Energy changes and bonding
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For the enthalpy change of atomisation for compounds, by definition, should the compound be in its standard state or a gas? I assumed it should be in its standard stae, but a book says that calculations are often incorrect as the enthalpy change of atomisation is for the compound to actually be a gas, which is not always the same as its standard state. (I am guessing here the book is trying to say the compound should technically be a gas but we use its standard state instead, which is why there are inaccuracies.) Which state should the compound actually be?
Jenkin says:
It is not often that one needs to deal with the enthalpy change of atomisation of a compound as distinct from an element. But there is no reason to treat compounds differently from elements.
This means that the definition refers to atomisation at 298 K from the stable state at 1 atm pressure. We need to remember that this is an arbitrary definition and there may be reasons for using a different temperature or pressure, provided they are stated.
Compounds are, of course, often liquid or solid at 298K and 1 atm so a value of DHat may include the enthalpy changes of fusion and/or vaporisation. These will usually be known so that a correction can be made if necessary.
When using bond energies to calculate the enthalpy change of atomisation of a compound, the enthalpy change will refer to the compound in gaseous form, although it may normally be a liquid (or a solid) under standard conditions.
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updated: 02 March 2007
