The common hybridization in alkynes is sp. Alkynes are compound that have triple bond between it's carbons.
Acetylene is linear, with a carbon–carbon bond distance of 120 pm and carbon–hydrogen bond distances of 106 pm.
Linear geometries characterize the H-C≡C-C and C-C≡C-C units of terminal and internal triple bonds, respectively as well. This linear geometry is responsible for the relatively small number of known cycloalkynes. Figure below shows a molecular model for cyclononyne in which the bending of the C-C≡C-C unit is clearly evident. Angle strain destabilizes cycloalkynes to the extent that cyclononyne is the smallest one that is stable enough to be stored for long periods. The next smaller one, cyclooctyne, has been isolated, but is relatively reactive and polymerizes on standing. In spite of the fact that few cycloalkynes occur naturally, they gained recent attention when it was discovered that some of them hold promise as anticancer drugs.
All of these trends can be accommodated by the orbital hybridization model. The bond angles are characteristic for the sp3, sp2, and sp hybridization states of carbon and don’t require additional comment. The bond distances, bond strengths, and acidities are related to the s character in the orbitals used for bonding. s Character is a simple concept, being nothing more than the percentage of the hybrid orbital contributed by an s orbital. Thus, an sp3 orbital has one quarter s character and three quarters p, an sp2 orbital has one third s and two thirds p, and an sp orbital one half s and one half p. We then use this information to analyze how various qualities of the hybrid orbital reflect those of its s and p contributors. Take C-H bond distance and bond strength, for example. Recalling that an electron in a 2s orbital is, on average, closer to the nucleus and more strongly held than an electron in a 2p orbital, it follows that an electron in an orbital with more s character will be closer to the nucleus and more strongly held than an electron in an orbital with less s character. Thus, when an sp orbital of carbon overlaps with a hydrogen 1s orbital
to give a C-H bond, the electrons are held more strongly and the bond is stronger and shorter than electrons in a bond between hydrogen and sp2-hybridized carbon. Similar reasoning holds for the shorter C-C bond distance of acetylene compared to ethylene, although here the additional π bond in acetylene is also a factor.
The pattern is repeated in higher alkynes as shown when comparing propyne and propene. The bonds to the sp-hybridized carbons of propyne are shorter than the corresponding bonds to the sp2 hybridized carbons of propene.
Acetylene is linear, with a carbon–carbon bond distance of 120 pm and carbon–hydrogen bond distances of 106 pm.
Linear geometries characterize the H-C≡C-C and C-C≡C-C units of terminal and internal triple bonds, respectively as well. This linear geometry is responsible for the relatively small number of known cycloalkynes. Figure below shows a molecular model for cyclononyne in which the bending of the C-C≡C-C unit is clearly evident. Angle strain destabilizes cycloalkynes to the extent that cyclononyne is the smallest one that is stable enough to be stored for long periods. The next smaller one, cyclooctyne, has been isolated, but is relatively reactive and polymerizes on standing. In spite of the fact that few cycloalkynes occur naturally, they gained recent attention when it was discovered that some of them hold promise as anticancer drugs.
All of these trends can be accommodated by the orbital hybridization model. The bond angles are characteristic for the sp3, sp2, and sp hybridization states of carbon and don’t require additional comment. The bond distances, bond strengths, and acidities are related to the s character in the orbitals used for bonding. s Character is a simple concept, being nothing more than the percentage of the hybrid orbital contributed by an s orbital. Thus, an sp3 orbital has one quarter s character and three quarters p, an sp2 orbital has one third s and two thirds p, and an sp orbital one half s and one half p. We then use this information to analyze how various qualities of the hybrid orbital reflect those of its s and p contributors. Take C-H bond distance and bond strength, for example. Recalling that an electron in a 2s orbital is, on average, closer to the nucleus and more strongly held than an electron in a 2p orbital, it follows that an electron in an orbital with more s character will be closer to the nucleus and more strongly held than an electron in an orbital with less s character. Thus, when an sp orbital of carbon overlaps with a hydrogen 1s orbital
to give a C-H bond, the electrons are held more strongly and the bond is stronger and shorter than electrons in a bond between hydrogen and sp2-hybridized carbon. Similar reasoning holds for the shorter C-C bond distance of acetylene compared to ethylene, although here the additional π bond in acetylene is also a factor.
The pattern is repeated in higher alkynes as shown when comparing propyne and propene. The bonds to the sp-hybridized carbons of propyne are shorter than the corresponding bonds to the sp2 hybridized carbons of propene.
