![]() In other words, you would expect the enthalpy change of hydrogenation of cyclohexa-1,3-diene to be exactly twice that of cyclohexene - that is, -240 kJ mol -1. If the ring had two double bonds in it initially ( cyclohexa-1,3-diene), exactly twice as many bonds would have to be broken and exactly twice as many made. Other bonds have to be made, and this releases energy.īecause the bonds made are stronger than those broken, more energy is released than was used to break the original bonds and so there is a net evolution of heat energy. Where does this heat energy come from? When the reaction happens, bonds are broken (C=C and H-H) and this costs energy. The negative sign shows that heat is evolved. Help! "Enthalpy change" can be translated as "heat evolved or absorbed". In other words, when 1 mole of cyclohexene reacts, 120 kJ of heat energy is evolved. The enthalpy change during this reaction is -120 kJ mol -1. The hydrogenation equation could be written: In this case, then, each corner represents CH 2. In the cyclohexane case, for example, there is a carbon atom at each corner, and enough hydrogens to make the total bonds on each carbon atom up to four. The structures of cyclohexene and cyclohexane are usually simplified in the same way that the Kekulé structure for benzene is simplified - by leaving out all the carbons and hydrogens. Note: cyclohexane: six carbons in a ring, but the ane ending means NO C=C bond. The "CH" groups become CH 2 and the double bond is replaced by a single one. When hydrogen is added to this, cyclohexane, C 6H 12, is formed. Note: If you are a bit shaky on names: cyclohexene: hex means six carbons, cyclo means in a ring, ene means with a C=C bond. ![]() Cyclohexene, C 6H 10, is a ring of six carbon atoms containing just one C=C. In order to do a fair comparison with benzene (a ring structure) we're going to compare it with cyclohexene. If, for example, you hydrogenate ethene you get ethane: Hydrogenation is the addition of hydrogen to something. This is all so simple that you could understand it even if you had never done any! Help! It doesn't matter whether you've done any thermochemistry sums recently or not. This is most easily shown using enthalpy changes of hydrogenation. Every time you do a thermochemistry calculation based on the Kekulé structure, you get an answer which is wrong by about 150 kJ mol -1. Real benzene is a lot more stable than the Kekulé structure would give it credit for. Real benzene is a perfectly regular hexagon. In real benzene all the bonds are exactly the same - intermediate in length between C-C and C=C at 0.139 nm. That would mean that the hexagon would be irregular if it had the Kekulé structure, with alternating shorter and longer sides. Note: "nm" means "nanometre", which is 10 -9 metre. The problem is that C-C single and double bonds are different lengths. Note: Follow these links to get details about the addition reactions of ethene, or the substitution reactions of benzene.īenzene is a planar molecule (all the atoms lie in one plane), and that would also be true of the Kekulé structure. Instead, it usually undergoes substitution reactions in which one of the hydrogen atoms is replaced by something new. īecause of the three double bonds, you might expect benzene to have reactions like ethene - only more so!Įthene undergoes addition reactions in which one of the two bonds joining the carbon atoms breaks, and the electrons are used to bond with additional atoms.īenzene rarely does this. Because carbon atoms form four bonds, that means you are a bond missing - and that must be attached to a hydrogen atom.Īlthough the Kekulé structure was a good attempt in its time, there are serious problems with it. In this case, each carbon has three bonds leaving it. You have to count the bonds leaving each carbon to work out how many hydrogens there are attached to it. In diagrams of this sort, there is a carbon atom at each corner. This diagram is often simplified by leaving out all the carbon and hydrogen atoms! Each carbon atom has a hydrogen attached to it. The carbons are arranged in a hexagon, and he suggested alternating double and single bonds between them. Kekulé was the first to suggest a sensible structure for benzene. Bonding in benzene - the Kekulé structure ![]()
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