Topic 9
Chapter 18
Chemical
Thermodynamics
Chemical
Thermodynamics
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First Law of Thermodynamics
• You will recall from Chapter 6 that
energy cannot be created nor
destroyed.
• Therefore, the total energy of the
universe is a constant.
• Energy can, however, be converted
from one form to another or transferred
from a system to the surroundings or
vice versa.
Chemical
Thermodynamics
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Spontaneous Processes
• Spontaneous processes
are those that can
proceed without any
outside intervention.
• The gas in vessel B will
spontaneously effuse into
vessel A, but once the
gas is in both vessels, it
will not spontaneously
return to vessel B.
Chemical
Thermodynamics
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Spontaneous Processes
Processes that are
spontaneous in one
direction are
nonspontaneous in
the reverse
direction.
Chemical
Thermodynamics
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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.
Chemical
Thermodynamics
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Reversible Processes
In a reversible
process the system
changes in such a
way that the system
and surroundings
can be put back in
their original states
by exactly reversing
the process.
Chemical
Thermodynamics
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Irreversible Processes
• Irreversible processes cannot be undone by
exactly reversing the change to the system.
• Spontaneous processes are irreversible.
Chemical
Thermodynamics
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Entropy
• Entropy (S) is a term coined by Rudolph
Clausius in the 19th century.
• Clausius was convinced of the
significance of the ratio of heat
delivered and the temperature at which
it is delivered, q .
T
Chemical
Thermodynamics
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Entropy
• Entropy can be thought of as a measure
of the randomness of a system.
• It is related to the various modes of
motion in molecules.
Chemical
Thermodynamics
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Entropy
• Like total energy, E, and enthalpy, H,
entropy is a state function.
• Therefore,
S = Sfinal  Sinitial
Chemical
Thermodynamics
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Entropy
For a process occurring at constant
temperature (an isothermal process), the
change in entropy is equal to the heat that
would be transferred if the process were
reversible divided by the temperature:
qrev
S =
T
Chemical
Thermodynamics
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Second Law of Thermodynamics
The second law of thermodynamics
states that the entropy of the universe
increases for spontaneous processes,
and the entropy of the universe does
not change for reversible processes.
Chemical
Thermodynamics
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Second Law of Thermodynamics
In other words:
For reversible processes:
Suniv = Ssystem + Ssurroundings = 0
For irreversible processes:
Suniv = Ssystem + Ssurroundings > 0
Chemical
Thermodynamics
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Second Law of Thermodynamics
These last truths mean that as a result
of all spontaneous processes the
entropy of the universe increases.
Chemical
Thermodynamics
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Entropy on the Molecular Scale
• Ludwig Boltzmann described the concept of
entropy on the molecular level.
• Temperature is a measure of the average
kinetic energy of the molecules in a sample.
Chemical
Thermodynamics
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Entropy on the Molecular Scale
• Molecules exhibit several types of motion:
– Translational: Movement of the entire 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.
Chemical
Thermodynamics
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Entropy on the Molecular Scale
• 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 of the
thermodynamic system.
Chemical
Thermodynamics
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Entropy on the Molecular Scale
• Each thermodynamic state has a specific number of
microstates, W, associated with it.
• Entropy is
S = k lnW
where k is the Boltzmann constant, 1.38  1023 J/K.
Chemical
Thermodynamics
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Entropy on the Molecular Scale
• The change in entropy for a process,
then, is
S = k lnWfinal  k lnWinitial
lnWfinal
S = k ln
lnWinitial
• Entropy increases with the number of
Chemical
microstates in the system.
Thermodynamics
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Entropy on the Molecular Scale
• The number of microstates and,
therefore, the entropy tends to increase
with increases in
– Temperature.
– Volume.
– The number of independently moving
molecules.
Chemical
Thermodynamics
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Entropy and Physical States
• Entropy increases with
the freedom of motion
of molecules.
• Therefore,
S(g) > S(l) > S(s)
Chemical
Thermodynamics
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Solutions
Generally, when
a solid is
dissolved in a
solvent, entropy
increases.
Chemical
Thermodynamics
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Entropy Changes
• In general, entropy
increases when
– Gases are formed from
liquids and solids;
– Liquids or solutions are
formed from solids;
– The number of gas
molecules increases;
– The number of moles
increases.
Chemical
Thermodynamics
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Third Law of Thermodynamics
The entropy of a pure crystalline
substance at absolute zero is 0.
Chemical
Thermodynamics
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Standard Entropies
• These are molar entropy
values of substances in
their standard states.
• Standard entropies tend
to increase with
increasing molar mass.
Chemical
Thermodynamics
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Standard Entropies
Larger and more complex molecules have
greater entropies.
Chemical
Thermodynamics
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Entropy Changes
Entropy changes for a reaction can be
estimated in a manner analogous to that by
which H is estimated:
S = nS(products) — mS(reactants)
where n and m are the coefficients in the
balanced chemical equation.
Chemical
Thermodynamics
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Entropy Changes in Surroundings
• Heat that flows into or out of the
system changes the entropy of the
surroundings.
• For an isothermal process:
Ssurr =
qsys
T
• At constant pressure, qsys is simply
H for the system.
Chemical
Thermodynamics
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Entropy Change in the Universe
• The universe is composed of the system
and the surroundings.
• Therefore,
Suniverse = Ssystem + Ssurroundings
• For spontaneous processes
Suniverse > 0
Chemical
Thermodynamics
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Entropy Change in the Universe
• Since Ssurroundings =
qsystem
and qsystem = Hsystem
T
This becomes:
Hsystem
Suniverse = Ssystem +
T
Multiplying both sides by T, we get
TSuniverse = Hsystem  TSsystem
Chemical
Thermodynamics
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Gibbs Free Energy
• TSuniverse is defined as the Gibbs free
energy, G.
• When Suniverse is positive, G is
negative.
• Therefore, when G is negative, a
process is spontaneous.
Chemical
Thermodynamics
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Gibbs Free Energy
1. If G is negative, the
forward reaction is
spontaneous.
2. If G is 0, the system
is at equilibrium.
3. If G is positive, the
reaction is
spontaneous in the
reverse direction.
Chemical
Thermodynamics
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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
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Free Energy Changes
At temperatures other than 25°C,
G° = H  TS
How does G change with temperature?
Chemical
Thermodynamics
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Free Energy and 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 the entropy
term.
Chemical
Thermodynamics
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Free Energy and Temperature
Chemical
Thermodynamics
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Free Energy and Equilibrium
Under any conditions, standard or
nonstandard, the free energy change
can be found this way:
G = G + RT lnQ
(Under standard conditions, all concentrations are 1 M,
so Q = 1 and lnQ = 0; the last term drops out.)
R is the gas constant (8.314 J/K•mol)
T is the absolute temperature (K)
Q is the reaction quotient (Products / Reactants)
Chemical
Thermodynamics
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Free Energy and Equilibrium
• At equilibrium, Q = K, and G = 0.
• The equation becomes
0 = G + RT lnK
• Rearranging, this becomes
G = RT lnK
or,
-G
K = e RT
Chemical
Thermodynamics
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G0 =  RT lnK
Chemical
Thermodynamics
18.6
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Energetics of Ionic Bonding
By accounting for all
three energies
(ionization energy,
electron affinity, and
lattice energy), we
can get a good idea
of the energetics
involved in such a
process.
Chemical
Thermodynamics
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Covalent Bond Strength
• Most simply, the strength of a bond is
measured by determining how much energy
is required to break the bond.
• This is the bond enthalpy.
• The bond enthalpy for a Cl-Cl bond, D(Cl-Cl),
is measured to be 242 kJ/mol.
Chemical
Thermodynamics
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Bond Enthalpy and Bond Length
• We can also measure an average bond
length for different bond types.
• As the number of bonds between two atoms
increases, the bond length decreases.
Chemical
Thermodynamics
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Average Bond Enthalpies
• This table lists the
average bond
enthalpies for many
different types of
bonds.
• Average bond
enthalpies are
positive, because
bond breaking is an
endothermic process.
Chemical
Thermodynamics
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Average Bond Enthalpies
NOTE: These are
average bond
enthalpies, not
absolute bond
enthalpies; the C-H
bonds in methane,
CH4, will be a bit
different than the C-H
bond in chloroform,
CHCl3.
Chemical
Thermodynamics
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Enthalpies of Reaction
• Yet another way to
estimate H for a
reaction is to compare
the bond enthalpies of
bonds broken to the
bond enthalpies of the
new bonds formed.
• In other words,
Hrxn = (bond enthalpies of bonds broken) (bond enthalpies of bonds formed)
Chemical
Thermodynamics
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Enthalpies of Reaction
CH4 (g) + Cl2 (g) 
CH3Cl (g) + HCl (g)
In this example, one C-H
bond and one Cl-Cl
bond are broken; one
C-Cl and one H-Cl
bond are formed.
Chemical
Thermodynamics
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Enthalpies of Reaction
So, H = Reactants - Products
H = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)]
= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]
= (655 kJ) - (759 kJ)
= -104 kJ
Chemical
Thermodynamics
© 2009, Prentice-Hall, Inc.