Thermodynamics pt 1:
Introduction to Spontaneity, Entropy,
and Gibbs Free Energy
SUGGESTED HW: Ch 23: 7, 11, 13, 17, 21
Intro to Thermodynamics
• Some things happen without influence, some things don’t.
• For example, decay just happens, without input. But creation requires
work. Water flows downhill, but you need a pump to force water uphill.
• Iron exposed to air will rust (Fe2O3). But rusted iron will not re-convert to
Fe(s) and O2. These are examples of irreversible processes.
X
Spontaneity
• What determines the direction of a process?
• The first law of thermodynamics tells us that
∆U = q + w
• This means that if a reaction occurs, the total energy of the universe is
unchanged.
• In this lecture, we address the word “if”.
• Why do some reactions occur, whereas others don’t?
• Let’s first begin by determining the criteria for a spontaneous process
Spontaneity
• A spontaneous process is any process that occurs without external
influence. ALL SPONTANEOUS PROCESSES ARE IRREVERSIBLE
• Work must be done by the surroundings to return the system to
the original state, but this leaves the surroundings permanently
changed.
• Thermodynamics allows us to determine if a process will occur, and in
which direction. Kinetics tells us how fast the reaction will go.
Hot
Cold
Spontaneity
• Spontaneous changes need not be fast.
• Ex. Diamonds spontaneously convert to graphite, but this process takes
centuries.
• What dictates the tendency of a process to spontaneously occur?
• When this question was first addressed in the 1860’s, it was thought that
the only criteria for spontaneity was that a reaction be exothermic.
Spontaneity
• Since exothermic processes are “energetically
downhill” processes, it was a logical assertion.
• However, this was quickly proved to be
incorrect.
• The dissolution of NaCl(s) in water is spontaneous, but endothermic
𝑁𝑎𝐶𝑙 𝑠 →
𝑁𝑎+
𝑎𝑞 +
𝐶𝑙 −
𝑎𝑞
∆𝐻𝑜
𝑘𝐽
= +3.9
𝑚𝑜𝑙
• Some spontaneous processes are temperature dependent. For example,
ice spontaneously melts at any temperature greater than 0oC.
𝐻2 𝑂 𝑠 → 𝐻2 𝑂 𝐿
∆𝐻𝑓𝑢𝑠 = +6.01
𝑘𝐽
𝑚𝑜𝑙
Entropy
• What is the common pattern with all spontaneous change?
• Spontaneous changes lead to increases in disorder.
• Expansion of gases creates a randomized, less ordered system
• Liquid water is much less ordered than ice. Ice atoms are held in
place, liquid atoms tumble around.
• Dissolving a salt in water yields ions that are free to move about
randomly
• The cooling of a hot block in air results in energy transfer to
surrounding air molecules, which increases their kinetic energy and
leads to more random motion and collisions
• This disorder is called ENTROPY (S)
Entropy
• The more disordered a system, the larger its entropy.
• Entropy, denoted S, is a state function, so it depends only on the initial
and final states, not the path taken.
• An increase in disorder represents a positive change in entropy (ΔS > 0),
while increases in order are negative (ΔS < 0)
• Suppose a system undergoes a process in which it changes from an initial
state (1) to a final state (2). The heat transferred during this process DOES
depend on the path
• To relate ΔS to heat, we consider a reversible path between the states.
Reversible Processes
• Imagine we have a reversible process. In a reversible process, the
direction of a process can be reversed by an infinitesimally small change in
one variable (Attaining a reversible condition is theoretical)
• Ex. Water at exactly 0oC. The tiniest change in pressure at constant
temperature (isothermal) will cause the water to move either toward
the solid or liquid phase.
• Because this pressure change is so small, no work is done, and thus,
none is required to reverse the pressure change.
• Whenever a chemical system is in equilibrium, we can go reversibly
between states without input of work or energy.
H2O(s)
H2O(L)
0oC
Example: Entropy Change for “Reversible” Processes
• For an isothermal process, like a phase change:
∆𝑆𝑠𝑦𝑠 =
𝑞𝑟𝑒𝑣
𝑇
• The subscript “rev” indicates that the transfer of heat is reversible,
so the system is in equilibrium.
• The heat term in the numerator accounts for the proportionality
between thermal transfer and disorder.
• The temperature term in the denominator accounts for the
disorder that already exists in the system.
Second Law of Thermodynamics
• THE ENTROPY OF THE UNIVERSE IS CONTINUALLY INCREASING.
∆𝑆𝑢𝑛𝑖𝑣 = ∆𝑆𝑠𝑦𝑠 + ∆𝑆𝑠𝑢𝑟𝑟
𝑟𝑒𝑣𝑒𝑟𝑠𝑖𝑏𝑙𝑒 𝑝𝑟𝑜𝑐𝑒𝑠𝑠: ∆𝑆𝑢𝑛𝑖𝑣 = ∆𝑆𝑠𝑦𝑠 + ∆𝑆𝑠𝑢𝑟𝑟 = 0
𝑖𝑟𝑟𝑒𝑣𝑒𝑟𝑠𝑖𝑏𝑙𝑒 𝑝𝑟𝑜𝑐𝑒𝑠𝑠: ∆𝑆𝑢𝑛𝑖𝑣 = ∆𝑆𝑠𝑦𝑠 + ∆𝑆𝑠𝑢𝑟𝑟 > 0
• For any irreversible process in which the system becomes more ordered,
the increase in disorder of the surroundings must be greater in
magnitude, and visa versa. The universe can NEVER become more
ordered.
• Considering the example of rust:
4𝐹𝑒 𝑠 + 3𝑂2 𝑔 → 2𝐹𝑒2 𝑂3 𝑠
∆𝐻𝑟𝑥𝑛
𝑘𝐽
= −825.50
𝑚𝑜𝑙
• ΔSsys is NEGATIVE. Why?
• Most combination reactions have negative entropy because you are
reducing the number of free species. Here, we have taken 7 total
moles of reactant and formed 2 moles of product
• Gases have much higher entropies than solids. Here, we have
consumed a gas to form a solid.
• ΔSsurr is POSITIVE because the reaction is highly exothermic. The
thermal energy gained by the surrounding atmosphere causes a high
degree of disorder in the surrounding gas molecules.
• The disorder to the surroundings caused by this process MUST be
greater than the order obtained by the system.
∆𝑆𝑠𝑦𝑠 + ∆𝑆𝑠𝑢𝑟𝑟 > 0
Molecular Interpretation of Entropy
• When we have a process than reduces the total number of free species, or
changes phase from gas to liquid/solid or liquid to solid, we limit the
motion of the molecules (i.e. the number of ways they can release energy)
• There are three types of motion: translational, vibrational, rotational.
The number of ways molecules can move are its degrees of freedom
𝑉
molecule moves from
one place to another
Translational (full movement)
Degrees of Freedom
Free motion
Vibrational
motion, restricted
rotational &
translational
motion
Vibrational motion
only
• Gases, being the least ordered, have the most ways of dissipating
thermal energy. Hence, they have the highest entropy.
Determine the sign of ΔSsys
• A(g) + 2B(g) ---> AB2(s)
negative
• H2O(s) ---> H2O (L)
positive
• NaCl(s) ---> Na+(aq) + Cl-(aq)
positive
• FeCl2(s) + H2(g) ---> Fe(s) + 2HCl(g)
positive
• A(g) + 2B(g) ---> C(g)
negative
H2O (L)
Third Law of Thermodynamics
• All molecular motion stops at 0oK (absolute zero). Therefore, S=0, and
the molecules arrange themselves in perfect order.
• The plot below shows a heating curve of entropy. The sharp increases at
phase boundaries is due to the added degrees of freedom
Calculations of Entropy Changes of Reactions
• Standard molar entropies, So (J/mol K) are
shown to the right .
1.
2.
3.
4.
Unlike enthalpies of formation, entropies
are NOT zero for elemental forms of
substances
Gases > Liquids > Solids
For the same phase, entropy increases
with molar mass
For the same phase and same molar mass,
entropy increases with the number of
atoms in the molecule.
𝑜 =
∆𝑆𝑟𝑥𝑛
𝑛𝑆 𝑜 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 −
𝑛𝑆 𝑜 (𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠)
Examples
• Which would you expect to have the higher molar entropy?
• H2O(L) or H2O(g)
• CO2(g) or H2O(g)
• CO(g) or CO2(g)
• Zn(s) or Li(s)
• NaClO4(s) or He(g)
Example
• Calculate the standard entropy change:
𝑁2 𝑔 + 3𝐻2 (𝑔) → 2𝑁𝐻3 (𝑔)
• Make sure equation is balanced. Use stoichiometric coefficients and
values from the table.
𝐽
𝐽
− (1)191.5 + 3 130.6
= −198.3 J/K
𝑚𝑜𝑙 𝐾
𝑚𝑜𝑙 𝐾
products
reactants
∆𝑆 𝑜 = 2 192.5
So What is the Criteria for Spontaneity?
• We have seen that spontaneous processes increase the entropy of the
universe
• ΔHsys does not have to be negative, and ΔSsys does not have to be positive.
• This brings us back to the initial question: What is the criteria of a
spontaneous process?
• Let’s use ΔH and ΔS concurrently to derive an expression
Math Time: Derivation of Gibbs Free Energy
∆𝑺𝒖𝒏𝒊𝒗 = ∆𝑺𝒔𝒚𝒔 + ∆𝑺𝒔𝒖𝒓𝒓
• If the surroundings include “everything else”, then we can assert that for
any process occurring in the system, the surroundings are large enough
that their temperature and pressure are constant.
∆𝑆𝑢𝑛𝑖𝑣
𝑞𝑠𝑦𝑠
= ∆𝑆𝑠𝑦𝑠 −
𝑇
∆𝑆𝑢𝑛𝑖𝑣 = ∆𝑆𝑠𝑦𝑠 −
Gibbs Free Energy
∆𝐻𝑠𝑦𝑠
𝑇
𝑇∆𝑆𝑢𝑛𝑖𝑣 = 𝑇∆𝑆𝑠𝑦𝑠 − ∆𝐻𝑠𝑦𝑠
−𝑇∆𝑆𝑢𝑛𝑖𝑣 = ∆𝐻𝑠𝑦𝑠 − 𝑇∆𝑆𝑠𝑦𝑠
∆𝐺 = ∆𝐻 − 𝑇∆𝑆
Sign of Gibbs Free Energy Dictates Direction of
Reaction
• If ΔG is negative, the reaction is spontaneous in the forward direction
• If ΔG is zero, the reaction is at equilibrium
• If ΔG is positive, the reaction is spontaneous in the reverse direction
What is Gibbs Free Energy?
• As you would imagine, it is very difficult to directly calculate ΔSuniv.
• However, the Gibbs Free Energy (-TΔSuniv) allows us to relate it to ΔH and
ΔS of the system. Hence, by following the 2nd law of thermodynamics, ΔG
tells us about the spontaneity of a process
• Physically speaking, ΔG is the maximum useful work that can be done by a
system on the surroundings at temperature T.
• In other words, all of the internal energy U of a system not accounted for
by ΔG will be lost as heat.
• When ΔG is positive, this value represents the minimum work that must be
done to the system to force the reaction to proceed.
Now We See That Spontaneity Depends on
Enthalpy AND Entropy
∆𝐺 = ∆𝐻 − 𝑇∆𝑆
Dictates if a process is
energetically favored
Dictates if a process
is entropically
favored