Hybridization in the Oxygen atom (you will have to imagine the points on the electron arrows!) 2p || | | hybridizes to 2sp3 || || | | 2s || 1s || 1s || Each hybrid orbital contains 1 or 2 electrons. If the atom has fewer than 5 valence electrons, the second electron in the "s" atomic orbital must be promoted into one of the other atomic orbitals at the same time as the orbitals are being hybridized. Hybridization in the Carbon atom 2p | | hybridizes to 2sp3 | | | | 2s || 1s || 1s || Hybridization in the Boron atom 2p | 2p __ hybridizes to 2sp3 | | 2s || 1s || 1s || There is one unoccupied, unhybridized p orbital (actually, the orbital does not exist since it has no electron). Hybridization in the Beryllium atom 2p 2p __ __ hybridizes to 2sp3 | | 2s || 1s || 1s ||
What is interesting is that the hybridized valence orbitals yield exactly the same number of paired
and unpaired as the Lewis dot diagram!
Take a look at page 234 in the text book. In these diagrams, the smaller side of the hybrid atomic
orbitals are left out of the diagram for sake of clarity.
If the atom has...
2 valence electrons, it will hybridize an "s" and a "p" orbital to make 2 sp hybrid orbitals.
180o apart
3 valence electrons, it will hybridize 1 "s" and 2 "p" orbitals to make 3 sp2 hybrid orbitals.
120o apart
4 valence electrons, it will hybridize 1 "s" and 3 "p" orbitals to make 4 sp3 hybrid orbitals.
109.5o apart
The hybrid orbitals are still atomic orbitals, not bonding orbitals. To make a bond, they still must
overlap with the atomic orbital of another atom. However, the hybridization process occurs concurrently
with the bonding process. Atoms do not hybridize their orbitals until they make covalent bonds. As well,
atoms that are making only one bond do not hybridize, at least, we have no evidence that they do.
The hybridization process must allow the atom to lower its potential energy as the bonding process proceeds
by allowing maximum angular separation of bonds. When atoms do not need to do this, they don't.
How real are hybrid orbitals?
They seem awfully convenient. My first-year text book answers this question by stating that hybrid
atomic orbitals are no less real than non-hybridized atomic orbitals. You can take that any way you
want!
Multiple Bonds
There are two types of bonding orbital. Read pages 236 and 239 in the text book - carefully!
1) Bonding orbitals made from and end-to-end overlap of atomic orbitals: sigma bonds.
These can occur with s orbitals, p orbitals, and hybrid orbitals.
Both atomic orbitals must be oriented parallel to the axis of the bond.
The shape of the orbital depends on the types of atomic orbitals involved
(see the "Atomic Orbitals, Bonding Orbitals" diagram above).
2) Bonding orbitals made form a side-by-side overlap of atomic orbitals: pi bonds
The must occur with 2 non-hybridized p orbitals
The orbitals must be oriented perpendicular to the axis of the bond.
The shape of a pi bond is that of two regions of electron density, both parallel to
the axis of the bond, but on either side of the bond.
Since the electron density is not directly in between the two nuclei, pi bonds are not
as good as sigma bonds at holding atoms together.
Since pi bonds require non-hybridized p orbitals, the atoms remaining atomic orbitals will
be in a less hybridized state, so their bonds will be in a lower molecular shape family.
e.g., oxygen with 1 pi bond and 3 sigma bonds will have 3 2sp2 orbitals and 1 2p
orbitals. Thus, its 3 hybridized orbitals will be in a trigonal planar orientation,
rather than tetrahedral.
Non-hybridized "p" orbitals remain in their original orientation.
Molecular Shape Family
Since non-hybridized "p" orbitals remain in their original orientation, the molecular shape family of
a molecule is determined only by the number of sigma bonds and lone pair electrons around the central atom.
C2H6 has only single bonds, so the Carbons are 2sp3 hybridized and tetrahetral.
Double Bonds: page 237
Double bonds are formed by 1 sigma bond and 1 pi bond.
Since the two non-hybridized "p" orbitals must remain parallel to each other in the pi bond, the
molecule cannot rotate around the double bond the way in which a single (sigma) bond can rotate.
C2H4 has a double bond, so the Carbons are 2sp2 hybridized and planar triangular.(fig. 13, pg 237)
Triple Bond: page 239
Triple bonds are formed by 1 sigma bond and 2 pi bonds. The two pi bonds are formed from two sets
of non-hybridized "p" orbitals oriented at right angles to each other. Thus, the two pi bonds are
oriented such that one is above and below the bond axis while the other is on both sides of the
bond axis.
C2H2 has a triple bond, so the Carbons are 2sp hybridized and linear. (fig. 16, pg 239)
Lone Pairs
In atoms where hybridization occurs, lone pairs always are hybridized orbitals in which there were
two electrons before bonding. If the atom does not hybridize, lone pairs are just regular atomic
orbitals.
Do Atoms Always Hybridize When Bonding?
The simple answer is no, since molecules like H2S have bond angles similar to regular atomic orbitals
(92o in H2S). For some reason, perhaps the Sulfur is so large and the Hydrogens so small that the
two bonds do not influence each other significantly, the Sulfur in H2S does not hybridize its
atomic orbitals when making this molecule.