Hess's Law

Edited by Jamie (ScienceAid Editor), Administrator, Taylor (ScienceAid Editor)


Introduction to Hess's Law

Hess's law is based on the first law of thermodynamics which says that energy can not be created or destroyed but can be converted from one form to another. Mr Hess used this principle and related it to chemical reactions, creating, Hess's Law:

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Hess's Law says that the enthalpy change if a reaction depends only on the initial and final states of the reaction and is independent of the route by which the reaction may occur. In practical terms this means that you can do other reactions to get the energy for the main one.

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Using Cycles

If you regularly commute somewhere by car, you won't have just one route. What if one of the roads is closed? Likewise, there are several different ways you could go, and this very principle can be used in Chemistry since Mr. Hess said it doesn't matter which way we go - the energy will be the same.

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  1. 1
    The first way of going about it would be by using combustion where you are given the enthalpies of combustion for the reactants and the products, and then just burn them. See the example below
    Hess's cycle using combustion
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  2. 2
    We can also do a similar thing with formation. This is defined as the creation of a molecule from the elements that make it up in their standard states. Therefore this means H2 is used in the formation, not H. The equation for the formation of ethanoic acid:
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2C(s) + O2(s) + 2H2(g) ==>> CH3COOH(l)

But remember that the enthalpy of formation for an element is 0, because it is already in its standard state.

Bond Enthalpies

The mean bond enthalpy is the amount of energy required to break a particular bond. For example, in methane there are 4 C-H bonds. The two following reactions show these being broken:

CH4(g) ==>> CH3(g) + H(g) |>> ~H = 423 kJ mol-1
CH4(g) ==>> C(g) + 4H(g) |>> ~H = 1664 kJ mol-1

So the mean bond enthalpy for the C-H bond is 1664/4 = 416.

You can now use these bond enthalpies to calculate the energy in the bonds of all molecules.


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Categories : Physical

Recent edits by: Administrator, Jamie (ScienceAid Editor)

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