Name __________________________________
Chemical Thermodynamics Quiz (157 points)
G = H - TS
G = G + RT ln k
R = 1.9872 cal /mol K
R = 8.31 J/mol K
1.
Multiple Choice 2pts each
Which reaction has the largest positive entropy change per mole of CO formed?
(A)
CO2(g) 
 CO(g) + 1/2O2(g)
(B)
CH3OH(g) 
 CO(g) + 2H2(g)
(C)
CO2(g) + NO(g) 
 CO(g) + NO2(g)
(D)
CO2(g) + H2(g) 
 CO(g) + H2O(l)
2.
When Al2O3(s) is formed from the elements at standard conditions, the values of ΔH0 and ΔG0 at 298 K
are
–1617 kJ·mol–1 and –1577 kJ·mol–1, respectively. The standard entropy of formation per mole,
in joules per degree, will be
(A)
–135
(B)
–157
(C)
–93.3
(D)
–0.0933
(E) +15.7
3.
Vaporization of a liquid is an example of a process for which
(A)
ΔH, ΔS, and ΔG are positive at all temperatures.
(B)
ΔH and ΔS are positive.
(C)
ΔG is negative at low temperatures, positive at high temperatures.
(D)
ΔH = ΔS
4.
A particular chemical reaction has a negative ΔH and negative ΔS. Which statement is correct?
(A)
The reaction is spontaneous at all temperatures.
(B)
The reaction is nonspontaneous at all temperatures.
(C)
The reaction becomes spontaneous as temperature increases.
(D)
The reaction becomes spontaneous as temperature decreases.
5.
The hydrolysis of ATP in the body: ATP + H2O  ADP + H2PO4-1
is a source of energy. For this reaction the K = 2.23 x 105. Which reaction quotient for this reaction
would be found in a system capable of doing the most work?
(A)
Q = 5.0 x 10 5
(B)
Q = 5.0 x 10 15
(C)
Q = 5.0 x 10-5
(D)
Q=1
6.
A spontaneous reaction will occur at any temperature when
(A)
ΔH is positive and ΔS is positive.
(B)
ΔH is negative and ΔS is negative.
(C)
ΔH is positive and ΔS is negative.
(D)
ΔH is negative and ΔS is positive.
7.
For this process at 25 °C:
H2O(g)  H2O(l)
(A)
ΔH is negative and ΔS is negative.
(B)
ΔH is negative and ΔS is positive.
(C)
ΔH is positive and ΔS is positive.
(D)
ΔH is positive and ΔS is negative.
8.
Which change is likely to be accompanied by the greatest increase in entropy?
(A)
N2(g) + 3H2(g)  2NH3(g) (at 25 °C)
(B)
Ag+(aq) + Cl–(aq)  AgCl(s) (at 25 °C)
(C)
CO2(s)  CO2(g) (at –70 °C)
(D)
H2O(g)  H2O(l) (at 100 °C)
9.
Under which conditions does nitrogen have the largest entropy per mole?
(A)
N2(s) at 50 K and l atm
(B)
N2(l) at 70 K and l atm
(C)
N2(g) at 80 K and 1 atm
(D)
N2(g) at 80 K and 0.5 atm
10.
What is the standard enthalpy of combustion of C2H6 in kJ·mol–1?
H2(g) + 1/2O2(g)  H2O(l)
C2H4(g) + H2(g)  C2H6(g)
C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(l)
(A)
–1275 kJ
(B)
–31561 kJ
(C)
–1561 kJ
(D)
+1834 kJ
Thermochemical Data
ΔH0= –286 kJ
ΔH0= –137 kJ
ΔH0= –1412 kJ
11.
More heat is derived from cooling one gram of steam at 100 °C to water at 50 °C than from cooling
one gram of liquid water at 100 °C to 50 °C because
(A)
water is a poor thermal conductor.
(B)
the steam is hotter than the water.
(C)
the steam occupies a greater volume than the water.
(D)
the density of water is greater than that of steam.
(E)
the heat of condensation is evolved.
12.
Copper is more reactive chemically than is gold. What is the best experimental evidence for this?
(A)
The copper atom may lose as many as two electrons, whereas the gold atom may lose as many as
three electrons.
(B)
Gold is alloyed with copper.
(C)
The copper ion has less tendency to lose electrons than has the gold ion.
(D)
Gold oxide requires less heat per gram mole for decomposition than does copper oxide.
(E)
Copper loses electrons more easily than does gold.
13.
For which process is the entropy change per mole the largest at constant temperature?
(A)
H2O(l)  H2O(g)
(B)
H2O(s)  H2O(l)
(C)
H2O(s)  H2O(g)
(D)
H2O(l)  H2O(s)
14.
A spontaneous reaction will occur at any temperature when
(A)
ΔH is positive and ΔS is positive.
(B)
ΔH is positive and ΔS is negative.
(C)
ΔH is negative and ΔS is negative.
(D)
ΔH is negative and ΔS is positive.
15.
Consider ice in equilibrium with liquid water at 273 K. Which of the following relationships is correct
for G(s), the free energy per mole of ice and G(l), the free energy per mole of the liquid?
(A)
G(s) is less than G(l)
(B)
G(s) is greater than G(l)
(C)
G(s) equals 0, G(l) equals 0
(D)
G(s) equals G(l); neither equals 0
16.
A certain reaction has negative values of both H and S. Therefore, the reaction
(A)
must be spontaneous at all temperatures.
(B)
cannot be spontaneous at any temperature.
(C)
will be spontaneous only at low temperatures.
(D)
will have a positive free energy at any temperature.
17) The value of ∆E for a system that performs 13 kJ of work on its surroundings and loses 9 kJ of heat is
__________ kJ.
A) 22
B) -22
C) -4
D) 4
E) -13
18) When a system __________, E is always negative.
A) absorbs heat and does work
B) gives off heat and does work
C) absorbs heat and has work done on it
D) gives off heat and has work done on it
E) none of the above is always negative.
19) The internal energy can be increased by __________.
(a) transferring heat from the surroundings to the system
(b) transferring heat from the system to the surroundings
(c) doing work on the system
A) a only
B) b only
C) c only
D) a and c
E) b and c
20) Which of the following statements is false?
A) Internal energy is a state function.
B) Enthalpy is an intensive property.
C) The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to
the enthalpy change for the reverse reaction.
D) The enthalpy change for a reaction depends on the state of the reactants and
products.
E) The enthalpy of a reaction is equal to the heat of the reaction.
21) A sample of calcium carbonate [CaCO3 (s)] absorbs 45.5 J of heat, upon which the temperature of the
sample increases from 21.1 °C to 28.5 °C. If the specific heat of calcium carbonate is 0.82 J/g-K, what is
the mass (in grams) of the sample?
A) 3.7
B) 5.0
C) 7.5
D) 410
E) 5.0  103
22) For which one of the following reactions is H°rxn equal to the heat of formation of the product?
A) N2 (g) + 3H2 (g)  2NH3 (g)
B) (1/2) N2 (g) + O2 (g)  NO2 (g)
C) 6C (s) + 6H (g)  C6H6 (l)
D) P (g) + 4H (g) + Br (g)  PH4Br (l)
E) 12C (g) + 11H2 (g) + 11O (g)  C6H22O11 (g)
23) Of the following, Hf° is not zero for __________.
A) Sc (g)
B) Si (s)
C) P4 (s, white)
D) Br2 (l)
E) Ca (s)
24) The combustion of titanium with oxygen produces titanium dioxide:
Ti (s) + O2 (g)  TiO2 (s)
When 2.060 g of titanium is combusted in a bomb calorimeter, the temperature of the calorimeter
increases from 25.00 °C to 91.60 °C. In a separate experiment, the heat capacity of the calorimeter is
measured to be 9.84 kJ/K. The heat of reaction for the combustion of a mole of Ti in this calorimeter is
__________ kJ/mol.
A) 14.3
B) 19.6
C) -311
D) -0.154
E) -1.52  104
25) When a system is at equilibrium, __________.
A) the reverse process is spontaneous but the forward process is not
B) the forward and the reverse processes are both spontaneous
C) the forward process is spontaneous but the reverse process is not
D) the process is not spontaneous in either direction
E) both forward and reverse processes have stopped
26) A reversible process is one that __________.
A) can be reversed with no net change in either system or surroundings
B) happens spontaneously
C) is spontaneous in both directions
D) must be carried out at low temperature
E) must be carried out at high temperature
27) Which of the following statements is false?
A) The change in entropy in a system depends on the initial and final states of the
system and the path taken from one state to the other.
B) Any irreversible process results in an overall increase in entropy.
C) The total entropy of the universe increases in any spontaneous process.
D) Entropy increases with the number of microstates of the system.
28) Which one of the following correctly indicates the relationship between the entropy of a system and the
number of different arrangements, W, in the system?
A) S = kW
B)
S=
C)
S=
D) S = klnW
E) S = Wk
29) Of the following, the entropy of gaseous __________ is the largest at 25°C and 1 atm.
A) H2
B) C2H6
C) C2H2
D) CH4
E) C2H4
30) The equilibrium position corresponds to which letter on the graph of G vs. f (course of reaction) below?
A)
B)
C)
D)
E)
A
B
C
D
E
31) For the reaction
2C4H10 (g) + 13O2 (g)  8CO2 (g) + 10H2O (g)
H° is -125 kJ/mol and S° is +253 J/K · mol. This reaction is __________.
A) spontaneous at all temperatures
B) spontaneous only at high temperature
C) spontaneous only at low temperature
D) nonspontaneous at all temperatures
E) unable to determine without more information
32) With thermodynamics, one cannot determine __________.
A) the speed of a reaction
B) the direction of a spontaneous reaction
C) the extent of a reaction
D) the value of the equilibrium constant
E) the temperature at which a reaction will be spontaneous
33) If
A)
B)
C)
D)
E)
G° for a reaction is greater than zero, then __________.
K=0
K=1
K>1
K<1
More information is needed.
34) Given the thermodynamic data in the table below, calculate the equilibrium constant (at 298 K) for the
reaction:
2SO2 (g) + O2 (g)
2SO3 (g)
A)
B)
C)
D)
E)
2.37  1024
1.06
1.95
3.82  1023
More data are needed.
35) The value of G° at 100.0°C for the formation of calcium chloride from its constituent elements:
Ca (s) + Cl2 (g) ? CaCl2 (s)
is __________ kJ/mol. At 25.0°C for this reaction, H° is -795.8 kJ/mol, G° is -748.1 kJ/mol, and S°
is -159.8 J/K.
A) -855.4
B) -736.1
C) 5.88  104
D) -779.8
E) 1.52  104
Problem A
2 NO(g) + O2(g)  2 NO2(g)
H°= -114.2 kJ, S°= -146.5 J K-1
The reaction represented above is one that contributes significantly to the formation of photochemical smog.
(a) Calculate the quantity of heat released when 72.1 g of NO(g) is converted to NO2(g).
(b) For the reaction at 25C, the value of the standard free-energy change, G, is -70.4 kJ.
(i) Calculate the value of the equilibrium constant, Keq, for the reaction at 25C.
(ii) Indicate whether the value of G would become more negative, less negative, or remain unchanged
as the temperature is increased. Justify your answer.
(c) Use the data in the table below to calculate the value of the standard molar entropy, S, for O2(g) at 25C.
Standard Molar Entropy,
S (J K-1 mol-1)
NO(g)
210.8
NO2(g)
240.1
Problem B
WO3(s) + 3 H2(g)  W(s) + 3 H2O(g)
Tungsten is obtained commercially by the reduction of WO3 with hydrogen according to the equation above.
The following data related to this reaction are available:
WO3(s) H2O(g)
Hf (kilocalories/mole) –200.85 –57.8
Gf (kilocalories/mole) –182.48 –54.6
(a)
What is the value of the equilibrium constant for the system represented above?
(b)
Calculate S at 25C for the reaction indicated by the equation above.
(c)
Find the temperature at which the reaction mixture is in equilibrium at when all components have a
pressure of 1 atmosphere.
Problem C
(a) When liquid water is introduced into an evacuated vessel at 25C, some of the water vaporizes. Predict how
the enthalpy, entropy, free energy, and temperature change in the system during this process. Explain the
basis for each of your predictions.
(b) When a large amount of ammonium chloride is added to water at 25C, some of it dissolves and the
temperature of the system decreases. Predict how the enthalpy, entropy, and free energy change in the
system during this process. Explain the basis for each of your predictions.
(c) If the temperature of the aqueous ammonium chloride system in part (b) were to be increased to 30C,
predict how the solubility of the ammonium chloride would be affected. Explain the basis for each of your
predictions.
Problem D
An experiment is to be performed to determine the standard molar enthalpy of neutralization of a strong acid by
a strong base. Standard school laboratory equipment and a supply of standardized 1.00 molar HCl and
standardized 1.00 molar NaOH are available.
(a) What equipment would be needed?
(b) What measurements should be taken?
(c) Without performing calculations, describe how the resulting data should be used to obtain the standard
molar enthalpy of neutralization.
(d) When a class of students performed this experiment, the average of the results was -55.0 kilojoules per
mole. The accepted value for the standard molar enthalpy of neutralization of a strong acid by a strong base
is -57.7 kilojoules per mole. Propose two likely sources of experimental error that could account for the
result obtained by the class.
Problem E
Ehtylene oxide, C2H4O, is a hydrocarbon commonly used in the production of plastics.
(a) Write the balanced equation for the complete combustion of ethylene oxide, which yields CO2(g) and
H2O(l).
(b) Calculate the volume of air at 30C and 1.00 atmosphere that is needed to burn completely 10.0 grams
of ethylene oxide. Assume that air is 21.0 percent O2 by volume. (PV = nRT , R = 0.0821 L
atm/(molK) , n = moles)
(c) The heat of combustion of ethylene oxide is –1300.5 kJ for 1.00 mol C2H4O. Calculate the heat of
formation, Hf, of ethylene oxide given that Hf =of H2O (l) = - 283.3 kJ/mol and Hf of CO2(g) = 393.5 kJ/mol.
(d) Assume that all of the heat evolved in burning 30.0 grams of ethylene oxide is transferred to 8.00
kilograms of water (specific heat = 4.18 J/g K) at 25.0 C, calculate the final temperature of the water.
Short Answer (5 pts each)
1) What is the normal boiling point for formic acid, HCOOH?
Hf (kJ/mol)
S (J/mol K)
HCOOH (l)
-410
130
HCOOH (g)
-363
251
2) For a particular reaction the equilibrium constant is 1.50 x 10 –2 at 370 C. H is + 16.0 kJ. What is the
S for the reaction?
3) Using the relationship ln(K) = -H/RT + S/R Show that for a system at equilibrium, the equilibrium
will shift to the right for an endothermic process when the temperature is increased.
4) Consider the following reaction at 800K. N2(g) + 3F2(g)  2 NF3(g) An equilibrium mixture contains the
following partial pressures: PN2 = 0.021 atm, PF2 = 0.063 atm, PNF3 = 0.48 atm. Calculate the G at 800 K
5) Given the following reactions and H values,
P4(s) + 6Cl2(g)  4 PCl3(g)
Ha = -1225.6 kJ
P4(s) + 5O2(g)  P4O10(s)
Hb = -2967.3 kJ
PCl3(g) + Cl2(g)  PCl5(g)
Hc = - 84.2 kJ
PCl3(g) + ½ O2(g)  Cl3PO(g)
Hd = - 285.7 kJ
Calculate H for:
P4O10(s) + 6PCl5(g)  10 Cl3PO(g)
6) What is the final temperature when 4.00 g of 373 K aluminum is added to 9.05 g of water that is 293 K?
The specific heat of aluminum is 0.891 J/g K, water is 4.184 J/gK
7) a. Why do we generally expect H and E to be nearly the same for all reactions in which the reactants and
products are in an aqueous solution?
b. Consider two aqueous double replacement reactions, one producing a weak electrolyte and the other an
insoluble gas (comes out of solution). For which of these reactions would the difference in values of E and
H be likely to be greater and why?
c. The energy change in a system that occurs at constant pressure is _________, while the energy change
in a system that occurs at constant volume is _______. ( Answer bank = H, E)
8) The enthalpy change for the conversion of white phosphorous to red phosphorus is –17.6 kj/mol. Will one
mole of red phosphorous or one mole of white phosphorus give off more heat when burned?