Name: ________________________
Hour: ____ Date: ___________
AP Chemistry: 17HW
Directions: Complete the following problems.
Calculate the pH of each of the following solutions.
1A.
i. 0.20 M propanoic acid
ii. 0.30 M potassium propanoate (KCH3CH2COO)
iii. a mixture of 0.20 M propanoic acid and 0.30 M KCH3CH2COO
1B.
i. 0.40 M hydroxylamine (HONH2)
ii. 0.30 M HONH3Cl
iii. a mixture containing 0.40 M HONH2 and 0.30 M HONH3Cl
ANSWERS:
1Ai. 2.79
1Aii. 9.18
1Aiii. 5.06
1Bi. 9.82
1Bii. 3.28
1Biii. 6.17
1C.
a mixture containing 0.20 M nitrous acid and 0.50 M lithium nitrite
1D.
a mixture containing 0.80 M hydrofluoric acid and 0.60 M sodium fluoride
1E.
a buffer soln containing 19 g of benzoic acid and 43 g of sodium benzoate in 260 mL of solution
1F.
a buffer solution in which 95 g of ammonium bromide are added to 2.00 L of a 0.85 M solution of
ammonia (assume a negligible volume change upon adding the ammonium bromide)
ANSWERS:
1C.
3.74
1D. 3.04
1E.
4.48
1F. 9.50
2A. Because it is tied to the respiratory system, the carbonate buffer system is considered to be the most
important buffer system in regulating human blood pH. Determine the ratio of the concentration of
carbonic acid to that of bicarbonate ion at a blood pH of 7.41.
2B. The phosphate buffer system helps regulate the pH of fluids in cells. Determine the ratio of the
concentration of H2PO4– to that of HPO42– at a pH of 7.17.
3. Consider the titration of 50.0 mL of 0.25 M chloric acid by 0.25 M sodium hydroxide. Determine the pH of the
mixture after the following volumes of base have been added. Assume additive volumes.
i. 0.0 mL
ii. 35.0 mL
iii. 49.3 mL
iv. 50.0 mL
v. 50.7 mL
ANSWERS:
2A. 1 : 11
2B. 0.92 : 1
3i. 0.60
3ii. 1.36
3iii.
3iv.
2.71
7.00
3v. 11.24
Calculate the pH at the equivalence point for the following titrations.
4A.
50. mL of 0.12 M benzoic acid titrated by 0.10 M potassium hydroxide
4B.
50. mL of 0.15 M ethylamine titrated by 0.20 M chloric acid
5. Determine the pH of a mixture of 50.0 mL of 0.40 M acetic acid and 96.2 mL of 0.20 M sodium hydroxide.
ANSWERS:
4A. 8.47
4B.
5.94
5.
6.15
6A.
A few drops of a weakly acidic indicator HIn (Ka = 1.0 x 10–10) are placed in a nitric acid solution. The
indicator itself is green in color, while its conjugate base is yellow.
i. State the initial color of the solution. Explain.
ii. The solution is titrated with a strong base. At what pH will the color change? Justify your answer.
iii. What color will the solution be after excess base has been added? Explain.
6B. An indicator has a pKa of 4.25. If the indicator solution changes color when 13% of the indicator has been
converted into its conjugate base, determine the pH at which this change becomes visible.
Write a dissolution equation and the corresponding solubility product expression for each of the following solids.
7A.
AgCl
7B.
Au2CO3
Co(OH)3
SrSO4
Ba3(PO4)2
PbI2
ANSWER:
6B.
3.42
Find the Ksp value for each of the following.
8A.
barium oxalate, BaC2O4, given that its solubility is 0.034 g/L
8B.
lead(II) chloride, given that [Pb2+] in a saturated solution is 0.016 M
8C.
the ionic compound M3X2, which has a molar mass of 308 g and a solubility of 1.9 x 10–4 g/L
Calculate the solubility of each compound, in moles per liter.
9A.
aluminum phosphate
9B.
a solution of cobalt(III) hydroxide buffered at a pH of 4.5
ANSWERS:
8A.
8B.
2.3 x 10–8
1.6 x 10–5
8C.
9.7 x 10–30
9A.
9B.
7.9 x 10–10 M
5.1 x 10–16 M
9C.
a solution of cobalt(III) hydroxide buffered at a pH of 11.5
Determine the solubility of silver chromate in each of the following solutions.
10A.
0.10 M silver nitrate
10B.
0.20 M sodium chromate
11A. Which of the substances in Q9 show increased solubility as the solution becomes more acidic? If the
solubility increases, write a “proper” reaction equation for the reaction that occurs.
11B. In each of the following groups, choose the substance whose solubility will depend on pH.
i.
CaF2
CaBr2
iii.
Fe(ClO3)2
FeCO3
ii.
Ni(OH)2
NiCl2
iv.
Pb(NO3)2
Pb(CN)2
11C. Briefly state what the answers to Q11B have in common and state how their solubilities will vary with pH.
ANSWERS:
9C.
5.1 x 10–37 M
10A.
1.1 x 10–10 M
10B.
1.7 x 10–6 M
AP Chemistry: Aqueous-Equilibrium Constants
Ka Values for Selected Acids at 25oC
Acid Name
Formula
Ka1
Acetic
Benzoic
Carbonic
Formic
Hydroazoic
Hydrocyanic
Hydrofluoric
Hypochlorous
Hypoiodous
Nitrous
Phenol
Phosphoric
Propanoic
CH3COOH
C6H5COOH
H2CO3
HCOOH
HN3
HCN
HF
HClO
HIO
HNO2
C6H5OH
H3PO4
CH3CH2COOH
1.8 x 10–5
6.3 x 10–5
4.3 x 10–7
1.8 x 10–4
1.9 x 10–5
4.9 x 10–10
6.8 x 10–4
3.0 x 10–8
2.3 x 10–11
4.5 x 10–4
1.3 x 10–10
7.1 x 10–3
1.3 x 10–5
Kb Values for Selected Bases at 25oC
Base Name
Formula
Kb
Ammonia
Aniline
Ethylamine
Hydroxylamine
Methylamine
Trimethylamine
NH3
C6H5NH2
CH3CH2NH2
HONH2
CH3NH2
(CH3)3N
1.8 x 10–5
4.3 x 10–10
6.4 x 10–4
1.1 x 10–8
4.4 x 10–4
6.4 x 10–5
Ksp Values for Selected Ionic Solids at 25oC
Substance Name
Formula
Ksp
Aluminum phosphate
Cobalt(III) hydroxide
Iron(III) hydroxide
Lead(II) chloride
Mercury(I) chloride
Silver chromate
Silver iodide
Silver phosphate
AlPO4
Co(OH)3
Fe(OH)3
PbCl2
Hg2Cl2
Ag2CrO4
AgI
Ag3PO4
6.3 x 10–19
1.6 x 10–44
4.0 x 10–38
1.6 x 10–5
1.1 x 10–18
1.1 x 10–12
1.5 x 10–16
1.8 x 10–18
Ka2
Ka3
5.6 x 10–11
6.2 x 10–8
4.5 x 10–13