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DBL%20Hendrix%20small.png College chemistry, 1983

Derek Lowe The 2002 Model

Dbl%20new%20portrait%20B%26W.png After 10 years of blogging. . .

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: derekb.lowe@gmail.com Twitter: Dereklowe

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October 26, 2010

Enthalpy and Entropy Again

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Posted by Derek

Earlier this year, I wrote here about using calorimetry in drug discovery. Years ago, people would have given you the raised eyebrow if you'd suggested that, but it's gradually becoming more popular, especially among people doing fragment-based drug discovery. After all, the binding energy that we depend on for our drug candidates is a thermodynamic property, and you can detect the heat being given off when the molecules bind well. Calorimetry also lets you break that binding energy down into its enthalpic (delta-H) and entropic (T delta-S) components, which is hard to do by other means.

And there's where the arguing starts. As I mentioned back in March, one idea that's been floating around is that better drug molecules tend to have more of an enthalpic contribution to their binding. Very roughly speaking, enthalpic interactions are often what med-chemists call "positive" ones like forming a new hydrogen bond or pi-stack, whereas entropic interactions are often just due to pushing water molecules off the protein with some greasy part of your molecule. (Note: there are several tricky double-back-around exceptions to both of those mental models. Thermodynamics is a resourceful field!) But in that way, it makes sense that more robust compounds with better properties might well be more enthalpically-driven in their binding.

But we do not live in a world bounded by what makes intuitive sense. Some people think that the examples given in the literature for this effect are the only decent examples that anyone has. At the fragment conference I attended the other week, though, a speaker from Astex (a company that's certainly run a lot of fragment optimization projects) said that they're basically not seeing it. In their hands, some lead series are enthalpy-driven as they get better, some are entropy-driven, and some switch gears as the SAR evolves. Another speaker said that they, on the other hand, do tend to go with the enthalpy-driven compounds, but I'm not sure if that's just because they don't have as much data as the Astex people do.

So as far as I'm concerned, the whole concept that I talked about in March is still in the "interesting but unproven" category. We're all looking for new ways to pick better starting compounds or optimize leads, but I'm still not sure if this is going to do the trick. . .

Comments (17) + TrackBacks (0) | Category: Analytical Chemistry | Drug Assays | Life in the Drug Labs


COMMENTS

1. barry on October 26, 2010 8:36 AM writes...

what I would want from calorimetry (having never used it in screening) is stoichiometry. I want to know that my "hit" binds one-to-one with the target, rather than relying on aggregation--or sticking indiscriminately all over the protein surface

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2. CYTIRPS on October 26, 2010 8:52 AM writes...

It is likely to stay in the "interesting but unproven" category forever. One could do all the in vitro enthalpy and entropy measurement but only the in vivo numbers matter. One of the deadliest mistakes in the drug industry is that most of the management believe that biochemical assays of isolated molecular target could replace phenotypic assay. The central idea of "enthalpy over entropy" is that entropic contribution could change significantly from in vitro to in vivo, such as viscosity, pH, dielectric constant and hetero protein-protein interactions. If one has to bet between an enthalpic driven compound and an entropic driven compound, one should develop the enthalpic driven compound.

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3. Brian on October 26, 2010 9:46 AM writes...

Although slightly related, is using calorimetry to determine if the molecule has a exothermic onset when it is melted. Frequently, as process chemists, we modified the synthesis conditions such that, the reaction temperature was an arm's length away from the temperature onset of an exothermic event. I remember seeing a picture of a steel ball that had burst when the molecule inside had reached the critical temperature. If the molecule had been tested by DSC before getting to us, it would have been eliminated from consideration of development.

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4. mad on October 26, 2010 9:52 AM writes...

Does binding affinity always correlate with lower delta G? Certainly selectivity does not.

My initial thought is thermodynamics is too broad a measure, especially if breaking it down to delta H and delta S doesnt add more information.

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5. DLIB on October 26, 2010 11:49 AM writes...

I hope by " better " you're not talking about the binding constant

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6. LeeH on October 26, 2010 1:58 PM writes...

The problem with calorimetry is that it's often much harder to do than performing the assay itself. In addition, it only gives you the binding constant and not a functional readout. That might be OK for an enzyme inhibition assay, but not much else.

I'm skeptical about the assertion that enthalpy is more important than entropy. It's just not sensible thermodynamically. Heat is king.

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7. HTSguy on October 26, 2010 3:10 PM writes...

@Barry: Both isothermal calorimetry and SPR (e.g. Biacore) are often used to determine stoichiometry of binding for just the reasons you mention.

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8. sepisp on October 27, 2010 2:47 AM writes...

A professor (from UKu) had a point about this, in form of "should I add hydroxyls or not". Namely, often you want to increase selectivity by adding a hydrophilic group, and you hope that it will bind to a corresponding group in the protein. The problem is then that the molecule is in water solution before binding, and the molecule-water interaction is not necessarily less favorable than molecule-protein interaction. That is to say, by adding a hydroxyl you don't automatically gain anything. Furthermore, you'll have to make the binding entropically palatable also.

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9. azetidine on October 27, 2010 4:46 AM writes...

"As I mentioned back in March, one idea that's been floating around is that better drug molecules tend to have more of an enthalpic contribution to their binding."

Is it possible to have an argument that could be more hand-waving than this? It seems like today medicinal chemists are reaching for any argument to predict activity other than actually making the molecules.

Thermodynamics deals with the bulk properties of molecules. You cannot use thermodynamic equations to describe one-on-one interactions of molecules, because that occurs on a quantum level. If someone wants to use quantum calculations based on the interactions of the standing-wave models of the electron "clouds", I'm all for it. There is no such thing as delta-H, delta-S, or delta-G in this situation, so please stop using them.

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10. Derek Lowe on October 27, 2010 6:53 AM writes...

Azetidine, first off, you're missing the point. The idea is to find some way to distinguish between the molecules that are already made. We're not talking calculations here. If you have two one-micromolar lead compounds, and it turns out that one of them has a lot better chance of being optimized than the other (because of its thermodynamic profile), then wouldn't you want to know that? As I said, there's some doubt about whether that prediction can be made, but if it can, we need to know.

Second, you're incorrect, as far as I know, in your statement that there's no such thing as Gibbs free energy for a single molecular interaction. Are you saying that there's no heat given off when a single molecule binds? That there's no change in entropy when a single water molecule is displaced, or a single ligand stops freely rotating? We're not measuring single molecules (yet), but we're measuing a lot of one-on-one interactions simultaneously, and a thermodynamic approach is completely appropriate.

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11. Lionel C. (France) on October 27, 2010 7:20 AM writes...

I would like to know, for whose that are using ITC, if they have one ITC in their company (lab), if they obtain ITC value through collaborations, or if you send your compounds to external companies (if this kind of "companies for services" exist; it exist?).

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12. Willia_A_Nelson on October 27, 2010 8:54 AM writes...

Most of the comments here remind me of how poorly Physical Chemistry is both taught and "learned". A little background reading here (see link below) might add some insight to this somewhat diffuse commentary discussion.

http://www.chem.missouri.edu/Tannerjj/bchem/fbdd-pdfs/DeAevedoDias-thermoptn-druginteract.pdf

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13. processchemist on October 27, 2010 9:14 AM writes...

@azetidine

Search on wiki "stastical mechanics" and you'll find how the classical H and G can be derived by the behavior of a set of particles.

The problem here is that, thermodinamically, crystals of proteins and ligands and proteins-ligand complex in a cell are probably quite different objects...

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14. Wavefunction on October 27, 2010 9:24 AM writes...

-Thermodynamics deals with the bulk properties of molecules

Yes, but that's what statistical thermodynamics is for, to relate bulk thermodynamic quantities to microscopic situations. Delta H and delta G undoubtedly operate on a molecular scale, and there are theories to describe how much average entropy you would get for instance from displacing a single water molecule from a hydrophobic protein active site.

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15. azetidine on October 27, 2010 9:32 AM writes...

Derek and others, you make very good points. It certainly is true that there is heat and entropy exchange during single-molecule interactions. But: the macroscopic determinations of delta-G are an average of all of the 10^20 (for a mmol) single interactions. We have no way of knowing where on the Boltzmann distribution the productive interactions lie; they may be close to the average but they also may be far from the average. Furthermore, there is no way of knowing whether the productive single-molecule delta-G tracks with the average. The only way to really know would be to do (impossible) quantum calculations of the electron-standing-wave interactions.

My larger point is that ever since medicinal chemistry started, the hit rate for synthesized compounds has been very low and will always be low. That's just the nature of the beast, and not something for individual chemists to feel bad about (although clueless managers think otherwise). The result has been to generate lots of reasons to make one compound over another, so you can have an "excuse" to defend why you made that compound. Unfortunately, the predictive nature of all these paradigms is worse than poor. An enlightened manager would throw all of these away and let the chemists do what they do best and slog their way toward the lead. Certainly all this technology has not increased the rate of new drug approvals.

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16. DLIB on October 27, 2010 11:08 AM writes...

Agree with Willia!!!! I also think university O chem lab courses might benefit to some exposure to calorimetry -- DSC at least and ITC if possible -- that way it won't be so foreign to them when they graduate and magically become med chemists. Sadly the toolset that people get comfortable with is set early on.

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17. Paul on February 11, 2011 2:35 AM writes...

I've run some isothermal calorimetry measurements myself, so I can say it's pretty great for biophysical measurements. It DOES tell you the stoichiometry of the binding, particularly if the first binding constant is much stronger than the second/third etc. You essentially get points for a titration curve, then fit the curve model for one-site or two-site binding models. I've yet to try my hand at SPR, but from the critiques and reviews it looks very useful as well.

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