For the initial part of this chapter, we will be looking at a specific type of cases where the system undergoing the change is held at constant pressure. This, in fact, represents the way the majority of reactions are done in our chemical experience.
We define a term called enthalpy is the energy transferred between a system and the surroundings under constant pressure. We cannot measure the absolute enthalpy of a system but we can measure the change in enthalpy for a process. If the Initial and final states have enthalpies and , respectively then the change in enthalpy for the process can be defined as
This definition is only good if H is a state function, i.e., the change in enthalpy depends only on the initial and final states, not on the process itself.
If we simply measure the heat evolved (or absorbed) during a process, we will not have a measure of the change in enthalpy since, in general, q is not a state function. It depends on how the process occurs.
We need to restrict our process measurements to specific conditions in order to be able to measure enthalpy directly. In this case, if we maintain constant pressure and then simply measure the heat transferred as a result of the process qp we will have a measure of the change in enthalpy. The subscript p refers to the fact that this heat was measured with the system held at constant pressure conditions.
Thus, we have simply,
ΔH = qp
the heat measured at constant pressure.
For chemical reactions we can write
ΔH = ∑pHp – ∑rHr
where subscripts p and r refer to products and reactants, respectively. If ΔH is negative, then there is a net flow of heat from chemical system to surroundings (heat is released during the reaction). This kind of process is called Exothermic (EXO sounds like exit).
If ΔH is positive, then net flow of heat from surroundings to chemical system (heat absorbed). Endothermic (ENDO sounds like enter)
|CH4(g) + 2O2(g) CO2(g) + 2H2O(l)||ΔH = -890 kJ/mol (EXOTHERMIC)|
|H2(g) + I2(g) 2 HI(g)||ΔH = 52.2 kJ/mol (ENDOTHERMIC)|