Assignment 6 (A)

CHAPTER 17

ACID-BASE EQUILIBRIA AND SOLUBILITY EQUILIBRIA

1- In which one of the following solutions will acetic acid have the greatest percent ionization?

A.0.1 M CH3COOH

B.0.1 M CH3COOH dissolved in 1.0 M HCl

C.0.1 M CH3COOH plus 0.1 M CH3COONa

D.0.1 M CH3COOH plus 0.2 M CH3COONa

2- Which of the following would increase the Ka for HAc?

A.decrease the pH of the solution

B.add some NaAc

C.add some NaOH

D.add some H2O

E.None of the above ─ the Ka is a constant at constant temperature.

3- Which one of the following salts would form an acidic solution when dissolved in water.

A. NaBr B. NaF C. NaHSO4 D. CaCl2

4- Which one of the following combinations cannot be a buffer solution?

A. HCN and KCN B. NH3 and (NH4)2SO4 C. HNO3 and NaNO3

D. HF and NaFE. HNO2 and NaNO2

5- Calculate the pH of a buffer solution which contains 0.25 M benzoic acid (C6H5CO2H) and 0.15 M sodium benzoate (NaC6H5CO2). Given Ka = 6.5  10 ─ 5.

A. 3.97 B. 4.83 C. 4.19 D. 3.40 E. 4.41

6- 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

7- Calculate the pH of a buffer solution prepared by dissolving 0.20 mol of sodium cyanate (NaCNO) and 1.0 mol of cyanic acid (HCNO) in enough water to make 1.0 liter of solution. Ka(HCNO) = 2.0  10 ─ 4

A. 0 B. 3.0 C. 3.7 D. 4.4 E. 5.0

8- For which type of titration will the pH be basic at the equivalence point?

A. strong acid vs. strong baseB. strong acid vs. weak base

C. weak acid vs. strong baseD. all of the aboveE. none of the above

9- 50.00 mL of 0.10 M HNO2 (nitrous acid) was titrated with 0.10 M KOH solution. After 25.00 mL of KOH solution was added, the pH in the titration flask will be: Given Ka = 4.5  10 ─ 4

A. 2.17 B. 3.35 C. 2.41 D. 1.48 E. 7.00

10- Calculate the pH at the equivalence point for the titration of 0.20 M HCl versus 0.20 M NH3. Kb = 1.8  10 ─ 5

A. 2.87 B. 4.98 C. 5.12 D. 7.00 E. 11.12

11- What is the pH at the equivalence point in the titration of 100 mL of 0.10 M HCl with 0.10 M NaOH?

A. 1.0 B. 6.0 C. 7.0 D. 8.0 E. 13.0

12- What is the pH at the equivalence point in the titration of 100 mL of 0.10 M HCN with 0.10 M NaOH?

A. 3.0 B. 6.0 C. 7.0 D. 11.0 E. 12.0

13- The solubility of lead iodide is 0.064 g/100 mL at 20oC. What is the solubility product for PbI2?

A. 1.1  10 ─ 8B. 3.9  10 ─ 6C. 1.1  10 ─ 11

D. 2.7  10 ─ 12E. 1.4  10 ─ 3

14- The molar solubility of MgCO3 is 1.8  10 ─ 4 mol/L. What is Ksp for this compound?

A. 1.8  10 ─ 4B. 3.6  10 ─ 4C. 1.3  10 ─ 7

D. 3.2  10 ─ 8E. 2.8  10 ─ 14

15- The molar solubility of tin iodide (SnI2) is 1.28  10 ─ 2 mol/L. What is Ksp for this compound?

A. 8.4  10 ─ 6B. 1.28  10 ─ 2C. 4.2  10 ─ 6

D. 1.6  10 ─ 4E. 2.1  10 ─ 6

16- The solubility product for CrF3 is Ksp = 6.6  10 ─ 11. What is the solubility of CrF3 in moles per liter of solution?

A. 1.6  10 ─ 3 MB. 1.3  10 ─ 3 MC. 6.6  10 ─ 11 M

D. 2.2  10 ─ 3 ME. 1.6  10 ─ 6 M

17- The solubility product for BaSO4 is 1.1  10 ─ 10. Calculate the molar solubility of BaSO4.

A. 5.5  10 ─ 11 mol/LB. 1.0  10 ─ 5 mol/LC. 2.1  10 ─ 5 mol/L

D. 1.0  10 ─ 10 mol/LE. 2.2  10 ─ 10 mol/L

18- The Ksp for Ag3PO4 is 1.8  10 ─ 18. Calculate the molar solubility of Ag3PO4.

A. 1.6  10 ─ 5 M B. 2.1  10 ─ 5 MC. 3.7  10 ─ 5 M

D. 7.2  10 ─ 1 ME. 1.8  10 ─ 1 M

19- Which of the following would decrease the Ksp for PbI2?

A.Lower the pH of the solution.

B.Add a solution of Pb(NO3)2.

C.Add a solution of KI.

D.None of the above ─ the Ksp of a compound is constant at constant temperature.

20- Calculate the minimum concentration of Mg2+ that must be added to 0.10 M NaF solution in order to initiate a precipitate of magnesium fluoride. Ksp = 6.9  10 ─ 9

A. 1.4  107 MB. 6.9  10 ─ 8 MC. 6.9  10 ─ 9 M

D. 1.7  10 ─ 7 ME. 6.9  10 ─ 7 M

21- Calculate the minimum concentration of Cr3+ that must be added to 0.095 M NaF solution in order to initiate a precipitate of chromium fluoride. For CrF3, Ksp = 6.6  10 ─ 11

A. 0.023 MB. 0.032 MC. 7.7  10 ─ 8 MD. 2.9  10 ─ 9 ME. 6.9  10 ─ 10 M

22- Will a precipitate form (yes or no) when 10.0 mL of 1.5  10 ─ 3 M Pb(NO3)2 is added to 40.0 mL of 3.0  10 ─ 4 M

Na2SO4? Ksp(PbSO4) = 1.6  10-8

A. no because ion product < KspB. no because ion product > Ksp

C. yes because ion product < KspD. yes because ion product > Ksp

23- Will a precipitate of MgF2 form when 200 mL of 1.9  10 ─ 3 M MgCl2 solution is added to 300 mL of

1.4  10 ─ 2 M NaF? Ksp (MgF2) = 6.9  10 ─ 9

A. yes, Q > KspB. no, Q < KspC. no, Q = KspD. yes, Q < Ksp

24- Will a precipitate form when 20.0 mL of 1.1  10 ─ 3 M Ba(NO3)2 solution is added to 80.0 mL of 8.4  10 ─ 4 M Na2CO3? Answer yes or no, and identify the precipitateif there is one.

A. yes: the precipitate is Ba(NO3)2 B. yes: the precipitate is NaNO3

C. yes: the precipitate is BaCO3D. yes: the precipitate is Na2CO3

25- TRUE-FALSE QUESTIONS

1- In the titration of acetic acid (CH3COOH), the pH at the equivalence point is above 7.0.T

2- The percent dissociation of a weak acid HA is greater in a solution containing the salt NaA, than it is in a solution of the weak acid only. F

3- All indicators are weak acids that are one color in acidic solution and another color in basic solution.F

4- Formic acid, Ka = 1.7  10 ─ 4, is a stronger acid than nitrous acid, Ka = 4.5  10 ─ 4. F

5- For any conjugate acid-base pair, it is true that Kb = KaKw .F

6- Solutions of the following salts have basic pH values, CH3CO2Na, KCN, and Na2CO3. T