Orbital Hybridisation

[ed_saifa]

Hi,
Can anyone please explain this concept to me?

Thanks heaps

[Toothpaste]

Bump. I need explanations on this same thing too. Yay if you help.

All I know is hybrid orbital = an orbital produced from the combination of two or more atomic orbitals.
Sigma bond? What.
orbital? WHAT?

, , hybrid orbitals = for elements in the second period?
WHAT'S THEIR PURPOSE!?

I'm missing the basics.

And what makes a resonance structure more 'stable'? Charges, bonds, anything else?

Thanks.

[bucket]

I'm pretty confused about these too...
Haha
I'm pretty sure a orbital is symmetrical and a orbital is not.

[bucket]

I just found this
http://studio4learning.tv/sciences/chemistry/chemical-bonds/orbital-overlap/

[Collin Li]

The reason why we require hybridisation is because the simple model of s and p orbitals is insufficient to explain molecular orbitals.

Take CH4 for example. It is tetrahedal, all C-H bonds equally spaced apart. This cannot be possible with 1 s orbital and 3 p orbitals (from the central carbon atom), since the energy levels are different in s and p orbitals. We need to think of an orbital that is 1/4 s, and 3/4 p, i.e: an sp3 orbital. There are 4 of these, and they are basically a 'hybrid' of s and p orbitals (they look like a p orbital with a spherical bulge). Furthermore, they have equivalent energy states, which is consistent with the observation that CH4 is a tetrahedral molecule with C-H bonds equally spaced apart.

[Collin Li]

And what makes a resonance structure more 'stable'? Charges, bonds, anything else?

It's because the formal representation of the molecule is not the reality. Molecules where multiple resonance structures can be drawn out are actually a hybrid of all these resonance states. For example, benzene does not actually alternate its double bonds, but instead has a 'delocalised' zone where the electrons are free to roam throughout the ring (which explains the alternative representation of benzene with the circle in the ring, rather than the Lewis convention of drawing 3 conjugated double bonds).

Now, the reason why this is more stable is because the electrons in the molecule are more delocalised, and less confined in the same packet of space. This helps to lower the energy of the electrons (lets it find a more even 'equilibrium'), hence stabilising the molecule. To help improve your conceptualisation: does it make sense for a high energy electron (whizzing around) to stay in one tiny packet of space, or does it seem 'better off' when given more room to move around? It should make intuitive sense that the more delocalised an electron is, the more likely that state is (and hence more stable).

I'm not sure if this was what you were asking though. Did you mean to ask what features help a particular resonance state contribute more stability to the overall structure? Usually it involves keeping like charges away from each other - have you studied electrophiles and nucleophiles yet? There are some groups that prefer electrons more than others for example, or prefer positive charge more than others (due to electronegativity). Resonance states which place the charge nearer to the group that prefers it are particularly stable, and contribute markedly to the stability of the overall molecule.

If there are any more questions, feel free to ask. I almost forgot why this was the case, thanks for reviving my memory.

[mark_alec]

since the energy levels are different in s and p orbitals.

Not true. Ignoring the effects of other electrons in a system, s and p (and indeed d f g h...) are all degenerate. The energy only depends on the principle quantum number, n.

[Collin Li]

Clearly I was talking about a carbon atom with 6 electrons (and 4 electrons in n=2), not a monoelectronic atom or ion. The point is that they exist in different energy levels, and hence the C-H bonds cannot be equally spaced apart.