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AP*
Chapter 17
Spontaneity, Entropy,
and Free Energy
AP Learning Objectives
 LO 2.15 The student is able to explain observations regarding the solubility of
ionic solids and molecules in water and other solvents on the basis of particle
views that include intermolecular interactions and entropic effects. (Sec 17.1)
 LO 5.3 The student can generate explanations or make predictions about the
transfer of thermal energy between systems based on this transfer being due to
a kinetic energy transfer between systems arising from molecular collisions. (Sec
17.3)
 LO 5.12 The student is able to use representations and models to predict the
sign and relative magnitude of the entropy change associated with chemical or
physical processes. (Sec 17.1-17.3, 17.5)
 LO 5.13 The student is able to predict whether or not a physical or chemical
process is thermodynamically favored by determination of (either quantitatively
or qualitatively) the signs of both Ho and So, and calculation or estimation of
Go when needed. (Sec 17.4, 17.6)
 LO 5.14 The student is able to determine whether a chemical or physical process
is thermodynamically favorable by calculating the change in standard Gibbs free
energy. (Sec 17.4, 17.6)
AP Learning Objectives
 LO 5.15 The student is able to explain how the application of external energy
sources or the coupling of favorable with unfavorable reactions can be used to
cause processes that are not thermodynamically favorable to become favorable.
(Sec 17.6, 17.9)
 LO 5.16 The student can use Le Châtelier’s principle to make qualitative
predictions for systems in which coupled reactions that share a common
intermediate drive formation of a product. (Sec 17.6)
 LO 5.17 The student can make quantitative predictions for systems involving
coupled reactions that share a common intermediate, based on the equilibrium
constant for the combined reaction. (Sec 17.6)
 LO 5.18 The student can explain why a thermodynamically favored chemical
reaction may not produce large amounts of product (based on consideration of
both initial conditions and kinetic effects), or why a thermodynamically
unfavored chemical reaction can produce large amounts of product for certain
sets of initial conditions. (Sec 17.1, 17.6-17.8)
AP Learning Objectives
 LO 6.25 The student is able to express the equilibrium constant in terms of Go
and RT and use this relationship to estimate the magnitude of K and,
consequently, the thermodynamic favorability of the process. (Sec 17.8)
Section 17.1
Spontaneous Processes and Entropy
AP Learning Objectives, Margin Notes and References
 Learning Objectives



LO 2.15 The student is able to explain observations regarding the solubility of ionic solids and molecules in water
and other solvents on the basis of particle views that include intermolecular interactions and entropic effects.
LO 5.12 The student is able to use representations and models to predict the sign and relative magnitude of the
entropy change associated with chemical or physical processes.
LO 5.18 The student can explain why a thermodynamically favored chemical reaction may not produce large
amounts of product (based on consideration of both initial conditions and kinetic effects), or why a
thermodynamically unfavored chemical reaction can produce large amounts of product for certain sets of initial
conditions.
 Additional AP References

LO 5.12 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.1
Spontaneous Processes and Entropy
Thermodynamics vs.
Kinetics
 Domain of Kinetics
 Rate of a reaction depends
on the pathway from
reactants to products.
 Thermodynamics tells us
whether a reaction is
spontaneous based only
on the properties of
reactants and products.
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Section 17.1
Spontaneous Processes and Entropy
Spontaneous Processes and Entropy
 Thermodynamics lets us predict the direction in
which a process will occur but gives no information
about the speed of the process.
 A spontaneous process is one that occurs without
outside intervention.
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Section 17.1
Spontaneous Processes and Entropy
CONCEPT CHECK!
Consider 2.4 moles of a gas contained in a 4.0 L bulb
at a constant temperature of 32°C. This bulb is
connected by a valve to an evacuated 20.0 L bulb.
Assume the temperature is constant.
a) What should happen to the gas when you open
the valve?
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Section 17.1
Spontaneous Processes and Entropy
CONCEPT CHECK!
Consider 2.4 moles of a gas contained in a 4.0 L bulb
at a constant temperature of 32°C. This bulb is
connected by a valve to an evacuated 20.0 L bulb.
Assume the temperature is constant.
b) Calculate ΔH, ΔE, q, and w for the process you
described above.
All are equal to zero.
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Section 17.1
Spontaneous Processes and Entropy
CONCEPT CHECK!
Consider 2.4 moles of a gas contained in a 4.0 L bulb
at a constant temperature of 32°C. This bulb is
connected by a valve to an evacuated 20.0 L bulb.
Assume the temperature is constant.
c) Given your answer to part b, what is the
driving force for the process?
Entropy
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Section 17.1
Spontaneous Processes and Entropy
The Expansion of An Ideal Gas Into an Evacuated Bulb
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Section 17.1
Spontaneous Processes and Entropy
Entropy
 The driving force for a spontaneous process is an
increase in the entropy of the universe.
 A measure of molecular randomness or disorder.
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Section 17.1
Spontaneous Processes and Entropy
Entropy
 Thermodynamic function that describes the number
of arrangements that are available to a system
existing in a given state.
 Nature spontaneously proceeds toward the states
that have the highest probabilities of existing.
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Section 17.1
Spontaneous Processes and Entropy
The Microstates
That Give a
Particular
Arrangement
(State)
Section 17.1
Spontaneous Processes and Entropy
Positional Entropy
 A gas expands into a vacuum to give a uniform
distribution because the expanded state has the
highest positional probability of states available to
the system.
 Therefore: Ssolid < Sliquid << Sgas
Section 17.1
Spontaneous Processes and Entropy
CONCEPT CHECK!
Predict the sign of ΔS for each of the following,
and explain:
+ a) The evaporation of alcohol
– b) The freezing of water
– c) Compressing an ideal gas at constant
temperature
+ d) Heating an ideal gas at constant
pressure
+ e) Dissolving NaCl in water
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Section 17.2
Entropy and the Second Law of Thermodynamics
AP Learning Objectives, Margin Notes and References
 Learning Objectives

LO 5.12 The student is able to use representations and models to predict the sign and relative magnitude of the
entropy change associated with chemical or physical processes.
 Additional AP References

LO 5.12 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.2
Entropy and the Second Law of Thermodynamics
Second Law of Thermodynamics
 In any spontaneous process there is always an increase
in the entropy of the universe.
 The entropy of the universe is increasing.
 The total energy of the universe is constant, but the
entropy is increasing.
Suniverse = ΔSsystem + ΔSsurroundings
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Section 17.2
Entropy and the Second Law of Thermodynamics
ΔSsurr
 ΔSsurr = +; entropy of the universe increases
 ΔSsurr = -; process is spontaneous in opposite direction
 ΔSsurr = 0; process has no tendency to occur
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Section 17.3
The Effect of Temperature on Spontaneity
AP Learning Objectives, Margin Notes and References
 Learning Objectives


LO 5.3 The student can generate explanations or make predictions about the transfer of thermal energy between
systems based on this transfer being due to a kinetic energy transfer between systems arising from molecular
collisions.
LO 5.12 The student is able to use representations and models to predict the sign and relative magnitude of the
entropy change associated with chemical or physical processes.
 Additional AP References


LO 5.3 (see Appendix 7.2, “Thermal Equilibrium, the Kinetic Molecular Theory, and the Process of Heat”)
LO 5.12 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.3
The Effect of Temperature on Spontaneity
CONCEPT CHECK!
For the process A(l)
A(s), which direction involves
an increase in energy randomness? Positional
randomness? Explain your answer.
As temperature increases/decreases (answer for both),
which takes precedence? Why?
At what temperature is there a balance between energy
randomness and positional randomness?
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Section 17.3
The Effect of Temperature on Spontaneity
CONCEPT CHECK!
Describe the following as spontaneous/non-spontaneous/cannot tell,
and explain.
A reaction that is:
a) Exothermic and becomes more positionally random
Spontaneous
b) Exothermic and becomes less positionally random
Cannot tell
a) Endothermic and becomes more positionally random
Cannot tell
a) Endothermic and becomes less positionally random
Not spontaneous
Explain how temperature affects your answers.
Section 17.3
The Effect of Temperature on Spontaneity
ΔSsurr
 The sign of ΔSsurr depends on the direction of the heat
flow.
 The magnitude of ΔSsurr depends on the temperature.
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Section 17.3
The Effect of Temperature on Spontaneity
ΔSsurr
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Section 17.3
The Effect of Temperature on Spontaneity
ΔSsurr
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Section 17.3
The Effect of Temperature on Spontaneity
ΔSsurr
Heat flow (constant P) = change in enthalpy = ΔH
Ssurr
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H
= 
T
26
Section 17.3
The Effect of Temperature on Spontaneity
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Section 17.4
Free Energy
AP Learning Objectives, Margin Notes and References
 Learning Objectives


LO 5.13 The student is able to predict whether or not a physical or chemical process is thermodynamically favored
by determination of (either quantitatively or qualitatively) the signs of both Ho and So, and calculation or
estimation of Go when needed.
LO 5.14 The student is able to determine whether a chemical or physical process is thermodynamically favorable
by calculating the change in standard Gibbs free energy.
Section 17.4
Free Energy
Free Energy (G)
Suniv
G
= 
(at constant T and P )
T
 A process (at constant T and P) is spontaneous in the
direction in which the free energy decreases.
 Negative ΔG means positive ΔSuniv.
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Section 17.4
Free Energy
Free Energy (G)
 ΔG = ΔH – TΔS (at constant T and P)
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Section 17.4
Free Energy
CONCEPT CHECK!
A liquid is vaporized at its boiling point. Predict the signs of:
–
w
+
q
+
ΔH
+
ΔS
–
ΔSsurr
0
ΔG
Explain your answers.
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Section 17.4
Free Energy
EXERCISE!
The value of ΔHvaporization of substance X is 45.7 kJ/mol,
and its normal boiling point is 72.5°C.
Calculate ΔS, ΔSsurr, and ΔG for the vaporization of one
mole of this substance at 72.5°C and 1 atm.
ΔS = 132 J/K·mol
ΔSsurr = -132 J/K·mol
ΔG = 0 kJ/mol
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Section 17.4
Free Energy
Spontaneous Reactions
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Section 17.4
Free Energy
Effect of ΔH and ΔS on Spontaneity
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Section 17.5
Entropy Changes in Chemical Reactions
AP Learning Objectives, Margin Notes and References
 Learning Objectives

LO 5.12 The student is able to use representations and models to predict the sign and relative magnitude of the
entropy change associated with chemical or physical processes.
Section 17.5
Entropy Changes in Chemical Reactions
CONCEPT CHECK!
Gas A2 reacts with gas B2 to form gas AB at constant
temperature and pressure. The bond energy of AB is much
greater than that of either reactant.
Predict the signs of:
ΔH
ΔSsurr
–
+
ΔS
ΔSuniv
0
+
Explain.
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Section 17.5
Entropy Changes in Chemical Reactions
Third Law of Thermodynamics
 The entropy of a perfect crystal at 0 K is zero.
 The entropy of a substance increases with temperature.
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Section 17.5
Entropy Changes in Chemical Reactions
Standard Entropy Values (S°)
 Represent the increase in entropy that occurs when a
substance is heated from 0 K to 298 K at 1 atm pressure.
ΔS°reaction = ΣnpS°products – ΣnrS°reactants
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Section 17.5
Entropy Changes in Chemical Reactions
EXERCISE!
Calculate ΔS° for the following reaction:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Given the following information:
S° (J/K·mol)
Na(s)
51
H2O(l)
70
NaOH(aq)
50
H2(g)
131
ΔS°= –11 J/K
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Section 17.6
Free Energy and Chemical Reactions
AP Learning Objectives, Margin Notes and References
 Learning Objectives






LO 5.13 The student is able to predict whether or not a physical or chemical process is thermodynamically favored
by determination of (either quantitatively or qualitatively) the signs of both Ho and So, and calculation or
estimation of Go when needed.
LO 5.14 The student is able to determine whether a chemical or physical process is thermodynamically favorable
by calculating the change in standard Gibbs free energy.
LO 5.15 The student is able to explain how the application of external energy sources or the coupling of favorable
with unfavorable reactions can be used to cause processes that are not thermodynamically favorable to become
favorable.
LO 5.16 The student can use Le Châtelier’s principle to make qualitative predictions for systems in which coupled
reactions that share a common intermediate drive formation of a product.
LO 5.17 The student can make quantitative predictions for systems involving coupled reactions that share a
common intermediate, based on the equilibrium constant for the combined reaction.
LO 5.18 The student can explain why a thermodynamically favored chemical reaction may not produce large
amounts of product (based on consideration of both initial conditions and kinetic effects), or why a
thermodynamically unfavored chemical reaction can produce large amounts of product for certain sets of initial
conditions.
Section 17.6
Free Energy and Chemical Reactions
AP Learning Objectives, Margin Notes and References
 Additional AP References




LO 5.15 (see Appendix 7.11, “Non-Spontaneous Reactions”)
LO 5.16 (see Appendix 7.11, “Non-Spontaneous Reactions”)
LO 5.17 (see Appendix 7.11, “Non-Spontaneous Reactions”)
LO 5.18 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.6
Free Energy and Chemical Reactions
Standard Free Energy Change (ΔG°)
 The change in free energy that will occur if the reactants
in their standard states are converted to the products in
their standard states.
ΔG° = ΔH° – TΔS°
ΔG°reaction = ΣnpG°products – ΣnrG°reactants
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Section 17.6
Free Energy and Chemical Reactions
CONCEPT CHECK!
A stable diatomic molecule spontaneously
forms from its atoms.
Predict the signs of:
ΔH°
ΔS°
–
–
ΔG°
–
Explain.
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Section 17.6
Free Energy and Chemical Reactions
CONCEPT CHECK!
Consider the following system at
equilibrium at 25°C.
PCl3(g) + Cl2(g)
PCl5(g)
ΔG° = −92.50 kJ
What will happen to the ratio of partial
pressure of PCl5 to partial pressure of PCl3 if
the temperature is raised? Explain.
The ratio will decrease.
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Section 17.7
The Dependence of Free Energy on Pressure
AP Learning Objectives, Margin Notes and References
 Learning Objectives

LO 5.18 The student can explain why a thermodynamically favored chemical reaction may not produce large
amounts of product (based on consideration of both initial conditions and kinetic effects), or why a
thermodynamically unfavored chemical reaction can produce large amounts of product for certain sets of initial
conditions.
 Additional AP References

LO 5.18 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.7
The Dependence of Free Energy on Pressure
Free Energy and Pressure
G = G° + RT ln(P)
or
ΔG = ΔG° + RT ln(Q)
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Section 17.7
The Dependence of Free Energy on Pressure
CONCEPT CHECK!
Sketch graphs of:
1. G vs. P
2. H vs. P
3. ln(K) vs. 1/T (for both endothermic and
exothermic cases)
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Section 17.7
The Dependence of Free Energy on Pressure
The Meaning of ΔG for a Chemical Reaction
 A system can achieve the lowest possible free energy by
going to equilibrium, not by going to completion.
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Section 17.8
Free Energy and Equilibrium
AP Learning Objectives, Margin Notes and References
 Learning Objectives


LO 5.18 The student can explain why a thermodynamically favored chemical reaction may not produce large
amounts of product (based on consideration of both initial conditions and kinetic effects), or why a
thermodynamically unfavored chemical reaction can produce large amounts of product for certain sets of initial
conditions.
LO 6.25 The student is able to express the equilibrium constant in terms of Go and RT and use this relationship to
estimate the magnitude of K and, consequently, the thermodynamic favorability of the process.
 Additional AP References


LO 5.18 (see Appendix 7.11, “Non-Spontaneous Reactions”)
LO 6.25 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.8
Free Energy and Equilibrium
 The equilibrium point occurs at the lowest value of free
energy available to the reaction system.
ΔG = 0 = ΔG° + RT ln(K)
ΔG° = –RT ln(K)
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Section 17.8
Free Energy and Equilibrium
Change in Free Energy to Reach Equilibrium
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Section 17.8
Free Energy and Equilibrium
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Section 17.9
Free Energy and Work
AP Learning Objectives, Margin Notes and References
 Learning Objectives

LO 5.15 The student is able to explain how the application of external energy sources or the coupling of favorable
with unfavorable reactions can be used to cause processes that are not thermodynamically favorable to become
favorable.
 Additional AP References

LO 5.15 (see Appendix 7.11, “Non-Spontaneous Reactions”)
Section 17.9
Free Energy and Work
 Maximum possible useful work obtainable from a
process at constant temperature and pressure is equal
to the change in free energy.
wmax = ΔG
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Section 17.9
Free Energy and Work
 Achieving the maximum work available from a
spontaneous process can occur only via a hypothetical
pathway. Any real pathway wastes energy.
 All real processes are irreversible.
 First law: You can’t win, you can only break even.
 Second law: You can’t break even.
 As we use energy, we degrade its usefulness.
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