Very often chemical changes are accompanied by changes in the heat content (enthalpy, H) of the materials which are reacting and the change in heat content is shewn by a change in temperature.
A reaction that gives out energy and heats the surroundings is said to be exothermic. The energy content of the products, H2, is less than that of the reactants, H1. The enthalpy change ΔH is hence negative. (NB. Δ means a change in).
A reaction that takes in energy and cools the surroundings is endothermic. The energy content of the products, H2, is greater than that of the reactants, H1. The enthalpy change ΔH is hence positive.
In comparing enthalpy changes it is essential to ensure the conditions of the system are the same before and after the reaction because ΔH is affected by temperature, pressure and concentration of solutions. The standard conditions for temperature and pressure are 298K and 1 atmosphere respectively. Any enthalpy change measured under these conditions is described as a standard enthalpy change of reaction, ΔHΘ298. Also the substances involved are in their normal physical states.
The standard enthalpy change of reaction, ΔHΘr, 298 is the enthalpy change when molar quantities of reactants as specified in the equation react together under standard conditions.
The standard enthalpy change of combustion, ΔHΘc, 298 is the enthalpy change when 1 mole of a substance is completely burnt in oxygen under standard conditions.
E.g. C2H5OH(l) + 3O2(g) = 2CO2(g) + 3H2O(l) ΔHΘc, 298 = -1368 kJ mol-1
The standard enthalpy change of formation ΔHΘf, 298 = is the enthalpy change when 1 mole of a substance is formed from its elements in their standard states.
E.g. Mg(s) + ½ O2 = MgO(s) ΔHΘf, 298 = -602 kJ mol-1
The enthalpy change of formation of methane cannot be determined directly by experiment. It is possible, however, to determine the enthalpy changes of combustion of carbon, hydrogen and methane. The key idea is that the total enthalpy change for one route is the same as the total enthalpy change for an alternative route. This is one way of stating Hess's Law which says that the total enthalpy change for a reaction is independent of the route taken.
If we can measure ΔH2 and ΔH, we can find ΔH1. Using Hess's Law:
ΔH + ΔH2 = ΔH1
Hence: ΔH = ΔH1 - ΔH2
Enthalpy cycles are useful because they enable a value for an enthalpy change to be determined for a reaction which cannot be determined directly from experiment.
Using Enthalpy Change of Formation to Determine Enthalpy Change of Reaction
It is not possible to determine the enthalpy change for the reaction between silicon tetrachloride and water directly by experiment.
|ΔH1||= ΔHf [SiCl4(l)] + 2 x ΔHf [H2O(l)]|
|= -640.2 + 2(-285.9)|
|= -1212 kJ mol-1|
|ΔH2||= ΔHf [SiO2(s)] + 4 x ΔHf [HCl(g)]|
|= -910.9 + 4(-92.3)|
|= -1280.1 kJ mol-1|
Using Hess' Law:
|ΔH2||= ΔH2 - ΔH1|
|= -1280.1 -(-1212)|
|= -68.1 kJ mol-1|
Different fuels have different enthalpy changes of combustion.
|ΔHc kJ mol-1||-4163||-890||-726||-393||-286|
Why do they vary so much? In order to release energy, fuels must combine with oxygen.
The enthalpy change of combustion depends on two things:
The energy released during combustion comes from the making of bonds to oxygen. If methanol already has one bond made, it will give out less energy when it burns.
As a general rule, the more oxygen a fuel has in its molecule, the less energy it will give out when it burns.
The quantity of energy needed to break a particular bond in a molecule is called the bond dissociation enthalpy, or bond enthalpy for short.
E.g. H2 (g) -> 2H (g) ΔHΘ = + 436 kJ mol-1
This is a positive ΔH value because bond breaking is an endothermic change. Bond making is an exothermicchange.
The stronger a bond, the more difficult it is to break - the higher its bond enthalpy. Tables of bond enthalpies are average values because the exact value depends on the particular compound in which the bond is found. Other atoms and bonds around it influence the value. Double bonds have higher bond enthalpies than single bonds. Triple bonds are higher still. In general the higher the bond enthalpy the shorter the bond.
The bond-breaking and bond-making can be represented in an enthalpy cycle.
|ΔH1||= enthalpy change for bond breaking|
|= 4 x E(C-H) + 2 x E(O=O)|
|= 4 x 413 + 2 x 498|
|= + 2648 kJ mol-1|
|ΔH2||= enthalpy change for bond making|
|= - [ 2 x E(C=O) + 4 x E(O-H) ]|
|= - [ 2 x 805 + 4 x 464 ]|
|= - 3466 kJ mol-1|
|Hence enthalpy change of combustion ΔH||= ΔH1 + ΔH2|
|= +2648 - 3466|
|= -818 kJ mol-1|
This value is a little different from the standard enthalpy change of combustion of methane, -890kJ mol-1. This is because bond enthalpies are average values from a range of compounds and ΔH1 is a standard value, H2O(g) was used rather than H2O(l).
Often fuels are compared by their energy density. This is the energy you get per kilogramme of fuel. This is important to consider as the fuel may have to be carried around.
Part of this site was last updated on 21st January 2009.
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