CHEM126 – CHAPTER 17 – REGULAR HOMEWORK / SUBJECT I/II

 

CHAPTER 17 BASIC/REGULAR HOMEWORK:

SUBJECT I/II

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Question 1

 

Calculate ΔGº (in kJ) at 298 K for the following reaction:

2H2O(g) +2Cl2(g) → 4HCl(g) + O2(g)

Use the data in Appendix II (Tro); Appendix 4 (Zumdahl)
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Explanation

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Question 2

 

Is the above reaction under standard conditions spontaneous in the forward direction?
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Explanation

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Question 3

 

Calculate ΔHº (in kJ) for the following reaction at 298 K:

2CO2(g) +4H2O(l) → 2CH3OH(l) + 3O2(g)

Use the data in Appendix II (Tro); Appendix 4 (Zumdahl)
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Explanation

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Question 4

 

Calculate ΔSº (in J/K) at 298 K for :

2CO2(g) +4H2O(l) → 2CH3OH(l) + 3O2(g)

Use the data in Appendix II (Tro); Appendix 4 (Zumdahl)
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Explanation

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Question 5

 

Use ΔGº =ΔHº -TΔSº to calculate ΔG (in kJ) at 298 K for :

2CO2(g) +4H2O(l) → 2CH3OH(l) + 3O2(g)
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Explanation

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Question 6

 

Use ΔGº =ΔHº -TΔSº to calculate ΔG (in kJ) at 298 K for :

2CO2(g) +4H2O(l) → 2CH3OH(l) + 3O2(g)
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Explanation

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Question 7

 

Is the above reaction spontaneous in the forward direction at either of the temperatures? (Pick 2 answers.)
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Explanation

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Question 8

 

What is ΔGrxno (in kJ) at 1447 K for the following reaction?

2POCl3(g) → 2PCl3(g) + O2(g)

POCl3(g): ΔHfo = -592.7 kJ/mol and Sº = 324.6 J/K mol)
PCl3(g): ΔHfo = -287.0 kJ/mol and Sº = 311.7 J/K mol)
O2(g): ΔHfo = ? kJ/mol and Sº = 205.0 J/K mol)

Hint given in feedback.
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Explanation

Find ΔHº and ΔSº. ΔG =ΔHº -TΔSº.

Question 9

 

At what temperature (in K) does the above reaction become spontaneous?

Hint given in feedback.
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Explanation

Use ΔG =ΔHº -TΔSº. The temperature at which ΔG = 0 is the key.

Question 10

 

The above reaction is spontaneous in the forward direction
(Pick 2)
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Explanation

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Question 11

 

What is ΔGo (in kJ) at 491 K for the following reaction?

PbO(g) + CO2(g) → PbCO3(s)

PbO: ΔHfo = -219.0 kJ/mol and So = 66.5 J/K mol)
PbCO3(s): ΔHfo = -699.1 kJ/mol and So = 131.0 J/K mol)

CO2: ΔHfo = -393.5 kJ/mol and So = 213.6 J/K mol)

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Explanation

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Question 12

 

At what temperature (in K) does the above reaction become spontaneous?

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Explanation

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Question 13

 

The above reaction is spontaneous in the forward direction
(Pick 2)
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Explanation

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Question 14

 

What is ΔGo (in kJ) at 21 ºC for the phase change of nitromethane from the liquid to the gaseous state?

CH3NO2(l) → CH3NO2(g)

CH3NO2(l): ΔHo = -113.1 kJ/mol and So = 171.8 J/K mol)
CH3NO2(g): ΔHo = -74.7 kJ/mol and So = 274.4 J/K mol)

Hint given in feedback.

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Explanation

Find ΔH and ΔS for the phase change. ΔG =ΔH -TΔS. In calculating ΔG it does not matter what kind of process it is–reaction or phase change.

Question 15

 

Which state is more stable for nitromethane at the temperature in the previous question?
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Explanation

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Question 16

What is the boiling point of nitromethane (in ºC)?

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Explanation

 

Hint, find temperature in K and convert to ºC.goes here…

 

 

Question 17

 

At 298 K, ΔGo = – 6.36 kJ for the reaction:

2N2O(g) + 3O2(g) ↔ 2N2O4(g)

Calculate ΔG (in kJ) at 298 K when PN2O = 3.10 atm, PO2 = 0.0072 atm, and PN2O4= 0.276 atm.

Help given in feedback.

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Explanation

Hint, first find Q.

Question 18

 

Under the conditions in the previous question, in which direction is the reaction spontaneous?
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Explanation

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Question 19

 

At 298 K, ΔGo = + 8.68 kJ for the reaction:

ZnF2(s) ↔ Zn2+(aq) + 2F-(aq)

Under standard conditions what is the spontaneous direction (Recall, the o after ΔG means standard conditions, that is, some solid ZnF2, 1 molar concentrations for Zn2+(aq) and F-(aq), and 298 K if no other temperature is specified.)
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Explanation

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Question 20

 

Calculate ΔG (in kJ) at 298 K for some solid ZnF2, 0.029 M Zn2+ and 0.057 M F-(aq).
Hint given in feedback.

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Explanation

Hint, first find Q.

Question 21

 

Under the conditions of the previous question what is the spontaneous direction?
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Explanation

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Question 22

 

Endothermic reaction; decrease in entropy:
Calculate the equilibrium constant at 41 K for a reaction with ΔHo = 10 kJ and ΔSo = -100 J/K. (Don’t round unil the end. Using the exponent enlarges any round-off error.)

Hint given in feedback.

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Explanation

Hint, first calculate ΔGo at 41 K.

Question 23

 

Calculate the equilibrium constant at 145 K for the thermodynamic data in the previous question.

Notice that Keq is larger at the larger temperature for an endothermic reaction.

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Explanation

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Question 24

 

Endothermic reaction; increase in entropy
Calculate the equilibrium constant at 35 K for a reaction with ΔHo = 10 kJ and ΔSo = 100 J/K.

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Explanation

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Question 25

 

Calculate the equilibrium constant at 132 K for the thermodynamic data in the previous question.

Notice that Keq is dramatically larger for a larger temperature when there is a substantial positive increase in entropy.

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Explanation

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Question 26

 

Warm-up question Using data from the Appendix, calculate ΔGo (in kJ) at 65 oC for the reaction:

N2O(1 atm) + H2(1 atm) ↔ N2(1 atm) +H2O(l)

(Recall, all gases at 1 atm for standard conditions)
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Explanation

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Question 27

 

Calculate ΔG (in kJ) at 65 oC for the reaction:

N2O(0.0046 atm) + H2(0.28 atm) ↔ N2(352.3 atm) +H2O(l)
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Explanation

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Question 28

 

Assume that the ΔHo and ΔSo of vaporization do not change significantly with temperature. Calculate the vapor pressure of CH3OH at 61 oC (in atm).

CH3OH (l) ↔ CH3OH(g) . . . ΔHo = 38.0 kJ and ΔSo = 112.9 J/K

(Don’t round unil the end. Using the exponent enlarges any round-off error.) Hint given in feedback.
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Explanation

Hint,what is the relationship between Keq and the vapor pressure.

Question 29

 

Calculate the vapor pressure of Hg at 72 oC (in atm).

Hg(l) ↔ Hg(g) . . . ΔHo = 61.32 kJ and ΔSo = 98.83 J/K
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Explanation

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Question 30

 

Cyclopropane is a hydrocarbon that contains a ring of three carbon atoms with two hydrogen atoms bonded to each carbon (see below). (a) What are the steric number and the orbital hybridization for the carbon atoms? (b) What is the expected C—C—C bond angle? The actual bond angle is 60º. (c) Consequently, do you expect this bond to be under a lot of stress? (d) What does this imply about the expected C—C bond strength?

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Explanation

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Question 31

 

Calculate the average C—C bond strength in cyclopropane (in kJ/mol). Its combustion and the experimental enthalpy of reaction are:

C3H6(g) + 4.5O2(g) → 3CO2(g) + 3H2O(g) ΔHºrxn = -1,957.7 kJ/mol

Since all reactants and products are in the gaseous state, bond energies may be used to estimate the enthalpy of reaction. Do not use the standard C—C bond strength, leave the C—C bond strength as an unknown and solve for its value. For review see chapter 9 p. 392 (Tro) or chapter 13 p. 608 (Zumdahl).
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Explanation

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