Spontaneity Of Chemical Processes 557 Ch.12 Full Test Bank

CHAPTER 12

SPONTANEITY OF CHEMICAL PROCESSES

CHAPTER STUDY OBJECTIVES

1. Recognize the driving force behind all chemical change: dispersal of energy and matter.

KEY CONCEPTS: A process that is spontaneous in one direction is non-spontaneous in the opposite direction. Energy in the form of heat always flows from a warmer object to a colder one. The spontaneity of any process must be evaluated.

2. Predict the direction of change based on the entropy changes in the system and in the

surroundings.

SKILLS TO MASTER: Using Boltzmann’s equation to calculate entropy of a system; calculating entropy changes from the flow of heat at constant temperature

KEY CONCEPTS: The second law of thermodynamics states that in any spontaneous process, the direction of change is such that total entropy increases: ∆S_total > 0.

3. Understand and calculate entropies of pure substances.

SKILLS TO MASTER: Calculating absolute entropies; calculating entropy changes with concentration or pressure; calculating standard reaction entropies

KEY CONCEPTS: The third law of thermodynamics states that a pure, perfect crystal at 0ºK has zero entropy. Entropy increases as the temperature increases, when a solid melts, and when a liquid evaporates.

4. Predict the direction of spontaneous change using the reaction free energy.

SKILLS TO MASTER: Calculating the standard free energy change for a reaction from standard free energies of formation; calculating the standard free energy change for a reaction from standard enthalpies and entropies of reaction; calculating a reaction quotient; calculating the free energy of reaction under nonstandard conditions; predicting the influence of temperature on spontaneity

KEY CONCEPTS: The reaction free energy is negative for any spontaneous process at constant temperature and pressure. A reaction free energy is a function of concentrations (or partial pressures) and of the absolute temperature.

5. Apply thermodynamics to chemical reactions and phase changes.

SKILLS TO MASTER: Calculating vapour pressure using the Clausius–Clapeyron equation; calculating the enthalpy of vapourization from vapour pressure data

6. Describe thermodynamically some representative energetic processes that operate in living organisms.

KEY CONCEPTS: Adenosine triphosphate is used in the body to store energy. Cells obtain energy to drive otherwise non-spontaneous reactions via the dephosphorylation of ATP.

Multiple Choice QUESTIONS

1. Which of the following arrangements of 4 identical socks in a 4-drawer dresser has the least degree of disorder?

a) 1 pair in 1 drawer and 1 pair in another drawer

b) 2 pair in 1 drawer

c) 1 pair in 1 drawer and 1 sock in 2 other drawers

d) 1 sock in each drawer

e) 3 socks in one drawer, 1 sock in another

Difficulty: Easy

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

2. Considerable energy is required to vapourize water, yet water does evaporate as demonstrated by clothes drying when hung out. The reason for this spontaneous process is

a) energy is always conserved.

b) an increase in disorder.

c) energy is released to the surroundings.

d) water has a higher vapour pressure outside rather than inside of a house.

e) an increase in order.

Difficulty: Medium

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

3. The reaction shown below is exothermic:

2 NO₂ (g)  N₂O₄(g)

At low pressures it is NOT spontaneous, but at high pressures, significant amounts of product are formed. This is because

a) the heavier products cause greater disorder upon collisions.

b) the smaller number of product molecules leads to more disorder in the container.

c) to place the system under high pressure, the surroundings are more disordered.

d) the higher pressures make the reaction proceed more rapidly.

e) at low pressure more molecules are required to fill the same space.

Difficulty: Hard

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

4. The exothermic oxychlorination of ethylene to vinyl chloride is shown below:

CH₂CH_2(g) + HCl_(g) + O_2(g)  CH₂CHCl_(g) + H₂O_(g)

What are the signs for the entropy change of the system and surroundings, respectively?

a) +ve, +ve

b) +ve, -ve

c) –ve,-ve

d) –ve, +ve

Difficulty: Medium

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

Feedback: a) entropy change is –ve; b) entropy change is –ve, must recognize that while exothermic corresponds to a –ve enthalpy, this translates to a +ve entropy for the surroundings; c) exothermic corresponds to a –ve enthalpy, this translates to a +ve entropy for the surroundings; d) correct answer

5. What are the signs for the entropy change of the system and surroundings, respectively, for ethane burning, when the system is defined as ethane and oxygen?

a) <0, <0

b) <0, >0

c) 0, 0

d) >0, <0

e) >0, >0

Difficulty: Medium

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

6. What are the signs for the entropy change of the system and surroundings, respectively, for water evaporating at 370ºK, when the system is defined as water?

a) <0, <0

b) <0, >0

c) 0, 0

d) >0, <0

e) >0, >0

Difficulty: Medium

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

7. Given three identical cups and three identical marbles, which of the following arrangements has the greatest value of W?

a) one marble in each cup

b) two marbles in one cup, one marble in one of the other cups

c) all three marbles in one cup

d) all have the same values of W

e) two of them have the same value of W

Difficulty: Easy

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

8. Which of the following samples of substances has the smallest entropy change for fusion at its melting point, T_f?

Substance Sample size, (g) MM (g/mole) ∆H_f (J/mole) T_f (K)

a) CO₂ 50 44.011 7949.6 195

b) H₂O 25 18.01594 6010 273.15

c) C₆H₆ 80 78.11 9951 278.68

d) CH₃OH 50 32.04 3177 175.25

e) C₃H₈O₃ (glycerol) 95 92.09 8475 291.17

Difficulty: Hard

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

9. Which of the following selections has the molecules arranged in order of increasing standard entropy? (One mole of each substance is being compared.)

a) C(graphite), C₆₀(s), C(diamond)

b) Si(s), P₄(s), S­₈(s)

c) C₂H₆(g), C₂H₄(g), CH₄(g)

d) Polystyrene (made of 1000 monomer units), DNA (1000 base pairs), polyethylene (1000 monomer units)

e) CH₄(g), C₂H₆(g), C₂H₄(g)

Difficulty: Medium

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

10. The standard molar entropy of NO is 211 J K^-1 mol^-1. What is the standard entropy of 3 moles of NO at 10 bar?

a) 202 J K^-1 mol

b) 606 J K^-1

c) 624 J K^-1

d) 633 J K^-1

e) 606 J K^-1 mol^-1

Difficulty: Medium

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

Feedback: a) question is asking for standard entropy of 3 moles, this is per mol; b) correct answer; c) 3(211) – 9; d) fails to account for change in pressure from 1 bar; e) incorrect units

11. An important reagent in organic chemistry is thionyl chloride, SOCl₂. One synthetic route would be to react SO₂ with HCl in the gas phase to produce SOCl₂ and water. What is the entropy change for this reaction? (J/K•mole SOCl₂)? (S˚ = 309.7(SOCl₂), 248.2(SO₂), 186.9(HCl), 188.8(H₂O) in J/K•mole)

a) –123

b) 123

c) –246

d) –146

e) 61.5

Difficulty: Hard

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

12. One process water can undergo is dissociation, the equation for which is shown below:

H₂O _(l)  H⁺ (aq) + OH^- (aq)

As temperature increases to 90°C,

a) there will be no shift in the position of equilibrium.

b) more reactants will be present at equilibrium.

c) the extent of dissociation will be less.

d) the extent of dissociation will be greater.

e) the water will vapourize rather than dissociate.

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

13. A reaction will NEVER be spontaneous when

a) ∆G° > 0.

b) ∆H°< 0 and ∆S° > 0.

c) ∆H°< 0 and ∆S° < 0.

d) ∆G=0.

e) ∆G > 0.

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

14. Isomerization of hydrocarbons is important in the refining of petroleum. A simple example is that of butane, which has only two isomers:

Which of the following statements is most correct?

a) At standard temperatures, n-butane is the preferred isomer.

b) At temperatures less than 450, n-butane is the preferred isomer.

c) At temperatures greater than 450ºK, n-butane is the preferred isomer.

d) At ambient temperatures, a sample of n-butane and isobutene which was at equilibrium would contain barely any butane.

e) A 1:1 mixture of butane and isobutane is at equilibrium at room temperature.

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

15. Consider the following chemical reaction:

CaSO₄(s) + 2H₂O (g) 🡪 CaSO₄ – 2H₂O (s)

This reaction is exothermic.

a) The reaction is spontaneous at all temperatures.

b) The reaction is never spontaneous.

c) The reaction is spontaneous at low temperature.

d) The reaction is spontaneous at high temperature.

e) There is not sufficient information given to assess the spontaneity of the reaction.

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

16. Consider the following unbalanced reaction:

CO₂(g) + H₂O(g) 🡪 O₂(g) + C₃H₈(g)

This reaction is endothermic.

a) The reaction is spontaneous at all temperatures.

b) The reaction is never spontaneous.

c) The reaction is spontaneous at low temperature.

d) The reaction is spontaneous at high temperature.

e) There is not sufficient information given to assess the spontaneity of the reaction.

Difficulty: Medium

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

Feedback: Reaction must be balanced before spontaneity can be assessed. Balance reaction has 7 moles of gas converted to 6 moles, therefore entropy is –ve and reaction is never spontaneous.

17. Which of the following energy producing technologies has no effect on the environment?

a) nuclear power

b) geothermal power

c) hydroelectric power

d) solar power

e) No energy producing technology is free from effects on the environment.

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

18. A child plays with a deck of cards and leaves the room. You politely pick up and organize the cards. Which of the following is true during the action of organizing the cards?

a) You become more ordered.

b) The universe becomes more ordered.

c) The deck of cards becomes more disordered.

d) You become more disordered.

e) The ordering of a deck of cards is spontaneous.

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

19. Which of the following processes generate thermal pollution?

I. Hydroelectric generation of electricity

II. Nuclear generation of electricity

III. Coal powered generation of electricity

IV. Natural gas/oil powered generation of electricity

V. Wind farms

a) all of the above

b) I, II, III, and IV

c) III and IV

d) II, III, IV and V

e) II, III, IV

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

20. The process of storing energy for short term in the body is by making ATP. What type of bond is made or broken in this process?

a) making of a phosphate ester (P-O-P) bond

b) breaking of a phosphate ester bond

c) addition of water to ADP

d) the forming of a C=O bond

e) the forming of a C=C bond

Difficulty: Easy

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

21. How does your body perform non-spontaneous reactions?

a) coupling these reactions with spontaneous reactions

b) Non-spontaneous reactions do not occur in our body.

c) by directly using sugars or fats for energy for these reactions

d) by only using ADP

e) All reactions in the body are spontaneous.

Difficulty: Easy

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

22. The following reaction has a ΔG_rxn = 2872 kJ/mol. How is it that this reaction happens in plants on a regular basis?

6 CO_2(g) + 6 H₂O_(l) 🡪 C₆H₁₂O_6(aq) + 6 O_2(g)

a) ΔG says that the reaction is spontaneous.

b) This reaction is coupled with spontaneous reactions.

c) Energy is released to compensate for ΔG.

d) This reaction happens in humans not plants.

e) It happens is small amounts so that not much energy is needed.

Difficulty: Easy

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

ESSAY QUESTIONS

23. Liquids composed of different molecules have different degrees of order. Compare the liquid phases of H₂S and H₂O and suggest which is more “ordered.”

Difficulty: Easy

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

24. Ammonium nitrate, [NH₄][NO₃], dissolves spontaneously, even though the process is sufficiently endothermic to make it useful in “cold packs”. Explain this phenomenon from the order-disorder perspective.

Difficulty: Medium

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

25. State whether the disorder of each of the following systems increases or decreases in the stated process without consulting any tables:

i. 2 S _(s) + 3 O2 _(g)  2 SO3 _(g)

ii. CO2 _(g)  CO2 _(s)

iii. heating water from 5 °C to 75 °C

Difficulty: Easy

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

26. State whether the disorder of each of the following systems increases or decreases in the stated process without consulting any tables:

i. 6 CO_2(g) + 6 H₂O_(l)  C₆H₁₂O_6(aq) + 6 O_2(g)

ii. CH_4(g) + 2 O_2(g)  CO_2(g) + 2 H₂O_(l)

iii. CH₃CH₂OH_(l)  CH₃CH₂OH_(g)

Difficulty: Medium

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

27. Compare the vapourization of H₂S and H₂O. For which of these will you expect the greater increase in disorder and why?

Difficulty: Medium

Learning Objective: Recognize the driving force behind all chemical change: dispersal of energy and matter.

Section Reference: 12.1 Spontaneity

28. What is the total entropy change when 40.5 g of ice (∆Hfus = 6.01 kJ/mole) melts at 0^oC?

Difficulty: Medium

Feedback: convert grams to moles to determine q

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

29. What is the total entropy change when 43 ml of water (∆Hfus = 6.01 kJ/mole, density 0.9981 g ml^-1) freezes at 0^oC?

Difficulty: Medium

Feedback: convert grams to moles using molar mass and density then determine q transferred.

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

30. Consider a 1 kg block of ice melting in a 35^oC chamber (∆Hfus = 6.01 kJ/mole). Determine the total entropy change for the ice and the chamber.

Difficulty: Hard

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

Feedback: Convert kg to moles; calculate q for melting ice at 0^oC, entropy of melting ice is determined at 273.15^oK, entropy of surroundings at 308.15^oK, total entropy is the sum of the two.

31. Calculate the entropy change for the fusion of 1 mole of hydrogen sulphide, mp = -85.6°C and ∆Hfus = 2.39 kJ/mole.

Difficulty: Easy

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

32. Calculate the entropy change of the universe for 1 mole of solid water melting at -5°C.

Difficulty: Easy

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

33. What is the entropy change in the surroundings when the system is a 2.2 kg block of ice that melts in a swimming pool (surrounding) whose temperature is 31°C?

Difficulty: Medium

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

34. What is the entropy change in the surroundings when a 25 g piece of ice, the system, melts in the mouth of a human (surrounding) which maintains a temperature of 37°C?

Difficulty: Medium

Learning Objective: Predict the direction of change based on the entropy changes in the system and in the surroundings.

Section Reference: 12.2 Entropy: The Measure of Dispersal

35. Write the definition of the 3 Laws of Thermodynamics.

Difficulty: Easy

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

36. What is the entropy change for the reaction of graphite (S˚=5.74 J/mol•K) and hydrogen (S˚=130.68) to produce methane (S˚=186.2)?

Difficulty: Easy

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

37. In adjusting the acidity of swimming pool water, concentrated aqueous HCl is added. What is the entropy for the chloride ions after 0.5 L of 12.0 M HCl is added to a swimming pool of 9.6 x 104 L (about 25,000 gallons)? (S° Cl- _(aq) = 56.5 J/mole • K)

Difficulty: Medium

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

38. One mole of SF₆ (g) (S°= 291.82 J/mole•K) is at STP. When released into a large room, the final concentration is 1 PPM by volume. What is the entropy change for the SF₆ upon this expansion?

Difficulty: Hard

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

39. Determine the change in entropy and if the following unbalanced reaction is spontaneous or non-spontaneous with respect to entropy.

Ba(OH)₂•8H₂O_(s) + NH₄NO­_3(s) 🡪 Ba(NO₃)_2(s) + H­₂O_(l) + NH_3(aq)

Difficulty: Hard

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

40. An electric heater releases 1200 J of energy every second. How much does the entropy of the surroundings increase when the heater runs for 10 minutes in a room that is at 25°C?

Difficulty: Medium

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

41. Hydrogen sulphide is a toxic gas, which fortunately has the distinctive odour of rotten eggs. What is the entropy (J/K) of 0.01 mole of H₂S (S˚=205.8 J/K•mole) which occupies a volume of 10⁶ L at a temperature of 298°K?

Difficulty: Medium

Learning Objective: Understand and calculate entropies of pure substances.

Section Reference: 12.3 Entropies of Pure Substances

42. Calculate the free energy for the oxychlorination of ethylene to vinyl chloride, CH₂CHCl, under standard conditions.

ΔG˚ = 53.6(CH₂CHCl), –95.3(HCl), 68.49(CH₂CH₂), –228.7(H₂O) kJ/mol

The equation of reaction is:

CH₂CH_{2 (g)} + HCl _(g) + O_{2 (g)} 🡪 CH₂CHCl _(g) + H₂O _(g)

Difficulty: Easy

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

43. Calculate the free energy for the oxychlorination of ethylene to vinyl chloride, CH₂CHCl, under standard conditions at 100^oC given the thermodynamic data below collected at 25^oC:

ΔG˚ = 53.6(CH₂CHCl), –95.3(HCl), 68.49(CH₂CH₂), –228.7(H₂O) kJ/mol

ΔH˚ = 37.2(CH₂CHCl), –92.3(HCl), 52.4(CH₂CH₂), –241.8(H₂O) kJ/mol

The equation of reaction is:

CH₂CH_{2 (g)} + HCl _(g) + O_{2 (g)} 🡪 CH₂CHCl _(g) + H₂O _(g)

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

Feedback: This problem requires several steps. Must calculate ΔG^o and ΔH^o at 25^oC from data given and use this to determine ΔS^o. The values of ΔH^o and ΔS^o can then be used to find ΔG^o at temperature other than 25^oC.

44. The conversion of calcite,CaCO_3(s), to CaO_(s) and CO_2(g) is NOT favourable at ambient temperatures. At what temperature will this decomposition of CaCO₃ become spontaneous?

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

45. Determine if the following unbalanced reaction is spontaneous or non-spontaneous at standard conditions. (ΔG˚ = –857, SiO₂; 0, Si; 0, Al; –1582 kJ/mol)

SiO_2(s) + Al_(s) 🡪Al₂O_3(s) + Si_(s)

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

46. What is the reaction quotient for the following unbalanced cell reaction?

Ba(OH)₂•8H₂O_(s) + NH₄NO­_3(s) 🡪 Ba(NO₃)_2(s) + H­₂O_(l) + NH_3(aq)

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

47. Determine Gibb’s Free Energy and if the following unbalanced reaction is spontaneous or non-spontaneous at standard conditions.

Ba(OH)₂•8H₂O_(s) + NH₄NO­_3(s) 🡪 Ba(NO₃)_2(s) + H­₂O_(l) + NH_3(aq)

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

48. Determine Gibb’s Free Energy and if the following unbalanced reaction is spontaneous at 0^oC.

Ba(OH)₂•8H₂O_(s) + NH₄NO­_3(s) 🡪 Ba(NO₃)_2(s) + H­₂O_(l) + NH_3(aq)

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

Feedback: This is a multistep problem. ΔH^o and ΔS^o are determined for the reaction and used to calculate ΔG^o at 0^oC. Important to use temperature in Kelvin.

49. At what identical concentration of H⁺ and CH₃CO₂^- (∆G°_f = -372.5 kJ) will the dissociation of 1 M acetic acid, below, (∆G°_f (CH₃CO₂H (aq)) = -399.6 kJ) become spontaneous at room temperature?

CH₃CO₂H (aq) 🡪 H⁺(aq) + CH₃CO₂^-(aq)

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

50. The possible isomerization for ethanol to methyl ether is shown below:

What is a) the value of the reaction quotient when ∆G=0 and b) will the predominant species be ethanol or methyl ether?

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

51. Complete the following table for determining at what relative temperature a reaction will be spontaneous.

Difficulty: Easy

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

52. One of the more important industrial chemicals is hydrogen. One process for hydrogen production is called “steam reforming”, in which hydrocarbons react with water to give hydrogen and CO. The equation of reaction for reforming methane is written below:

CH_{4 (g)} + H₂O _(g) 🡪 CO _(g) + 3 H_{2 (g)}

a) Calculate the free energy change for this reaction at standard conditions.

b) Estimate the temperature at which the process becomes spontaneous.

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

53. Calcite, CaCO_3(s), can be converted to CaO_(s) and CO_2(g). Determine the pressure of CO₂ at 1150 K at which the reaction is no longer spontaneous.

Difficulty: Hard

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

Feedback: Multistep problem requiring determination of ΔG^o, the expression for the reaction quotient, Q, and solving for pressure of carbon dioxide.

54. One process for hydrogen production is called “steam reforming”, in which hydrocarbons react with water to give hydrogen and CO. The equation of reaction for reforming methane is written below.

CH_{4 (g)} + H₂O _(g) 🡪 CO _(g) + 3 H_{2 (g)}

What is the free energy change, ∆G, for reforming methane at 298^oK when the partial pressure of the products is 5.0 x 10-5 atm and the partial pressure of the reactants is 1.0 atm?

Difficulty: Medium

Learning Objective: Predict the direction of spontaneous change using the reaction free energy.

Section Reference: 12.4 Spontaneity and Free Energy

55. What are the signs of ΔH, ΔS, and ΔG for the process of ice melting at 10˚C and 1 atm?

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

56. Given the following chemical reaction at 298^oK:

3 O_­2(g) 🡪 2 O_3(g) ΔG˚ = 344 kJ/mole, what would be ΔG if the partial pressure of O₂ = 0.20 bar and O₃ = 0.0001 bar?

Difficulty: Medium

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

57. An explosive that is also used as fertilizer is ammonium nitrate. Calculate ΔG˚ for one of the decomposition reactions this fertilizer can undergo. (ΔG˚ = -183.9 (NH­₄NO₃), 103.7 (N₂O), -228.7 (H₂O) in kJ/mol)

NH₄NO_3(s) 🡪 N₂O_(g) + 2 H₂O_(g)

Difficulty: Easy

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

58. At 45˚C, what is the vapour pressure of Iodine?

(I_2(s): ΔH˚ = 0, S˚ = 116.1 J/mol•K) (I_2(g): ΔH˚ = 62.4 kJ/mol, S˚ = 260.7 J/mol•K)

Difficulty: Hard

Learning Objective: Apply thermodynamics to chemical reactions and phase changes.

Section Reference: 12.5 Some Applications of Thermodynamics

59. Estimate how many moles of ATP would be required to ride a bicycle at 35 km/hour for 1 hour, an activity that would require about 1500 kJ of energy, given

ATP 🡪 ADP + H₂O, ∆G° = -30.6 kJ

Difficulty: Easy

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

60. Glucose (MM = 180g/mole) is metabolized at a rate of about 38% efficiency. How many grams of glucose would supply the energy to ride a bicycle at 35 km/hour for 1 hour, an activity that would require about 1500 kJ of energy?

Difficulty: Hard

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

61. Palmitic acid, C₁₅H₃₁CO₂H undergoes complete combustion with ΔG_rxn ^o of –9790 kJ mol^-1. Using bond energies, estimate the entropy change associated with this reaction.

Difficulty: Hard

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

Feedback: Uses concepts covered in Chapter 3; requires student to correctly balance combustion rxn.

62. Chemical energy stored in glucose can be used to convert ADP to ATP; under normal physiological conditions the process, shown below, is 38.3 % efficient:

C₆H₁₂O₆ + 6O₂ + 36 ADP + 36 H₃PO₄ 🡪 6CO₂ + 36 ATP + 42 H₂O ∆G° =?

Given that the reaction: ADP + H₃PO₄ 🡪 ATP + H₂O requires 30.6 kJ mol^-1 determine ∆G° for the reaction of glucose and ADP above.

Difficulty: Hard

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

63. The formation of ATP from ADP requires 30.6 kJ/mol. The oxidation of the fat, palmitic acid, to carbon dioxide yields –9790 kJ/mol of energy. Yet your body is only able to produce 130 ATP units per 1 palmitic acid unit. What is the percent yield for this biological process?

Difficulty: Medium

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

64. A person exercises for 30 minutes and burns 800 kJ of free energy. How many moles of ATP (–30.6 kJ/mol) does this represent and based on 38% efficiency, how many grams of glucose were burned if 1 mole of glucose can make 36 ATPs?

Difficulty: Hard

Learning Objective: Describe thermodynamically some representative energetic processes that operate in living organisms.

Section Reference: 12.6 Bioenergetics

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