AP Chemistry
Unit 13 – Thermodynamics
Thermodynamics is the study of the relationship between forms of energy in chemical or physical
processes. In this unit we're going to focus on thermodynamic quantities of entropy, enthalpy, and
Gibb’s free energy. Entropy is a thermodynamic quantity that is a measure of the disorder in a system.
Gibbs free energy, which incorporates the enthalpy and entropy of a system, will be introduced. After
we've learned about Gibbs free energy, we'll see how it relates to the equilibrium constant both
qualitatively and quantitatively.
Objectives
13.1 Decide whether a process will be spontaneous or nonspontaneous.
13.2 Calculate thermodynamic quantities of entropy, S, enthalpy, H, and Gibbs free energy, G.
13.3 Apply thermodynamic calculations to descriptions of equilibrium.
Skills to Master:
a. Explain why enthalpy alone cannot be used to predict the direction of a reaction.
b. Describe the relationship between entropy and disorder.
c. Rank different states of matter in order of increasing entropy.
d. State the second law of thermodynamics.
e. Explain the role of entropy in reaction spontaneity.
f. Define the term "standard molar entropy."
g. Write an expression for the change of entropy of a reaction.
h. Determine the sign of DS for a phase transition.
i. Predict the sign of DS for a chemical reaction involving gases.
j. Calculate DS for a chemical reaction using tables of standard molar entropy values.
k. Determine the spontaneity of a reaction on the basis of the sign of its Gibbs free energy.
l. Define the term "standard free energy of formation."
m. Calculate the standard free energy of a reaction using the enthalpy, temperature, and entropy of
the reaction.
n. Predict the algebraic sign of DG° from the algebraic signs of DH° and DS°.
o. Explain the role of temperature in determining the sign of DG°.
p. Determine the temperature at which a spontaneous process becomes nonspontaneous, or vice
versa.
q. Calculate DG from DG°, T, and the reaction quotient Q.
r. Calculate DG° when K is known, or vice versa.
s. Use a free-energy diagram to predict the direction of a reaction at any ratio of products to
reactants.
Chapter 19 Problem Set: 1, 7, 21, 32, 43, 47, 53, 73, 75, 91
1
Unit 14 – Thermodynamics: The study of energy relationships (Chapter 19)
Podcast 14.1 Spontaneity
Spontaneous Reactions
• Naturally-occurring reactions (happens on its own) that favor the formation of ________
• Do not favor product formation nor provide _______________ product yield.
*Does NOT refer to Rate! Reactions can have parts of both.
*_______________ and _______________ have great impacts on spontaneous reactions
Free Energy, ∆G
• Energy available to do __________
• No reaction is 100% efficient
• Energy can only be obtained if the reaction occurs (spontaneous)
• Measure of the _________ of a system
• Law of Disorder: Processes will spontaneously move in the direction of __________
__________ or randomness
Entropy and Phase of Matter
• Entropy of __________ is greater than the entropy of a __________ or __________
• _________ entropy is greater than a _________ entropy
• Entropy for a reaction will _____________ when a solid changes into a liquid or a gas.
• Entropy also increases when a substance is ____________ ______ ________ .
• Entropy increases when the total # of __________ molecules is larger than the # of
_____________ molecules
• Temp Increase = ____________ Increase
The Meaning of the Signs of Thermodynamic Properties
2
Property
Positive (+)
Negative (-)
∆H
∆S
∆G
3
Heat Change, ∆H
Entropy, ∆S
Reaction Type
Example 1: H2O (s) ” H2O (l)
A. When is this process spontaneous?
B. When is this process non-spontaneous?
• It all depends on the ________________!
4
5
Podcast 13.2: Calculations in Thermodynamics
____________ measurement of Disorder of a system
• Symbol = ∆S
• Units = J/K or J/(K*mol)
• Standard Entropy = ∆So
at 25 oC and 101.3kPa
• A perfect crystal at 0 K would have ∆S=0
• ∆So = ____________ - _____________
Example 2:(use Appendix p. 558)
H2 (g) + Cl2 (g) – 2HCl (g)
Example 3: Calculate the ∆S for both the forward and reverse reactions for the synthesis of water.
Free Energy Calculations
• Gibbs Free Energy Change, _______
• Maximum amount of energy that can be coupled to another process to do useful work
• Relates to Entropy and Enthalpy changes
• Gibbs Free Energy Change, _______
• Maximum amount of energy that can be coupled to another process to do useful work
• Relates to Entropy and Enthalpy changes
∆G = _______ - _________
• If ∆G is negative, the process is______________
• If ∆G is positive, the process is ______________
• If ∆G is 0, the process is ____________________
Example 4: Calculate the Free Energy for the following reaction to determine if it is spontaneous at
25oC
C(s) + O2(g) ” CO2(g)
6
7
Calculations of Thermodynamics
1. Use the bond energies in the table provided to calculate the H for the reactions below. You
WILL need to draw the Lewis Structures to answer these correctly.
a. H2 (g) + Cl2 (g)  2HCl (g)
b. CH4 (g) + 2O2 (g)  2H2O (g) + CO2 (g)
c. 2N (g)  N2 (g)
2. Calculate the entropy change, So for the reaction shown below.
2H2 (g) + O2 (g)  2H2O (l)
Substance So (J mol-1 K-1)
O2 (g)
205.0
H2O (l)
69.9
H2 (g)
130.6
8
3. Calculate the entropy change, So, for the reaction shown below.
Al2O3 (s) + 3H2 (g)  2Al (s) + 3H2O(g)
Substance So (J mol-1 K-1)
Al2O3 (g)
51.0
H2O (g)
189.0
H2 (g)
130.6
Al (s)
28.0
4. Calculate the standard free energy change, ΔG0, for the complete combustion of methanol,
CH3OH (g), at 25 0C. Also, is this reaction spontaneous?
Substance Go (J mol-1 K-1)
O2 (g)
0
H2O (g)
-229.4
CO2 (g)
-394.4
CH3OH (g) -163.2
5. Calculate the standard free energy change, ΔG0, for the reaction shown below at 25 oC. Is this
reaction spontaneous?
2Al (s) + 3H2O(g)  Al2O3 (s) + 3H2 (g)
9
Podcast 13.3 Free Energy and Equilibrium
∆Go = ∆Ho – T∆ So can be used algebraically to solve for any unknown quantity ∆Ho, T, or ∆ So
Free energy at
Reaction Quotient
Non-standard
(compare to Keq )
Conditions
Free Energy and Equilibrium
At standard conditions, Q = Keq
∆Go = - RT lnKeq
Rearrange equation to solve for K
∆Go = negative, Keq > 1
∆Go = 0, Keq = 1
∆Go = positive, Keq < 1
Calculating Enthalpy, ∆H
1. Calorimetry: q = mC∆T
2. Hess’s Law : Add reactions
3. Hess’s Law : Σproducts – Σ reactants
4. Bond Energies : Bonds Broken – Bonds Formed
Calculating Free Energy, G
5. ∆Go = ∆Ho – T∆ So
6. ∆Go = -nF€o
7. ∆Go = - RT lnKeq
Example 5: Calculate ∆H, ∆S, and ∆Go for the following reaction at 298K: Use the appendix in your
book to find the individual values.
2SO2(g) + O2(g) ” 2SO3(g)
Example 6: Consider the ammonia synthesis reaction: N2(g) + 3H2(g) – 2NH3(g)
where ∆ G = -33.3 kJ/mol of N2 consumed at 25oC. Calculate the value for the equilibrium constant.
∆Go = - RT lnKeq
Example 7: Calculate the value of ∆Go for the following reaction at 389K where the [NH3] = 2.0 M,
[H2] = 1.25 M, and [N2] = 3.01 M
N2(g) + 3H2(g) – 2NH3(g)
10
THERMODYNAMICS REVIEW
0
2 Fe (s) + 32O2 (g) → Fe2O3 (s) ΔH f = -824 kJ mol
-1
1. Iron reacts with oxygen to produced iron(III) oxide, as represented by the equation above. A 75.0 gram
sample of Fe(s) is mixed with 11.5 L of O2 (g) at 2.66 atm and 298 K.
(a) Calculate the number of moles of each of the following before the reaction begins.
(i) Fe(s)
(ii) O2 (g)
(b) Identify the limiting reactant when the mixture is heated to produce Fe O
2
. Support your answer
3 (s)
with calculations.
(c) Calculate the number of moles of Fe2O3 (s) produced when the reaction proceeds to completion.
0
-1
(d) The standard free energy of formation, ΔG f , of Fe2O3 (s) is -740 kJ mol at 298 K.
0
(i) Calculate the standard entropy of formation, ΔS f , of Fe2O3 (s) at 298 K, Include units with your
answer.
(ii) Which is more responsible for the spontaneity of this reaction at 298 K, the standard enthalpy of
0
0
formation, ΔH f , or the standard entropy formation, ΔS f ? Justify your answer.
0
-1
The reaction represented below also produces iron(III) oxide. The value of ΔH for the reaction is -280 kJ mol of
Fe2O3 (s) formed.
2 FeO (s) + 12O2 (g) → Fe2O3 (s)
0
(e) Calculate the standard enthalpy of formation, ΔH f , for FeO (s). (Another Hint: Remember, questions
like this give information that may be vital in earlier parts… )
11
2003 FRQ
7. Answer the following questions that relate to the chemistry of nitrogen.
a. Two nitrogen atoms combine to form a nitrogen molecule, as represented by the following
equation
2 N(g)  N2 (g)
Using the table of average bond energies below, determine the enthalpy change, ΔH, for
the reaction.
Average Bond Energy (kJ mol-1)
160
420
950
Bond
N–N
N=N
N≡N
b. The reaction between nitrogen and hydrogen to form ammonia is represented below.
N2(g) + 3 H2(g)  2 NH3(g)
ΔHo = - 92.2 kJ
Predict the sign of the standard entropy change, ΔSo, for the reaction. Justify your answer.
8.
c.
The value of ΔGo for the reaction represented in part (b) is negative at low temperatures
but positive at high temperatures. Explain
d.
When N2(g) and H2(g) are placed in a sealed container at a low temperature, no
measurable amount of NH3(g) is produced. Explain.
Using the information in the table, answer the following questions about organic compounds.
Compound Name
Compound Formula
ΔHvapo (kJ mol-1)
Propane
Propanone
1-propanol
CH3CH2CH3
CH3COCH3
CH3CH2CH2OH
19.0
32.0
47.3
a.
b.
For propanone,
i.
draw the complete structural formula (showing all atoms and bonds)
ii.
predict the approximate carbon-to-carbon-to-carbon bond angle.
For each pair of compounds below, explain why they do not have the same value for their
standard heat of vaporization, ΔHvapo. (You must include specific information about both
compounds in each pair.)
i.
propane and propanone
12
ii.
propanone and 1-propanol
c. Draw the complete structural formula for an isomer of the molecule you drew in part a. i.
d.
Given the structural formula for propyne below,
i.
indicate the hybridization of the carbon atom indicated by the arrow in the
structure above;
ii.
indicate the total number of sigma (σ) bonds and the total number of pi (π) bonds
in the molecule.
13