Hybrid Orbitals

Part 1. sp³ Hybridization of Carbon

The simplest way to illustrate hybridization of atomic orbitals is to look at the case of methane, CH₄.Obviously, we need to make four bonds using the atomic valence orbitals of hydrogen and those of carbon. Two ways of representing the molecule are shown below. Note especially that the molecule is symmetrical, i.e. all four bonds are of equal length and all bond angles are equal to the perfect tetrahedral angle (109.5^o).

The electrons of the carbon atom are shown below. Note that we are only interested in the valence electrons (2s² 2p²).

 

It is apparent that the valence electrons are not equivalent to one another, and a non-symmetrical methane molecule would result if we used these electrons as-is. Thus, one of the 2s electrons is "promoted" to the 2p level, resulting in four half-filled atomic orbitals:

 

We could (theoretically) make four bonds using the four un-paired elecrons of the carbon plus the four electrons from the hydrogen, but the result would be a non-symmetrical methane molecule, since the four carbon electrons are not identical. Instead, we (conceptually) combine the s orbital with the three p orbitals to make four equivalent "sp³" hybrid orbitals. These new hybrid orbitals each will contain one electron, and each will have some s character and some p character. Shown below is one such sp3 hybrid orbital. Use the mouse to rotate it.

Of course, we need four such orbitals to make a methane molecule. They are arranged around the carbon atom as shown below. Use the mouse to rotate it and have a good look at it.

Finally, the electron from each hydrogen pairs up with the electron in one of the sp³ hybrids, forming a (two electron) single bond. The methane molecule thus looks like the following. Note the hydrogen atoms (white) around the carbon atom (grey) and note, by rotating the molecule, that all H-C-H bond angles are the same.

Cool, eh?

On to part 2, sp² hybridization

Chicken out