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Organic Chemistry

1.9 sp Hybrid Orbitals and the Structure of Acetylene

Organic Chemistry1.9 sp Hybrid Orbitals and the Structure of Acetylene

1.9 • sp Hybrid Orbitals and the Structure of Acetylene

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, H−C≡C−HH−C≡C−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 2s orbital hybridizes with only a single p orbital. Two sp hybrid orbitals result, and two p orbitals remain unchanged. The two sp orbitals are oriented 180° 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 1.16.

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 1.16 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 spsp σ bond. At the same time, the pz orbitals from each carbon form a pzpz π bond by sideways overlap, and the py orbitals overlap similarly to form a pypy π 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 1.17).

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 1.17 The structure of acetylene. The two carbon atoms are joined by one spsp σ bond and two pp 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 1.2.

Table 1.2 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, CH4 (sp3) C−H 439 105 109
Ethane, CH3CH3 (sp3) C−C (sp3) 377  90 153
(sp3) C−H 421 101 109
Ethylene, H2C=CH2 (sp2) C=CC=C (sp2) 728 174 134
(sp2) C−H 464 111 109
Acetylene, HC≡CHHC≡CH (sp) C≡CC≡C (sp) 965 231 120
(sp) C−H 558 133 106
Problem 1-13
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.
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