Review
Reflecting on Chapter 8
Summarize this chapter in the format of your
choice. Here are a few ideas to use as guidelines:
• Relate the microscopic properties of acids and
bases to their macroscopic properties.
• Identify conjugate acid-base pairs for selected
acid-base reactions, and compare their strengths.
• State the relationship among Ka , Kb , and Kw .
• Outline the relationship among [H3O+ ], pH,
[OH− ], and pOH.
• Describe two examples of buffer solutions in
your daily life, and explain how they function.
Reviewing Key Terms
For each of the following terms, write a sentence
that shows your understanding of its meaning.
hydronium ion (H3O+(aq)) conjugate acid-base
pair
monoprotic acids
polyprotic acids
ion product constant
pH
for water (Kw )
acid dissociation
pOH
constant (Ka )
percent dissociation
base dissociation
constant (Kb )
buffer solution
buffer capacity
acid-base titration curve equivalence point
Knowledge/Understanding
1. Give two examples of each of the following
acids and bases.
(a) Arrhenius acids
(b) Brønsted-Lowry bases
(c) Brønsted-Lowry bases that are not
Arrhenius bases
2. Classify each compound as a strong acid, strong
base, weak acid, or weak base.
(a) phosphoric acid, H3PO4 (used in cola
beverages and rust-proofing products)
(b) chromic acid, H2CrO4 (used in the
production of wood preservatives)
(c) barium hydroxide, Ba(OH)2 , a white, toxic
base (can be used to de-acidify paper)
(d) CH3NH2 , commonly called methylamine
(is responsible for the characteristic smell
of fish that are no longer fresh)
3. Write a chemical formula for each acid or base.
(a) the conjugate base of OH−
(b) the conjugate acid of ammonia, NH3
(c) the conjugate acid of HCO3−
(d) the conjugate base of HCO3−
4. Decide whether each statement is true or false,
and explain your reasoning.
(a) HBr is a stronger acid than HI.
(b) HBrO2 is a stronger acid than HBrO.
(c) H2SO3 is a stronger acid than HSO3− .
5. Arrange the following aqueous solutions in
order of pH, from lowest to highest:
2.0 mol/L HClO4 , 2.0 mol/L NaCl, 0.20 mol/L
CH3COOH, 0.02 mol/L HCl.
6. In each pair of bases, which is the stronger
base?
(a) HSO4−(aq) or SO42−(aq)
(b) S2−(aq) or HS−(aq)
(c) HPO42−(aq) or H2PO4−(aq)
(d) HCO3−(aq) or CO32−(aq)
7. (a) Use Appendix E to find the values of Ka for
hydrosulfuric acid, HS−(aq), and sulfurous
acid, HSO3−(aq) .
(b) Write equations for the base dissociation
constants of HS−(aq) and HSO3−(aq) .
(c) Calculate the value of Kb for each ion.
(d) Which is the stronger base, HS−(aq) or
HSO3−(aq) ? Explain.
8. While the pH of blood must be maintained
within strict limits, the pH of urine can vary.
The sulfur in foods, such as eggs, is oxidized in
the body and excreted in the urine. Does the
presence of sulfide ions in urine tend to
increase or decrease the pH? Explain.
9. Sodium methanoate, NaHCOO, and methanoic
acid, HCOOH, can be used to make a buffer
solution. Explain how this combination resists
changes in pH when small amounts of acid or
base are added.
10. Oxoacids contain an atom that is bonded to
one or more oxygen atoms. One or more of
these oxygen atoms may also be bonded to
hydrogen. Consider the following oxoacids:
HBrO3(aq) , HClO3(aq) , HClO4(aq) , and H2SO3(aq) .
(a) What factors are used to predict the
strengths of oxoacids?
(b) Arrange the oxoacids above in the order of
increasing acid strength.
Chapter 8 Acids, Bases, and pH • MHR
415
Inquiry
11. What is the pH of a mixture of equal
solution, what is the acid dissociation constant
for salicylic acid?
COOH
volumes of 0.040 mol/L hydrochloric acid
and 0.020 mol/L sodium hydroxide?
OH
12. Suppose that 15.0 mL of sulfuric acid just
neutralized 18.0 mL of 0.500 mol/L sodium
hydroxide solution. What is the concentration
of the sulfuric acid?
13. A student dissolved 5.0 g of vitamin C in
250 mL of water. The molar mass of ascorbic
acid is 176 g/mol, and its Ka is 8.0 × 10−5.
Calculate the pH of the solution. Note:
Abbreviate the formula of ascorbic acid
to HAsc.
14. Benzoic acid is a weak, monoprotic acid
(Ka = 6.3 × 10−5). Its structure is shown below.
Calculate the pH and the percent dissociation
of each of the following solutions of benzoic
acid. Then use Le Châtelier’s principle to
explain the trend in percent dissociation of the
acid as the solution becomes more dilute.
(a) 1.0 mol/L (b) 0.10 mol/L (c) 0.01 mol/L
COOH
Communication
19. List the oxoacids of bromine (HOBr, HBrO2 ,
HBrO3 , and HBrO4 ) in order of increasing
strength. What is the order of increasing
strength for the conjugate bases of these acids?
20. Consider the following acid-base reactions.
HBrO2(aq)+CH3COO−(aq)
CH3COOH(aq) + BrO−2(aq)
H2S(aq) + OH−(aq) HS−(aq) + H2O()
HS−(aq)+CH3COOH(aq)
H2S(aq) + CH3COO−(aq)
If each equilibrium lies to the right, arrange the
following compounds in order of increasing
acid strength: HBrO2 , CH3COOH, H2S, H2O.
21. Discuss the factors that can be used to predict
the relative strength of different oxoacids.
22. The formula of methyl red indicator can be
15. Hypochlorous acid, HOCl, is a weak acid that
is found in household bleach. It is made by
dissolving chlorine gas in water.
Cl2(g) + 2H2O() H3O+(aq) + Cl−(aq) + HOCl(aq)
(a) Calculate the pH and the percent
dissociation of a 0.065 mol/L solution
of hypochlorous acid.
(b) What is the conjugate base of hypochlorous
acid? What is its value for Kb ?
16. Calculate the pH of a 1.0 mol/L aqueous
solution of sodium benzoate. Note: Only the
benzoate ion affects the pH of the solution.
17. Calculate the pH of a 0.10 mol/L aqueous solu-
tion of sodium nitrite, NaNO2 . Note: Only the
nitrite ion affects the pH of the solution.
18. A student prepared a saturated solution of
salicylic acid and measured the pH of the
solution. The student then carefully evaporated
100 mL of the solution and collected the solid.
If the pH of the solution was 2.43, and 0.22 g
was collected after evaporating 100 mL of the
416 MHR • Unit 4 Chemical Systems and Equilibrium
abbreviated to HMr. Like most indicators,
methyl red is a weak acid.
HMr(aq) + H2O()
H3O+(aq) + Mr−(aq)
The change between colours (when the
indicator colour is orange) occurs at a pH
of 5.4. What is the equilibrium constant for
the reaction?
23. Gallic acid is the common name for
3,4,5-trihydroxybenzoic acid.
(a) Draw the structure of gallic acid.
(b) Ka for gallic acid is 3.9 × 10−5. Calculate Kb
for the conjugate base of gallic acid. Then
write the formula of the ion.
24. (a) Sketch the pH curves you would expect if
you titrated
• a strong acid with a strong base
• a strong acid with a weak base
• a weak acid with a strong base
(b) Congo red changes colour over a pH range of
3.0 to 5.0. For which of the above titrations
would Congo red be a good indicator to use?
(b) O2−/OH− and H2O/OH− 4.(a) H2S/HS− and NH3/NH4+
Making Connections
25. Citric acid can be added to candy to give a
sour taste. The structure of citric acid is
shown below.
O
H
O
O
H
O
H
C
H
O
C
C
C
C
C
H
O
H
O
H
H
(a) Identify the acidic hydrogen atoms that are
removed by water in aqueous solution. Why
do water molecules pull these hydrogen
atoms away, rather than other hydrogen
atoms in citric acid?
(b) Why does citric acid not form OH− ions in
aqueous solution, and act as a base?
(c) When citric acid and sodium hydrogen carbonate are used in bubble gum, the bubble
gum foams when chewed. Suggest a reason
why this happens.
26. (a) Imagine that you have collected a sample
of rainwater in your community. The pH of
your sample is 4.52. Unpolluted rainwater
has a pH of about 5.6. How many more
hydronium ions are present in your sample,
compared with normal rainwater? Calculate
the ratio of the concentration of hydronium
ions in your sample to the concentration of
hydronium ions in unpolluted rainwater.
(b) You have been invited to a community
meeting to explain your findings to local
residents. No one at the meeting has a
background in chemistry. In a paragraph,
write what you would say at this meeting.
(c) Suggest at least two possible factors that
could be responsible for the pH you measured. What observations would you want
to make, and what data would you want to
collect, to help you gain confidence that one
of these factors is responsible?
Answers to Practice Problems and Short Answers to
Section Review Questions
Practice Problems: 1.(a) chloride ion, Cl− (b) carbonate ion,
CO32− (c) hydrogen sulfate ion, HSO4− (d) hydrazine, N2H4
2.(a) nitric acid, HNO3 (b) water, H2O (c) hydronium ion,
H3O+ (d) carbonic acid, H2CO3 3.(a) HS−/H2S and H2O/OH−
(b) H2SO4/HSO4− and H2O/H3O+ 5.(a) 4.5 mol/L (b) 1.35 mol/L
(c) 0.02 mol/L (d) 0.0375 mol/L 6.(a) 3.1 mol/L
(b) 0.87 mol/L (c) 0.701 mol/L (d) 0.697 mol/L 7.(a) acidic
solution; [H3O+] = 0.479 mol/L (b) acidic solution;
[H3O+] = 1.98 mol/L 8. acidic solution; [H3O+] = 0.46 mol/L
9.(a) [H3O+] = 0.45 mol/L; [OH−] = 2.2 × 10−14 mol/L
(b) [OH−] = 1.1 mol/L; [H3O+] = 9.1 × 10−15 mol/L
10.(a) [H3O+] = 0.95 mol/L; [OH−] = 1.1 × 10−14 mol/L
(b) [OH−] = 0.024 mol/L ; [H3O+] = 4.2 × 10−13 mol/L
11. [HCl] = 0.18 mol/L 12. [Ca(OH)2] = 0.29 mol/L 13. acidic;
pH = 6.400 ; [OH−] = 2.51 × 10−8 mol/L 14. pOH = 2.1;
[OH−] = 8 × 10−3 mol/L 15. acidic;
[H3O+] = 1.9 × 10−5 mol/L; [OH−] = 5.2 × 10−10 ; pOH = 9.28
16. [H3O+] = [OH−] = 1.6 × 10−7 ; neutral
17. basic; pH = 8.19; [H3O+] = 6.5 × 10−9 mol/L;
[OH−] = 1.5 × 10−6 mol/L 18. [H3O+] = 1.9 × 10−3 mol/L;
[OH−] = 5.4 × 10−12 mol/L 19. pH = 2.41; 0.46% dissociation
20. Ka = 1.0 × 10−4 ; percent dissociation = 3.2%
21. pH = 2.65 22. Ka = 2.9 × 10−8 23. pH = 3.411 24. 1.2 g
25. pH = 4.82; [CO32−] = 4.7 × 10−11 mol/L 26. pH = 2.74
27. pH = 4.59; [HS−] = 2.6 × 10−5 mol/L 28. Ka = 1.0 × 10−14 ;
Ka2 =
[O2−][H3O+]
[OH−]
29. pH = 11.14 30. pH = 10.50
31. Kb = 1.6 × 10−6 32. [OH−] = 1.0 × 10−2 mol/L; pOH = 1.98
33. [OH−] = 3.8 × 10−3 mol/L; percent dissociation = 1.7%
34. [NH3] = 2.8 × 10−2 mol/L
35. C6H5O− > C6H5COO− > HCOO− > F− 36. NH4+
37. pH = 9.10 38. as an acid:
HSO3−(aq) + H2O() H3O+(aq) + SO32−(aq) ; as a base:
HSO3−(aq) + H2O() OH−(aq) + H2SO3(aq)
Section Review: 8.1:
H2PO4−(aq) + H3O+(aq) ;
1.(a) H3PO4(aq) + H2O(aq) H2PO4−(aq) + H2O(aq) HPO42−(aq) + H3O+(aq) ;
HPO42−(aq) + H2O(aq) PO43−(aq) + H3O+(aq)
HPO42−(aq) + H3O+(aq) ;
(b) as an acid: H2PO4−(aq) + H2O(aq) H3PO4(aq) + H2O(aq)
as a base: H2PO4−(aq) + H3O+(aq) (c) H3PO4(aq) is much stronger (although it is still a weak
acid) than H2PO4−(aq) . 2. C6H4NH2COO− 3.(a) B(OH)3/B(OH)4−
and H2O/H3O+ (b) weak 4.(a) weak acid
(b) strong acid (c) strong base (d) weak base
8.2: 1.(a) 4.43 (b) 2.70 × 10−10 (c) 9.57 (d) acidic (e) 3.9 × 10−11
(f) 2.6 × 10−4 (g) 3.59 (h) basic (i) 1.4 × 10−13 (j) 12.85 (k) 1.15
(l) basic (m) 8 × 10−6 (n) 5.1 (o) 1 × 10−9 (p) acidic 2. pH = 2.39
3. Ka = 2.5 × 10−4 4. pH = 1.70;
[HOOCCOO−] = 5.4 × 10−5 mol/L
5. [H2CO3] = 2.3 × 10−3 mol/L
8.3: 1. [OH−] = 7.6 × 10−3 mol/L; pH = 11.88 2. pH = 8.80
3. OBr− ; Kb = 3.6 × 10−6 8.4: 4. phenolphthalein 5. pH ∼ 5.0
Chapter 8 Acids, Bases, and pH • MHR
417