Camphor
Sample Examination # 1A
Chemistry 1412
1
CHEM 1412 Exam #1 Sample Exam # 1A
(chapters 12,13, and 14)
Name:
Score:
PART I - ( 3 points each) - Please write your correct answer next to each question
number. DO NOT CIRCLE
____1. In which colligative property(ies) does the value decrease as more solute is added?
a. boiling point
c. vapor pressure
b. freezing point
d. freezing point and vapor pressure
____2. What is the molarity of a solution prepared by dissolving 25.2 g CaCO3 in 600. mL
of a water solution?
a. 0.420 M
b. 0.567 M
c. 0.042 M
d. 0.325 M
____3. The solubility of nitrogen gas in water at a nitrogen pressure of 1.0 atm is
6.9 x 10-4 M. What is the solubility of nitrogen in water at a nitrogen pressure of
0.80 atm?
a. 5.5 x 10-4
b. 8.6 x 10-4
c. 3.7 x 10-3
d. 1.2 x 103
____4. What is the freezing point of an aqueous glucose solution that has 25.0 g of glucose,
C6H12O6, per 100.0 g H2O (Kf = 1.86°C /m)?
a. 0.258
b. -0.258
c. 2.58
d. -2.58
____5. What is the osmotic pressure in atm produced by a 1.20 M glucose (C6H12O6)
solution at 25°C?
a. 29.3
b. 4.89
c. 25.1
d. 36.0
____6. The vapor pressure of pure ethanol at 60°C is 349 mm Hg. Calculate vapor pressure
in mm Hg at 60°C for a solution prepared by dissolving 10.0 mol naphthalene
(nonvolatile) in 90.0 mol ethanol?
a. 600
b. 314
c. 34.9
d. 69.8
____7. Which statement is not correct regarding the function of a catalyst?
a. it lowers the activation energy
c. it affects the equilibrium constant
b. it affects the rate of a chemical reaction
d. it changes the rate constant of a reaction
____8. For first-order reactions the rate constant, k, has the units
a. M s-1
b. M-1 s-1
c. M-2 s-1
2
d. s-1
____ 9. For second-order reactions the slope of a plot of 1/[A] versus time is
a. k
b. k/[A]0
c. kt
d. -k
____10. If the reaction 2A + 3D  products is first-order in A and second- order in D,
then the rate law will have the form rate =
a. k[A]2[D]3
c. k[A]2[D]2
b. k[A][D]
d. k[A][D]2
____11. In the first-order reaction A  products, the initial concentration of A is 1.56 M and the
concentration is 0.869 M after 48.0 min. What is the value of the rate constant, k, in min-1?
a. 3.84 x 10-2
b. 2.92 x 10-2
c. 5.68 x 10-2
d. 1.22 x 10-2
____12. The following time and concentration data was obtained for the reaction; 2A  products
time (min)
0
1.1
2.3
4.0
[A], M
1.20
1.00
0.80
0.60
Refer to the table above. If the reaction is known to be first-order, determine the rate constant for the
reaction.
a. 0.17
b. 0.37
c. 0.49
d. 0.60
____13. Consider the reaction 2HI(g)  H2(g) + I2(g). What is the value of the equilibrium constant, Kc, if at
equilibrium , [H2] = 6.50 x 10-7 M, [I2] = 1.06 x 10-5 M, and [HI] = 1.87 x 10-5 M?
a. 3.68 x 10-7
b. 1.97 x 10-2
c. 1.29 x 10-16
d. 50.8
____ 14. For the elementary reaction NO3 + CO  NO2 + CO2
a.
b.
c.
d.
the molecularity is 2 and rate = k[NO3][CO]/[NO2][CO2]
the molecularity is 4 and rate = k[NO3][CO]/[NO3][CO]
the molecularity is 4 and rate = k[NO3][CO][NO2][CO2]
the molecularity is 2 and rate = k[NO3][CO]
____15. Given the following mechanism, determine which of the species below is a catalyst?
I) C + ClO2  ClO + CO
II) CO + ClO2  CO2 + ClO
III) ClO + O2  ClO2 + O
IV) ClO + O  ClO2
a.
ClO2
b.
CO2
c.
3
O
d. CO
____16. For the system CaO(s) + CO2(g)  CaCO3(s) the equilibrium constant expression is
a. [CO2]
b. 1 / [CO2]
c. [CaO] [CO2] / [CaCO3]
d. [CaCO3] / [CaO] [CO2]
____17. The value of Kp for the reaction 2NO2(g)  N2O4(g) is 1.52 at 319 K. What is the value of Kp at this
temperature for the reaction
N2O4(g)  2NO2(g)?
a. -1.52
c. 5.74 x 10-4
b. 1.23
d. 0.658
____18. The value of Kc for the reaction C(s) + CO2(g)  2CO(g) is 1.6. What is the equilibrium
concentration of CO if the equilibrium concentration of CO2 is 0.50 M?
a. 0.31
b. 0.80
c. 0.89
d.
0.75
____19. Consider the reaction below:
2SO3(g)  2SO2(g) + O2(g) , ∆H° = +198 kJ
All of the following changes would shift the equilibrium to the left except one. Which one would not
cause the equilibrium to shift to the left?
a removing some SO3
b. decreasing the temperature
c. increasing the container volume
d. adding some SO2
____ 20. For which of the following reactions is Kc equal to Kp?
a. N2O4(g)  2NO2(g)
c. H2(g) + Cl2(g)  2HCl(g)
b. 2SO3(g)  2SO2(g) + O2(g)
d. C(s) + CO2(g)  2CO(g)
PART II- ( 8 points each) Please show all your work .
21. What is the boiling point (in °C) of a solution prepared by dissolving 11.5 g of Ca(NO3)2
(formula weight = 164 g/mol) in 150 g of water? (Kb for water is 0.52°C/m)
22. A solution is prepared by dissolving 6.00 g of an unknown nonelectrolyte in enough water to make 1.00 L
of solution. The osmotic pressure of this solution is 0.750 atm at 25.0°C. What is the molecular weight of
the unknown solute (R = 0.0821 L·atm/K·mol)?
23. The rate constant for a particular reaction is 2.7 x 10-2 s-1 at 25°C and 6.2 x 10-2 s-1 at 75°C. What is the
activation energy for the reaction in kJ/mol? ( R = 8.314 J/mol.K)
4
24. Initial rate data were obtained for the following reaction: A(g) + 2B(g)  C(g) + D(g)
Experiment
1
2
3
initial
initial
[A], mol/L [B], mol/L
0.15
0.30
0.15
initial
rate
0.10
0.10
0.20
0.45
1.8
0.9
What are the rate law and k value for the reaction?
25. A mixture of 0.100 mol of NO, 0.0500 mol of H2, and 0.100 mol of H2O is placed in a 1.00-L vessel. The
following equilibrium is established:
2NO(g) + 2H2 (g)  N2(g) + 2H2O(g)
At equilibrium [NO] = 0.0620 M. Calculate the equilibrium concentrations of H2, N2, and H2O.
BONUS QUESTION - (10 points)
Using the following experimental data, determine; a) the rate law expression
b) the rate constant
c) the initial rate of this reaction when [A] = 0.60 M, [B] = 0.30 M, and [C] = 0.10 M
2 A + B2 + C  A2B + BC
Trial
Initial [A],M
Initial[B2],M
Initial[C], M
1
0.20
0.20
0.20
2.4x10-6
2
0.40
0.20
0.20
9.6x10-6
3
0.20
0.30
0.20
2.4x10-6
4
0.20
0.20
0.40
4.8x10-6
5
Initial rate M/s
1412 EX#1A Sample(key)
PART I
1. D
25.2/100 mol
2. A M = n/VL =
= 0.420 mol/L
600/1000 L
3. A S1/S2 = P1/P2  S2 = (0.6)(6.9x10-4) /1 = 5.52x10-4 M
4. D
Δ Tf = Kf. M.i = (1.86)[( 25.0/18) /0.100](1) = 2.58 0C
Tf = nTf - Δ Tf = 0.00 –2.58 = -2.58 0C
5. A
6. B
7. C
8. D
9. A
10. D
11. D
12. A
π = M.R.T.i = (1.2)(0.0821)(298)(1) = 29.35 atm
PA = XA . P0A = [ 90/(10+90)](349) = 314 mmHg
K = M (1-n) .s-1 = M(1-1) s -1 = s -1
ln[A]t = -kt + ln[A]0  ln[0.869] = -k(48) + ln[1.56]  k = 1.22x10-2 s-1
[H2] [I2]
(6.50x10-7)(1.06x10-5)
Kc = ------------- = ------------------------------- = 1.97x10-2
[HI]
(1.87x10-5)2
13 B
14. D
18. C
19. C
15. A
16. B
17. D
for reverse reaction n = - 1  K′ = (Kc)n = (Kc)-1 = ( 1.52)-1 = (1/1.52) = 0.658
20. C
PART II
21. Δ Tb = Kb. M.i = (0.52)[(11.5/164)/(0.150)] (3) = 0.73 0C,
22.
Tb = nTb + ΔTb = 100.00 + 0.73 = 100.73 0C
π = M.R.T.i  M = (0.750) / (0.0821)(298)(1)  M = 0.031 mol/L
M = n/VL  n = M.VL = ( 0.031 mol/L)(1.00L) = 0.031 mol
MW = grams/moles = 6.00 g/0.031 mol = 196 g/mol
OR
g.R.T.i
(6.00)(0.0821)(298)(1)
MW = ------------ = ---------------------------- = 196 g/mol
π . VL
(0.750)(1.00)
23. ln(k1/k2) = -(Ea/R)(1/T1 –1/T2)  ln (2.7x10-2/6.2x10-2) = -(Ea/8.314)(1/298 – 1/348)
6
 -0.8313 = -Ea (0.000058)  Ea = 14332.716 J/mol (divided by 1000)  Ea = 14 kJ/mol
24. A is second- order and B is first-order  rate = k[A]2[B]
rate
( 0.45 M s-1)
k = ------------ = ------------------------- = 200 M –2 s-1
[A]2[B]
(0.15 M)2 (0.10M)
[NO] = 0.100 mol/1L = 0.100 M , [H2O] = 0.100 mol/1L = 0.100 M , [H2] = 0.050 mol/1L = 0.050 M
2 NO + 2 H2
 N2
(0.100-2x) (0.050-2x)
+x
+
2 H2O
(0.100+2x)
0.100 –2x = 0.0620  2x = -0.062 – 0.100 = 0.038  x = 0.019 M
[H2] = 0.050 – 2x = 0.050 – 2(0.019) = 0.012 M , [H2O] = 0.100 + 2x = 0.100 + 2(0.0190 = 0.138 M
[N2] = x = 0.019 M
Bonus
rate = k[A]x [B]y [C]z
[B]y1
R1/R3 = ---------  (2.4x10-6/ 2.4x10-6) = 1 = (0.20/0.30)y  y = 0 , B is zero-order
[B]y3
[A]1x
R1/R2 = ---------  (2.4x10-6/9.6x10-6) = ( 0.20/0.40)x  (0.25) = (0.5)x  x =2 , A is second-order
[A]2x
[C]1z
R1/R4 = ---------  (2.4x10-6/4.8x10-6) = (0.20/0.40)z  (0.50) = (0.50)z  z =1 , C is first-order
[C]4z
(2.4x10-6 M s-1)
rate = k[A]2[C] , k = -------------------------  k = 3.0x10-4 M-2 s-1
(0.20 M)2 (0.20M )
rate = k[A]2[C] = (3.0x10-4 M-2 s-1) (0.60 M)2 (0.10 M)  rate = 1.08x10-5 M s-1
7
Sample Examination # 1B
Chemistry 1412
8
CHEM 1412 Exam #1 Sample Exam # 1B
(chapters 12,13, and 14)
Name:
Score:
PART I - ( 3 points each) - Please write your correct answer next to each question
number. DO NOT CIRCLE
_____ 1. The number of moles per of solute per one liter of solvent is called __________ .
a) molarity
b) molality
c) normality
d) none of these
____ 2. How many grams NaOH (40.0 g/mol) are required to make 250 mL of a 0.500 M
solution?
a) 5
b) 5000
c) 0.125
d) 125
_____3. If 200 mL of 1.60 M NaOH are diluted with water to a volume of 350 mL, the new
concentration of the solution is ……..
a) 0.257 M
b) 0.914 M
c) 2.8 M
d) 0.582 M
____ 4. Of the following salts, the one that is LEAST soluble in water is ……….
a) MgCl2
b) FeCl2
c) AgCl
d) CaCl2
____ 5. Which one of the following 0.15 m aqueous solutions lowers freezing point the most?
a) NaCl
b) C6H12O6
c) K2SO4
d) NaNO3
____ 6. What is the normality of a 2.0 M solution of phosphoric acid ?
a) 2
b) 3
c) 6
d) 0.67
____ 7. A catalyst
a) increases the yield of product
b) increases the energy of activation
c) decreases the enthalpy of the reaction
d) provides a new pathway which requires lower activation energy
_____8. The unit for a first order rate constant is ………
a) M s-1
____ 9.
b) M-1 s-1
c) s-1
d) M-2 s-1
If the rate of a reaction is second order with respect to component A, how will the rate
change if the concentration of A tripled?
a) It will double
c)It will be six times as great
b) It will tripled
d) It will be nine times as great
9
_____10. For the chemical reaction A + B  C , a plot of ln[A]t vs time is found to give a straight
line with a negative slope. What is the order of the reaction?
a) zero
b) first
c) second
d) third
_____11. The rate constant for the first order decomposition of C4H8 at 500 oC is
9.2x10-3 s-1. How long will it take for 10.0% of 0.100 M sample of C4H8
to decompose at 500 oC?
a) 12 sec
b) 0.0084 sec
c) 512 sec
d) none of these
_____ 12. For the hypothetical reaction A + 3 B  2C, the rate of appearance of C,  [C]/  t
may also be expressed as
a)  [C]/  t = - [A] /t
c)  [C]/  t = -2/3 [B]/ t
b)  [C]/  t = - 3/2 [B] /t
d)  [C]/  t = -1/2 [A]/ t
_____ 13. Consider the reaction 2HI  H2 + I2
time, sec:
[HI], M :
20
0.531
21
0.529
22
0.527
What is the rate of reaction of HI between the interval 21 sec and 20 sec?
a) 0.531 M/s
b) 0.002 M/s
_____ 14. For the reaction CaCO3(s)
c) 0.529 M/s
d) 0.527 M/s
CaO(s) + CO2 (g) , increasing the pressure on the
system at equilibrium causes
a) increased amount of CaCO3 and CaO
b) decreased the amount of CaO and CO2
c) increased the amount of CO2 and CaCO3
d) increased the amount of CaCO3 and CO2
_____ 15. Equilibrium is reached in all reversible reaction when the
a)
b)
c)
d)
forward reaction stops
reversed reaction stops
concentrations of reactants and the products become equal
rates of the opposing reactions become equal
_____ 16. The value of Kc for the following reaction is 1.60. C(s) + CO2(g)  2 CO(g)
What is the equilibrium concentration of CO if the equilibrium concentration of
CO2 is 0.50 M?
a) 0.79
b) 0.40
c) 0.894
10
d) 2.24
_____ 17. Phosgene, COCl2 , a poisonous gas decomposes according to the following equation;
COCl2 (g)  CO(g) + Cl2 (g)
If Kc = 0.083 at 900
what is the value of Kp?
oC,
a) 0.125
b) 8.0
c) 6.1
d) 0.16
_____ 18. Consider the two gaseous equilibria;
SO2(g) + ½ O2 (g) SO3 (g)
2SO3 (g)  2 SO2 (g) + O2 (g)
, K1
, K2
The value of the equilibrium constant s are related by
a) K2 = K1
b) K2 = (K1) -1
c) K2 = (K1) -2
d) K2 = (K1)2
_____ 19. Consider the reaction N2g) + O2(g)
2 NO(g) , Kc = 0.10 at 200 oC.
Starting with initial concentration of 0.04 mol/L of N2 and 0.040 mol/L of O2, calculate
the equilibrium concentration on NO in mol/L.
a) 5.4x10-3
b) 0.0096
c) 0.013
d) 1.6x10-4
_____ 20. At 700 K the reaction, 2 SO2 (g) + O2 (g)  2SO3 (g) , Kc = 4.3x106
has an equilibrium concentration of [SO2]o = 0.10M , [SO3]o = 1.0 M , [O2]o = 0.10 M.
a) The reaction mixture is at equilibrium
c) The reaction direction is reversed
b) The reaction direction is forward
d) none of the above is correct
PART II- ( 8 points each) Please show all your work .
21. At 35 oC and 75 oC the second order rate constants of a reaction are 2.50 x 10-5 and
3.26 x 10- 3 M-1 s-1 respectively. What is the enthalpy of activation (kJ/mole)?
(R = 8.314 J/mole K).
22. The equilibrium constant for the following reaction is 10.5 at 500 K. A system at equilibrium
has [CO] = 0.250 M and [H2] = 0.120 M. What is the concentration of [CH3OH]?
CO (g) + 2 H2 (g)  CH3OH (g)
11
23. HCl(g) initially at a partial pressure of 0.445 atm; is reacting with I2(s) ;
2 HCl(g) + I2 (s)  2 HI (g) + Cl2 (g) , Keq = 3.9x10 -33 , at 25 0C .
Calculate the final partial pressures at equilibrium.
24. A 1.450 g sample of an unknown organic compound , X, is dissolved in 15.0 g of toluene
( C7H8 = 92 g/mol) and the freezing point is lowered by 1.33 oC. What is the molecular weight
of the organic compound? (Kf = 5.12 oC/m).
25. Calculate the osmotic pressure in torr of a solution of 0.050 g of an unknown organic
compound in 10.0 mL of water at 25 oC. Molecular weight of the unknown organic compound
is 195 g/mol ? ( R = 0.0821 L.atm/ mol.K) ( 1 atm = 760 torr)
BONUS QUESTION - (10 points)
The following rate data were obtained at 25oC for the indicated reaction. 2 A + B
Exp.
[A] mol/L
[B] mol/L
1
2
3
0.10
0.10
0.20
0.10
0.20
0.20
Rate of formation of C
(M/min)
2.0 x 10-4
8.0 x 10-4
1.6 x 10-3
Calculate;
a) the rate law expression
b) the rate constant
c) the initial rate of this reaction when [A] = 0.40 M, [B] = 0.30 M
12
4 C?
1412 EX#1B Sample(key)
PART I
1. D
2. A
3.B
4. C
5. C
6. C
7. D
8. C
9. D
10. C
11. A
12. C
13. B
14. B
15. D
16. C
17. B
moles solute/liters solvent is not defined
( 0.5)(0.250)(40) = 5 g/mol
(1.60)(200)/350 = 0.914 M
AgCl is insoluble in water
K2SO4 has three ions 2K+ and SO42- and lowers the freezing point more.
N = nM = ( 3H)(2) = 6 Eq/L
Catalyst lowers the activation energy.
rate will increase by (3)2 = 9 times.
ln(0.09) = -9.2x10-3 t + ln (0.1)  solve for t = 12 sec
 [C]/  t = -1/3 [B]/ t = - ½ [A]/ t   [C]/  t = -2/3 [B]/ t
rate = -(0.29 – 0.531) M / (21 –20)sec = 0.002 M/s
reaction will shift to left, decreases the amounts of CO2 and CaO
18. C
19. C
20. B
Kc = [CO]2 / [CO2]  1.60 = [CO]2 / 0.50  [CO]2 = 0.80  [CO] = 0.894 M
Δ n = 2-1 = 1 , T = 900 + 273 = 1173 K
Kp = Kc (RT) Δ n = (0.083)(0.0821x 1173)1 = 8.0
Reverse and double K2 = ( K1)-2
Kc = [NO]2 / [N2][O2]  0.1 = [NO]2/ (0.04)(0.040)  solve for [NO] = 0.013 mol/L
Q = [SO3]2 / [SO2]2 [O2]  Q = (0.1)2 / (0.1)2(0.10  Q= 1000
Kc = 4.3x106 > Q = 1000 direction ( ) to forward, right- more products formed.
PART II
21.
-( 8.314) ln ( 2.50x10-5 / 3.26x10-3)
ln(k1/k2) = (-Ea/R)(1/T1 – 1/T2)  Ea =
( 1/ 308 – 1/ 348)
1
1
4.05x10
4.05x10
=
=
= 1.08x105 J/mol
= 109 kJ/mol
( 0.0032468-0.0028736)
0.000375
(divided by 1000)
22. .
Keq =
[CH3OH]
[CH3OH]
 10.5 =
[CO][
H2]2
 [CH3OH] = 0.0373 M
(
0.250)(0.120)2
13
23.
PCl2 . P2HI
Keq =
 3.9x10-33 =
P2HCl
(x) . ( 2x)2
=
2
(0.445 – 2x)
4 x3
= 3.9x10-33
(0.445)2
Ignore (Keq is too small)
4 x3 = (3.9x10-33) (0.445) = 7.72x10-34 
x3 = 1.93x10-34  x = 5.78x10-12 atm
PCl2 = x = 5.78x10-12 atm
PHCl = 0.445 atm – 2x = 0.445 atm
PHI = 2 x = 2(5.78x10-12 ) = 1.16x10-11 atm
24
(g) . (Kf) . (i)
MW =
(kg) . Δ T
25.
(1.450)(5.12)(1)
=
= 372 g/mol
(15.0/1000) (1.66)
π = M. R. T. I = [( 0.050/195)/(10.0/1000)] x (0.0821) x (25+273) x (1) = 0.63 atm
0.63 atm x 760 = 477 torr
Bonus
Rate = k[A]m [B]n
R1/R2 = [A]1m / [A]2 m  (2.0x10-4 / 8.0x10-4) = (0.10/0.20)m
(2/8) =(1/4) = (1/2)m  (1/2)2 = (1/2)m  m =2 ( order of B – second oreder)
R2/R3 = [A]2m [B]2n / [A]3m [B]3n  (8.0x10-4 / 1.6x10-3) = (0.10/0.20)n
(1/2) = (1/2) n  (1/2)1 = (1/2)n  n = 1 ( order of A – first oreder)
rate = k[A][B]2
rate
k=
=
[A]
[B] 2
2.0x10-4 M min-1
= 0.20 M-2 min-1
2
(0.10 M) (0.10 M)
rate = k[A][B] 2 = (0.20 M-2 min-1) ( 0.30 M)(0.40 M)2 = 7.2 x10-3 M min-1
14
Assignment 15 A
1- Write the equilibrium constant for the heterogeneous reaction
2NaHCO3(s)  Na2CO3(s) + CO2(g) + H2O(g).
a) PCO2PH2O
b) 1/[CO2][H2O]
c) [Na2CO3][CO2][H2O]/[NaHCO3]2
d) [Na2CO3][CO2][H2O]/[NaHCO3]
e) [Na2CO3][PCO2][H2O]
(The concentrations [Na2CO3(s)] and [NaHCO3(s)] are not included in the equilibrium constant
expression.)
2- Gaseous hydrogen iodide is placed in a closed 1.0-L container at 425°C, where it partially decomposes to
hydrogen and iodine: 2HI(g)  H2(g) + I2(g).
At equilibrium it is found that PHI = 3.53 × 10−3 atm; PH2 = 4.79 × 10−4 atm; and PI2 = 4.79 × 10−4 atm.
What is the value of Kp at this temperature?
a) 2.71 × 10−1
b) 5.43 × 101
c) 1.54 × 104
d) 6.50 × 10−5
e) 1.84 × 10−2
(You correctly substituted the partial pressures into the Kp expression.)
3- Consider the reaction N2O4(g)  2NO2(g). Determine the value of the equilibrium constant for this reaction
if an initial partial pressure of N2O4(g) of 0.0400 atm is reduced to 0.0055 atm at equilibrium.
(There is no NO2(g) present at the start of the reaction.)
a) 6.3
b) 1.2
c) 0.87
d) 13
e) 0.22
(The PNO2 is equal to twice the amount of PN2O4 that reacted.)
4- At 100°C the equilibrium constant for the reaction COCl2(g)  CO(g) + Cl2(g) has a value of
Kp = 2.19 × 10−10. Are the following mixtures of COCl2, CO, and Cl2 at equilibrium? If not, indicate the
direction that the reaction must proceed to achieve equilibrium.
(i) PCOCl2 = 5.00 × 10−2 atm; PCO = 3.31 × 10−6 atm; PCl2 = 3.31 × 10−6 atm
(ii) PCOCl2 = 3.50 × 10−3 atm; PCO = 1.11 × 10−5 atm; PCl2 = 3.25 × 10−6 atm
a) (i) not at equilibrium, right to left, (ii) equilibrium
b) (i) equilibrium, (ii) not at equilibrium, left to right
c) (i) not at equilibrium, left to right, (ii) equilibrium
d) (i) equilibrium, (ii) equilibrium
e) (i) equilibrium, (ii) not at equilibrium, right to left
(For reaction (i), Q = Kp. For reaction (ii), Q > Kp; thus, more reactants will form until Q = Kp.)
5- This question pertains to the equilibrium
2 POCl3(g)  2PCl3(g) + O2(g) for which Horxn = +508 kJ.
How will the equilibrium of the reaction shift if POCl3 is added to the reaction vessel?
a) The equilibrium will not shift in either direction.
b) The equilibrium will shift to generate more products.
c) The equilibrium will shift to generate more reactants.
(If moles of reactant are added, the reaction will shift toward products, thereby reducing the amount
of reactant and increasing the amount of products, until equilibrium is reestablished.)
6- Both the forward and reverse reactions of the following equilibrium are believed to be elementary steps:
CO(g) + Cl2(g)  COCl(g) + Cl(g)
At 25°C the rate constants for the forward and reverse reactions are 1.4 × 10−28 M−1 s−1 and
9.3 × 1010 M−1 s−1, respectively. What is the value for the equilibrium constant at 25°C?
a) 1.5 × 10−39
b) 6.6 × 1038
c) 1.3 × 10−17
(The equilibrium constant is the forward rate constant divided by the reverse rate constant.)
7- What is the expression for Kc for the reaction 2N2O(g) + O2(g)  4NO(g)?
a) [NO]4/[N2O]2[O2]
b) [NO]4/[N2O]2
c) [NO]/[N2O][O2]
d) [N2O]2[O2]/[NO]4
e) [N2O][O2]/[NO]
(Kc = [products]coefficients/[reactants]coefficients.)
8- For the gas-phase reaction (all components are in the gas phase) CO + 3H2  CH4 + H2O, which expression
represents Kp correctly?
a) PCH4PH2O/PCO27PH23
b) PCH4PH2O/PCOPH23
c) PCO3P>H2/PCH4PH2O
d) PCOPH23/PCH4PH2O
e) PCH4PH2O/PCO3PH2
(The partial pressure of the products is divided by the partial pressure of the reactants (with correct
power terms).
9- The proper expression for Kc for the reaction NiCO3(s) + 2H+(aq)  Ni2+(aq) + CO2(g) + H2O(l) is
a) [Ni2+]/[H+]2.
b) [Ni2+][CO2]/[H+]2.
c) [CO2].
d) [Ni2+]/[NiCO3].
e) [NiCO3]/[Ni2+].
(The solid NiCO3 and the pure liquid H2O are left out of the equilibrium constant expression.)
10- Calculate the equilibrium constant, Kp, for the reaction below if a 3.25-L tank is found to contain
0.343 atm O2, 0.0212 atm SO3, and 0.00419 atm SO2 at equilibrium.
2SO3(g)  2SO2(g) + O2(g)
a) 6.79 × 10−3
b) 8.78
c) 4.12 × 10−3
d) 2.43 × 102
e) 1.34 × 10−2
(The correct partial pressures were substituted into the Kp expression.)
11- A 2.21-L vessel was found to contain 4.18 × 10−2 mol of CO2, 2.81 × 10−2 mol of CO, and 8.89 × 10−3 mol of
O2 at 298 K. Is the system at equilibrium for the reaction 2CO2  2CO + O2? If not, which direction
must the reaction proceed to achieve equilibrium? (Kp = 1.2 × 10−13)
a) no, to the left
b) no, to the right
c) Yes
(Q > Kp; thus, reaction will proceed to the left to reduce Q.)
12- At 500 K the equilibrium constant for the reaction 2NO(g) + Cl2(g)  2NOCl(g) is Kp = 52.0.
An equilibrium mixture of the three gases has partial pressures of 0.0950 atm and 0.171 atm for NO and
Cl2, respectively. What is the partial pressure of NOCl in the mixture?
a) 0.845 atm
b) 0.283 atm
c) 8.02 × 10−2 atm
d) 5.45 × 10−3 atm
e) 2.97 × 10−5 atm
(You solved the Kp expression for PNOCl.)
13- Which one of the following statements is incorrect?
a) Adding reactants shifts the equilibrium to the right.
b) Adding products shifts the equilibrium to the left.
c) Removing a product shifts the equilibrium to the right.
d) Exothermic reactions shift the equilibrium to the left with increasing temperature.
e) Adding a catalyst shifts the equilibrium to the right.
(It has no effect, since it changes the rates of the forward and reverse reactions equally.)
14- Of the following equilibria, which one will shift to the left in response to a decrease in volume?
a) 2SO3(g)  2SO2(g) + O2(g)
b) H2(g) + Cl2(g)  2HCl(g)
c) 4Fe(s) + 3O2(g)  2Fe2O3(s)
d) N2(g) + 3H2(g)  2NH3(g)
e) CO2(g) + H2O(l)  H2CO3(aq)
(There are more moles of gaseous product.)
15- How many of the following factors affect the numerical value of K?
pressure , initial concentration, volume, temperature, chemical equation
a) 2
b) 4
c) 3
d) 1
e) 0
(Only temperature and the chemical equation affect the value of K.)
16- What is the relationship between the rate constants for the forward and reverse reactions and the equilibrium
constant for a process involving just elementary reactions (very simple primary reactions)?
a) K = kf + kr
b) K = kf − kr
c) K = kf/kr
d) K = kr/kf
e) K = kfkr
(K equals kf/kr.)
17- Suppose that the reactions A  B and B  A are both elementary processes with rate constants of
9.6 × 102 s−1 and 3.8 × 104 s−1, respectively.
(a) What is the value of the equilibrium constant for the equilibrium A  B?
(b) Which is greater at equilibrium, the concentration of A or the concentration of B?
a) 1.0, [A] > [B]
b) 4.0 × 101, [A] > [B]
c) 2.5 × 10−2, [B] > [A]
d) 2.5 × 10−2, [A] > [B]
e) 4.0 × 101, [B] > [A]
(Use Kc equals kf/kr. Since Kc < 1, there are more reactants present.)
18- If Kc = 0.0140 at 100.0 °C for the reaction:
2NOBr(g)  2NO(g) + Br2(g)
what is Kp at this same temperature for the reaction:
4NO(g) + 2Br2(g)  4NOBr(g)
a) 8.53
b) 71.4
c) 2.33
d) 5.44
(The change in the number of moles of gas is −2 for the new reaction; and you also remembered
that the temperature was 373.15 K, not 298.15 K.)
19- Consider the equilibrium system Fe3O4(s) + CO(g)  CO2(g) + 3FeO(s), a slightly endothermic reaction at
room temperature. Which of the following changes is incorrect?
a) Adding CO makes the equilibrium shift to the right.
b) Removing some FeO does not change the equilibrium.
c) Adding CO2 makes the equilibrium shift to the left.
d) Increasing the temperature above room temperature will drive the reaction to the left.
e) Adding more FeO does not change the position of the equilibrium.
(Endothermic reactions produce more products with an increase in temperature.)
20- Consider the following exothermic reaction:
N2(g) + 3H2(g)  2NH3(g)
Which of the following changes would not increase the amount of NH3 produced from given quantities
of N2 and H2?
a) increase in P
b) decrease in V
c) remove some NH3 and reestablish equilibrium
d) increase in T
e) none of these
(This is an exothermic reaction. An increase in temperature will limit the yield of product. (Note:
The system initially has no NH3 present so at least some will be formed.))
21- Consider the reaction at equilibrium as given below:
2SO3(g)  2SO2(g) + O2(g)
H° = +198 kJ
All of the following changes would shift the equilibrium to the left except one. Which one would not
cause the equilibrium to shift to the left?
a) removing some SO3
b) decreasing the container volume
c) decreasing the temperature
d) adding some SO2
e) adding a catalyst that speeds up the decomposition of SO3
(A catalyst speeds up the rate at which equilibrium is obtained, but does not alter the position of the
equilibrium.)
22- Consider the reaction below:
H° = −41 kJ
CO(g) + H2O(g)  CO2(g) + H2(g)
All of the following changes would shift the equilibrium to the right except one. Which one would not
cause the equilibrium to shift to the right?
a) adding some CO
b) removing some CO2
c) decreasing the container volume
d) decreasing the temperature
e) increasing the partial pressure of H2O
(This will increase the pressure. However, this will have no effect on the equilibrium since there is
an equal number of gas molecules on both sides of the equation.)
23- To an equilibrium mixture of 2SO2(g) + O2(g)  2SO3(g), some helium, an inert gas, is added at constant
volume. The addition of helium causes the total pressure to double. Which of the following is true?
a) The concentrations of all three gases are unchanged.
b) The number of moles of O2 increases.
c) [SO3] increases.
d) The number of moles of SO3 increases.
e) [SO2] increases.
(Since the He is not involved in the equilibrium, it has no effect on the equilibrium.)
24- A mixture is prepared with PCO = 0.035 atm, PCl2 = 0.015 atm, and PCOCl2 = 0.95 atm. It is known that Kp for
the equilibrium CO(g) + Cl2(g)  COCl2(g) is 1.2 × 103 at 400°C. Predict what will happen.
a) The reaction proceeds in the reverse direction until equilibrium is established.
b) The reaction occurs in the forward direction.
c) The reaction is at equilibrium, so no net reaction occurs.
d) It is impossible to predict without more information.
(Q > Kp. The production of reactants is favored.)
25- At 300.0 K, Kp = 54.3 for the reaction H2(g) + I2(g)  2HI(g).
If 1.0 mole of H2 and 1.0 mole of I2 are placed in a 5.0-L container, what would be the equilibrium
partial pressure of HI?
a) 0.79 atm
b) 0.88 atm
c) 3.9 atm
d) 7.7 atm
e) 1.6 atm
(The correct partial pressure terms were set up in terms of the variable x. The Kp expression was
solved, and the appropriate PHI was reported.)
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
1. Which of these species would you expect to have the lowest standard entropy (S°)?
A. CH4(g)
B. HF(g)
C. NH3(g)
D. H2O(g)
2. Which one of the following reactions would you expect to have highest ∆S°?
A. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
B. C2H2(g) + 5/2O2(g) → 2CO2(g) + H2O(g)
C. C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(g)
D. C2H6(g) + 7/2O2(g) → 2CO2(g) + 3H2O(g)
3.
Which response includes all the following processes that are accompanied by an increase in entropy?
1) 2SO2(g) + O2(g) → SO3(g)
2) H2O(l) → H2O(s)
3) Br2(l) → Br2(g)
4) H2O2(l) → H2O(l) + 1/2O2(g)
A. 1, 2, 3, 4
B. 1, 2
C. 2, 3, 4
D. 3, 4
E. 1, 4
1
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
4.
Which response includes all of the following processes that are accompanied by an increase in entropy?
1) I2(s) → I2(g)
2) 2I(g) → I2(g)
3) 2NH3(g) → N2(g) + 3H2(g)
4) Mg2+(aq) + 2OH–(aq) → Mg(OH)2(s)
A. 1, 2
B. 1, 3
C. 3, 4
D. 3
E. 2, 4
5.
Arrange these reactions according to increasing ∆S.
1) H2O(g) → H2O(l)
2) 2NO(g) → N2(g) + O2(g)
3) MgCO3(s) → MgO(s) + CO2(g)
A. 1 < 2 < 3
B. 2 < 3 < 1
C. 3 < 2 < 1
D. 2 < 1 < 3
E. 1 < 3 < 2
6. Which of the following is expected to have zero entropy?
I. N2(g) at 273 K
II. SiO2(s, amorphous) at 0 K
III. NaCl(s) perfectly ordered crystal at 25 K
IV. Na(s) perfectly ordered crystal at 0 K
A. I and IV
B. III and IV
C. I and II
D. I, II, and III
E. IV only
2
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
7.
Aluminum forms a layer of aluminum oxide when exposed to air which protects the bulk metal from further corrosion.
4Al(s) + 3O2(g) → 2Al2O3(s)
Using the thermodynamic data provided below, calculate ∆S° for this reaction.
A. 182.3 J/K·mol
B. 131.5 J/K·mol
C. –182.3 J/K·mol
D. –626.2 J/K·mol
E. –802.9 J/K·mol
8.
Determine ∆S° for the reaction SO3(g) + H2O(l) → H2SO4(l).
A. 169.2 J/K·mol
B. 1343.2 J/K·mol
C. –169.2 J/K·mol
D. –29.4 J/K·mol
E. 29.4 J/K·mol
3
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
9.
Calculate ∆S° for the reaction SO2(s) + NO2(g) → SO3(g) + NO(g).
A. 53.6 J/K·mol
B. –53.6 J/K·mol
C. –22.2 J/K·mol
D. 474.8 J/K·mol
E. –474.8 J/K·mol
10.
With respect to the system only, a reaction with ∆H < 0 and ∆S > 0 is predicted to be:
A. Spontaneous at all temperatures
B. Spontaneous at high temperatures only
C. Spontaneous at low temperatures only
D. Nonspontaneous at all temperatures
11. Which of the following is consistent with a spontaneous endothermic reaction?
A.
∆H > 0, ∆S < 0, ∆G < 0
B.
∆H > 0, ∆S > 0, ∆G < 0
C.
∆H < 0, ∆S < 0, ∆G < 0
D.
∆H < 0, ∆S > 0, ∆G > 0
E.
∆H > 0, ∆S < 0, ∆G > 0
4
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
12.
Ozone (O3) in the atmosphere can react with nitric oxide (NO):
O3(g) + NO(g) → NO2(g) + O2(g).
Calculate the ∆G° for this reaction at 25°C. (∆H° = –199 kJ/mol, ∆S° = –4.1 J/K·mol)
A. 1020 kJ/mol
B. –1.22 × 103 kJ/mol
C. 2.00 × 103 kJ/mol
D. –1.42 × 103 kJ/mol
E. –198 kJ/mol
13. For the reaction H2(g) + S(s) → H2S(g), ∆H° = –20.2 kJ/mol and ∆S° = +43.1 J/K·mol. Which of these
statements is true?
A. The reaction is only spontaneous at low temperatures.
B. The reaction is spontaneous at all temperatures.
C. ∆G° becomes less favorable as temperature increases.
D. The reaction is spontaneous only at high temperatures.
E. The reaction is at equilibrium at 25°C under standard conditions.
14. The normal boiling point of acetic acid is 118.1°C. If a sample of the acetic acid is at 125.2°C, predict the
signs of ∆H, ∆S, and ∆G for the boiling process at this temperature.
A. ∆H > 0, ∆S > 0, ∆G < 0
B. ∆H > 0, ∆S > 0, ∆G > 0
C. ∆H > 0, ∆S < 0, ∆G < 0
D. ∆H < 0, ∆S > 0, ∆G > 0
E. ∆H < 0, ∆S < 0, ∆G > 0
15.
Calculate the equilibrium constant for the decomposition of water
2H2(g) + O2(g)
2H2O(l)
at 25°C, given that ∆G°f (H2O(l)) = –237.2 kJ/mol.
A. 0.83
B. 6.3 × 10–84
C. 2.5 × 10–42
D. 1.6 × 1083
E. 4.7 × 105
5
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
16.
Nitrosyl chloride (NOCl) decomposes at elevated temperatures according to the equation
2NO(g) + Cl2(g). Calculate Kp for this reaction at 227°C. (∆H° = 81.2 kJ/mol, ∆S° = 128 J/K·mol)
2NOCl(g)
A. 1.59 × 10–2
B. 2.10 × 10–7
C. 62.8
D. 4.90 × 106
E. 3.20 × 109
17.
For the reaction 2 SO2(g) + O2(g) → 2 SO3(g), if initially P(SO2) = 1.2 atm, P(O2) = 1.8 atm, and P(SO3) = 2.1 atm, calculate ∆G for this reaction at
25°C. The following data is valid at 25°C:
A. –140.0 kJ/mol
B. –137.6 kJ/mol
C. –138.7 kJ/mol
D. 1,174.7 kJ/mol
E. –141.3 kJ/mol
6
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
18.
For the reaction 2NO(g) + O2(g) → 2NO2(g) if initially P(NO) = 1.5 atm, P(O2) = 1.4 atm, and P(NO2) = 2.0 atm, calculate ∆G for this reaction at
25°C. The following data is valid at 25°C:
A. –69.9 kJ/mol
B. –69.2 kJ/mol
C. 522.1 kJ/mol
D. –79.9 kJ/mol
E. –35.0 kJ/mol
19.
Determine the equilibrium constant (Kp) at 25°C for the reaction
CO(g) + H2O(g)
CO2(g) + H2(g).
∆G° = –28.5 kJ/mol
A. 2.9 × 10–60
B. 1.0 × 10–4
C. 1.2
D. 1.0 × 105
E. 3.4 × 1059
20.
The solubility product constant at 25°C for AgI(s) in water has the value 8.3 × 10–17. Calculate ∆Grxn at 25°C for the process AgI(s)
I– (aq) where [Ag+] = 9.1 × 10–9 and [I–] = 9.1 × 10–9.
A. +4.4 kJ/mol
B. +91.7 kJ/mol
C. 0.0 kJ/mol
D. –91.7 kJ/mol
E. –4.4 kJ/mol
7
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
Ag+(aq) +
21.
Predict the normal boiling point of triethylborane (C6H15B) using the following data:
A. 92°C
B. –21°C
C. 21°C
D. 365°C
E. 256°C
22.
For the reaction HCONH2(g)
NH3(g) + CO(g), Kc = 4.84 at 400 K. If ∆H° for this reaction is 29 kJ/mol, find Kc at 500 K.
A. 5.8
B. 0.17
C. 27
D. 0.88
E. 10.3
23.
In the gas phase, methyl isocyanate (CH3NC) isomerizes to acetonitrile (CH3CN),
H3C–N≡C (g)
H3C–C≡N (g)
with ∆H° = –89.5 kJ/mol and ∆G° = – 73.8 kJ/mol at 25°C. Find the equilibrium constant for this reaction at 100°C.
A. 1.68 × 10–10
B. 5.96 × 109
C. 2.16 × 1010
D. 4.63 × 10–11
E. 8.64 × 1012
8
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
24.
Using the thermodynamic data provided below, calculate the standard change in entropy when one mole of sodium sulfate is dissolved in water.
Will the solubility of sodium nitrate increase or decrease if the temperature of the system is increased?
A. –11.84 J/K·mol; solubility decreases with increasing temperature
B. –11.84 J/K·mol; solubility increases with increasing temperature
C. 11.84 J/K·mol; solubility decreases with increasing temperature
D. 11.84 J/K·mol; solubility increases with increasing temperature
E. None of the above
25. Assuming ∆S° and ∆H° do not vary with temperature, at what temperature will the reaction shown below
become spontaneous?
C(s) + H2O(g) → H2(g) + CO(s) (∆S° = 133.6 J/K·mol; ∆H° = 131.3 kJ/mol)
A. 670°C
B. 690°C
C. 710°C
D. 730°C
E. None of the above
26. Rubidium has a heat of vaporization of 69.0 kJ/mol at its boiling point (686°C). Calculate ∆S for this
process, Rb(l) → Rb(g), at 1 atm and 686°C.
A. 65.9 J/K·mol
B. 67.9 J/K·mol
C. 69.9 J/K·mol
D. 71.9 J/K·mol
E. None of the above
9
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
27. The free energy of formation of nitric oxide, NO, at 1000 K (roughly the temperature in an automobile
engine during ignition) is about 78 kJ/mol. Calculate the equilibrium constant Kp for the reaction N2(g) +
O2(g)
2NO(g) at this temperature.
A. Kp = 6.9 × 10–9
B. Kp = 7.1 × 10–9
C. Kp = 7.3 × 10–9
D. Kp = 7.5 × 10–9
E. None of the above
28. Predict the signs (–, +, or 0) of ∆H and ∆S, in that order, for the reaction: 6CO2(g) + 6H2O(g) → C6H12O6(g)
+ 6O2(g).
A. +, –
B. 0, +
C. +, +
D. –, –
E. None of the above
29. What is the free energy change for the reaction SiO2(s) + Pb(s) → PbO2(s) + Si(s)?
∆G°f (PbO2) = –217 kJ/mol
∆G°f (SiO2) = –856 kJ/mol
A. 619 kJ/mol
B. 639 kJ/mol
C. 659 kJ/mol
D. 679 kJ/mol
E. None of the above
30. The heat of vaporization of water is 2.27 kJ/g. What is ∆Svap per mole at the normal boiling point?
A. 170. J/K·mol
B. 150. J/K·mol
C. 130. J/K·mol
D. 110. J/K·mol
E. None of the above
10
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
31. The following reaction is nonspontaneous at 25°C:
Cu2O(s) → 2Cu(s) + 1/2O2(g), ∆G° = 141 kJ/mol
If ∆S° = 75.8 J/K·mol, what is the lowest temperature at which the reaction will be spontaneous?
A. 2160 K
B. 2260 K
C. 2360 K
D. 2460 K
E. None of the above
32.
Select True or False: The reaction 3H2(g) + N2(g)
5.0 mol H2 are mixed in a 2.5 L reactor.
True
2NH3(g), Kc = 9.0 at 350°C proceeds from right to left when 1.0 mol NH3, 5.0 mol N2, and
False
33. Consider the reaction CO(g) + 2H2(g)
∆G°f (CO) = –137.3 kJ/mol
∆G°f (CH3OH) = –166.3 kJ/mol
∆H°f (CO) = –110.5 kJ/mol
∆H°f (CH3OH) = –238.7 kJ/mol
S°(CO) = 197.9 J/K·mol
S°(CH3OH) = 126.8 J/K·mol
Calculate ∆G° at 25°C.
A. –29.0 kJ/mol
B. –31.0 kJ/mol
C. –33.0 kJ/mol
D. –35.0 kJ/mol
E. None of the above
CH3OH(l) at 25°C.
34. Select True or False: At a given temperature, O2(g) at 5 atm has a higher entropy per mole than O2(g) at 0.5
atm.
True False
11
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
35. ∆Hvap for ethanol is 38.7 kJ/mol at its boiling point (78°C). What is ∆Ssurr when 1.00 mole of ethanol is
boiled?
A. –310 J/K·mol
B. –210 J/K·mol
C. –110 J/K·mol
D. –100 J/K·mol
E. None of the above
36.
Sulfur can be separated from lead in the mineral galena, PbS(s), by "roasting" the ore in the presence of oxygen as shown in the following reaction:
2PbS(s) + 3O2(g) → 2PbO(s) + 2SO2(g)
Determine ∆G for the above reaction at 850°C.
A. –620 kJ/mol
B. –640 kJ/mol
C. –660 kJ/mol
D. –680 kJ/mol
E. None of the above
37.
Given the following data, calculate the boiling point of HCOOH (formic acid).
A. 115°C
B. 125°C
C. 145°C
D. 165°C
E. None of the above
12
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
38.
Using the thermodynamic data provided below, calculate Ka1 for H2CO3(aq) at 25°.
A. 5 × 10–5
B. 5 × 10–6
C. 5 × 10–7
D. 5 × 10–8
E. None of the above
39. Select True or False: Melting an ionic solid always results in an increase in entropy.
True False
40. Select True or False: Dissolving an ionic solid in water always results in an increase in entropy.
True False
13
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
Key
1. B
2. D
3. D
4. B
5. A
6. E
7. D
8. C
9. C
10. A
11. B
12. E
13. B
14. A
15. B
16. A
17. C
18. B
19. D
20. C
21. A
22. C
23. B
24. A
1
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
25. C
26. D
27. B
28. A
29. B
30. D
31. A
32. FALSE
33. A
34. FALSE
35. C
36. C
37. A
38. C
39. TRUE
40. FALSE
2
CHEM 1412. Chapter 19. Chemical Thermodynamics (Homework) Ky
Name: ________________________ Class: ___________________ Date: __________
Ch 17 Electrochemistry Practice Test
Multiple Choice
Identify the choice that best completes the statement or answers the question.
____
1) Which element is reduced in the reaction below?
Fe(CO)5 (l) + 2HI (g)  Fe(CO)4I2 (s) + CO (g) + H2 (g)
____
A) Fe
B) C
C) O
D) H
E) I
2) Which of the following reactions is a redox reaction?
____
(a) K2CrO4 + BaCl2  BaCrO4 + 2KCl
(b) Pb22+ + 2Br–  PbBr
(c) Cu + S  CuS
A) (a) only
B) (b) only
C) (c) only
D) (a) and (c)
E) (b) and (c)
3) Which substance is the oxidizing agent in the following reaction?
____
Fe2S3 + 12HNO3  2Fe(NO3)3 + 3S + 6NO2 + 6H2O
A) HNO3
B) S
C) NO2
D) Fe2S3
E) H2O
4) What is the coefficient of the permanganate ion when the following equation is balanced?
MnO4– + Br–  Mn2+ + Br2
____
(acidic solution)
A) 1
B) 2
C) 3
D) 5
E) 4
5) What is the coefficient of Fe3+ when the following equation is balanced?
CN– + Fe3+  CNO– + Fe2+
A)
B)
C)
D)
E)
(basic solution)
1
2
3
4
5
1
ID: A
Name: ________________________
____
____
____
ID: A
6) Which transformation could take place at the anode of an electrochemical cell?
A) Cr2O72–  Cr2+
B) F2 to F–
C) O2 to H2O
D) HAsO2 to As
E) None of the above could take place at the anode.
7) The purpose of the salt bridge in an electrochemical cell is to ________.
A) maintain electrical neutrality in the half-cells via migration of ions
B) provide a source of ions to react at the anode and cathode
C) provide oxygen to facilitate oxidation at the anode
D) provide a means for electrons to travel from the anode to the cathode
E) provide a means for electrons to travel from the cathode to the anode
8) Which transformation could take place at the cathode of an electrochemical cell?
A) MnO2  MnO4–
B) Br2  BrO3–
C) NO  HNO2
D) HSO4–  H2SO3
E) Mn2+  MnO4–
Table 20.1
Half Reaction
F2 (g) + 2e–  2F– (aq)
Cl2 (g) + 2e–  2Cl– (aq)
Br2 (l) + 2e–  2Br– (aq)
O2 (g) + 4H+ (aq) + 4e–  2H2O (l)
Ag+ + e–  Ag (s)
Fe3+ (aq) + e–  Fe2+ (aq)
I2 (s) + 2e–  2I– (aq)
Cu2+ + 2e–  Cu (s)
2H+ + 2e–  H2 (g)
Pb2+ + 2e–  Pb (s)
Ni2+ + 2e–  Ni (s)
Li+ + e–  Li (s)
E°(V)
+2.87
+1.359
+1.065
+1.23
+0.799
+0.771
+0.536
+0.34
0
–0.126
–0.28
–3.05
____
9) Which of the halogens in Table 20.1 is the strongest oxidizing agent?
A) Cl2
B) Br2
C) F2
D) I2
E) All of the halogens have equal strength as oxidizing agents.
____ 10) Which one of the following types of elements is most likely to be a good oxidizing agent?
A) alkali metals
B) lanthanides
C) alkaline earth elements
D) transition elements
E) halogens
2
Name: ________________________
ID: A
____ 11) Consider an electrochemical cell based on the reaction:
2H+ (aq) + Sn (s)  Sn2+ (aq) + H2 (g)
____ 12)
____ 13)
____ 14)
____ 15)
____ 16)
____ 17)
Which of the following actions would change the measured cell potential?
A) increasing the pH in the cathode compartment
B) lowering the pH in the cathode compartment
C) increasing the [Sn2+] in the anode compartment
D) increasing the pressure of hydrogen gas in the cathode compartment
E) Any of the above will change the measure cell potential.
What is the anode in an alkaline battery?
A) MnO2
B) KOH
C) Zn powder
D) Mn2O3
E) Pt
What is the cathode in an alkaline battery?
A) MnO2
B) KOH
C) Zn powder
D) Mn2O3
E) Pt
In a lead-acid battery, the electrodes are consumed. In this battery, ________.
A) the anode is Pb
B) the anode is PbSO4
C) the anode is PbO2
D) the cathode is PbSO4
E) the cathode is Pb
Cathodic protection of a metal pipe against corrosion usually entails ________.
A) attaching an active metal to make the pipe the anode in an electrochemical cell
B) coating the pipe with another metal whose standard reduction potential is less negative
than that of the pipe
C) attaching an active metal to make the pipe the cathode in an electrochemical cell
D) attaching a dry cell to reduce any metal ions which might be formed
E) coating the pipe with a fluoropolymer to act as a source of fluoride ion (since the latter is
so hard to oxidize)
One of the differences between a voltaic cell and an electrolytic cell is that in an electrolytic cell, ________.
A) an electric current is produced by a chemical reaction
B) electrons flow toward the anode
C) a nonspontaneous reaction is forced to occur
D) O2 gas is produced at the cathode
E) oxidation occurs at the cathode
What is the oxidation number of bromine in the HBrO molecule?
A) +1
B) +2
C) 0
D) –1
E) –2
3
Name: ________________________
ID: A
____ 18) What is the oxidation number of sulfur in the S 2O32– ion?
A) +2
B) +1
C) 0
D) –1
E) –2
____ 19) Which substance is the oxidizing agent in the reaction below?
Fe(CO)5 (l) + 2HI (g)  Fe(CO)4I2 (s) + CO (g) + H2 (g)
A) HI
B) Fe(CO)5
C) Fe(CO)4I2
D) CO
E) H2
____ 20) Which element is reduced in the reaction below?
Fe2+ + H+ + Cr2O72–  Fe3+ + Cr3+ + H2O
A) Cr
B) Fe
C) H
D) O
____ 21) Which element is oxidized in the reaction below?
I– + MnO4– + H+  I2 + MnO2 + cO
A) I
B) Mn
C) O
D) H
____ 22) What is the correct coefficient for the electrons in the following half-reaction:
Ni6+ + ___e–  Ni
A) 6
B) 1
C) 2
D) 3
E) 5
____ 23) In the galvanic cell using the redox reaction below, the reduction half-reaction is ________.
A)
B)
C)
D)
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
+ 2e–  Cu
Zn  Zn2+ + 2e–
Cu2+  Cu + 2e–
Zn + 2e–  Zn2+
Cu2+
4
Name: ________________________
ID: A
____ 24) In the electrochemical cell using the redox reaction below, the oxidation half reaction is ________.
Sn4+ (aq) + Fe (s)  Sn2+ (aq) + Fe2+ (aq)
A)
+ 2e–  Sn2+
B) Fe  Fe2+ + 2e–
C) Sn4+  Sn2+ + 2e–
D) Fe + 2e–  Fe2+
E) Fe + 2e– ? Sn2+
____ 25) The standard cell potential (E°) of a voltaic cell constructed using the cell reaction below is 0.76 V:
Sn4+
Zn (s) + 2H+ (aq)  Zn2+ (aq) + H2 (g)
With PH2 = 1.0 atm and [Zn2+] = 1.0 M, the cell potential is 0.53 V. The concentration of H + in the cathode
compartment is ________ M.
A) 1.3  10–4
B) 1.7  10–8
C) 1.1  10–2
D) 7.7  103
E) 1.3  10–11
____ 26) The standard cell potential (E° cell) for the reaction below is +0.63 V. The cell potential for this reaction is
________ V when [ Zn 2+] = 3.0 M and [Pb2+] = 2.0  10–4 M.
Pb2+ (aq) + Zn (s)  Zn2+ (aq) + Pb (s)
A) 0.51
B) 0.86
C) 0.40
D) 0.75
E) 0.63
____ 27) A voltaic cell is constructed with two Zn 2+–Zn electrodes, where the half-reaction is
Zn2+ + 2e–  Zn (s)
E° = –0.763 V
The concentrations of zinc ion in the two compartments are 4.50 M and 1.11  10–2 M, respectively. The cell
emf is ________ V.
A) –1.88  10–3
B) –309
C) 0.0772
D) 0.154
E) –0.761
5
Name: ________________________
ID: A
____ 28) The standard emf for the cell using the overall cell reaction below is +2.20 V:
2Al (s) + 3I2 (s)  2Al3+ (aq) + 6I– (aq)
____ 29)
____ 30)
____ 31)
____ 32)
____ 33)
The emf generated by the cell when [Al3+] = 3.5  10–3 M and [I–] = 0.30 M is ________ V.
A) 2.20
B) 2.28
C) 2.12
D) 2.36
E) 2.23
The electrolysis of molten AlCl3 for 2.50 hr with an electrical current of 12.0 A produces ________ g of
aluminum metal.
A) 90.7
B) 0.373
C) 2.80  10–3
D) 10.1
E) 30.2
How many seconds are required to produce 5.00 g of aluminum metal from the electrolysis of molten AlCl 3
with an electrical current of 15.0 A?
A) 27.0
B) 9.00
C) 1.19E  103
D) 2.90  105
E) 3.57  103
How many minutes will it take to plate out 2.19 g of chromium metal from a solution of Cr 3+ using a current of
19.5 amps in an electrolyte cell?
A) 10.4
B) 625
C) 208
D) 3.47
E) 31.2
What current (in A) is required to plate out 1.22 g of nickel from a solution of Ni 2+ in 0.50 hour?
A) 65.4
B) 8.02  103
C) 2.22
D) 12.9
E) 4.46
How many grams of copper will be plated out by a current of 2.3 A applied for 35 minutes to a 0.50 M solution
of copper (II) sulfate?
A) 1.6
B) 3.2
C) 1.8  10–2
D) 3.6  10–2
E) 0.019
6
ID: A
Ch 17 Electrochemistry Practice Test
Answer Section
MULTIPLE CHOICE
1) ANS:
OBJ:
2) ANS:
OBJ:
3) ANS:
OBJ:
4) ANS:
OBJ:
5) ANS:
OBJ:
6) ANS:
OBJ:
7) ANS:
OBJ:
8) ANS:
OBJ:
9) ANS:
OBJ:
10) ANS:
OBJ:
11) ANS:
OBJ:
12) ANS:
OBJ:
13) ANS:
OBJ:
14) ANS:
OBJ:
15) ANS:
OBJ:
16) ANS:
OBJ:
17) ANS:
OBJ:
18) ANS:
OBJ:
19) ANS:
OBJ:
20) ANS:
OBJ:
21) ANS:
OBJ:
D
20.1; G2
C
20.1; G2
A
20.1; G2
B
20.2; G2
B
20.2; G2
E
20.3; G2
A
20.3; G2
D
20.3; G2
C
20.4; G3
E
20.4; G3
E
20.6; G2
C
20.7; G2
A
20.7; G2
A
20.7; G2
C
20.8; G2
C
20.9; G2
A
20.1; G2
A
20.1; G2
A
20.1; G2
A
20.1; G2
A
20.1; G2
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 2
REF: Page Ref: 20.2
PTS: 1
DIF: 2
REF: Page Ref: 20.2
PTS: 1
DIF: 1
REF: Page Ref: 20.3
PTS: 1
DIF: 1
REF: Page Ref: 20.3
PTS: 1
DIF: 1
REF: Page Ref: 20.3
PTS: 1
DIF: 1
REF: Page Ref: 20.4
PTS: 1
DIF: 1
REF: Page Ref: 20.4
PTS: 1
DIF: 1
REF: Page Ref: 20.6
PTS: 1
DIF: 1
REF: Page Ref: 20.7
PTS: 1
DIF: 1
REF: Page Ref: 20.7
PTS: 1
DIF: 1
REF: Page Ref: 20.7
PTS: 1
DIF: 2
REF: Page Ref: 20.8
PTS: 1
DIF: 1
REF: Page Ref: 20.9
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
PTS: 1
DIF: 1
REF: Page Ref: 20.1
1
ID: A
22) ANS:
OBJ:
23) ANS:
OBJ:
24) ANS:
OBJ:
25) ANS:
OBJ:
26) ANS:
OBJ:
27) ANS:
OBJ:
28) ANS:
OBJ:
29) ANS:
OBJ:
30) ANS:
OBJ:
31) ANS:
OBJ:
32) ANS:
OBJ:
33) ANS:
OBJ:
A
20.2; G2
A
20.3; G2
D
20.3; G2
A
20.6; G4
A
20.6; G4
C
20.6; G4
B
20.6; G4
D
20.9; G4
E
20.9; G4
A
20.9; G4
C
20.9; G4
A
20.9; G4
PTS: 1
DIF: 2
REF: Page Ref: 20.2
PTS: 1
DIF: 2
REF: Page Ref: 20.3
PTS: 1
DIF: 2
REF: Page Ref: 20.3
PTS: 1
DIF: 3
REF: Page Ref: 20.6
PTS: 1
DIF: 2
REF: Page Ref: 20.6
PTS: 1
DIF: 2
REF: Page Ref: 20.6
PTS: 1
DIF: 2
REF: Page Ref: 20.6
PTS: 1
DIF: 2
REF: Page Ref: 20.9
PTS: 1
DIF: 2
REF: Page Ref: 20.9
PTS: 1
DIF: 2
REF: Page Ref: 20.9
PTS: 1
DIF: 2
REF: Page Ref: 20.9
PTS: 1
DIF: 2
REF: Page Ref: 20.9
2
Chapter 20
MULTIPLE CHOICE. Choose the one alternative that best completes the statement or answers the question.
1) The gain of electrons by an element is called __________.
1)
A) oxidation
B) reduction
C) sublimation
D) fractionation
E) disproportionation
2) __________ is reduced in the following reaction:
2)
Cr2 O72- + 6S2 O32- + 14H+ → 2Cr3+ + 3S4 O6 2- + 7H2O
A) S2O3 2-
B) Cr3+
D) Cr2O7 2-
C) H+
E) S4O6 2-
3) __________ is the oxidizing agent in the reaction below.
3)
Cr2 O7 2- + 6S2 O3 2- + 14H+ → 2Cr3+ + 3S4 O62- + 7H2 O
A) H+
B) Cr3+
C) S2 O3 2-
D) S4 O6 2-
E) Cr2 O72-
4) Which substance is serving as the reducing agent in the following reaction?
4)
14H+ + Cr2 O7 2- + 3Ni → 3Ni2+ + 2Cr3+ + 7H2O
A) H2O
B) Ni
C) Cr2 O7 2-
D) Ni 2+
E) H+
5) Which substance is the reducing agent in the reaction below?
5)
Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2 O
A) PbSO4
B) PbO2
C) H2 SO4
D) Pb
E) H2O
26) What is the oxidation number of chromium in Cr2 O 7 ion ?
A) +12
B) +14
C) +6
6)
D) +3
E) +7
7) What is the oxidation number of potassium in KMnO4 ?
A) +3
B) 0
C) +1
7)
D) -1
1
E) +2
18) What is the oxidation number of manganese in the MnO 4 ion?
A) +4
B) +2
C) +7
8)
D) +1
E) +5
9) What is the oxidation number of manganese in MnO2 ?
A) +2
B) +7
C) +1
9)
D) +4
E) +3
10) __________ electrons appear in the following half-reaction when it is balanced.
10)
S4O6 2- → 2S2O3 2A) 2
B) 4
C) 6
D) 3
E) 1
11) The balanced half-reaction in which chlorine gas is reduced to the aqueous chloride ion is a
__________ process.
11)
A) two-electron
B) six-electron
C) four-electron
D) one-electron
E) three-electron
12) The balanced half-reaction in which dichromate ion is reduced to chromium metal is a __________
process.
12)
A) two-electron
B) six-electron
C) twelve-electron
D) four-electron
E) three-electron
13) The balanced half-reaction in which dichromate ion is reduced to chromium(III) ion is a
__________ process.
A) three-electron
B) six-electron
C) twelve-electron
D) two-electron
E) four-electron
2
13)
14) The balanced half-reaction in which sulfate ion is reduced to sulfite ion is a __________ process.
14)
A) six-electron
B) one-electron
C) four-electron
D) two-electron
E) three-electron
15) The electrode at which oxidation occurs is called the __________ .
15)
A) cathode
B) oxidizing agent
C) reducing agent
D) voltaic cell
E) anode
16) The half-reaction occurring at the anode in the balanced reaction shown below is __________.
16)
3MnO4 - (aq) + 24H+ (aq) + 5Fe (s) → 3Mn2+ (aq) + 5Fe3+ (aq) + 12H2 O (l)
A) 2MnO4 - (aq) + 12H+ (aq) + 6e- → 2Mn2+ (aq) + 3H2 O (l)
B) MnO4 - (aq) + 8H+ (aq) + 5e- → Mn2+ (aq) + 4H2 O (l)
C) Fe (s) → Fe3+ (aq) + 3eD) Fe2+ (aq) → Fe3+ (aq) + eE) Fe (s) → Fe2+ (aq) + 2e17) In a voltaic cell, electrons flow from the __________ to the __________.
17)
A) cathode, anode
B) salt bride, anode
C) anode, cathode
D) anode, salt bridge
E) salt bridge, cathode
18) The reduction half reaction occurring in the standard hydrogen electrode is __________.
A) 2H+ (aq) + 2OH- → H2 O (l)
B) O2 (g) + 4H+ (aq) + 4e- → 2H2 O (l)
C) 2H+ (aq, 1M) + Cl2 (aq) → 2HCl (aq)
D) H2 (g, 1 atm) → 2H+ (aq, 1M) + 2eE) 2H+ (aq, 1M) + 2e- → H2 (g, 1 atm)
3
18)
19) 1V = __________.
A) 1 J/s
19)
B) 96485 C
C) 1 J/C
D) 1 amp · s
E) 1 C/J
20) The more __________ the value of E°red, the greater the driving force for reduction.
20)
A) exothermic
B) extensive
C) endothermic
D) negative
E) positive
Table 20.2
Half-reaction
Cr3+ (aq) + 3e- → Cr (s)
Fe2+ (aq) + 2e- → Fe (s)
Fe3+ (aq) + e- → Fe2+ (s)
Sn4+ (aq) + 2e- → Sn2+ (aq)
E° (V)
-0.74
-0.440
+0.771
+0.154
21) The standard cell potential (E°cell ) for the voltaic cell based on the reaction below is __________ V.
21)
Sn2+ (aq) + 2Fe3+ (aq) → 2Fe2+ (aq) + Sn4+ (aq)
A) +0.617
B) +1.39
C) +1.21
D) -0.46
E) +0.46
22) The standard cell potential (E°cell ) for the voltaic cell based on the reaction below is __________ V.
22)
Cr (s) + 3Fe3+ (aq) → 3Fe2+ (aq) + Cr3+ (aq)
A) +1.57
B) -1.45
C) +1.51
D) +3.05
E) +2.99
23) The standard cell potential (E°cell ) for the voltaic cell based on the reaction below is __________ V.
23)
2Cr (s) + 3Fe2+ (aq) → 3Fe (s) + 2Cr3+ (aq)
A) -0.16
B) +3.10
C) +0.83
D) +2.80
E) +0.30
24) The standard cell potential (E°cell ) for the voltaic cell based on the reaction below is __________ V.
3Sn4+ (aq) + 2Cr (s) → 2Cr3+ (aq) + 3Sn2+ (aq)
A) +1.94
B) -0.59
C) +0.89
4
D) -1.02
E) +2.53
24)
25) The relationship between the change in Gibbs free energy and the emf of an electrochemical cell is
given by __________.
-nF
A) ΔG =
ERT
B) ΔG =
-nF
E
C) ΔG =
-E
nF
25)
D) ΔG = -nFE
E) ΔG = -nRTF
26) The standard cell potential (E°cell ) of the reaction below is +0.126 V. The value of ΔG° for the
26)
reaction is __________ kJ/mol.
Pb (s) + 2H+ (aq) → Pb2+ (aq) + H2 (g)
A) -24
B) +24
C) -12
D) +12
E) -50
27) The standard cell potential (E°cell ) for the reaction below is +0.63 V. The cell potential for this
27)
reaction is __________ V when [ Zn2+ ] = 1.0 M and [Pb2+ ] = 2.0 × 10-4 M.
Pb2+ (aq) + Zn (s) → Zn2+ (aq) + Pb (s)
A) 0.74
B) 0.41
C) 0.85
D) 0.63
E) 0.52
28) The standard cell potential (E°cell ) for the reaction below is +1.10 V. The cell potential for this
28)
reaction is __________ V when the concentration of [Cu2+ ] = 1.0 × 10-5 M and [Zn2+ ] = 1.0 M.
Zn (s) + Cu2+ (aq) → Cu (s) + Zn2+ (aq)
A) 0.95
B) 1.25
C) 1.10
D) 0.80
E) 1.40
29) The lead-containing reactant(s) consumed during recharging of a lead-acid battery is/are
__________.
A) PbO2 (s) only
B) Pb (s) only
C) PbSO4 (s) only
D) both PbO2 (s) and PbSO4 (s)
E) both Pb (s) and PbO2 (s)
5
29)
30) Galvanized iron is iron coated with __________.
30)
A) zinc.
B) chromium.
C) phosphate.
D) magnesium.
E) iron oxide.
31) Corrosion of iron is retarded by __________.
31)
A) the presence of salts
B) low pH conditions
C) high pH conditions
D) both the presence of salts and high pH conditions
E) both the presence of salts and low pH conditions
32) How many minutes will it take to plate out 2.19 g of chromium metal from a solution of Cr3+ using
a current of 35.2 amps in an electrolyte cell __________ ?
A) 17.3
B) 115
C) 346
D) 1.92
E) 5.77
33) What current (in A) is required to plate out 1.22 g of nickel from a solution of Ni 2+ in 1.0 hour
__________ ?
A) 65.4
B) 2.34
C) 1.11
D) 12.9
32)
33)
E) 4.01 × 103
34) How many grams of Ca metal are produced by the electrolysis of molten CaBr2 using a current of
34)
30.0 amp for 10.0 hours __________ ?
A) 22.4
B) 448
C) 0.0622
D) 112
E) 224
35) How many grams of CuS are obtained by passing a current of 12 A through a solution of CuSO4
35)
for 15 minutes __________ ?
A) 3.6
B) 7.1
C) 14
D) 1.8
E) 0.016
36) How many seconds are required to produce 1.0 g of silver metal by the electrolysis of a AgNO3
36)
solution using a current of 30 amps __________ ?
A) 3.7 × 10-5
B) 60
C) 3.2 × 103
D) 30
E) 2.7 × 104
37) How many grams of copper will be plated out by a current of 2.3 A applied for 25 minutes to a
0.50-M solution of copper(II) sulfate __________ ?
A) 0.019
B) 2.2
C) 0.036
6
D) 1.1
E) 1.8 × 10-2
37)
Answer Key
Testname: CHAPTER 20 PRACTICE QUESTIONS
1) B
2) D
3) C
4) B
5) D
6) C
7) C
8) C
9) D
10) A
11) A
12) C
13) B
14) D
15) E
16) C
17) C
18) E
19) C
20) E
21) A
22) C
23) E
24) C
25) D
26) A
27) E
28) A
29) C
30) A
31) C
32) E
33) C
34) E
35) A
36) D
37) D
7
Page 1
1.
The dissolution of water in octane (C8H18) is prevented by ___________.
a. London dispersion forces between octane molecules
b. hydrogen bonding between water molecules
c. dipole-dipole attraction between octane molecules
d. ion-dipole attraction between water and octane molecules
e. repulsion between like-charged water and octane molecules
2.
Which combination cannot constitute a solution?
a. gaseous solvent, gaseous solute
b. gaseous solvent, solid solute
c. liquid solvent, gaseous solute
d. solid solvent, liquid solute
e. solid solvent, gaseous solute
3.
Hydration is a specific example of the phenomenon known generally as
____________.
a. dissolution
b. disordering
c. solvation
d. condensation
e. dilution
4.
The dissolution of gases in water is virtually always exothermic because
____________.
a. one of the endothermic steps of the three-step dissolution process is
unnecessary (separation of solute particles)
b. the exothermic step of the three-step dissolution process is
unnecessary
c. gases react exothermically with water
d. neither of the endothermic steps in the three-step dissolution
process is necessary
e. all three steps in the three-step dissolution process are exothermic
5.
Spontaneous dissolution processes can be endothermic ____________.
a. if they are accompanied by another process that is exothermic
b. if they are accompanied by an increase in order
c. if they are accompanied by an increase in disorder
d. if the solvent is a gas and the solute is a solid
e. if the solvent is water and the solute is a gas
6.
When argon is placed in a container of neon, the argon spontaneously
disperses throughout the neon because ______________.
a. large attractive forces between argon and neon atoms
b. hydrogen bonding
c. a decrease in energy occurs when the two mix
d. the dispersion of argon atoms produces an increase in disorder
e. solvent-solute interactions
Page 2
7.
The phrase "like dissolves like" refers to the fact that ____________.
a. gases can only dissolve other gases
b. polar solvents dissolve polar solutes and nonpolar solvents dissolve
nonpolar solutes
c. solvents can only dissolve solutes of similar molar mass
d. condensed phases can only dissolve other condensed phases
e. polar solvents dissolve nonpolar solutes and vice versa
8.
A saturated solution _______________.
a. contains as much solvent as it can hold
b. contains no double bonds
c. contains dissolved solute in equilibrium with undissolved solid
d. will rapidly precipitate if a seed crystal is added
e. cannot be attained
9.
Pairs of liquids that will mix in all proportions are called
___________________ liquids.
10. The principal reason for the extremely low solubility of NaCl in benzene
(C6H6) is the ________________.
a. great strength of solvent-solvent interactions
b. great strength of solute-solvent interactions
c. great strength of solute-solute interactions
d. great weakness of solute-solvent interactions
e. c and d
11. Which one of the following substances would be the most soluble in CCl4?
a. CH3CH2OH
b. H2O
c. NH3
d. C10H22
e. NaCl
12. The lowest value of the Henry's Law constant (kH) would be obtained with
________ as the solvent at a temperature of _______ K when the solute is
methane gas (CH4).
a. C5H12, 301
b. C6H6, 322
c. C6H6, 349
d. H2O, 301
e. H2O, 349
Page 3
13. Pressure has an appreciable effect on the solubility of _____________ in
liquids.
a. gases only
b. solids only
c. liquids only
d. all of the above
e. solids and liquids
14. The solubility of Ar in water at 25øC is 1.6 x 10-3 mol/L when the
pressure of the Ar above the solution is 1.0 atm. The solubility of Ar
at a pressure of 2.5 atm is ______________ mol/L.
a. 1.6 x 103
b. 6.4 x 10-4
c. 4.0 x 10-3
d. 7.5 x 10-2
e. 1.6 x 10-3
15. Which of the following choices has the compounds correctly arranged in
order of increasing solubility in water? (Least soluble to most
soluble.)
a. CCl4 < CHCl3 < NaNO3
b. CH3OH < CH4 < LiF
c. CH4 < NaNO3 < CHCl3
d. LiF < NaNO3 < CHCl3
e. CH3OH < CCl4 < CHCl3
16. Which component of air is the primary problem in a condition known as
the bends?
a. O2
b. CO2
c. He
d. N2
e. CO
17. A solution contains 28% phosphoric acid by mass. This means that
___________.
a. 1 mL of this solution contains 28 g of phosphoric acid
b. 1 L of this solution has a mass of 28 g
c. 100 g of this solution contains 28 g of phosphoric acid
d. 1 L of this solution contains 28 mL of phosphoric acid
e. the density of this solution is 2.8 g/mL
Page 4
18. A solution is prepared by dissolving 23.7 g of CaCl2 in 375 g of water.
The density of the resulting solution is 1.05 g/mL. The concentration
of CaCl2 (mass %) is _____________.
a. 5.94
b. 6.32
c. 0.0632
d. 0.0594
e. 6.24
19. The concentration of KBr in a solution prepared by dissolving 2.21 g of
KBr in 897 g of water is ______________ molal.
a. 2.46
b. 0.0167
c. 0.0207
d. 2.07 x 10-5
e. 0.0186
20. The concentration (molal) of lead nitrate in 0.726 M Pb(NO3)2 (density
1.202 g/mL) is
a. 0.476
b. 1.928
c. 0.755
d. 0.819
e. 0.650
21. A solution was prepared by dissolving 15.0 g of NH3 in 250 g of water.
The density of the resulting solution was 0.974 g/mL. The mole fraction
of NH3 in the solution is _____________.
a. 0.0640
b. 0.0597
c. 0.940
d. 0.922
e. 16.8
22. A solution was prepared by dissolving 23.7 g of CaCl2 in 375 g of water.
The density of the resulting solution was 1.05 g/mL. The concentration
of Cl- in this solution is __________ M.
a. 0.214
b. 0.562
c. 1.12
d. 1.20
e. 6.64 x 10-2
Page 5
23. What is the density (g/mL) of an aqueous solution of sodium chloride
that has a molar concentration of 2.22 M and is 11.0% sodium chloride by
mass?
a. 0.998
b. 1.20
c. 4.95
d. 1.18
e. 2.22
24. A 0.100 m solution of which one of the following solutes will have the
lowest vapor pressures?
a. KClO4
b. Ca(ClO4)2
c. Al(ClO4)3
d. sucrose
e. NaCl
25. The magnitudes of Kf and of Kb depend on the identity of the
_______________.
a. solute
b. solvent
c. solution
d. solvent and on temperature
e. solute and solvent
26. Adding solute to a solution decreases the ____________ of the solution.
a. freezing point
b. osmotic pressure
c. boiling point
d. vapor pressure
e. freezing point and vapor pressure
27. The vapor pressure of pure ethanol at 60øC is 349 torr. Raoult's Law
predicts that a solution prepared by dissolving 10.0 mmol naphthalene
(nonvolatile) in 90.0 mmol ethanol will have a vapor pressure of
__________ torr.
a. 34.9
b. 314
c. 600
d. 279
e. 69.8
Page 6
28. What is the freezing point (øC) of a solution prepared by dissolving
11.3 g of Ca(NO3)2 (formula weight = 164 g/mol) in 115 g of water? The
molal freezing point depression constant for water is 1.86 øC/m.
a. -3.34
b. -1.11
c. 3.34
d. 1.11
e. 0.00
29. After swimming in the ocean for several hours, swimmers noticed that
their fingers appeared to be very wrinkled. This is an indication that
seawater is ___ relative to the fluid in cells.
a. isotonic
b. hypertonic
c. hypotonic
d. none of these
e. supertonic
30. A 1.35 m aqueous solution of compound X had a boiling point of 101.4øC.
Which one of the following could be compound X? (The boiling point
elevation constant for water is 0.52 øC/m.)
a. CH3CH2OH
b. C6H12O6
c. Na3PO4
d. KCl
e. CaCl2
31. A solution containing 100 g unknown liquid and 900 g water has a
freezing point of -3.33øC. Given Kf = 1.86øC/m for water, the molecular
weight of the unknown liquid is _________ amu.
a. 69.0
b. 333
c. 619
d. 161
e. 62.1
32. A solution prepared by dissolving 0.60 g of nicotine (a nonelectrolyte)
in water to make 12 mL of solution has an osmotic pressure of 7.55 atm
at 25øC. The molecular weight of nicotine is ______ g/mol.
a. 28
b. 43
c. 50
d. 160
e. 0.60
Page 7
33. Which of the following aqueous solutions will have the highest boiling
point?
a. 0.10 m Na2SO4
b. 0.20 m glucose
c. 0.25 m sucrose
d. 0.10 m NaCl
e. 0.10 m SrSO4
34. Determine the freezing point (øC) of a 0.015 molal aqueous solution of
MgSO4. Assume i = 2.0 for MgSO4.
a. -0.056
b. -0.028
c. -0.17
d. -0.084
e. 0.000
Page 1
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
21.
22.
23.
24.
25.
26.
27.
28.
29.
30.
31.
32.
33.
34.
b
b
c
a
c
d
b
c
miscible
e
d
e
a
c
a
d
c
a
c
c
b
c
d
c
b
e
b
a
b
d
e
d
a
a