A bond by any other name...: How the simple definition of a hydrogen bond gives us a glimpse into the heart of chemistry

Basic hydrogen bonding between two water molecules,
with the central hydrogen shared between two oxygens

A few years ago, a committee organized by the International Union of Pure and Applied Chemistry (IUPAC) - the international body of chemists that defines chemical terms and guides the lexicon of the field - met to debate and finalize the precise definition of a hydrogen bond.

Defining a hydrogen bond in the year 2011? That was odd to say the least. Hydrogen bonds have been known for at least seventy years now. It was in the 1940s that Linus Pauling defined them as foundational elements of protein structure; the glue that holds molecules including life-giving biological molecules like proteins and DNA together. Water molecules form hydrogen bonds with each other, and this feature accounts for water's unique properties. Whether it's sculpting the shape and form of DNA, governing the properties of materials or orchestrating the delicate dance of biochemistry performed by enzymes, these interactions are essential. Simply put, without hydrogen bonds, life would be impossible. No wonder that they have been extensively studied in hundreds of thousands of molecular structures since Pauling highlighted them in the 1940s. Today no chemistry textbook would be considered legitimate without a section on hydrogen bonding. The concept of a hydrogen bond has become as familiar and important to a chemist as the concept of an electromagnetic wave or pressure is to a physicist. 

What the devil, then, were chemists doing defining them in 2011?

It turns out that behind the story of defining hydrogen bonds lies a paradigm that takes us into the very guts of the nature and philosophy of chemistry. It leads us to ask what a chemical bond is in the first place; drill down deep enough into it, and it starts appearing like one of those existential questions a Zen monk would ask about life or the soul.

Remarkably enough, in spite of the debate and redefinitions, the basic features of a hydrogen bond are not controversial. A hydrogen bond is a bond that exists between a hydrogen and two other non-hydrogen atoms. Those two other atoms are most commonly oxygen and nitrogen, but that's where the consensus seems to end and the endless debate begins.

The basic reason for the debate has to do with the very definition of a bond as laid out by pioneering chemists like Gilbert Newton Lewis, Irving Langmuir and Linus Pauling in the 1920s. Loosely speaking a bond is a force of attraction between two atoms. It was Lewis who made the groundbreaking suggestion that a chemical bond results from the donation, acceptance or sharing of electrons. Pauling codified this definition into a rigorous principle using the laws of quantum mechanics. He showed that using quantum mechanics, you could prove that the sharing of electrons between two atoms leads to a net lowering of energy between them: this sounds logical since if there were no lowering of energy, there would be no incentive for two atoms to come close to each other.

So far so good. Everyone accepts this basic fact about chemical bonds. Treatises have been written about the various bonds between different atoms seen in a bewildering array of metallic, organic and inorganic molecules; Pauling's seminal book, "The Nature of the Chemical Bond", described such interactions in loving detail. In this book he showed that bonds can be covalent or ionic; the former involve a symmetric sharing of electrons while the latter involve an asymmetric give and take. There are also many intermediate cases. Furthermore, whether a bond is ionic or covalent depends on the electronegativity of the atoms involved. Electronegativity is the innate property of an atom to hold on closely to its electrons. Fluorine for instance is the most electronegative element of all, holding on to its electrons in a fiery embrace. Metals on the other hand are happy to give away their electrons. Not surprisingly, bonds between metals and fluorine are readily formed.

It is in the context of electronegativity that it's easy to understand hydrogen bonds. A hydrogen atom, just like a metal, is happy to give away its lone electron and form a positive ion. Oxygen and nitrogen atoms are happy to accept this electron. Thus, a hydrogen bond forms when a hydrogen atom is shared between two oxygen atoms (or one oxygen and one nitrogen, or two nitrogens). Because of the difference in electronegativity, the hydrogen acquires a slightly positive partial charge, and the other atoms acquire a slightly negative partial charge (represented in the picture above). Think of two oxygen atoms with a hydrogen between them acting as a bridge: that's a hydrogen bond between them. Because a hydrogen atom is small and has only a single electron it can achieve the feat of forming this bridge; what is fascinating is that this is partly the result of a unique quantum mechanical phenomenon called the tunnel effect.

All this is rather uncontroversial. The debate starts when we ask what criteria we use to actually identify such hydrogen bonds. Chemists usually identify bonds by looking at experimental structures of molecules. The most common method of acquiring these structures is by x-ray diffraction. Bouncing x-rays off a crystallized molecule allows us to accomplish the molecular equivalent of taking a photograph. The molecules under question can be anything from common salt to complicated proteins. These structures form a cornerstone of chemistry, and they are not just of academic importance but have also been immensely useful in the design of new drugs and materials.

When chemists acquire the crystal structure of a molecule, they will usually do something very simple: they will measure the distance between various atoms. The hallmark of a chemical bond is a short distance between two atoms; think of it as two people embracing each other and reducing the distance between themselves. How short, exactly? Short enough to be less than the sum of Van der Waals radii of the two atoms. Imagine the Van der Waals radius of an atom as a kind of safe perimeter which an atom uses to hold other atoms at bay. Each atom has a unique Van der Waals radius defining its safe perimeter which is measured in the unit angstroms (1 Å = 10^-10 meters). Hydrogen for instance has a radius of 1.2 Å, nitrogen has a radius of 1.5 Å. If a chemist measures the distance between a hydrogen and nitrogen in a crystal structure as being less than 2.7 Å (1.2 + 1.5), she would usually declare that there's a bond between the two atoms. If the three atoms in a hydrogen bond satisfy this geometric criteria, then one could legitimately claim a bond between them.

But this simple procedure turns out to be far trickier to accept than we can imagine. Here's the problem: a crystallized molecule is a complex assembly of hundreds or even thousands of atoms, all of which are jostling for space. The eventual formation of the crystal is not a zero-sum game; some atoms have to give way so that others can be happy. As the pioneering chemist Jack Dunitz has succinctly noted, "Atoms have to go somewhere". Thus, some atoms are simply forced together against their will, just like people in a cramped train compartment who are forced together against their will. Now, just like those cramped people, the distance between two such cramped atoms will be unnaturally short, often shorter than the the sum of their Van der Waals radii. The key term here is "unnaturally": the distance is short not because the atoms actually want to be close, but because they are forced to be so. Can one then say that there is a bond between them? And how short does this distance need to be in order to be called a bond?

There's another conundrum that emerged as thousands of crystal structures of molecules started to appear during the last few decades. While the "classical" hydrogen atoms bridging oxygens and nitrogens were well known, short distances also started appearing between hydrogen and atoms like carbon and chlorine. The purported hydrogen bond zoo started to get populated with even more exotic creatures. For instance chemists started noticing close contacts between hydrogens and the flat face of benzene rings: these hydrogens are attracted to the ghostly electron clouds lying on top of the rings. Sometimes you did not even need a benzene ring; two doubly bonded carbon atoms would appear as magnets for hydrogens and pull them close. And these hydrogens in turn again would not just be bonded to oxygens or nitrogens as their second partners; they could be bonded to carbons, or even to metallic elements. Experimental observation forced the concept of hydrogen bonds to be extended not just to other bonafide atoms but to more abstract entities like clouds of electrons.

An illustration of two kinds of hydrogen bonds:
conventional ones with oxygen atoms in blue
and those with a benzene ring in red

Two overriding questions thus emerged from these decades-long exercise of observation and analysis: Can one claim a bond between two atoms in a molecule - specifically a hydrogen bond in case of hydrogen, oxygen or nitrogen - simply based on the distance between them? And what atoms exactly constitute a hydrogen bond?

The IUPAC committee in 2011 seems to have exorcised both demons with a single stroke. They defined a hydrogen bond as any kind of bond between a hydrogen and two other atoms, provided that the other two atoms are more electronegative than hydrogen itself. That always included oxygen and nitrogen, but now the definition was expanded to include carbon and halogens (chlorine, fluorine, bromine and iodine) as well as ghostly entities like electron clouds on top of benzene rings. As far as the distance criterion was concerned, IUPAC made a wise decision in allowing it to be fairly variable and not simply limited to a hard cutoff within the sum of Van der Waals radii.

The IUPAC decision illuminates three very important aspects of chemistry. Firstly, it tells us that many chemical concepts are fuzzy and only approximately defined. For instance, although we have defined the distance between the hydrogen and other two atoms in a hydrogen bond as being less than the sum of their Van der Waals radii, in reality these distances are best derived as a statistical distribution. Some hydrogen bonds are very short, some are very long, and most are in the middle: even this classification has been a hotbed of contention, and you will find arguments about the nature and energies of "short" and "weak" hydrogen bonds in the scientific literature that approach the vigor of gladiatorial contests. The ones that are short usually involve nitrogen and oxygen, while the long ones involve the elements included in the expansive IUPAC definition. This fuzzy reality extends to other concepts in chemistry such as aromaticity, polarizability and acid and base behavior. Unlike concepts in physics like wavelength and electrical conductivity, concepts in chemistry can be much more malleable. However this malleability is not a bug but a feature, since it allows you to capture a vast amount of chemical diversity in a central idea.

The second key aspect of chemistry that the IUPAC statement reveals is the fact that scientific concepts don't have to always be rigorously defined in order to be useful. The definition of a hydrogen bond goes to the heart of chemistry as a science based on models. A hydrogen bond is a bond, but it's really a model based on our ideas of the myriad ways in which molecules interact. The key feature of a model is utility, not reality. In my own work for instance, I can recognize and exploit hydrogen bonds between proteins and drugs and improve their directionality and strength, all without worrying one bit about their precise, rigorous definition. Similarly polymer chemists can create new forms of hydrogen-bonded plastics with novel functions without caring much whether their hydrogen bonds conform to some well-defined ideal. Whatever the debate about hydrogen bonds may be, my view of them is similar to Potter Stewart's: I know them when I see them. Or, if I wanted to offer a cheesy but still quotable line from one of "The Matrix" movies, a hydrogen bond is merely a phrase; what matters is the connection it implies.

Ultimately, this fuzzy yet useful view of hydrogen bonding makes an important statement about the philosophy of chemistry. It tells us that chemistry has its own philosophy based on models and utility that is independent of its roots in physics. As the chemist Roald Hoffmann memorably put it, chemical concepts like aromaticity and electronegativity start to "wilt at their edges" when examined too closely, and Hoffmann could have been talking about the hydrogen bond there. The physics-based view of a hydrogen bond would involve writing down the Schrodinger equation for these bonds and rigorously solving it to obtain quantities like the energy and geometry of the bond. But that exercise is too reductionist; it does not actually help me understand the nature and variety of these bonds. It won't tell me why water is liquid and why it flows from mountains, why it's a universal solvent and why it's an existential life-giver. Chemistry proclaims its own emergent language and philosophy, and while the material entities of which molecules are composed are grounded in physics, the molecules themselves belong squarely in the domain of chemistry.

The fact that chemists are still debating hydrogen bonds means that chemistry remains an exciting and vibrant discipline, one in which even basic concepts are still being sculpted and fine-tuned. It tells us that for the foreseeable future chemistry will continue to be a science without limits, a country with open borders. The hydrogen bond should give chemists an opportunity to celebrate the very soul of their work.

The fundamental philosophical dilemma of chemistry

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 (dictated by the sharing of electrons), hydrogen bonds (dictated by weak electrostatic interactions), strong charge-charge interactions (dictated by attraction between unlike charges), hydrophobic effects (dictated by the interaction between 'water-loving' and 'water-hating' parts of molecules) etc. 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, my day job involves looking at the crystal structures of proteins involved in disease and then designing small organic molecules which bind to and block such proteins. For binding to their target protein, 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 by simply looking at the crystal structure, 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 in the small molecule is. What I could do would be to replace the atoms comprising the hydrogen bond with non hydrogen-bonding atoms and then look at the change in the affinity of the resulting molecule for the protein, either computationally or experimentally. Unfortunately this change also impacts other properties of the molecules; its molecular weight, its hydrophobicity, its steric or spatial 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 and we know that even a tiny change in the energy of binding can correspond to a substantial loss of effective blocking of a protein, 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, a concept which has been at the heart of the scientific method. 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.

The roots of the problem run even deeper. At the most elemental level, all interactions between molecules are simply a function of one of the four fundamental forces known in nature - the electromagnetic force. Of the four basic forces, gravity is too weak to play a role, while the strong and weak nuclear forces don't usually apply to molecular interactions since such interactions only involve the sharing of electrons. It is the electromagnetic force that is thus ascendant in mediating every single molecular interaction in the universe. When we divide this force up into hydrogen bonds, electrostatic interactions, hydrophobic interactions etc. what we are doing is imposing an artificial division on an indivisible fundamental force, purely for our convenience. It's a bit like the parable of the blind men and the elephant - there is only one electromagnetic force, just like there is only one elephant, but each of us describing that force divides it up into multiple flavors. No wonder then that we are led astray when we think we are doing a controlled experiment, since whenever we think we are varying one flavor or another we are actually varying the same basic parameter and not its independent components. That is because there are no independent components in the true sense of the term.

This inability to perform the truly controlled experiment is thus 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 in the essay is that we are often misled by ‘short’ distances between atoms 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 (figure above) 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. 

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 those between aromatic (benzene) rings. I am always wary when I see a benzene ring from a small molecule nicely sandwiched between another 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 or groups of atoms attached to 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. 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. It is why so much of complicated chemistry is still trial and error, why observations on one system cannot be easily extrapolated to another, and why even supercomputers are not yet able to nail down the precise balance of forces that dictate the structure and function of specific molecules. In a nutshell, the fundamental dilemma is why chemists are always ignorant and why chemistry will therefore always be endlessly fascinating.

On Freeman Dyson, cadmium estimation and the joy of chemistry

Cadmium sulfide (the solution has a much
deeper, brilliant yellow color)

Freeman Dyson who is a hero of mine and who I have had the privilege of meeting at length several times is someone who has done a lot of interesting things during a long and fruitful life. Not only has he made brilliant, foundational contributions to many areas of mathematics, physics, engineering and astronomy but he has also written books showcasing elegant, wide-ranging, poetic prose.

If you look at Dyson's astonishingly diverse writings or talks it is easy to mistake him as a philosopher of science, a big picture thinker striding on the widest possible canvas.

But as he himself said to me during a long, one-on-one, intellectually sparkling lunch discussion last year, a lot of people think that talking about the big picture automatically makes one a philosopher. Dyson has indeed written about an amazing variety of big picture topics, from the origins of life to the colonization of space. But he also maintains that he has always been first and foremost a problem solver, someone who is much more interested in details than in grand theorizing. Whatever philosophy he manages to weave is built on a foundation of solving specific technical problems. This is partly evidenced by his work on highly technical engineering projects ranging from nuclear spaceships to nuclear reactors.

This quality would make Dyson quite comfortable in the company of chemists, since chemistry by its very nature is more a problem-solving discipline rather than a philosophical pursuit like cosmology or evolutionary biology. I was curious to know Dyson’s views on chemistry since while he has had thoughts on almost every imaginable topic, I could not really find anything by him on chemistry except for a review in Physics Today of Nobel Laureate Roald Hoffmann’s fine book “The Same and Not the Same”.

Then I remembered an excellent collection of interviews with Dyson that I found a few years ago on a website called Web of Stories. This website is a must-see for history of science enthusiasts. It features interviews with scores of famous scientists, humanists and artists from diverse disciplines. These interviews are real gems. The best thing about them is that they are long and detailed instead of two minute sound bytes; each interview lasts in total for several hours and covers most significant events in the interviewees’s life, so you get a real feel for the life and work of leading thinkers. In one of Dyson’s interviews I was delighted to find this:

“I was going to say about chemistry that Roald Hoffmann whom I got to know quite recently, who is a chemist who writes poetry and is a great character, he has the same attitude toward chemistry that I do. I mean it is the beauty of the details rather than any over-arching theory. In that way it’s very different from physics, and I had a taste for it. My taste is always more for the details than for the big picture…

…I learned chemistry from Christopher Longuet-Higgins who was already much more of an expert and more excited about chemistry than Eric James (Dyson’s chemistry teacher at Winchester College). And I remember Christopher bringing to Winchester some crystals of stannic iodide which he had made, which is the most marvelous stuff. It is a brilliant scarlet colour and it makes these beautiful scarlet crystals, and they’re also extremely heavy. If you have a little bottle full of it, it feels like lead. So that kind of chemistry I found delightful, just the sort of details of the actual stuff, rather than the theory that lay behind it.

I remember the joy when, here in Princeton, Willard Libby came on a visit once and brought along another little bottle of chemicals, which also was very heavy, and that was barium xenate, which was barium xenon oxide, which of course was an absolute revelation because nobody imagined that xenon could have compounds, being an inert gas. And it was sometime in the 1950s these compounds were discovered, and barium xenate is just such an ordinary stuff. It’s a sort of heavy white crystals which are completely stable, they don’t show any signs of anything strange and there it is. If you heat them up of course the xenon comes bubbling off…”

So there, I think Dyson would have felt right at home as a chemist. Chemists revel in solving particular problems. It’s not that they don’t care about general principles. Every once in a while someone like Linus Pauling or Alfred Werner comes along and ties together disparate threads into a neat superstructure. But the vast majority of chemical work consists of celebrating the diversity of structure and function inherent in natural and synthetic molecules. The point Dyson makes is an obvious and important one that is often lost among the clichéd caricatures of chemists bubbling frothy liquids and crystallizing colorful solids that one often sees in literature and cinema. But it’s precisely these bubbling liquids and colorful solids that endow chemistry with a palpable reality that’s often missing from more theoretical sciences.

A personal digression. I remember an episode from an undergraduate chemistry lab where we were supposed to estimate two unknown metal ions from a solution. After trying out every test in the book we could only detect copper. The other ion remained a mystery and we finally threw up our hands. That’s when the instructor revealed his trick. It turns out that the method of copper estimation that we were using involved turning the solution highly acidic with hydrochloric acid. With a smile on his face, the instructor put a single drop of the concentrated acidic solution in a large flask and then filled the flask to the brim with water, diluting the initial solution by at least a factor of ten thousand. Our eyes were glued to the flask as he passed hydrogen sulfide gas into the solution. And then, starting from the bottom and rising to the top, the flask filled up with the most beautiful yellow color that I have ever seen; it’s a sight that I will never forget. What happens is that cadmium is precipitated as cadmium sulfide only in dilute acidic solutions while for copper it’s the opposite. All our tests for detecting copper in concentrated acid missed the hidden cadmium, until it was ready to be unmasked by simple dilution.

Every chemist is familiar with this feeling of discovering something unknown, no matter how trivial or important, that actually exists; all the better if it has a brilliant scarlet or full-throated yellow color, as is often the case in chemistry. Dyson is right that there is something unmistakably reassuring, in-your-face - real – about holding a vial of something that was previously considered impossible.

The great organic chemist Robert Burns Woodward spoke of this very real quality of chemistry that helped him choose between chemistry and mathematics when he was a teenager:

“The fact is that I have always been very fond of mathematics- for one short period, I even toyed with the possibility of abandoning chemistry in its favour. I enjoyed immensely both its conceptual and formal beauties, and the precision and elegance of its relationships and transformations. Why then did I not succumb to its charms? For two reasons, I believe:

FIRST, because by and large, mathematics lacks the sensuous elements which play so large a role in my attraction to chemistry. I love crystals, the beauty of their form- and their formation; liquids, dormant, distilling, sloshing!; the fumes; the odors- good and bad; the rainbow of colours; the gleaming vessels, of every size, shape and purpose. Much as I might think about chemistry, it would not exist for me without these physical, visual, tangible, sensuous things.

SECOND, while in mathematics, presumably one’s imagination may run riot without limit, in chemistry, one’s ideas, however beautiful, logical, elegant, imaginative they may be in their own right, are simply without value unless they are actually applicable to the one physical environment we have- in short, they are only good if they work! I personally very much enjoy the very special challenge which this physical restraint on fantasy presents”

Woodward’s second reason for studying chemistry goes to the heart of the nature of empirical science, what Richard Feynman called “imagination in a straitjacket”. It’s a philosophy that drives all of science including the most rarefied theoretical ideas. Physicists often like to tell the story of how Einstein felt that “something had snapped inside” him when he saw the predictions of general relativity confirmed by observations of the perihelion of mercury. He surely must have felt the rare, once-in-a-lifetime satisfaction of a great theoretical construct being validated by a real observation that could be boiled down to a single number. We tend to think of Einstein as the great scientist-philosopher, but there he was, being ecstatic about a technical detail that was a crucial part of his magnum opus. Observing barium xenate or precipitating cadmium is not quite as momentous as confirming the theory of relativity, but I can readily imagine feeling a shiver down my spine if I had been presented with that sort of chemical evidence. Evidence that seemed to defy the impossible but which I could nonetheless hold in my hand and keep in my closet. That’s the joy of chemistry.

This is an updated version of a previous post on Scientific American.

ChemCoach Carnival: What I do

I am late to the party, but SeeArrOh's ChemCoach Carnival has given me a chance to indulge in some narcissistic self-promotion. There many great entires on his blog so you should take a look. Here's my pitch.

Your current job.

I am an organic chemist turned molecular modeler at a small biotech startup in Cambridge, MA. I spend as much time looking at synthetic strategies, building block procurement, target selection and assays as I spend building models. I also spend a lot of time thinking about how my work fits within the broader boundaries of science.

What you do in a standard "work day."

As a lot of others scientists on this thread have emphasized, one of the great things about our job is that there is no “standard work day”. I am the lone modeler in a small startup so that requires me to wear several hats. I am as involved in discussing synthesis and assays as I am in docking small molecules to proteins or running molecular dynamics simulations. In addition I also need to occasionally look up building block availablity, talk to database and informatics specialists and arrange for presentations from outside vendors. The point is that in drug discovery and especially in a small outfit, you must be adaptable and be able to accomplish multiple, diverse tasks. This kind of capability makes you a valued member of the team, especially in a small company where your voice will be heard by everyone, from the intern to the CEO. It’s also a terrific learning experience in general.

What kind of schooling / training / experience helped you get there?

I have a doctorate in organic chemistry, although frankly that is just the means to an end. I switched from synthesis to modeling because I was clumsy in the lab and because I was interested in many different fields of science. I don’t regret my choice at all. Modeling allowed me to indulge interests in physics, chemistry, computer science and biology. I would say that if you have diverse scientific interests, modeling and simulation in a general sense would excellent career choices for you. If you are planning for a career in drug discovery or biotechnology, I would encourage you to soak up as much knowledge from diverse fields of chemistry and biology as possible. You won’t regret it.

This would also be the place for me to sneak in my favorite pitch regarding the history and philosophy of science. A simple piece of advice: study it. Science is done by fascinating human beings with all their flaws and triumphs. Your experiments are not being done in a vacuum. Reading up on the history of your discipline will give you the feeling of participating in a grand, unbroken thread of discovery going back to the Greeks. Even if you may not be a world-famous scientist or are not doing earth-shattering research, the simple fact that you are exploring the same laws of physics and chemistry that world-famous scientists once did will put you in their league and inspire a feeling of kinship. Studying the history of science will convince you that there are many who empathize and who have shared the same sense of despair and triumph that you do. Study the history of science, and you will know that you are not alone.

How does chemistry inform your work?

It is all-pervasive in my work. When I say I am a “molecular modeler”, I mean that in the broadest sense of the term. For me all of chemistry is largely about models, whether the models consist of structures scribbled on a hood or three-dimensional protein images built on a computer screen. A lot of people think computation in drug discovery is all about building regression models and writing fancy algorithms. But what it’s really about is data interpretation, and pretty much all this data is chemical. I cannot stress how important it is for a molecular modeler to understand chemistry, especially organic and physical chemistry. What has turned out to be an intractable problem has often proved amenable to a solution found in the principles of basic organic chemistry. In addition you have to have a real feel for structure-activity relationships and the basic physiochemical properties of functional groups. Useful numbers from thermodynamics and kinetics should ideally roll off your tongue like French verbs. A knowledge of statistics is also important. My background in organic chemistry is much more important than any facility with programming or knowledge of particular software that I may have picked up on the way. Those things you can learn, but the bedrock of your work will always be chemistry, even when it's operating behind the scenes.

Finally, a unique, interesting, or funny anecdote about your career*

Well, when I said I was clumsy in the lab I was thinking about the time I actually dropped a rotavap on the floor and broke it. My advisor, a generous man, said that maybe I was not quite cut out for working in the lab. It was the only time in my life that an embarrassing accident gently pointed out by a wise future advisor has fortuitously decided the trajectory of my career.

Rookie mistakes in molecular modeling: Part 1

Molecular modeling as a general approach is no longer utilized only by experts but has reached the masses. Improved hardware and software capabilities combined with easy-to-use graphical user interfaces have enabled experimental chemists of all kinds to build models of molecules and perform relatively sophisticated calculations on them. Calculations which once required supercomputers can now be routinely done on desktops by organic, inorganic and biological chemists who can use the results to explain, support and predict chemical phenomena. In the coming years we can be confident that we will witness an increasing use of modeling techniques by experimentalists.

An unfortunate (but probably not unexpected) consequence of this ease of use is that it has also become easier to make mistakes while building molecular structures. The main source of errors arises during the translation of 2D chemical structures to their 3D counterparts using some energy minimization protocol. The apparently simple process of drawing a 3D-worthy 2D structure is trickier than it sounds and is therefore quite prone to error. Conformation which was not as important when drawing in 2D is suddenly of overriding importance and it's relatively easy to get it wrong.

As a modeler who has interacted closely with experiment, I have come across a number of rookie mistakes which I have seen myself and others make over the years. Sometimes these mistakes don't matter too much for the final result but sometimes they can completely change it. So I thought I would make a short list of easy-to-avoid errors which may provide checks on modeling structures. In part 1 I will describe mistakes commonly seen during the simple building of structures. Part 2 will deal with interactions with experimentalists.

1. Getting the ionization state wrong: I put this rookie mistake at the top because it's remarkable how many times I have seen even experienced modelers make it. Always remember; amines are protonated at physiological pH while carboxylic acids are deprotonated. The reason why getting this right is important is because it can completely change results from protocols like docking. Just think of the difference a protonated vs unprotonated carboxylate makes for binding to a protein. Also, many modeling algorithms use force fields which are dominated by electrostatic interactions; the wrong protonation state can therefore make a world of difference. A corollary of the ionization state problem results when replacing atoms. For instance you may have a protonated amine which you then want to turn into an alcohol by replacing the N with a O. Unfortunately the atom does change but not the ionization state, and you end up with a weird positively charged doubly bonded oxygen. On a related note, it goes without saying that you shouldn't charge up inappropriate atoms such as those which are conjugated to aromatic systems. The best way to overcome these issues is to simply display charges for all heteroatoms in your final structure.

2. Getting the stereochemistry wrong: The CIP rules were taught to us because they really matter. Here's a typical stereochemical mistake: You construct a structure in 2D and come across a stereocenter. You may even build that stereocenter with the right absolute (R or S) stereochemistry. And then you attach something else to that center and forget to recheck the stereochemistry which may have changed because of the change in CIP priority. The simplest way to make sure about stereochemistry is to always have the program display all absolute stereochemistry for the final structure.

2. Forgetting basic conformational rules: This mistake is most commonly made when converting a 2D structure into a 3D structure. The problem is that when you build a 2D structure, your placement of bonds and angles is somewhat ad hoc based on the rather random way in which you are conveniently rotating and viewing the structure. When you then suddenly convert 2D to 3D, you may end up with axial substituents on six-membered rings, syn-pentane or eclipsing interactions between substituents, funky substructures like non-planar aromatic rings resulting from strain or in the worst cases, even boats for cyclohexanes. Here's another common pitfall: You may try to close a ring by building an unrealistic long bond between two initially separated atoms, thinking that when you then minimize this structure the program will take care of the bond by shortening it to its standard length. This usually happens, but in the process some other parts of your molecule gets messed up. Again, judicious inspection can avoid most of these issues.

3. Cis and trans: The process of building unrealistic bonds between distant atoms and then simply minimizing a structure that I just mentioned can sometimes result in amide bonds becoming cis and this is important enough to be listed as a separate point. This is also a common consequence of importing 2D files in SDF format (which lack hydrogens) and asking a program to add hydrogens. The same thing can happen with double bonds.

4. Forgetting basic chemistry: This mistake has more to do with forgetting basic rules of bonding and chemistry than with modeling. Occasionally you may do things like exceeding the allowed valency of an atom, putting a double bond at a bridgehead carbon (violating Bredt's rule), generating antiaromatic rings, forgetting Baldwin's rules for ring closure, building a vinyl amine or a geminal amino alcohol...and generally creating all sorts of unstable and "impossible" molecules. The problem is that your program won't always raise red flags notifying you about these errors so you need to remember your chemistry and make sure you don't recommend some wacky molecules to make to the synthetic chemist (one of the constant sources of friction between experimentalists and modelers arises from the latter forgetting what's synthetically feasible and stable).

Ultimately, the path to a well-constructed molecule simply depends on being vigilant and judiciously checking your final structure. Remember the well-worn adage; computers don't know any chemistry whatever and they are only as good as the code that goes into them. Nothing can trump a sound knowledge of basic chemical principles.

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Does bond cleavage require or release energy?

Back from a severe flu spell with a question worth pondering. A commenter on Derek's blog posed the following question, which I think could be a nice trick question for undergrads (or grads for that matter).

Question: "Looks like a bunch of chemists are pooling here. Let me ask an undergrad question as to why bond cleavage in chemistry requires energy whereas bond cleavage in biology releases energy, i.e. from ATP to ADP."

Here's the answer which I could think of on the spur of the moment:

Answer: "Bond cleavage always requires energy, even in biology. The question is whether the net reaction is energetically favorable, which it is in case of ATP. That gives the illusion of energy 'release' by bond cleavage. There are such cases in organic chemistry too; for instance the cleavage of a bond that expels nitrogen as a leaving group is usually very favorable (nitrogen is sometimes called "the world's best leaving group"). That does not mean that you don't need energy to cleave this bond, it's just that the unusual stability of nitrogen makes the net reaction favorable.

Fair enough?