One of the delights of wandering around an undergraduate chemistry laboratory is discussing the unexpected, if not the outright impossible, with students. The >100% yield in a reaction is an example. This is sometimes encountered (albeit only briefly) when students attempt to recrystallise a product from cyclohexane, and get an abundant crop of crystals when they put their solution into an ice-bath to induce the crystallisation. Of the solvent of course! I should imagine 1000% yields are possible like this.
What the students are not expecting is that cyclohexane has such a high melting point, higher than that of water! n-Octane for example melts at -57°C (and most of us have seen those travelogues in the antarctic where the petrol tanks need to be warned to prevent freezing), so why is that of cyclohexane so much higher? That it might be strange is shown by the melting points of the series:
- benzene, +5.5°C
- cyclohexadiene, -89°C
- cyclohexene, -97°C
- cyclohexane, +6.5°C.
Benzene one might explain because it famously stacks in a herring-bone fashion, with the relatively electropositive hydrogen attracted to the π-cloud on the face.
Clearly, this explanation cannot hold for cyclohexane, which has no π-face. What does the crystal look like?
If one inspects the structure closely, one can find quite a few H…H contacts at about 2.4Å and they are arranged in a particularly rigid three-dimensional manner. The maximum attractive force resulting from van der Waals, or dispersion interactions between two hydrogens is thought to occur at ~2.4Å. Perhaps cyclohexane is a prime (possibly THE prime) example of the influence of this (under-rated) interaction? A molecule covered in Velcro no less. By the way, can you spot the connection with the previous post?
Postscript: Below is a so-called non-covalent-analysis (NCI) of cyclohexane as packed into a crystal lattice. The coordinates are obtained from a neutron diffraction structure. The green regions indicate weakly attractive zones.