Assignment 17 A
1- Does the pH increase, decrease, or remain the same on addition of each of the following? (i) NaNO2 to
a solution of HNO2, (ii) HCl to a solution of NaC2H3O2
a) (i) increase, (ii) remain the same
b) (i) increase, (ii) increase
c) (i) decrease, (ii) decrease
d) (i) decrease, (ii) increase
e) (i) increase, (ii) decrease
(The NO2− causes the formation of more HNO2, which lowers the pH. The added H+ from the HCl
reacts with the OH− from the hydrolysis of the C2H3O2− ion, increasing the [H+] and decreasing the
pH.)
2- Calculate the pH of a buffer that is 0.20 M in formic acid and 0.15 M in sodium formate (Ka = 1.8 × 10−4).
a) 0.82
b) 8.33
c) 3.62
d) 2.4 × 10−4
e) 0.70
(The hydrogen-ion concentration equals 2.4 × 10−4 M.)
3- A 50.0-mL sample of 0.50M acetic acid, HC2H3O2, is titrated with a 0.150 M NaOH solution. Calculate the
pH after 25.0 mL of the base have been added (Ka = 1.8 × 10−5).
a) 3.99
b) 3.92
c) 0.55
d) 2.62
e) 2.52
(This is the pH of the buffer formed by the acid-base reaction.)
4- If the molar solubility of CaF2 at 35°C is 1.24 × 10−3 mol/L, what is Ksp at this temperature?
a) 7.63 × 10−9
b) 3.81 × 10−9
c) 1.54 × 10−6
d) 1.91 × 10−9
(The [Ca2+] = 1.24 × 10−3 M and the [F−] = 2.48 × 10−3 M. The Ksp expression is [Ca2+][F−]2.)
5- Which of the following would you not expect to be more soluble in acid than in pure water?
a) AgCl
b) CuCN
c) AlPO4
d) BaCO3
e) FeS
(In order for a solid to be more soluble in acid than in pure water, the anion of the solid compound
must be a weak base. Chloride is a neutral anion.)
6- If 0.1 M aqueous solutions of the following pairs of substances are mixed together, which pair could result in
the formation of a precipitate?
a) NaOH, HCl
b) KCl, Al(NO3)3
c) Na2S, FeCl3
d) NaOH, KNO3
e) Ni(NO3)2, Mg(NO3)2
(Iron will form a precipitate with a basic sulfide solution.)
7- Which one of the following changes is incorrect?
a) CaCl2 added to an HCl solution does not change the pH.
b) NaBr added to a solution of HBr will raise the pH.
c) NaC2H3O2 added to a solution of HC2H3O2 raises the pH.
d) HNO3 added to CaC2H3O2 lowers the pH.
e) KNO3 added to an HCl solution does not change the pH.
(It does nothing to the pH, since HBr is a strong acid.)
8- How many moles of sodium hypobromite, NaBrO, should be added to 1.00 L of 0.200 M hypobromous acid,
HBrO (Ka = 2.5 × 10−9), to form a buffer solution of pH 8.80? Assume that no volume change occurs
when the NaBrO is added.
a) A buffer solution with this pH cannot be formed.
b) 0.13 mol
c) 0.32 mol
d) 0.20 mol
(The [BrO−] = Ka[acid]/[H+].)
9- Calculate the molar concentration of bromide ions in a saturated solution of mercury(II) bromide
(Ksp = 8.0 × 10−20).
a) 2.8 × 10−10 M
b) 2.0 × 10−20 M
c) 2.7 × 10−7 M
d) 5.4 × 10−7 M
(The Ksp equals 4S3, where S equals the molar solubility. The concentration of bromide ion equals
2S.)
10- Calculate the solubility in mol/L of Cu(OH)2 at pH = 7.0. Ksp = 2.2 × 10−20.
a) 4.4 × 10−6
b) 2.1 × 10−4
c) 5.5 × 10−7
d) 2.2 × 10−6
e) 1.77 × 10−7
(The hydroxide-ion concentration at pH 7.0 is 1 × 10−7 M.)
11- Ksp for AgI = 8.3 × 10−17, and Kf for Ag(CN)2− = 1.0 × 1021. Calculate the equilibrium constant for the
following reaction:
AgI(s) + 2CN−(aq)  Ag(CN)2−(aq) + I−(aq)
a) 1.2 × 10−5
b) 8.3 × 104
c) 8.3 x 20−38
d) 8.3 × 10−17
e) 1.0 × 1021
(The reaction of interest is the algebraic sum of the dissolution of AgI and the formation reaction of
Ag(CN)2−. Therefore, the equilibrium constant for the reaction of interest is the product of Ksp × Kf.)
12- Calculate the minimum pH needed to precipitate Ni(OH)2 so completely that the concentration of Ni2+ is less
than 1.0 µg/L (1.0 part per billion (ppb)) (Ksp = 1.6 × 10-14).
a) 10.10
b) 9.49
c) 7.97
d) 3.01
e) 10.99
(The maximum allowable hydroxide-ion concentration is 9.7 × 10−4.)
13- In the course of various qualitative analysis procedures, the following mixture is encountered: Mg2+ and K+.
Suggest how this mixture might be separated.
a) add dilute HCl
b) It is not possible to separate the ions.
c) add (NH4)2HPO4 to a basic solution
d) add (NH4)2S at pH 8
e) add 0.2 M HCl and H2S
(This will precipitate only the Mg2+.)
14- Consider the titration of 50. mL of 0.217 M HN3 (Ka = 2.6 × 10−5) with 0.183 M NaOH. Calculate the pH of
the solution after the addition of 29.7 mL of NaOH solution.
a) 7.00
b) 4.51
c) 2.62 ,.
d) 4.59
e) 2.88
(This is the pH of the buffer formed at the half-neutralization point in the titration.)
15- Calculate the pH of a solution that is both 0.50 M CH3COOH and 0.50 M CH3COONa
(Ka (CH3COOH) is 1.76 × 10−5).
a) 7.00
b) 2.52
c) 4.75
d) 11.48
e) 9.26
(When the concentrations of the salt and acid are equal, the pH equals the pKa of the acid.)
16- What is the pH of a buffer prepared from 0.30 M formic acid and 0.15 M potassium formate
(Ka = 1.8 × 10−4)?
a) 2.13
b) 4.04
c) 3.44
d) 3.74
e) 1.87
(The dissociation of formic acid is negligible.)
17- How many moles of sodium hypobromite, NaBrO, should be added to 1.00 L of 0.200 M hypobromous acid,
HBrO (Ka = 2.5 × 10−9), to form a buffer solution of pH 8.80? Assume that no volume change occurs
when the NaBrO is added.
a) 0.20 mol
b) 0.13 mol
c) 0.32 mol
d) A buffer solution with this pH cannot be formed.
(The [BrO−] = Ka[acid]/[H+].
18- A buffer contains 0.30 M acetic acid and 0.20 M sodium acetate. What is the pH of the buffer as prepared
and after 0.030 mol/L of a strong acid or 0.030 mol/L of a strong base are added (Ka = 1.8 × 10−5)?
a) 4.57, 4.46 with acid, 4.68 with base
b) 4.57, 4.56 with acid, 4.56 with base
c) 4.92, 5.03 with acid, 4.81 with base
d) 4.56, 4.68 with acid, 4.46 with base
e) 4.74, 4.61 with acid, 4.88 with base
(The added H+ reacted with the acetate to lower the pH slightly, and the added OH− reacted with
the acetic acid to raise the pH slightly.)
19- Estimate the pH at the equivalence point of an HOAc solution if 25.5 mL of this solution required 37.5 mL
of 0.175 M NaOH to reach the equivalence point (Ka = 1.8 × 10−5).
a) 4.31
b) 12.2
c) 7.00
d) 8.88
e) 2.07
(NaOAc is the salt of a weak acid, so its solution would be slightly basic.)
20- Calculate the pH of a solution formed by adding 50.0 mL of 6.0 M NH3 to 75.0 mL of 1.0 M HCl.
(Kb (NH3) = 1.8 × 10−5).
a) 4.27
b) 5.22
c) 9.26
d) 9.73
e) 8.79
(This is the pH of the buffer formed after the acid-base reaction.)
21- The solubility of Mg(OH)2 is 1.4 × 10−4 mol/L. Determine the Ksp for Mg(OH)2.
a) 2.7 × 10−12
b) 3.9 × 10−8
c) 1.1 × 10−11
d) 5.5 × 10−12
e) 2.0 × 10−8
(The hydroxide-ion concentration is double the magnesium-ion concentration.
The Ksp = [Mg2+][OH−] 2.)
22- Calculate the number of mg of silver in 250 mL of a saturated solution of Ag2CO3 (Ksp = 8.1 × 10−12).
a) 14 mg
b) 6.8 mg
c) 27 mg
d) 8.7 mg
e) 3.4 mg
(The Ksp is 4x3, where x is the molar solubility. The mass in milligrams of Ag equals the
concentration (2x) converted to mg of Ag in 250 mL.)
23- Will Mn(OH)2 precipitate from solution if the pH of a 0.050 M solution of MnCl2 is adjusted to 8.0
(Ksp = 1.8 × 10−11)?
a) no
b) yes
(Q = 5 × 10−14 < Ksp.)
24- In which case will a precipitate form?
(i) a 0.050 M MnCl2 solution with pH = 8.00 (does Mn(OH)2 precipitate?) Ksp = 1.9 × 10−13.
(ii) 100. mL of 0.010 M AgNO3 is added to 20 mL of 0.050 M Na2SO4 (does Ag2SO4 precipitate?)
Ksp = 1.4 × 10−5.
a) i, yes; ii, no
b) i, no; ii, yes
c) i, no; ii, no
d) i, yes; ii, yes
(Q for (i) is 5.0 × 10−14 compared to Ksp = 1.9 × 10−13, while Q for (ii) is 5.8 × 10−7 compared to
Ksp = 1.4 × 10−5.)
25- Consider a solution containing 0.181 M lead (II) ions and 0.174 M mercury(II) ions. Calculate the maximum
concentration of sulfide ions that can be in solution without precipitating any lead ions.
Ksp: lead sulfide is 3.4 × 10−28; mercury(II) sulfide is 4.0 × 10−53.
a) 3.3 × 10−52 M
b) 2.0 × 10−27 M
c) 1.9 × 10−27 M
d) 2.2 × 10−52 M
(This is the Ksp for PbS divided by the lead concentration.)