For constant temperature and pressure processes, work is done by the system to expand or contract. To study the heat change accompanying chemical reactions at constant pressure and constant temperature, a new term called enthalpy is used.
Enthalpy (H) is a more general measure of the energy of a system. It takes into account the internal energy of a system as well as its pressure and volume.
H = U + PV
Like in the case of internal energy, change in enthalpy is identified.
Change in enthalpy is equal to heat supplied to the system
Therefore changes in enthalpy can be measured by changes in heat energy. Enthalpy is also defined as heat flow.
Heat (q) is a type of energy that is related to temperature but is not the same as temperature, since the energy may be used for something other than increasing temperature. For example, the heat energy required to melt 1 mole of molecules is the heat of fusion (H-fus), but no temperature change is involved in this phase change. The heat energy required to change 1 mole of liquid to 1 mole of gas is the heat of vaporization (H-vap). Similarly, no temperature change is involved in this or any other phase change.
It also requires different amounts of heat energy to change the temperature different types of molecules. The energy required to increase the temperature of 1 mole (n) of substance 1 °C is its heat capacity (cp). Similarly, the energy required to increase the temperature of 1 g of substance 1 °C is specific heat. Therefore the heat (q) required for a temperature change (T) can be calculated from
q = n*(cp)*T where (cp = heat capacity; n = number of moles)
Chemical reactions also involve a change in enthalpy (∆H). The ∆Hrxn is the energy per mole of reaction (q/mol). Exothermic reactions release energy into the surrounding and are designated by a negative sign on the H. Endothermic reactions take energy from the surroundings and are designated by a positive value of H.
The value of H can be determined experimentally by measuring the temperature change of the surroundings. Calorimetry is an experiment to measure q by measuring temperature change. The heat capacity of the calorimeter (the constant pressure surroundings in the experiment) is used to convert temperature change to heat energy. Since the heat capacity (CP) refers to the whole calorimeter, heat can be calculated from
q = (cp)T
The value of ∆H can also be calculated theoretically. Since the energy change in a chemical reaction comes from making and breaking bonds, the value of ∆H can be calculated from the energy of the bonds, bond energy. Energy is released (–∆H) when bonds are made and energy is absorbed (+∆H) when bonds are broken.
The value of ∆H can also be calculated from the ∆H values of other reactions. Because H is a state function, the path does not affect the final result. Thus ∆H values of a series of known reactions can be mixed to obtain the ∆H of an unknown reaction. The ∆H value depends on how the reaction is written. If the reaction is reversed, the sign on ∆H is changed. If the stoichiometric coefficients of a reaction are multiplied by some factor, ∆H is multiplied by the same factor. If reactions are added together, so are the ∆H values. Hess's law states the ∆H of a reaction that is the sum of other reactions is the sum of the ∆H values of those reactions.