Thermodynamics Part 5 - Spontaneity

Thermodynamics Thermodynamics = the study of energy changes that accompany physical and chemical changes. Enthalpy (H): the total energy “stored” within a substance Enthalpy Change (ΔH): a comparison of the total enthalpies of the product & reactants. ΔH = H products - H reactants

Exothermic vs. Endothermic Exothermic reactions/changes: release energy in the form of heat; have negative ΔH values. H 2 O (g)  H 2 O (l) ΔH = kJ Endothermic reactions/changes: absorb energy in the form of heat; have positive ΔH values. H 2 O (l)  H 2 O (g) ΔH = kJ

Reaction Pathways Changes that involve a decrease in enthalpy are favored! EndothermicExothermic time EaEa EaEa P P RR

Entropy Entropy (S): the measure of the degree of disorder in a system; in nature, things tend to increase in entropy, or disorder. ΔS = S products – S reactants All physical & chemical changes involve a change in entropy, or ΔS. (Remember that a high entropy is favorable)

Entropy

Driving Forces in Reactions Enthalpy and entropy are DRIVING FORCES for spontaneous reactions (rxns that happen at normal conditions) It is the interplay of these 2 driving forces that determines whether or not a physical or chemical change will actually happen.

Free Energy Free Energy (G): relates enthalpy and entropy in a way that indicates which predominates; the quantity of energy that is available or stored to do work or cause change.

Free Energy ΔG = ΔH – TΔS Where: ΔG = change in free energy (kJ) ΔH = change in enthalpy (kJ) T = absolute temp (K) ΔS = change in entropy (kJ/K)

Free Energy ΔG: positive (+) value means change is NOT spontaneous ΔG: negative (-) value means change IS spontaneous

Relating Enthalpy and Entropy to Spontaneity ExampleΔHΔHΔSΔSSpontaneity 2K + 2H 2 O  2KOH + H 2 -+always spon. H 2 O(g)  H 2 O(l) lower temp. H 2 O(s)  H 2 O(l) higher temp. 16CO 2 +18H 2 O  2C 8 H O 2 +-never spon.

Example #1 For the decomposition of O 3 (g) to O 2 (g): 2O 3 (g)  3O 2 (g) ΔH = kJ/mol ΔS = °C a) Calculate ΔG for the reaction. ΔG = ( kJ/mol ) – (298 K )( KJ/mol·K ) ΔG = kJ

Example #1 For the decomposition of O 3 (g) to O 2 (g): 2O 3 (g)  3O 2 (g) ΔH = kJ/mol ΔS = °C b) Is the reaction spontaneous? YES

Example #1 For the decomposition of O 3 (g) to O 2 (g): 2O 3 (g)  3O 2 (g) ΔH = kJ/mol ΔS = °C c) Is ΔH or ΔS (or both) favorable for the reaction? Both ΔS and ΔH are favorable (both are driving forces)

Example #2 What is the minimum temperature (in °C) necessary for the following reaction to occur spontaneously? Fe 2 O 3 (s) + 3CO(g)  2Fe(s) + 3CO 2 (g) ΔH = kJ/mol; ΔS = J/K·mol (Hint: assume ΔG = kJ/mol) ΔG = ΔH – TΔS = (144.5) – (T)(0.0243) T ≈ 5950 K T = 5677 °C