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Ch. 17 Kimball Spring 2019
Discussion Worksheet - Chapter 19
Chemical Thermodynamics
Dr Cabirac Lecture
Match.
G 2nd Law of Thermodynamics a. ΔSuniv = 0
E Entropy b. ΔSuniv < 0
A Process at equilibrium c. ΔSuniv > 0
C Spontaneous process d. A particular combination of motions and locations of the
atoms and molecules of a system at a particular instant.
B Nonspontaneous process e. Thermodynamic state function that is associated with the
extent of disorder in a system or how energy of a reaction
gets distributed across the molecules.
D microstate f. S = k(ln W) where k = 1.38 x 10-23 J/K
F Boltzmann’s equation g. In any spontaneous process, the entropy of the universe
increases (ΔSuniv = ΔSsys + ΔSsurr).
2. Determine whether each of the following process is spontaneous or non-spontaneous.
A. dissolving table salt in hot soup spontaneous
B. climbing A Mountain nonspontaneous
C. opening a bottle of perfume to allow the fragrance to spread through the room spontaneous
D. separating helium and neon from a mixture of gases nonspontaneous
E. water solidifying to ice at 3 °C nonspontaneous
3. For each of the following pairs, identify which one has a greater number of microstates
(Ω):
a. CH4or C2H6
b. H2O(l) or H2O(s)
c. 1 mol O2 occupying 1L a 25°C or 1 mol O2 occupying 2L at 25°C
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Ch. 17 Kimball Spring 2019
4. Predict whether each of the following reactions would have a positive or negative entropy change:
A. S(s) + O2 (g) → SO2 (g) ΔS > 0
B. H2O(g) → H2O(l) ΔS < 0
C. PCl3 (g) + Cl2 (g) → PCl5 (g) ΔS < 0
D. 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2 (g) ΔS > 0
E. N2 (g) → 2N(g) ΔS > 0
5. Use data from Appendix C to calculate ΔH°, ΔS°, and ΔG° at 298 K for the following
reaction, and determine whether the reaction is spontaneous at 298K.
2P(g) + 10HF(g) → 2PF5 (g) + 5H2 (g)
ΔH° = - 1135.5 kJ/mol
ΔS° = -807.00 J/molK
ΔG° = -895.0 kJ/mol
6. For each of the following reactions, calculate ΔG° at 298K. Determine under what
conditions the reaction is spontaneous. If the reaction is not spontaneous under standard
conditions, at what temperature would it become spontaneous?
a. N2 (g) + 3F2 (g) → 2NF3 (g) ΔH° = -249 kJ/K ΔS° = -278 J/K
ΔG° = - 166.16 kJ/mol
spontaneous becomes nonspontaneous where ΔG=0, T = 895 K
b. N2 (g) + 3Cl2 (g) → 2NCl3 (g) ΔH° = 460 kJ ΔS° = -275 J/K
ΔG° = 541.95 kJ/mol
nonspontaneous at all T
c. N2F4 (g) → 2 NF2 (g) ΔH° = 85 kJ ΔS° = 198 J/K
ΔG° = 26.0 kJ.mol
Nonspontaneous becomes spontaneous where ΔG=0, T = 429 K
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Ch. 17 Kimball Spring 2019
7. Consider the following Bronstead acid-base reaction at 25°C:
HF(aq) + Cl-(aq) ⇌ HCl(aq) + F-(aq)
A. Predict whether K will be greater or smaller than unity.
HCl is the stronger acid, so equilibrium favors the reactans. K < 1
B. Does ΔS° or ΔH° make a greater contribution to ΔG°?
There are approx.. the same number of aqueous particles on both sides of rxn, so ΔS will be
small, therefore ΔH will contribute more
C. Is ΔH° likely to be positive or negative?
H – F bond is stronger than H – Cl bond, so calculating ΔH using bond energies would show
ΔH > 0
8. Use the data in Appendix C to estimate the boiling point of methanol, CH3OH(l) .
boiling point T = 33K or 65°C
9. Using the data from Appendix C, write the equilibrium constant expression and calculate
the free energy change and the equilibrium constant for these reactions at 298K.
a. C2H5OH (g) ⇌ C2H4 (g) + H2O (g)
ΔG° = 8.04 kJ/mol
K = 0.039
b. CaCO3 (s) ⇌ CaO (s) + CO2 (g)
ΔG° = 130.18 kJ/mol
K = 1.8 x 10-23
10. Consider the reaction
2NO (g) + O2 (g) ⇌ 2NO2 (g)
a. Using data in Appendix C, calculate ΔG° at 298 K. ΔG° = - 69.74 kJ/mol
b. Calculate ΔG at 298 K if the reaction mixture consists of 2.75 atm of NO, 1.75 atm
of O2 , and 0.115 atm of NO2.
ΔG = -86.9 kJ/mol
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