Pause for thought

The naming, structural isomerism and physical properties of the alkanes and alkenes have all been covered in the previous sub-topic. This sub-topic is concerned with the chemical reactions of the alkanes and the alkenes. Despite the fact that their main use is as fuels the alkanes are remarkably unreactive chemically (their old name was the paraffins which means ‘little activity’). The lack of reactivity of alkanes is not just due to their realtively strong bond enthalpies and non-polar nature. The fact that they are virtually non-polar (due to the similar electronegativities of carbon and hydrogen and their often symmetrical shapes) does explain why they do not attract electrophiles or nucleophiles. However it is not true that just because they are relatively non-polar they should be unreactive per se. Many non polar molecules, e.g. oxygen, fluorine, silicon tetrachloride and boron trifluoride are extremely reactive. Equally their bond enthalpies are not that remarkably strong. Although the carbon to hydrogen bond has a relatively high value of approximately 414 kJ mol^-1, the carbon-carbon single bond is in the region of 346 kJ mol^-1 which is about average for bond enthalpies.

One of the real reasons why alkanes are so relatively unreactive is that carbon atoms are unable to ‘expand their octet’ to form what is technically known as a hypervalent molecule. Of course carbon atoms do contain empty 3d orbitals but these are much higher in energy than the occupied 2s and 2p (or hybridized combinations of 2s and 2p) orbitals so they cannot be utilized. Contrast this with a silicon atom which has the electronic configuration 1s²2s²2p⁶3s²3p². Silicon can utilize its empty d orbitals and expand its octet. This explains why tetrachloromethane, CCl₄, is almost completely inert, particularly in the presence of water, whereas silicon tetrachloride, SiCl₄, reacts vigorously with water.

SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(aq)

Unlike a carbon atom, a silicon atom can act as a Lewis acid and accept a non-bonding pair of electrons from a water molecule as it expands its octet in the process (see right). The fact that carbon cannot do this is as much responsible for the inactivity of alkanes as its bond enthalpies and lack of bond polarity.

Some books also state that because carbon is a much smaller atom relative to silicon then the water molecule cannot approach it to react. This may be true in the case of tetrachlormethane but is not really true in the case of methane where the four hydrogen atoms take up much less room than the four chlorine atoms in tetrachloromethane and yet methane, unlike silane, SiH₄ (which undergoes spontaneous combustion in air and decomposes above 420 ^oC), is still relatively unreactive.

One of the most obvious differences between alkanes and unsaturated compounds such as alkenes is that the unsaturated compounds burn in air with a much more yellow and smoky flame.

Although this is not a definitive way to distinguish between alkanes and alkenes – bromine water in the absence of ultraviolet light does that - it is a very noticeable difference. Why should this be so?

One proposed solution would be that the alkene needs more oxygen to burn than alkanes and that as the amount of oxygen in the surrounding air is limited alkenes do not burn so completely. However a quick glance at the relevant equations will show that this is not the case. For example, one mol of ethane requires 3.5 mol of oxygen to combust completely whereas one mol of ethene requires only 3 mol of oxygen for complete combustion. According to this the ethene should burn more completely in air.

C₂H₆(g) + 3.5O₂(g) → 2CO₂(g) + 3H₂O(l)

C₂H₄(g) + 3O₂(g) → 2CO₂(g) + 2H₂O(l)

Perhaps it is because the C=C bond enthalpy is stronger than the C-C bond enthalpy and so it does not break so readily. If this was the case we might expect the enthalpy of combustion of ethane to be much greater than the enthalpy of combustion of ethene, particularly as more water is formed (although there are less C-H bonds to break).

It is greater but only by about 10% per mole. In fact when 1.0 g of ethene burns in a plentiful supply of oxygen the heat given out is only slightly less than when 1.0 g of ethane completely combusts.

Approximately the same amount of heat is also given out with 1.0 g of ethyne, C₂H₂(g), which requires less oxygen (2.5 mol per mol of ethyne). Ethyne burning in air used to be used to provide light in miners’ and cavers’ lamps (see left) and on early automobiles as the flame is so yellow. However when burned in oxygen the heat evolved from the very hot blue flame is used to weld metal, as in oxy acetylene welding (see right).

It seems as though the smokiness is somehow connected with the degree of unsaturation but it is not really clear why. You can demonstrate this by burning a small amount of benzene in air (you will need to actually burn methylbenzene as the use of benzene is banned). The flame is not only extremely yellow and smoky but small pieces of black soot float around in the atmosphere afterwards. An unsaturated liquid such as benzene or methylbenzene takes longer to ignite than a gas as it needs to be vaporized first. The ratio of oxygen to hydrocarbon required to combust benzene completely is much lower at 7.5:6, compared to the much higher ratio of 21:6 when burning ethane. On this basis it is perhaps surprising that benzene is so smoky when it requires so much less oxygen than ethane when it is burned in air. Another example in chemistry of where a simple observation is not so easy to explain?