Single, Double, and Triple Bonds   Recently updated !


Single, Double, and Triple Bonds
Comparison of single, double, and triple bonds.

Single, double, and triple bonds are three types of covalent bonds mainly involving nonmetals. Atoms form these bonds as a way of obtaining the most stable electron configuration, according to the octet rule. Since metals usually need more than three electrons to achieve this, they less commonly form these types of bonds. Here is a closer look at single, double, and triple bonds, along with examples of each type and their properties.

Review of Covalent Bonding

Irving Langmuir first described covalence in his 1919 article “The Arrangement of Electrons in Atoms and Molecules” in the Journal of the American Chemical Society. According to Langmuir, covalence is the number of pairs of electrons shared between an atom and its neighbor.

  • Two atoms form a bond to increase their stability, which results in a loss of energy. In other words, covalent bond formation is an exothermic process.
  • Covalent bond formation occurs between the valence electrons of two atoms.
  • Maximum stability occurs when atoms achieve the nearest noble gas configuration. A filled shell is most stable, followed by a half-filled shell.
  • Whether an atom forms a single, double, or triple bond depends on how many electrons it needs to achieve the most stable electron configuration.

Single Bond

A single bond is a covalent bond that occurs when two atoms share one electron pair. Atoms that form this type of bond are one electron away from a noble gas configuration, so elements participating in single bonds are hydrogen and the halogens, with each other or with other elements. There are some exceptions. The notation for a single bond is a single dash between the atoms, such as H-H or Cl-Cl.

Examples of single bonds are H2 (hydrogen, H-H), F2 (fluorine, F-F), some other diatomic molecules, hydrochloric acid (HCl, H-Cl), methane (CH4), and NH3 (ammonia).

Usually, a single bond is a sigma bond, although the bond in diboron (B2) is a pi bond. A sigma bond forms by head-on overlapping of σ orbitals. Unlike in double and triple bonds, atoms are free to rotate around a single bond.

Double Bond

A double bond forms when two atoms share two electron pairs or six electrons. The symbol for this is a double dash or equal sign between the two atoms, like O=O. Carbon and members of the oxygen family of elements (the chalcogens) participate in double bonds.

Examples of double bonds are O2 (oxygen, O=O), CO2 (carbon dioxide, O=C=O), and C2H2 (ethylene, H-C=C-H).

The double bond consists of one sigma (σ) bond and one pi (π) bond. A pi bond forms by sideways overlapping of p orbitals.

Triple Bond

A triple bond forms when two atoms share three electron pairs. The symbol for the triple bond is a triple dash, as in NN. The most common triple bond occurs between two carbon atoms in the alkynes. Nitrogen also forms triple bonds with itself and with carbon.

Examples of molecules with triple bonds include nitrogen (N2, NN), carbon monoxide (CO, CO), acetylene (C2H2, H-CC-H), and cyanogen (C2N2, NC-CN).

A triple bond consists of one sigma bond and two pi bonds.

Comparing Single, Double, and Triple Bonds

Single BondDouble BondTriple Bond
Valence ElectronsShare 1 pair
(2 electrons)
Share 2 pairs
(4 electrons)
Share 3 pairs
(6 electrons)
Bond LengthLongestIntermediateShortest
Bond StrengthWeakestIntermediateStrongest
ReactivityLowestIntermediateHighest
Rotation Around BondYesNoNo
OrbitalsOne sigmaOne sigma, one piOne sigma, two pi
NotationSingle dash (C-C)Double dash (C=C)Triple dash (CC)

References

  • McMurry, John (2016). Chemistry (7th ed.). Pearson. ISBN 978-0-321-94317-0.
  • Miessler, Gary L.; Tarr, Donald Arthur (2004). Inorganic Chemistry. Prentice Hall. ISBN 0-13-035471-6.
  • Pauling, L. (1960). The Nature of the Chemical Bond. Cornell University Press.
  • Pyykkö, Pekka; Riedel, Sebastian; Patzschke, Michael (2005). “Triple-Bond Covalent Radii”. Chemistry: A European Journal11 (12): 3511–20. doi:10.1002/chem.200401299
  • Weinhold, F.; Landis, C. (2005). Valency and Bonding. Cambridge. ISBN 0-521-83128-8.