The great simplification
in this localized hybrid picture is that each C-H bond involves
only one hybrid orbital from carbon, and what the other three Sp
orbitals are doing is irrelevant. Bonds can be considered one at
a time, and it is not necessary to throw all five atoms into one
great mathematical pot. Consequently calculations of electron density
and energy are greatly simplified. The same Sp
hybridization can be used for ethane, HC-CH
and for a great many other carbon compounds. In ethane, shown at
the left, three of the four Sp3 hybrid atomic orbitals on each carbon
are combined with atomic orbitals of hydrogen, as in methane, and
the fourth is combined with one sp3 from the other carbon atom.
The sp3 hybrid orbitals extend out farther from the nucleus than
the is orbitals of hydrogen do, so a C-C bond is longer than a C-H
bond: 1.54 Å versus 1.09 Å. Bond angles throughout the
molecule still have tetrahedral values of 109.5.
The two ends of the molecule can rotate freely around the C-C bond,
but the most stable arrangement of hydrogen atoms by a small amount
of energy is that shown at the lower left. The hydrogen atoms are
"staggered" so that the hydrogen atoms on one carbon atom are as
far as possible from the hydrogens on the other carbon atom.