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Chapter 19. Chemical Thermodynamics
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 19.4 Reversible Flow of Heat 19.1 Spontaneous Processes
Figure 19.5 An Irreversible Process 19.1 Spontaneous Processes
Figure 19.6 Possible Arrangements of Two Gas 19.3 Molecular Interpretation of Entropy
Molecules in Two Flasks
Activities: Section:
Molecular Motions 19.3 Molecular Interpretation of Entropy
Mixing of Gases 19.3 Molecular Interpretation of Entropy
Animations: Section:
Air Bags 19.1 Spontaneous Processes
Movies: Section:
Formation of Water 19.5 Gibbs Free Energy
3-D Models: Section:
Oxygen 19.1 Spontaneous Processes
Ice 19.1 Spontaneous Processes
Water 19.3 Molecular Interpretation of Entropy
Chemical Thermodynamics
275
Other Resources
Further Readings: Section:
Demystifying Introductory Chemistry. Part 4: 19.1 Spontaneous Processes
An Approach to Reaction Thermodynamics
through Enthalpies, Entropies, and Free
Thermodynamics
Entropy: Conceptual Disorder 19.2 Entropy and the Second Law of
Thermodynamics
Spontaneous Assembly of Soda Straws 19.2 Entropy and the Second Law of
Thermodynamics
Give Them Money: The Boltzmann Game, a 19.3 Molecular Interpretation of Energy
Classroom or Laboratory Activity Modeling
Entropy Changes and the Distribution of
Live Demonstrations: Section:
Entropy, Disorder, and Freezing 19.2 Entropy and the Second Law of
Thermodynamics
A Chemical Hand Warmer 19.2 Entropy and the Second Law of
Thermodynamics
Chapter 19
276
Chapter 19. Chemical Thermodynamics
Common Student Misconceptions
• Students often believe that a spontaneous process should occur very quickly. They do not appreciate
the difference between kinetics and thermodynamics.
Teaching Tips
• Equation 19.15 is valid only when H° and S° do not significantly change with temperature and
pressure.
Lecture Outline
19.1 Spontaneous Processes
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• Chemical thermodynamics is concerned with energy relationships in chemical reactions.
• We consider enthalpy.
• We also consider entropy in the reaction.
• Recall the first law of thermodynamics: energy is conserved.
E= q + w
Chemical Thermodynamics
277
Seeking a Criterion for Spontaneity
• To understand why some processes are spontaneous we must look at the ways in which the state of a
system might change.
• Temperature, internal energy, and enthalpy are state functions.
• Heat transferred between a system and the surroundings, as well as work done on or by a system,
are not state functions.
Reversible and Irreversible Processes
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• A reversible process is one that can go back and forth between states along the same path.
• The reverse process restores the system to its original state.
• The path taken back to the original state is exactly the reverse of the forward process.
• Remove the partition, and the gas expands to fill the space.
• No P-V work is done on the surroundings.
• w = 0
• Now use the piston to compress the gas back to the original state.
• The surroundings must do work on the system.
• w > 0
• A different path is required to get the system back to its original state.
• Note that the surroundings are NOT returned to their original conditions.
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• Each flask contains gas at 0.5 atm.
• Therefore, the gas does no work and heat is not transferred.
• Why does the gas expand?
• Why is the process spontaneous?
19.2 Entropy and the Second Law of Thermodynamics
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Entropy Change
• Entropy, S, is a thermodynamic term that reflects the disorder, or randomness, of the system.
• The more disordered, or random, the system is, the larger the value of S.
• Entropy is a state function.
• It is independent of path.
• For a system, S = Sfinal – Sinitial.
• If S > 0 the randomness increases, if S < 0 the order increases.
The Second Law of Thermodynamics
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• The second law of thermodynamics:
• any irreversible process results in an increase in total entropy while any reversible process results
in no overall change in entropy.
• This explains why spontaneous processes have a direction.
• In any spontaneous process, the entropy of the universe increases.
• The change in entropy of the universe is the sum of the change in entropy of the system and the
change in entropy of the surroundings.
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• For a reversible process:
Suniv = Ssystem + Ssurroundings = 0
19.3 Molecular Interpretation of Entropy
Expansion of a Gas at the Molecular Level
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• Gas expansion into a vacuum is a spontaneous process.
• Consider two flasks connected by a stopcock.
• Track the movement of two gas molecules as they move around.
• Before opening the stopcock: both molecules are confined to the left flask.
• After opening the stopcock: the molecules move randomly throughout the entire apparatus.
Boltzmann’s Equation and Microstates
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• Statistical thermodynamics is a field that uses statistics and probability to link the microscopic and
macroscopic worlds.
• Entropy may be connected to the behavior of atoms and molecules.
• Envision a microstate: a snapshot of the positions and speeds of all molecules in a sample of a
particular macroscopic state at a given point in time.
• Consider a molecule of ideal gas at a given temperature and volume.
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• The Boltzmann equation shows how entropy (S) relates to W.
S = k lnW, where k is Boltzmann’s constant (1.38 10–23 J/K).
• Entropy is thus a measure of how many microstates are associated with a particular macroscopic
state.
Molecular Motions and Energy
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• When a substance is heated, the motion of its molecule increases.
• The higher the temperature, the faster the molecules move.
• Hotter systems have broader distribution of molecular speeds.
• Consider a sample of ideal gas.
• The molecules move around the container.
• They also show three kinds of more complex motion:
system.
Making Qualitative Predictions About S
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• In most cases, an increase in the number of microstates (and thus entropy) parallels an increase in:
• temperature
• volume
• number of independently moving particles.
Chemical Thermodynamics
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• Some must be used to hydrate the ions.
• Thus this example involves both ordering and disordering.
• The disordering usually predominates (for most salts).
• Consider the reaction of NO(g) with O2(g) to form NO2(g):
2NO(g) + O2(g) → 2NO2(g)
• The total number of gas molecules decreases.
The Third Law of Thermodynamics
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• In a perfect crystal at 0 K there is no translation, rotation, or vibration of molecules.
• Therefore, this is a state of perfect order.
• Third law of thermodynamics: The entropy of a perfect pure crystal at 0 K is zero.
• Entropy will increase as we increase the temperature of the perfect crystal.
19.4 Entropy Changes in Chemical Reactions
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• Absolute entropy can be determined from complicated measurements.
• Values are based on a reference point of zero for a perfect crystalline solid at 0K (the 3rd law).
• Standard molar entropy, S° is the molar entropy of a substance in its standard state.
• Similar in concept to H°.
• Units: J/mol-K.
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Entropy Changes in the Surroundings
• For an isothermal process,
• For a reaction at constant pressure,
• qsys = H
19.5 Gibbs Free Energy
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• For a spontaneous reaction the entropy of the universe must increase.
• Reactions with large negative H values tend to be spontaneous.
• How can we use S and H to predict whether a reaction is spontaneous?
• The Gibbs free energy, (free energy), G, of a state is:
G = H – TS
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“Josiah Willard Gibbs and Wilhelm Ostwald: A Contrast in Scientific Style” from Further Readings
35
“J. Willard Gibbs (1839–1903): A Modern Genius” from Further Readings
36
Chemical Thermodynamics
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• The equilibrium position in a spontaneous process is given by the minimum free energy available to
the system.
• The free energy decreases until it reaches this minimum value.
Standard Free-Energy Changes
• We can tabulate standard free energies of formation, ∆G°f.
• Standard states are pure solid, pure liquid, 1 atm (gas), 1 M concentration (solution), and ∆G°f = 0
for elements.
FORWARD REFERENCES
• Standard conditions (Table 19.3) will be brought up throughout Chapter 20.
• ∆G as a measure of reaction spontaneity will be linked to the cell potential in Chapter 20
(section 20.5).
• Entropy of the chelate effect will be further discussed in Chapter 23 (section 23.3).
19.6 Free Energy and Temperature
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• The sign of G tells us if the reaction is spontaneous, at a given temperature and pressure condition.
• Focus on G = H – TS.
• If H < 0 and –TS < 0:
• G will always be < 0.
• At a temperature less than 0 °C:
• H > TS
• G > 0
• The melting of ice is not spontaneous when the temperature is less than 0 oC.
• At a temperature greater than 0 °C:
• H < TS
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• Thermodynamics gives us the direction of a spontaneous process; it does not give us the rate of
the process.
19.7 Free Energy and the Equilibrium Constant
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Free Energy Under Nonstandard Conditions
• Recall that G° and Keq (equilibrium constant) apply to standard conditions.
• Recall that G and Q (equilibrium quotient) apply to any conditions.
• It is useful to determine whether substances will react under specific conditions:
G = G° + RTlnQ
Relationship Between ∆G˚ and K
• At equilibrium, Q = Keq and G = 0, so:
Driving Nonspontaneous Reactions
• If G > 0, work must be supplied from the surroundings to drive the reaction.
• Biological systems often use one spontaneous reaction to drive another nonspontaneous reaction.
• These reactions are coupled reactions.
• The energy required to drive most nonspontaneous reactions comes from the metabolism of foods.
• When ATP is converted back to ADP the energy released may be used to “drive” other reactions.
FORWARD REFERENCES
• Units of Gibb’s free energy change in calculations linking G with the cell potential will be
discussed in Chapter 20 (section 20.5).
• Gibb’s free energy change will be used to derive the Nernst equation in Chapter 20 (section
20.6).
• Equation 19.17 will be used in Chapter 22 (section 22.5) to calculate the equilibrium constant
“Sodium Chloride (1 1 Unit Cell)” 3-D Model from Instructor’s Resource CD/DVD
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Further Readings:
1. James N. Spencer, Richard S. Moog, and Ronald J. Gillespie, “Demystifying Introductory Chemistry.
Part 4: An Approach to Reaction Thermodynamics through Enthalpies, Entropies, and Free Energies of
Atomization,” J. Chem. Educ., Vol. 73, 1996, 631–636.
2. Raymond S. Ochs, “Thermodynamics and Spontaneity,” J. Chem. Educ., Vol. 73, 1996, 952–954.
7. John J. Fortman, “Pictorial Analogies III: Heat Flow, Thermodynamics, and Entropy,” J. Chem. Educ.,
Vol. 70, 1993, 102–103.
8. L. Glasser, “Order, Chaos, and All That!” J. Chem. Educ., Vol. 66, 1989, 997–1001.
9. John P. Lowe, “Entropy: Conceptual Disorder,” J. Chem. Educ., Vol. 65, 1988, 403–406.
13. Douglas K. Russell, “The Boltzmann Distribution,” J. Chem. Educ., Vol. 73, 1996, 299–300.
14. Travis Thoms, “Periodic Trends for the Entropy of Elements,” J. Chem. Educ., Vol. 72, 1995, 16.
15. Sidney Rosen, “J. Willard Gibbs (1839–1903): A Modern Genius,” J. Chem. Educ., Vol. 60, 1983,
593–594.
16. Robert J. Deltete and David L. Thorsell, “Josiah Willard Gibbs and Wilhelm Ostwald: A Contrast in
Scientific Style,” J. Chem. Educ., Vol. 73, 1996, 289–295.
Chapter 19
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Live Demonstrations:
1. Lee R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “A Chemical Hand Warmer,” Chemical
Demonstrations: A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
pp. 99–100.
