CHM 1046: General Chemistry and
Qualitative Analysis
Unit 14:
Chemical
Thermodynamics
Dr. Jorge L. Alonso
Miami-Dade College
– Kendall Campus
Miami, FL
Textbook Reference:
•Chapter # 15 (sec.12-17)
Chemical
•Module # 3 (sec.Thermodynamics
VIII-XII)
First Law of
Thermodynamics
• The law of conservation of
energy: energy cannot be
created nor destroyed. (James
Joule in 1843 )
• Energy can, however, be converted
from one form to another or transferred
from a system to the surroundings or
vice versa.
E = q + w
w
=
PV
Esys + Esurr = 0
E = q + PV
Esys = -Esurr
E = q + RT
• Therefore, the total energy of
the universe is a constant.
Chemical
Thermodynamics
Spontaneity
E1
First Law of
Thermodynamics
Stone
Do all processes that loose
energy occur
spontaneously (by
themselves, without
external influence)??????
+ Work
E2
Second Law of
Thermodynamics
- (work + heat)
E = E2 – E1
Chemical
Thermodynamics
Spontaneous Processes
• can proceed without any outside intervention.
{Spontaneity}
Processes that
are spontaneous
in one direction
are
nonspontaneous
in the reverse
direction.
Chemical
Thermodynamics
Spontaneous Processes
• Processes that are spontaneous at one temperature
may be nonspontaneous at other temperatures.
• Above 0C it is spontaneous for ice to melt.
• Below 0C the reverse process is spontaneous.
Spontaneous @ T < 0ºC
Is the
spontaneity of
melting ice
dependent on
anything?
Spontaneous @ T > 0ºC
Chemical
Thermodynamics
Spontaneity vs. Speed
C diamond
C graphite
Thermodynamics vs. Kinetics
Chemical
Thermodynamics
Irreversible Processes
E1
Stone
• Irreversible processes cannot be undone
by exactly reversing the change to the
system.
• Heat energy is lost to dissipation and that
energy will not be recoverable if the
process is reversed.
+ Work • Spontaneous processes are irreversible.
Reversible Processes
In a reversible process the system changes
in such a way that the system and
Chemical
E2
surroundings can be put back in theirThermodynamics
original
- (work + heat) states by exactly reversing the process.
Entropy (S)
Entropy is a measure of the energy
that becomes dissipated and unavailable
(friction, molecular motion = heat).
• Entropy (S) is a term coined by Rudolph Clausius in the 1850’s.
Clausius chose "S" in honor of Sadi Carnot (who gave the first successful
theoretical account of heat engines, now known as the Carnot cycle, thereby laying the foundations of the
second law of thermodynamics).
• Clausius was convinced of the significance of the ratio of heat
delivered and the temperature at which it is delivered,
Entropy (S) =
q
T
Chemical
Thermodynamics
Entropy (S)
Gas
• Entropy can be thought of
as a measure of the
randomness (disorder) of
a system.
• It is related to the various
modes of motion in
molecules.
• Like total energy, E, and
enthalpy, H, entropy is a
state function.
• Therefore,
S = Sfinal  Sinitial
{Entropy.WaterBoiling}
Liquid
Solid
E
N
T
R
O
P
Y
Chemical
Thermodynamics
Second Law of
Thermodynamics
• the entropy of the universe increases for
spontaneous (irreversible) processes.
Suniv = Ssystem + Ssurroundings > 0
• the entropy of the universe does not
change for reversible processes.
Suniv = Ssystem + Ssurroundings = 0
Chemical
Thermodynamics
Second Law of Thermodynamics
All spontaneous processes cause the
entropy of the universe to increase.
So what is our fate as a result of the
second law operating in our Universe?
ENTROPIC DOOM!
Chemical
Thermodynamics
Entropy on the Molecular Scale
• Molecules exhibit several types of motion (Kinetic energies):
 Translational: Movement of a molecule from one place to another.
 Vibrational: Periodic motion of atoms within a molecule.
 Rotational: Rotation of the molecule on about an axis or rotation about
 bonds.
• Boltzmann envisioned the motions of a sample of molecules at a
particular instant in time.
 This would be akin to taking a snapshot of all the
molecules.
 He referred to this sampling as a microstate (W) of the
thermodynamic system
• Entropy is …….
S = k ln W
…..where k is the Boltzmann constant, 1.38  1023 J/K.
Chemical
Thermodynamics
Entropy on the Molecular Scale
S = k ln W
…..where k is the Boltzmann constant, 1.38  1023 J/K.
• The number of microstates (W) and, therefore,
the entropy (S) tends to increase with increases
in which variables….
Temperature (T).
Volume (V).
The number of independently moving
molecules ().
Chemical
Thermodynamics
Entropy Changes
• In which of the following does Entropy increase & WHY?…….
 Gases are formed from liquids and solids.
•
Entropy increases with the freedom of motion of molecules.
Heat
H2O (l)
H2O(g)
S(g) > S(l) > S(s)
{*Entropy&PhaseOfMatter}
 Liquids or solutions are formed from solids.
CaCl2 (s)
H2O
Ca
2+
-
(aq) + 2Cl (aq)
{*EntropySolutions.KMnO4(aq)}
 The number of gas molecules (or moles) increases.
2 H2O (l)
Electricity
2 C8H18 (l) + 25 O2 (g)
2 H2 (g) + O2(g)
16 CO2(g) + 18 H2O(g)
 gas= 34-25 = +9/2 = 4.5 C8H18
Chemical
Thermodynamics
Third Law of
Thermodynamics
The entropy (S) of a pure crystalline substance at
absolute zero (-273°C) is 0.
Chemical
Thermodynamics
Standard Entropies
• Standard entropies tend to increase with increasing
molar mass.
• Larger and more complex molecules have greater
entropies (greater ways to execute molecular
motions)
Chemical
Thermodynamics
{*Entropy&MolecuarSize}
{Entropy&Temp}C7H15 @ 500 K S=921J/nK vs, @ 200 K
Standard Entropies (298 K)
from Absolute Entropies (0K)
Absolute Entropy (S)
Solid
Liquid
Gas
@ - 237°C (0 K), S = 0
q = mcT
Standard Entropy (S˚)
@ 25°C (298 K), S = ????
H°vap
q mcT CT
S  

T
T
T
S 
T  298 K
CT
T
T 0

 0
298
C
dT
T
S
q = mcT
S°
Calculate the sum of all the infinitesimally
small changes in entropy as T varies from
T=0  T= 298, by taking its Integral.
H°fus
q = mcT
Chemical
Thermodynamics
298 Temp (K)
Entropy Changes in the System
S°syst
=
S°rxn @ T
Entropy changes for a reaction (= system) can be
estimated in a manner analogous to that by which H
is estimated:
S
o
298
  nS
o
products
  mS
o
reactants
where n and m are the coefficients in the balanced
chemical equation.
Chemical
Thermodynamics
Entropy Changes
in the System
Problem: Calculate the standard
entropy changes for the following
reaction at 25oC.
N2 (g) + 3 H2 (g)  2 NH3 (g)
S° = nS°(prod) - mS°(react)
2(192.5) – [(191.5)+3(130.6)]
S° = - 198.3 J/
Chemical
Thermodynamics
Thermodynamic Changes in
Systems (Chem. Reactions)
Hrxn =  Hf (products) -  Hf (reactants)
So298   Soproducts   Soreactants
☺
Grxn =  Gf (products)   Gf (reactants)
Appendix 1 (CHM 1046 Module): notice S° is in J not kJ.
Chemical
Thermodynamics
Entropy Changes in the Surroundings
What in a chemical reaction causes
entropy changes in the surroundings?
• Heat (q) that flows into or out
of the system changes the
entropy of the surroundings:
q
q
q
q
System
Ssurr ∝ - (qsys)
• For an isothermal process:
(qsys)
Ssurr=
T
• At constant pressure, qsys is
simply H for the system.
q
q
q
Surroundings
Ssurr=
Hsys
T
Chemical
Thermodynamics
Entropy Change in the Universe
Problem: Calculate the Suniv for the synthesis of ammonia @ 25oC.
N2 (g) + 3 H2 (g)  2 NH3 (g)
Suniv =
H°rxn = - 92.6 kJ/mol
Ssyst or rxn
Ssurr
+
nS(prod) - mS(react)
Ssurr =
2(192.5) – [(191.5)+3(130.6)]

S°syst = - 199 J/K·mol
Suniv =
Suniv =
- 198.3 J/K·mol
113 J/K·mol
-Hsys
T
 (92.6 kJ / mol x 1000 J / kJ)
298 K
Ssurr = 311 J/K·mol
+
Chemical
311 J/K·mol
Thermodynamics
Entropy Change in the Universe
Suniv =
Ssyst or rxn
• Then:
Suniv = Ssyst +
+
Ssurr
• Since: Ssurr =
-Hsys
T
Hsystem
T
Multiplying both sides by T,
 TSuniv =  TSsyst + Hsyst
TSuniv = Hsyst  TSsyst
TSuniv is defined as the
Gibbs (free) Energy, G.
J. Willard Gibbs
Chemical
USA, 1839-1903
Thermodynamics
Gibbs Free Energy (G)
TSuniv = Guniv = Hsys  TSsys
☺
• When Suniv is positive, G is negative.
• When G is negative, the process is spontaneous.
Gibbs Energy (-TΔS) measures the "useful" or process-initiating work
obtainable from an isothermal, isobaric thermodynamic system. Technically, the
Gibbs free energy is the maximum amount of non-expansion work whichChemical
can be
extracted from a closed system or this maximum can be attained onlyThermodynamics
in a
completely reversible process.
Free Energy Changes
At temperatures other than 25°C,
G° = H  TS ☺
How does G change with temperature?
• There are two parts to the free energy equation:
 H— the enthalpy term
 TS — the entropy term
• The temperature dependence of free energy, then comes from
Chemical
the entropy term.
Thermodynamics
Spontaneity: Enthalpy & Entropy
G° = H  TS
Spontaneous @ all T
NonSpontaneous @ all T
Spontaneous @ low T
Spontaneous @ high T
Chemical
Thermodynamics
Spontaneity: Enthalpy & Entropy
G = H(
{Entropy Driven Reactions}
NH4NO3(s)
H2O
NH4
+
(aq)
+
NO3-(aq)
Enthalpy
2 H2O(g)
n = 2-3 = -1
Entropy
H = +
S = +
(-TS)
H = -
S = (+TS)
{Enthalpy Driven Reaction}
2 H2(g) + O2 (g)
TS)
{EntropySyst+Surr.FormationOfWater}
{Entropy & Enthalpy Driven Reaction}
H = Na2CO3(s) + HCl(aq)
NaCl(aq) + CO2 (g)
S
=+
Chemical
Thermodynamics
(-TS)
G° = HT(S
Problems
Product
Reactant
(-763)
(.3549)
– (-804)
– (.2219)
+41
(-T) +.1330
TiCl4(l)  TiCl4(g)
G° = H 
0 =TS
(131.3kJ)  T(.133kJ)
T = 987
Chemical
Thermodynamics
Standard Free Energy Changes
Analogous to standard enthalpies of
formation are standard free energies of
formation, G.
f
G = nGf (products)  mGf (reactants)
where n and m are the stoichiometric
coefficients.
Chemical
Thermodynamics
Standard Free Energy Changes
2 C6H6 (l) + 15 O2 (g)
12 CO2(g) + 6 H2O(g)
Calculate the standard free energy changes for the above
reaction @ 25 °C.
Grxn = nGf(prod)  mGf (react)
[12 CO2(g) + 6 H2O(g)] – [2 C6H6 (l) + 15 O2(g)]
[12(-394) + 6(-229)] – [2(125) + 15 (0)]
Grxn = -
6352 J/mol · K
Standard Molar
Gibbs Energy of
Formation (G°f)
CO2 (g)
-394
H2O (g)
-229
C6H6 (l)
125
Chemical
Thermodynamics
Fourth Law of
Thermodynamics: Emergence
Complex emergent systems spontaneously (-G) arise when energy flows through a
collection of many interacting particles, resulting in new patterns of complex
behaviors that are much more than the sum of the individual parts (+S)
The formation of these complex patterns in emergent systems is more efficient in the
dissipation of energy (-H), thus speeds up the increase of entropy in the universe.
G = H -TS
A precise definition of emergence and a useful mathematical formulation of this
phenomenon remains elusive.
C  f [n, i, E(t)]
Examples: cells forming living organisms; stars forming galaxies; neurons forming
conscious brain.
Chemical
Thermodynamics
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