CHEMICAL BONDING II. MOLECULAR ORBITAL THEORY. DO NOW. Pick up handout. Get out homework handout. HOMEWORK ANSWERS #1. a . BrF 5 VE = 42; 6 e- pairs: 5 bonded pairs and 1 lone pair; shape: square pyramid b . C 2 H 2 VE = 10; 2 e- pairs: 2 bonded pairs; shape: linear
Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.
MOLECULAR ORBITAL THEORY
shape: square pyramid
shape: trigonal planar
Problems with LE Model
1. Electrons are not always localized as in the VSEPR theory; therefore resonance must be added and explained as best possible.
2. Molecules containing unpaired electrons are not easily dealt with using the localized model.
3. Magnetism is easily described for molecules using the MO theory. (Oxygen is paramagnetic which is unexplained by the localized electron model.)
4. Bond energies are not easily related using the localized model.
Be sure you know these terms. It will be tough to fill out the diagrams without them.
An atomic orbital hold an ATOM’s electrons
A molecular orbital is made from atomic orbitals in a MOLECULE.
The two hydrogen electrons
- Each H entered on the outside with its lone 1s electron.
- As they approach each other, their two atomic orbitals blend to form two molecular orbitals.
- One MO is of high energy and one MO is of low energy.
- The electrons will choose the LOW energy route.
- The electrons occupy the lower energy level and thus a bond is formed.
Lower energy bonding orbital
The two elements involved in the bonding is this example are “A” and “B”.
Diatomic Helium, He2
One helium atom’s two electrons enter on the left,
the other helium’s two electrons on the right.
4. Electrons lower their energies by being attracted by both nuclei.
5. Labels on the molecular orbitals indicate their symmetry (shape), the parent atomic orbital, and whether they are bonding or antibonding.
6. Molecular electron configurations can be written just like atomic electron configurations. H2 = σ1s2.
7. Each molecular orbital can hold two electrons with opposite spins.
8. Orbitals are conserved. The number of molecular orbitals is always the same and the atomic orbitals used to make them.
1. Write a capital “F” on the bottom left and one on the bottom right for the fluorine atoms.
4. Fill in the bonding 1s bonding (σ)orbital with two electrons and the anti-bonding (σ*)orbital with two electrons.
7. Start with the blue pi bonds first – fill them with two electrons each. (total 4 electrons)