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1.9: sp Hybrid Orbitals and the Structure of Acetylene

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    In addition to forming single and double bonds by sharing two and four electrons, respectively, carbon can also form a triple bond by sharing six electrons. To account for the triple bond in a molecule such as acetylene, \(\mathrm{H}-\mathrm{C} \equiv \mathrm{C}-\mathrm{H}\), we need a third kind of hybrid orbital, an sp hybrid. Imagine that, instead of combining with two or three \(p\) orbitals, a carbon 2 s orbital hybridizes with only a single \(p\) orbital. Two \(s p\) hybrid orbitals result, and two \(p\) orbitals remain unchanged. The two \(s p\) orbitals are oriented \(180^{\circ}\) apart on the right-left \((x)\) axis, while the \(p\) orbitals are perpendicular on the up-down \((y)\) axis and the in-out \((z)\) axis, as shown in Figure \(\PageIndex{1}\).

    The orientation of two s p orbitals at 180 degrees on the x-axis. The p orbital in the second image is perpendicular to the y and z axis.
    Figure \(\PageIndex{1}\): sp Hybridization. The two sp hybrid orbitals are oriented 180° away from each other, perpendicular to the two remaining p orbitals (red/blue).

    When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp–sp σ bond. At the same time, the pz orbitals from each carbon form a pz–pz \(\pi\) bond by sideways overlap, and the py orbitals overlap similarly to form a py–py \(\pi\) bond. The net effect is the sharing of six electrons and formation of a carbon–carbon triple bond. Each of the two remaining sp hybrid orbitals forms a σ bond with hydrogen to complete the acetylene molecule (Figure \(\PageIndex{2}\)).

    The formation of carbon-carbon triple bond from two sp-hybridized carbon atoms. The space-filling model, chemical structure, and ball and stick model of acetylene are shown.
    Figure \(\PageIndex{2}\): The structure of acetylene. The two carbon atoms are joined by one sp–sp σ bond and two p–p π bonds.

    As suggested by sp hybridization, acetylene is a linear molecule with H–C–C bond angles of 180°. The C–H bonds have a length of 106 pm and a strength of 558 kJ/mol (133 kcal/mol). The C–C bond length in acetylene is 120 pm, and its strength is about 965 kJ/mol (231 kcal/mol), making it the shortest and strongest of any carbon–carbon bond. A comparison of sp, sp2, and sp3 hybridization is given in Table \(\PageIndex{1}\).

    Table \(\PageIndex{1}\) Comparison of C−C and C−H Bonds in Methane, Ethane, Ethylene, and Acetylene
    Molecule Bond Bond strength Bond length (pm)
        (kJ/mol) (kcal/mol)
    Methane, \(\mathrm{CH}_4\) \(\left(s p^3\right) \mathrm{C}-\mathrm{H}\) 439 105 109
    Ethane, \(\mathrm{CH}_3 \mathrm{CH}_3\) \(\left(s p^3\right) C-C\left(s p^3\right)\) 377 90 153
      \(\left(s p^3\right) \mathrm{C}-\mathrm{H}\) 421 101 109
    Ethylene, \(\mathrm{H}_2 \mathrm{C}=\mathrm{CH}_2\) \(\left(s p^2\right) \mathrm{C}=\mathrm{C}\left(s p^2\right)\) 728 174 134
      \(\left(s p^2\right) \mathrm{C}-\mathrm{H}\) 464 111 109
    Acetylene, \(\mathrm{HC} \equiv \mathrm{CH}\) \((s p) \mathrm{C} \equiv \mathrm{C}(s p)\) 965 231 120
      \((s p) \mathrm{C}-\mathrm{H}\) 558 133 106
    Exercise \(\PageIndex{1}\)

    Draw a line-bond structure for propyne, CH3C≡CH. Indicate the hybridization of the orbitals on each carbon, and predict a value for each bond angle.

    Answer

    The CH3 carbon is sp3; the triple-bond carbons are sp; the C≡C−C and H−C≡C bond angles are approximately 180°.

    The wedge-dash structure of propyne.

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