Unit 9 Summative Assessment Practice
Show your work for each question in the space provided. Examples and equations may be included in
your responses where appropriate. For calculations, clearly show the method used and the steps involved
in arriving at your answers. You must show your work to receive credit for your answer. Pay attention to
significant figures.
CO(g) + 2 H2(g) → CH3OH(g)
1.
ΔH° = –90.2 kJ/molrxn
Answer the following questions based on the reaction represented by the equation shown above.
The values of the standard molar entropies of the substances involved in the reaction are given in
the following table.
Substance
S° (J/(K·mol))
CO(g)
197.7
H2(g)
130.6
CH3OH(g)
239.9
(a) Use the data in the table to calculate the value of the standard entropy change, ΔS°, in units of
J/(K·molrxn) for the reaction. Show your calculations in the space below.
(b) Circle one of the following choices that best describes the thermodynamic favorability of this
reaction.
Not favored
at any T
Favored
at all T
Favored
at high T
Favored
at low T
(c) Calculate the value of the standard free energy change, ΔG°, in units of kJ/molrxn, for the
reaction at 298 K. Show your calculations in the space below.
(d) Is the reaction thermodynamically favorable at a temperature of 298 K? Justify your answer.
2.
Chemical Equation for the Reaction
ΔH° (kJ/molrxn)
ΔS° (J/ K·molrxn)
CO2(g) + 2 NH3(g) → CO(NH2)2(s) + H2O(l)
–133.3
–424.6
Answer the following questions related to the information shown above.
(a) Use particle-level reasoning to explain why the sign of the standard entropy change, ΔS°, is
negative for this reaction.
(b) This reaction is thermodynamically favorable at 298 K. Which of the following represents the
driving force for this reaction? (Circle one of the following choices.)
ΔH° only
ΔS° only
both ΔH° and ΔS°
(c) Justify your choice in part (b) in terms of the standard free energy change, ΔG°, and the signs
of ΔH° and ΔS° for this reaction.
Although this reaction is thermodynamically favorable at 298 K, it is observed that the reaction
occurs at a very slow rate.
(d) A common explanation for the slow reaction rate is the following. The value of
(
Ea
ΔG°
) for this reaction has a relatively (
low
high ) magnitude.
When a suitable catalyst is added to the reaction mixture at 298 K, it is observed that the reaction
rate increases.
(e) If a suitable catalyst is added to the reaction mixture, the magnitude of the activation
energy (Ea) for the reaction will (
decrease
increase
remain the same
).
(f) If a suitable catalyst is added to the reaction mixture, the value of the standard free energy
change (ΔG°) for the reaction will (
decrease
increase
remain the same
).
N2O4(g) → 2 NO2(g)
3.
Answer the following questions related to the reaction represented by the equation shown above.
The value of the standard free energies of formation of the substances involved in the reaction are
given in the following table.
Substance
G of (kJ/mol)
N2O4(g)
99.8
NO2(g)
51.3
(a) Use the data in the table to calculate the value of the standard free energy change, ΔG°, in units
of kJ/molrxn for the reaction. Show your calculations in the space below.
(b) Is the reaction thermodynamically favorable under standard conditions? Justify your answer.
(c) Calculate the value of the equilibrium constant Kp under standard conditions at a temperature of
298 K. Show your calculations in the space below.
4.
Temperature
Kp
300 K
0.0024
500 K
410
The value of the equilibrium constant Kp for a certain chemical reaction varies with temperature, as
shown in the table above. Use the information in the table to answer the following questions about
this chemical reaction.
(a) Circle the correct choices for the substances that are favored at equilibrium in this reaction at
temperatures of 300 K and 500 K. Assume that ΔH° and ΔS° are independent of temperature.
Temperature
Are Reactants or Products Favored at Equilibrium?
300 K
reactants
products
500 K
reactants
products
(b) Identify the sign of ΔG° for this reaction at 300 K and at 500 K. Assume that ΔH° and ΔS° are
independent of temperature.
Sign of ΔG° for the Reaction at 300 K
negative
positive
Sign of ΔG° for the Reaction at 500 K
negative
positive
(c) Identify the signs of ΔH° and ΔS° for this reaction. Assume that ΔH° and ΔS° are independent
of temperature.
ΔH°
negative
positive
ΔS°
negative
positive
(d) Justify your choices in part (c) in terms of the Gibbs free energy equation (ΔG° = ΔH° – TΔS°)
and your answer to part (b).
(e) Circle one of the following choices that best describes the thermodynamic favorability of this
reaction.
Not favored
at any T
Favored
at all T
Favored
at high T
Favored
at low T
5.
Aluminum hydroxide, Al(OH)3(s), is produced according to the equation below.
Overall Reaction:
4 Al(s) + 3 O2(g) + 6 H2O(l) → 4 Al(OH)3(s)
Two half-reactions and standard reduction potentials are given in the following table.
Half-Reaction
Balanced Chemical Equation
Al(OH)3(s) + 3 e– → Al(s) + 3 OH–(aq)
1
2
O2(g) + 2 H2O(l) + 4 e– → 4 OH–(aq)
E° (V)
–2.31
+0.40
(a) In the table below, write the modified versions of half-reactions 1 and 2 that will produce the
overall reaction when they are added together.
Modified Version
of Half-Reaction 1
Modified Version
of Half-Reaction 2
Overall Reaction
4 Al(s) + 3 O2(g) + 6 H2O(l) → 4 Al(OH)3(s)
(b) Calculate the standard cell potential, E°, in units of volts, for the overall reaction. Show your
calculations in the space below.
(c) Based on your answer to part (b), calculate the value of the standard free energy change, ΔG°,
in units of kJ/molrxn, for the overall reaction. Show your calculations in the space below.
5.
(continued)
Overall Reaction:
4 Al(s) + 3 O2(g) + 6 H2O(l) → 4 Al(OH)3(s)
ΔH° = –3393 kJ/molrxn
(d) The standard enthalpy change, ΔH°, for the overall reaction is –3393 kJ/molrxn. Based on this
information and your answer to part (c), calculate the standard entropy change, ΔS°, in units
of J/(K·molrxn), for the overall reaction at 298 K. Show your calculations in the space below.
____________________________________________________________________________________
Reaction
1
6.
ΔG°
(kJ/molrxn)
Chemical Equation
CaCO3(s) → CaO(s) + CO2(g)
+131.1
2
C(s) + O2(g) → CO2(g)
–394.4
3
CaCO3(s) + C(s) + O2(g) → CaO(s) + 2 CO2(g)
?
Answer the following questions related to the information shown above.
(a) Use the information in the table to calculate the value of the standard free energy change, ΔG°,
for reaction 3 in units of kJ/molrxn. Show your calculations in the space below.
(b) Reaction 1 is (
not favorable
favorable
) under standard conditions.
Reaction 2 is (
not favorable
favorable
) under standard conditions.
Reaction 3 is (
not favorable
favorable
) under standard conditions.
7.
Half-Reaction
E° (V)
Al3+(aq) + 3 e– → Al(s)
–1.66
Ni2+(aq) + 2 e– → Ni(s)
–0.26
A student sets up a standard galvanic cell, with an Al(s) electrode immersed in 1.0 M Al(NO3)3(aq)
and a Ni(s) electrode immersed in 1.0 M Ni(NO3)2(aq). Assume that the temperature is 25°C.
Two half-reactions and standard reduction potentials are given in the table above.
(a) Write the balanced net ionic equation for the reaction that occurs as this galvanic cell operates.
(b) Calculate the value of the standard cell potential, E°, in units of volts, for the chemical reaction
that occurs as this cell begins to operate. Show your calculations in the space below.
(c) Identify the anode and the cathode in this galvanic cell.
anode
Al(s)
Ni(s)
cathode
(d) On the diagram above, draw an arrow to indicate the direction
of electron flow in the wire as the galvanic cell operates.
(e) A solution of potassium nitrate,
KNO3(aq), is used in the salt bridge.
The particle diagram shown at right
represents an expanded view of the
center portion of the salt bridge.
Draw arrows in this diagram to indicate
the direction of movement for the K+(aq)
ions and the NO3–(aq) ions in the salt bridge
as the cell begins to operate.
Al(s)
Ni(s)
7.
(continued)
(f) As this galvanic cell operates over time, will the mass of the Al(s) electrode decrease, increase,
or remain the same? Justify your answer.
A second galvanic cell is set up with identical conditions to the first galvanic cell, except for a
change in the concentration of Al3+(aq) in the half-cell container as shown below.
Galvanic Cell #1
Galvanic Cell #2
[Al3+]
[Ni2+]
[Al3+]
[Ni2+]
1.0 M
1.0 M
0.010 M
1.0 M
(g) Calculate the value of the reaction quotient, Q, for galvanic cell #2 at the moment that the cell
begins to operate. Show your calculations in the space below.
(h) Do you predict that the cell potential for galvanic cell #2 will be less than, greater than, or
equal to the value of the cell potential calculated in part (b) for galvanic cell #1?
Justify your answer.
8.
An external direct-current power supply is connected to two graphite electrodes immersed in a
container of molten PbBr2, as shown in the diagram above. As the cell operates, Br 2(g) is produced
at one electrode and Pb(l) is formed at the other electrode. Two half-reactions and standard
reduction potentials are given in the following table.
Half-Reaction
E° (V)
Br2 + 2 e– → 2 Br–
1.07
Pb2+ + 2 e– → Pb
–0.13
(a) Write a balanced net ionic equation for the electrolysis reaction that occurs in the cell.
(b) Write the net ionic equation for the half-reaction that occurs at each electrode in this
electrolytic cell.
Half-Reaction that Occurs at the Anode
Half-Reaction that Occurs at the Cathode
(c) Calculate the value of the standard cell potential (E°), in units of volts, for the electrolysis
reaction that occurs in this cell. Show your calculations in the space below.
8.
(continued)
(d) Based on your answer to part (c), calculate the value of the standard free energy change, ΔG°,
in units of kJ/molrxn, for the electrolysis reaction that occurs in this cell. Show your calculations
in the space below.
(e) An electric current of 3.00 amperes passes through the sample of molten PbBr 2 for a period of
7.50 minutes. Calculate the mass, in grams, of Pb that is produced during this time period.