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Chapter 19
Chemical Thermodynamics
19.1
Chemical Thermodynamics
Area of chemistry that studies energy relationships
How fast is the reaction?
Kinetics -- energy
How far does the reaction proceed?
Equilibrium – energy
O Thermodynamics will not predict kinetics
O Thermodynamics will predict equilibrium
Spontaneous Process
One that proceeds without any outside
assistance; spontaneous in one direction but
not the opposite
Examples:
O Na and Cl2 make NaCl
O Brick falls off ledge
O Gas expands into container
O Rusting nail
Spontaneous Process
O Does not mean it occurs at a rate that you
can see or notice
O Rusting nail  very slow
O Ice melting  pretty fast
O Energy is lost in spontaneous process
O Can be endothermic or exothermic
O irreversible
Reversible Process
O Both the process and surroundings can be
simultaneously returned to their initial state
after the process has been completed
O Ideal
O Infinitesimal changes – small ΔT & heat flow
O all real processes are irreversible
19.2
Entropy
Associated with system randomness
O S
Phase changes
O State function
occur at constant T
ΔHfus
O If isothermal (constant T)
ΔS =
qrev
Tconstant
Crash course chemistry – entropy
https://www.youtube.com/watch?v=ZsY4WcQOrfk
ΔHvap
2nd Law of Thermodynamics
entropy of the system and surroundings will
increase for an irreversible process
ΔSuniverse > 0
ΔSuniverse = 0
for spontaneous process
for reversible process
Example
The normal boiling point of ethanol is 78.3ºC and the
molar enthalpy of vaporization is 38.56 kJ/mol. What is
the change in entropy of the system when 68.3 g of
ethanol at 1 atm is condensed to liquid?
-163 J/K
Suggested Homework
pp. 833-836
(1, 9, 13, 15, 23, 25)
19.3
Molecular Interpretation
O KMT – ideal gas
O Average KE direction proportional to
absolute temperature (higher T = more KE)
O More KE = more motion
O Translational
O Vibrational
O Rotational
Molecular Interpretation
O Energy of ideal gas depends only on
temperature
Microstates
snapshot of molecules at point in time; single
possible state of arrangements and energies
of molecules
Molecular Interpretation
Entropy is a measure of the number of
microstates of the system.
Entropy increases with the number of
microstates of the system.
Isothermally increase volume  more arrangements
Increase temperature (fixed volume)  more kinetic energies
Increase molecules (not ideal)  more kinetic energies
Qualitative Predictions
To predict the change in entropy consider what is
happening to:
1. Temperature
2. Volume
3. Number of moving particles
Examples:
O Melting
O Evaporation
O number of gas molecules increase
3rd Law of Thermodynamics
The entropy of a pure
crystalline substance
at absolute zero is
zero.
Example
Predict an increase or decrease in entropy.
1. CO2(s)  CO2(g)
2. CaO(s) + CO2(g)  CaCO3(s)
3. HCl(g) + NH3(g)  NH4Cl(s)
4. 2 SO2(g) + O2(g)  2 SO3(g)
Increase
Decrease
Decrease
Decrease
Example
Which sample has the highest entropy?
Assume 1 mole of sample unless otherwise specified.
1. H2(g) at STP or at 100°C and 0.5 atm
2. Water at freezing temperature or room
temperature
3. H2(g) or SO2(g) at STP
4. 1 mol H2(g) or 2 mol H2(g) at STP
100°C and 0.5 atm
Room temperature
SO2(g)
2 mol H2(g)
Suggested Homework
pp. 834-836
(3, 4, 29, 33, 39, 41)
Entropy in Reactions
ΔS° = Σ n S° (products) - Σ m S° (reactants)
O Measure heat capacity to determine ΔS
O Standard molar entropy, S°, p.817/Appendix C
Notable Trends:
O Increase with molar mass
O Increase with number of atoms in substance
Entropy in Reactions
O If entropy of reaction system decreases,
assume entropy gained by surroundings is
larger
O Entropy of universe must increase if
spontaneous
O If exothermic, entropy of surroundings
always increases (unless isolated system)
Example
Calculate ΔS° for the synthesis of ammonia.
The values of S° are 192.5 kJ/mol for ammonia,
191.5 kJ/mol for nitrogen, and 130.6 kJ/mol for
hydrogen.
-198.3 kJ/mol (decreases)
Suggested Homework
pp. 836-837
(49)
Gibbs Free Energy
O Spontaneous processes that decrease the
system’s entropy are always exothermic
O Entropy and enthalpy must be related
O New state function
Gibbs Free Energy
G=H–TS
ΔG° = ΔH° – TΔS°
Gas @ 1 atm
Solution @ 1 M
25°C
Gibbs Free Energy
ΔSuniv = ΔSsys + ΔSsurr
ΔSuniv = ΔSsys + -ΔHsys/T
If spontaneous
ΔSuniv + positive
ΔGsys - negative
-TΔSuniv = ΔHsys - TΔSsys
ΔGsys = ΔHsys - TΔSsys
Gibbs Free Energy
ΔG sign
Negative Reaction is spontaneous in forward
direction
(supply work to make it go in reverse)
Zero
Reaction is at equilibrium
Positive
Reaction is spontaneous in reverse
direction
(supply work to make it go forward)
Gibbs Free Energy
O Free energy changes
until it reaches a
minimum value
O Free energy always
decreases for a
spontaneous
process
Free Energy in Reactions
ΔG° = Σ n Gf° (products) - Σ m Gf° (reactants)
O Standard free energies of formation, Gf° are in
Appendix C
O Can use free energies of formation to
calculate free energy of reaction
O Can also use Hess’s law
Example
A reaction occurs at 298K and has ΔH° = 24.6
kJ and ΔS° = 132 J/K. Calculate ΔG° and
determine if the reaction is spontaneous under
these conditions.
-14.7 kJ
Example
P4(g) + 6 Cl2(g)  4 PCl3(g)
The free energy of formation for P4 is 24.4 kJ/mol and
for PCl3 is -269.6 kJ/mol. What is the free energy
change for this reaction?
-1102.8 kJ
Example
Would you expect the standard free energy
change for the combustion of propane gas,
C3H8(g), to be more or less negative than the
standard enthalpy change?
More negative
Free Energy & Temperature
O Enthalpy and entropy only change small
amounts with temperature
O T in free energy equation dominates
Free Energy & Temperature
ΔH
ΔS
-TΔS
ΔG
-
+
-
-
Always spontaneous
+
+
+
Never spontaneous
-
-
+
+ or -
Spontaneous at low T
+
+
-
+ or -
Spontaneous at high T
Depends on
bond
energies
and IMF
Depends on
microstates
Remember:
Can be very,
very slow
Suggested Homework
pp. 834, 837-838
(5, 6, 8, 53, 57, 59, 61)
Free Energy & Equilibrium
For nonstandard conditions:
ΔG = ΔG° + RT ln Q
O If at equilibrium, Q = 1 and ln Q = 0
0 = ΔG° + RT ln K
K = e -ΔG°/RT
If free energy is
negative, K is
greater than 1
Important Nonspontaneous
Reactions
Must find a way to do work on the reaction
O In body, glucose combustion is very exergonic
(negative free energy)
O Plants use light to complete photosynthesis
O ATP  ADP is exergonic
O Need to convert ADP back to ATP
O Couple the reaction with ones that are spontaneous
Bozeman – life requires free energy
https://www.youtube.com/watch?v=JBmykor-2kU
Example
Indicate if ΔG increases, decreases, or stays the
same when the partial pressure of hydrogen is
increased in the following reactions:
N2(g) + H2(g)  2 NH3(g)
2 HBr(g)  H2(g) + Br2(g)
2 H2(g) + C2H2(g)  C2H6(g)
Example
Hydrogen gas and bromine liquid reaction to form
gaseous hydrogen bromide. The free energy of
formation for hydrogen bromide is -53.22 kJ/mol.
1. Calculate the standard free energy change for the
reaction
2. Calculate the equilibrium constant for the reaction
at 298 K
-106.4 kJ/mol
4 x 1018
Example
Hydrogen gas and bromine liquid reaction to form
gaseous hydrogen bromide. The equilibrium constant
for the reaction is 4 x 1018.
1. Calculate the free energy change at 200°C
2. Calculate the free energy change if the partial
pressure of HBr is 1.50 atm and hydrogen is 0.500
atm at 298 K. The standard free energy change is
106.4 kJ/mol
-2 x 105 J/mol
-104 kJ/mol
Suggested Homework
pp. 834, 837-838
(7, 73, 75, 79)
20.5
Free Energy and Redox
O Voltaic cells are spontaneous
E° = E°red (red) - E°red (ox)
O Positive value of E° is spontaneous
O Activity series
20.5
Free Energy and EMF
ΔG° = -n F E°
O n is the moles of electrons
O Faraday’s constant is charge on 1 mole of eO F = 96,485 C/mol or J/V-mol
O Positive E  negative ΔG
ΔG° = -RT ln K
Example
Use standard reduction potentials in Table
20.1 (p. 857) to calculate ΔGº and the
equilibrium constant at 298 K for the following
reaction.
4 Ag(s) + O2(g) + 4 H+(aq)  4 Ag+(aq) + 2 H2O(l)
Example
Use standard reduction potentials in Table
20.1 (p. 857) to calculate ΔGº and the
equilibrium constant at 298 K for the following
reaction.
1
2
2 Ag(s) + O2(g) + 2 H+(aq)  2 Ag+(aq) + H2O(l)
Suggested Homework
p. 887
(51, 53)
Integrative Exercise p.831