I've been involved in a mailing list discussion that I wanted to open up to a wider audience in drug discovery, so here goes. We spend our time (well, a lot of it, when we're not filling out forms) trying to get compound to bind well to our targets. And that binding is, of course, all about energy: the lower the overall energy of the system when your compound binds, relative to the starting state, the tighter the binding.
That energy change can be broken down (all can all chemical free energy changes) into an enthalpic part and an entropic part (that latter one depends on temperature, but we'll assume that everything's being done at a constant T and ignore that part). Roughly speaking, the enthalpic component is where you see effects of hydrogen bonds, pi-pi stacking, and other such "productive" interactions, and the entropic part is where you're pushing water molecules and side chains around - hydrophobic interactions and such.
That's a gross oversimplification, but it's a place to start. It's important to remember that these things are all tangled together in most cases. If you come in with a drug molecule and displace a water molecule that was well-attached to your binding pocket, you've broken some hydrogen bonds - for which you'll pay in enthalpy. But you may well have formed some, too, to your molecule - so you'll get some enthalpy term back. And by taking a bound water and setting it free, you'll pick up some good entropy change, too. But not all waters are so tightly bound - there are a few cases where they're actually at a lower entropy state in a protein pocket then they are out in solution, so displacing one of those actually hurts you in entropy. Hmm.
And as I mentioned here, you have the motion of your drug molecule to consider. If it goes from freely rotating to stuck when it binds (as it may well), then you're paying entropy costs. (That's one reason why tying down your structure into a ring can help so dramatically, when it helps at all). And don't forget the motion of the protein overall - if it's been flopping around until it folds over and clenches down on your molecule, there's another entropy penalty for you, which you'd better be able to make up in enthalpy. And so on.
There's been a proposal, spread most vigorously by Ernesto Freire of Johns Hopkins, that drug researchers should use calorimetry to pick compounds that have the biggest fraction of their binding from enthalpic interactions. (That used to be a terrible pain to do, but recent instruments have made it much more feasible). His contention is that the "best in class" drugs in long-lived therapeutic categories tend to move in that direction, and that we can use this earlier in our decision-making process. People doing fragment-based drug discovery are also urged to start with enthalpically-biased fragments, so that the drug candidate that grows out from them will have a better chance of ending up in the same category.
One possible reason for all this is that drugs that get most of their binding from sheer greasiness, fleeing the water to dive into a protein's sheltering cave, might not be so picky about which cave they pick. There's a persistent belief, which I think is correct, that very hydrophobic compounds tend to have tox problems, because they're often just not selective enough about where they bind. And then they tend to get metabolized and chewed up more, too, which adds to the problem.
And all that's fine. . .except for one thing: is anyone actually doing this? That's the question that came up recently, and (so far), for what it's worth, no one's willing to speak up and say that they are. Perhaps all this is a new enough consideration that all the work is still under wraps. But it will be interesting to see if it holds up or not. We need all the help we can get in drug discovery, so if this is real, then it's welcome. But we also don't need to run more assays that only confuse things, either, so it would be worth knowing if drug-candidate calorimetry falls into that roomy category, too. Opinions?