Nitrogen's intermolecular bonding

[sludger13]

Hi again, I have a couple of questions:
1) Nitrogen (N₂) exists liquid and solid. The molecules bond with van der Waals forces. What's the difference between liquid and solid state? Just quantity of van der Waals forces?

2) When is nitrogen ready for freezing? When atoms have such a heat kinetic energy, that they bring near enough and the induced electric field is strong enough to induce dipoles in another molecules? Or where is the breaking point?

3) Dipole's interaction is attractive. What's the repulsive interactions, in order atoms don't collapse?

4) One more question here (LINK). I hope I'm not so much annoying with that. :shy:
Thanks for every single advice or remark.

 

[Borek]

Why do you limit your question to nitrogen? The answer would be mostly identical for every other substance that can exist as a gas (by which I mean it doesn't decompose before becoming gaseous, think glucose for example).

 

[sludger13]

Of course, I don't understand the behaviour of many other substances with similar properties.
Some substances have obvious freezing transition, as covalent bonds are creating. Glucose doesn't belongs there, with its molecular (or ionic) crystal structure (I don't know, maybe glucose can't even get liquid, due to pyrolysis).

 

[Borek]

There are no covalent bonds created during freezing.

 

[sludger13]

I thought maybe some inorganic substances are covalent bonded in solid state (e.g. graphite) and some of those covalent bonds disappear as the atoms are sufficiently distant (the orbitals are no longer overlapping).

 

[DrDu]

The van der Waals equation describes the liquid-gas phase transition of gasses like nitrogen qualitatively correctly:
http://en.wikipedia.org/wiki/Van_der_Waals_equation
It assumes attractive induced dipole-dipole interactions ~1/V^2 and a repulsive hard sphere repulsion at short distances (mainly due to Pauli exclusion principle for the electrons
http://en.wikipedia.org/wiki/Pauli_exclusion
).