see Covalent (disambiguation).
A covalent bond forming H
2 (right) where two hydrogen atoms share the two electrons
A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. The stable balance of attractive and repulsive forces between atoms when they share electrons is known as covalent bonding.[1] For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration.
Covalent bonding includes many kinds of interactions, 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", in essence, means that the atoms share "valence", such as is discussed in valence bond theory.
In the molecule H
2, the hydrogen atoms share the two electrons via covalent bonding.[5] Covalency is greatest between atoms of similar electronegativities. Thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons over more than two atoms is said to be delocalized.
Contents
- 1 History
- 2 Physical properties of covalent compounds (polar and non-polar)
- 3 Polarity of covalent bonds
- 4 Subdivision of covalent bonds
- 5 See also
- 6 References
- 7 Sources
- 8 External links
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 by indicating electrons shared 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 that a given atom shares with its neighbors."[6]
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.[7] He introduced the Lewis notation or electron dot notation or 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 bonds and triple bonds. An alternative form of representation, not shown here, has bond-forming electron pairs represented as solid lines.
Lewis proposed that an atom forms enough covalent bonds to form a full (or closed) outer electron shell. In the methane diagram shown here, the carbon atom has a valence of four and is, therefore, surrounded by eight electrons (the octet rule), four from the carbon itself and four from the hydrogens bonded to it. Each hydrogen has a valence of one and is surrounded by two electrons (a duet rule) - its own one electron plus one from the carbon. The numbers of electrons correspond to full shells in the quantum theory of the atom; the outer shell of a carbon atom is the n=2 shell, which can hold eight electrons, whereas the outer (and only) shell of a hydrogen atom is the n=1 shell, which can hold only two.
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 (molecular hydrogen) in 1927.[8] 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.
Physical properties of covalent compounds (polar and non-polar)
| Physical properties |
Covalent compounds |
| States (at room temperature) |
Solid, liquid, gas |
| Electrical conductivity |
Usually none |
| Boiling point and Melting point |
Varies, but usually lower than ionic compounds |
| Solubility in water |
Varies, but usually lower than ionic compounds |
| Thermal conductivity |
Usually low |
Polarity of covalent bonds
Covalent bonds are affected by the electronegativity of the connected atoms. Two atoms with equal electronegativity will make nonpolar covalent bonds such as H–H. An unequal relationship creates a polar covalent bond such as with H−Cl.
Subdivision of covalent bonds
There are three types of covalent substances: individual molecules, molecular structures, and macromolecular structures. Individual molecules have strong bonds that hold the atoms together, but there are negligible forces of attraction between molecules. Such covalent substances are usually gases, for example, HCl, SO2, CO2, and CH4. In molecular structures, there are weak forces of attraction. Such covalent substances are low-boiling-temperature liquids (such as ethanol), and low-melting-temperature solids (such as iodine and solid CO2). Macromolecular structures have large numbers of atoms linked in chains or sheets (such as graphite), or in 3-dimensional structures (such as diamond and quartz). These substances have high melting and boiling points, are frequently brittle, and tend to have high electrical resistivity. Elements that have high electronegativity, and the ability to form three or four electron pair bonds, often form such large macromolecular structures.[9]
See also
- Bonding in solids
- Bond order
- Coordinate Covalent Bond, also known as Dipolar Bond or Dative Covalent Bond
- Covalent Bond Classification (or LXZ notation)
- Covalent radius
- Disulfide bond
- Hybridization
- Hydrogen bond
- Ionic bond
- Linear combination of atomic orbitals
- Metallic bonding
- Noncovalent bonding
- Resonance (chemistry)
- Shared pair
References
- ^ Campbell, Neil A.; Brad Williamson; Robin J. Heyden (2006). Biology: Exploring Life. Boston, Massachusetts: Pearson Prentice Hall. ISBN 0-13-250882-6. Retrieved 2012-02-05.
- ^ March, Jerry (1992). Advanced organic chemistry: reactions, mechanisms, and structure. John Wiley & Sons. ISBN 0-471-60180-2.
- ^ Gary L. Miessler; Donald Arthur Tarr (2004). Inorganic chemistry. Prentice Hall. ISBN 0-13-035471-6.
- ^ Merriam-Webster – Collegiate Dictionary (2000).
- ^ "Chemical Bonds". Hyperphysics.phy-astr.gsu.edu. Retrieved 2013-06-09.
- ^ 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.
- ^ 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.
- ^ W. Heitler and F. London, Zeitschrift für Physik, vol. 44, p. 455 (1927). English translation in Hettema, H. (2000). Quantum chemistry: classic scientific papers. World Scientific. pp. 140–. ISBN 978-981-02-2771-5. Retrieved 2012-02-05.
- ^ Stranks, D. R.; M. L. Heffernan, K. C. Lee Dow, P. T. McTigue, G. R. A. Withers (1970). Chemistry: A structural view. Carlton, Victoria: Melbourne University Press. p. 184. ISBN 0-522-83988-6.
Sources
- "Covalent bonding — Single bonds". chemguide. 2000. Retrieved 2012-02-05.
- "Electron Sharing and Covalent Bonds". Department of Chemistry University of Oxford. Retrieved 2012-02-05.
- "Chemical Bonds". Department of Physics and Astronomy, Georgia State University. Retrieved 2012-02-05.
External links
- Covalent Bonds and Molecular Structure
- Structure and Bonding in Chemistry—Covalent Bonds
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Chemical bonds
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|
Intramolecular
("strong") |
Covalent
bonds |
| By symmetry |
- Sigma (σ)
- Pi (π)
- Delta (δ)
- Phi (φ)
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| By multiplicity |
- 1 (single)
- 2 (double)
- 3 (triple)
- 4 (quadruple)
- 5 (quintuple)
- 6 (sextuple)
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| Miscellaneous |
- Agostic bond
- Bent bond
- Dipolar bond
- Pi backbond
- Hapticity
- Conjugation
- Hyperconjugation
- Antibonding
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| Resonant bonds |
- Electron-deficient
- Hypercoordination
- Aromaticity
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| Metallic bonds |
|
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| Ionic bonding |
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Intermolecular
("weak") |
van der Waals
forces |
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Hydrogen
bond |
- Low-barrier
- Resonance-assisted
- Symmetric
- Dihydrogen bonds
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Other
noncovalent |
- Mechanical bond
- Halogen bond
- Aurophilicity
- Intercalation
- Stacking
- Cation–pi bond
- Anion–pi bond
- Salt bridge
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Concepts in organic chemistry
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- Aromaticity
- Covalent bonding
- Functional groups
- Nomenclature
- Organic compounds
- Organic reactions
- Organic synthesis
- Publications
- Spectroscopy
- Stereochemistry
- List of organic compounds
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