Driving Forces and Gibbs Free Energy
Some reactions are spontaneous because they give off
energy in the form of heat (
H < 0). Others are spontaneous because they lead to an increase in the
disorder of the system (S > 0). Calculations of H and S can be used to probe the driving force behind a particular reaction.
| Practice Problem 5: Calculate N2(g) + 3 H2(g) |
2 NH3(g)What happens when one of the potential driving forces behind a chemical reaction is favorable and the other is not? We can answer this question by defining a new quantity known as the Gibbs free energy (G) of the system, which reflects the balance between these forces.
The Gibbs free energy of a system at any moment in time is defined as the enthalpy of the system minus the product of the temperature times the entropy of the system.
G = H - TS
The Gibbs free energy of the system is a state function because it is defined in terms of thermodynamic properties that are state functions. The change in the Gibbs free energy of the system that occurs during a reaction is therefore equal to the change in the enthalpy of the system minus the change in the product of the temperature times the entropy of the system.
G
= H - (TS)
If the reaction is run at constant temperature, this equation can be written as follows.
G
= H - TS
The change in the free energy of a system that occurs during a reaction
can be measured under any set of conditions. If the data are collected under
standard-state conditions, the result is the standard-state free energy of reaction
(Go).
Go
= Ho - TSo
The beauty of the equation defining the free energy of a system is its ability to determine the relative importance of the enthalpy and entropy terms as driving forces behind a particular reaction. The change in the free energy of the system that occurs during a reaction measures the balance between the two driving forces that determine whether a reaction is spontaneous. As we have seen, the enthalpy and entropy terms have different sign conventions.
| Favorable | Unfavorable | |
| Ho
< 0 |
Ho
> 0 |
|
| So
> 0 |
So
< 0 |
The entropy term is therefore subtracted from the enthalpy term when
calculating Go
for a reaction.
Because of the way the free energy of the system is defined, Go is negative
for any reaction for which Ho is negative and So is positive. Go is therefore negative for any
reaction that is favored by both the enthalpy and entropy terms. We can therefore conclude
that any reaction for which Go is negative should be favorable, or spontaneous.
| Favorable, or spontaneous reactions: | Go
< 0 |
Conversely, Go is positive for any reaction for which Ho is positive
and So
is negative. Any reaction for which Go is positive is therefore unfavorable.
| Unfavorable, or non-spontaneous reactions: | Go
> 0 |
Reactions are classified as either exothermic (H < 0) or endothermic
(H > 0) on the
basis of whether they give off or absorb heat. Reactions can also be classified as exergonic
(G < 0) or endergonic
(G > 0) on the
basis of whether the free energy of the system decreases or increases during the reaction.
When a reaction is favored by both enthalpy (Ho < 0) and
entropy (So
> 0), there is no need to calculate the value of Go to decide whether the reaction should
proceed. The same can be said for reactions favored by neither
enthalpy (Ho
> 0) nor entropy (So
< 0). Free energy calculations become important for reactions favored by only one of
these factors.
| Practice Problem 6: Calculate NH4NO3(s) + H2O(l)
Use the results of this calculation to determine the value of |
The Effect of Temperature on the Free Energy of a Reaction
The balance between the contributions from the enthalpy and entropy terms to the free energy of a reaction depends on the temperature at which the reaction is run.
| Practice Problem 7: Use the values of N2(g) + 3 H2(g) |
The equation used to define free energy suggests that the entropy term will become more important as the temperature increases.
Since the entropy term is unfavorable, the reaction should become less favorable as the temperature increases.
| Practice Problem 8: Predict whether the following reaction is still spontaneous at 500C: N2(g) + 3 H2(g) Assume that the values of |
Standard-State Free Energies of Reaction
Go
for a reaction can be calculated from tabulated standard-state free energy data. Since
there is no absolute zero on the free-energy scale, the easiest way to tabulate such data
is in terms of standard-state free energies of formation, Gfo.
As might be expected, the standard-state free energy of formation of a substance is the
difference between the free energy of the substance and the free energies of its elements
in their thermodynamically most stable states at 1 atm, all measurements being made under
standard-state conditions.
Interpreting Standard-State Free Energy of Reaction Data
We are now ready to ask the obvious question: What does the value of Go tell us
about the following reaction?
N2(g) + 3 H2(g) 2 NH3(g)
By definition, the value of Go for a reaction measures the difference between
the free energies of the reactants and products when all components of the reaction are
present at standard-state conditions.
Go
therefore describes this reaction only when all three components are present at 1 atm
pressure.
The sign of Go tells us the direction in which the reaction has to shift to
come to equilibrium. The fact that Go is negative for this reaction at 25oC means that
a system under standard-state conditions at this temperature would have to shift to the
right, converting some of the reactants into products, before it can reach equilibrium.
The magnitude of Go for a reaction tells us how far the standard state is from
equilibrium. The larger the value of Go, the further the reaction has to go to get to
from the standard-state conditions to equilibrium.
Assume, for example, that we start with the following reaction under standard-state conditions, as shown in the figure below.
N2(g) + 3 H2(g) 2 NH3(g)
The value of G at that moment in time will be equal to the standard-state free energy
for this reaction, Go.
When Qp = 1:
As the reaction gradually shifts to the right, converting N2
and H2 into NH3, the value of G for the reaction will decrease. If we could find
some way to harness the tendency of this reaction to come to equilibrium, we could get the
reaction to do work. The free energy of a reaction at any moment in time is therefore said
to be a measure of the energy available to do work.
The Relationship Between Free Energy and Equilibrium Constants
When a reaction leaves the standard state because of a change in the ratio
of the concentrations of the products to the reactants, we have to describe the system in
terms of non-standard-state free energies of reaction. The difference between Go and G for a reaction is
important. There is only one value of Go for a reaction at a given temperature, but there
are an infinite number of possible values of G.
The figure below shows the relationship between G for the following
reaction and the logarithm to the base e of the reaction quotient for the
reaction between N2 and H2 to form NH3.
N2(g) + 3 H2(g) 2 NH3(g)
Data on the left side of this figure correspond to relatively small values
of Qp. They therefore describe systems in which there is far more
reactant than product. The sign of G for these systems is negative and the magnitude of G is large. The system is
therefore relatively far from equilibrium and the reaction must shift to the right to
reach equilibrium.
Data on the far right side of this figure describe systems in which there
is more product than reactant. The sign of G is now positive and the magnitude of G is moderately large. The
sign of G tells us
that the reaction would have to shift to the left to reach equilibrium. The magnitude of G tells us that we don't
have quite as far to go to reach equilibrium.
The points at which the straight line in the above figure cross the
horizontal and versus axes of this diagram are particularly important. The straight line
crosses the vertical axis when the reaction quotient for the system is equal to 1. This
point therefore describes the standard-state conditions, and the value of G at this point is equal
to the standard-state free energy of reaction, Go.
When Qp = 1:
The point at which the straight line crosses the horizontal axis describes
a system for which G
is equal to zero. Because there is no driving force behind the reaction, the system must
be at equilibrium.
When Qp = Kp:
The relationship between the free energy of reaction at any moment in time
(G) and the
standard-state free energy of reaction (Go) is described by the following equation.
G
= Go + RT
ln Q
In this equation, R is the ideal gas constant in units of J/mol-K, T is the temperature in kelvin, ln represents a logarithm to the base e, and Q is the reaction quotient at that moment in time.
As we have seen, the driving force behind a chemical reaction is zero (G = 0) when the reaction
is at equilibrium (Q = K).
0 = Go
+ RT ln K
We can therefore solve this equation for the relationship between Go and K.
Go
= - RT ln K
This equation allows us to calculate the equilibrium constant for any reaction from the standard-state free energy of reaction, or vice versa.
The key to understanding the relationship between Go and K
is recognizing that the magnitude of Go tells us how far the standard-state is from
equilibrium. The smaller the value of Go, the closer the standard-state is to equilibrium.
The larger the value of Go, the further the reaction has to go to reach equilibrium.
The relationship between Go and the equilibrium constant for a chemical reaction is
illustrated by the data in the table below.
Values of Go and K for Common Reactions at 25oC
| Reaction | Go (kJ) |
K | |||||
| 2 SO3(g) | |
2 SO2(g) + O2(g) | 141.7 | 1.4 x 10-25 | |||
| H2O(l) | |
H+(aq) + OH-(aq) | 79.9 | 1.0 x 10-14 | |||
| AgCl(s) + H2O | |
Ag+(aq) + Cl-(aq) | 55.6 | 1.8 x 10-10 | |||
| HOAc(aq) + H2O | |
H+(aq) + OAc-(aq) | 27.1 | 1.8 x 10-5 | |||
| N2(g) + 3 H2(g) | |
2 NH3(g) | -32.9 | 5.8 x 105 | |||
| HCl(aq) + H2O | |
H+(aq) + Cl-(aq) | -34.2 | 1 x 106 | |||
| Cu2+(aq) + 4 NH3(aq) | |
Cu(NH3)42+(aq) | -76.0 | 2.1 x 1013 | |||
| Zn(s) + Cu2+(aq) |  |
Zn2+(aq) + Cu(s) | -211.8 | 1.4 x 1037 |
| Practice Problem 9: Use the value of N2(g) + 3 H2(g) |
The equilibrium constant for a reaction can be expressed in two ways: Kc and Kp. We can write equilibrium constant expressions in terms of the partial pressures of the reactants and products, or in terms of their concentrations in units of moles per liter.
For gas-phase reactions the equilibrium constant
obtained from Go
is based on the partial pressures of the gases (Kp). For reactions in
solution, the equilibrium constant that comes from the calculation is based on
concentrations (Kc).
| Practice Problem 10: Use the following standard-state free energy of formation data to calculate the acid-dissociation equilibrium constant (Ka) at for formic acid: Compound
HCO2(aq) -372.3 H+(aq) 0.00 HCO2-(aq) -351.0 |
The Temperature Dependence of Equilibrium Constants
Equilibrium constants are not strictly constant because they change with temperature. We are now ready to understand why.
The standard-state free energy of reaction is a measure of how far the standard-state is from equilibrium.
Go
= - RT ln K
But the magnitude of Go depends on the temperature of the reaction.
Go
= Ho - TSo
As a result, the equilibrium constant must depend on the temperature of the reaction.
A good example of this phenomenon is the reaction in which NO2 dimerizes to form N2O4.
2 NO2(g) N2O4(g)
This reaction is favored by enthalpy because it forms a new bond, which makes the system more stable. The reaction is not favored by entropy because it leads to a decrease in the disorder of the system.
NO2 is a brown gas and N2O4 is colorless. We can therefore monitor the extent to which NO2 dimerizes to form N2O4 by examining the intensity of the brown color in a sealed tube of this gas. What should happen to the equilibrium between NO2 and N2O4 as the temperature is lowered?
For the sake of argument, let's assume that there is no significant change
in either Ho
or So
as the system is cooled. The contribution to the free energy of the reaction from the
enthalpy term is therefore constant, but the contribution from the entropy term becomes
smaller as the temperature is lowered.
Go
= Ho - TSo
As the tube is cooled, and the entropy term becomes less important, the net effect is a shift in the equilibrium toward the right. The figure below shows what happens to the intensity of the brown color when a sealed tube containing NO2 gas is immersed in liquid nitrogen. There is a drastic decrease in the amount of NO2 in the tube as it is cooled to -196oC.
| Practice Problem 11: Use
values of 2 NO2(g) |
N2O4(g)
The Relationship Between Free Energy and Cell Potentials
The value of G for a reaction at any moment in time tells us two things. The sign of G tells us in what
direction the reaction has to shift to reach equilibrium. The magnitude of G tells us how far the
reaction is from equilibrium at that moment.
The potential of an electrochemical cell is a measure of how far an oxidation-reduction reaction is from equilibrium. The Nernst equation describes the relationship between the cell potential at any moment in time and the standard-state cell potential.
Let's rearrange this equation as follows.
nFE = nFEo - RT ln Q
We can now compare it with the equation used to describe the relationship between the free energy of reaction at any moment in time and the standard-state free energy of reaction.
G
= Go + RT
ln Q
These equations are similar because the Nernst equation is a special case of the more general free energy relationship. We can convert one of these equations to the other by taking advantage of the following relationships between the free energy of a reaction and the cell potential of the reaction when it is run as an electrochemical cell.
| Practice Problem 12: Use the relationship
between |