In theoretical chemistry, the charge-shift bond is a proposed new class of chemical bonds that sits alongside the three familiar families of covalent, ionic, and metallic bonds where electrons are shared or transferred respectively.[1][2] The charge shift bond derives its stability from the resonance of ionic forms rather than the covalent sharing of electrons which are often depicted as having electron density between the bonded atoms. A feature of the charge shift bond is that the predicted electron density between the bonded atoms is low. It has long been known from experiment that the accumulation of electric charge between the bonded atoms is not necessarily a feature of covalent bonds.[3]
An example where charge shift bonding has been used to explain the low electron density found experimentally is in the central bond between the inverted tetrahedral carbon atoms in [1.1.1]propellanes. Theoretical calculations on a range of molecules have indicated that a charge shift bond is present, a striking example being fluorine, F2, which is normally described as having a typical covalent bond.[2] The charge shift bond (CSB) has also been shown to exist at the cation-anion interface of protic ionic liquids (PILs).[4] The authors have also shown how CSB character in PILs correlates with their physicochemical properties.
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I'm going to draw a molecule of benzene. And then we're going to think about if anything interesting might happen with that molecule. So let me draw it. So we have 6 carbons in a ring. 1, 2, 3, 4, 5 and 6 carbons in a ring. What's interesting about benzene, why it's different than cyclohexane, is that it has these 3 double bonds in the ring. So let's say we have these two carbons are double bonded to each other, these two carbons are double bonded, and then these two carbons over here are double bonded. Actually, I'll draw the hydrogens here just so that we remember that they're there. But I'll it in a subtle color. So this carbon right here is going to be bonded to how many hydrogens? It has 1, 2, 3 valence electrons already used up. So it's going to have one bonded to one hydrogen. This one right here, same thing. Bonded to one hydrogen. So it has 4 valence electrons. This one, same thing. I think you see the pattern. That each of these are-- they have three bonds to carbons, one single bond to two carbons, and then one extra double bond. And then the fourth bond is to hydrogen. So let me just draw all of the hydrogens here. I'm doing it in this dark color so we don't have to pay too much attention to it. Now this right here, this is benzine. And you're going to see a lot about benzine in the future. But in this video, we're going to study or try to understand a particularly interesting property of benzene, and that's resonance. And it's not a property of just benzene, it's a property of many organic molecules. But benzene is kind of the most fun version. So let's think about what might happen with this molecule right here. So I have this electron. Let me do that in a different color. I have this-- let me do it in this blue. I have this electron over here. What if this electron moved over to this carbon over here? So this carbon is still going to have the other electron in the bond. It's just going to kind of pivot around a little bit. So that electron moves over there. Now this carbon doesn't need 5 electrons, so this electron goes to that carbon right over there. Now this carbon doesn't need 5 electrons, so that electron goes back to the original carbon that lost that first electron. So at the end of the day, everyone has kind of broken even. If this happened, we might end up with a structure that looks like this. And I'll draw a two-way arrow because we can actually go in both directions. So let me draw just the carbon chain. So 1 carbon, 2 carbons, 3 carbons, 4 carbons, 5 carbons and 6 carbons. And then, over here, we had the double bond over here, but now it's moved over here. So now the double bond-- and actually, let me do it in a blue color so we see the difference. So now the double bond is over here. This blue electron has moved over there. This blue electron has moved up here. Actually, let me color code it, so it makes it very clear. So let's say this is a green electron. Now the green electron has moved from this carbon over to that carbon. We can imagine that it's done that. Then you would have this magenta carbon or this magenta electron that was with this carbon, but now it's moved over to this carbon over here. And now the double bond has shifted as well. That's what this arrow showed. We'll stick with the blue carbon over there. That blue carbon has moved down to the original carbon. And now the double bond has shifted over here. So we essentially have a very similar, really, a very similar molecule. This is actually just a rotated version of that. But we have these double bonds that could keep flipping back and forth between this position and that position over there. They can just keep on doing it. They can just keep flipping either backwards or forwards. And the reality of benzene is that it's actually never in either this structure or this structure. It's always, actually, in something right in between. The reality of benzene actually looks something more like this. And I'll just draw it without drawing all the carbons and the hydrogens. And obviously, in this case, let me draw the hydrogens here since I drew the hydrogens up here. This had the hydrogens over here. Don't want to forget those. If you ever forget them they're implicit. I want to draw the hydrogens. But if we just look at the overall ring, we know that the carbons and the hydrogens are implicit. The actual structure of benzene is actually in between that and that. In reality, you kind of have a half double bond between all of the carbons. So the reality is, is that it looks something like this. So you have half a double bond there, half a double bond there, half a double bond over here, half a double bond over here, and then half a double bond over here. And then we're almost done. And then, half a double bond over here. The reality of benzene is that these electrons are actually spinning around the whole ring. It's not flip-flopping between this structure and this structure. The actual structure, the lower energy state structure, is this right here. Now these Lewis diagrams or actually, I haven't drawn all of the Lewis electrons. But these are considered contributing structures. And you often draw these when you're doing reaction mechanisms. But the reality is, is that resonance, the resonance of these positions creates-- the reality of benzene is that it's actually sitting in this intermediate position. Now, this doesn't happen only with benzene. Another example, there's going to be many examples. But just so that we're familiar with maybe two of the best examples, another example that you'll see a lot in the context of resonance is the carbonate ion. So carbonate ion. You have a double bond to one oxygen and then you have single bonds to two other oxygens. And those two other oxygens have extra electrons. So if I were to draw this oxygen over here it has 1, 2, 3, 4, 5, 6 valence-- or actually, I should say, it has 7 valence. Let me make it very clear. So it has 1, 2, 3, 4, 5, 6, 7 valence electrons. It has one extra electron, so it has a negative charge. And the same is true for this one. It has 1, 2, 3, 4, 5, 6, 7 valence electrons. One extra. So it has a negative charge. If you were to just look at this, I guess you could call it this resonance structure or this contributing structure, you'd say hey, maybe this oxygen-- and this oxygen here is neutral, so it has 6 valence electrons. 1, 2, 3, 4, 5, 6. Maybe, just maybe, one of these electrons can be given to the carbon and then the carbon would lose an electron to this guy on top. So maybe you could imagine a situation where this electron right here gets given to the carbon. And when that gets given to the carbon, the carbon releases-- it all happen simultaneously. The carbon releases this electron and it goes back up to that oxygen over there. And so what's that going to look like if that were to happen? So if that were to happen, now our structure will look like this. We have a carbon. Now this carbon only has a single bond up here. And then we have our oxygen. The oxygen, it had its 6 valence electrons. 1, 2, 3, 4, 5, 6 valence electrons. But now it got this extra blue one. And now it got this extra blue one, so now it has 7 valence electrons, and it has a negative charge. Now this oxygen over here gave one of its electrons to the carbon. Now it is bonded with it. So now the carbon has a double bond. I'll actually do it in that color. Has a double bond with this oxygen down here. It gave an electron, so now it only has 6 valence electrons. 1, 2, 3, 4, 5, 6. And it is now neutral. And this oxygen over here, nothing really new happened to it. I could just copy and paste it. So let me copy and then let me paste it. So this one is just sitting right like that. But you could imagine a situation where this oxygen right here, then all of a sudden-- and it could have come from this oxygen up here or it could come from this oxygen right here. This oxygen says hey, I have an extra electron. Let me give it to the carbon. And then the carbon releases a double bond with one of the other oxygens. In this case, it would be this one. Let me draw it. So maybe this electron right here gets given to the carbon. Forms a double bond. Then the carbon can let go of an electron. And so this electron right here goes back to this oxygen. And so what happens? So if that were to happen, our structure looks like this. We have a carbon single bonded to an oxygen up here that has 1, 2, 3, 4, 5, 6, 7 valence electrons. That hasn't changed in-- we could call it this resonance reaction, or however you want to call it. So it still has a negative charge. We have this guy down here. He took his electron back. So now he has 7 valence electrons again. So 1, 2, 3, 4, 5, 6, 7 valence electrons again. And I can even show the one that he got back. That one's in purple. So he now has a negative charge. And this guy now gave an electron to the carbon. So he forms a double bond, a new double bond. So this guy forms a double bond with the carbon. He gave an electron, so he only has 1, 2, 3, 4, 5, 6 valence electrons, and is now neutral. Now these can all keep swapping between each other. You can even go from this structure to that structure. You can actually go from any one of these structures to any of the others. And the reality of the carbonate ion, let me write this down. This is the carbonate ion. The reality of it is that its true structure is some place in between all of these. So the true structure of a carbonate ion would look like this. You would have a carbon and you'd have three oxygens. They have at least one single bond with each of those three oxygens. And then you have 1/3 and then you have-- actually, I should say, you have 1/3 of a double bond with each of them. This is a 1/3 of a bond. This is not standard notation, but this is essentially going. 1/3 of the time, the electron is on that bond. And then the other 2/3 of the time, each of these oxygens have an extra electron. You could imagine almost having a negative 2/3 charge. Now, people normally draw one of these structures because this is a nice-- kind of you're dealing with whole numbers. But the reality of carbonate ions is that it's experiencing this resonance. That the electrons are actually always floating in between these forms. Actually floating across all of these bonds. And that actually, makes this molecule more stable. This is at a lower energy state than any of these forms. And the same thing is true with benzene. This right here, where we're in between these two structures, is actually at a lower energy state, a more stable state than either of these forms.
Valence bond description
The valence bond view of chemical bonding that owes much to the work of Pauling is familiar to many, if not all, chemists. The basis of Pauling's description of the chemical bond is that an electron pair bond involves the mixing, resonance, of one covalent and two ionic structures. In bonds between two atoms of the same element, homonuclear bonds, Pauling assumed that the ionic structures make no appreciable contribution to the overall bonding. This assumption followed on from published calculations for the hydrogen molecule in 1933 by Weinbaum and by James and Coolidge[5] that showed that the contribution of ionic forms amounted to only a small percentage of the H−H bond energy. For heteronuclear bonds, A−X, Pauling estimated the covalent contribution to the bond dissociation energy as being the mean of the bond dissociation energies of homonuclear A−A and X−X bonds. The difference between the mean and the observed bond energy was assumed to be due to the ionic contribution. The calculation for HCl is shown below.[5]
| Actual H−H | Actual Cl−Cl | H−Clcov Covalent bond energy H−Cl, arithmetic mean (H−H) and (Cl−Cl) |
H−Clact Actual H−Cl |
"Ionic contribution" H−Clact – H−Clcov | |
|---|---|---|---|---|---|
| Bond dissociation energy(kcal mol−1) | 103.5 | 57.8 | 80.6 | 102.7 | 22.1 |
The ionic contribution to the overall bond dissociation energy was attributed to the difference in electronegativity between the A and X, and these differences were the starting point for Pauling's calculation of the individual electronegativities of the elements. The proponents of charge shift bond bonding re−examined the validity of Pauling's assumption that ionic forms make no appreciable contribution to the overall bond dissociation energies of homonuclear bonds. What they found using modern valence bond methods was that in some cases the contribution of ionic forms was significant, the most striking example being F2, fluorine, where their calculations indicate that the bond energy of the F−F bond is due wholly to the ionic contribution.[2]
Calculated bond energies
The contribution of ionic resonance structures has been termed the charge−shift resonance energy, REcs, and values have been calculated for a number of single bonds, some of which are shown below:[2]
| Covalent contribution kcal mol−1 |
REcs kcal mol−1 |
% REcs contribution | |
|---|---|---|---|
| H−H | 95.8 | 9.2 | 8.8 |
| Li−Li | 18.2 | 2.8 | 13.1 |
| H3C−CH3 | 63.9 | 27.2 | 30.2 |
| H2N−NH2 | 22.8 | 43.8 | 65.7 |
| HO−OH | –7.1 | 56.9 | 114.3 |
| F−F | –28.4 | 62.2 | 183.9 |
| Cl−Cl | –9.4 | 48.7 | 124.1 |
| H−F | 33.2 | 90.8 | 73.2 |
| H−Cl | 57.1 | 34.9 | 37.9 |
| H3C−Cl | 34.0 | 45.9 | 57.4 |
| H3Si−Cl | 37.0 | 65.1 | 63.8 |
The results show that for homonuclear bonds the charge shift resonance energy can be significant, and for F2 and Cl2 show it is the attractive component whereas the covalent contribution is repulsive. The reduced density along the bond axis density is apparent using ELF, electron localization function, a tool for determining electron density.[2][6]
The bridge bond in a propellane
The bridge bond (inverted bond between the bridgehead atoms which is common to the three cycles) in a substituted [1.1.1]propellane has been examined experimentally.[7] A theoretical study on [1.1.1]propellane has shown that it has a significant REcs stabilisation energy.[8]
Factors causing charge shift bonding
Analysis of a number of compounds where charge shift resonance energy is significant shows that in many cases elements with high electronegativities are involved and these have smaller orbitals and are lone pair rich. Factors that reduce the covalent contribution to the bond energy include poor overlap of bonding orbitals, and the lone pair bond weakening effect where repulsion due to the Pauli exclusion principle is the main factor.[2] There is no correlation between the charge−shift resonance energy REcs and the difference between the electronegativities of the bonded atoms as might be expected from the Pauling bonding model, however there is a global correlation between REcs and the sum of their electronegativities which can be accounted for in part by the lone pair bond weakening effect.[2] The charge-shift nature of the inverted bond in [1.1.1]propellanes has been ascribed to the Pauli repulsion due to the adjacent "wing" bonds destabilising of the covalent contribution.
Experimental evidence for charge-shift bonds
The interpretation of experimentally determined electron density in molecules often uses AIM theory. In this the electron density between the atomic nuclei along the bond path are calculated, and the bond critical point where the density is at a minimum is determined. The factors that determine the type of chemical bond are the Laplacian and the electron density at the bond critical point. At the bond critical point a typical covalent bond has significant density and a large negative Laplacian. In contrast a "closed shell" interaction as in an ionic bond has a small electron density and a positive Laplacian.[2] A charge shift bond is expected to have a positive or small Laplacian. Only a limited number of experimental determinations have been made, compounds with bonds with a positive Laplacian are the N–N bond in solid N2O4,[9][10] and the (Mg−Mg)2+ diatomic structure.[11][disputed ]
References
- ^ Sini, Gjergji; Maitre, Philippe; Hiberty, Philippe C.; Shaik, Sason S. (1991). "Covalent, ionic and resonating single bonds". Journal of Molecular Structure: THEOCHEM. 229: 163–188. doi:10.1016/0166-1280(91)90144-9. ISSN 0166-1280.
- ^ a b c d e f g h Shaik, Sason; Danovitch, David; Wei, Wu & Hiberty, Phillippe.C. (2014) [1st. Pub. 2014]. "Chapter 5: The Valence Bond Perspective of the Chemical Bond". In Frenking, Gernod & Shaik, Sason (eds.). The Chemical Bond. Wiley-VCH.[failed verification]
- ^ Dunitz, Jack D.; Seiler, Paul (1983). "The absence of bonding electron density in certain covalent bonds as revealed by x-ray analysis". Journal of the American Chemical Society. 105 (24): 7056–7058. doi:10.1021/ja00362a007. ISSN 0002-7863.
- ^ Patil, Amol Baliram; Bhanage, Bhalchandra Mahadeo (2016). "Modern ab initio valence bond theory calculations reveal charge shift bonding in protic ionic liquids". Physical Chemistry Chemical Physics. 18 (23): 15783–15790. Bibcode:2016PCCP...1815783P. doi:10.1039/C6CP02819E.
- ^ a b The Nature of the Chemical bond, L. Pauling, 1940, 2d edition, pp. 49−59, Oxford University Press
- ^ Shaik, Sason; Danovich, David; Silvi, Bernard; Lauvergnat, David L.; Hiberty, Philippe C. (2005). "Charge−Shift Bonding—A Class of Electron-Pair Bonds That Emerges from Valence Bond Theory and Is Supported by the Electron Localization Function Approach". Chemistry: A European Journal. 11 (21): 6358–6371. doi:10.1002/chem.200500265. ISSN 0947-6539. PMID 16086335.
- ^ Messerschmidt, Marc; Scheins, Stephan; Grubert, Lutz; Pätzel, Michael; Szeimies, Günter; Paulmann, Carsten; Luger, Peter (2005). "Electron Density and Bonding at Inverted Carbon Atoms: An Experimental Study of a [1.1.1]Propellane Derivative". Angewandte Chemie International Edition. 44 (25): 3925–3928. doi:10.1002/anie.200500169. ISSN 1433-7851. PMID 15892137.
- ^ Shaik, Sason; Danovich, David; Wu, Wei; Hiberty, Philippe C. (2009). "Charge-shift bonding and its manifestations in chemistry". Nature Chemistry. 1 (6): 443–449. Bibcode:2009NatCh...1..443S. doi:10.1038/nchem.327. ISSN 1755-4330. PMID 21378912.
- ^ Messerschmidt, Marc; Wagner, Armin; Wong, Ming Wah; Luger, Peter (2002). "Atomic Properties of N2O4 Based on Its Experimental Charge Density". Journal of the American Chemical Society. 124 (5): 732–733. doi:10.1021/ja011802c. ISSN 0002-7863. PMID 11817931.
- ^ Tsirelson, Vladimir G.; Shishkina, Anastasia V.; Stash, Adam I.; Parsons, Simon (2009). "The experimental and theoretical QTAIMC study of the atomic and molecular interactions in dinitrogen tetroxide" (PDF). Acta Crystallographica Section B. 65 (5): 647–658. doi:10.1107/S0108768109028821. hdl:20.500.11820/5fa0a31e-7a10-466e-a0f3-239f685217e6. ISSN 0108-7681. PMID 19767687.
- ^ Platts, James A.; Overgaard, Jacob; Jones, Cameron; Iversen, Bo B.; Stasch, Andreas (2011). "First Experimental Characterization of a Non-nuclear Attractor in a Dimeric Magnesium(I) Compound". The Journal of Physical Chemistry A. 115 (2): 194–200. Bibcode:2011JPCA..115..194P. doi:10.1021/jp109547w. ISSN 1089-5639. PMID 21158464.
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This page was last edited on 13 May 2024, at 13:37