[Fireredburn1]

Whenever we talk about resonance we always associate it with overlap of p orbitals as well as delocalisation of electrons into a conjugate system.

We know that aryl halides, vinyl halides are unreactive to hydrolysis due to the partial double bond of the carbon-halide bond present inside. Many textbooks have explained in terms that the p orbital of the halides has overlapped with the p orbitals of carbon and thus giving the C-halide bond a partial double bond character.

Does that mean we expect the halides to be sp2 hybridised instead of sp3 since there must be an unhybridised p orbital to overlap with adjacent p orbital of carbon? Adding the fact that we know the nitrogen atom in pyrrole is sp2 hybridised instead of sp3 to favour delocalisation and therefore aromaticity, does that mean the halide atom to be sp2 hybridised to favour delocalisation too?



Likewise can be expect the nitrogen in amides, oxygen in carboxylate ion, halide in vinyl halides to be sp2 hybridised since all involves resonance and delocalisation of lone pairs into the system?

[orgopete]

I agree with the premise, but I don't agree with the generalization. The non-bonded electron can and do interact with neighboring pi electrons. However, they don't interact to the same extent for all atoms. The electrons of a nitrogen can interact much more strongly than those of a vinyl chloride. Enamines can react with a variety of electrophiles to give new carbon adducts. Enol ethers also interact, but not to the same extent. Hydrolysis of an enol ether occurs with protonation occurring on carbon. I don't know whether this is a thermodynamic result and possibly a faster O-protonation occurs, but is reversible. Since the electrons of a chlorine are held much more tightly (making HCl a strong acid), I would expect it would be more difficult to see them react, at least they will be less reactive.

[Fireredburn1]

Today my tutor told me that terminal atoms do not hybridise. This may explain why an unhybridised p orbital is able to overlap with adjacent p orbitals of the carbon to form a C-halogen bond with a partial double bond.

Is it true that terminal atoms do not hybridise?

[clarkstill]

That is categorically not true... The examples are extremely numerous (i.e. every molecule I can think of), but include:

MeCN - the terminal nitrogen is sp hybridized, with a linear orientation of the lone pair relative to the carbon

Formaldehyde - the oxygen is "terminal" but definitely sp2 hybridized, with lone pairs in the plane of the double bond

etc.

Did he maybe mean that terminal hydrogens don't hybridize?  Otherwise he's talking out of his... well, you know.

[phth]

Hybridization comes from the constructive/destructive wave interferences of the atomic wave functions s p_x p_y p_z (orbitals).  When add/subtract them they form a new function that can be graphed.  The result is the 4 sp³ orbitals.  Dont think of it as a sudden flip flop of shapes, it slowly changes as functional groups do.

How does this apply to conjugation with adjacent groups? It depends.  Adjacent electrons change the energies of each other depending on the strength of the effect.  We know pure sp² is aromatic because when aromatic the energy levels are degenerate (4n+2 electrons).  This effect is lessened, and the electrons will start to constructively interfere differently. The electrons will fluctuate between each hybridization spending more time in one compared to the other (average in between both).

This is why a primary amine has different pKa compared to a secondary amine or a tertiary amine, etc.