🎉 Announcing Numerade's $26M Series A, led by IDG Capital!Read how Numerade will revolutionize STEM Learning ## Chapter 17 ## Aqueous Ionic Equilibrium ## Educators + 4 more educators ### Problem 1 What is the pH range of human blood? How is human blood maintained in this pH range? Md M. Auburn University Main Campus ### Problem 2 What is a buffer? How does a buffer work? How does it neutralize added acid? Added base? Shahriar K. Auburn University Main Campus ### Problem 3 What is the common ion effect? Md M. Auburn University Main Campus ### Problem 4 What is the Henderson-Hasselbalch equation, and why is it useful? Keenan M. University of Miami ### Problem 5 What is the pH of a buffer when the concentrations of both buffer components (the weak acid and its conjugate base) are equal? What happens to the pH when the buffer contains more of the weak acid than the conjugate base? More of the conjugate base than the weak acid? Md M. Auburn University Main Campus ### Problem 6 Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong acid is added to the buffer? What happens when a small amount of strong base is added? Shahriar K. Auburn University Main Campus ### Problem 7 How do you use the Henderson-Hasselbalch equation to calculate the pH of a buffer containing a base and its conjugate acid? Specifically, how do you determine the correct value for pKa? Md M. Auburn University Main Campus ### Problem 8 What factors influence the effectiveness of a buffer? What are the characteristics of an effective buffer? Shahriar K. Auburn University Main Campus ### Problem 9 What is the effective pH range of a buffer (relative to the pKa of the weak acid component)? Md M. Auburn University Main Campus ### Problem 10 Describe acid-base titration. What is the equivalence point? Keenan M. University of Miami ### Problem 11 The pH at the equivalence point of the titration of a strong acid with a strong base is 7.0. However, the pH at the equivalence point of the titration of a weak acid with a strong base is above 7.0. Explain. Md M. Auburn University Main Campus ### Problem 12 The volume required to reach the equivalence point of an acid-base titration depends on the volume and concentration of the acid or base to be titrated and on the concentration of the acid or base used to do the titration. It does not, however, depend on the strength or weakness of the acid or base being titrated. Explain. Keenan M. University of Miami ### Problem 13 In the titration of a strong acid with a strong base, how do you calculate these quantities? a. initial pH b. pH before the equivalence point c. pH at the equivalence point d. pH beyond the equivalence point Md M. Auburn University Main Campus ### Problem 14 In the titration of a weak acid with a strong base, how do you calculate these quantities? a. initial pH b. pH before the equivalence point c. pH at one-half the equivalence point d. pH at the equivalence point e. pH beyond the equivalence point Keenan M. University of Miami ### Problem 15 The titration of a diprotic acid with sufficiently different pKa's displays two equivalence points.Why? Md M. Auburn University Main Campus ### Problem 16 In the titration of a polyprotic acid, the volume required to reach the first equivalence point is identical to the volume required to reach the second one. Why? Keenan M. University of Miami ### Problem 17 What is the difference between the endpoint and the equivalence point in a titration? Md M. Auburn University Main Campus ### Problem 18 What is an indicator? How can an indicator signal the equivalence point of a titration? Keenan M. University of Miami ### Problem 19 What is the solubility-product constant? Write a general expression for the solubility constant of a compound with the general formula AmXn. Md M. Auburn University Main Campus ### Problem 20 What is molar solubility? How do you obtain the molar solubility of a compound from Ksp? Keenan M. University of Miami ### Problem 21 How does a common ion affect the solubility of a compound? More specifically, how is the solubility of a compound with the general formula AX different in a solution containing one of the common ions (A+ or X-) than it is in pure water? Explain. Md M. Auburn University Main Campus ### Problem 22 How is the solubility of an ionic compound with a basic anion affected by pH? Explain. Keenan M. University of Miami ### Problem 23 For a given solution containing an ionic compound, what is the relationship between Q, Ksp, and the relative saturation of the solution? Md M. Auburn University Main Campus ### Problem 24 What is selective precipitation? Under which conditions does selective precipitation occur? Keenan M. University of Miami ### Problem 25 In which of these solutions does HNO2 ionize less than it does in pure water? a. 0.10 M NaCl b. 0.10 M KNO3 c. 0.10 M NaOH d. 0.10 M NaNO2 Md M. Auburn University Main Campus ### Problem 26 A formic acid solution has a pH of 3.25. Which of these substances raises the pH of the solution upon addition? Explain your answer. a. HCl b. NaBr c. NaCHO2 d. KCl Keenan M. University of Miami ### Problem 27 Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. a solution that is 0.20 M in HCHO2 and 0.15 M in NaCHO2 b. a solution that is 0.16 M in NH3 and 0.22 M in NH4Cl Md M. Auburn University Main Campus ### Problem 28 Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. a solution that is 0.195 M in HC2H3O2 and 0.125 M in KC2H3O2 b. a solution that is 0.255 M in CH3NH2 and 0.135 M in CH3NH3Br Keenan M. University of Miami ### Problem 29 Calculate the percent ionization of a 0.15 M benzoic acid solution in pure water and in a solution containing 0.10 M sodium benzoate. Why does the percent ionization differ significantly in the two solutions? Md M. Auburn University Main Campus ### Problem 30 Calculate the percent ionization of a 0.13 M formic acid solution in pure water and also in a solution containing 0.11 M potassium formate. Explain the difference in percent ionization in the two solutions. Keenan M. University of Miami ### Problem 31 Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. 0.15 M HF b. 0.15 M NaF c. a mixture that is 0.15 M in HF and 0.15 M in NaF Md M. Auburn University Main Campus ### Problem 32 Solve an equilibrium problem (using an ICE table) to calculate the pH of each solution. a. 0.18 M CH3NH2 b. 0.18 M CH3NH3Cl c. a mixture that is 0.18 M in CH3NH2 and 0.18 M in CH3NH3Cl Keenan M. University of Miami ### Problem 33 A buffer contains significant amounts of acetic acid and sodium acetate. Write equations that demonstrate how this buffer neutralizes added acid and added base. Md M. Auburn University Main Campus ### Problem 34 A buffer contains significant amounts of ammonia and ammonium chloride. Write equations that demonstrate how this buffer neutralizes added acid and added base. Keenan M. University of Miami ### Problem 35 Use the Henderson-Hasselbalch equation to calculate the pH of each solution in Problem 27. Md M. Auburn University Main Campus ### Problem 36 Use the Henderson-Hasselbalch equation to calculate the pH of each solution in Problem 28. Keenan M. University of Miami ### Problem 37 Use the Henderson-Hasselbalch equation to calculate the pH of each solution. a. a solution that is 0.135 M in HClO and 0.155 M in KClO b. a solution that contains 1.05% C2H5NH2 by mass and 1.10% C2H5NH3Br by mass c. a solution that contains 10.0 g of HC2H3O2 and 10.0 g of NaC2H3O2 in 150.0 mL of solution Md M. Auburn University Main Campus ### Problem 38 Use the Henderson-Hasselbalch equation to calculate the pH of each solution. a. a solution that is 0.145 M in propanoic acid and 0.115 M in potassium propanoate b. a solution that contains 0.785% C5H5N by mass and 0.985% C5H5NHCl by mass c. a solution that contains 15.0 g of HF and 25.0 g of NaF in 125 mL of solution Keenan M. University of Miami ### Problem 39 Calculate the pH of the solution that results from each mixture. a. 50.0 mL of 0.15 M HCHO2 with 75.0 mL of 0.13 M NaCHO2 b. 125.0 mL of 0.10 M NH3 with 250.0 mL of 0.10 M NH4Cl Catherine S. Missouri State University ### Problem 40 Calculate the pH of the solution that results from each mixture. a. 150.0 mL of 0.25 M HF with 225.0 mL of 0.30 M NaF b. 175.0 mL of 0.10 M C2H5NH2 with 275.0 mL of 0.20 M C2H5NH3Cl Keenan M. University of Miami ### Problem 41 Calculate the ratio of NaF to HF required to create a buffer with pH = 4.00. Md M. Auburn University Main Campus ### Problem 42 Calculate the ratio of CH3NH2 to CH3NH3Cl concentration required to create a buffer with pH = 10.24. Keenan M. University of Miami ### Problem 43 What mass of sodium benzoate should you add to 150.0 mL of a 0.15 M benzoic acid solution to obtain a buffer with a pH of 4.25? (Assume no volume change.) Md M. Auburn University Main Campus ### Problem 44 What mass of ammonium chloride should you add to 2.55 L of a 0.155 M NH3 solution to obtain a buffer with a pH of 9.55? ( Assume no volume change.) Keenan M. University of Miami ### Problem 45 A 250.0-mL buffer solution is 0.250 M in acetic acid and 0.250 M in sodium acetate. a. What is the initial pH of this solution? b. What is the pH after addition of 0.0050 mol of HCl? c. What is the pH after addition of 0.0050 mol of NaOH? Md M. Auburn University Main Campus ### Problem 46 A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. a. What is the initial pH of this solution? b. What is the pH after addition of 150.0 mg of HBr? c. What is the pH after addition of 85.0 mg of NaOH? George M. Numerade Educator ### Problem 47 For each solution, calculate the initial and final pH after the addition of 0.010 mol of HCl. a. 500.0 mL of pure water b. 500.0 mL of a buffer solution that is 0.125 M in HC2H3O2 and 0.115 M in NaC2H3O2 c. 500.0 mL of a buffer solution that is 0.155 M in C2H5NH2 and 0.145 M in C2H5NH3Cl Md M. Auburn University Main Campus ### Problem 48 For each solution, calculate the initial and final pH after the addition of 0.010 mol of NaOH. a. 250.0 mL of pure water b. 250.0 mL of a buffer solution that is 0.195 M in HCHO2 and 0.275 M in KCHO2 c. 250.0 mL of a buffer solution that is 0.255 M in CH3CH2NH2 and 0.235 M in CH3CH2NH3Cl Keenan M. University of Miami ### Problem 49 A 350.0-mL buffer solution is 0.150 M in HF and 0.150 M in NaF. What mass of NaOH does this buffer neutralize before the pH rises above 4.00? If the same volume of the buffer was 0.350 M in HF and 0.350 M in NaF, what mass of NaOH is neutralized before the pH rises above 4.00? Md M. Auburn University Main Campus ### Problem 50 A$100.0-\mathrm{mL}$. buffer solution is$0.100 \mathrm{M}$in$\mathrm{NH}_{3}$and$0.125 \mathrm{M}$in$\mathrm{NH}_{4} \mathrm{Br}$. What mass of$1 \mathrm{HCl}$does this buffer neutralize before the$\mathrm{pH}$falls below 9.007 If the same volume of the buffer were$0.250 \mathrm{M}$in$\mathrm{NH}_{3}$and$0.400 \mathrm{M}$in$\mathrm{NH}_{4} \mathrm{Br},$what mass of$\mathrm{HCl}$is neutralized before the pH falls below$9.00 ?\$

Keenan M.
University of Miami

### Problem 51

Determine whether the mixing of each pair of solutions results in a buffer.
a. 100.0 mL of 0.10 M NH3; 100.0 mL of 0.15 M NH4Cl
b. 50.0 mL of 0.10 M HCl; 35.0 mL of 0.150 M NaOH
c. 50.0 mL of 0.15 M HF; 20.0 mL of 0.15 M NaOH
d. 175.0 mL of 0.10 M NH3; 150.0 mL of 0.12 M NaOH
e. 125.0 mL of 0.15 M NH3; 150.0 mL of 0.20 M NaOH

Katie M.

### Problem 52

Determine whether the mixing of each pair of solutions results in a buffer.
a. 75.0 mL of 0.10 M HF; 55.0 mL of 0.15 M NaF
b. 150.0 mL of 0.10 M HF; 135.0 mL of 0.175 M HCl
c. 165.0 mL of 0.10 M HF; 135.0 mL of 0.050 M KOH
d. 125.0 mL of 0.15 M CH3NH2; 120.0 mL of 0.25 M CH3NH3Cl
e. 105.0 mL of 0.15 M CH3NH2; 95.0 mL of 0.10 M HCl

Keenan M.
University of Miami

### Problem 53

Blood is buffered by carbonic acid and the bicarbonate ion.Normal blood plasma is 0.024 M in HCO3- and 0.0012 M H2CO3(pKa1 for H2CO3 at body temperature is 6.1).

a. What is the pH of blood plasma?
b. If the volume of blood in a normal adult is 5.0 L, what mass of HCl can be neutralized by the buffering system in blood before the pH falls below 7.0 (which would result in death)?
c. Given the volume from part b, what mass of NaOH can be neutralized before the pH rises above 7.8?

Md M.
Auburn University Main Campus

### Problem 54

The fluids within cells are buffered by H2PO4- and HPO42-.

a. Calculate the ratio of HPO42- to H2PO4- required to maintain a pH of 7.1 within a cell.

b. Could a buffer system employing H3PO4 as the weak acid and H2PO4- as the weak base be used as a buffer system within cells? Explain.

Keenan M.
University of Miami

### Problem 55

Which buffer system is the best choice to create a buffer with pH = 7.20? For the best system, calculate the ratio of the masses of the buffer components required to make the buffer.

Md M.
Auburn University Main Campus

### Problem 56

Which buffer system is the best choice to create a buffer with
pH = 9.00? For the best system, calculate the ratio of the masses
of the buffer components required to make the buffer.
HF>KF HNO2>KNO2
NH3>NH4Cl HClO>KClO

Keenan M.
University of Miami

### Problem 57

A 500.0-mL buffer solution is 0.100 M in HNO2 and 0.150 M in
KNO2. Determine whether each addition would exceed the
capacity of the buffer to neutralize it.
a. 250.0 mg NaOH
b. 350.0 mg KOH
c. 1.25 g HBr
d. 1.35 g HI

Md M.
Auburn University Main Campus

### Problem 58

A 1.0-L buffer solution is 0.125 M in HNO2 and 0.145 M in
NaNO2. Determine the concentrations of HNO2 and NaNO2 after
a. 1.5 g HCl
b. 1.5 g NaOH
c. 1.5 g HI

Keenan M.
University of Miami

### Problem 59

The graphs labeled (a) and (b) are the titration curves for two
equal-volume samples of monoprotic acids, one weak and one
strong. Both titrations were carried out with the same concentration
of strong base.

Md M.
Auburn University Main Campus

### Problem 60

Two 25.0-mL samples, one 0.100 M HCl and the other 0.100 M
HF, are titrated with 0.200 M KOH.
a. What is the volume of added base at the equivalence point for
each titration?
b. Is the pH at the equivalence point for each titration acidic,
basic, or neutral?
c. Which titration curve has the lower initial pH?
d. Sketch each titration curve.

Keenan M.
University of Miami

### Problem 61

Two 20.0-mL samples, one 0.200 M KOH and the other 0.200 M
CH3NH2, are titrated with 0.100 M HI.
a. What is the volume of added acid at the equivalence point for
each titration?
b. Is the pH at the equivalence point for each titration acidic,
basic, or neutral?
c. Which titration curve has the lower initial pH?
d. Sketch each titration curve.

Md M.
Auburn University Main Campus

### Problem 62

The graphs labeled (a) and (b) are the titration curves for two
equal-volume samples of bases, one weak and one strong. Both
titrations were carried out with the same concentration of
strongacid.

i. What is the approximate pH at the equivalence point of each
curve?
ii. Which graph corresponds to the titration of the strong base
and which one to the weak base?

Keenan M.
University of Miami

### Problem 63

Consider the curve shown here for the titration of a weak monoprotic
acid with a strong base and answer each question.
a. What is the pH, and what is the volume of added base at the
equivalence point?
b. At what volume of added base is the pH calculated by working
an equilibrium problem based on the initial concentration
and Ka of the weak acid?
c. At what volume of added base does pH = pKa?
d. At what volume of added base is the pH calculated by working
an equilibrium problem based on the concentration and Kb of
the conjugate base?
e. Beyond what volume of added base is the pH calculated by
focusing on the amount of excess strong base added?

Md M.
Auburn University Main Campus

### Problem 64

Consider the curve shown here for the titration of a weak base
with a strong acid and answer each question.

a. What is the pH, and what is the volume of added acid at the
equivalence point?
b. At what volume of added acid is the pH calculated by working
an equilibrium problem based on the initial concentration
and Kb of the weak base?
c. At what volume of added acid does pH = 14 - pKb?
d. At what volume of added acid is the pH calculated by working
an equilibrium problem based on the concentration and Ka of
the conjugate acid?
e. Beyond what volume of added acid is the pH calculated by
focusing on the amount of excess strong acid added?

Keenan M.
University of Miami

### Problem 65

Consider the titration of a 35.0-mL sample of 0.175 M HBr with
0.200 M KOH. Determine each quantity.
a. the initial pH
b. the volume of added base required to reach the equivalence point
c. the pH at 10.0 mL of added base
d. the pH at the equivalence point
e. the pH after adding 5.0 mL of base beyond the equivalence point

Md M.
Auburn University Main Campus

### Problem 66

A 20.0-mL sample of 0.125 M HNO3 is titrated with 0.150 M
NaOH. Calculate the pH for at least five different points on the
titration curve and sketch the curve. Indicate the volume at the

Keenan M.
University of Miami

### Problem 67

Consider the titration of a 25.0-mL sample of 0.115 M RbOH with
0.100 M HCl. Determine each quantity.
a. the initial pH
b. the volume of added acid required to reach the equivalence
point
c. the pH at 5.0 mL of added acid
d. the pH at the equivalence point
e. the pH after adding 5.0 mL of acid beyond the equivalence point

Md M.
Auburn University Main Campus

### Problem 68

A 15.0-mL sample of 0.100 M Ba(OH)2 is titrated with 0.125 M
HCl. Calculate the pH for at least five different points on the
titration curve and sketch the curve. Indicate the volume at the

Keenan M.
University of Miami

### Problem 69

Consider the titration of a 20.0-mL sample of 0.105 M HC2H3O2
with 0.125 M NaOH. Determine each quantity.
a. the initial pH
b. the volume of added base required to reach the equivalence point
c. the pH at 5.0 mL of added base
d. the pH at one-half of the equivalence point
e. the pH at the equivalence point
f. the pH after adding 5.0 mL of base beyond the equivalence point

Md M.
Auburn University Main Campus

### Problem 70

A 30.0-mL sample of 0.165 M propanoic acid is titrated with
0.300 M KOH. Calculate the pH at each volume of added base:
0mL, 5 mL, 10 mL, equivalence point, one-half equivalence
point, 20 mL, 25 mL. Sketch the titration curve.

Keenan M.
University of Miami

### Problem 71

Consider the titration of a 25.0-mL sample of 0.175 M CH3NH2
with 0.150 M HBr. Determine each quantity.
a. the initial pH
b. the volume of added acid required to reach the equivalence
point
c. the pH at 5.0 mL of added acid
d. the pH at one-half of the equivalence point
e. the pH at the equivalence point
f. the pH after adding 5.0 mL of acid beyond the equivalence point

Teesta D.
University of Pittsburgh - Main Campus

### Problem 72

A 25.0-mL sample of 0.125 M pyridine is titrated with 0.100 M
HCl. Calculate the pH at each volume of added acid: 0 mL,
10mL, 20 mL, equivalence point, one-half equivalence point,
40mL, 50 mL. Sketch the titration curve.

Keenan M.
University of Miami

### Problem 73

Consider the titration curves (labeled a and b) for equal volumes
of two weak acids, both titrated with 0.100 M NaOH.

i. Which acid solution is more concentrated?
ii. Which acid has the larger Ka?

Md M.
Auburn University Main Campus

### Problem 74

Consider the titration curves (labeled a and b) for equal volumes
of two weak bases, both titrated with 0.100 M HCl.

i. Which base solution is more concentrated?
ii. Which base has the larger Kb?

Keenan M.
University of Miami

### Problem 75

A 0.229-g sample of an unknown monoprotic acid is titrated with
0.112 M NaOH. The resulting titration curve is shown here.
Determine the molar mass and pKa of the acid.

Md M.
Auburn University Main Campus

### Problem 76

A 0.446-g sample of an unknown monoprotic acid is titrated with
0.105 M KOH. The resulting titration curve is shown here.
Determine the molar mass and pKa of the acid.

LE
Leon E.

### Problem 77

A 20.0-mL sample of 0.115 M sulfurous acid (H2SO3) solution is
titrated with 0.1014 M KOH. At what added volume of base solution
does each equivalence point occur?

Md M.
Auburn University Main Campus

### Problem 78

A 20.0-mL sample of a 0.125 M diprotic acid (H2A) solution is titrated
with 0.1019 M KOH. The acid ionization constants for the
acid are Ka1 = 5.2 * 10-5 and Ka2 = 3.4 * 10-10. At what added
volume of base does each equivalence point occur?

Keenan M.
University of Miami

### Problem 79

Methyl red has a pKa of 5.0 and is red in its acid form and yellow
in its basic form. If several drops of this indicator are placed in a
25.0-mL sample of 0.100 M HCl, what color does the solution
appear? If 0.100 M NaOH is slowly added to the HCl sample, in
what pH range will the indicator change color?

Md M.
Auburn University Main Campus

### Problem 80

Phenolphthalein has a pKa of 9.7. It is colorless in its acid form
and pink in its basic form. For each of the pH values, calculate
[In-]/[HIn] and predict the color of a phenolphthalein solution.
a. pH = 2.0
b. pH = 5.0
c. pH = 8.0
d. pH = 11.0

Keenan M.
University of Miami

### Problem 81

Referring to Table 17.1, pick an indicator for use in the titration
of each acid with a strong base.
a. HF b. HCl c. HCN

Md M.
Auburn University Main Campus

### Problem 82

Referring to Table 17.1, pick an indicator for use in the titration
of each base with a strong acid.
a. CH3NH2 b. NaOH c. C6H5NH2

Keenan M.
University of Miami

### Problem 83

Write balanced equations and expressions for Ksp for the
dissolution of each ionic compound.
a. BaSO4 b. PbBr2 c. Ag2CrO4

Md M.
Auburn University Main Campus

### Problem 84

Write balanced equations and expressions for Ksp for the
dissolution of each ionic compound.
a. CaCO3 b. PbCl2 c. AgI

Keenan M.
University of Miami

### Problem 85

Refer to the Ksp values in Table 17.2 to calculate the molar
solubility of each compound in pure water.
a. AgBr b. Mg(OH)2 c. CaF2

Md M.
Auburn University Main Campus

### Problem 86

Refer to the Ksp values in Table 17.2 to calculate the molar
solubility of each compound in pure water.
a. MX (Ksp = 1.27 * 10-36)
b. Ag2CrO4
c. Ca(OH)2

Keenan M.
University of Miami

### Problem 87

Use the given molar solubilities in pure water to calculate Ksp for
each compound.
a. MX; molar solubility = 3.27 * 10-11 M
b. PbF2; molar solubility = 5.63 * 10-3 M
c. MgF2; molar solubility = 2.65 * 10-4 M

Md M.
Auburn University Main Campus

### Problem 88

Use the given molar solubilities in pure water to calculate Ksp for
each compound.
a. BaCrO4; molar solubility = 1.08 * 10-5 M
b. Ag2SO3; molar solubility = 1.55 * 10-5 M
c. Pd(SCN)2; molar solubility = 2.22 * 10-8 M

Keenan M.
University of Miami

### Problem 89

Two compounds with general formulas AX and AX2 have
Ksp = 1.5 * 10-5. Which of the two compounds has the higher
molar solubility?

Md M.
Auburn University Main Campus

### Problem 90

Consider the compounds with the generic formulas listed and
their corresponding molar solubilities in pure water. Which compound
has the smallest value of Ksp?
a. AX; molar solubility = 1.35 * 10-4 M
b. AX2; molar solubility = 2.25 * 10-4 M
c. A2X; molar solubility = 1.75 * 10-4 M

Keenan M.
University of Miami

### Problem 91

Refer to the Ksp value from Table 17.2 to calculate the solubility of
iron(II) hydroxide in pure water in grams per 100.0 mL of
solution.

Md M.
Auburn University Main Campus

### Problem 92

The solubility of copper(I) chloride is 3.91 mg per 100.0 mL of
solution. Calculate Ksp for CuCl.

Keenan M.
University of Miami

### Problem 93

Calculate the molar solubility of barium fluoride in each liquid
or solution.
a. pure water
b. 0.10 M Ba(NO3)2
c. 0.15 M NaF

Shahriar K.
Auburn University Main Campus

### Problem 94

Calculate the molar solubility of MX (Ksp = 1.27 * 10-36) in
each liquid or solution.
a. pure water
b. 0.25 M MCl2
c. 0.20 M Na2X

Keenan M.
University of Miami

### Problem 95

Calculate the molar solubility of calcium hydroxide in a solution
buffered at each pH.
a. pH = 4
b. pH = 7
c. pH = 9

Md M.
Auburn University Main Campus

### Problem 96

Calculate the solubility (in grams per 1.00 * 102 mL of solution)
of magnesium hydroxide in a solution buffered at pH = 10. How
does this compare to the solubility of Mg(OH)2 in pure water?

Keenan M.
University of Miami

### Problem 97

Is each compound more soluble in acidic solution or in pure
water? Explain.
a. BaCO3
b. CuS
c. AgCl
d. PbI2

Md M.
Auburn University Main Campus

### Problem 98

Is each compound more soluble in acidic solution or in pure
water? Explain.
a. Hg2Br2
b. Mg(OH)2
c. CaCO3
d. AgI

Keenan M.
University of Miami

### Problem 99

A solution containing sodium fluoride is mixed with one containing
calcium nitrate to form a solution that is 0.015 M in NaF
and 0.010 M in Ca(NO3)2. Does a precipitate form in the mixed
solution? If so, identify the precipitate.

Md M.
Auburn University Main Campus

### Problem 100

A solution containing potassium bromide is mixed with one containing
lead acetate to form a solution that is 0.013 M in KBr and
0.0035 M in Pb(C2H3O2)2. Does a precipitate form in the mixed
solution? If so, identify the precipitate.

Keenan M.
University of Miami

### Problem 101

Predict whether a precipitate forms if you mix 75.0 mL of a
NaOH solution with pOH = 2.58 with 125.0 mL of a 0.018 M
MgCl2 solution. Identify the precipitate, if any.

Md M.
Auburn University Main Campus

### Problem 102

Predict whether a precipitate forms if you mix 175.0 mL of a
0.0055 M KCl solution with 145.0 mL of a 0.0015 M AgNO3
solution. Identify the precipitate, if any.

Keenan M.
University of Miami

### Problem 103

Potassium hydroxide is used to precipitate each of the cations
from their respective solution. Determine the minimum concentration
of KOH required for precipitation to begin in each case.
a. 0.015 M CaCl2
b. 0.0025 M Fe(NO3)2
c. 0.0018 M MgBr2

Md M.
Auburn University Main Campus

### Problem 104

Determine the minimum concentration of the precipitating agent
on the right you need to cause precipitation of the cation from
the solution on the left.
a. 0.035 M Ba(NO3)2; NaF
b. 0.085 M CaI2; K2SO4
c. 0.0018 M AgNO3; RbCl

Keenan M.
University of Miami

### Problem 105

A solution is 0.010 M in Ba2+ and 0.020 M in Ca2+.

a. If sodium sulfate is used to selectively precipitate one of the
cations while leaving the other cation in solution, which
cation precipitates first? What minimum concentration of
Na2SO4 will trigger the precipitation of the cation that
precipitates first?

b. What is the remaining concentration of the cation that precipitates
first, when the other cation begins to precipitate?

Md M.
Auburn University Main Campus

### Problem 106

A solution is 0.022 M in Fe2+ and 0.014 M in Mg2+.

a. If potassium carbonate is used to selectively precipitate one of
the cations while leaving the other cation in solution, which
cation precipitates first? What minimum concentration of
K2CO3 will trigger the precipitation of the cation that
precipitates first?

b. What is the remaining concentration of the cation that
precipitates first, when the other cation begins to precipitate?

Keenan M.
University of Miami

### Problem 107

A solution is made 1.1 * 10-3 M in Zn(NO3)2 and 0.150 M in
NH3. After the solution reaches equilibrium, what concentration
of Zn2+(aq) remains?

Nicholas W.

### Problem 108

A 120.0-mL sample of a solution that is 2.8 * 10-3 M in AgNO3 is
mixed with a 225.0-mL sample of a solution that is 0.10 M in NaCN.
After the solution reaches equilibrium, what concentration of
Ag+(aq) remains?

Keenan M.
University of Miami

### Problem 109

Use the appropriate values of Ksp and Kf to find the equilibrium
constant for the reaction:

Md M.
Auburn University Main Campus

### Problem 110

Use the appropriate values of Ksp and Kf to find the equilibrium
constant for the reaction:

Keenan M.
University of Miami

### Problem 111

A 150.0-mL solution contains 2.05 g of sodium benzoate and 2.47 g
of benzoic acid. Calculate the pH of the solution.

Md M.
Auburn University Main Campus

### Problem 112

A solution is made by combining 10.0 mL of 17.5 M acetic acid
with 5.54 g of sodium acetate and diluting to a total volume of
1.50 L. Calculate the pH of the solution.

Keenan M.
University of Miami

### Problem 113

A buffer is created by combining 150.0 mL of 0.25 M HCHO2
with 75.0 mL of 0.20 M NaOH. Determine the pH of the buffer.

Md M.
Auburn University Main Campus

### Problem 114

A buffer is created by combining 3.55 g of NH3 with 4.78 g of HCl
and diluting to a total volume of 750.0 mL. Determine the pH of
the buffer.

Keenan M.
University of Miami

### Problem 115

A 1.0-L buffer solution initially contains 0.25 mol of NH3 and
0.25 mol of NH4Cl. In order to adjust the buffer pH to 8.75,
should you add NaOH or HCl to the buffer mixture? What mass
of the correct reagent should you add?

Md M.
Auburn University Main Campus

### Problem 116

A 250.0-mL buffer solution initially contains 0.025 mol of
HCHO2 and 0.025 mol of NaCHO2. In order to adjust the buffer
pH to 4.10, should you add NaOH or HCl to the buffer mixture?
What mass of the correct reagent should you add?

Keenan M.
University of Miami

### Problem 117

In analytical chemistry, bases used for titrations must often be
standardized; that is, their concentration must be precisely
determined. Standardization of sodium hydroxide solutions can
be accomplished by titrating potassium hydrogen phthalate
(KHC8H4O4), also known as KHP, with the NaOH solution to be
standardized.

a. Write an equation for the reaction between NaOH and KHP.

b. The titration of 0.5527 g of KHP required 25.87 mL of an
NaOH solution to reach the equivalence point. What is the
concentration of the NaOH solution?

Md M.
Auburn University Main Campus

### Problem 118

A 0.5224-g sample of an unknown monoprotic acid was titrated
with 0.0998 M NaOH. The equivalence point of the titration
occurs at 23.82 mL. Determine the molar mass of the unknown
acid.

Keenan M.
University of Miami

### Problem 119

A 0.25-mol sample of a weak acid with an unknown pKa is
combined with 10.0 mL of 3.00 M KOH, and the resulting
solution is diluted to 1.500 L. The measured pH of the solution is
3.85. What is the pKa of the weak acid?

Md M.
Auburn University Main Campus

### Problem 120

A 5.55-g sample of a weak acid with Ka = 1.3 * 10-4 is combined
with 5.00 mL of 6.00 M NaOH, and the resulting solution
is diluted to 750 mL. The measured pH of the solution is 4.25.
What is the molar mass of the weak acid?

Keenan M.
University of Miami

### Problem 121

A 0.552-g sample of ascorbic acid (vitamin C) is dissolved in
water to a total volume of 20.0 mL and titrated with 0.1103 M
KOH. The equivalence point occurs at 28.42 mL. The pH of the
solution at 10.0 mL of added base was 3.72. From this data,
determine the molar mass and Ka for vitamin C.

Md M.
Auburn University Main Campus

### Problem 122

Sketch the titration curve from Problem 121 by calculating the
pH at the beginning of the titration, at one-half of the
equivalence point, at the equivalence point, and at 5.0 mL beyond
the equivalence point. Pick a suitable indicator for this titration
from Table 17.1.

Keenan M.
University of Miami

### Problem 123

One of the main components of hard water is CaCO3. When hard
water evaporates, some of the CaCO3 is left behind as a white
mineral deposit. If a hard water solution is saturated with
calcium carbonate, what volume of the solution has to evaporate
to deposit 1.00 * 102 mg of CaCO3?

Md M.
Auburn University Main Campus

### Problem 124

Gout-a condition that results in joint swelling and pain-is
caused by the formation of sodium urate (NaC5H3N4O3) crystals
within tendons, cartilage, and ligaments. Sodium urate precipitates
out of blood plasma when uric acid levels become abnormally
high. This can happen as a result of eating too many rich foods
and consuming too much alcohol, which is why gout is sometimes
referred to as the 'disease of kings.' If the sodium concentration in
blood plasma is 0.140 M and Ksp for sodium urate is 5.76 * 10-8,
what minimum concentration of urate results in precipitation?

Keenan M.
University of Miami

### Problem 125

Pseudogout, a condition with symptoms similar to those of gout
(see Problem 124), is caused by the formation of calcium
diphosphate (Ca2P2O7) crystals within tendons, cartilage, and
ligaments. Calcium diphosphate precipitates out of blood plasma
when diphosphate levels become abnormally high. If the calcium
concentration in blood plasma is 9.2 mg/dL and Ksp for calcium
diphosphate is 8.64 * 10-13, what minimum concentration of
diphosphate results in precipitation?

Md M.
Auburn University Main Campus

### Problem 126

Calculate the solubility of silver chloride in a solution that is
0.100 M in NH3.

Keenan M.
University of Miami

### Problem 127

Calculate the solubility of CuX in a solution that is 0.150 M in
NaCN. Ksp for CuX is 1.27 * 10-36.

Md M.
Auburn University Main Campus

### Problem 128

Aniline, C6H5NH2, is an important organic base used in the manufacture of dyes. It has Kb=4.3*10-10. In a certain manufacturing process, it is necessary to keep the concentration of C6H5NH3 + (aniline's conjugate acid, the anilinium ion) below 1.0 * 10-9 M in a solution that is 0.10 M in aniline. Find the
concentration of NaOH required for this process.

Keenan M.
University of Miami

### Problem 129

The Kb of hydroxylamine, NH2OH, is 1.10 * 10-8. A buffer solution
is prepared by mixing 100.0 mL of a 0.36 M hydroxylamine
solution with 50.0 mL of a 0.26 M HCl solution. Determine the
pH of the resulting solution.

Md M.
Auburn University Main Campus

### Problem 130

A 0.867-g sample of an unknown acid requires 32.2 mL of a 0.182 M barium hydroxide solution for neutralization. Assuming the acid is diprotic, calculate the molar mass of the acid.

Keenan M.
University of Miami

### Problem 131

A 25.0-mL volume of a sodium hydroxide solution requires 19.6mL of a 0.189 M hydrochloric acid for neutralization. A 10.0-mL volume of a phosphoric acid solution requires 34.9 mL of the sodium hydroxide solution for complete neutralization. Calculate the concentration of the phosphoric acid solution.

Md M.
Auburn University Main Campus

### Problem 132

Determine the mass of sodium formate that must be dissolved in 250.0 cm3 of a 1.4 M solution of formic acid to prepare a buffer solution with pH = 3.36.

Keenan M.
University of Miami

### Problem 133

What relative masses of dimethyl amine and dimethyl ammonium chloride do you need to prepare a buffer solution of pH = 10.43?

Md M.
Auburn University Main Campus

### Problem 134

You are asked to prepare 2.0 L of a HCN>NaCN buffer that has a pH of 9.8 and an osmotic pressure of 1.35 atm at 298 K. What masses of HCN and NaCN should you use to prepare the buffer?
(Assume complete dissociation of NaCN.)

Keenan M.
University of Miami

### Problem 135

What should the molar concentrations of benzoic acid and sodium benzoate be in a solution that is buffered at a pH of 4.55and has a freezing point of -2.0 C?
(Assume complete dissociation of sodium benzoate and a density of 1.01 g>mL for
the solution.)

Md M.
Auburn University Main Campus

### Problem 136

Derive an equation similar to the Henderson-Hasselbalch
equation for a buffer composed of a weak base and its
conjugate acid. Instead of relating pH to pKa and the relative
concentrations of an acid and its conjugate base (as the
Henderson-Hasselbalch equation does), the equation should
relate pOH to pKb and the relative concentrations of a base and
its conjugate acid.

Keenan M.
University of Miami

### Problem 137

Soap and detergent action is hindered by hard water so that
laundry formulations usually include water softeners-called
builders-designed to remove hard-water ions (especially Ca2+
and Mg2+) from the water. A common builder used in North
America is sodium carbonate. Suppose that the hard water used
to do laundry contains 75 ppm CaCO3 and 55 ppm MgCO3 (by
mass). What mass of Na2CO3 is required to remove 90.0% of
these ions from 10.0 L of laundry water?

Md M.
Auburn University Main Campus

### Problem 138

A 0.558-g sample of a diprotic acid with a molar mass of 255.8 g/mol
is dissolved in water to a total volume of 25.0 mL. The solution is
then titrated with a saturated calcium hydroxide solution.

a. Assuming that the pKa values for each ionization step are sufficiently
different to see two equivalence points, determine the
volume of added base for the first and second equivalence points.

b. The pH after adding 25.0 mL of the base is 3.82. Find the
value of Ka1.

c. The pH after adding 20.0 mL past the first equivalence point
is 8.25. Find the value of Ka2.

Keenan M.
University of Miami

### Problem 139

When excess solid Mg(OH)2 is shaken with 1.00 L of 1.0 M
NH4Cl solution, the resulting saturated solution has pH = 9.00.
Calculate the Ksp of Mg(OH)2.

Md M.
Auburn University Main Campus

### Problem 140

What amount of solid NaOH must be added to 1.0 L of a 0.10 M
H2CO3 solution to produce a solution with [H+] = 3.2 * 10-11 M?
There is no significant volume change as the result of adding the
solid.

Keenan M.
University of Miami

### Problem 141

Calculate the solubility of Au(OH)3 in (a) water and (b) 1.0 M nitric acid solution (Ksp = 5.5 * 10-46).

Md M.
Auburn University Main Campus

### Problem 142

Calculate the concentration of I-in a solution obtained by
shaking 0.10 M KI with an excess of AgCl(s).

Keenan M.
University of Miami

### Problem 143

What volume of 0.100 M sodium carbonate solution is required
to precipitate 99% of the Mg from 1.00 L of 0.100 M magnesium
nitrate solution?

Md M.
Auburn University Main Campus

### Problem 144

Find the solubility of CuI in 0.40 M HCN solution. The Ksp of CuI is 1.1 * 10-12, and the Kf for the Cu(CN)2-complex ion is 1 * 1024.

Keenan M.
University of Miami

### Problem 145

Find the pH of a solution prepared from 1.0 L of a 0.10 M solution of Ba(OH)2 and excess Zn(OH)2(s). The Ksp of Zn(OH)2 is 3 * 10-15, and the Kf of Zn(OH)4 2- is 2 * 10-15.

Md M.
Auburn University Main Campus

### Problem 146

What amount of HCl gas must be added to 1.00 L of a buffer solution that contains [acetic acid] = 2.0 M and [acetate] = 1.0 M in order to produce a solution with pH = 4.00?

Keenan M.
University of Miami

### Problem 147

Without doing any calculations, determine if pH = pKa, pH 7 pKa,
or pH 6 pKa. Assume that HA is a weak monoprotic acid.
a. 0.10 mol HA and 0.050 mol of A- in 1.0 L of solution
b. 0.10 mol HA and 0.150 mol of A- in 1.0 L of solution
c. 0.10 mol HA and 0.050 mol of OH- in 1.0 L of solution
d. 0.10 mol HA and 0.075 mol of OH- in 1.0 L of solution

Md M.
Auburn University Main Campus

### Problem 148

A buffer contains 0.10 mol of a weak acid and 0.20 mol of its
conjugate base in 1.0 L of solution. Determine whether or not
each addition exceeds the capacity of the buffer.

a. adding 0.020 mol of NaOH
b. adding 0.020 mol of HCl
c. adding 0.10 mol of NaOH
d. adding 0.010 mol of HCl

Keenan M.
University of Miami

### Problem 149

Consider three solutions:

i. 0.10 M solution of a weak monoprotic acid
ii. 0.10 M solution of strong monoprotic acid
iii. 0.10 M solution of a weak diprotic acid

Each solution is titrated with 0.15 M NaOH. Which quantity is the same for all three solutions?

a. the volume required to reach the final equivalence point
b. the volume required to reach the first equivalence point
c. the pH at the first equivalence point
d. the pH at one-half the first equivalence point

Md M.
Auburn University Main Campus

### Problem 150

Equal volumes of two monoprotic acid solutions (A and B) are titrated with identical NaOH solutions. The volume needed to reach the equivalence point for solution A is twice the volume required to reach the equivalence point for solution B, and the pH at the equivalence point of solution A is higher than the pH at the equivalence point for solution B. Which statement is true?

a. The acid in solution A is more concentrated than in solution B
and is also a stronger acid than that in solution B.

b. The acid in solution A is less concentrated than in solution B
and is also a weaker acid than that in solution B.

c. The acid in solution A is more concentrated than in solution B
and is also a weaker acid than that in solution B.

d. The acid in solution A is less concentrated than in solution B
and is also a stronger acid than that in solution B.

Keenan M.
University of Miami

### Problem 151

Describe the solubility of CaF2 in each solution compared to its solubility in water.

a. in a 0.10 M NaCl solution
b. in a 0.10 M NaF solution
c. in a 0.10 M HCl solution

Md M.
Auburn University Main Campus

### Problem 152

Name a compound that you could add to a solution of each of the following compounds to make a buffer. Explain your reasoning in complete sentences.

a. acetic acid
b. sodium nitrite
c. ammonia
d. potassium formate

Keenan M.
University of Miami

### Problem 153

Derive the Henderson-Hasselbalch equation as a group. Take turns having each group member write and explain the next step in the derivation.

Check back soon!

### Problem 154

With group members acting as atoms or ions, act out the reaction that occurs when HCl is added to a buffer solution composed of HC2H3O2 and NaC2H3O2. Write out a script for a narrator that describes the processes that occur, including how the buffer keeps the pH approximately the same even though a strong acid is added.

Keenan M.
University of Miami

### Problem 155

A certain town gets its water from an underground aquifer that contains water in equilibrium with calcium carbonate limestone.

a. What is the symbol for the equilibrium constant that describes calcium carbonate dissolving in water? What is the numerical value?

b. Calculate the molar solubility of calcium carbonate.

c. If you were a resident of this town and an entire coffee cup of water (about 200 mL) evaporated on your desk over spring break, how many grams of calcium carbonate would be left behind?

d. If you wanted to clean out your coffee cup, would it be better to use an acidic or a basic cleaning solution? Why?

Md M.
Auburn University Main Campus

### Problem 156

Have each group member look up the Ksp for a different compound. Calculate the molar solubility. Do the numerical values suggest that the compound is soluble or insoluble? Compare answers with the solubility rules from Chapter 8, and have each group member present his or her findings to the group.

Keenan M.
University of Miami

### Problem 157

A base is known to be one of the three listed in the table. You are given a sample of the base and asked to identify it. To do so, you dissolve 0.30 g of the base in enough water to make 25.0 mL of the basic solution. You then titrate the solution with 0.100 M HCl and record the pH as a function of the added acid, resulting in the titration curve that follows. Examine the table and the titration curve and answer the questions.

Md M.
Auburn University Main Campus