A recent report in Nature Chemistry by some german chemists at Ludwig-Maximilians University suggests that our current models of chemical bonding may be inadequate to explain the behavior of all molecules. The scientists chose to study a small organic molecule known as chlorotrinitromethane, which stands out right away to any student of organic chemistry as a molecule that would possess an incredibly lopsided dipole moment. This publication is important because it is only through a study of chemical bonding that we can understand a molecules reactivity. Chemists such as myself have trouble predicting a molecules properties without knowing how the different atoms are all interconnected through chemical bonds. To appreciate the results of this study, it’s necessary to understand what a dipole moment really means.
I’ve written before here on Associated Content about the implications of electronegativity, but we have to dive a little deeper to understand this particular Nature paper. All organic molecules are held together with covalent bonds. A covalent bond is nothing more complicated than a pair of electrons which form a link between two separate atoms. Bonds are the “glue” that make building molecules possible. Ideally, the two electrons within each covalent bond are equally shared between the two atoms that they connect. Each atom has approximately half the electron density, and neither atom has any more or less. However, in practice you will find that most molecules have an unequal sharing of the electrons. You can usually predict this behavior by looking at the Periodic Table of Elements. Atoms which are found towards the upper right hand corner of the Table are more “electronegative” than atoms found in lower left hand corner. Electronegative atoms are more greedy for the electrons, and take more than their fair 50 percent share. This leaves the other atom with a deficiency of electrons (otherwise known as a slight positive charge, as the electrons are negative, and the deficiency of a negative charge leads to a positive charge).
In the case of molecule under examination, chlorotrinitromethane, there is one carbon at the center of the molecule and the four substituents are one chlorine and three nitro, or -NO2 groups. You can see the molecular structure of chlorotrinitromethane in the image I have attached to this article. Each of the three nitro group contains three electronegative atoms (one nitrogen and two oxygens), each of which is drawing the electron density of the carbon-nitro bonds onto the nitro groups. This leaves the center carbon atom with not near enough electron density to make it happy. As a result, the bond between the carbon and the chlorine – which would normally place most of its electron density on the chlorine, as the chlorine is more electronegative than carbon – is sucked back almost entirely onto the carbon, as a last-ditch attempt to try and compensate for all the the electron density that the nitro groups are sucking off the carbon. The resulting cloud of electron density from this bond therefore hovers over the carbon, leaving little to make up the carbon-chlorine bond.
As a result, the bond cannot extend as far as it normally would. While a normal carbon-halogen bond is in the realm of 1.85 Angstroms long (one Angstrom being 1×10^-10 meters), X-ray analysis of this chlorotrinitromethane molecule showed the bond to be only 1.69 Angstroms long. Not a very significant difference to most people – a fraction of a nanometer difference – but still, chemically speaking, it is incredibly significant. A chlorine that displays a partial positive charge (as this one does) is almost unheard of, and the chlorotrinitromethane isn’t that complicated of a molecule; that this discovery has waited so long to be found speaks volumes about our so-called “laws” of bonding and chemical behavior. It just goes to show, there is at least one exception to every rule.
This discovery should go on to form the foundation of other important discoveries. The assembly of molecules into larger structures and the reactivity of certain atoms within the molecule now has to be reconsidered, as what we once thought to be true turns out, in some cases, to be quite false. This is an important lesson for all scientists: never become so adhered to dogma and tradition that you have trouble accepting an instance outside the normal realm of expectation.