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Energy and
Equilibrium: The
Laws of
Thermodynamics
Chapter 7.7
Energy
• Substances tend react to achieve the lowest energy
state.
• Most chemical reactions are exothermic.
• Doesn’t work for things like ice melting.
• An ice cube must absorb heat to melt, but it melts
anyway. Why?
Introduction
• Thermodynamics examines the relationship between
heat and work.
• Entropy is a measure of how energy is spread out
among the atoms and molecules of a system.
• Spontaneity is the notion of whether or not a process
can take place unassisted.
• Free energy is a thermodynamic function that relates
enthalpy and entropy to spontaneity.
• Free energy can also be related to equilibrium
constants.
The Concept of Entropy
• The degree of randomness or disorder.
• S not s!!
• The First Law of Thermodynamics. The energy of the universe
is constant.
• The Second Law of Thermodynamics. The entropy of the
universe increases in any change.
– Drop a box of marbles.
– Watch your room for a week.
Absolute Entropies
The Concept of Entropy
• Entropy, S.
ΔU = ΔH = 0
– The greater the
number of
configurations of the
microscopic particles
among the energy
levels in a particular
system, the greater the
entropy of the system.
•
Third Law of Thermodynamics.
•
Standard molar entropy.
– The entropy of a pure perfect crystal at 0 K is zero.
– Found in Data Tables.
ΔS = [ npS (products) - nrS (reactants)]
ΔS > 0 spontaneous
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Spontaneous Process
• A process that occurs in a system left to itself.
– Once started, no external action is necessary to make the process
continue.
• A non-spontaneous process will not occur
without external action continuously applied.
• The reverse of any spontaneous process is always
nonspontaneous
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
H2O(s)  H2O(l)
Two Factors
• Exothermic reactions tend to be spontaneous.
• negative H.
• Reactions where the entropy of the products is
greater than reactants tend to be
spontaneous.
• Positive S.
• A change with positive S and negative H is
always spontaneous.
• A change with negative S and positive H is
never spontaneous.
Other Possibilities
• An endothermic reaction with an increase in
entropy like melting ice.
–Spontaneous at high temperature.
–Nonspontaneous at low temperature.
Other Possibilities
• Temperature affects entropy.
• Higher temperature, higher entropy.
• For an exothermic reaction with a decrease in
entropy (like rusting).
– Spontaneous at low temperature.
– Nonspontaneous at high temperature.
–Enthalpy driven.
Standard Free Energy Change, ΔG
• The standard free energy of formation, ΔGf .
– The free energy change for a reaction in which a
substance in its standard state is formed from its
elements in reference forms in their standard
states.
• The standard free energy of reaction, ΔG .
ΔG = [ np ΔGf (products) - nr ΔGf (reactants)]
–Entropy driven.
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Gibbs Free Energy
Criteria for Spontaneous Change
• The energy free to do work is the change in
Gibbs free energy.
 Gº = Hº - T Sº (T must be in Kelvin)
• All spontaneous reactions release free energy.
• So G <0 for a spontaneous reaction.
ΔGsys < 0 (negative), the process is
spontaneous.
ΔGsys = 0 (zero), the process is at
equilibrium.
ΔGsys > 0 (positive), the process is nonspontaneous.
Problems
G= H-T S
S
H
+ -
G
Spontaneous?
-
At all Temperatures
+ + ?
-
-
?
-
+ +
At high temperatures,
“entropy driven”
At low temperatures,
“enthalpy driven”
• Using the information on page 799 and pg 800
determine if the following changes are
spontaneous at 25ºC.
• 2H2S(g) + O2(g)
2H2O(l) + 2S(rhombic)
• At what temperature does it become
spontaneous?
Not at any temperature,
Reverse is spontaneous
2H2S(g) + O2(g)
2H2O(l) + 2S
• We find Hf° for each component
–H2S = -20.1 kJ
–H2O = -285.8 kJ
2H2S(g) + O2(g)
• We find S for each component
–H2S = 205.6 J/K
–H2O = 69.94 J/K
O2 = 0 kJ
S = 0 kJ
O2 = 205.0 J/K
S = 31.9 J/K
• Then Products - Reactants
• Then Products - Reactants
•
•
H =[2 (-285.8) + 0]
- [2 (-20.1) + 1(0)] = -531.4 kJ
2H2O(l) + 2S
S=[2 (69.94) + 2(31.9) ]
- [2 (205.6) + 205] = -412.5 J/K
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•
•
•
•
•
•
•
•
2H2S(g) + O2(g)
2H2O(l) + 2S
G= H-T S
G = -531.4 kJ - 298K (-412.5 J/K)
G = -531.4 kJ - -122925 J
G = -531.4 kJ - -123 kJ
G = -408.4 kJ
Spontaneous
Exergonic- it releases free energy.
At what temperature does it become
spontaneous?
There’s Another Way
• There are tables of standard free energies of
formation compounds.
 Gºf is the free energy change in making a
compound from its elements at 25º C and 1
atm.
• for an element Gºf = 0
• Look them up.
 Gº= Gºf(products) - Gºf(reactants)
• Check the last problems.
Spontaneous
• It becomes spontaneous when G = 0
• That’s where it changes from positive to
negative.
• Using 0 = H - T S and solving for T
• 0- H=-T S
• - H = -T
S
• T = H = -531.4 kJ = -531400 J
= 1290 K
S
-412.5 J/K -412.5 J/K
Criteria for Spontaneous Change
Every chemical reaction consists of both a
forward and a reverse reaction.
The direction of spontaneous change is the direction
in which the free energy decreases.
Summary
Summary (Continued)
• A spontaneous change is one that occurs by itself
without outside intervention.
• The third law of thermodynamics states that the entropy
of a pure, perfect crystal at 0 K can be taken to be zero.
• The direction of spontaneous change is that in which
total entropy increases.
• The free energy change, G, is equal to -T ( S), and it
applies just to the system itself, without regard for the
surroundings.
• The standard free energy change, Go, can be calculated
by substituting standard enthalpies and entropies of
reaction and a Kelvin temperature into the Gibbs
equation, or, by combining standard free energies of
formation.
• The condition of equilibrium is one for which
G=
0.
• The value of Go is by itself often sufficient to determine
how a reaction will proceed.
• Values of Gof, Hof, and So are generally tabulated for
25oC.
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