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.