order of filling in hybridized orbitals.?

Question

In an sp hibrydized orbial we have 2 hybridized and 2 unhybridized orbitals. What is the order of filling? Do we start filling the hybridized orbitals first or the unhybridized?
E.g. if an atom has 5 valence elctrons, we put 1 on each orbital. but where does the 5th one go?

Answer 1

Keep in mind one thing: hybridization is not "real," it's just one pretty not-right way we use to visualize molecular bonding.

As such, we don't use hybridization until atoms are bonded to one another, so filling hybrid orbitals in one atom at a time is not how it's done.

I want to wax about my understanding of theories of bonding for a moment. Skip to the last full paragraph for my answer to your question.

There is nothing in M.O. theory that energetically needs orbitals to "hybridize", either when atoms are bonded to other atoms or not. Atomic orbitals do combine into molecular orbitals, but that is not a "hybridization," it's an overlap. M.O. theory doesn't contain any hybrids.

Valence bond theory and M.O. theory are two different theories. Try not to confuse the two. VB says nothing at all about orbital energies; it only deals with geometry and electron localization. The hybridization of orbitals is simply a trick to use when drawing or calculating structures that happens to work, and that's all. We don't draw hybridized orbitals because of energy considerations, but rather because hybridization helps visualize geometry. In some calculations, VB theory actually seems to describe molecules better than MO theory. Neither theory is "right" all the time, and sometimes both are wrong.

Discussion over.

That being said, convention (definition) says that when the central atom in a compound is hybridized, the sigma bonds are filled first, with two electrons each. Any leftover electrons are put into pi-bonding and nonbonding orbitals in whatever way optimizes formal charges on the various atoms.

These are the rules for filling out Lewis structures too.

Make sense?

Answer 2

You only use hybrids in molecules, not isolated atoms. That is because in the isolated atoms, s and p are at different energies and there is nothing to make them mix. If you are looking at hybrids used in bonding, you need to count the electrons from both ends of the bond.

For example, in ammonia you have sp3 hybridized nitrogen, and you have eight electrons altogether. In each of the bonds, you have two electrons, and the two left over give you your lone pair on nitrogen.

If you like, you can think of one electron in an N-H bond as coming from hydrogen and the other one from nitrogen, but since all electrons are identical all that matters is the total number.

Answer 3

You need to look at the total pairs (add the lone and bonding).

total pairs=4 you have sp3 hybridization
total pairs=3 you have sp2 hybridization
total pairs=2 you have sp hybridization

You fill in the orbitals the same way you do atomic orbitals, using all the same rules (Aufbau, Pauli and Hund).

Notes:
*you only hybridize the central atom
*draw the lewis structure first and determin your pairs.
*think pairs not valence electrons
*Anytime you have four pairs ;you have sp3 hybrid