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CHEM 1032 – PRACTICE EXAM III – CLASS – SPRING 2017
(You may need a periodic table. Useful information appears on page 4.)
Deuterium, 21 D, is an isotope of hydrogen, 11 H. Deuterium oxide, D2O, commonly known as heavy water
(HW), has: 2 D2O ⇌ D3O+ + OD– where KHW = 8.9 x 10–16 and pD = –log[D3O+]. Use this information to
answer questions 1 – 5.
1. How many electrons are present in the D+ cation?
A. 0
B. 1
C. 2
D. 3
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2. When a pure sample of D2O is at equilibrium, which of the following is true?
A. [D3O+] < [OD–]
B. [D3O+] > [OD–]
C. [D3O+] = [OD–]
D. not enough info.
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C. 7.00
D. 6.48
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C. 2.7 x 10–9 M
D. 3.0 x 10–8 M
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D. 1.1 x 10–15
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3. What is the value of pD in pure D2O?
A. 15.05
B. 7.52
4. What is the value of [OD–] in pure D2O?
A. 3.3 x 10–7 M
B. 1.0 x 10–7 M
5. What is the value of the equilibrium constant for: D3O+ + OD– ⇌ 2 D2O ?
A. –8.9 x 10–16
B. 1.1 x 1015
C. 0
The following statements are either true = answer A, or false = answer B.
6. Every Brønsted-Lowry acid is also a Lewis acid.
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7. Every Arrhenius base is also a Lewis base.
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8. Every Brønsted-Lowry acid is also an Arrhenius acid.
____
9. Conjugate acids of weak bases produce more acidic solutions than conjugate acids of
strong bases.
____
10. The percent ionization of a weak acid in water increases as the concentration of the acid
decreases.
____
11. Aqueous solutions of Fe(NO3)3 have a pH of 7.00.
____
12. Aqueous solutions of NaHSO4 have a pH less than 7.00.
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13. Aqueous solutions of KHCO3 have a pH greater than 7.00.
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14. The pH of a 0.100 M solution of HC2H3O2(aq) is 1.00.
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15. The pH of a 0.100 M solution of NaC2H3O2(aq) is greater than 7.00.
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16. For the acid-base reaction: H2PO4–(aq) + HCO3–(aq) ⇌ HPO42–(aq) + H2CO3(aq), which of the following is
correct?
A. acid = HCO3– conjugate base = H2CO3
base = H2PO4– conjugate acid = HPO42–
B. acid = H2PO4– conjugate base = HCO3–
base = HCO3– conjugate acid = HPO42–
C. acid = HPO42– conjugate base = H2PO4–
base = H2CO3 conjugate acid = HCO3–
D. acid = H2PO4– conjugate base = HPO42–
base = HCO3– conjugate acid = H2CO3
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17. For the reaction: Zn2+(aq) + 4 NH3(aq)  [Zn(NH3)4]2+(aq), which is true?
A. Zn2+ is acting as a Brønsted-Lowry base
C. Zn2+ is acting as a Lewis acid
B. NH3 is acting as a Lewis acid
D. NH3 is acting as a Brønsted-Lowry base
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18. You are given 50.0 mL of a 0.100 M solution of HCl(aq) and 50.0 mL of a 1.00 x 10–4 M solution of
HNO3(aq). You then mix these two acids together. What will be the pH of the mixture?
A. 5.00
B. 4.00
C. 2.50
D. 1.30
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Citric acid is a triprotic acid with the formula H3C6H5O7, where Ka1 = 7.5 x 10–4, Ka2 = 1.7 x 10–5, and
Ka3 = 4.0 x 10–7. Sodium salts of this acid are: Na3C6H5O7, Na2HC6H5O7, and NaH2C6H5O7. Use this
information to answer questions 19 – 23.
19. You are asked to prepare a citrate buffer of pH = 5.00. What acid-base combination would you use?
A. acid = H3C6H5O7 base = Na3C6H5O7
C. acid = H3C6H5O7 base = NaH2C6H5O7
B. acid = NaH2C6H5O7 base = Na2HC6H5O7
D. acid = Na2HC6H5O7 base = Na3C6H5O7
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20. If the pH 5.00 buffer prepared has an acid molarity of 0.100 M, what molarity of base must be used?
A. 0.170 M
B. 1.70 M
C. 5.89 x 10–2 M
D. 1.00 x 10–13 M
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21. The base used to prepare the buffer of pH 5.00 is a solid and is added to 100.0 mL of the 0.100 M
acid. How much base is required? (Atomic masses: Na = 23.0, H = 1.01, C = 12.0, O = 16.0 g/mol).
A. 3.3 g
B. 4.4 g
C. 4.0 g
D. 3.6 g
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22. If 100.0 mL of pure water is added to 100.0 mL of this buffer of pH 5.00, what will be the
approximate pH of the resulting solution?
A. 7
B. 6
C. 5
D. 4
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D. 9.70
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23. What is the pH of a 0.100 M solution of Na3C6H5O7?
A. 4.30
B. 5.94
C. 8.06
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pH
A 0.0962 g sample of an acid of molar mass = 150.0 g/mol was dissolved in 30.0 mL of water and then
titrated with 0.0610 M NaOH solution. The pH of the mixture was monitored during the titration
resulting in the curve shown below. Use this information to answer questions 24 – 27.
32.0
31.0
30.0
29.0
28.0
27.0
26.0
25.0
24.0
23.0
22.0
21.0
20.0
19.0
18.0
17.0
16.0
15.0
14.0
13.0
12.0
11.0
10.0
9.0
8.0
7.0
6.0
5.0
4.0
3.0
2.0
1.0
0.0
25.00
24.00
23.00
22.00
21.00
20.00
19.00
18.00
17.00
16.00
15.00
14.00
13.00
12.00
11.00
10.00
9.00
8.00
7.00
6.00
5.00
4.00
3.00
2.00
1.00
0.00
Volume NaOH (mL)
24. The acid being titrated could be:
A monoprotic or diprotic
C. monoprotic or triprotic
B. diprotic or triprotic
D. diprotic
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25. What are the approximate values of the 1st and 2nd equivalence points?
A. 10.0 mL and 20.5 mL
C. 11.5 mL and 21.5 mL
B. 10.5 mL and 21.0 mL
D. 9.0 mL and 19.0 mL
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C. 4.8 and 8.5
D. 3.0 and 23.0
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D. 8.00
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26. What are the approximate values of pKa1 and pKa2?
A. 8.5 and 18.5
B. 4.8 and 13.6
27. Calculate the pH of the solution at the first equivalence point.
A. 8.75
B. 8.50
C. 8.25
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Nickel(II) hydroxide, has a Ksp value of 5.5 x 10–16. Use this information to answer questions 28 –32.
28. What is the chemical formula of nickel(II) hydroxide?
A. Ni·2H2O
B. Ni2O
C. Ni(OH)2
D. Ni3(OH)2
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D. 1.7 x 10–8 M
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D. 6.4
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29. What is the molar solubility of nickel(II) hydroxide?
A. 1.2 x 10–8 M
B. 6.5 x 10–6 M
C. 5.2 x 10–6 M
30. What is the approximate pH of this saturated solution?
A. 9.0
B. 8.7
C. 7.0
31. Which of the following, when added to a saturated solution of nickel(II) hydroxide, will
increase its solubility?
A. H2O(l)
B. NaOH(aq)
C. NiCl2(aq)
D. HCl(aq)
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32. In order to prepare a saturated solution of nickel(II) hydroxide, which of the following
two reagents would you mix together?
A. Ni(NO3)2(aq) and H2O(l)
C. Ni(s) and NaOH(aq)
B. NiO(s) and H2(g)
D. NiCl2(aq) and CaO(aq)
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Useful information:
Kw = [H3O+][OH–] = 1.0 x 10–14
pH = –log[H3O+]
In pure water @ 25 C: [H3O+] = [OH–] = 1.00 x 10–7
 [base] 

pH = pKa + log 
 [acid] 
pOH = –log[OH–]
pH + pOH = 14.00
KaKb = Kw
pKa + pKb = 14.00