Chapter 17
THE DIRECTION OF CHEMICAL
CHANGE
Direction of spontaneous
change
 If a process is found to be spontaneous, the
reverse process is non spontaneous.
 Both spontaneous and non spontaneous
processes are possible but only spontaneous
process will occur naturally. Non spontaneous
processes require the system to be acted
upon in some way.
 The distinction between spontaneous and
non-spontaneous processes is formalized in
the Second Law of Thermodynamics.
Entropy and disorder
 Second Law of Thermodynamics: In any cyclic
process the entropy will either increase or
remain the same.
 A more precise way to characterize entropy is
to say that it is a measure of the "multiplicity"
associated with the state of the objects. If a
given state can be accomplished in many
more ways, then it is more probable than one
which can be accomplished in only a few
ways.
For a glass of water the number of molecules is
astronomical. The jumble of ice chips may look more
disordered in comparison to the glass of water which looks
uniform and homogeneous. But the ice chips place limits on
the number of ways the molecules can be arranged. The
water molecules in the glass of water can be arranged in
many more ways; they have greater "multiplicity" and
therefore greater entropy.
Class Practice
 Calculate the change in entropy of a large
tank of water when a total of 100J of energy
is transferred to it reversibly as heat at 20⁰C.
Entropy of vaporization
/fusion
 The entropy of vaporization of a substance,
∆Svap , is the difference in molar entropies of
the vapor and liquid phases of a substance at
a stated temperature.
 The entropy of fusion, ∆S fus is the difference
in the molar entropies of the liquid and solid
phases of a substance at a stated
temperature.
 The second law of thermodynamics states
that the entropy of an isolated system tends
to increase.
Class practice
 Calculate the change in molar entropy when
water vaporizes at its boiling point.
 The third law of thermodynamics states that
the entropy of a perfect crystal approaches 0
as the absolute temperature approaches 0.
Reaction entropy
 The standard reaction entropy, ∆S ⁰f is the
difference between the standard molar
entropies of the products and the reactants.
 ∆S ⁰f =∑nS⁰m (products) - ∑nS⁰m (reactants)
 n is the various stochiometric coefficients.
 The standard reaction entropy is usually
positive(an increase in entropy) if there is a
net production of gas in a reaction; it is
usually negative( a decrease) if there is a net
consumption of gas.
Class Practice
 Calculate the standard reaction entropy for
N₂(g) +3H₂(g)  2NH₃(g)
Entropy of the surrounding
 Whenever we talk of entropy we must take into
account the sum of the changes in the system,
∆S, and the surrounding ∆S surr
 Total entropy change= entropy change of system
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+ entropy change of surrounding
∆Stot = ∆S + ∆S surr
Entropy change of surrounding=heat transferred
to surroundings/temperature of surroundings
=enthalpy change of system/temp. of the
surrounding
∆S surr = - ∆H/T
 A chemical reaction is spontaneous if it is
accompanied by an increase in the total
entropy of the system and the surroundings.
Spontaneous exothermic reactions are
common because they release heat that
increases the entropy of the surroundings.
 Endothermic reactions are spontaneous only
if the reaction mixture undergoes a large
increase in entropy.
Class Practice
 Is the dissolution of ammonium nitrate to
form a dilute aqueous solution spontaneous
at 25⁰C ?
Free Energy
 The total entropy change ∆Stot is the sum of
the entropy changes in the system ,∆S and its
surroundings , ∆S surr
 ∆Stot = ∆S + ∆S surr
∆S surr = - ∆H/T
 ∆Stot = ∆S - ∆H/T
 ∆G = -T∆Stot = ∆H-T ∆S
 ∆G is Gibbs free energy
 The free energy change is a measure of the
change in the total entropy of a system and
its surroundings at constant pressure;
spontaneous processes are accompanied by a
decrease in free energy.
Free energy and equilibrium
 If ∆G<0 process is spontaneous
 If ∆G>0 the reverse of the process is
spontaneous
 If ∆G=0 process is at equilibrium
 ∆G=0 at 0⁰C signifies the fact that water and
ice are in equilibrium at that temperature.
Predicting the boiling point
 ∆G = ∆H-T ∆S
 We want the temperature at which ∆G=0
 ∆Hvap-Tb ∆Svap =0
 Tb =∆Hvap / ∆Svap
 The melting point follows the same formula
Standard free energy
 A thermodynamically stable compound is a
compound with a negative free energy of
formation. Such a compound has no
tendency to decompose into its elements.
 A thermodynamically unstable compound is a
compound with a positive standard free
energy of formation. Such a compound has a
tendency to decompose into its elements.
 The standard free energy of formation of a
substance is the standard free energy of
reaction per mole of compound when it is
prepared from its elements in their most
stable forms. The sign of ∆G⁰f tells us
whether a compound is stable or unstable
with respect to its elements.
Using free energies of
formation
 ∆G⁰f =∑n ∆G⁰f (products) - ∑n ∆G⁰f (reactants)
 Standard free energies of formation are
combined to calculate the standard free
energy of the reaction.
Equilibrium constants
 The reaction free energy is related to the
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

composition of the reaction mixture by
∆Gr = ∆G⁰r + RTlnQ
At equilibrium ∆Gr = 0 and Q=K
0 = ∆G⁰r + RTlnK
Rearranging the equation
∆G⁰r = -RTlnK
K<1 when ∆G⁰r >0
K>1 when ∆G⁰r < 0
 The equilibrium constant is related to the
standard free energy of reaction by
 ∆G⁰r = -RTlnK
 K is expressed in terms of numerical values of
the partial pressures of gases Kp and as the
numerical values of the molar concentrations
of solutes Kc.
Class Practice
 Calculate Kp at 25⁰C for the equilibrium
N2O4(g)↔2NO2(g)
Effect of temperature
 The free energy increases with temperature
for reactions with a negative ∆Sr⁰ and
decreases with temperature for reactions with
a positive ∆Sr⁰ .
Home work
 Page 786
 17.46,17.52, 17.71, 17.74