9.6 End-of-Chapter Material

Exercises

1. Classify each of the following as spontaneous or nonspontaneous processes:

a)  the browning of a cut apple slice on a snack tray over time

b)  the rolling of a ball uphill

c)  the formation of diamond from the graphite in your pencil

d)  the melting of ice cubes in a glass of water you are holding

 

2. State the second law of thermodynamics.

 

3. What sign would you expect for ΔS for the following processes?

a)  water freezing in a lake during a cold Alberta winter

b)  AB(s) + CD(s) → AC(g) + BD(g)

c)  a balloon with a fixed amount of gas stretched to a larger volume

d)  the sublimation of dry ice

e)  3 E2F2(g) → E6F6(g)

 

4. Which of the following would you expect to have the higher standard molar entropy, S⁰?

a)  Br2() or Br2(g)

b)  NO2(g) or NO(g)

c)  C2H6(g) or C5H12(g)

 

5. Calculate the ΔS⁰ of the following reactions using the values listed in the appendix.

a)  4 NH3(g) + 5 O2(g) → 6 H2O(g)  + 4 NO(g)

b)  2 HgO(s) → 2 Hg()  + O2(g)

c)  3 FeCl2(s) + KNO3(s) + 4 HCl (aq) → 3 FeCl3(s) + KCl(s) + NO(g) + 2 H2O()

6. Draw a diagram showing the approximate entropy change of water from -100⁰C to 250⁰C.

 

7. A chemical reaction has ΔH⁰ = -43.5 kJ and ΔS⁰ = -65.8 J/K. Calculate ΔG⁰ at 298 K. Is this a spontaneous process?

 

8. Using data from the appendix, determine ΔG⁰ for all reactions listed in question 5.

 

9. Find the standard Gibbs energy change for the reaction

CaCO3(s) → CaO(s) + CO2(g)

The ΔGf° values for the three components of this reaction system are CaCO3(s): –1128 kJ mol–1; CaO(s): –603.5 kJ/mol; CO2(g): –137.2 kJ/mol.[1]

 

10. Using the appendix data, determine the approximate temperature at which the following processes are at equilibrium:

a)  CH2Cl2() → CH2Cl2(g)

b)  CH3OH() → CH3OH(g)

 

11. Determine if the following reactions would be spontaneous at any temperature, nonspontaneous at any temperature, spontaneous at low temperature but not high temperature, or spontaneous at high temperature but not low temperature:

 

a)  ΔH⁰ = -750 kJ and ΔS⁰ = 250 J/K

b)  ΔH⁰ = -100 kJ and ΔS⁰ = -300 J/K

c)  ΔH⁰ = 95 kJ and ΔS⁰ = -70 J/K

d)  ΔH⁰ = 400 kJ and ΔS⁰ = 500 J/K

Answers

1.

a)  spontaneous

b) nonspontaneous

c)  nonspontaneous

d)  spontaneous

3.

a)  ΔS = –

b)  ΔS = +

c)  ΔS = +

d)  ΔS = +

e)  ΔS = –

5.

a)

ΔS⁰ = ∑nS⁰(products) – ∑mS⁰(reactants)

ΔS⁰ = (6 S⁰ H2O(g)  + 4 S⁰ NO(g)) – (4 S⁰ NH3(g)  + 5 S⁰ O2(g))

ΔS⁰ = [(6 x 188.8 J/mol K) + (4 x 210.8 J/mol K)] – [(4 x 192.8 J/mol K) + (5 x 205.2 J/mol K)]

ΔS⁰ = 178.8 J/mol K

b)

ΔS⁰ = ∑nS⁰(products) – ∑mS⁰(reactants)

ΔS⁰ = (2 S⁰ Hg() + S⁰ O2(g)) – (2 S⁰ HgO)

ΔS⁰ = [(2 x 75.9 J/mol K) + (205.2 J/mol K)] – (2 x 70.3 J/mol K)

ΔS⁰ = 216.4 J/mol K

c)

ΔS⁰ = ∑nS⁰(products) – ∑mS⁰(reactants)

ΔS⁰ = (3 S⁰ FeCl3(s) + S⁰ KCl(s) + S⁰ NO(g) + 2 S⁰ H2O())- (3 S⁰ FeCl2(s) + S⁰ KNO3(s) + 4 S⁰ HCl (aq))

ΔS⁰ = [(3 x 142.3 J/mol K) + (82.6 J/mol K) + (210.8 J/mol K) + (2 x 70.0 J/mol K)] – [(3 x 118.0 J/mol K) + (133.1 J/mol K) + (4 x 56.5 J/mol K)]

ΔS⁰ = 147.2 J/mol K

7.

ΔG⁰ = ΔH⁰ – TΔS

ΔG⁰ = -43.5 kJ – [(298 K)(-65.8 J/K)]

ΔG⁰ = -23.9 kJ

ΔG⁰ is negative therefore this is a spontaneous process.

9.

Δ = (–603.5 –137.2) – (–1128) kJ/mol = +130.9 kJ/mol, indicating that the process is not spontaneous under standard conditions (i.e., solid calcium carbonate will not form solid calcium oxide and CO2 at 1 atm partial pressure at 25°C.)

11.

a)  spontaneous at all temperatures

b)  spontaneous at low temperatures, nonspontaneous at high temperatures

c)  nonspontaneous at all temperatures

d)  spontaneous at high temperatures, nonspontaneous at low temperatures


  1. Question and solution from Chem1 Virtual Textbook, Stephen Lower/CC-BY-SA-3.0

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9.6 End-of-Chapter Material by David W. Ball and Jessie A. Key is an adapted version of the open textbook Introductory Chemistry – 1st Canadian and is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.

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