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Hybridization and Hypervalent molecules

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Qube
#1
May5-14, 02:55 PM
PF Gold
P: 442
Are species such as sulfuric acid and phosphoric acid better represented as charge-separated species rather than as species with no formal charge on any of the constituent atoms?

Also are the central atoms of trigonal bipyramidal molecules (e.g. phosphorous) hybridized as sp3d2 or do such species actually not use their empty d-orbitals while acting as the central atom?
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DrDu
#2
May5-14, 04:20 PM
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P: 3,593
Quite generally, d-orbitals do not play a role in hybridization in main group elements, although they have some influence as polarisation functions.
This also means that the bonding e.g. between S and O in sulfuric acid has a high ionic contribution.
Qube
#3
May5-14, 04:22 PM
PF Gold
P: 442
So basically all textbooks until very recently have been wrong about both these issues?

DrDu
#4
May6-14, 01:07 AM
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P: 3,593
Hybridization and Hypervalent molecules

Yes, bonding in the higher main group elements has been accessible to more and more precise quantum chemical calculations since the 1960ies, but most text books perpetuated ideas from Paulings books.
If you search the forum, you will find similar threads where I gave some references.
cgk
#5
May6-14, 09:45 AM
P: 426
OP, additionally to there not being d-hybridization, these so called hypervalent compounds often have bonds with highly ionic character, as you suspected (i.e., are better written with formal charges). For example, one point in http://dx.doi.org/10.1021/ct400687b is that SO3 can be well understood as just having three real covalent sigma bonds. While there are three additional bond orbitals, they are closer to delocalized pi-lone pairs on oxygen than real pi-bonds---they have only a small SO bonding component (and the pi-bonds are not correlated with each other, as the common rotating-double-bond resonance structure would suggest).

The same situation happens in many other compounds of higher main group elements, too. Despite what many textbooks still teach, they do not behave like their first-row counterparts at all---mainly in the sense that in these additional "multiple bonds", the pi bonds do not represent electrons shared to an more-or-less equal degree between two atoms. E.g. in SO3 above, 1.7 electrons of the pi bond sit on O, and only 0.3 on S. If you treat those as "almost one electron on S and almost one on O, equally shared", you are bound to run into problems...
Qube
#6
May6-14, 10:05 AM
PF Gold
P: 442
I'll have to raise this issue with my chem teacher. He's a bit old school and scoffs often at people "playing with their wave functions." Not sure how this will go across with him. Other than that the man's brilliant, if a bit outdated on the quantum mechanics stuff.

Probably a good way to earn some brownie points by raising a chemical quandary (at least a quandary for him!)

Also have you guys heard of the paper published in the 1960s which referred to d-functions in quantum wave calculations but was instead mistakenly taken by a lot of people to refer to d-orbitals (instead of d-functions)?
DrDu
#7
May6-14, 12:00 PM
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Quote Quote by Qube View Post
Also have you guys heard of the paper published in the 1960s which referred to d-functions in quantum wave calculations but was instead mistakenly taken by a lot of people to refer to d-orbitals (instead of d-functions)?
This may relate to the distinction between d-orbitals in hybrids and polarisation functions I mentioned before.
However, I don't know about a specific paper from the 1960ies.

If your teacher is old school and doesn't trust in quantum mechanical calculations then you could look out for some experiments which probe the bonding in main group compounds, like femtosecond pump probe spectroscopy.
However, I fear that understanding these experiments is way beyond high school level.
Qube
#8
May6-14, 08:28 PM
PF Gold
P: 442
Quote Quote by DrDu View Post
This may relate to the distinction between d-orbitals in hybrids and polarisation functions I mentioned before.
However, I don't know about a specific paper from the 1960ies.

If your teacher is old school and doesn't trust in quantum mechanical calculations then you could look out for some experiments which probe the bonding in main group compounds, like femtosecond pump probe spectroscopy.
However, I fear that understanding these experiments is way beyond high school level.
Good, since I'm not in HS nor am I at the HS level.

What are your thoughts on someone not trusting quantum mechanics? Are they fit to be teachers?

Also the paper I was talking about was this paper:

In 1990, Magnusson published a seminal work definitively excluding the role of d-orbital hybridization in bonding in hypervalent compounds of second-row elements. This had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds (or else unreasonably high energies and distorted geometries result), and the contribution of the d-function to the molecular wavefunction is large. These facts were historically interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency.[6]
http://pubs.acs.org/doi/abs/10.1021/ja00178a014
Qube
#9
May6-14, 08:52 PM
PF Gold
P: 442
Quote Quote by DrDu View Post
Yes, bonding in the higher main group elements has been accessible to more and more precise quantum chemical calculations since the 1960ies, but most text books perpetuated ideas from Paulings books.
Dammit Pauling, wrong about Vitamin C and wrong about bonding.
DrDu
#10
May7-14, 01:39 AM
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Quote Quote by Qube View Post
Dammit Pauling, wrong about Vitamin C and wrong about bonding.
:-)
Nevertheless a bit unfair, as his books are full of interesting stuff and insight. He arrived at many descriptions at a time before quantum chemical calculations were possible due to the lack of computers.
It is not surprising that many things had to be revised.
The astonishing thing is rather that teachers don't care about the results of theoretical chemistry of the last 50 years or so. Mostly because it didn't appear in their own curriculum.
sludger13
#11
May7-14, 08:57 AM
P: 75
Can anybody explain the difference between orbital role in hybridization and orbital role as polarization function? What is the interpretation of orbital polarization? Is it just electromagnetic quantum interaction achievable with mathematical description, giving the observed results? Or could it be interpreted with e.g. Coulomb law?

Anyway, the presented result of polarization function is (inexpertly) similar to hybridization - different shape of final orbital:
Attached Thumbnails
polarize.gif  
DrDu
#12
May7-14, 09:20 AM
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P: 3,593
Of course, but the admixture of d-type polarization functions is only some percent. So this is far from forming a set of hybrid orbitals each containing an equal amount of d-functions and all participating in bonding.
If you express the d-functions in terms of atomic d-orbitals, they contain also 4- 5- and higher d-orbitals.
Hence, all in all, the virtual unoccupied orbitals still are mainly of d-type.

Btw, to get a reasonable description of bonding in some sulfur compounds, you even have to include f-type polarization functions
Qube
#13
May7-14, 09:19 PM
PF Gold
P: 442
Quote Quote by DrDu View Post
:-)
Nevertheless a bit unfair, as his books are full of interesting stuff and insight. He arrived at many descriptions at a time before quantum chemical calculations were possible due to the lack of computers.
It is not surprising that many things had to be revised.
The astonishing thing is rather that teachers don't care about the results of theoretical chemistry of the last 50 years or so. Mostly because it didn't appear in their own curriculum.
Well I'm not sure where my teacher really stands on QM stuff; I think he appreciates it but he doesn't go into depth with it with us, instead choosing to emphasize stereochemistry (I got two open stereochemistry questions right now in the forums) and Bronsted-Lowry acid/base chemistry. He accepts MO theory; I just came across an old test question in which he discusses N2O5 and how it's actually ionic (ascertained through solid state evidence). He has at length discussed the wave and particle properties of the electrons. He just doesn't get into the calculation part of QM, instead poo-pooing faculty members who try to describe to him how to calculate resonance energy. Understandable, given that his undergrad math skills according to himself were non-existent and he had a 2.8 GPA and nearly got kicked out his pHD program.

He's come a long way I might say!

Perhaps I'll explore these topics more in-depth in a physical chemistry course (if I choose to take one).
DrDu
#14
May8-14, 01:33 AM
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Quote Quote by Qube View Post

Perhaps I'll explore these topics more in-depth in a physical chemistry course (if I choose to take one).
I doubt you will learn much about bonding there. Usually it is about thermodynamics, kinetics and spectroscopy.

You would have to take a course on theoretical chemistry.


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