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Derek Lowe The 2002 Model

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Derek Lowe, an Arkansan by birth, got his BA from Hendrix College and his PhD in organic chemistry from Duke before spending time in Germany on a Humboldt Fellowship on his post-doc. He's worked for several major pharmaceutical companies since 1989 on drug discovery projects against schizophrenia, Alzheimer's, diabetes, osteoporosis and other diseases. To contact Derek email him directly: Twitter: Dereklowe

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« Argumentum ad Crumenam | Main | What's So Special About Ribose? »

July 7, 2009

Another Thing We Don't Know

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Posted by Derek

Hydrogen bonds are important. There, that should be an sweepingly obvious enough statement to get things started. But they really are - hydrogen bonding accounts for the weird properties of water, for one thing, and it's those weird properties that are keeping us alive. And leaving out the water (a mighty big step), internal hydrogen bonding is still absolutely essential to the structure of large biological molecules - proteins, complex carbohydrates, DNA and RNA, and so on.

But we don't understand hydrogen bonds all that well, dang it all. It's not like we're totally ignorant of them, for sure, but there are a lot of important things that we don't have a good handle on. One of these may just have been illustrated by this paper in Nature Structural and Molecular Biology by a group from Scripps. They've been working on understanding the fact that all hydrogen bonds are not created equal. By carefully going through a lot of protein mutants, they have evidence for the idea that H-bonds that form in polar environments are weaker than ones that form in nonpolar ones.

That makes sense, on the face of it. One way to think of it is that a hydrogen bond in a locally hydrophobic area is the only game in town, and counts for more. But this work claims that such bonds can be worth as much as 1.2 kcal/mole more than the wimpier ones, which is rather a lot. Those kinds of energy differences could add up very quickly when you're trying to understand why a protein folds up the way it does, or why one small molecule binds more tightly than another one.

Do we take such things into account when we're trying to compute these energies? Generally speaking, no, we do not - well, not yet. If these folks are right, though, we'd better start.

Update: note that the paper itself doesn't suggest that this is a new idea - they reference work going back to 1963 (!) on the topic. What they're trying to do is put more real numbers into the mix. And that's what my last paragraph above is trying to state (and perhaps overstate): it's difficult to account for these thing computationally, since they vary so widely, and since we don't have that good a computational handle on hydrogen bonds in general. The more real world data that can be fed back into the models, the better.

Comments (7) + TrackBacks (0) | Category: In Silico


1. happydog on July 7, 2009 7:24 AM writes...

I hate to tell you Derek, but this isn't exactly new to computational chemists. It's been known for a while that if you remove hydrogen bonds (or salt bridges) from solvent, you end up paying and entropic penalty overall, but the resulting H-bonds are significantly stronger. Hendsch and Tidor have been studying this for the past fifteen years or so. It was also something I investigated during my PhD thesis studying lipid/protein simulations. Of course, whenever I tried to point this out to med. chemists in industry, they always looked at me like I had sprouted another head for mentioning such an absurd idea. I guess it takes work from Scripps or an Ivy League lab before you guys will start to take things seriously.

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2. RB Woodweird on July 7, 2009 7:46 AM writes...

happydog sez: "Of course, whenever I tried to point this out to med. chemists in industry, they always looked at me like I had sprouted another head for mentioning such an absurd idea."

I don't think they disbelieved you, and they probably did have an intuitive understanding that what you were selling was correct. But without some predictive way to use the model it is just pretty thoughts.

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3. Ty on July 7, 2009 8:26 AM writes...

Did it really take an extensive structural study in a Nature S&MB paper and/or 15 years of dedicated computational work to realize this? Why is TFA much more strongly acidic in dichloromethane than in methanol (or water)? What's so special about 'hinge binding' in kinase inhibitors? Why are nitrogen heterocycles so ubiquitous in drug (candidate) molecules? I'll just answer the last question. Because they can make hydrogen bonds in the hydrophobic pocket without paying too much desolvation penalty. Most medicinal chemist must have seen a certain ring nitrogen, not necessarily a better H-bond acceptor than, say, carbonyl oxygen, swinging the potency by >100~1000-fold.

Of course knowing by intuition and experience is one thing and proving it is another... yet, still, is a collection of crystal structures and/or some ab initio calculation really a superior proof?

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4. retread on July 7, 2009 8:29 AM writes...

Voet "Biochemistry" chapter 2. "Water is a chemically reactive liquid with such extraordinary physical properties that, if chemists had discovered it in recent times, it would undoubtedly have been classified as an exotic substance."

[ Proc. Natl. Acad. Sci. vol. 98 pp. 10533 - 10540 '01 ] "Despite the construction of hundreds of model force fields for use in simulations, the great advances in computational technology, and the development of powerful ab initio molecular dynamics methods we remain unable to accurately calculate the properties of liquid water (e.g. heat capacity, density, dielectric constant, compressibility) over wide ranges in conditions. We do not yet have a satisfactory molecular description of how a proton moves in the liquid. We do not fully understand the molecular nature of the surfaces of either ice or liquid water. Although it is clear that the hydrogen bond network and its fluctuations and rearrangement dynamics determine the properties of the liquid, no experimental studies are available showing detalied information about this process (without considerable interpretation)."

Hopefully things have improved since all this was written.

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5. KwadGuy on July 7, 2009 9:14 AM writes...

This is very well known among computational chemists and has been for a very long time. Hydrogen bonds between the ligand and the receptor that replace interactions with water in the unbound form are usually worth very little, or else are unfavorable. Hydrogen bonds formed in a solvent occluded region (buried) are generally worthwhile to binding.

This is modifying ligands to improve the number of hydrogen bonds in solvent exposed regions almost NEVER provide the boost to binding expected, while those made to regions of the ligand deep in the binding pocket DO usually provide the intuitive boost in binding.

As I said, this is well well well known among (good) computational chemists and (some) (good) medicinal chemists.

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6. Curious Wavefunction on July 7, 2009 9:28 AM writes...

I agree. Computational chemists and physical biochemists have known for a while that hydrogen bonds in hydrophobic environments are worth more those in polar environments. Hydrogen bonds formed by proteins are usually exchanged for those with water and thus as someone pointed aout, are worth very little. Physical biochemists have also known this for a while; for instance see this article on the diffiiculty of rationalizing alpha helix formation on the basis of h-bond energetics.

The problem is of course that it is still not easy to engineering the subangstrom features of hydrogen bonds in structure based design.

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7. mthomson on July 7, 2009 11:07 PM writes...

It should also be noted that in the same issue of Nature Structural and Molecular Biology a group from Scripps also retracted another crystal structure. How many structures retracted by Scripps in the last couple of years?

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