Chemical Thermodynamics

Objectives:

After having completed this unit, students should be able to:

1. Define or explain and give an example of each of the following terms.
   1. spontaneous process
   2. state function (see pg 167)
   3. system
   4. surroundings
   5. reversible and irreversible processes
2. Explain how entropy is related to randomness or disorder, positional probability, and the number of microstates available to the system.
3. Predict whether a process is spontaneous.
4. Predict the relative number of microstates available to a system under varying conditions, or using your understanding of the relative number of microstates available to systems predict whether entropy of a system increases or decreases.
5. Predict the sign of the entropy change for a given process, phase change, or chemical reaction.
6. Calculate ΔS° for a chemical reaction using the tabulated S° values of the reactants and products and a balanced equation.
7. State and explain the second law of thermodynamics in terms of entropy. Understand the relationship of the entropy change in the system, surroundings, and universe to the second law.
8. Explain the temperature dependence of entropy and predict how spontaneity can be affected by increasing or decreasing temperature.
9. Given any two of the following variables, ΔS, ΔH or T calculate the third variable.
10. Define free energy and relate it to spontaneity.
11. Given ΔH, ΔS, and T, calculate the change in Gibbs Free Energy and then predict whether a process will be spontaneous.
12. Define and write the equation for the standard free energy of formation, ΔG_f° of a substance.
13. Calculate ΔG°, ΔH° and ΔS° for a reaction from tabulated data.
14. Given temperature and any two of the following variables, ΔS, ΔH or ΔG,  calculate the third variable
15. Explain how ΔG changes with a change in the concentration or partial pressure of reactants or products.
16. Calculate ΔG given the concentration or partial pressures of reactants and products.
17. Explain equilibrium in terms of minimum free energy and show how the value of K is related to free energy.
18. Calculate ΔG° from K and vice-versa.
19. Predict how ΔG° will change with temperature given ΔH°  and ΔS°.
20. Explain the relationship between the maximum work obtainable and the free energy change.
21. Explain the relationship between kinetics and thermodynamics.

Reading and Homework Exercises

Table of Contents from: OpenStax Chemistry 2e:

• Chapter 5, section 5.3  (Review of Enthalpy and 1^st Law of Thermodynamics)
• Chapter 16 – Thermodynamics

5.3 Enthalpy

Question 1:

Question 2:

Question 3:

For the reaction:

2 C₆H₆ (l) + 15 O₂ (g) → 12 CO₂ (g) + 6 H₂O(g)

Using the data in Appendix G, answer the following questions:

Question 4:

Question 5:

Given the following data:

2Fe(s) + 3/2 O₂(g) —-> Fe₂O₃(s)   ΔH°= -822.2 kJ/mol

CO(g) + 1/2 O₂(g) —-> CO₂(g)     ΔH°= -383.3 kJ/mol

Determine ΔH° for the reaction

Fe₂O₃(s) + 3CO(g)   —-> 3CO₂(g) + 2Fe(s)

16 Introduction

16.1 Spontaneity

16.2 Entropy

Question 6:

Question 7:

Question 8:

Question 9:

16.3 The Second and Third Laws of Thermodynamics

Question 10:

Question 11:

Question 12:

For the reaction C₂H₂(g) + 2 H₂(g) → C₂H₆(g), answer the following question.  Use data in Appendix G

Question 13:

For the reaction 2 BaO (s) → 2 Ba (s) + O₂ (g), answer the following question. Use data in Appendix G

Question 14:

16.4 Free Energy

Question 15:

Question 16:

Question 17:

 

The boiling of an unknown compound X can be represented as follows:

X (l) → X (g)

ΔH° = 42.8 kJ/mol and ΔS° = 110.9 J/mol⋅K

Question 18:

Consider the reaction 3 CH₄ (g) → C₃H₈ (g) + 2 H₂ (g)

The reaction occurs at 25°C.

Using the data in Table 1, answer the following questions:

Table 1

Formula	ΔHf° (kJ/mol)	S° (J/mol⋅K)
CH4 (g)	-74.6	186.3
C3H8 (g)	-103.8	270.3
H2 (g)	0	130.7

Question 19:

Question 20:

Question 21:

Given the following data:

2Fe(s) + 3/2 O₂(g) —-> Fe₂O₃(s)   ΔG°= -742.2 kJ/mol

CO(g) + 1/2 O₂(g) —-> CO₂(g)     ΔG°= -257.2 kJ/mol

Determine ΔG° for the reaction

Fe₂O₃(s) + 3CO(g)   —-> 3CO₂(g) + 2Fe(s)

Question 22:

Question 23:

Consider the reaction NH₄Cl(g) → NH₃ (g) +  HCl (g)

The reaction occurs at 298 K.

Using the data in Table 2, answer the following questions:

Table 2

Formula	ΔHf° (kJ/mol)	ΔGf° (kJ/mol)
NH3 (g)	-46.1	-16.5
HCl (g)	-92.3	-95.3
NH4Cl(g)	-314.4	-203

Question 24:

Question 25:

Consider the reaction Ca(OH)₂ (s) ↔ Ca²⁺ (aq) + 2 OH^– (aq)

The reaction occurs at 25°C.  For [Ca²⁺] = 1.5 M and pOH = 8.00 , and using ΔG_f^° data from Appendix G answer the following questions:

Question 26:

Consider the reaction HC₂H₃O₂ (aq) ↔ H⁺ (aq) + C₂H₃O₂^– (aq)

The reaction occurs at 25°C.  Using the Ka data in Appendix H, answer the following questions:

Question 27: