Lone pairs and hybridization

[Mdr1985]

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Simply stated...

It takes energy to hybridize, generally the 'payoff' is the ability to bond, which makes the molecule more stable...

So, why do line pairs hybridize? There's no payoff that I can see.

As an example, why is oxygen sp3 hybridized instead of simply sp with two non hybridized p orbitals?

Thanks!

 

[SN2ed]

Thread moved. These types of questions belong in the Study Q&A forum.

 

[LingoLaine]

It takes energy to hybridize

I'm not sure this statement is true. Where'd you get that info?

 

[Mdr1985]

My Ochem professor...

 

[ljc]

Don't you mean O2 is sp2 hybridized, not sp3? Or were you talking about something else? (2 lone pairs, 1 double bond)

Anyways, as far as I understand it, the payoff comes from geometric realignment. In its unhybridized state, oxygen has a filled 2s orbital, 1 filled 2p orbital, and 2 half filled 2p orbitals. One of these 2p orbitals could form a sigma bond, the other a pi bond, and the last be the lone pair. This however would result in orbitals that are at 90 degree angles from each other (along X, Y, and Z axes). By hybridizing, you form 120 degree angles, lowering the overall energy of the molecule. Sure you have to promote the 2s orbital, which costs energy, but you're lowering the energy of 3 2p orbitals as a result, which must result in a net gain or it would not happen.

Someone correct me if I'm talking out of my ass, but the way I learned hybridization was that it was meant to optimize 3-dimensional shape as a means of minimizing energy.

 

[jsp132]

Don't you mean O2 is sp2 hybridized, not sp3? Or were you talking about something else? (2 lone pairs, 1 double bond)

Anyways, as far as I understand it, the payoff comes from geometric realignment. In its unhybridized state, oxygen has a filled 2s orbital, 1 filled 2p orbital, and 2 half filled 2p orbitals. One of these 2p orbitals could form a sigma bond, the other a pi bond, and the last be the lone pair. This however would result in orbitals that are at 90 degree angles from each other (along X, Y, and Z axes). By hybridizing, you form 120 degree angles, lowering the overall energy of the molecule. Sure you have to promote the 2s orbital, which costs energy, but you're lowering the energy of 3 2p orbitals as a result, which must result in a net gain or it would not happen.

Someone correct me if I'm talking out of my ass, but the way I learned hybridization was that it was meant to optimize 3-dimensional shape as a means of minimizing energy.

yup how i learned it also due to vsepr