Henry Rzepa's blog

Tag: cyclohexane

Tetrahedral carbon and cyclohexane.
Following the general recognition of carbon as being tetrahedrally tetravalent in 1869 (Paterno) and 1874 (Van’t Hoff and Le Bell), an early seminal exploitation of this to the conformation of cyclohexane was by Hermann Sachse in 1890.[cite]10.1002/cber.189002301216 [/cite] This was verified when the Braggs in 1913[cite]10.1098/rspa.1913.0084[/cite], followed by an oft-cited article by Mohr in 1918,[cite]10.1002/prac.19180980123[/cite] established the crystal structure of diamond as comprising repeating rings in the chair conformation.^† So by 1926, you might imagine that the shape (or conformation as we would now call it) of cyclohexane would be well-known. No quite so for everyone!

When The Journal of the Imperial College Chemical Society (Volume 6) was brought to my attention, I found an article by R. F Hunter;

He proceeds to argue as follows:
1. The natural angle subtended at a tetrahedral carbon is 109.47°.
2. “The internal angle between the carbon to carbon valencies of a six-membered ring cyclohexane will, if the ring is uniplanar, be … 120°.
3. “When the cyclohexane ring is prepared … we must therefore have the pushing apart of two of the valencies”.
4. The object of the experiments commenced in this College in 1914 was “to find what effect the pushing apart of the valencies …must have on the angle between the remaining pair of valencies“.
5. You do wonder then why the assumption highlighted in red above was never really questioned during the twelve-year period of investigating angles around tetrahedral carbon.
The article itself is quite long, reporting the synthesis of many compounds in search of the postulated effect. Of course around twenty years later, Derek Barton used the by then generally accepted conformation of cyclohexane to explain reactivity in what become known as the theory of conformational analysis.

These two articles dating from 1926, and probably thought lost to science, show how some ideas can take decades to have any influence, whilst others can take root very much more quickly.

Postscript added in 2022. In the news recently has been Lonsdaleite, another allotrope of carbon. The saturated diamond structure has all the tetrahedral carbon forming six-membered rings in the chair^† conformation. In Lonsdaleite,[cite]10.1063/1.1841236[/cite] there are boat as well as chair rings!

^†The chair cyclohexane structure is easily discerned from Figure 7 in the Braggs’ paper![cite]10.1098/rspa.1913.0084[/cite]
August 22, 2018
The melting points from benzene to cyclohexane: a prime example of dispersion forces in action?
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:
1. benzene, +5.5°C
2. cyclohexadiene, -89°C
3. cyclohexene, -97°C
4. 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.

The crystal structure of benzene. Click for 3D

Clearly, this explanation cannot hold for cyclohexane, which has no π-face. What does the crystal look like?

Crystal structure of cyclohexane. Click for 3D

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.

Click for 3D.
December 30, 2010

Tag: cyclohexane

Tetrahedral carbon and cyclohexane.

The melting points from benzene to cyclohexane: a prime example of dispersion forces in action?