Field of Science

DMSO to SAHA: Questions of toxicity

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I alluded in the last post to Breslow and Friend's discovery of DMSO as a HDAC (histone deacetylase) targeting agent. The story of how such a simple molecule was made into a putative drug deserves a closer look. The Nature Biotechnology paper that was cited in the last post contains most of the details. Suffice it to say here that the following points were noted.

DMSO was postulated to bind to a metal ion. Since Zn-heteroatom interactions are well known, another known sulfur binding moiety (look here for Whistling's list of zinc binding organic motifs), hydroxamic acid, was modified after some SAR studies to give the end molecule. One of the interesting steps in the lead optimization was the simple observation that it was unlikely that there were two metal ions in the same active site which bound to two hydroxamic acid groupings. So instead, one hydroxamic acid grouping was converted into a phenyl ring to gain binding affinity from hydrophobic binding. As the picture above shows, the drug now nicely fits into the active site, with the expected metal-hydroxamic acid interaction.

A journal club meeting member realised that SAHA's structure was similar to that of TSA (trichostatin A shown above), which was known to modulate HDACs. Thus came the hypothesis, later validated, that SAHA binds to HDACs. The details however are not yet clear, as HDAC interaction leads to a cascade of different events, and there are a dozen HDACs which are crucial for transcription.

There is one paragraph in the paper however, which I don't completely agree with. The authors were testing a few analogs, and came up with a branched structure with two hydrophobic rings that was much more potent than SAHA. However, it was later found to be quite toxic. While the authors make a good point that increased potency (increased binding affinity) can translate into greater toxicity because of continuous target activation, they also say:
"SAHA has hit the happy medium. It is potent enough to be useful and tolerated in patients. If the dosing is intermittent, such as not to maintain a continuous ‘therapeutic’ level of SAHA, it can be released from the binding site periodically so as to allow the deacetylation activities in cells. This is a general consideration, which could well be true of many other medicinal compounds. Thus, it is probably a mistake for medicinal chemists to set out first to find the most potent compound they can achieve in a series and then to look at any question of toxicity, as is often done."
The first thought that came to my mind after reading this was; this can't be quite right. Aren't some of the best-selling drugs of all time both highly potent and not particularly toxic? Toxicity and potency cannot always coincide.

I think there are two issues here. First is the fact that 'toxicity is in the eye of the beholder'. Cancer drugs are some of the most toxic drugs ever produced, but they are tolerated because of their essential requirement. So toxicity is a relative phenomenon whose impact has to be judged based on the disease. Secondly, even for very potent compounds, toxicity will depend on the exact nature of the target modulated, and its role in life processes. There may be a target like HDAC which plays a central role in the crucial process of transcription. You wouldn't want to keep this target activated/deactivated for too long. On the other hand, there might be another target which could be activated or deactivated for relatively longer periods because of its less sensitive nature and role in metabolism. So while the opinion of Breslow et al. is not wrong, it misses the other big part of the picture- the protein- on whose nature toxicity will depend too.

Least favourite NMR solvent, but a swell liquid nonetheless

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Coronene has a poll on what everyone's least favourite NMR solvent is. I have not used all the solvents on the list, but among the ones I have used, DMSO was probably my least favourite, as it also seems to be for others taking the poll.

In my case, the reason was specific. I needed to do NOESY spectroscopy on a degassed, sealed sample. And even ice kept freezing up the DMSO when we were connecting it to the vacuum. For CDCl3, we could happily use lig. N2 and be happy, but not so for DMSO; we had to patiently do the degassing for almost half an hour. In addition, DMSO is viscous and radically changes the tumbling of molecules, causes aggregation and hydrogen bonding etc. Needless to say, one has to re-experiment with his/her NOESY parameters.

On the other hand, it is instructive to see what a great molecule DMSO is for biological applications. It can be used as a drug-delivery vehicle, although this use is debated. DMSO is widely used as a cryo-protectant, and the serine protease inhibitor PMSF is dissolved in DMSO as it will naturally get inactivated in water.

Ronald Breslow and the late Charlotte Friend also discovered that this simple molecule can induce cell differentiation and later, it was found to be a HDAC (histone deacetylase) inhibitor, a class of molecules that's heavily researched these days. Interaction with HDACs intimately interferes with transcription. For example, the recent celebrity 'wine molecule' resveratrol is purported to increase expression of the HDAC protein SIRT1, possibly affecting longevity.
Incidentally, Breslow's latest paper is an article in Nature Biotechnology, in which DMSO is used as the starting lead to develop SAHA, a HDAC inhibitor which is now in Phase 2 clinical trials as an anticancer agent.

DMSO also "enters the cell and kills the herpes virus". Apparently, this is an 'indisputable fact', and seems to be the basis of scented DMSO creams as the picture above shows. More interesting talk for the cocktail party; guys, beware of the DMSO girl.

A remarkably simple molecule with some remarkable (and occasionally annoying) properties, this DMSO.

The passing of Gossypium extraordinarium

Whenever I hear the name "cotton", the first thing that comes to my mind is "...and wilkinson". The second thing, apart from the cotton we all love to wear, is the dreaded group theory, for which Cotton's again was the pioneering text.

I could never really make sense of group theory. But today, the author of my (and millions of others') favourite advanced inorganic textbook has passed into history, and I regret his loss along with everyone else.

"Cotton and Wilkinson" has the same relation to its discipline as "Samuelson" and "The Feynman Lectures on Physics" have to theirs. Cotton was a truly prolific chemist. I am not qualified to judge whether his discovery of metal-metal quadruple bonds merits a Nobel or not, but I suspect that he would more than deserve a Coreyesque lifetime achievement prize. His output, honours, and contributions to research and teaching are immense. F. Albert Cotton, RA (Rest Assured) rather than RIP. Your legacy will live. My dog-eared copy of "C & W" and thousands of others bear witness to that.

P.S. Since we are on the topic, I have to say that my all-time favourite descriptive introductory inorganic text is J D Lee's "Concise Inorganic Chemistry", an affair that dates back to early college days. The book is so engaging and informative that it makes for great bedtime reading. For a more advanced treatment, I also love Huheey.

All transition metals exist only in +1 state (or a close approximation thereof)

Yes, I know; reading this statement might make people think that megagallons of ethanol are coursing through my arteries. It sounds quite counterintuitive. But I have become convinced that this is the case. I have not explored the matter in exhaustive detail, and some questions still linger. The explanation can get long-winded and I will reserve it for sometime when my own thoughts are clearer. But here's the summary.

There is a very simple reason why all transition metals will prefer only the +1 state in their complexes; Pauling's electroneutrality principle.
More simply put, consider this; won't it be improbable for a charge of +2 or +3 to accumulate on a metal ion? There will be repulsive destabilization. We know that charge tries to disperse itself and stabilize any system. So this positive charge will try to redistribute itself and get delocalized, thereby making the system close to electrostatically neutral. Where can it do this? Well, we don't want to forget 'the other half'- the ligands. What this means is that it's actually the ligands that are getting oxidised by helping to delocalize the positive charge. So even in a compound like AuCl3, the Au has a +1 charge, and it's really the chlorides that are getting oxidised.

The whole root of the problem is that chemists' definition of oxidation state actually is quite arbitrary, based on qualitative notions of electronegativity. We think that since chlorine is more electronegative than gold, it will carry all the negative charge. But Pauling's electroneutrality principle tells us otherwise.

However, this is not something we usually think about. Consider the simple ammonium cation. Organic chemists depict it with the positive charge on nitrogen. But it's a well-known fact that the positive charge in fact is distributed over the hydrogens. Any system will try to stabilize itself by dispersing charge as much as it can. That's what the ammonium cation does, nicely dispersing the positive charge over the hydrogens.

We might say; but how can halogens get oxidised? They are supposed to be among the most strongly oxidising elements. But note, that this statement applies to diatomic halogen molecules, not to halogen ions in solution.

You can take my word that even the fiercely electronegative CF3- anion in the complex [Cu(CF3)4]- is oxidised, and that the copper exists as Cu(I) in this state. Even a look at the formula suggests otherwise; the Cu must be existing as Cu(III). But you don't have to take my word for it. Consider this publication in Angewandte Chemie. There was also a small debate, published in another issue of the same journal, which further combs some fine points and definitions and clarifies things. High level calculations strongly suggest that the Cu exists as Cu(I).

High level calculations, but more logically, the electroneutrality principle and basic concepts of electrostatics.

The point is, many things in chemistry are defined either arbitrarily, or with respect to some standard. Based on the existence of another chemical species, many things can change for a given species. Given the right conditions, almost anything can get oxidised or reduced. Given the right conditions, anything can be made stable. After all, nobody had thought that stable carbocations could exist. But George Olah opened up an entire new field of carbocation chemistry with his pioneering work. In chemistry, there are no absolutes. Which again proves why it is so interesting.

Actually, it's simple to think about. Usually we assume that the metal is changing its oxidation state. Instead of that, let's adopt the other perspective; the metal makes the ligand change its oxidation state. The result is the same. We still see AuCl3 in solution. But it's really the Cl- that has changed, and not the Au.

Some questions still linger; then why do all these metals seem to exist in different oxidation states. More importantly, how can one prove this contention by experiment? After all, ours is an experimental science.

But even in the absence of definite experiments, if one ad hoc concept is taken for granted, then another ad hoc concept surely cannot be worse. Especially if it sounds more logical and convincing. Like all metals existing only in +1 oxidation state (or something close to that).

Important addendum: I should mention that when I say +1, the actual number is better defined as "close to +1". Calculations show that it can change from +0.7 to about +1.2 (this bears on the question d_orbital asked in the comments) depending on the ligand environment. Thus, in hemoglobin for example, in the oxidised and reduced form, Fe won't exist as exactly +1 both of the times. Its oxidation state will fluctuate between decimal places close to one.
And, as Mitch quite rightly says, in most of our daily life, metals have an unequivocal oxidation state of 0!!

Nature adores a vacuum?

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Well, ok, nature does not actually adore a vacuum, but what are we then to make of Bruce Berne's newest PNAS piece, in which he talks about water molecules in protein active sites that are so horribly unhappy that they would almost let a vacuum take their place?

The article comes on the heels of Friesner et al.'s comprehensive JMC 06 Glide XP article, which was an impressively comprehensive instance of how computational chemists can take into consideration the detailed physical chemistry of protein-ligand/drug interactions. As usual, electrostatics, van der Waals interactions, and desolvation among other things are taken care of. As usual, it turns out that in most earlier studies, the 'other half'-water-had not been considered. No wonder attempts to explain diabolically strong interactions like that of streptavidin-biotin (the strongest protein-ligand interaction known) resulted in drastic underestimates of the binding energy. With their latest framework, Berne et al. seek to remedy this situation. They make a couple of elementary sounding but oft neglected and important points:

1. Water is not going to be happy in a place in a protein where it is surrounded by a hydrophobic enclosure.

2. It is not going to be happy if it cannot make its full complement of hydrogen bonds with the protein/with other water molecules.

3. Even if it can make such bonds, it is not going to be happy if there are only one or two configurations in which it can do so.

4. Clearly, such water molecules are going to be more than happy to be kicked out by ligand atoms which can form their full complement of h-bonds with protein residues. These bonds, if they form part of a ring or if they are in close vicinity, have been called correlated hydrogen bonds. As per the authors' interpretation, its is these correlated hydrogen bonds that give special binding stability to examples like the streptavidin-biotin complex.

Obviously, like all life forms, water aspires to that highest ideal- freedom.

In the PNAS article, Berne et al. examine three cases, including streptavidin and COX 2 binding, in which water seems to be simply aching to get of the protein cavity. In the case of streptavidin, five water molecules form a metastable ring as depicted above. They do manage to form bonds with each other and with protein residues, but they are hopelessly trapped on top and bottom by hydrophobic parts. In this case, as explained in the JMC piece, the atom-atom pairwise addition principle does not work for hydrophobic enclosures. Clearly, the sum of the parts is not equal to the whole here, and the effect of two hydrophobic atoms, one each on top and bottom, is different from the calculated pairwise sum of effects for two atoms .

The word "metastable" is the best word I can think of for describing this unhappy situation for water. Surreal, simmering, patient, but waiting to become free.

In the case of COX 2, I saw a few statements which neither me nor my advisor could make sense of at first. Aristotle notwithstanding, it seems that there are actually some solvation situations which are almost as bad as a vacuum. Or even worse? Consider what the authors have to say:
The Cox-2 active site was found to contain no persistent hydration sites and is in fact entirely devoid of solvent in 80% of the simulation, despite the cavity sterically accommodating approximately seven water molecules. The high excess chemical potential of the binding-cavity solvent is due to an inability of the water molecules to make hydrogen bonds with the surrounding hydrophobic protein residues and other water molecules...An extreme case of hydrophobic enclosure is observed in the Cox-2 binding cavity, where no energetically stable solvent configurations appear to exist; insertion of ligand hydrophobic groups into such a region of persistent vacuum will result in substantially larger free-energy liberation than would be expected if the binding cavity were treated as solvated.

Well...I second the study; water, that humble but amazingly complex liquid, has often been neglected and underestimated in studies of protein-drug-ligand interactions.

Almost concomitantly, Essex et al. have published a report in JACS, in which they examined the nature of water molecules in active sites. They find that most of the water molecules are displaceable. Not surprisingly, they conclude that the binding affinity of the ligand depends upon the nature and environment of the water molecules. It's a good statistical study.

A central point that is coming to light through such studies is that the absolute hydrophobicity/hydrophilicity of a ligand certainly does not constitute the whole story. This may well be kept in mind by medicinal chemists who aim at modulating these properties of drug-like molecules from SAR studies. Sometimes, a single well-positioned ligand atom which displaces one or two extremely metastable water molecules can lead to orders of magnitude of binding affinity, which cannot be predicted from absolute characteristics of the ligand.

As usual, quality, not quantity, matters.

Another nice 3 page review I found in the rarely read Chemical Biology and Drug Design again makes the case. The review explores the thermodynamic aspects of drug-protein binding, and says among other things:
As the major contributors to the binding enthalpy are polar groups, a common misconception is that enthalpically driven compounds must be highly polar and that consequently their bioavailability will be compromised. In fact, what is often observed experimentally is that compounds with the same number of polar groups have vastly different binding enthalpies...To generate a favorable binding enthalpy, it is not the number of polar groups that matters but the quality of their interactions with the target. It is better to have few groups that establish strong interactions than a large number of groups mostly paying the desolvation penalty.

Again, by a "few groups", all that is needed is two functional groups that can kick out a metastable water molecule, and your job might be done.

CETP resolved

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It is a bit ironic that a group from the same company that saw its drug sink has now come up with a nice crystal structure of the protein to which the drug binds. (Nat. Struc. Mol. Biol., 14, 2, 2007, 106-113)

CETP (Cholesteryl Ester Binding Protein) was resolved by the group at 2.2 A. It contains a long hydrphobic tunnel (~60 A) which sequesters its hydrophobic cargo and shields it from the aqueous environment of the blood. The entrances of the tunnel are blocked by the polar head group of a lipid- oleaylphosphatidyl choline. Remarkably, the group managed to crystallize the protein with four different molecules, all of which are important in lipid metabolism- cholesteryl oleate, triglycerides, and phoshatidyl choline. The overall shape of the protein is like a boomerang, and the authors surmise that the shape could help to latch on to an HDL particle.

It is sad that Torcetrapid failed. The binding of the drug had been intensively studied, and its potency had been evaluated with a high degree of confidence. In the end, the drug sank in the one foggy area where human prediction still fails- toxicity and side effects. The company has redeemed itself to some extent with this structure and hopefully, the structure will lead them to a safer version or analog of torcetrapib. Till then, it does no harm to study the structure for the sake of beauty alone.