MOLECULAR ORBITAL (MO) THEORY
MO theory only concerns with valence electrons because molecular bonding is solely driven by them (Lewis theory & Atomic Orbital theory also only concern with and valence e-)
Most quizzes and tests will test you on drawing molecular orbital diagrams.
So, right off the bat, here are the steps to accurately draw them:
1. Determine if the element has Z ≤ 7 or Z > 7 (Z = atomic number)
Use MO diagram on the left if Z ≤ 7, and on the right if Z > 7 (also be able to draw both of them)
it is only the "σ2pz" - underlined red - that switches places with the two "π" orbitals
2. Then fill in the atomic orbitals on either side with the valence electrons (e-).
For example, two carbon atoms in the C2 molecule would have 4 valence e- each, so draw 2e- for the 2s orbital and the other 2e- for the 2p orbital using Hund's rule.
Do this for both the left side and right side as there are two carbon atoms.
This is an MO diagram for C2. On either side is the atomic orbital for each of the two carbon atoms (2s and 2p). The middle section is the molecular orbitals.
3. After filling in the atomic orbitals for each of the atoms with e- (in this case, carbon atoms), you now fill in the molecular orbitals from lower energy level (bonding MO) first to higher energy (antibonding MO).
The total # of e- that fills the two molecular orbitals (σ2s & σ2s*) should match the total # of e- in the 2s orbitals, which is 4 e- in this case.
We have a total of 4e- from the 2s orbitals we can use to fill in the molecular orbitals in the bottom midsection (σ2s & σ2s*).
Fill each MO with 2e-, which totals 4e- as well.
4. Likewise, let's move up to the 2p orbitals and their corresponding molecular orbitals in the top midsection.
We also have a total of 4e- in the 2p orbitals, which we can use to fill up the "π2px" & "π2py" molecular orbitals. Now the two π orbitals have a total 4e- as well.
TWO IMPORTANT RULES FOR MO DIAGRAMS:
1. the number of atomic orbitals = the number of molecular orbitals
^see further explanation
2. the number of e- in atomic orbitals must match the number of e- in molecular orbitals.
^see further explanation
DIAMAGENTIC VS. PARAMAGNETIC
Easy definitions here.
Diamagnetic= all electrons are paired
Paramagnetic = at least one electron unpaired
When we determine diamagnetic vs paramagnetic, we are looking specifically at tbe molecular orbitals which are in the midsection.
^H2 molecule for instance is DIAMAGNETIC, because all of the existing electrons are fully paired.
^H2- molecule is PARAMAGNETIC, because we see one electron that is unpaired in the σ1s* orbital
BOND ORDER
Again, bonding MO does not have the aesterisk (*), whereas the antibonding MO does,
meaning σ1s is the bonding MO and σ1s* is the antibonding MO as shown below.
Bond order for the H2- molecule based on the formula provided:
(2 electrons - 1 electron) / 2 = 0.5 --> (2 e- from σ1s, and 1 e- from σ1s*)
Thus, "0.5" is the answer.
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