Transcript Chemical Thermodynamics presentation 1

Chemistry and Thermodynamics
• Physics Helps us understands chemistry (and
biology etc.)
First Law of Thermodynamics
• 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.
• But there is more going on than just the
First Law
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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
Equilibrium state is
the most likely
Slide 11-33
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 = Sfinal  Sinitial
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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:
Suniv = Ssystem + Ssurroundings = 0
For irreversible processes:
Suniv = Ssystem + Ssurroundings > 0
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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.
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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)
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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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Entropy Changes
– Recall entropy increases
if the number of gas
molecules increases;
– In this reaction the
number of gas
molecules decrease so
this is an example of a
decrease in entropy.
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Entropy decreases when:
a.
b.
c.
d.
gases are formed from liquids.
liquids are formed from solids.
solids are formed from gases.
the number of gas molecules increases
during a chemical reaction.
Entropy decreases when:
a.
b.
c.
d.
gases are formed from liquids.
liquids are formed from solids.
solids are formed from gases.
the number of gas molecules increases
during a chemical reaction.
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.
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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
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Gibbs Free Energy
• Gibbs Free Energy is a quantity that helps tell
us whether a process is spontaneous
• G= H-TS
• G = H  TS
• When ΔG is negative Suniverse is positive,
• Therefore, when G is negative, a process is
spontaneous.
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Upon heating, limestone (CaCO3) decomposes to CaO
and CO2. Speculate on the sign of H and S for this
process.
CaCO3(s)
CaO(s)
+
CO2(g)
Upon heating, limestone (CaCO3) decomposes to CaO
and CO2. Speculate on the sign of H and S for this
process.
CaCO3(s)
CaO(s)
+
CO2(g)
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 (products)  mG a (rectants)
f
f
where n and m are the stoichiometric
coefficients.
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Sample Exercise 19.6 Calculating Free-Energy Change from ΔH°, T, ΔS°
Calculate the standard free energy change for the
formation of NO(g) from N2(g) and O2(g) at 298 K:
N2(g) + O2(g) → 2 NO(g)
given that ΔH° = 180.7 kJ and ΔS° = 24.7 J/K. Is the
reaction spontaneous under these circumstances?
Calculate the standard free energy change for the formation of
NO(g) from N2(g) and O2(g) at 298 K:
N2(g) + O2(g) → 2 NO(g)
given that ΔH° = 180.7 kJ and ΔS° = 24.7 J/K. Is the reaction
spontaneous under these circumstances?
Analyze: We are asked to calculate ΔG° for the
indicated reaction (given ΔH°, ΔS° and T) and to predict
whether the reaction is spontaneous under standard
conditions at 298 K.
Plan:
To calculate ΔG°, we use Equation, ΔG° = ΔH° – TΔS°.
To determine whether the reaction is spontaneous
under standard conditions, we look at the sign of ΔG°.
Solution
Because ΔG° is positive, the reaction is not spontaneous
under standard conditions at 298 K.
Note: We had to convert the units of the T ΔS term to kJ
sp that it could be added to the ΔH term whose units
are kJ
Free Energy and Temperature
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