Transcript Chapter 19 Chemical Thermodynamics
Topic 9 Chapter 18 Chemical Thermodynamics
Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
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.
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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
.
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Spontaneous Processes Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.
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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.
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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.
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Irreversible Processes • Irreversible processes cannot be undone by exactly reversing the change to the system.
• Spontaneous processes are irreversible .
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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, .
T
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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.
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Entropy • Like total energy,
E
, and enthalpy,
H
, entropy is a state function .
• Therefore,
S
=
S
final
S
initial Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
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:
S
=
q
rev
T
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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.
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Second Law of Thermodynamics In other words: For reversible processes:
S
univ =
S
system +
S
surroundings = 0 For irreversible processes:
S
univ =
S
system +
S
surroundings > 0 Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Second Law of Thermodynamics These last truths mean that as a result of all spontaneous processes the entropy of the universe increases .
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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.
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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.
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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.
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Entropy on the Molecular Scale • Each thermodynamic state has a specific number of microstates,
W
, associated with it.
• Entropy is where
k S
=
k
ln
W
is the Boltzmann constant, 1.38 10 23 J/K.
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Entropy on the Molecular Scale • The change in entropy for a process, then, is
S
=
k
ln
W
final
k
ln
W
initial
S
=
k
ln ln
W
final ln
W
initial • Entropy increases with the number of microstates in the system.
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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.
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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 © 2009, Prentice-Hall, Inc.
Solutions Generally, when a solid is dissolved in a solvent, entropy increases.
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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.
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Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0.
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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.
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Standard Entropies Larger and more complex molecules have greater entropies.
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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 © 2009, Prentice-Hall, Inc.
Entropy Changes in Surroundings • Heat that flows into or out of the system changes the entropy of the surroundings.
• For an isothermal process:
S
surr =
q
sys
T
• At constant pressure,
q
sys
H
for the system.
is simply Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Entropy Change in the Universe • The universe is composed of the system and the surroundings.
• Therefore,
S
universe =
S
system +
S
surroundings • For spontaneous processes
S
universe > 0 Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Entropy Change in the Universe • Since
S
surroundings and
q
system
q
system
T
=
H
system This becomes:
S
universe
=
S
system +
H
system
T
Multiplying both sides by
T
, we get
T
S
universe =
H
system
T
S
system Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Gibbs Free Energy •
T
S
universe energy,
G
.
is defined as the Gibbs free • When
S
negative.
universe is positive,
G
is • Therefore, when
G
is negative, a process is spontaneous.
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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.
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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
coefficients.
are the stoichiometric Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Free Energy Changes At temperatures other than 25 ° C,
G
° =
H
T
S
How does
G
change with temperature?
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Free Energy and Temperature • There are two parts to the free energy equation:
H
— the enthalpy –
T
S
— the entropy term term • The temperature dependence of free energy, then comes from the entropy term.
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Free Energy and Temperature Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Free Energy and Equilibrium Under any conditions, standard or nonstandard, the free energy change can be found this way:
G
=
G
+
RT
ln
Q
(Under standard conditions, all concentrations are 1
M
, so
Q
= 1 and ln
Q
= 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 © 2009, Prentice-Hall, Inc.
Free Energy and Equilibrium • At equilibrium,
Q
=
K
, and
G
= 0.
• The equation becomes 0 =
G
+
RT
ln
K
• Rearranging, this becomes
G
=
RT
ln
K
or,
K
= e
G
RT
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G 0
=
RT
ln
K
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.
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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.
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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.
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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.
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Average Bond Enthalpies NOTE: These are
average
bond enthalpies, not absolute bond enthalpies; the C-H bonds in methane, CH 4 , will be a bit different than the C-H bond in chloroform, CHCl 3 .
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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,
H rxn
= (bond enthalpies of bonds broken) (bond enthalpies of bonds formed) Chemical Thermodynamics © 2009, Prentice-Hall, Inc.
Enthalpies of Reaction CH 4 (
g
) + Cl 2 (
g
) CH 3 Cl (
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.
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Enthalpies of Reaction So,
H
H
= [
D
= Reactants - Products (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.