Gibbs Free Energy (ΔG) represents energy associated with a chemical reaction that can be used to do work.
Concept #1: Gibbs Free Energy
Example #1: The reduction of iron (III) oxide with hydrogen produces iron metal and can be written as follows:
If our Gibbs Free Energy is less than zero then a reaction is spontaneous.
Practice: If ∆G is small and positive which of the following statements is true?
If Gibbs Free Energy is less than zero then a reaction will be spontaneous and favor both products and the forward direction.
Practice: Nitrogen gas combines with fluorine gas to form nitrogen trifluoride according to the reaction below at 25oC:
N2 (g) + 3 F2 (g) → 2 NF3 (g) ΔHo = -249.0 kJ ΔSo = -278 J/K
Calculate ΔGo and state if the reaction favors reactants or products at standard conditions.
To determine if a compound will either freeze or vaporize we must first determine their normal freezing and boiling point by assuming Gibbs Free Energy is zero.
Example #2: For mercury, ΔHvap = 58.5 kJ/mol and ΔSvap = 92.9 J/Kmol at 25°C. Does mercury boil at 350°C and 1 atm pressure?
Besides our standard equation for Gibbs Free Energy we have another that connects it to our equilibrium constant K:
Example #3: The chemical reaction, 2 NO2 (g) + Br2 (g), has a ΔS°= 135 J/molK and ΔH°=926 kJ/mol. Calculate the temperature when Keq = 4.50 x 105.
Under nonstandard conditions the pressure, temperature and molarities will be different from 1.0 atm, 25oC and 1.0 M respectively and so we use the following equation for Gibbs Free Energy:
Example #4: Calculate ΔGrxn at 25°C under the conditions shown below for the following reaction.
3 Cl2(g) → 2 Cl3 (g) ΔG° = +31.6 kJ
Example #5: For the reaction: N2 (g) + 2 O2 (g) ⇌ 2 NO2 (g), ΔG° = 75,550 J/mol at 175 K and ΔG° = 41,875 J/mol at 225 K.