Why Do Double Bonds Not Allow Bond Rotation?

[Cheman]

Why in organic chemistry do double bonds not allow rotation of bonds, whilst single bonds do?

Thanks.

 

[Sirus]

Picture bonds in three dimensions to explain things like this. A single bond will be created by the area of intersection between the axial s-orbitals, which point toward each other. A double bond, on the other hand, is created by the overlap of parallel vertically-oriented p-orbitals. If you rotate the atoms relative to each other, the s-orbital continue to overlap as before, while the p-orbitals take on what you could call a staggered position (when viewed from the side), so overlap does not exist.

 

[Cheman]

Ok, so they rotate polarised light - my questions now are:

a) Why? How do they cause it to rotate?

b) Why do they turn it the opposite ways?

Thanks.

 

[Cheman]

Sorry - put that on the wrong one of my posts!

The question on THIS post should be "So, what happens if you try to rotate a double bond? Does it resist rotation or is it still possible?"

Thanks.

 

[movies]

Rotating the double bond would eliminate the overlap between the p-orbitals, as Sirus described. You would need enough energy to break the C-C pi bond, which isn't common at normal temperatures. It is possible, however, at high temperatures.

 

[Cheman]

Why is it though that the p orbital electrons are the ones which become delocalised and form the pi bonds after the other electrons in the outer shell have formed sigma bonds with 2 hydrogens and the other carbon? After, surely we would expect the electron with the highest energy to react first ie the p orbital - not the s orbitals?

Thanks.

 

[movies]

The electrons are trying to adopt the lowest possible energetic configuration though, right? S orbitals are lower energy than p orbitals, so the lower energy molecular orbitals will have a larger contribution from the s atomic orbitals. In the case of an alkene the last p orbital is what is "left over."