University of Texas, Arlington

CHEM 1442

1442 Final Review

Properties of Solutions

Useful Information:

R = 0.08206 L.atm/mol.K = 8.314 J/mol.K

1 atm = 760 torr = 760 mm Hg

Psoln = (Psolv) (Xsolv)

T_f = imK_f ( or T_b )

 Psoln = (Psolv) ( Xsolute)

P_total = P_AX_a + P_bX_b (two volatile liquids)

π = MRT ( R = 0.0821 Latm/mol K)

1. An aqueous solution of sulfuric acid (H₂SO₄) has a concentration of 11.0 M and a density of 1.59 g/mL. What is the mass percent of H₂SO₄ in this solution? (The molar mass of H₂SO₄ is 98.1 g/mol; the molar mass of H₂O is 18.02 g/mol.)

2. Rank the substances listed below in order of increasing solubility in water:
   1. CH₃CH₂CH₂CH₃
   2. CH₃OH
   3. CH₃CH₂CH₂CH₂OH

a)        I < II < III

b)        I < III < II

3. II < I < III
4. II < III < I
5. III < II < I

3. What is the vapor pressure of a solution formed by dissolving 0.500 mol of a nonvolatile nonelectrolyte in 2.50 mol of methanol at 20°C? (The vapor pressure of pure methanol is 88.7 torr at 20°C.)

1. 14.8 torr
2. 17.7 torr

c)                                                73.9 torr

4. 63.4 torr
5. 106 torr

4. At a given temperature, you have a mixture of benzene (vapor pressure of pure benzene = 745 torr) and toluene (vapor pressure of pure toluene = 290. torr). The mole fraction of benzene in the solution is 0.425. Assuming ideal behavior, calculate the mole fraction of toluene in the vapor above the solution.

a)        0.425

b)        0.517

c)        0.566

d)        0.638

e)        0.345

5. When acetone is mixed with water, the ?H_soln is found to be exothermic. How does the actual vapor pressure of the solution compare to that predicted by Raoult’s law?

a)        The actual vapor pressure of the solution is greater than that predicted by Raoult’s law.

b)        The actual vapor pressure of the solution is less than that predicted by Raoult’s law.

c)        The actual vapor pressure is the same as that predicted by Raoult’s law.

6. Which of the aqueous  below has the lowest freezing point?

1. 0.12 m Na₃PO₄
2. 0.25 m KCl

c)        0.22 m CaCl₂

4. 0.30 m NaCl
5. 0.40 m glucose (a nonelectrolyt

7. Pure cyclohexane (C₆H₁₂) boils at 81.0°C. What is the boiling point of a solution in which 40.0 g of bromine (Br₂, a nonelectrolyte) is dissolved in 500 g of cyclohexane? (K_b for cyclohexane = 2.80 °C/m)

a)        79.6° C

b)        83.8° C

c)        81.7° C

d)        78.2° C

e)        82.4° C

8. A 100.0 mL aqueous solution containing 1.73 g of an unknown nonvolatile nonelectrolyte generates an osmotic pressure of 2.72 atm at 25.0°C. What is the molar mass of the unknown compound?

a)        156 g/mol

2. 78.0 g/mol
3. 312 g/mol
4. 342 g/mol
5. 196 g/mol

9. Under what conditions is oxygen gas most soluble in water?

a) high pressure and low temperature

2. high pressure and high temperature

3. low pressure and low temperature
4. low pressure and high temperature

10. What is the boiling point of a solution containing 30.0 g of K₂SO₄ (molar mass = 174.3 g/mol) dissolved in 100.0 g of water, assuming complete dissociation? (K_b for H₂O is 0.51 °C kg/mol.)

a)   101.8°C

b) 102.6°C c)   106.1°C

d)   103.5°C

e)   100.8°C

11. The lattice energy of NaI is 686 kJ/mol and its heat of solution is -7.6 kJ/mol. Calculate the hydration energy of NaI(s).

1. +15.2 kJ/mol
2. -678 kJ/mol

c)                                            -694 kJ/mol

4. +678 kJ/mol
5. +694 kJ/mol

12. When a nonvolatile solute is added to a volatile solvent, the solution vapor pressure                      , the boiling point                                                                                                                                                                               , the freezing point                                                    , and the osmotic pressure across a semipermeable membrane        .

1. decreases, increases, decreases, decreases
2. increases, increases, decreases, increases c)increases, decreases, increases, decreases

d)                                            decreases, decreases, increases, decreases

e)                                            decreases, increases, decreases, increases

13. A 0.20 M solution of MgSO₄ has an observed osmotic pressure of 7.8 atm at 25° C. Determine the observed van’t Hoff factor for this experiment.

a)1.2

b)                                            1.4

c)                                            1.6

d)                                            1.8

e)2.0

Chapter 16 Kinetics

The formula for deriving rate law: [exp1/exp2]^a[exp1/exp2]^b = rate1/rate2

Useful Information:

ln [A]_o / [A_t] = kt 1^st order                                                            1 / [A]_t – 1 / [A]_o = kt 2^nd

order

t        =     1         

R = 8.314 J/mol.K                                                        ln k₂ / k₁ = ­ E_a /R (1 / T₂ – 1 / T₁)

14. Given the initial rate data for the reaction A + B → C, determine the rate equation for the reaction.

[A] (mol/L)                      [B] (mol/L)                            rate of appearance of C

0.033                                                 0.015         0.010

0.033                                                 0.030         0.020

0.099                                                 0.015         0.090

1. rate = k [A] [B]
2. rate = k [A]²

c)        rate = k [A]²[B]

4. rate = k [A]³[B]²
5. rate = k [A]³[B]

15. Consider the following reaction: N₂(g)      +     3H₂(g) ?              2NH₃(g)

If hydrogen is being consumed at a rate of 0.082 M/s, what is the rate of production of ammonia?

a)      0.014 M/s

b)      0.12 M/s

c)       0.041 M/s

d)   0.055 M/s

e)      0.027 M/s

16. If a reaction is third order with respect to a certain reactant, then doubling the initial concentration of that reactant with everything else remaining the same will cause the initial rate of the reaction to:

1. double
2. triple
3. increase by a factor of 6

d)        increase by a factor of 8

e)        remain the same

17. Consider the gas-phase decomposition of N₂O₅: N₂O₅(g) →                        2NO₂g) + 1/2 O₂(g)

This is known to be a first-order reaction with a half-life of 102 s at 70°C. How long would it take for 35.0% of a sample of N₂O₅ to decompose at 70°C?

a)        63.4 s

2. 154 s
3. 135 s

d)        66.3 s

e)        614 s

18. Consider the reaction

A → products

If a plot of the concentration of A vs. time gives a straight line, then

1. this is a first order reaction.
2. this is a second order reaction.
3. this is a third order reaction.

d)        this is a zero order reaction.

e)        None of the above can be true, since a plot of the concentration of A vs. time can never be a straight line.

19. The reaction:

A ? B + C

is second order in A. When [A]_o = 0.100 M, the reaction is 20.0% complete in 40.0 minutes. What is the half- life of this reaction?

1. 124 min
2. 11.1 min
3. 10.0 min
4. 100. min

e)        160. min

20. Suppose that the hypothetical reaction 2A₂ + B₂ ? 2C proceeds by the following mechanism:

What is the molecularity of step 2, and what is the role of X?

a)        unimolecular, catalyst

b)        bimolecular, intermediate

3. bimolecular, catalyst
4. termolecular, intermediate
5. termolecular, catalyst

21. Increasing the temperature of a reaction results in an increase in the rate of the reaction. Why is this true?

a) At higher temperatures, more of the colliding molecules possess the minimum energy required for a reaction to occur.

2. Increasing the temperature lowers the activation energy.
3. Increasing the temperature lowers the rate constant.
4. Choices a and b are both correct.
5. Choices a, b, and c are all correct.

22. A certain first-order reaction has a rate constant k = 2.1 x 10⁵ s^–1 at 355 K. If the activation energy for this reaction is 135 kJ/mol, calculate the rate constant at 550. K.

a)3.3 x 10⁵

b)            7.2 x 10⁵

c)2.1 x 10⁵

d)            4.9 x 10⁵

e)                                    2.3 x 10¹²

23. Consider the following diagram that corresponds to the reaction W ? Z

What would be the activation energy for the reverse reaction, Z ? W, and would that reaction be endothermic or exothermic?

1. A, exothermic
2. A, endothermic
3. B, exothermic

d)        B, endothermic

e)        C, endothermic

24. The reaction           2H₂O₂ → 2H₂O + O₂        has the following mechanism:

H₂O₂ + I^— → H₂O + IO^—

H₂O₂ + IO^— → H₂O + O₂ + I^—

Which species is the catalyst in this reaction? a)H₂O

b)                                            I^—

3. H₂O₂
4. IO^—
5. This reaction does not have a catalyst.

25. A certain reaction has the experimentally determined rate law:

rate = k [A]² [B]²

What are the units of the rate constant for this reaction? a)L/mol.s

b)                L²/mol².s c)mol/L.s

d)                L³/mol³.s e)mol³/L

Chapter 17 Equilibrium

Useful Information:                               R = 0.08206 L^.atm/mol^.K = 8.314 J/mol^.K K_p = K_c(RT)^{?n gas} ( R= .0821 L atm/molK)

26. Consider the following reaction at a certain temperature:

Ni(s) + 4CO(g) Ni(CO)4(g)

An equilibrium mixture contains 0.50 mol Ni, 0.020 mol CO, and 0.010 mol Ni(CO)4, all in a 2.0 L container. What is the value of Kc for this reaction?

a)        5.0 x 10⁵

b)        2.0 x 10⁶

c)        2.0 x 10³

d)        6.3 x 10⁴

e)        1.3 x 10⁵

27. Consider the following reaction:

CO(g) + 2H2(g) CH3OH(g)                           ?H° = -91 kJ

If the temperature of this reaction is increased, what happens to the magnitude of Kc?

a)        Kc increases.

b)        Kc decreases.

3. Kc remains constant.
4. More information is needed.

28. Determine the value of Kc at 1000°C for the equilibrium below:

C(s) + 2H2(g) CH4(g)                     Kp = 0.262 at 1000°C a)                                                                        2.18 x 10³

b)        2.86 x 10³

c)        2.40 x 10^-5

d)        2.51 x 10^-3

e)        27.4

29. At a particular temperature, Kc = 6.5 x 10² for

2NO(g) + 2H2(g) N2(g) + 2H2O(g)

What is the value of Kc at the same temperature for the reaction below?

2N2(g) + 4H2O(g) 4NO(g) + 4H2(g)

a)        1.5 x 10^-3

b)        4.2 x 10⁵

c)        1.3 x 10³

d)        7.7 x 10^-4

e)        2.4 x 10^-6

30. What is the equilibrium constant expression for Kc for the following reaction?

Na2O2(s) + CO2(g) Na2CO3(s) + ½ O2(g)

1

[Na2CO3 ](

[O2 ])

a)                                                                                                Kc =                         2            

[Na2O2 ][CO2 ]

4. Kc =

[Na2O2 ][CO2 ]

1

[Na2CO3 ]( 2 [O2 ])

[Na CO

][O

]1/ 2

b)                                                                                                Kc =         2       3       2      

[Na2 O2 ][CO2 ]

5. Kc =

[Na2 O2 ][CO2 ]

[O ]1 / 2

c)        Kc =      2       

[CO2 ]

31. Hydrogen sulfide decomposes according to the following reaction, for which Kc = 9.30 x 10^-8 at 700°C: 2 H2S(g)                              2 H2(g) + S2(g)

If 0.50 mol of H2S is placed in a 2.0 L container, what is the equilibrium concentration of H2 at 700°C?

a)        3.6 x 10^-3 M

b)        1.5 x 10^-4 M

c)        3.0 x 10^-4 M

d)        4.5 x 10^-3 M

e)        2.3 x 10^-3 M

32. At 100°C, Kp= 60.6 for the reaction

2NOBr(g)       2NO(g)    + Br2(g)

In a given experiment, 0.10 atm of each gas (NOBr, NO, and Br2) is placed in a container. Which statement below is

correct?

a)        Q < K, and the reaction will proceed in a forward direction.

1. 2. Q < K, and the reaction will proceed in a reverse direction.
   3. Q > K, and the reaction will proceed in a forward direction.
   4. Q > K, and the reaction will proceed in a reverse direction.
   5. Q = K, and the reaction is at equilibrium.

33. 0.100 mol of HI(g) was placed in a 1.00 L container. When the reaction shown below had reached equilibrium at 1000 K, the concentration of H2(g) was found to be 0.0135 mol/L. What is the value of the equilibrium constant Kc at 1000 K?

2HI(g)             H2(g)   +    I2(g)

a)        0.0342

b)        0.00250

c)        0.00182

d)        0.0244

e)        0.156

34. The reaction below, which occurs in automobile engines, is an important contributor to smog. N2(g) + O2(g)   2NO(g)        ?H° = +181 kJ

Which of the following actions will decrease the equilibrium yield of NO(g)?

1. 1. Increasing the total pressure by compression
   2. Decreasing the total pressure by expansion
   3. Increasing the total pressure by addition of He(g)
   4. Increasing the temperature
   5. Decreasing the temperature

a)                                                                                                                                                                                                          IV only

b)        V only

3. I and V
4. II and IV
5. I, III, and V

Chapter 18 Acid-Base Equilibria

Useful Information:

pH = - log[H⁺]           pOH = -log[OH^-]           pH + pOH = 14              [H⁺][OH^-] = 1 X 10^-14

35. What is the best definition of a Brønsted-Lowry base?

1. electron-pair acceptor
2. electron-pair donor

c)        proton acceptor

4. proton donor
5. produces hydroxide ions in aqueous

36. If the concentration of hydroxide ion in a certain solution is 5.8 x 10^-3 M, what is the pH of the solution?

37. What is the pH of 0.035 M HClO4?

38. What is the pH of 0.025 M barium hydroxide?

39. A 0.075 M solution of a certain weak monoprotic acid has a pH of 3.69. What is Ka for this acid? a)          3.2 x 10^-6

b)        5.6 x 10^-7

c)        9.1 x 10^-5

d)        4.4 x 10^-6

e)        3.7 x 10^-5

40. What is the pH of 0.10 M benzoic acid?

(Benzoic acid, C6H5COOH, is a monoprotic acid with Ka = 6.3 x 10^-5.)

41. Trimethylamine, (CH3)3N, is a weak base with Kb = 6.3 x 10^-5. What is the pH of 0.15 M trimethylamine?

42. Carbonic acid, H2CO3, is a diprotic acid with the following Ka values:

Ka1 = 4.5 x 10^-7

Ka2 = 4.7 x 10^-11

43. Phenol, C6H5OH, is a weak acid with Ka = 1.0 x 10^-10. What is the value of Kb for its conjugate base phenoxide, C6H5O^-? a)          2.5 x 10^-5

b)        1.0 x 10^-3

c)        9.0 x 10^-3

d)        9.0 x 10^-4

e)        1.0 x 10^-4

44. What is the pH of a 0.10 M solution of NaCN at 25°C? (Ka = 4.9 x 10^-10 for HCN at 25°C.)

45. Which of the salts below, when dissolved in pure water, will cause the pH to go up?

1. NaI
2. NH4Cl
3. KCN

1. 2 only
2. 2 and 3

c)        3 only

4. 1 and 3
5. 1 and 2

46. Consider the following equilibrium: 2H₂ (g) + X₂ (g)  2H₂X (g)

Increasing the pressure by decreasing the volume will cause:

a)                  the reaction to occur to produce H₂X.

2. the reaction to occur to produce H₂ and X₂.
3. the reaction to occur to produce H₂ but no more X₂.
4. no reaction to occur.
5. X₂ to dissociate.

47. Which of the following is true for a system whose equilibrium constant is relatively small?

1. It will take a short time to reach equilibrium.
2. It will take a long time to reach equilibrium.

c)The equilibrium lies to the left.

4. The equilibrium lies to the right.
5. Two of these.

48. What concentration of acetic acid (K_a = 1.80 × 10^­5) has the same pH as that of 5.00 × 10^­3 M HCl?

a) 3.60 × 10^­3 M

b) 5.00 × 10^­3 M

c)                  1.39 M

d) 5.00 M

e) none of these

49. Which of the following would give the highest pH when dissolved in water to form a 0.10 M

solution?

1. a strong acid
2. a weak acid

c)the potassium salt of a weak acid

4. the potassium salt of a strong acid
5. the ammonium salt of a strong acid

1. KCl, NH₄Cl, HNO₃, KOH, NaC₂H₃O₂
2. HNO₃, KCl, NH₄Cl, KOH, NaC₂H₃O₂
3. NH₄Cl, HNO₃, KCl, KOH, NaC₂H₃O₂

d)                  HNO₃, NH₄Cl, KCl, NaC₂H₃O₂, KOH

e) none of these

Chapter 19 Ionic Equilibria in Aqueous Solution

51. Which equation below best describes the reaction that occurs when 0.01 mol HCl(g) is added to a 1.0 L solution containing

0.10 M nitrous acid (HNO2) and 0.10 M sodium nitrite (NaNO2)?

4. HCl(aq) + HNO2(aq) ? ClNO(aq) + H2O(l)

52. Which of the following will give a buffer solution when equal volumes of the two  are mixed?

0. 1. 0.10 M HNO3 and 0.10 M NaNO3
   2. 0.10 M HC2H3O2 and 0.10 M NaC2H3O2
   3. 0.10 M HC2H3O2 and 0.20 M NaOH
   4. 0.10 M HNO3 and 0.20 M NaC2H3O2
1. II only
2. I and II

c)        II and IV

4. I, II, and IV
5. II, III, and IV

53. Determine the pH of a buffer that is formed by mixing 25 mL of 0.15 M acetic acid, HC2H3O2, with 10 mL of 0.10 M sodium acetate, NaC2H3O2.  (The Ka of acetic acid is 1.8 x 10^-5.)

e)        5.32

54. 100.0 mL of a buffer which is 0.25 M in HCN and 0.25 M in KCN has 20.0 mL of 1.0 M HCl added to it. What is the pH after the HCl has been added? (Ka = 4.9 x 10^-10 for HCN)

a)    9.24

b)   5.64

c)    4.69

d)   9.31

e) 8.36

55. Consider the titration of 50.0 mL of 0.100 M HCl with 0.200 M NaOH. What is the pH after the addition of 20.0 mL of NaOH?

56. Consider the titration of 75.0 mL of 0.200 M HF with 0.15 M KOH. What is the pH after adding 75.0 mL of KOH? (Ka for HF is 3.5 x 10^-4.)

a)    2.98

b)   3.58

c)    3.46

d) 3.93

e)    3.33

57. Consider the titration of 50.0 mL of 0.30 M HC2H3O2 (Ka = 1.8 x 10^-5) with 0.20 M NaOH. What is the pH at the equivalence point?

58. A certain saturated solution of copper(II) phosphate, Cu3(PO4)2, has the following concentration of ions:

[Cu²⁺] = 5.0 x 10^-8 M

What is the value of Ksp for copper(II) phosphate?

a)        1.7 x 10^-15

b)        9.8 x 10^-29

c)        1.7 x 10^-8

d)        1.4 x 10^-37

e)        9.0 x 10^-38

59. Ksp for PbF2 is 3.6 x 10^-8. What is the molar solubility of PbF2 in pure water?

a) 2.1 x 10^-3 M

b)   1.3 x 10^-4 M

c)    1.9 x 10^-4 M

d)   2.6 x 10^-3 M

e)    3.6 x 10^-8 M

60. What is the molar solubility of CaF2 in 0.050 M Ca(NO3)2? (Ksp for CaF2 is 3.9 x 10^–11.)

a)        2.1 x 10^-4 M

b)        1.4 x 10^-5 M

c)        1.4 x 10^-6 M

d)        2.8 x 10^-5 M

e)        6.2 x 10^-6 M

61. Ksp for silver sulfate, Ag2SO4, is 1.2 x 10^-5. If 100 mL of 0.080 M AgNO3 is added to 100 mL of 0.015 M K2SO4, which statement best describes what happens to the resulting solution?

1. The solution is supersaturated, and a precipitate will form.
2. The solution is supersaturated, but a precipitate will not form.

c) The solution is saturated, and no precipitate will form.

d) The solution is unsaturated, and no precipitate will form.

62. Which of the salts below will become more soluble as the pH is lowered?

1. 1. Ag2CO3
   2. PbCl2
   3. MgF2

1. none of these
2. I

c)                  I and III

4. III
5. I, II, and III

Chapter 20 Thermodynamics

?H_reaction = ∑ H _products ­ ∑ H _reactants ( Also works for S and G)

?G = ? H ­ T?S                        ?G = ­RT lnK ( R = 8.314 J/mol.K)

?G = ?G^o + RT lnQ

63. Which of the following statements is true?

increase.

1. In any spontaneous process, the entropy of the system always increases.
2. All spontaneous processes are exothermic.
3. Any spontaneous process is accompanied by a positive free energy change.
4. In order for a process to be spontaneous, the process must be exothermic and the entropy of the system must

e)None of the statements above are true.

64. Which of the following reactions involve(s) an increase in the entropy of the system?
    1. CaCO3(s) → CaO(s)  + CO2(g)
    2. 2Hg(s) + O2(g) → 2HgO(s)
    3. C6H12O6(s) → 2C2H5OH(l) + 2CO2(g)
    4. CS2(l) → CS2(g)

1. I and II
2. II only
3. III and IV
4. I and IV

e) I, III, and IV

65. Under what conditions is the absolute entropy of a pure perfectly ordered crystalline substance equal to zero?

1. at 25°C and 1 atm pressure
2. at 0°C and 1 atm pressure

c) at 0 K

4. under any conditions where two phases are present in equilibrium
5. when ?H°f and ?G°f are both zero

66. Consider the thermochemical equation below:

NH3(g) + HCl(g) → NH4Cl(s)     ?H° = -176.0 kJ

When is this reaction expected to be spontaneous?

a)    It is spontaneous only at relatively high temperatures.

b) It is spontaneous only at relatively low temperatures.

3. It is always spontaneous at any temperature.
4. It is never spontaneous at any temperature.

67. What are the signs of ?H, ?S, and ?G for freezing water at 25°C?

68. Determine the normal boiling point of chloroform (CHCl3), given that its enthalpy of vaporization (?H°vap) is 29.24 kJ/mol, and its entropy of vaporization is 87.5 J/(K^.mol).

a)    -75°C

b)   122°C

3. 15°C
4. 47°C

e) 61°C

69. Use the table of thermodynamic data below to determine the value of ?G° at 25 °C for the reaction: N2O(g) + NO2(g) → 3NO(g)

a)        -50,900 kJ

b)        105 kJ

c)        -4130 kJ

d)        -27.3 kJ

5. 237 kJ

70. Consider the following reaction:

NO(g) + O3(g)  →   NO2(g) + O2(g)                ?G° = -198 kJ at 25°C

What is the value of ?G at 25°C when each gas is at the partial pressure specified below?

1.0 x 10^-6 atm NO

1.0 x 10^-7 atm O3

3.0 x 10^-6 atm NO2

0.20 atm O2

a) -159 kJ

2. -167 kJ
3. -198 kJ
4. -236 kJ

e)    +32 kJ

71. Given the following free energies of formation at 25°C:

?G°f C2H2(g)                     209.2 kJ/mol

C2H6(g)      -32.9 kJ/mol

Determine Kp at 25°C for the reaction:    C2H2(g) + 2H2(g) C2H6(g) a)                                     9.07 x 10^-1

b)        97.2

c)        1.24 x 10³¹

d)        2.72 x 10⁴²

5. Not enough information is given.

Chapter 21 Electrochemistry

Ecell = E oxidation + E reduction

it = nFe       i = amps        t= seconds n= moles an element

e= moles of electrons transferred in the reaction 1 C = J/V

Reduction Potentials at 25°C                                            E° (V)                                                                                                E°(V)

F₂(g) + 2 e^–

?

2 F^–(aq)

2.87

H₂O₂(aq) + 2 H⁺(aq) +  2

e–

?2 H₂O(l) 1.78

5 e^–

?

Mn²⁺(aq) + 4 H₂O(l)

1.51

Cl₂(g) + 2 e^–

?

2 Cl^–(aq)

1.36

6 e^–

?

2 Cr³⁺(aq) + 7 H₂O(l)

1.33

O₂(g) + 4 H⁺(aq) + 4 e^–

?

2 H₂O(l)

1.23

Br₂(l) + 2 e^–

?

2 Br^–(aq)

Sn⁴⁺(aq) + 2 e^–

? Sn²⁺(aq) 0.15

2 H⁺(aq) + 2 e^–

? H₂(g) 0.00

Pb²⁺(aq) + 2 e^–

?

Pb(s)

–0.13

Ni²⁺(aq) + 2 e^–

?

Ni(s)

–0.26

Cd²⁺(aq) + 2 e^–

?

Cd(s)

–0.40

Fe²⁺(aq) + 2 e^–

?

Fe(s)

–0.45

Zn²⁺(aq) + 2 e^–

?

Zn(s)

–0.76

2 H₂O(l) + 2 e^–

Useful Information:                            R = 0.08206 L^.atm/mol^.K = 8.314 J/mol^.K

1 F  =  96,485 C/mol e^–

72. Which of the reactions below is a redox reaction?
    1. K2CrO4(aq) + BaCl2(aq) → BaCrO4(s) + 2KCl(aq)
    2. 2HNO3(aq) + Ba(OH)2(aq) → Ba(NO3)2(aq) + 2H2O(l)
    3. Cu(s) + S(s) → CuS(s)

1. I
2. II

c)        III

4. I and III
5. none of these

73. Which of the species below cannot function as an oxidizing agent?

1. CrO₃
2. I₂
3. Cu²⁺
4.  

e) Ni

74. What is the oxidation number of C in sodium oxalate, Na2C2O4?

75. Which element can have both positive and negative oxidation states?

1. Be
2. Cu
3. Ne
4. F

e)          P

76. Which element is being reduced in the following reaction?

2H2SO4(aq) + 2NaBr(s) → Br2(l) + SO2(g) + Na2SO4(aq) + 2H2O(l)

a)          H

b)          S

3. O
4. Na
5. Br

77. Which substance is the oxidizing agent in the reaction below?

H2O(l)

1. Zn(s)
2. H⁺(aq)

4. Zn²⁺(aq)
5. N2O(g)

78. Balance the following redox reaction, which occurs in acidic solution, using the smallest whole- number coefficients.

When this equation is appropriately balanced, what is the coefficient on water, and on which side of the equation does water appear?

1. 1, product side
2. 4, reactant side
3. 3, reactant side

d)          7, reactant side

e)          4, product side

79. Balance the following redox reaction, which occurs in basic solution, using the smallest whole- number coefficients:

When this equation is appropriately balanced, what is the coefficient on water, and on which side of the equation does water appear?

1. 6, on the right side
2. 2, on the right side
3. 3, on the left side

d)        3, on the right side

e)        9, on the right side

80. Which of the following is true of the redox reaction in a galvanic cell?

a)        ?G > 0,  E < 0,  and  K < 1

b)        ?G < 0,  E > 0,  and  K > 1

c)        ?G < 0,  E < 0,  and  K < 1

d)        ?G > 0,  E > 0,  and  K > 1

e)        ?G < 0,  E > 0,  and  K < 1

81. Which of the following is the best oxidizing agent under standard conditions at 25°C?

1. Sn⁴⁺ (aq)
2. I^– (aq)

c)          Fe³⁺ (aq)

4. Cd²⁺ (aq)
5. Br^– (aq)

82. Which of the following is the best reducing agent under standard conditions at 25°C?

a)          Sn⁴⁺ (aq)

b)          I^– (aq)

3. Fe³⁺ (aq)
4. Cd²⁺ (aq)
5. Br^– (aq)

83. Which of the following transformations could take place at the cathode of an electrochemical cell?

e)          NO →   HNO2

84. A galvanic cell is constructed using a Ni(s) / Ni²⁺(aq) half-cell and a Cd(s) / Cd²⁺(aq). The two  are connected with a salt bridge, and the Ni(s) electrode is connected to the Cd(s) electrode with a wire. Which statement below is correct?

a)          The electrons move through the wire towards the nickel electrode, which is the anode.

b)          The electrons move through the wire towards the nickel electrode, which is the cathode.

3. The electrons move through the wire towards the cadmium electrode, which is the anode.
4. The electrons move through the wire towards the cadmium electrode, which is the cathode.
5. Electrons will not move through the wire, because this is not a spontaneous reaction.

85. Consider a galvanic cell constructed by connecting a standard Zn/Zn²⁺ half cell to a standard Cu/Cu²⁺ half cell. Which of the following statements is false?

electrode.

1. The concentration of Cu²⁺(aq) decreases during discharge.
2. The copper electrode is the cathode.
3. Electrons flow through the external circuit from the zinc electrode to the copper

4. Reduction occurs at the copper electrode during discharge.

e)        The mass of the zinc electrode increases during discharge.

86. What is the cell potential for the following galvanic cell at standard conditions and 25°C?

Ag(s) | Ag⁺(aq) || Br2(l) | Br^–(aq) | Pt(s)

a)    1.89 V

b)   2.69 V

c) 0.29 V

d) -0.51 V

e)    0.51 V

87. When a Ag/Ag⁺ half-cell is connected to a Co/Co²⁺ half-cell, the resulting galvanic cell has the following overall reaction:

2Ag⁺(aq) + Co(s) → 2Ag(s) + Co²⁺(aq)

The standard potential for this galvanic cell is 1.08 V. What is the standard reduction potential for Co²⁺ at 25°C?

a)        -0.28 V

b)        +0.28 V

c)        -1.88 V

d)        +1.88 V

e)        +0.52 V

88. A certain galvanic cell is constructed by connecting an Fe³⁺/Fe²⁺ half cell with an I2/I^– half cell. What is the potential of the cell at 25°C, given the following ion concentrations?

[Fe³⁺] = 1.5 M [Fe²⁺] = 0.10 M [I^–] = 1.5 M

a)          0.31 V

b)          0.091 V

c)          0.24 V

d)          0.51 V

e)          0.17 V

89. A concentration cell can be prepared using a half cell consisting of a Cu(s) electrode in a 0.30 M solution of Cu²⁺(aq) and a half cell consisting of a Cu(s) electrode in a 0.10 M solution of Cu²⁺(aq). What is the cell potential at 25°C and which electrode is the cathode?

a)  0.014 V, Cu(s) electrode in the 0.30 M solution of Cu²⁺(aq)

1. 2. 0.028 V, Cu(s) electrode in the 0.30 M solution of Cu²⁺(aq)
   3. –0.014 V, Cu(s) electrode in the 0.10 M solution of Cu²⁺(aq)
   4. –0.028 V, Cu(s) electrode in the 0.10 M solution of Cu²⁺(aq)
   5. 0.014 V, Cu(s) electrode in the 0.10 M solution of Cu²⁺(aq)

90. What is the standard free energy change, ?G°, for the reaction below at 25°C?

2MnO4–(a) + 10 Cl^–(aq) + 16H⁺(aq) → 2Mn²⁺(aq) + 5Cl2(g) + 8H2O(l)

1. -280 kJ
2. -730 kJ
3. -72 kJ
4. -29 kJ

e) -140 kJ

91. Use the table of standard reduction potentials to determine the equilibrium constant for the following reaction at 25°C:

2Fe³⁺(aq)   +    H2O2(aq) → 2Fe²⁺(aq)                    +   O2(g) +             2H⁺(aq)

a) 2.3 x 10²

b) 7.6 x 10⁴⁹

c)    15

d) 1.3 x 10³⁴

e)    0.030

92. Given the reaction

Determine the number of electrons involved in this reaction.

a)            10

2. 8
3. 6
4. 4
5. 2

93. What quantity of charge is required to reduce 40.0 g of CrCl3 to chromium metal? a)      2.45 × 10⁴ C

b)            7.31 × 10⁴ C

c)            2.20 × 10⁴ C

d)            9.65 × 10⁴ C

e)            none of these

Chapter 12 Intermolecular Attractions

1. Intermolecular attractions from strongest to weakest
   1. Ion­Ion­ Ionic compounds (metal/non­metal)
   2. Hydrogen Bonding­ Compounds containing H attached to N, O, or F
   3. Dipole­Dipole­ polar compounds
   4. London Dispersion Forces­ All compounds exhibit these, however they are most important with non­polar compounds.

2. Questions pertaining to intermolecular attractions:
   1. Which compound has the highest boiling point, melting point, and highest heat of vaporization corresponds to     the compound with the strongest intermolecular attractions.

1. 2. Which compound has the highest vapor pressure corresponds to lowest intermolecular attractions.

3. If two compounds are both non­polar, the compound with the greatest molecular mass has the greater London Dispersion Forces.

4. If two compounds are both ionic, the compound which contains the ions with the greatest charge has the greater intermolecular attraction. If ions have the same charge, the smallest ions have the greatest intermolecular attractions.

Free Response:

Substance                                           Melting Point ^oC

H₂                                                           ­259

C₃H₈                                                       ­190

HF                                                          ­92

CsI                                                           621

LiF                                                           870

SiC                                                          >2000

1. Discuss how the trend in melting points of the substances tabulated above can be explained in terms of the types of attractive forces and / or bonds in these substances.

2. For any pairs of substances that have the same kind(s) of attractive forces and/or bonds, discuss the factors that cause variations in the strengths of the forces and /or bonds.

Practice Test Chapter 12

94. Which one of the following would have hydrogen bonding as one of its intermolecular attractions?

(A) CH₄ (B) HCl (C) CH₃OH (D) H₂ ( means answer is C)

95. Which one of the following has dispersion forces as the only intermolecular attraction?

(A) CH₃OH (B) NH₃ (C) H₂S (D) Kr

96. The substance with the greatest enthalpy(heat) of vaporization is:

(A) I₂ (B) Br₂ (C) Cl₂ (D) F₂

97. Which one of the following would have the highest boiling point?

(A) H₂O (B) CO₂ (C) CH₄ (D) Kr

98. Which has the highest boiling point?

(A) N₂ (B) Br₂(C) H₂ (D) Cl₂

99. Which of the following decreases as the strength of the intermolecular forces increases?

1. heat of vaporization
2. boiling point
3. extent of deviation from the ideal gas law

(D) vapor pressure of a liquid

100. Given the following boiling points: (A) water, 100 (B) methanol, 64.9 (C) ethanol, 78.5 (D) ethylene glycol, 198. Which one of the above liquids has the highest vapor pressure at room temperature?

101. The boiling point of the halogens increases going from F₂ to I₂. What type of intermolecular forces are responsible for this trend?

(A) Permanent dipole (B) Hydrogen bonding (C) Ion­ion interaction

(D) London dispersion forces (E) ion­dipole forces

102. Which of the following boils at the lowest temperature?

(A) CF₄ (B) HF (C) HCl (D) KI (E) SiF₄

103. Which has the lowest boiling point?

(A) C₂H₅OH (B) H₂O (C) NH₃ (D) CHCl₃ (E) CH₄

104. Which of the following has the highest boiling point?

(A) CF₄ (B) CCl₄ (C) CBr₄ (D) CI₄

105. On a relative basis, the weaker the intermolecular forces in a substance:

1. the greater the heat of vaporization
2. the more it deviates from ideal gas behavior

(C) the greater is its vapor pressure

(D) the higher is its melting point

106. At room temperature, CsF is expected to be:

1. gas (B) conducting solid (C) liquid (D) brittle solid (E) soft solid

107. Which of the following statements is incorrect?

(A) molecular solids have high melting points

2. the binding forces in a molecular solid include London dispersion forces

(C) molten ionic solids are non­conducting

(D) all statements are correct

108. Which has the least H­bonding?

(A) NH₃ (B) H₂O (C) HF (D) CH₄

109. A certain solid that is very hard, has a high melting point, and is non­conducting unless melted is most likely:

1. I₂ (B) NaCl (C) CO₂ (D) H₂O (E) Cu

110. The triple point of a substance can be defined as:

(A) the point at which solid, liquid, and vapor are all in equilibrium

2. the point at which vapor pressure of a solid is 1 atm
3. the point at which liquid first starts to condense as the temperature of a vapor is lowered
4. the point at which the density of solid and liquid are equal
5. none of the above