In particular, make sure that you can see why the first 4 appears in the expression "4(+464)". Before you go on, make sure that you can see why every single number and arrow on this diagram is there. This obviously looks more confusing than the cycles we've looked at before, but apart from the extra enthalpy change of vaporization stage, it isn't really any more difficult. To do this you have to supply 41 kJ mol -1. You cannot apply bond enthalpies to this. To see how this fits into bond enthalpy calculations, we will estimate the enthalpy change of combustion of methane - in other words, the enthalpy change for this reaction: But for calculation purposes, it isn't something you need to worry about. You may well have to know the difference between a bond dissociation enthalpy and a mean bond enthalpy, and you should be aware that the word mean (or average) is used in two slightly different senses. So don't expect calculations using mean bond enthalpies to give very reliable answers. That means that if you use the C-H value in some calculation, you can't be sure that it exactly fits the molecule you are working with. So data tables use average values which will work well enough in most cases. The bond enthalpy of, say, the C-H bond varies depending on what is around it in the molecule. In fact, tables of bond enthalpies give average values in another sense as well, particularly in organic chemistry. Mean bond enthalpies are sometimes referred to as "bond enthalpy terms". That means that many bond enthalpies are actually quoted as mean (or average) bond enthalpies, although it might not actually say so. The average bond energy is therefore +1662/4 kJ, which is +415.5 kJ per mole of bonds. That comes to +1662 kJ and involves breaking 4 moles of C-H bonds. In the methane case, you can work out how much energy is needed to break a mole of methane gas into gaseous carbon and hydrogen atoms. In cases like this, the bond enthalpy quoted is an average value. And the strength of a bond is affected by what else is around it. Every time you break a hydrogen off the carbon, the environment of those left behind changes. However, if you took methane to pieces one hydrogen at a time, it needs a different amount of energy to break each of the four C-H bonds. It contains four identical C-H bonds, and it seems reasonable that they should all have the same bond enthalpy. What happens if the molecule has several bonds, rather than just 1? Consider methane, CH 4. The bond dissociation enthalpy for the H-Cl bond is +432 kJ mol -1. As an example of bond dissociation enthalpy, to break up 1 mole of gaseous hydrogen chloride molecules into separate gaseous hydrogen and chlorine atoms takes 432 kJ. The point about everything being in the gas state is essential you cannot use bond enthalpies to do calculations directly from substances starting in the liquid or solid state. So you can take all these terms as being interchangeable. These days, the term "bond enthalpy" is normally used, but you will also find it described as "bond energy" - sometimes in the same article. One of the most confusing things about this is the way the words are used.
0 Comments
Leave a Reply. |
Details
AuthorWrite something about yourself. No need to be fancy, just an overview. ArchivesCategories |