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So far the orbital overlaps we studied were all linearly shaped.
But what about other shapes, like planar triangular and tetrahedral?
There are no orbitals directed at 120° and 109.5° angles.
To explain orientations like that, we must study the way atomic orbitals of the same atom interact with each other.
Lets take an example--beryllium hydride (BeH2).
The orbital diagram for the valence shell of beryllium is:
Experimental observation shows that this molecule is linear, therefore the Be-H bonds must point in opposite directions.
For this to happen, each of the beryllium orbitals must contain only one electron.
To accomplish this, beryllium unpairs the electrons, putting one of them in the p orbital, like so:
Because of this switch, the 2s and the 2p orbitals overlap, forming 2 sp hybrid orbitals where the hydrogens can bond.
Other hybrids can form by combining an s orbital with two or three p orbitals.
When two p orbitals are mixed with an s orbital, a set of three sp2 hybrid orbitals are formed, all extending 120° away from each other (planar triangular).
When three p orbitals are mixed with an s orbital, a set of four sp3 hybrid orbitals are formed, all extending 109.5° away from each other (tetrahedral).
When the central atom has more than an octet hybrid orbitals can involve d orbitals.
The most common orbitals involving d orbitals are sp3d and sp3d2.
They extend out to form trigonal bipyramidal and octahedral structures, respectively.