Thermodynamics :  Gibbs Free Energy

Gibbs Free Energy (G) - The energy associated with a chemical reaction that can be used to do work.  The free energy of a system is the sum of its enthalpy (H) plus the product of the temperature (Kelvin) and the entropy (S) of the system:

Free energy of reaction (G)

Standard-state free energy of reaction (G)

Standard-state conditions

Measurements are also generally taken at a temperature of 25C (298 K)

Standard-State Free Energy of Formation (Gf)

Recall from the enthalpy notes that reactions can be classified according to the change in enthalpy (heat):

  • Endothermic - absorbs heat, H > 0
  • Exothermic - releases heat,H > 0

Reactions can also be classified according to the change in the free energy of the reaction:

  • Endergonic - NON-SPONTANEOUS, G > 0
  • Exergonic - SPONTANEOUS, G < 0
Summary
Favorable
Conditions
 
Unfavorable
Conditions
H < 0 H > 0
S > 0 S > 0

SPONTANEOUS: G is negative (G< 0)

NON-SPONTANEOUS: Gis positive (G > 0)

EQUILIBRIUM: G = 0


Compound   Hf   S
NH4NO3(s) -365.56 151.08
NH4+(aq) -132.51 113.4
NO3-(aq) -205.0 146.4

CalculateH,S, andG for the above reaction to determine whether the reaction is spontaneous or not.

First let's calculateHf.  Note that in the above reaction, one mole of NH4NO3 dissociates in water to give one mole each of NH4+ and NO3-:

Next, let's calculateS:

Now we can plug in these

values we've calculated into the free energy equation.

 


Temperature and Free Energy

Sample Calculations


Compound   Hf   S
N2(g)
0
191.61
H2(g)
0
130.68
NH3(g)
-46.11
192.45

1) CalculateH andS for the above reaction.  Explain what each of the signs mean.

H is negative which is favorable.

S is negative which is unfavorable.

2) Predict whether the above reaction is spontaneous at 25C.

G is negative, so the reaction is SPONTANEOUS.

3) Predict whether the above reaction is spontaneous at  500C.

Gis positive, so the reaction is NOT SPONTANEOUS.

 


Free energy and Equilibrium Constants

The following equation relates the standard-state free energy of reaction with the free energy of reaction at any moment in time during a reaction (not necessarily at standard-state conditions):

Reaction quotient (Qc or Qp) - The mathematical product of the concentrations (or partial pressures) of the products of a reaction divided by the mathematical product of the concentrations (or partial pressures) reactants of a reaction AT ANY MOMENT IN TIME.

Note: When Qc = Kc (or when Qp = Kp), a reaction is at equilibrium.

It was stated earlier that whenG = 0, a reaction is at equilibrium.  Let's consider the above reaction at equilibrium:

If we move RTlnK to the opposite side by subtracting it from both sides, we get the following reaction which relates the free energy of a reaction to the equilibrium constant of a reaction:


Summary

SPONTANEOUS
 
NON-SPONTANEOUS
G < 0
K > 1
G > 0
K < 1

The magnitude ofG measures how far a reaction is from equilibrium.  The larger the value ofG, the further the reaction is from equilibrium and the further the reaction must shift to reach equilibrium.  In reactions in which enthalpy is favorable and entropy is unfavorable, the reaction becomes less spontaneous (G increases) until eventually the reaction is not spontaneous (whenG > 0).  As the magnitude ofGchanges, so does the equilibrium constant. K.

 


Free energy and Cell potentials

Cell potential - A measure of the driving force behind an electrochemical reaction, reported in volts.  The potential of an electrochemical cell measures how far an oxidation-reduction reaction is from equilibrium.

The Nernst equation relates the standard-state cell potential with the cell potential of the cell at any moment in time:

If we rearrange the equation, we get:

This equation is very similar to the equation that relates the standard-state free energy of reaction with the free energy of reaction at any moment in time during a reaction:

We can convert these equations to get the following:

This shows that the free energy of a oxidation-reduction reaction is directly proportional to the cell potential of the reaction.
 

See notes for cell potentials and the Nernst equation.

Next:  "Electrochemistry:  Oxidation and Reduction"