Download The basic trigonal bipyramidal molecular

The basic trigonal bipyramidal molecular
geometry. There are bonds from the central atom
to five atoms. The five electron pairs assume an
arrangement that will minimize the repulsion
between electrons. You can visualize two threesided pyramids, except one is up-side-down, with
their bases together.
The two vertical bonds are in the axial position,
while the remaining three bonds are in the
equatorial position. This arrangement results in the
lowest possible energy.
In a “see-saw” molecular geometry, there is one
lone pair in the equatorial position. It is 120
degrees away from the other two bonding pairs of
electrons in the equatorial positions, and therefore
lower in energy than placing the lone pair in an
axial position.
Mike Jones – Pisgah High School – Canton NC
In a T-shaped molecule, the lone pairs are directed
towards the “equator”. This arrangement places
each lone pair at right angles to the axial bonding
pairs, and 120 degrees away from the other lone pair
in the equatorial position, and 120 degrees away
from the bonding pair in the equatorial position.
This results in lower overall energy.
If the lone pairs were in the axial positions, the molecule
would be trigonal planar. But the lone pairs would be 90
degrees from all three of the bonding pairs. This would
result in greater repulsion and a system at a higher energy
than a T-shaped molecule. Therefore, the T-shaped
molecule is preferable.
In a linear molecule, all three lone pairs are directed
towards the “equator”. This arrangement places
each lone pair at right angles to the axial bonding
pairs, and 120 degrees away from the other lone
pairs in the equatorial position. This results in lower
overall energy.
Mike Jones – Pisgah High School – Canton NC
Review exercises. Answer on notebook paper. Make your answers complete and to the point.
1. Explain why the bond angles in the trigonal bipyramidal electron pair geometry are 90 degrees and
120 degrees.
2. Explain the difference between the terms “axial” and “equatorial” as it applies to molecular geometry.
3. A student suggested that PCl5 could have bond angles of 72 degrees. Describe the arrangement of the
electron pairs in the molecule that the student was imagining? How would you explain to that student
why the bond angles are not 72 degrees?
4. What allows phosphorous to have five bonds, anyway? After all, there are only three unpaired
electrons in the 3p-sublevel. (Hint: Draw the electron energy diagram of phosphorous.)
5. Assuming that phosphorous can have five bonds, explain why nitrogen, which is in the same family
on the periodic table, cannot have five bonds.
6. Consider BeCl2, SCl2 and XeCl2. Determine which one(s) are linear and explain why.
7. Explain why ICl3 is T-shaped and not trigonal planar.
8. Explain why in a T-shaped molecule the two electrons pairs in the axial positions are not
exactly 180 degrees apart, while both are slightly less than 90 degrees from the one
equatorial bonding pair.
9. Draw the Lewis diagrams for CCl4, SCl4 and XeCl4 and give the molecular geometry of each.
Explain what is different about each of the molecules that determines the molecular geometry.
10. Consider the behavior of lone pairs of electrons as you explain why the ICl4- ion has a molecular
geometry of square planar and not see-saw.