Chapter 17 Practice Test

Chapter 17 Practice Test 1. Which one of the following is a buffer solution? A. 0.40 M HCN and 0.10 KCN

B. 0.20 M CH3COOH

C. 1.0 M HNO3 and 1.0 M NaNO3 D. 0.10 M KCN E. 0.50 M HCl and 0.10 NaCl

-4 -5 2. Which of the following is the most acidic solution? (Ka HNO2 = 4.5 x 10 and Ka CH3COOH = 1.8 x 10 )

A. 0.10 M CH3COOH and 0.10 M CH3COONa

B. 0.10 M CH3COOH

C. 0.10 M HNO2

D. 0.10 M HNO2 and 0.10 M NaNO2

E. 0.10 M CH3COONa

3. 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? 8 [Ka(HOCl) = 3.2  10 ] A. 4.10 B. 7.00 C. 7.19 D. 7.49 E. 7.80

4. Calculate the pH of a buffer solution that contains 0.25 M benzoic acid (C6H5CO2H) and 0.15 M sodium benzoate 5 (C6H5COONa). [Ka = 6.5  10 for benzoic acid] A. 3.97 B. 4.83 C. 4.19 D. 3.40 E. 4.41

5. You are asked to go into the lab and prepare an acetic acid - sodium acetate buffer solution with a pH of 4.00 ± 0.02. What

molar ratio of CH3COOH to CH3COONa should be used? A. 0.18 B. 0.84 C. 1.19 D. 5.50 E. 0.10

6. What is the net ionic equation for the reaction that occurs when small amounts of hydrochloric acid are added to a HOCl/NaOCl buffer solution? + + A. H + H2O  H3O B. H+ + OCl  HOCl C. HOCl  H+ + OCl + + D. H + HOCl  H2OCl

E. HCl + HOCl  H2O + Cl2

7. Over what range of pH is a HOCl  NaOCl buffer effective? (Ka HOCl = 3.2 x 10-8) A. pH 2.0 – pH 4.0 B. pH 7.5 – pH 9.5 C. pH 6.5 – pH 8.5 D. pH 6.5 – pH 9.5 E. pH 1.0 – pH 14.0

8. Assuming equal concentrations of conjugate base and acid, which one of the following mixtures is suitable for making a buffer solution with an optimum pH of 9.29.3? 5 A. CH3COONa / CH3COOH (Ka = 1.8  10 ) 10 B. NH3 / NH4Cl (Ka = 5.6  10 ) 8 C. NaOCl / HOCl (Ka = 3.2  10 ) 4 D. NaNO2 / HNO2 (Ka = 4.5  10 ) E. NaCl / HCl

9. You have 500.0 mL of a buffer solution containing 0.30 M acetic acid (CH3COOH) and 0.20 M sodium acetate (CH3COONa). 5 What will the pH of this solution be after the addition of 20.0 mL of 1.00 M NaOH solution? [Ka = 1.8  10 ] A. 4.65 B. 4.71 C. 4.56 D. 4.84 E. 5.07

10. In which one of the following solutions will acetic acid have the greatest percent ionization?

A. 0.1 M CH3COOH D. 0.1 M CH3COOH plus 0.1 M CH3COONa

B. 0.1 M CH3COOH dissolved in 0.1 M HCl E. 0.1 M CH3COOH plus 0.2 M CH3COONa

C. 0.1 M CH3COOH dissolved in 0.2 M HCl 11. The pH at the equivalence point of a titration may differ from 7.0 due to A. the initial concentration of the solution. D. the initial pH of the unknown. B. the indicator used. E. hydrolysis of the salt formed.

C. the self-ionization of H2O.

12. For which type of titration will the pH be basic at the equivalence point? A. Strong acid vs. strong base. D. All of the above. B. Strong acid vs. weak base. E. None of the above. C. Weak acid vs. strong base.

5 13. Calculate the pH at the equivalence point for the titration of 0.20 M HCl with 0.20 M NH3 (Kb = 1.8  10 ). A. 2.87 B. 4.98 C. 5.12 D. 7.00 E. 11.12

6 14. Methyl red is a common acid-base indicator. It has a Ka equal to 6.3  10 . Its un-ionized form is red and its anionic form is yellow. What color would a methyl red solution have at pH = 7.8? A. green B. red C. blue D. yellow E. violet

–4 2 2 16. For PbCl2 (Ksp=2.4  10 ), will a precipitate form if 0.10 L of 3.0  10 M Pb(NO3)2 is added to 400mL of 9.0  10 M NaCl?

A. Yes, because Q > Ksp. C. No, because Q = Ksp.

B. No, because Q < Ksp. D. Yes, because Q < Ksp.

4 17. The molar solubility of magnesium carbonate is 1.8  10 mol/L. What is Ksp for this compound? A. 1.8  104 B. 3.6  104 C. 1.3  107 D. 3.2  108 E. 2.8  1014

o 18. The solubility of strontium carbonate is 0.0011 g/100 mL at 20 C. Calculate the Ksp value for this compound. A. 7.5  105 B. 1.5  104 C. 5.6  109 D. 7.5  106 E. 1.5  103

11 19. The solubility product for chromium(III) fluoride is Ksp = 6.6  10 . What is the molar solubility of chromium(III) fluoride? A. 1.6  103 M B. 1.2  103 M C. 6.6  1011 M D. 2.2  103 M E. 1.6  106 M

12 20. Calculate the silver ion concentration in a saturated solution of silver(I) carbonate (Ksp = 8.1  10 ). 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

21. Which of the following would decrease the Ksp for PbI2? ` A. Lowering the pH of the solution C. Adding a solution of KI

B. Adding a solution of Pb(NO3)2 D. None of the above

22. Calculate the minimum concentration of Mg2+ that must be added to 0.10 M NaF in order to initiate a precipitate of magnesium 9 fluoride. (For MgF2 , Ksp = 6.9  10 .) 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

3 –4 23. Will a precipitate form (yes or no) when 50.0 mL of 1.2  10 M Pb(NO3)2 are added to 50.0 mL of 2.0  10 M Na2S? If so, identify the precipitate.

A. Yes, the precipitate is PbS. D. Yes, the precipitate is Pb(NO3)2.

B. Yes, the precipitate is NaNO3. E. No, a precipitate will not form.

C. Yes, the precipitate is Na2S.

2+ + 8 24. Solid sodium iodide is slowly added to a solution that is 0.0050 M Pb and 0.0050 M Ag . (Ksp (PbI2) = 1.4  10 ; Ksp (AgI) = 8.3  1017) a. What compound will precipitate first? + b. Calculate the Ag concentration when PbI2 just begins to precipitate. c. What percent of Ag+ remains in solution at this point?

4 25. A 50.0 mL sample of 2.0  10 M CuNO3 is added to 50.0 mL of 4.0 M NaCN. The formation constant of the complex ion 2– 9 Cu(CN)3 is 1.0  10 . What is the copper(I) ion concentration in this system at equilibrium?