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There are two main ways to look at hydrogen bonding. The first is electrostatic, where the electronegativity of the atoms is used to describe the interaction. Your argument about chlorine being more electronegative than nitrogen is a good one suggesting that the electrostatic argument is only part of the story, and there is at least one study that suggests ...
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Though it does go against your intuition, you've actually mentioned the answer in your question. Stibane has a higher boiling point than ammonia/azane on account of van der Waals interactions (owing to the larger size of the antimony atom).
Our teacher had actually posed this question to us during my first year of high-school. All of us were incredulous ...
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I think there are a variety of qualitative ways of looking at this:
Perhaps the most obvious is that $\ce{H2O}$ can form a greater number of hydrogen bonds due to having an equal numbers of hydrogen bond acceptors and donors. Each of the hydrogen atoms can be hydrogen bond acceptors; each of the lone pairs on the oxygen can be donors. In $\ce{HF}$ however ...
17
It appears that perhaps water may have symmetric hydrogen bonds between oxygen atoms, but it would take at least 60 GPa of pressure to make water to bond like that.
Article named very descriptively "Compression of Ice to 210 gigapascals: Infrared Evidence for a Symmetric Hydrogen-Bonded Phase" proves it:
Protonated and deuterated ices ($\ce{H2O}$ and $\...
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This question crossed my mind during a lab today because we were making $\ce{Ni(dmg)2}$, which is shown below.
Well first, this is charged in some sense, but it is net-neutral. Also, the configuration shown above is only one way to draw it, and it seems as if everything should be symmetrical and thus that hydrogen should be equally shared between the two ...
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I don't think there is any such traditional definition requiring $\ce{N}$, $\ce{O}$ or $\ce{F}$.
For example, in table 7 and the discussion thereof in Hydrogen Bonding Annual Review of Physical Chemistry Vol. 22: 347-385 the hydrogen bonds in the following species are discussed:
$\ce{ClHCl^-}$
$\ce{BrHBr^-}$
$\ce{IHI^-}$
$\ce{BrHCl^-}$
as well as related ...
15
If I understood you correctly, you are talking about the peptide bond nitrogens ($\ce{R-C(=O)-\color{red}{N}H-R}$). This is, reduced to its significant chemical functional group an amide, more precisely a carboxylic amide.
The amide nitrogen technically has a lone pair and thus technically could function as a hydrogen bond acceptor when viewed alone. ...
14
N, O, F have atomic numbers $Z=7, 8, 9$, respectively. Chlorine has $Z = 17$ which is much larger. Consequently, the atom is larger as well and it is more diffuse. Equivalently, the lone pairs in chlorine are at the 3-level which is too high. Because the hydrogen bond isn't a real bond but a dipole-dipole attraction and because the force between two dipoles ...
14
According to the IUPAC gold book a van der Waals force is:
The attractive or repulsive forces between molecular entities (or between groups within the same molecular entity) other than those due to bond formation or to the electrostatic interaction of ions or of ionic groups with one another or with neutral molecules. The term includes: dipole–dipole, ...
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Well, it turns out that this is a very active area of research. I will only summarize what I understand to be true about the covalent nature of the hydrogen bond, so I'm sure the explanation could be more detailed and potentially more accurate in some places (I hope someone gives a more detailed answer), but here's what I've got.
As you said, it has been ...
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It is safe to say that there will always be intermolecular forces at play. At the time where you will consider these you should already have a good idea about the molecules involved in your system.
Based on the composition and molecular structures you can make certain assumptions. In a molecule it is straight forward to estimate (bond) polarities based on ...
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@Dissenter, you really need to stop asking questions like this. Once again you point out that the emperor has no clothes! :)
It turns out that the mechanism of simple proton transfer is an active area of investigation. There is not a single yes or no answer that would adequately address your questions. But if this truth were told in the introductory ...
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Preamble:
I will conveniently avoid entering into a debate as to what a hydrogen bond is, given the numerous theories which attempt to do so from various perspectives.
I will, however, attempt to use the relatively cheap and easy natural-bond orbital (NBO) analysis method to examine the possible existence of an intramolecular hydrogen bond (HBond) in the ...
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It can indeed hydrogen bond, this can (to a certain extent) be observed by NMR. Due to the internal hydrogen bonding, the phenolic proton is surprisingly consistent, and shows no dependence on concentration.
Image reproduced from Basic one- and two-dimensional NMR spectroscopy, Wiley, 2010.
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At Wikipedia you find this structure (drawn by Benjah-bmm27 on Wikimedia Commons):
So, the answer (c) is correct. Yomen Atassi correctly stated that in such a hydrogen bond the two electronegative partners and the hydrogen prefer a linear arrangement, as this maximizes the orbital overlap for the hydrogen bond. That the configuration (c) is preferred over (...
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The strength of a hydrogen bond somewhat depends on the $\ce{X-H\bond{...}X}$ angle that the hydrogen-bonding hydrogen forms with the two electronegative elements $\ce{X}$. In our case, carboxylic acids or alcohols, $\ce{X} = \ce{O}$ so the angle is $\ce{O-H\bond{...}O}$. The ideal angle for this fragment is $180^\circ$.
As you have drawn for carboxylic ...
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You are right, in a sense.
A more accurate statement would be that as the temperature decreases, the average lifetime of a hydrogen bond increases. Thus, on average the liquid/solid/molecule will be more stable. This is because, as you point out, hydrogen bonds are weak enough to be broken by collisions due to thermal motions (translation and rotation). ...
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In Schaefer's paper,[1] which you linked, the authors describe nitromalonamide as being the compound with the lowest calculated barrier to intramolecular proton transfer, or tautomerisation.
Earlier this month (Feb 2019), Perrins and Wu used NMR isotope shifts to show that nitromalonamide possesses a symmetric hydrogen bond.[2] This is "the first case, to ...
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NO, Water is not only the example for hydrogen bonding. It's most common example for basics.
Hydrogen bonds occur in DNA, proteins, polymers, etc. But for introductory level it is easy to explain with water.
DNA: Hydrogen bonding between guanine and cytosine, one of two types of base pairs in DNA.
Polymer: Para-aramid structure
Images: Wikipedia
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Molecules with inter-molecular hydrogen bonds tend to associate with one another, while molecules with intra-molecular hydrogen bonds tend to associate with themselves.
Ortho- and para-nitrophenol are often used as examples of molecules with intra- and inter-molecular hydrogen bonds respectively (see picture).
Because of their molecular structures the ...
9
Let me answer the second question first, because the answer is so much easier: Of course the hydrogen bond strength depends among others upon the carbon oxygen is attached to. Consider phenol and methanol or methanol and formic acid.
The guess as to which hydrogen bond is stronger is — I have to admit it — my speculation. But I think that methanol should be ...
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True, Indigo can form hydrogen bonds, but with what? Is it water? Nope.
It makes hydrogen bonds with itself.
This makes the molecule incapable of bonding with water. Not only that, but the molecule itself is quite symmetric, with oppositely oriented polar bonds, that cancel out each others' dipole moments. Notice the molecule's similarity to an ace of ...
answered Jun 5 '17 at 15:59
Pritt says Reinstate Monica
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8
I know that the standard textbook answer is that a polar protic
solvent is able to better stabilize/solvate carbocations through
hydrogen bonding.
In an $\ce{S_{N}1}$ reaction, a polar protic solvent can stabilize the leaving group through both hydrogen bonding and solvent dielectric effects. While these same solvent dielectric effects can also ...
8
Yes. Temperature disrupts bond of all kinds. Heat up a protein hot enough and you can even disrupt its primary structure - the linear sequence of amino acids, and amino acids are held together through covalent bonds. This is why prion contaminated organisms and instruments must be heated to extremely, extremely high temperatures. The question is whether or ...
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The interactions do not limit to only dipole-dipole, which means positive charge is not everything.
What's special about $\ce{NF3}$? It has no hydrogens. As a result, it cannot form hydrogen bond, therefore its boiling point (or condense point) is far lower than the others'.
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It is simple. Longer the chains higher the similarity between them. In other words the OH group are getting "diluted".
It is more difficult for the donor to find the accepting counterpart, and the number density of the hydrogen bonds, both for volume and mass, gets lower.
The interaction between molecules becomes dominated by VdW forces.
To visualize this ...
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I'm not sure whether the term "half-life" means the same as "lifetime" but I found a very interesting website about hydrogen bonding in water (with a lot of references) where it is stated that the lifetime of a hydrogen bond is in the range of 1 - 20 ps (it gives the following reference for this: F. N. Keutsch and R. J. Saykally, Water clusters: Untangling ...
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Short answer: the ketone oxygen can participate in hydrogen bonding
The compound you mention, methoxymethane, is an ether not a ketone.
Ethers don't hydrogen bond very well because the oxygen isn't very polarized.
On the other hand, the carbonyl in a ketone is polarized. We can draw resonance structures that show this polarization.
(image source)
...
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Disclaimer: I am not a computational chemist, I am completely new to this stuff. Just a humble student of @pentavalentcarbon
However, my thought is that if you have interacting OH groups, the hydrogen bonding would make the eclipsed conformations more stable. Is this be the case? Why or why not? For example, a small molecule containing two carbons, each ...
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The $\mathrm pK_\mathrm a$ of the $\ce{CH2}$ group in a dicarbonyl compound is roughly $10$, whereas the $\mathrm pK_\mathrm a$ of a carboxylic acid is roughly $5$. For example, dimethyl malonate has a $\mathrm pK_\mathrm a$ of $13$ whereas acetic acid has a $\mathrm pK_\mathrm a$ of $4.7$. So, the C–H bond acidity is not likely to be under consideration in ...
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