(Polar molecules, Non-polar molecules, etc.)

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Postby Brian_Wu_3B » Thu Dec 10, 2020 1:19 pm

Why is it that with 5 regions of election density and two lone pairs, the shape is t shaped instead of trigonal planar. I'm imagining a molecule with lone pairs on the top and bottom, and three atoms bounded to the central atom on its equator (is that the right word?). Instead, in t shaped, the lone pairs are next to each other on the equator and the rest of the bounded atoms form a tshape. The molecular shape I imagined trigonal planar. Is this incorrect?

Stuti Pradhan 2J
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Re: AX3E2

Postby Stuti Pradhan 2J » Thu Dec 10, 2020 7:02 pm

T-shaped does not have lone pairs in the axial position (top and bottom) but instead, has them on the equatorial position. This is because when there is only 1 lone pair, making the atom on the top a lone pair would cause the more repulsive lone pair to interact with three atoms (all the atoms in the equatorial plane). Since you want the least amount of interaction with the lone pair, you would replace one of the atoms on the equatorial plane instead, so it would only interact with the 2 atoms on the axial plane at 90 degrees. Again, this is because you want to create the shape with the least repulsion because that makes it more stable.

For the second lone pair, making any of the axial atoms from the seesaw shape a lone pair will result in repulsions from the first lone pair on the axial atoms and from the second pair on the equatorial atoms, so all the atoms will experience a significant amount of repulsion. Instead, replacing another equatorial atom with a lone pair will only affect the axial atoms which are already being affected by repulsions from the first lone pair, so this minimizes overall repulsions. Therefore, the molecule would be T-shaped, not trigonal planar.

Hope this helps!

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Re: AX3E2

Postby HilaGelfer_2H » Thu Dec 10, 2020 7:10 pm


For AX3E2 the central atom has five regions of electron density. Therefore, its electron arrangement must be trigonal bipyramidal. When choosing the lone pairs you have to choose lone pairs that minimize electron pair repulsions. By placing our lone pairs in the equatorial position, it minimizes the electron pair repulsions along the whole atom and the lone pairs push the bonding pairs closer together resulting in a T-shape.

I hope this helps :)

Tobie Jessup 2E
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Re: AX3E2

Postby Tobie Jessup 2E » Thu Dec 10, 2020 7:28 pm

I was also confused on this shape, but there was a good explanation in the textbook that helped me out!

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Re: AX3E2

Postby EmilyGillen_1A » Thu Dec 10, 2020 8:02 pm

It took me a second to really understand why the shape it T instead of trigonal planar, but it comes down to the fact that e- repel against each other at all times. The repulsion between electrons means the shapes that form are the shapes that allow electrons to be the farthest away from each other as possible.
With 5 regions of electron density, the shape would be trigonal bipyramidal, with 2 e- on the "top" and "bottom" and 3 electron in the axis in between. Now imagine you have to take away 2 e- areas (to form the shape that accounts for only bonded pairs). Ask yourself, which configuration places the atoms farthest apart - removing the top and bottom e-'s (leaving 3 e- on the same plane) or removing 2 e- from the same plane and leaving the "top" and "bottom" as-is (making a T shape)?

If you can visualize that, you see that leaving 3 e- on the same plane (trigonal planar) pushes them closer than they would be if you took out some e- from the same plane (T shape). This means a T shape is more favorable because it spaces out the e- onto different planes and puts them farther apart than a trigonal planar shape would.

I hope that makes sense!

Yichen Fan 3A
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Re: AX3E2

Postby Yichen Fan 3A » Thu Dec 10, 2020 8:36 pm

I still don't quite understand why the molecule will form T-shape instead of trigonal planar. I know that if there is one lone pair it will form seesaw because the electrons don't want to be at axial forming three 90 degree bonds with atoms, but when there are two lone pairs won't they try to stay far away from each other like square planar?

Chinmayi Mutyala 3H
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Re: AX3E2

Postby Chinmayi Mutyala 3H » Thu Dec 10, 2020 9:54 pm

This is what I got from the lecture. If the lone pairs are on the top and bottom, they'll be 90 degrees away from all the other bonds which is not the best for them to be as far away as possible. You want them to be as far away are possible so the lone pairs are kind angled in the axis in three dimensions instead of directly facing up or down while the bonds are in a T-shape on the plane. I'm not sure if that was a clear explanation but just googling the shape helped me understand it better.

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Re: AX3E2

Postby Brian_Wu_3B » Fri Dec 11, 2020 9:34 am

Thank you! It makes a lot of sense now.

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