Thoughts with Richard Bleil
There is a saying that teaching is the art of lying to students without doing any real harm. Take, for example, science. We have habit of teaching the disciplines of science in such a way that it looks like one complete beautiful package, like a gift under a tree. But what we teach, it should be remembered, is not “truth”. We are teaching the models that best encompass the totality of observation and experiment to date, and any one of these models can fail right down to the fundamental laws of thermodynamics. If you look closely enough, really look, you’ll see that there are huge gaping holes in what we think we understand.
Take, for example, the hydrogen bond. Despite its unfortunate name, the hydrogen bond is one of the three known primary non-bonding intermolecular forces. These are the forces that, for example, hold elements and compounds to one another to form liquids and solids. They’re not chemical bonds, but without them, we would have gases but no other states of matter. Something has to hold those chemicals together.
The hydrogen bond is the strongest of the three bonds, and is critical for biochemical structures, recognition and processes, and yet it is very poorly understood. In fact, there is an active debate as to if it really is an independent intermolecular force (most of us feel it is) or if it’s just a really strong form of the second strongest force, the dipole-dipole force. The dipole-dipole non-bonding intermolecular force is easy to understand, because it’s just the attraction for the slightly negative portion of one molecule for the partial positive portion of another. It’s like a little magnet, but hydrogen bond is different. Somehow. We just don’t know how.
The way I was taught about hydrogen bonds, and the way I still teach it today (at least it presents the hydrogen bond as different and unique) is that you must have a hydrogen directly bonded to an element in a compound that has at least one lone pair of electrons and that is very electronegative. Electronegativity is a measure of how much attraction an element has for electrons, so to have hydrogen (not very electronegative) bonded to a very electronegative element gives is a very strong partial positive charge. It’s not fully charged, but it’s an element with a very positive nature. A lone-pair of electrons are electrons that do not participate in bonding. These are two electrons (hence “pair”) that occupy an orbital near that electronegative element. An orbital is just a region of space near an element with a high electron density because the electrons enjoy hanging out there. So in a hydrogen bond, the hydrogen is not attracted to the element so much as it is to that orbital with the lone-pair electrons.
Is this true? I don’t know, but here’s the definite lie. I was taught, and I still teach today (if you can find my lecture on the topic) that on the periodic chart, there are only three elements that fit both of these criteria (high electronegative and at least one lone pair of electrons): nitrogen, oxygen and fluorine. But this is a lie.
If you closely examine the periodic chart, you’ll realize that chlorine also matches these criteria. It has three lone pair electron orbitals (like fluorine), and its electronegativity is equal to that of nitrogen, so why doesn’t chlorine have hydrogen bonding? This question has haunted me for well over forty years now. It just makes no sense to me and, being a scientist, I’ve been chasing this question for, well, most of my life.
For a long time, I felt like it night have something to do with the concept of the hard and soft Lewis base. Basically, a Lewis base has those lone-pair electrons and donates them to do some pretty advanced chemistry, but the important thing is that a hard Lewis base has electrons that are held tightly and are very difficult to distort, while soft Lewis bases are easily distorted. If the hydrogen in a hydrogen bond is attracted to the lone-pair electrons, wouldn’t it make sense that it would distort this electron cloud by attracting the electrons to it? And if this is true, then wouldn’t it make sense that you need a soft (easily distorted) electron cloud? But that’s wrong. I don’t know why, but Fluorine is the hardest Lewis base on the periodic chart. Chlorine has electron clouds more easily distorted, and yet it’s Fluorine that forms the hydrogen bond.
I’ve had the concept of the hard/soft Lewis Base for a couple of decades now but have never been able to get past the fact that, in my mind, it’s just opposite of what I’m thinking it should be. So, I’m wrong. Another question to ponder and keep me awake at night and drive away my wife from me and into the arms of another man with much less on his mind.
Driving home from the Drive-In theater where I work (but next week is our last one for the season) on Saturday, I was mulling this question over in my mind. I work with a delightful young woman interested in chemistry as her college major, and we often discuss science and life. We happened to be pondering the hydrogen bond on Saturday which brought it to the forefront of my mind. And it struck me. A brand-new thought on how and why the hydrogen bond might form. No, I’m not going to go into it, but as it turns out, I have the software package (that cost me thousands of dollars) to perform the necessary calculations, but I needed a book for reference on the way that I think it’s happening (a five-hundred-dollar book). So today, I’ve kept my computer running the preliminary calculations continuously (it’ll probably take about a month’s worth of calculations to complete) on a topic that will never be published.
See, to publish, you must have resources that an individual like me simply doesn’t have. The American Chemical Society controls these resources (research journals) and they are priced far out of reach for individuals like me who are not affiliated with a major research institution. No paper will be taken seriously unless it’s published in one of these journals, but they won’t publish papers without a current bibliography which you have to build by having, ready for this? Access to those publications. Will my research pay off? I hope so, but if so, it’ll only be to finally settle the question in my own mind. Kind of sad, really.
Bleil's Banter 