| Image may be NSFW. Clik here to view.  |
| The classic potential energy curve of chemistry hides a fundamental truth: bonds mean short distances, but short distances don't mean bonds |
Every field has its set of great philosophical dilemmas. For physics it may be the origin of the fundamental constants of nature, for biology it might be the generation of complexity by random processes. Just like physics and biology chemistry operates on both grand and local scales, but the scope of its fundamental philosophical dilemmas sometimes manifests itself in the simplest of observations.
For me the greatest philosophical dilemma in chemistry is the following: It is the near impossibility of doing controlled experiments on the molecular level. Other fields also suffer from this problem, but I am constantly struck by how directly one encounters it in chemistry.
Let me provide some background here. Much of chemistry is about understanding the fundamental forces that operate within and between molecules. These forces come in different flavors: strong covalent bonds, weak and strong hydrogen bonds, electrostatic interactions, weak multipolar interactions, hydrophobic effects. The net interaction or repulsion between two molecules results from the sum total of these forces, some of which may be attractive and others might be repulsive. Harness these forces and you can control the structure, function and properties of molecules ranging from those used for solar capture to those used as breakthrough anticancer drugs.
Here’s how the fundamental dilemma manifests itself in the control of all these interactions: it is next to impossible to perform controlled experiments that would allow one to methodically vary one of the interactions and see its effect on the overall behavior of the molecule. In a nutshell, the interactions are all correlated, sometimes intimately so, and it can be impossible to change one without changing the other.
The fundamental dilemma is evident in many simple applications of chemistry. For instance, as someone involved in structure-based drug design on a daily basis, I am used to carefully looking at the x-ray crystal structures of small molecules bound to proteins of biological interest. These small molecules exploit many different interactions including hydrogen bonds, charge-charge interactions and hydrophobic effects to bring about a net lowering of their interaction energy with the protein. The lower this interaction or free energy the better the interaction. Unfortunately, while one can visualize the geometry of the various interactions, it is very difficult to say anything about their energies, for to do so would entail varying an interaction individually and looking at its effects on the net energy. Crystal structures thus can be very misleading when it comes to making a statement about how tightly a small molecule binds to a protein.
Let’s say I am interested in knowing how important a particular hydrogen bond with an amide in the small molecule is. What I could do would be to replace the amide with a non hydrogen-bonding group and then look at the affinity, either computationally or experimentally. Unfortunately this change also impacts other properties of the molecules; its molecular weight, its hydrophobicity, its steric interactions with other molecules. Thus, changing a hydrogen bonding interaction also changes other interactions, so how can we then be sure that any change in the binding affinity came only from the loss of the hydrogen bond? The matter gets worse when we realize that we can’t even do this experimentally; in my colleague Peter Kenny’s words, an individual interaction between molecules such as a hydrogen bond is not really an experimental observable. What you see in an experiment is only the sum total, not the dissection into individual parts.
There have of course been studies on ‘model systems’ where the number of working parts is far less than those in protein-bound small molecules, and from these model systems we have gotten a good sense of the energies of typical hydrogen bonds, but how reliably can we extend the results of these systems to the particular complex system that we are studying? Some of that extrapolation has to be a matter of faith. Also, model systems usually provide a ranges of energies rather than a single value (say from 2-5 kcal/mol for a hydrogen bond) and we know that even a change of 1.8 kcal/mol can correspond to a substantial 10 fold change in binding affinity, so the margin of error entrusted to us is slim indeed.
It is therefore very hard, if not impossible, to pin down a change in binding affinity resulting from a single kind of interaction with any certainty, because changing a single interaction potentially changes all interactions; it is impossible to perform the truly controlled experiment. Sometimes these changes in other interactions can be tiny and we may get lucky, but the tragedy is that we can’t even calculate with the kind of accuracy we would like, what these tiny increments or reductions might be. The total perturbation of a molecule’s various interactions remains a known unknown.
This inability to perform the truly controlled experiment – a device that lies at the very foundations of modern science – is what I call the great philosophical dilemma of chemistry. The dilemma not only makes the practical estimation of individual interactions very hard but it leads to something even more damning: the ability to even call an interaction an 'interaction' or 'bond' in the first place. This point was recently driven home to me through an essay penned by one of the grand old men of chemistry and crystallography – Jack Dunitz. Dunitz’s point is that we are often misled by ‘short’ distances observed in crystal structures. We ascribe these distances to ‘attractive interactions’ and even ‘bonds’ when there is little evidence that these distances are actually attractive.
Let’s backtrack a bit to fundamentals. The idea of ascribing a short distance to an attractive interaction comes from the classic van der Waals potential energy curve that is familiar to anyone who has taken a college chemistry class. The minimum of this curve corresponds to both the shortest distance (called the van der Waals distance) between two molecules and the lowest energy, typically taken to signify a bond. However this leads to a false equivalence that seems to flow both ways: van der Waals distances correspond to bonds and bonds correspond to van der Waals distances.
In reality the connection only flows one way. Bonds do correspond to short distances but short distances do not necessarily correspond to bonds. So then why do we observe short distances in molecules in the first place? Again, Dunitz said it very succinctly in a previous review: simply because ‘Atoms have to go somewhere’. The fact is that a crystal structure is the net result of a complex symphony of attractive and repulsive interactions, a game of energetic musical chairs if you will. At the end, when the dust has settled everyone has to find a chair, even if it means that two people might end up uncomfortably seated on the same chair. Thus, when you see a short distance between two atoms in a crystal, it does not mean at all that the interaction between them is attractive. It could simply mean that other interactions between other atoms are attractive and that those two atoms have simply then settled where they find a place, even if the interaction between them may be repulsive.
How repulsive can it be? Dunitz gives the example of a carboxylic acid crystal where two oxygens have settled next to each other within van der Waals distance but whose interaction with each other is predicted to be repulsive to the order of thousands of kcal/mol (for reference, energies of typical covalent bonds are usually in the dozens of kcal/mol). The reason the crystal does not blow itself apart in spite of this interaction is of course because the other interactions make up for the repulsion.
How repulsive can it be? Dunitz gives the example of a carboxylic acid crystal where two oxygens have settled next to each other within van der Waals distance but whose interaction with each other is predicted to be repulsive to the order of thousands of kcal/mol (for reference, energies of typical covalent bonds are usually in the dozens of kcal/mol). The reason the crystal does not blow itself apart in spite of this interaction is of course because the other interactions make up for the repulsion.
The message here is clear: it is folly to describe an interaction as ‘attractive’ simply because the distance is short. This applies especially to weaker interactions like stacking interactions between aromatic rings. I am always wary when I see a benzene ring from a small molecule nicely sandwiched between a benzene ring in a protein and hear the short distance between the two described as a ‘stacking interaction’. Does that mean there is actually an attractive stacking interaction between the two? Perhaps, but maybe it means simply that there was no other place for the benzene ring to be. How could I test my hypothesis? Well, I know that varying the substituents on benzene rings is known to vary their energies of interaction with other benzene rings. So I ask the chemist to make some substituted versions of that benzene ring. But hold on! Based on the previous discussion, I just remembered that varying the substituents is not going to just change the stacking energy; it’s also going to change other qualities of the ring that mess up the other interactions in the system. It’s that problem with performing controlled experiments all over again - welcome to the fundamental dilemma of chemistry.
The fundamental dilemma is why it is so hard to understand individual interactions in chemical systems, let alone exploit them for scientific or commercial gain. We see it in a myriad of chemical experiments, from investigating the effects of structural changes on the rates of simple chemical reactions to investigating the effects of structural changes on the metabolism of a drug. We can’t change one component without changing every other component. There may be cases where these other changes might be minuscule, but in reality the belief that they may be minuscule in a particular case will always remain a matter of faith than of fact.
The fundamental dilemma then is why drug design, materials designs and every other kind of molecular design in chemistry is so tricky. In a nutshell, it’s why chemists are always ignorant and why chemistry is therefore always interesting.
Image may be NSFW.The fundamental dilemma then is why drug design, materials designs and every other kind of molecular design in chemistry is so tricky. In a nutshell, it’s why chemists are always ignorant and why chemistry is therefore always interesting.
Clik here to view.