Second Major Exam
KFUPM                                                                          CHEM 102
Chemistry Department                                                  Fall 05/06 (Term 052)
                                                                                       April 23, 2006

Name: ____________________ID #: _____________ Section #: _____________



1.
A solution is prepared by mixing 500. mL of 0.10 M NaOCl and 500. mL of 0.20 M HOCl. What is the pH of this solution?   [Ka(HOCl) = 3.2 × 10–8]
A.
4.10
B.
7.00
C.
7.19
D.
7.49
E.
7.80


2.
Calculate the pH of a solution prepared by mixing equal volumes of 0.15 M HNO3 and 0.40 M ammonia, NH3.
Kb(NH3) =1.8 × 10–5.
A.
7.95
B.
8.57
C.
8.98
D.
9.21
E.
9.47


3.
Which one of the following statements is NOT TRUE?
A.
The pH at the equivalence point is 7 when titrating strong acid vs strong base.
B.
The pH at the equivilance point is >7 when titrating weak acid with a strong base.
C.
The solubility of CaF2 in acidic solution is more than that in pure water.
D.
The solubility of AgCl in acidic solution is more than that in pure water.
E.
A ligand is a Lewis base.


4.
Calculate the pH of the solution resulting from the addition of 10.0 mL of 0.10 M NaOH to 50.0 mL of 0.10 M HCN (Ka = 4.9 × 10–10) solution.
A.
5.15
B.
8.71
C.
5.85
D.
9.91
E.
13.0


5.
What mass of sodium fluoride, NAF, must be added to 250. mL of a 0.100 M HF solution to give a buffer solution having a pH of 3.50? (Ka(HF) = 7.1 × 10–4)
A.
0.49 g
B.
1.5 g
C.
3.4 g
D.
2.3 g
E.
0.75 g


6.
For PbCl2 (Ksp = 2.4 × 10–4), will a precipitate of PbCl2 form when 0.10 L of 3.0 × 10–2 M Pb(NO3)2 is added to 400 mL of 9.0 × 10–2 M NaCl?
A.
Yes, because Q > Ksp.
B.
No, because Q < Ksp.
C.
No, because Q = Ksp.
D.
Yes, because Q < Ksp.


7.
Calculate the silver ion concentration in a saturated solution of silver(I) carbonate, Ag2CO3, (Ksp = 8.1 × 10–12).
A.
5.0 × 10–5 M
B.
2.5 × 10–4 M
C.
1.3 × 10–4 M
D.
2.0 × 10–4 M
E.
8.1 × 10–4 M


8.
Which one of the following statements is TRUE?
A.
The anode is the electrode at which reduction takes place.
B.
In a galvanic cell, elctrons flow from the anode to the cathode.
C.
The potential of a hydrogen electode with H+ of 0.10M and H2 gas of 1.0 atm is 0 V.
D.
Salt bridge is used to allow free passage for electrons.
E.
The cathode is the electrode at which oxidation takes place.


9.
Calculate the minimum concentration of Mg2+ that must be added to 0.10 M NaF in order to initiate a precipitate of magnesium fluoride. (For MgF2 , Ksp = 6.9 × 10–9.)
A.
1.4 × 107 M
B.
6.9 × 10–9 M
C.
6.9 × 10–8 M
D.
1.7 × 10–7 M
E.
6.9 × 10–7 M


10.
Considering the following processes, which one of the choices below includes the 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) + (1/2)O2(g)
A.
1, 2, 3, 4
B.
1, 2
C.
2, 3, 4
D.
3, 4
E.
1, 4


11.
HI has a normal boiling point of –35.4°C, and its ΔHvap is 21.16 kJ/mol. Calculate the molar entropy of vaporization (ΔSvap).
A.
598 J/K·mol
B.
68.6 J/K·mol
C.
75.2 J/K·mol
D.
0.068 J/K·mol
E.
89.0 J/K·mol


12.
Considering the following reaction:
HgO(s) → Hg(l) + (1/2)O2(g),              ΔH° = 90.84 kJ/mol.
estimate the temperature at which this reaction will become spontaneous under standard state conditions.
    S°(Hg) = 76.02 J/K·mol
    S°(O2) = 205.0 J/K·mol
    S°(HgO) = 70.29 J/K·mol
A.
108 K
B.
430 K
C.
620 K
D.
775 K
E.
840 K


13.
For the reaction H2(g) + S(s) → H2S(g), ΔH° = –20.2 kJ/mol and ΔS° = +43.1 J/K·mol. Which one of the following statements is TRUE?
A.
The reaction is only spontaneous at low temperatures.
B.
The reaction is spontaneous at all temperatures.
C.
ΔG° becomes more positive 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.
At 1500°C, the equilibrium constant for the reaction CO(g) + 2H2(g) CH3OH(g) has the value
Kp = 1.4 × 10–7. Calculate ΔG° for this reaction at 1500°C. (R= 8.31 J/mol.K).
A.
105 kJ/mol
B.
1.07 kJ/mol
C.
–233 kJ/mol
D.
–105 kJ/mol
E.
233 kJ/mol


15.
Given the following notation for an electrochemical cell
    Pt(s) | H2(g) | H+(aq) || Ag+(aq) | Ag(s),
what is the balanced overall (net) cell reaction?
A.
2H+(aq) + 2Ag+(aq) → H2(g) + 2Ag(s)
B.
H2(g) + 2Ag(s) → H+(aq) + 2Ag+(aq)
C.
2H+(aq) + 2Ag(s) → H2(g) + 2Ag+(aq)
D.
H2(g) + Ag+(aq) → H+(aq) + Ag(s)
E.
H2(g) + 2Ag+(aq) → 2H+(aq) + 2Ag(s)


16.
A certain electrochemical cell has the following cell reaction:
    Zn + HgO → ZnO + Hg
Which is the half-reaction occurring at the anode?
A.
HgO + 2e– → Hg + O2–
B.
Zn2+ + 2e– → Zn
C.
Zn → Zn2+ + 2e–
D.
Hg + 2e–   → HgO
E.
ZnO + 2e– → Zn


17.
Calculate the value of E°cell for the following reaction:
    2Au(s) + 3Ca2+(aq) → 2Au3+(aq) + 3Ca(s)
Given that:
Au3+ + 3e- → Au,         Eo = 1.53 V
Ca2+ + 2e-→ Ca,          Eo = -2.84 V
A.
–4.37 V
B.
–1.37 V
C.
–11.6 V
D.
1.37 V
E.
4.37 V


18.
Consider the following standard reduction potentials in acid solution:
 
E°(V)
Al3+ + 3e– → Al(s)
–1.66
AgBr(s) + e– → Ag(s) + Br–
+0.07
Sn4+ + 2e– → Sn2+
+0.14
Fe3+ + e– → Fe2+
+0.77

The strongest reducing agent among those shown above is
A.
Fe3+.
B.
Fe2+.
C.
Br–.
D.
Al3+.
E.
Al.


19.
Calculate the cell voltage for the following redox reaction:
    Cu2+ (0.010 M) + H2(1 atm) → Cu(s) + 2H+( pH = 7.0)
    Cu2+ + 2e- →   Cu,        Eo = 0.34 V
A.
0.19 V
B.
–0.01 V
C.
0.34 V
D.
0.69 V
E.
0.49 V


20.
The overall formation constant for HgI42- is 1.0X1030. That is

                             
What is the concentration of Hg2+ in 500 mL of a solution that was originally 0.010M Hg2+ and 0.78M I-? The reaction is
         Hg2+ (aq)   + 4I- (aq) Δ HgI42- (aq)
A.
2.8X10-32
B.
0 M
C.
3.3X10-32
D.
1.0X10-32
E.
1.7X10-29


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