(Polar molecules, Non-polar molecules, etc.)
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pi bonds have 2 areas of overlapping e- density because the p-orbitals overlap side-by-side, so if the atoms in a pi bond were to rotate, then the bond would break. today in lecture, Dr. Lavelle drew the pi bonds like two figure 8s on top of each other, so that both lobes of each pi bond overlapped.
The way I think of it is that sigma bonds have electrons that can interact completely because they're directly overlapped, so all electrons are equally distributed and consequently can rotate. Pi bonds are more "stiff" because they can only interact peripherally since the electron repulsion coming from the sigma bond gets in the way.
Sigma bonds only have one field of electron sharing, so the atoms involved in a sigma bond can rotate independently of one another on that sigma-bond axis and still remain bonded. However, because pi bonds have two areas of electron sharing, no rotation is allowed: If an atom tried to rotate on one of the electron-sharing axises, the other one would be broken.
Sigma bonds can rotate because they are only connected at one point, making the connecting atoms able to move around. Pi bonds cannot rotate because they are connected at two points and the bonds would break if they tried to rotate. If you tried to twist it any way, at least one of the bonds would break.
if you want a more visual example of pi bonds think about Dr.Lavelle's example in lecture with the 2 chalk pieces between his two pointer fingers. When he moved his finger, one of the chalks fell. The chalks represent the pi bonds' regions of electron density, and rotating would break one of the two regions of electron density that make up a pi bond.
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