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3.10: Bonding in Ethyne

  • Page ID
    221780
  • Objectives

    After completing this section, you should be able to

    1. use the concept of sp hybridization to account for the formation of carbon-carbon triple bonds, and describe a carbon-carbon triple bond as consisting of one σ bond and two π bonds.
    2. list the approximate bond lengths associated with typical carbon-carbon single bonds, double bonds and triple bonds. [You may need to review Sections 1.7 and 1.8.]
    3. list the approximate bond angles associated with sp3-, sp2- and sp‑hybridized carbon atoms and predict the bond angles to be expected in given organic compounds. [If necessary, review Sections 1.6, 1.7 and 1.8.]
    4. account for the differences in bond length, bond strength and bond angles found in compounds containing sp3-, sp2- and sp‑hybridized carbon atoms, such as ethane, ethylene and acetylene.
    Key Terms

    Make certain that you can define, and use in context, the key term below.

    • sp hybrid orbital
    Study Notes

    The bond angles associated with sp3-, sp2- and sp‑hybridized carbon atoms are approximately 109.5°, 120° and 180°, respectively.

    Bonding in acetylene

    Finally, the hybrid orbital concept applies well to triple-bonded groups, such as alkynes and nitriles. Consider, for example, the structure of ethyne (another common name is acetylene), the simplest alkyne.

    Bond line drawing of ethyne (acetylene)

    ethyne
    (acetylene)

    This molecule is linear: all four atoms lie in a straight line. The carbon-carbon triple bond is only 1.20Å long. In the hybrid orbital picture of acetylene, both carbons are sp-hybridized. In an sp-hybridized carbon, the 2s orbital combines with the 2px orbital to form two sp hybrid orbitals that are oriented at an angle of 180°with respect to each other (eg. along the x axis). The 2py and 2pz orbitals remain non-hybridized, and are oriented perpendicularly along the y and z axes, respectively.

    Carbon has one lone pair in the 2 s orbital and two unpaired electrons in the 2 p orbital. After hybridization, there two unpaired electrons in the s p orbital, one unpaired electron in the 2 py orbital, and one unpaired electron in the 2 p z orbital.

    The 2 p y orbital is perpendicular to the s p orbitals. and the 2 P z orbital is perpendicular to the plane of the page.

    The C-C sigma bond is formed by the overlap of one sp orbital from each of the carbons, while the two C-H sigma bonds are formed by the overlap of the second sp orbital on each carbon with a 1s orbital on a hydrogen. Each carbon atom still has two half-filled 2py and 2pz orbitals, which are perpendicular both to each other and to the line formed by the sigma bonds. These two perpendicular pairs of p orbitals form two pi bonds between the carbons, resulting in a triple bond overall (one sigma bond plus two pi bonds).

    depiction of sigma bonding framework in ethyne.svg depiction of pi bonding framework in ethyne.svg
    sigma bonding in ethylene pi bonding in ethylene

    Acetylene is said to have three sigma bonds and two pi bonds. The carbon-carbon triple bond in acetylene is the shortest (120 pm) and the strongest (965 kJ/mol) of the carbon-carbon bond types. Because each carbon in acetylene has two electron groups, VSEPR predicts a linear geometry and and H-C-C bond angle of 180o.

    The carbon-carbon bond length is 120 pm while the carbon hydrogen length is 106 pm. The bond angle is 180 degrees.

    Comparison of C-C bonds Ethane, Ethylene, and Acetylene

    Molecule Bond Bond Strength (kJ/mol) Bond Length (pm)
    Ethane, CH3CH3 (sp3) C-C (sp3) 376 154
    Ethylene, H2C=CH2 (sp2) C=C (sp2) 728 134
    Acetylene, HC≡CH (sp) C≡C (sp) 965 120

    Notice that as the bond order increases the bond length decreases and the bond strength increases.

    The hybrid orbital concept nicely explains another experimental observation: single bonds adjacent to double and triple bonds are progressively shorter and stronger than ‘normal’ single bonds, such as the one in a simple alkane. The carbon-carbon bond in ethane (structure A below) results from the overlap of two sp3 orbitals.

    carbon-carbon bond orbitals in methane, A.svg carbon-carbon bond orbitals in propene, B.svg carbon-carbon bond orbitals in propyne, C.svg
    A B C

    In propene (B), however, the carbon-carbon single bond is the result of overlap between an sp2 orbital and an sp3 orbital, while in propyne (C) the carbon-carbon single bond is the result of overlap between an sp orbital and an sp3 orbital. These are all single bonds, but the single bond in molecule C is shorter and stronger than the one in B, which is in turn shorter and stronger than the one in A.

    The explanation here is relatively straightforward. An sp orbital is composed of one s orbital and one p orbital, and thus it has 50% s character and 50% p character. sp2 orbitals, by comparison, have 33% s character and 67% p character, while sp3 orbitals have 25% s character and 75% p character. Because of their spherical shape, 2s orbitals are smaller, and hold electrons closer and ‘tighter’ to the nucleus, compared to 2p orbitals. Consequently, bonds involving sp + sp3 overlap (as in alkyne C) are shorter and stronger than bonds involving sp2 + sp3 overlap (as in alkene B). Bonds involving sp3-sp3overlap (as in alkane A) are the longest and weakest of the group, because of the 75% ‘p’ character of the hybrids.

    Hybridization Summary

    • A single bond is a sigma bond.
    • A double bond is made up of a sigma bond and a pi bond.
    • A triple bond is made up of a sigma bond and two pi bonds.
    • Sigma bonds are made by the overlap of two hybrid orbitals or the overlap of a hybrid orbital and a s orbital from hydrogen.
    • Pi bonds are made by the overlap of two unhybridized p orbitals.
    • Lone pair electrons are usually contained in hybrid orbitals.

    The hybrid orbitals used (and hence the hybridization) depends on how many electron groups are around the atom in question. An electron group can mean either a bonded atom or a lone pair. Molecular geometry is also decided by the number of electron groups so it is directly linked to hybridization.

    # of Electron Groups Hybrid Orbital Used Example Basic Geometry Basic Bond Angle
    2 sp example of sp hybridized carbon.svg Linear 180o
    3 sp2 example of sp2 hybridized carbon.svg Trigonal Planar 120o
    4 sp3 example of sp3 hybridized carbon.svg Tetrahedral 109.5o

    Exercises

    1) For the molecule acetonitrile:

    Bond line drawing of acetonitrile.

    a) How many sigma and pi bonds does it have?

    b) What orbitals overlap to form the C-H sigma bonds?

    c) What orbitals overlap to form the C-C sigma bond?

    d) What orbitals overlap to form the C-N sigma bond?

    e) What orbitals overlap to the form the C-N pi bonds?

    f) What orbital contains the lone pair electrons on nitrogen?

    Solutions

    1)

    a) 5 sigma and 2 pi

    b) An sp3 hybrid orbital from carbon and an a s orbital from hydrogen.

    c) An sp3 hybrid orbital from one carbon and an a sp3 orbital from the other carbon.

    d) An sp hybrid orbital from carbon and an a sp orbital from nitrogen.

    e) An py and pz orbital from carbon and an py and pz orbital from nitrogen.

    f) An sp hybrid orbital.

    Questions

    Q1.9.1

    1-Cyclohexyne is a very strained molecule. By looking at the molecule explain why there is such a intermolecular strain using the knowledge of hybridization and bond angles.

    Solutions

    S1.9.1

    The alkyne is a sp hybridized orbital. By looking at a sp orbital, we can see that the bond angle is 180°, but in cyclohexane the regular angles would be 109.5°. Therefore the molecule would be strained to force the 180° to be a 109°.

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