Chemical Forums
Chemistry Forums for Students => Organic Chemistry Forum => Topic started by: user11 on June 10, 2014, 07:13:56 AM
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What is the exact problem (the correct problem) with Molecular orbital theory and the case With Methane CH4? Is this the exact and correct problem that gave rise to the Hybridization theory:
The 2s orbital from carbon and 1S from hydrogen should combine 2 molecular orbitals which would (in theory) have a lower energy than the 6 resulting MO’s from combining Carbons 3x2P orbitals with 3x1S orbitals from hydrogen. Since the combination of 2S and 1s should have a lower energy we would expect that they would be completely filled with electrons (2 electrons in each), such that one of the P orbitals would be vacant and therefore there would be a missing covalent bond? My organic chemistry book by Clayden states that “No one hydrogen atom has more or fewer electrons than any other – they are all equivalent” so I’m trying to decipher what that means.
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The problem here is that 2s orbital of carbon has 2 electrons so they cant form a bond with 1s of hydrogen and then only 2p orbitals of carbon has any electrons in it and the last one is empty. So it cant form 4 bonds. There are two possible solutions to this problem:
a) you move one electron from 2s to 2p of carbon thus having 4 electrons ready for bonding. However, in this case, the s and p orbitals doesnt have same energy which means, when they form the bond, there will be one bond shorter that the other 3 ones (because the s orbital of carbon is more electronegative than the p orbitals)
b) you form 4 hybrid sp3 orbitals which all have the same energy and have 1 electron each. Those electrons can then form the bonds with hydrogen electrons and because they all have same energy then the bonds are same which is how is in real world.
hope I answered your question
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There's nothing wrong with MO theory to explain the structure of methane. In fact, as in many others cases MO theory actually is supported by experimental data much better than the hybridization (VB) approach.
For methane, see this link: http://www.users.csbsju.edu/~frioux/h2bond/MethaneMOBonding.pdf
Honestly, it's strange to me that hybridization is still taught as gospel in undergraduate organic chemistry classes. It has some historic significance and I guess is far easier to understand than MO theory, but I think in the long run it does students a disservice to make them believe that hybridization is the ONE TRUTH when it comes to structure of carbon-based molecules.
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Hybridization is wrong in some aspects (like the spectrum of methane), but it keeps bonds between atom pairs only instead of the whole methane molecule for instance - and from what I see on this forum, bonds between atom pairs help to explain reaction mechanisms, don't they?
You know, the diagrams with arrows indicating where electrons jump in molecules and radicals during a reaction. Too complicated for me, but apparently you chemists can use them to make sensible predictions. Until equivalent diagrams exist with molecular orbital (or do they exist?) atom-to-atom bonds are a useful notion, and it results from the outdated hybridization and co.
Hey, that may be a research topic if it doesn't exist yet! Make the equivalent diagrams but using molecular orbitals. Guaranteed headache.
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I'm not saying it should not be taught to undergrads, but the way it is taught in organic chemistry classes I think gives the false impression that it's the only molecular bonding model available. No doubt, many organic chemists think this way as a result, when in fact the hybridization model doesn't agree with experimental (spectroscopic) evidence in many cases, even for simple molecules like methane. Practically speaking, it probably doesn't matter, because if hybridization and atom-atom bonding schemes are functional enough to allow organic chemists to predict the course of chemical reactions and provide some understanding of molecular structure - good enough! But from a pedagogical standpoint, I think it's important for students to understand that all bonding models are just that - bonding models. Some of them work better to explain some things, and some of them work better to explain other things. There's a historical reason why hybridization arose as a concept, but also it has numerous and significant failings. An appreciation of this fact makes for more well-rounded chemists, even if the finer details of more robust models aren't exactly understood.
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Hybridization is wrong in some aspects (like the spectrum of methane), but it keeps bonds between atom pairs only instead of the whole methane molecule for instance - and from what I see on this forum, bonds between atom pairs help to explain reaction mechanisms, don't they?
You know, the diagrams with arrows indicating where electrons jump in molecules and radicals during a reaction. Too complicated for me, but apparently you chemists can use them to make sensible predictions. Until equivalent diagrams exist with molecular orbital (or do they exist?) atom-to-atom bonds are a useful notion, and it results from the outdated hybridization and co.
Hey, that may be a research topic if it doesn't exist yet! Make the equivalent diagrams but using molecular orbitals. Guaranteed headache.
1) Hybridization is also ''Molecular Orbital Theory'', it's just a tool to qualitatively guess the structure of a compound and has doesn't even exist in any computational chemistry formalism.
2) The structure of methane is predicted right by QC programs which don't use Hybridization
3) The Energy of a system is invariant under a unitary transformation of the orbital coefficients which means that there are infinitely many possible MO's. There are procedures to ''localize'' them and ''interpret'' from a chemists view. no Hybridization needed.
4) The ''curly arrow'' approach fails miserably when it comes to excited states and correlation diagrams are needed.
in summary, hybridization is not really more than a tool to remember the structure of carbon atoms in different environments.
I think it's important for students to understand that all bonding models are just that - bonding models.
yea but I think one could emphasize the difference between a model where atomic structure with nuclei and correlated electrons and QM are used and a model using potatoes and sticks.
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Interesting replies - So What does it take to be a good organic chemist? Which topics or ares are important to focus on? how important is it to completely understand Molecular orbital and Hybridization theories for succes in organic chemistry?
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If you're primarily interested in synthesis, probably not much specific knowledge about this topic is needed. However a physical organic chemist (i.e., studying optical, electronic or other physical physical properties of organic molecules) would need to have a good understanding of these bonding theories.
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I'm not saying it should not be taught to undergrads, but the way it is taught in organic chemistry classes I think gives the false impression that it's the only molecular bonding model available.
I can't think of one popular intro organic textbook these days that DOESN'T include both "hybridization" theory and "MO" theory.
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I'd have to check my organic chemistry textbook to see how deeply into molecular orbital theory it went. I'm sure it touched on Huckel treatments, which are based on MO theory, due to the importance of aromatic rings, but a rigorous comparison of the strengths and weakness of VB, MO, and hybridization bonding models for carbon-based molecules? I'd be surprised.
Regardless, what's in a textbook and what is actually taught in lecture are two different things. I am fairly certain I don't remember being taught how to apply molecular orbital theory, or god forbid a real symmetry treatment, to determine the structure and electron spectroscopy of an organic molecule. This might have something to do with the fact that physical chemistry is taught AFTER organic chemistry in most undergraduate chemistry programs (which itself, I believe, is due mostly to the fact that organic chemistry is a requirement for premedical students - which gets to the fact that I've always felt that chemistry curricula should be tied in some way to the future plans of chemistry majors... but I digress).
At the risk of committing the fallacy of gross generalization, I came out of my undergraduate program having very little appreciation of the connection between organic chemistry and physical chemistry. They were for all intents and purposes two completely different worlds. It would be nice if organic chemistry courses drew more on core physical and general chemistry courses like reaction rates, structural models, thermodynamics, and so forth, that are learned both the year after and the year before, rather than focusing almost exclusively on memorizing reaction mechanisms and how to form functional groups. I get that the latter is sort of the point of it all - but it'd still be nice if the two basic sides of chemistry were connected better in undergraduate chemistry programs.
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So What does it take to be a good organic chemist?
The answer to that question wont help anybody. If you ask somebody how to be cool and you act accordingly, you won't be cool. if you do anything a good organic chemists says what you should do, you won't be a good organic chemist. To be a good scientist means to be extremely fascinated by the things you investigate. If you think you want to understand what happens on a molecular level and it fascinates you, learn what we know so far. If you don't, don't learn it and keep cooking.
I am fairly certain I don't remember being taught how to apply molecular orbital theory, or god forbid a real symmetry treatment, to determine the structure and electron spectroscopy of an organic molecule. This might have something to do with the fact that physical chemistry is taught AFTER organic chemistry in most undergraduate chemistry programs
How long does your undergraduate program take? We had Physical, Organic and Inorganic Chemistry right from the beginning...
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In most American universities, a four year chemistry program consists of 1 year of general chemistry, then 1 year of organic chemistry, then 1 year of physical chemistry, then 1 year of some elective (inorganic chemistry often, but sometimes an advanced form of one of the other courses). There are also usually some lab courses built into the curriculum. There are also options often for advanced placement (skipping general chemistry, say). In my limited experience, what is taught in each of these classes is pretty conserved.
Just for kicks I checked out my organic chemistry textbook (Solomons 6th Ed) to see what coverage of molecular orbital theory is provided. The first chapter does have a few sections (spanning about 4 pages) on quantum mechanics, atomic orbitals, and molecular orbitals - the section on molecular orbitals is 11 paragraphs, about half of which deal with the hydrogen molecule, and none of which address applications to carbon-based molecules (there is a brief 1 page interlude into MO theory later in the book when discussing the electronic structure of benzene). The introductory section on hybridization, immediately following, is 4 pages by itself. The first section covers (topically enough) methane. Hybridization is presented as an excellent model of methane, and nowhere are weaknesses of the hybridization model addressed, not with respect to methane specifically or to organic molecules in general. Were I a student being introduced to organic chemistry for the first time, my take away from reading chapter 1 is that hybridization is a perfect theoretical model of the structure of organic compounds, and I would certainly have no inkling of any weaknesses in the model or any idea that molecular orbital theory or other models are in better agreement with experimental data in many cases.
Anyway, I didn't mean to make a crusade out of this. I just think chemistry professors need to do a better job of explaining the limitations of chemical theories to their students. I know if you start going too much into exceptions to rules, you can end up confusing people who are new to the subject area, but on the other hand students should at least be aware that limitations exist. Otherwise, you end up with students who think they're doing something wrong with they encounter a case where the simple model they learned about fails.
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... that molecular orbital theory or other models are in better agreement with experimental data in many cases.
I am not wanting to make you a defender of MO theory, but you have virtually addressed the poster's question. I think we all get that Pauling created hybridization to make methane symmetrical and tetrahedral. I also get that it cannot be correct as there is no such thing as 'hybridized emissions'. How is MO theory better for methane?
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@orgopete
Did you look at the link I provided above? The photoionization data on methane is consistent with an MO model but not a hybridization model of methane.
Nevertheless, molecular orbitals are mathematical constructs only. They are an approximation of the real, expected location of electrons in molecules - a good approximation in most cases, as it turns out, but an approximation nonetheless. Of course, all chemical structure models are approximations, and we judge their relative quality solely by how good their predictions agree with experimental data. In the case of the hybridization model applied to methane, the agreement with structure is good but the agreement with spectroscopy is poor. The agreement of MO theory with structure is good AND the agreement with spectroscopy is good. The tradeoff is simplicity of the model. If you are interested in understanding molecular geometry only, hybridization is a fine approach for organic molecules and it has the added benefit of being easy to understand and apply in many common bonding situations, and this is probably why it is taught heavily in organic chemistry courses. But if spectroscopic states of electrons is a concern, as it is for many electronic materials, a hybridization model is simply and purely inaccurate.
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The agreement of MO theory with structure is good...
I knew the photoelectron argument. That is what I was implying with the absence of hybridized emissions (and there is only one kind of orbital, sp3).
I can't ask Pauling, but as I recall from his paper, isn't structure the reason for creating hybridization? Why should Pauling create hybridized orbitals if MO theory matched the structure of methane so well?
By the way, I have seen MO touted over VBT, but I had always been looking for a clear explanation. I often found something like it was a better fit. I just wanted to know what that meant. I am not an advocate for either.
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The first question is relatively easy. Valence Bond slightly predates MO theory. When Linus Pauling proposed orbital hybridization as an outshoot of VB theory, MO theory was still in its infancy and wouldn't be applied to real molecules for several more years. Even beyond that, determining the tetrahedral shape of methane with hybridization/VB is a lot easier than determining it with MO theory, especially without the more recent mathematical tools like group theory and symmetry treatments.
The second question I think is predicated on a false impression that MO theory is somehow superior to VB theory. This isn't really the case. It's true that MO theory is more complex and is far more developed than VB, but this doesn't mean it's always better. VB approaches are still used today in many theoretical calculations of organic molecules. Ironically enough, while I complain about organic chemistry courses doing a disservice to students by glossing over MO theory, I could just as easily complain about physical chemistry courses doing the same thing by glossing over VB theory as a mere historical footnote.
That said, I do think MO theory has a far wider range of applications than VB theory. VB theory uses only atomic orbitals with valence (bonding) electrons to form localized chemical bonds, with most electron density located between neighboring atoms. This may be appropriate in most cases, and indeed theoretical treatments using VB and MO theories very frequently give, within a margin of error, similar results for chemical structure, reactivity, and electron density. VB theory is, I believe, computationally simpler (because only valence orbitals are involved), and so it is often relied upon for large or complex molecules, or where a qualitative result is sufficient. Just as with VB theory, an MO calculation will also tend to show that most electron density is localized between adjacently bonded atoms - although because of its better stage of development, results from modern MO calculations are usually more detailed and offer a better quantitative match to experimental data in many cases. This is especially true in molecules in which delocalized electrons are clearly important, because MO theory allows for the possibility that (indeed, is based on the assumption that) electrons are located in orbitals that span the entire molecule. Understanding spectroscopic behavior is an additional exclusive strength of MO theory, because of its intimate relationship with molecular symmetry and its ability to actually predict electronic transitions.
So, were I to sum it up generally in a sentence: VB theory is a good qualitative approach that does a very good job of describing molecular structure and reactivity in most cases, whereas MO theory does a better job when quantitative predictions of molecular characteristics are important. MO theory also succeeds in some specific situations that VB is simply not well suited, such as open shell or parametric molecules, geometries when lone pairs are involved, and so on.
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MO theory also succeeds in some specific situations that VB is simply not well suited, such as open shell or parametric molecules, geometries when lone pairs are involved, and so on.
Hartree-Fock based calculations often fail miserably with open-shell transition metal complexes... The Quantum chemistry workhorse at the moment is DFT anyway...
Are there actual calculations done with VB at the moment? I mean I can treat Valence electrons only in MO-theory (what ever you mean by that) as well.
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Are there actual calculations done with VB at the moment? I mean I can treat Valence electrons only in MO-theory (what ever you mean by that) as well.
Absolutely: go to Google scholar, type in "valence bond approach", and you will see dozens of papers that use a valence bond model. Honestly, I think most serious computational chemists view the two models as opposite ends of a spectrum, one (MO) treating electrons as too delocalized and one (VB) treating electrons as too localized. Many modern quantum chemical methods draw on both approximations to get to a better result that's somewhere in the middle. But it's been quite some time since I've been involved in serious research based on quantum chemical calculations, so you'll have to pardon me if my memory of these issues is a little fuzzy. Honestly, I have never done modern VB calculations so I do not know what developments are incorporated into them. I can only speak of VB theory in a general way. Rigorous comparisons between modern VB and modern MO quantum chemical methods are beyond my level of expertise.
DFT is certainly great and is commonly used by professionals, but it's not very useful toward conceptualizing bonding, especially for students. There simply isn't enough time to teach such advanced quantum chemical techniques in a basic undergraduate curriculum.
(What I meant by treating valence electrons only in VB theory is that bonds are assumed to be formed only by adjacent atomic orbitals that are occupied by valence electrons. This is very different from molecular orbital theory, where molecular orbitals are created by combining atomic orbitals spanning the entire molecule, and usually not just those that have valence electrons - although you are right that simple molecular orbital treatments, such as those in homework problems in intro physical chemistry courses - often make the approximation that valence atomic orbitals are the only ones that need to be considered for a qualitative picture of what's going on. In principle, any atomic orbital of the appropriate symmetry, no matter how far away from the valence orbitals, will impact the energy of the molecular orbital, but the degree of interaction of two orbitals scales inversely with the difference in energy between them, so in practice low energy filled shells can be ignored. An example is an MO treatment of a diatomic, in which the 1s orbitals are neglected. A student drawing an MO diagram can do this and still have a nice qualitative understanding of the magnetic and bonding properties of the molecule. A quantum chemist will not neglect these low-lying orbitals, though, because they do impact the actual energy of the molecular orbitals to a small degree.)