[PicturesOfLilly]

Hi all. These are some questions that I've gathered from quite a bit of study that I'm still baffled by... and feel I should understand before going any further. Don't feel that you have to answer all of them. Anything is of help.

1. Regarding the benzene ring, my chemistry book, it says "we can imagine the 6 valence electrons, one from each carbon atom, belonging to the whole molecule, instead of being localised in the three double bonds". I'm not sure where it's getting the '6' from. I mean I know that carbon has 6 electrons, but I don't think that's got anything to do with it. Given that carbon is tetravalent, I would have thought of it as there being 4 valence electrons for each carbon atom... so that would be 4 by 6, which would be 20 valence electrons in the entire structure.

2. This question is to do with sub orbital. I know there's a total number of 18 electrons accounted for within the orbitals of the third energy level; 3s2, 3p6, 3d10; and I've shaded these elements a certain colour on my periodic table. What I'm asking about here is about the orbital of the fourth energy levels. I want to shade these in my periodic table. I think 4d orbital is the elements from 39 to 48 in the d block, but I want to know which ones are in the 4f energy level?

3. In the case of the water molecule, can one of these lone pairs be thought of as the electrons within 2s², with the other lone pair being thought of as the electrons within 2p_x²? Likewise I'm wondering about the ammonium molecule, and whether I can think of the one lone pair on the N atom as being the electrons in the 2s² orbital? If I were to ask about O₂, would it be that (on either of the oxygen atoms) that one of the lone pairs is from 2p_x², and the other lone pair is from 2s²?

4. Following on from the reasons why elements Cr & CU have different electronic configurations than what one might expect due to the fact that p, d and f sublevels are more stable if they are either half or completely filled. So my question is, if this is the case, then how come the same isn't true for carbon and fluorine?

5. What's the deal with Boron Trichloride only achieving an outer electronic configuration of 6? If it were ionic bonds I could understand, but it's covalent bonds! From what I can see, it would need another chlorine atom to join in in order for there to be another shared pair of electrons, thereby giving the octet. Of course that would make it tetravalent; which it's not! In thinking of it in terms of sigma and pi bonds, I can see how the half filled x orbital would form a sigma bond, but this way of thinking about it doesn't help me out either! The same issue with Beryllium chloride.

Thank you

[Meter]

1. I think it means that the electron in the p-orbital perpendicular to the plane of each carbon can be delocalized around the ring.

3. Yes, you can say that the lone pair in nitrogen is in the 2s orbital (as per the Aufbau principle). However, we like to say that the 2s orbitals and the 2p orbitals hybridize to form four sp3 hybridized orbitals where one of the sp3 orbitals contain a lone pair. Similarly, with oxygen you also achieve four sp3 hybridized orbitals where two contain lone pairs, but you could say that the lone pair is contained in 2p_x if you wanted to (as this is the case in elemental oxygen).

[mjc123]

1. 3 valence electrons on each carbon are involved in sigma bonds - two to C and one to H. The fourth is in the p orbital perpendicular to the ring plane. These orbitals can overlap giving rise to the delocalisation. (Oh, and 6 x 4 = 24!)

2. You are right about the 4d elements. The 4f elements are the actinides (in the f block, geddit?) It should be easy to find a list of the electron configurations of the elements, and see where a particular subshell is being filled up.

3. The sp³ description is preferable, as the bond angles are closer to the tetrahedral value than to the 90° one would expect for pure p orbitals. (And it's the ammonia molecule.)

4. Where do you get the extra electron for C and F? For Cr and Cu it comes from the 4s orbital, which has 1 electron rather than 2, e.g. Cu is 4s¹ 3d¹⁰ rather than 4s² 3d⁹.

5. B has only 3 valence electrons, so it can only form 3 bonds (in a neutral compound). It can form a dative bond with Cl^- to give the BCl₄^- ion, i.e. it acts as a Lewis acid. Don't get fixated with the octet. You can't pluck electrons out of thin air to make it.

[Orcio_87]

3. In the case of the water molecule, can one of these lone pairs be thought of as the electrons within 2s2, with the other lone pair being thought of as the electrons within 2px2? Likewise I'm wondering about the ammonium molecule, and whether I can think of the one lone pair on the N atom as being the electrons in the 2s2 orbital? If I were to ask about O2, would it be that one of the lone pairs is from 2px2, and the other lone pair is from 2s2?

Isn't simplier to imagine two lone pairs in H₂O as two sp₃ oxygen hybrydized atomic orbitals? (of course it's simplification)

As of O₂ - no, both pairs are 2s atomic orbitals.

[mjc123]

Correction, I should have said the 4f elements are the lanthanides. Actinides are 5f! D'oh!

[PicturesOfLilly]

Thanks for taking the time to answer

4. Where do you get the extra electron for C and F? For Cr and Cu it comes from the 4s orbital, which has 1 electron rather than 2, e.g. Cu is 4s¹ 3d¹⁰ rather than 4s² 3d⁹.

I understand, but shouldn't Fluorine be [1s², 2s¹, 2p⁶]? Do you see what I mean?

[mjc123]

The electron configurations of the transition metals result from 4s and 3d being very close in energy, so the energy required to promote from 4s to 3d is compensated by the stabilisation of a full or half-full 3d subshell. There is a bigger energy difference between 2s and 2p in the first-row elements, so 2s¹ 2p⁶ is definitely higher energy than 2s² 2p⁵.

[PicturesOfLilly]

3. Yes, you can say that the lone pair in nitrogen is in the 2s orbital (as per the Aufbau principle). However, we like to say that the 2s orbitals and the 2p orbitals hybridize to form four sp3 hybridized orbitals where one of the sp3 orbitals contain a lone pair. Similarly, with oxygen you also achieve four sp3 hybridized orbitals where two contain lone pairs, but you could say that the lone pair is contained in 2p_x if you wanted to (as this is the case in elemental oxygen).

But which is it? Is it black or white? It's easier for me to think of it as if these electrons are from the 2s orbital, as they are closer to the nucleus and wouldn't be involved in bonding.