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Covalent bond

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Covalent bond between two atoms of hydrogen.

A covalent bond is a type of chemical bond produced by the joining of two atoms that essentially "share" electrons between them.[1] There is no sharp boundary between ionic bond and covalent bond.[2]


Any atom, other than that of a noble gas, has an incomplete outer electron "shell." This is an inherently unstable configuration which can resolve in one of two ways:

  1. An atom may shed, or acquire, various electrons in order to achieve an electron configuration similar to that of the nearest noble gas. The result is an ion.
  2. Two atoms may share electrons. An electron so shared spends part of its time completing a stable electron configuration for each of the two atoms that share it. This kind of sharing is a covalent bond.

Covalent bonds are found both in molecules and in polyatomic ions. In the latter case, even the mutual sharing of electrons is not sufficient to bring stability to the electron configurations of all atoms. In that case, the entire structure might still have a deficit or excess of electrons, giving it an ionic valence.


A covalent bond may be single, double, or triple.[3] (Quadruple covalent bonds are unknown to modern chemistry.) The descriptor refers to the number of electrons that the two atoms share between them.

Double and triple bonds impose drastic geometrical constraints on any molecules of which they are a part. For example, double bonds between carbon atoms force other substances bound to these atoms to attach at 120 degrees on each side of the bond, and in the same plane. Triple bonds force the other substances to bind diametrically opposite to the bond.

Resonance structures

Main Article: Resonance structure

Any given atom (most typically of carbon) can always have a single bond to one atom and a double bond to another. But often the second and third atoms are of the same chemical element, and those atoms in turn have the same types of bonds (or no covalent bonds at all) to other atoms. In such a case, no reasonable ground exists to assume that one atom is any more tightly bound than the other--and in fact, chemists have discovered that the overall bindings of the various atoms have equal strength.

Current theory holds that resonance takes place between the two alternative schemes for binding the atoms together. Resonance does not mean that the molecule, or polyatomic ion, oscillates between two binding schemes. Rather, the atoms will form bonds that are intermediate between single and double bonds. Such a scheme is a resonance structure, and accomplishes all the goals of a covalent bond.

Ozone (O3), the carboxylic group of carboxylic acids, and the aromatic ring structure of benzene are three of the most common resonance structures known to chemistry.

Dative covalent bonds

The poisonous gas carbon monoxide is a typical example of dative covalent bond.

Often, when atoms bond together, both the electrons in a covalent bond came from just one of the atoms involved.[4] This kind of bond is named dative bond, coordinate bond[4] or dipolar bond. To represent what atom is donating the electrons an arrow is sometimes used indicating the direction of the donor to the receptor as shown in the figure at the right. Two oxygen electrons bind to two carbon electrons in covalent bonds, but also the oxygen atom provides more two electrons to complete the outer electron "shell", forming a dative bond indicated by the arrow.

Polar covalent bonds

Hydrogen chloride

A polar covalent bond or polar bond is a covalent bond in which the bonding electrons spend more time near one of the atoms than the other.[5] Covalent bonds are affected by the electronegativity of the connected atoms. Two atoms with different electronegativity will make a polar bond such as with Hydrogen chloride (H−Cl). If there is a difference in the relative ability to attract electrons large enough, an ionic bond is formed.[6] It is possible to consider the polar covalent bond as intermediate between a nonpolar covalent bond such as with the Hydrogen molecule (H-H), and an ionic bond, for instance NaCl.[5]


  1. Silberberg, Martin S (2010). Principles of General Chemistry (2nd ed.). Boston: McGraw-Hill. p. 280. ISBN 978–0–07–351108–5. 
  2. Huheey, James E.; Keiter, Ellen A.; Keiter, Richard L (1993). Inorganic Chemistry: Principles of Structure and Reactivity (4th ed.). New York: Harper Collins College Publishers. p. 92. ISBN 0-06-042995-X. 
  3. Brown, Lawrence S.; Holme, Thomas A (2011). Chemistry for Engineering Students (2nd ed.). Belmont, CA: Brooks/Cole. p. 211. ISBN 978-1-4390-4791-0. 
  4. 4.0 4.1 Conoley, Chris; Hills, Phil (2008). Chemistry (3rd ed.). London: Harper Collins Publishers. p. 76. ISBN 978-0-00-726748-4. 
  5. 5.0 5.1 Ebbing, Darrell D.; Gammon, Steven D (2009). General Chemistry (9th ed.). Boston: Houghton Mifflin Company. p. 345. ISBN 978-0-618-85748-7. 
  6. Brown, Theodore L.; LeMay Jr., H. Eugene; Bursten, Bruce E.; Murphy, Catherine J.; Woodward, Patrick (2009). Chemistry: The Central Science (11th ed.). Upper Saddle River, NJ: Pearson Education, Inc.. p. 308. ISBN 978-0-13-600617-6.