Heat in Chemical Reactions and Processes Thermochemistry Lecture 2

Enthalpy Symbol : H Units : J Heat content of a system at constant pressure Chemists are interested in how enthalpy changes (ΔH) during reactions and processes Is heat absorbed or released ΔHrxn = H products – H reactants

Enthalpy of Reactions If heat is absorbed If heat is released ΔHrxn is positive meaning the Hproducts >Hreactants Endothermic reaction If heat is released ΔHrxn is negative meaning the Hreactants > Hproducts Exothermic reaction Note: ΔHrxn = q

Enthalpy of Phase Changes ΔHvap = heat required to vaporize one mole of a liquid H2O (l)  H2O (g) ΔHvap = 40.7 KJ Since condensation is opposite of vaporization ΔHcond = -40.7 KJ ΔHfus = heat required to melt one mole of a solid H2O (s)  H2O (l) ΔHvap = 6.01 KJ Since freezing is opposite of melting ΔHsolid = -6.01 KJ

Practice Questions Is the following endothermic or exothermic? Explain why. NH3 (g)  NH3 (l)

Practice Question 2 Is the following endothermic or exothermic? Explain why. 4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) + 1625 kJ ΔHrxn = -1625 KJ Exothermic

Thermochemical equations Balanced chemical equation Includes: physical state of reactants and products (s, l, g) ΔH Example: 4 Fe (s) + 3 O2 (g)  2 Fe2O3 (s) ΔHrxn = -1625 KJ

Calculating Enthalpy Change Some reactions are too slow to measure the enthalpy change Therefore, chemists theoretically determine enthalpy Hess’s Law: if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction

Applying Hess’s Law Determine ΔH for the following reaction: 2 S(s) + 3O2(g)  2SO3(g) If: S(s) + O2(g)  SO2(g) ΔH = -297 KJ 2SO3(g)  O2(g) + 2SO2 ΔH = 198 KJ

Determine ΔH for the following reaction: 2H2O2(l)  O2(g) + 2H2O(l) If: 2H2 (g) + O2(g)  2H2O(l) ΔH = -572 KJ H2 (g) + O2(g)  H2O2(l) ΔH = -188 KJ

Standard enthalpy (heat) of formation ΔH that accompanies the formation of one mole of the compound made from its constituents Every free element in its standard state is assigned a ΔHf of exactly 0.0 KJ. Example: H2(g) + S(s)  H2S (g) ΔHf = -21 KJ

Practice Problem Use the standard enthalpies of formation to calculate ΔHrxn for the combustion of methane. CH4(g) + O2(g)  CO2(g) + 2H2O(l)

Practice Problem Use the standard enthalpies of formation to calculate ΔHrxn for the following reaction: CaCO3 (s)  CaO (s) + CO2 (g)