[kekie]

So, I heard that the lone pairs on nitrogens in thymine can delocalize/partisipate in resonance, while the lone pairs on the nitrogen of, for instance, pyridine, can't.

For reference,
Pyridine;

Thymine;

 

I would guess it's because the electron config of the nitrogen on the pyridine is like so;



And that the electron config of the nitrogens on the thymine are like so;


Because each of the nitrogens are bonded to three other things, so it has to have three unpaired electrons at the top.

(Lone pair in red, electrons participating in sigma bonds in blue)

In the second config (Thymine) the lone pair is in a p-orbital, letting it participate in resonance.

Am I correct?

[kriggy]

I think the difference is that in pyridine, when you would add the electron pair into the system, you would disturb the aromaticity, however in thymine by addit the 2 pairs the system becomes aromatic.
edit: and btw why do you think the orbitals are like this? I mean, there is 1 s orbital and 3 p orbitals in nitrogen valence shell. So your schemes are not correct. The hybrid SP2 orbitals are made of combination on 1 s and 2p orbitals so if you create 3 SP2 orbitals then you have only 1 p orbital left. But you have 3 p orbitals adn 3 sp2 orbitals and 1 s orbital in valence shell.

[Nescafe]

So, I heard that the lone pairs on nitrogens in thymine can delocalize/partisipate in resonance, while the lone pairs on the nitrogen of, for instance, pyridine, can't.

For reference,
Pyridine;

Thymine;

 

I would guess it's because the electron config of the nitrogen on the pyridine is like so;



And that the electron config of the nitrogens on the thymine are like so;


Because each of the nitrogens are bonded to three other things, so it has to have three unpaired electrons at the top.

(Lone pair in red, electrons participating in sigma bonds in blue)

In the second config (Thymine) the lone pair is in a p-orbital, letting it participate in resonance.

Am I correct?

I thought this had to do with the fact that the electron pair on the nitrogen of pyridine is in a pi orbital which is perpendicular to the orbitals of the ring and that is why, due to the geometry it does not delocalize within the ring.

[kekie]

That might be the case- I don't know.

You tell me - the above idea was just my guess.

Also, you don't need to quote the entire post like that. 

[Dan]

I thought this had to do with the fact that the electron pair on the nitrogen of pyridine is in a pi orbital which is perpendicular to the orbitals of the ring and that is why, due to the geometry it does not delocalize within the ring.

This is correct. The N lone pair in pyridine cannot overlap with the π system - it's geometrically impossible.

[goops]

In pyridene, the  N is sp2 hybridised. If it tries to delocalize its e-s to form a pi bond, it gets sp hybridised. =N= tends to get linear due to sp hybridisation, this will distort the geometry and creates stress in the compound which further leads to unstability. So it prefers not delocalising its e-s.
Hope it would help.. 

[Sam_]

Here's your problem:

(Lone pair in red, electrons participating in sigma bonds in blue)

Single atom orbitals like sp² and sp³ are no longer very useful for analysis once you start considering molecular orbitals.

Rather than trying to assign the nitrogen's electrons to nitrogen atomic orbitals, consider the molecular orbitals of the pyridine system that nitrogen can participate in.  The pyridine nitrogen is sp² hybridized, so it contributes 2 sp² orbitals to form a σ bond with each of the neighboring carbon atoms, one sp² orbital is left for the lone pair, and it contributes its unhybridized p orbital to the aromatic π system.  The nitrogen contributes one valence electron to each of the σ bonding orbitals, one to the π system, and two to the nonbonding lone pair orbital.  The core electrons remain in the 1s orbital.

Even in the case of the nonbonding lone pair orbital, though, if you calculated the energies of the various orbitals, you would find some contributions from other atomic orbitals.

[kekie]

Thanks!