Rules for Using Thermochemical Equations
Thermochemical equations follow some easy-to-remember rules that make them useful for applications that will be used later in this module.
1. If a certain process has an enthalpy change DH, the reverse of that process has an enthalpy change of -DH.
H2O (s) H2O (l) DH = 6.00 kJ
For example, melting one mole of ice to liquid water requires the input of 6.00 kJ of enthalpy. Thus the liquid water has 6.00 kJ more enthalpy than the ice. If the reverse process, freezing the water to ice, is to occur, the water has to lose that enthalpy. So freezing one mole of liquid water to ice has an enthalpy change of -6.00 kJ.
Incidentally, as you can see, enthalpy and other thermodynamic changes depend a great deal on the phases of the reactants and products in a thermochemical equation. Therefore, it is imperative that the phases always be carefully specified when you are writing thermochemical equations.
Note: A process with a positive DH is called endothermic. A process with a negative DH is called exothermic. Thus, the reverse of an endothermic process is exothermic, and the reverse of an exothermic process is endothermic.
2. Multiplying a thermochemical equation by a constant also multiplies the thermodynamic quantity by that constant.
If it takes 6.00 kJ of enthalpy to melt one mole of ice, then it will take 2*6.00 or 12.0 kJ to melt two moles of ice, and 0.5*6.00 or 3.00 kJ to melt 0.5 moles of ice.
3. The thermodynamic quantity for the reaction applies as the equation is written. So, it can be used as a stoichiometric ratio with any of the reactants or products in the reaction. For, example, in the reaction:
H2 (g) + Cl2 (g) 2 HCl (g) DHº = -183 kJ
You can use -183 kJ/mol H2, -183 kJ/mol Cl2 or -183 kJ/2 mol HCl as stoichiometric ratios in a factor-label problem.
What is the enthalpy change when 3.5 g of H2 (g) reacts with Cl2 (g) to form HCl (g)?