Chapter 16 test

1.

1 point

Calculate the molar solubility of Mg(OH)2 in a solution with pH = 12.20.
(Ksp (Mg(OH)2 = 1.2 x 10-11)

3.5 x 10-6 M

4.8 x 10-8 M

. 2.4 x 10-6 M

1.4 x 10-4 M

7.6 x 10-10 M

2.

1 point

What is the optimum pH of a sodium formate/formic acid buffer? (For formic acid, Ka = 1.7  10–4.)

3.57

3.77

3.97

4.07

none of the above

3.

1 point

Which one of the following combinations cannot function as a buffer solution?

HNO3 and NaNO3

HF and NaF

HNO2 and NaNO2

HCN and KCN

NH3 and (NH4)2SO4

4.

1 point

What is the pH at the equivalence point in the titration of 100 mL of 0.10 M HCN (Ka = 4.9  10–10) with 0.10 M NaOH?

7.0

11.0

6.0

3.0

12.0

5.

1 point

Select True or False: The equation shown here is the net reaction that occurs when a strong base is added to a CO32–/HCO3– buffer solution (for carbonic acid, Ka1 = 4.2  10–7, Ka2 = 2.4  10–8): HCO3– + OH–  CO32– + H2O

True

False

6.

1 point

A titration of an acid and base to the equivalence point results in a noticeably acidic solution. It is likely this titration involves

a strong acid and a strong base.

a weak acid and a strong base.

a weak acid and a weak base (where Ka equals Kb).

a strong acid and a weak base.

7.

1 point

Calculate the pH of a buffer solution that contains 0.25 M benzoic acid (C6H5CO2H) and 0.15M sodium benzoate (C6H5COONa). [Ka = 6.5  10–5 for benzoic acid]

3.97

4.41

3.40

4.83

4.19

8.

1 point

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

Strong acid vs. strong base.

Strong acid vs. weak base.

Weak acid vs. strong base.

All of the above.

None of the above.

9.

1 point

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

1.0

8.0

6.0

13.0

7.0

10.

1 point

25.0 mL of a 0.100 M solution of NH3 is titrated with 0.150M HCl. After 10.0 mL of the HCl has been added, the resultant solution is:

Acidic and before the equivalence point

Basic and before the equivalence point

Acidic and after the equivalence point

Neutral and at the equivalence point

Basic and after the equivalence point

11.

1 point

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

8.71

13.0

5.85

9.91

5.15

12.

1 point

The pH at the equivalence point of a titration may differ from 7.0 due to

the initial concentration of the standard solution.

hydrolysis of the salt formed

the indicator used

the self-ionization of H2O.

the initial pH of the unknown

13.

1 point

Consider a buffer solution prepared from HOCl and NaOCl. Which is the net ionic equation for the reaction that occurs when NaOH is added to this buffer?

OH– + HOCl  H2O + OCl–

Na+ + HOCl  NaCl + OH–

H+ + HOCl  H2 + OCl–

OH– + OCl–  HOCl + O2–

NaOH + HOCl  H2O + NaCl

14.

1 point

Calculate the molar solubility of AgBr in a 0.25M solution of NH3(aq) (Ksp (AgBr) = 7.7 x 10-13 ; Kf (Ag(NH3)2+) = 1.5 x 107

8.8 x 10-7 M

8.4 x 10-4 M

2.5 x 10-1 M

9.7 x 102 M

3.4 x 10-3 M

15.

1 point

Select True or False: The solubility of a salt increases as its Ksp increases.

True

False

16.

1 point

Which of the following yields a buffer solution when equal volumes of the two solutions are mixed?

0.15M HBr and 0.15M NaBr

0.10M HF and 0.10M NaF

0.15 M HNO3 and 0.15 M NaNO3

0.10M HCl and 0.10 M NaCl

0.10M HClO4 and 0.10 NaClO4

17.

1 point

Select True or False: The equation shown here is the net reaction that occurs when a strong acid is added to a CO32–/HCO3– buffer solution (for carbonic acid, Ka1 = 4.2  10–7, Ka2 = 2.4  10–8): CO32– + H+  HCO3– and HCO3– + H+  H2CO3

True

False

18.

1 point

Select True or False: All indicators are weak acids that are one color in acidic solution and another color in basic solution.

True

False

19.

1 point

Select True or False: The pH of a solution that is 0.20 M CH3COOH and 0.20 M CH3COONa should be higher than the pH of a 0.20 M CH3COOH solution.

True

False

20.

1 point

Methyl red is a common acid-base indicator. It has a Ka equal to 6.3  10–6. Its un-ionized form is red and its anionic form is yellow. What color would a methyl red solution have at pH = 7.8?

violet

yellow

red

green

blue