worked solutions: chapter 9 acids and bases

Worked solutions: Chapter 9 Acids and bases

© Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008.

This page from the Chemistry: For use with the IB Diploma Programme SL Teacher’s Resource may be reproduced for classroom use.

Section 9.1 Exercises

1 a HClO4(aq) + H2O(l) → ClO4–(aq) + H3O+(aq)

b CH3NH2(aq) + H2O(l) → CH3NH3+(aq) + OH–(aq)

2 a HBr, HSO4–, NH4

+, HClO4, H2CO3

b I–, PO43–, NH3, SO4

2–, H2O

3 a HSO4–(aq) + H2O(l) → H2SO4(aq) + OH–(aq)

b HSO4–(aq) + H2O(l) → SO4

2–(aq) + H3O+(aq)

4 As an acid: HCO3–(aq) + OH–(aq) → CO3

2–(aq) + H2O(l)

As a base: HCO3–(aq) + HCl(aq) → H2CO3(aq) + Cl–(aq)

5 a HCO3–/CO3

2– and HCl/Cl–

b HSO4–/SO4

2– and NH4+/NH3

c HCl/Cl– and H3O+/H2O

6 a Base

b Base

c Neither

d Base

e Acid

7 H2SO4 ionizes when placed in water to produce HSO4–, SO4

2– and H3O+ ions. The presence of these ions allows the solution to conduct electricity.

H2SO4(l) + H2O(l) → HSO4–(aq) + H3O+(aq)

HSO4–(aq) + H2O(l) → SO4

2–(aq) + H3O+(aq)

8 Acids: HF, NH4+, H3O+

Bases: Cl–, CO32–, NH3, SO4

2–

Both acid and base: HSO4–, HPO4

2–


1 Both bases and alkalis are proton acceptors; however, an alkali is a base that dissolves in water.

2 Not all acids are dangerous; for example, vinegar is a dilute solution of a weak acid, and foods such as lemons and limes contain acids. A concentrated solution of lemon juice is unlikely to be dangerous; however, strong acids such as hydrochloric acid and sulfuric acid should always be treated with caution. Concentrated solutions of strong acids are most certainly dangerous. The concentration of H3O+ ions in solution governs the danger of the solution.




3 a 2Al(s) + 6HF(aq) → 2AlF3(aq) + 3H2(g)

b 2KHCO3(s) + H2SO4(aq) → K2SO4(aq) + 2H2O(l) + 2CO2(g)

c Fe2O3(s) + 6HNO3(aq) → 2Fe(NO3)3(aq) + 3H2O(l)

d Ca(OH)2(aq) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)

e NH3(aq) + HNO3(aq) → NH4NO3(s)

4 You could add a few drops of an acid–base indicator such as litmus indicator to a sample of the solution. If the litmus indicator turns blue, the solution is a base; if it turns red, it is an acid.

You could add a piece of magnesium ribbon to a sample of the solution. If a reaction occurs the solution is an acid. Similarly a reaction with calcium carbonate would indicate that the solution is an acid.

5 a pH Bromothymol blue Phenolphthalein Methyl orange i 3 yellow colourless red ii 8 blue colourless yellow iii 11 blue pink yellow

b Bromothymol blue is yellow for pH < 7.

Methyl orange is yellow for pH > 4.

∴ pH range is 4–7.


1 A concentrated acid solution has a greater number of H3O+ ions in a given volume of solution than a dilute solution. A 6.0 mol dm–3 solution of HCl would be considered to be concentrated, while a 0.1 mol dm–3 solution of HCl is dilute.

2 An acid such as HCl is a strong acid, whereas an acid such as CH3COOH is a weak acid. If the concentrations of a strong acid and a weak acid are equal, the strong acid will have a greater H3O+ concentration than the weak acid.

3 A strong acid dissociates in water to a greater extent than does a weak acid.

HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)

CH3COOH(aq) + H2O(l) ⇔ H3O+(aq) + CH3COO–(aq)

4 a HCl is a strong acid. It is therefore completely dissociated in water to produce many H+ and Cl– ions. Many ions produce high conductivity. CH3COOH is a weak acid, and is only partially dissociated. It has fewer H+ and CH3COO– ions, and therefore lower conductivity.

b HNO3 is a strong monoprotic acid, so there is 1 H+ ion per molecule. H2SO4 is a strong diprotic acid, with 2 H+ ions per molecule; therefore, there are more ions in a given volume of an H2SO4 solution (higher H+ concentration), and it has higher conductivity.




5 a NaOH completely dissociates: NaOH(s) ⎯⎯ →⎯ OH2 Na+(aq) + OH–(aq)

NH3 partially dissociates: NH3(g) + H2O(l) ⇔ NH4+(aq) + OH–(aq)

There are more ions in a given volume of NaOH, so it has better conductivity and higher OH– concentration.

b KOH completely dissociates: KOH(s) ⎯⎯ →⎯ OH2 K+(aq) + OH–(aq)

Na2CO3 completely dissociates: Na2CO3(s) ⎯⎯ →⎯ OH2 2Na+(aq) + CO32–(aq)

There are more ions in a given volume of solution for Na2CO3, so it has higher conductivity.

CO3– is a weak base, producing a low concentration of OH– ions in solution.

CO32–(aq) + H2O(l) ⇔ HCO3

–(aq) + OH–(aq)

So the concentration of OH– ions is higher for KOH.

6 a HNO3(aq) + H2O(l) → H3O+(aq) + NO3–(aq)

b HCN(g) + H2O(l) ⇔ H3O+(aq) + CN–(aq)

c KOH(s) ⎯⎯ →⎯ OH2 K+(aq) + OH–(aq)

d KF(s) H2O⎯ → ⎯ K+(aq) + F–(aq) and F–(aq) + H2O(l) ⇔ HF(aq) + OH–(aq)

7 a Good conductor. Complete dissociation of a strong base to yield many ions in solution.

KOH(s) ⎯⎯ →⎯ OH2 K+(aq) + OH–(aq)

b Good conductor. Complete dissociation of a strong acid to yield many ions in solution.

HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq)

c Poor conductor. Partial dissociation of a weak base to yield few ions in solution.

NH3(g) + H2O(l) ⇔ NH4+(aq) + OH–(aq)

d Poor conductor. Partial dissociation of a weak acid to yield few ions in solution.

H2CO3(aq) + H2O(l) ⇔ HCO3–(aq) + H3O+(aq)

8 An amphiprotic substance can act as either an acid or a base, i.e. it can donate or accept an H+ ion.

9 The electrical conductivity of solutions of equal concentrations could be measured. The solution with the higher conductivity contains the stronger base.

10 The strength of two different acids could be compared by: 1 measuring the conductivity of solutions of equal concentrations—the solution with the

higher conductivity contains the stronger acid

2 comparing the rate of reaction of solutions of equal concentration with a carbonate compound, such as calcium carbonate—the stronger acid will react at a greater rate

3 comparing the rate of reaction of equal concentration solutions with magnesium ribbon—the stronger acid will react at a greater rate




4 measuring the temperature change (very accurately) of the reaction between solutions of equal concentration with a solution of NaOH—the stronger acid will produce a greater temperature change than the weaker acid.

5 measuring the pH of solutions of equal concentration. The stronger acid will have a lower pH.

11 Acid rain occurs in areas of emission of oxides of nitrogen and sulfur, which occur in high-temperature situations such as in motor vehicles and at industrial sites, both of which are more common in areas of high population density and highly industrialized areas.


1 A neutral solution will have a pH of about 7, while an acidic solution will have a pH < 7.

2 a Tomato juice is more acidic than rainwater.

b Milk is more acidic than seawater.

c Vinegar is more acidic than orange juice.

3 a Seawater is more alkaline than rainwater.

b Black coffee is more alkaline than orange juice.

c Household ammonia is more alkaline than brass polish.

4 Solution B will have a pH that is 3 pH units higher than solution A, so its pH will be 5 (higher pH means less acidic).

5 A solution with a pH 100 times more acidic than pH of 12, will have a pH of 10 (lower pH means more acidic).

6 A drop of 2 pH units = 102 = 100-fold change in acidity, so urine is 100 times more acidic than blood.

7 [H+(aq)] of a solution with a pH = 2 is 10–2 mol dm–3

8 pH = 5

9 Sodium hydroxide is a stronger base than ammonia. Sodium hydroxide dissociates completely in water, whereas ammonia undergoes partial dissociation, so the concentration of OH– is higher (and [H3O+] is lower) in a 1.0 mol dm–3 solution of sodium hydroxide than a 1.0 mol dm–3 solution of ammonia, and the pH is higher.

NaOH(s) ⎯⎯ →⎯ OH2 Na+(aq) + OH–(aq)

NH3(g) + H2O(l) ⇔ NH4+(aq) + OH–(aq)




Chapter 9 Review questions

1 A Brønsted–Lowry acid is a H+ donor; a Lewis acid is an electron pair acceptor. HCl is a Brønsted–Lowry acid with an H that can be donated, while BF3 is a Lewis acid that can accept an electron pair (B is electron deficient).

2. a H2O: Brønsted–Lowry acid, Brønsted–Lowry base; Lewis base

b NH3: Brønsted–Lowry base; Lewis base

c HCl: Brønsted–Lowry acid; H+ is a Lewis acid and Cl– is a Lewis base

d F–: Brønsted–Lowry base; Lewis base

e BF3: Lewis acid

f OH–: Brønsted–Lowry base; Lewis base

3 a ClO3–, S2–, NH3, OH–

b H2CO3, H2S, HClO4, H3PO4

4 a H2C2O4/HC2O4– and H3O+/H2O

b H2O/OH– and HCN/CN–

c CH3COOH/CH3COO– and HS–/S2–

d HSO4–/SO4

2– and HF/F–

5 a A

b B

c B

d AB

e A

f AB

g A

h A

i N

6 Acid or base Use Sulfuric acid Car batteries Sodium carbonate Washing powders Magnesium hydroxide Antacid powders Ammonia Window cleaner Phosphoric acid Fertilizer manufacture Zinc hydroxide Watch batteries




7 a HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq) b Ca(OH)2(s) H2O⎯ → ⎯ Ca2+(aq) + 2OH–(aq)

c H2SO4(l) + H2O(l) → HSO4–(aq) + H3O+(aq)

HSO4–(aq) + H2O(l) ⇔ SO4

2–(aq) + H3O+(aq)

d H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)

e Ba(OH)2(aq) + 2HCl(aq) → BaCl2(aq) + 2H2O(l)

8 a Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

b The reaction does not involve the transfer of a proton from one species to another. (This is in fact a redox reaction involving transfer of electrons.)

9 a i It does not dissociate readily or completely in water.

ii It is able to donate or accept H+ ions.

b HCrO4–(aq) + OH–(aq) → CrO4

2–(aq) + H2O(l)

HCrO4–(aq) + HCl(aq) → H2CrO4(aq) + Cl–(aq)

10

11 Measure the conductivity of the two solutions. If the unknown solution has a higher

conductivity than ethanoic acid, it is more acidic; if lower, it is less acidic.

12 a H2SO4, HSO4–, SO4

2–, H2O, H3O+

b H2O

c SO42–

13 NaOH, NH3, KCl, H2CO3, HSO4–, HCl

14 HCl is a strong, monoprotic acid: HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)

H2SO4 is a strong, diprotic acid: H2SO4(aq) + H2O(l) → HSO4–(aq) + H3O+(aq) and

HSO4–(aq) + H2O(l) → SO4

2–(aq) + H3O+(aq)

Greater concentration of H3O+ ions produced, hence lower pH than that of HCl.




CH3COOH is a weak, monoprotic acid: CH3COOH(aq) + H2O(l) ⇔ CH3COO–(aq) + H3O+(aq)

Lower concentration of H3O+ ions produced, hence higher pH than that of HCl.

15 a Ba(OH)2: there are 3 ions per unit of Ba(OH)2.

b NH3: only partially dissociates in water as it is a weak base.

c HCl (complete dissociation) and NaOH (complete dissociation).

16 a Tomato juice is the most acidic.

b Milk is most the acidic.

17 a Seawater is the most alkaline.

b Household ammonia is the most alkaline.

18 pH = 4

19 a Solution Y has an [H+] that is 100 times greater (more acidic) than solution X; it has pH = 5.

b Solution Y is acidic.

20 Solution B is 103 = 1000 times more alkaline than solution A.

Chapter 9 Test

Part A: Multiple-choice questions

Question Answer Explanation 1 A Like strong acids, weak acids react with carbonates, but more slowly, and they

conduct electricity to a much smaller extent than strong acids. 2 B The volume has changed from 10 cm3 to a total of 100 cm3. There has been a

tenfold decrease in the concentration of the hydrogen ion and this will lead to an increase of 1 pH unit.

3 D There is an increase in 2 pH units, thus there will be a decrease in the hydrogen ion concentration by a factor of 100.

4 D A strong acid will conduct electricity to a greater extent than a weak acid. A strong acid will have a higher concentration of H+ ions.

5 C H2O has the ability to both donate and accept protons; thus it has the ability to act as both an acid and a base.

6 C The substance with the highest pH will be the one which is a strong base, which is NaOH.

7 A A Brønsted–Lowry base is defined as a proton (H+) acceptor. 8 B Weak acids have higher pHs, react more slowly with both carbonates and metals

and require the same number of moles of base for neutralization as strong acids. 9 A The least acidic solution will have the highest pH and the lowest concentration of

H+. Thus, P is the least acidic, followed by S, then R and Q. 10 D In equation 1, H2SO4 is an acid and HSO4

– is the conjugate base. In equation 2, HSO4

– is an acid and its conjugate base is SO42–.




Part B: Short-answer questions

1 a Both NaOH and NH3 are capable of accepting protons (H+). (1 mark)

b i NaOH has greater conductivity because it is fully dissociated, therefore it has a greater number of ions to conduct.

(1 mark) ii Because NaOH is a stronger base than NH3, it would have pH >11.

(2 marks)

2 a The solution is alkaline because [OH–] > [H+]. (1 mark)

b HCO3– is acting as a base because it is accepting a proton (H+).

(2 marks)

3 a A strong acid is fully ionized or dissociated in solution:

HCl (aq) → H+(aq) + Cl–(aq)

A weak acid is only partly ionized or dissociated in solution:

H2CO3(aq) ⇔ H+(aq) + HCO3–(aq)

(4 marks) b The ratio of H+ ions in HCl : H2CO3 is 10 000 : 1.

(2 marks)

Part C: Data-based question

1 Strong acid: Acid 1—acid with high conductivity

Weak acid: Acid 2—acid with lower conductivity

A strong acid is fully dissociated into ions. As acid concentration increases, the number of ions in a given volume increases, and so does the conductivity.

A weak acid is only partially dissociated, producing fewer ions in solution. As acid concentration increases, the number of ions increases initially until an equilibrium is established. Then the concentration of ions becomes constant and the conductivity remains constant.

(6 marks) b Both reactions produce gas (H2). The acid that reacts more quickly, producing more

bubbles in a given time, is the strong acid. (1 mark)




Part D: Extended-response question

a Many explanations could be given, but some examples are:

Mg + H2SO4 → MgSO4 + H2

CaO + H2SO4 → CaSO4 + 2H2O

2NaOH + H2SO4 → Na2SO4 + 2H2O

CaCO3 + H2SO4 → CaSO4 + H2O + CO2 (3 marks)

b i A Brønsted–Lowry acid is an H+ donor and a Brønsted–Lowry base is an H+ acceptor. (1 mark)

ii NH3 + H2O ⇔ NH4+ + OH–

pH paper will turn blue; pH value of 10–12 . (3 marks)

c Acid Conjugate base Base Conjugate acid i H2SO4 HSO4

– HNO3 H2NO3+

ii H2O OH– CH3CH2NH2 CH3CH2NH3+

(2 mark)

d H2SO4 is a stronger acid than HNO3. This can be seen when H2SO4 gives a proton to HNO3 in equation c part i.

(2 marks)

e A strong acid is completely dissociated; a weak acid is only partially dissociated.

To distinguish between a strong acid and a weak acid, experiments could be performed using solutions of equal concentration.

Strong acid will give lower pH value or will have a higher conductivity or will give a faster reaction with a carbonate or a reactive metal.

(4 marks)

f i Acid A is stronger. It is 10 × 10 × 10 × 10 = 10 000 times more acidic than B.

This is because a difference of 1 pH unit = 10-fold difference in acidity. (3 marks)

ii To produce a solution of pH 6, a base could be added, or water could be added. (2 marks)

worked solutions: chapter 9 acids and bases

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