Chapter 18: Thermodynamics
Renee Y. Becker
Valencia Community College
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Thermodynamics
Thermodynamics: The study of interconversion of heat
and other forms of energy
Spontaneous process: Process that proceeds on its own
without external influence (-G)
Non-spontaneous: Needs continuous external influence
to take place (+G)
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Thermodynamics
Entropy (S): Disorder, molecular randomness
S = Sfinal - Sinitial
When disorder increases +S
When disorder decreases -S
Enthalpy (H): Heat flow
In to the system +H
Out of the system -H
Gibbs Free-Energy: Measure of spontaneity
G = H-TS
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Example 1:
Predict the sign of S in the system for each of the
following
a)
H2O(g)  H2O(l)
b)
I2(g)  2I-(g)
c)
CaCO3(s)  CaO(s) + CO2(g)
d)
Ag+(aq) + Br-(aq)  AgBr(s)
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Example 2:
Which of the following reactions has an
increase in entropy?
1. H2O(g)  H2O(l)
2. H2O(l)  H2O(g)
3. H2O(g)  H2O(s)
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Entropy and temperature
3rd Law of Thermodynamics
a) The entropy of a perfectly ordered
crystalline substance at 0 K is zero
b) As the temperature increases, the KE
increases, Molecular motion increases,
entropy increases
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Standard Molar Entropies and Standard Entropies of Reaction
Standard Molar Entropy, S
The entropy of one mole of the pure substance at 1 atm
pressure and a specific temperature usually 25C
Standard Entropy of Reaction, S
Entropy change for a chemical reaction
S = Sproducts - Sreactants
Based on 1 mole of substance so you have to multiply S by the
number of moles present
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Standard Entropy of Reaction, S
aA + bB  cC + dD
S = [c S(C) + d S(D)] – [a S(A) + b S(B)]
Units: coefficients are moles
S = J/K mol
S = J/K
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Example 3:
Calculate the standard entropy of reaction at 25C
for the decomposition of calcium carbonate
CaCO3(s)  CaO(s) + CO2(g)
Substance
CaCO3(s)
CaO(s)
CO2(g)
S (J/K mol)
92.9
39.7
213.6
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Entropy and the Second Law of Thermodynamics
1st Law of Thermodynamics
In any process, spontaneous or nonspontaneous,
the total energy of a system and its surroundings is
constant
2nd Law of Thermodynamics
In any spontaneous process, the total entropy of a
system and its surroundings always increases
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Entropy and the Second Law of Thermodynamics
Stotal = Ssystem + Ssurroundings
if S > 0 spontaneous
if S< 0 non spontaneous
if S = 0 equilibrium
Ssurr = -H / T
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Entropy and the Second Law of Thermodynamics
a)
b)
Exothermic reaction: H<0, because the surroundings
gain heat (entropy increases), heat is lost from the system
Endothermic reaction: H>0, surroundings lose heat
(entropy decreases), and system gains the heat
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Example 4:
Which of the following reactions is endothermic?
1. N2O4(g)  2 NO2(g)
H = +57.1 kJ
2. 2 NO2(g)  N2O4(g)
H = -57.1 kJ
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Free-Energy
Free energy, G
G = H – TS
G = H - TS
if G < 0 spontaneous Only direct relationship that
if G > 0 nonspontaneous
always holds up!
if G = 0 equilibrium
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Example 5:
Consider the decomposition of gaseous N2O4
N2O4(g)  2 NO2(g)
H = +57.1 kJ
S = +175.8 J/K
a) Is this reaction spontaneous under standard-state
conditions at 25C?
b) Estimate the temperature at which the reaction
becomes spontaneous
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Standard Free-Energy Changes for Reactions
1. Standard State Conditions: Solids, liquids, and
gases in pure form at 1 atm pressure, Solutes at
1M concentration, specified temperature, usually
25 celsius
2. Standard Free Energy Change, G: The change
in free energy that occurs when reactants in their
SS are converted to products in their SS
3. G = H - TS
G = H - TS
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Example 6:
Consider the thermal decomposition of calcium
carbonate
CaCO3(s)  CaO(s) + CO2(g)
a)
b)
c)
H = 178.3 kJ
S = 160.4 J/K
Calculate the standard free energy change for this reaction
at 25C
Will a mixture of solid CaCO3, CaO, and gaseous CO2 at 1
atm pressure react spontaneously at 25C?
Assuming that H and S are independent of
temperature, estimate the temperature at which the
reaction becomes spontaneous
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Standard Free Energies of Formation
1. Standard Free Energy of Formation, Gf
The free energy change for formation of one mole of
the substance in its standard state from the most
stable form of its constituent elements in their
standard states
2. Gf measures the substances thermodynamic
stability with respect to its constituent elements
3. -Gf are stable and do not decompose to their
constituent elements under standard state conditions
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Standard Free Energies of Formation
4.
+Gf are thermodynamically unstable with respect to their
constituents elements
a) There is no point in trying to synthesize a substance that
has a +Gf because it would degrade into it’s constituents
b) You would need to synthesize it at different temperatures
and or pressures or start with different starting materials that
has a reaction with a -Gf
5.
G = Gf(products) - Gf(reactants)
6. General reaction: aA +bB  cC + dD
G = [cGf(C) + dGf(D)] – [aGf(A) + bGf(B)]
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Example 7:
Calculate the standard free energy change for the reaction
of calcium carbide with water. Might this reaction be
used for the synthesis of acetylene (C2H2)?
CaC2(s) + 2 H2O(l)  C2H2(g) + Ca(OH)2(s)
Gf (CaC2) = -64.8 kJ/mol
Gf (H2O(l)) = -237.2 kJ/mol
Gf (C2H2) = 209.2 kJ/mol
Gf (Ca(OH)2) = -898.6 kJ/mol
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Free Energy Changes and Composition of the Reaction Mixture
Standard state conditions are unrealistic, the reaction
itself will change the temperature and pressure, so
what can we use to calculate the free energy
change under non-standard state conditions?
G = G + RT ln Q
R = gas constant
T = temperature in Kelvins
Q = reaction quotient (Qc or Qp)
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Example 8:
Calculate the Free energy change for the formation of
ethylene (C2H4) from carbon and hydrogen at 25C
when the partial pressures are 100 atm H2 and 0.10 atm
C2H4
2 C(s) + 2 H2(g)  C2H4(g)
G = 68.1 kJ/mol
Is the reaction spontaneous in the forward or reverse
direction?
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Free Energy and Chemical Equilibrium
1. When the RM is mostly reactants
Q<<1
RT lnQ <<0
G<0
2. When the RM is mostly products
Q>>1
RT lnQ >>0
G>0
3. G= -RT ln K
K = equilibrium constant Kc or Kp
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Example 9:
Calculate the Kp at 25C for the reaction
CaCO3(s)  CaO(s) + CO2(g)
G= 130.5 kJ/mol
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Example 10:
Which of the following can always predict the
spontaneity of a reaction?
1. H
2. S
3. G
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