Topic 2: Acids and bases

Turn every volume into litres and every amount into moles before anything else, and scale an aliquot back up to the whole flask.

A lesson for this topic

Quick checks: multiple choice

Question 1

In the Brønsted-Lowry model, an acid is a species that

Question 2

What is the conjugate base of the hydrogensulfate ion, HSO₄⁻?

Question 3

Which species can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base in water?

Question 4

Consider the reaction HCO₃⁻(aq) + NH₃(aq) ⇌ CO₃²⁻(aq) + NH₄⁺(aq). Which species acts as the Brønsted-Lowry acid in the forward reaction, and what is its conjugate base?

Question 5

Phosphoric acid, H₃PO₄, is described as a triprotic acid because

Question 6

Ethanoic acid, CH₃COOH, has four hydrogen atoms in each molecule, yet it is a monoprotic acid. Why?

Question 7

In which reaction does water act as a Brønsted-Lowry base?

Question 8

A strong acid is an acid that

Question 9

Which solution is both weak and concentrated?

Question 10

Methanoic acid ionises in water: HCOOH(aq) + H₂O(l) ⇌ HCOO⁻(aq) + H₃O⁺(aq). Which is its acidity constant expression?

Question 11

Acidity constants of four weak monoprotic acids at 25 °C
Acid	Ka at 25 °C
methanoic acid	1.8 × 10⁻⁴
benzoic acid	6.3 × 10⁻⁵
ethanoic acid	1.8 × 10⁻⁵
hypochlorous acid	3.0 × 10⁻⁸

The table gives Ka values for four weak monoprotic acids. A student makes a 0.10 mol L⁻¹ solution of each at 25 °C. Which lists the solutions in order of increasing pH?

Question 12

Acidity constants of three weak acids at 25 °C
Acid	Ka at 25 °C
hydrofluoric acid, HF	6.8 × 10⁻⁴
ethanoic acid, CH₃COOH	1.8 × 10⁻⁵
hydrocyanic acid, HCN	6.2 × 10⁻¹⁰

The table gives Ka values for three weak acids at 25 °C; hydrochloric acid is a strong acid. A student makes 0.10 mol L⁻¹ solutions of four sodium salts. Which solution has the highest pH?

Question 13

Which equation explains why a solution of ammonium chloride, NH₄Cl, is acidic?

Question 14

Which 0.10 mol L⁻¹ solution has a pH above 7 at 25 °C?

Question 15

A solution of baking soda, sodium hydrogencarbonate, is slightly basic. Which equation shows the hydrogencarbonate ion acting as a Brønsted-Lowry base in water?

Question 16

pH of 0.10 mol L⁻¹ solutions of three unlabelled salts at 25 °C
Solution	pH
X	11.6
Y	7.0
Z	5.1

Three bottles in a Geraldton school laboratory have lost their labels. They hold 0.10 mol L⁻¹ solutions of sodium carbonate, potassium nitrate and ammonium chloride. The table shows the pH of each at 25 °C. Which matches the solutions to the salts?

Question 17

Two acid solutions tested at 25 °C
Solution	pH	Electrical conductivity (mS cm⁻¹)
0.10 mol L⁻¹ hydrochloric acid	1.00	39
0.10 mol L⁻¹ ethanoic acid	2.88	0.5

A student measures the pH and electrical conductivity of two acid solutions, as shown in the table. Which explains why the ethanoic acid conducts much less well?

Question 18

A student dilutes a 0.10 mol L⁻¹ ethanoic acid solution tenfold with water at 25 °C. What happens to the Ka of ethanoic acid and to the percentage of its molecules that are ionised?

Question 19

Which mixture, with both substances at 0.20 mol L⁻¹, forms a buffer solution?

Question 20

A buffer contains ethanoic acid and sodium ethanoate: CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq). A few drops of hydrochloric acid are added. Which describes what happens?

Question 21

Buffer P contains 0.50 mol L⁻¹ ethanoic acid and 0.50 mol L⁻¹ sodium ethanoate. Buffer Q contains 0.050 mol L⁻¹ of each. Both have the same pH. Which statement is correct?

Question 22

A buffer contains ammonia and ammonium chloride: NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq). A small amount of sodium hydroxide solution is added. Which describes the response of the buffer?

Question 23

Human blood is kept near pH 7.4 by a buffer based on carbonic acid. Which is the conjugate acid-base pair in this buffer?

Question 24

What is the pH of 0.0250 mol L⁻¹ hydrochloric acid at 25 °C?

Question 25

What is the pH of 0.0150 mol L⁻¹ sodium hydroxide solution at 25 °C? Kw = 1.00 × 10⁻¹⁴.

Question 26

A water treatment technician in Kalgoorlie makes 0.00350 mol L⁻¹ barium hydroxide solution, Ba(OH)₂. What is its pH at 25 °C? Kw = 1.00 × 10⁻¹⁴.

Question 27

What is the pH of 0.0400 mol L⁻¹ sulfuric acid at 25 °C? Assume that both protons of each H₂SO₄ are fully donated.

Question 28

A 10.0 mL sample of nitric acid with a pH of 1.50 is diluted with distilled water to a total volume of 250.0 mL. What is the pH of the diluted solution at 25 °C?

Question 29

In a Bunbury school laboratory, 25.0 mL of 0.120 mol L⁻¹ hydrochloric acid is mixed with 30.0 mL of 0.150 mol L⁻¹ sodium hydroxide solution at 25 °C. What is the pH of the mixture? Kw = 1.00 × 10⁻¹⁴.

Question 30

A student mixes 20.0 mL of 0.0500 mol L⁻¹ sulfuric acid with 30.0 mL of 0.0500 mol L⁻¹ potassium hydroxide solution at 25 °C. Assume both protons of each H₂SO₄ are fully donated. What is the pH of the mixture? Kw = 1.00 × 10⁻¹⁴.

Question 31

A water sample from a Swan River estuary site has a pH of 9.40 at 25 °C. What is its hydroxide ion concentration? Kw = 1.00 × 10⁻¹⁴.

Question 32

Solution J has a pH of 3 and solution K has a pH of 5 at 25 °C. How does [H₃O⁺] in J compare with [H₃O⁺] in K?

Question 33

Water self-ionises: 2H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq), with Kw = 1.00 × 10⁻¹⁴ at 25 °C. Which is true of pure water at 25 °C?

Question 34

A student keeps adding distilled water to a sodium hydroxide solution of pH 12 at 25 °C. What happens to the pH?

Question 35

What mass of sodium hydroxide (M = 40.00 g mol⁻¹) must be dissolved in water to make 500.0 mL of solution with a pH of 12.30 at 25 °C? Kw = 1.00 × 10⁻¹⁴.

Question 36

Acid-base indicators such as methyl orange are best described as

Question 37

In an acid-base titration, the end point is

Question 38

Vinegar, a solution of ethanoic acid, is titrated with sodium hydroxide solution. Indicator ranges: methyl orange 3.1 to 4.4, methyl red 4.4 to 6.2, bromothymol blue 6.0 to 7.6, phenolphthalein 8.3 to 10.0. Which indicator is the most suitable?

Question 39

A student titrates household ammonia solution with hydrochloric acid. The pH at the equivalence point is 5.3. Indicator ranges: methyl orange 3.1 to 4.4, methyl red 4.4 to 6.2, bromothymol blue 6.0 to 7.6, phenolphthalein 8.3 to 10.0. Which indicator changes colour closest to the equivalence point?

Question 40

The graph shows a titration curve recorded with a pH meter as 0.100 mol L⁻¹ sodium hydroxide is added to a 25.00 mL aliquot of an acid. What does the curve show?

Question 41

Which rinsing procedure gives the most accurate results in an acid-base titration?

Question 42

A student rinses her pipette with distilled water, but not with the acid, before taking 20.00 mL aliquots of hydrochloric acid to titrate against standard sodium hydroxide solution. What is the effect on her result?

Question 43

Which substance is suitable for use as a primary standard?

Question 44

Titres recorded for 20.00 mL aliquots
Titration	Titre (mL)
Rough	25.10
1	24.30
2	24.65
3	24.25
4	24.33

A student's titres are in the table. She uses titres that agree within 0.10 mL as her concordant results. What titre should she use in her calculation?

Question 45

A 25.00 mL aliquot of sodium hydroxide solution is titrated with 0.0500 mol L⁻¹ sulfuric acid; the mean titre is 21.40 mL. The reaction is H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l). What is the concentration of the sodium hydroxide?

Question 46

A student dissolves a sample of sodium carbonate and makes it up to 250.0 mL. A 25.00 mL aliquot needs 18.60 mL of 0.1050 mol L⁻¹ hydrochloric acid, with methyl orange. The reaction is Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g). M(Na₂CO₃) = 105.99 g mol⁻¹. What mass of sodium carbonate was dissolved?

Question 47

A 0.500 g sample of crushed limestone from a quarry near Perth is added to 50.00 mL of 0.200 mol L⁻¹ hydrochloric acid (an excess). When the reaction stops, the leftover acid needs 22.80 mL of 0.100 mol L⁻¹ sodium hydroxide. Reactions: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) and HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). M(CaCO₃) = 100.09 g mol⁻¹. Assuming only the calcium carbonate reacts with acid, what is its percentage by mass in the sample?

Question 48

A student dissolves 1.500 g of crude aspirin (C₉H₈O₄, a monoprotic acid, M = 180.15 g mol⁻¹) and makes it up to 250.0 mL. 25.00 mL aliquots need a mean titre of 19.85 mL of 0.0400 mol L⁻¹ sodium hydroxide. Aspirin reacts with NaOH in a 1 to 1 ratio. What is the percentage purity of the aspirin?

Question 49

A technician dilutes 10.00 mL of concentrated hydrochloric acid to 500.0 mL with distilled water. 20.00 mL aliquots of the diluted acid react with a mean titre of 22.40 mL of 0.1000 mol L⁻¹ sodium carbonate solution. The reaction is Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g). What is the concentration of the concentrated acid?

Question 50

A student dilutes 25.0 mL of 0.500 mol L⁻¹ hydrochloric acid to 250.0 mL with distilled water. She then mixes 40.0 mL of the diluted acid with 60.0 mL of 0.0400 mol L⁻¹ sodium hydroxide solution at 25 °C. What is the pH of the final mixture? Kw = 1.00 × 10⁻¹⁴.

Question 51

50.0 mL of hydrochloric acid of pH 1.00 is mixed with 50.0 mL of nitric acid of pH 2.00 at 25 °C. What is the pH of the mixture?

Question 52

A student makes a primary standard by dissolving 1.576 g of oxalic acid dihydrate, H₂C₂O₄.2H₂O (M = 126.07 g mol⁻¹), and making it up to 250.0 mL. 20.00 mL aliquots need a mean titre of 21.15 mL of sodium hydroxide solution. The reaction is H₂C₂O₄(aq) + 2NaOH(aq) → Na₂C₂O₄(aq) + 2H₂O(l). What is the concentration of the sodium hydroxide?

Question 53

A student dissolves 1.392 g of an unknown solid diprotic acid, H₂A, in water and makes the solution up to 250.0 mL. 25.00 mL aliquots are titrated with 0.1000 mol L⁻¹ sodium hydroxide solution, which reacts with both protons of each H₂A molecule; the mean titre is 23.58 mL. What is the molar mass of the acid?

Question 54

Acidity constants of five weak acids at 25 °C
Acid	Ka at 25 °C
hydrogensulfate ion, HSO₄⁻	0.010
carbonic acid, H₂CO₃	4.3 × 10⁻⁷
hypochlorous acid, HOCl	3.0 × 10⁻⁸
ammonium ion, NH₄⁺	5.6 × 10⁻¹⁰
hydrogencarbonate ion, HCO₃⁻	4.8 × 10⁻¹¹

The table gives Ka values at 25 °C. A student makes 0.10 mol L⁻¹ solutions of sodium hydrogensulfate (NaHSO₄), ammonium chloride (NH₄Cl), sodium hypochlorite (NaOCl) and sodium carbonate (Na₂CO₃). Which lists the solutions in order of increasing pH?

Question 55

Four mixtures are made at 25 °C. Assume both protons of sulfuric acid are fully donated. Once any reaction on mixing is complete, which mixture is a buffer solution?

Question 56

A student has 50.0 mL of hydrochloric acid with a pH of 1.00 at 25 °C. What volume of 0.0500 mol L⁻¹ barium hydroxide solution, Ba(OH)₂, must she add so that the pH of the mixture is 2.00?

Question 57

A 0.9000 g sample of a garden fertiliser is warmed with excess sodium hydroxide, and all the ammonia released is bubbled into 50.00 mL of 0.2000 mol L⁻¹ hydrochloric acid (an excess). This solution is made up to 250.0 mL, and 25.00 mL aliquots need a mean titre of 16.72 mL of 0.02500 mol L⁻¹ sodium hydroxide, with methyl orange (colour change range pH 3.1 to 4.4) so that NH₄⁺ does not react. Reactions: (NH₄)₂SO₄(s) + 2NaOH(aq) → Na₂SO₄(aq) + 2NH₃(g) + 2H₂O(l), NH₃(g) + HCl(aq) → NH₄Cl(aq) and HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). M((NH₄)₂SO₄) = 132.14 g mol⁻¹. Assuming all the ammonia comes from ammonium sulfate, what is the percentage of ammonium sulfate in the fertiliser by mass?

Question 58

To find the percentage of calcium carbonate in a limestone sample, a student reacts the sample with a measured excess of standard hydrochloric acid, then titrates the leftover acid with standard sodium hydroxide. She rinses the burette with distilled water but not with the sodium hydroxide before filling it. How does this affect her calculated percentage of calcium carbonate?

Short and extended answer questions

Ethanoic acid in a white wine vinegarExtended answer · 8 marks

A student at a Bunbury school pipettes 25.00 mL of white wine vinegar into a 250.0 mL volumetric flask and makes it up to the mark with distilled water. She titrates 20.00 mL aliquots of the diluted vinegar with 0.1050 mol L⁻¹ sodium hydroxide solution. Her titres are in the table; concordant titres agree within 0.10 mL. The reaction is CH₃COOH(aq) + NaOH(aq) → CH₃COONa(aq) + H₂O(l). Molar mass of CH₃COOH = 60.05 g mol⁻¹.

Titres of 0.1050 mol L⁻¹ sodium hydroxide for 20.00 mL aliquots of diluted vinegar
Titration	Titre (mL)
Rough	15.20
1	14.60
2	14.95
3	14.65
4	14.57

(a) Calculate the mean of the concordant titres. [1]

(b) Calculate the concentration of ethanoic acid in the original, undiluted vinegar. [3]

mol L⁻¹

(d) The student used phenolphthalein (8.3 to 10.0) rather than methyl orange (3.1 to 4.4). Explain why phenolphthalein is the better choice for this titration. [2]

0 words · a sentence or two

Calcium carbonate in an antacid tabletExtended answer · 7 marks

A student crushes an antacid tablet (1.250 g) and adds 50.00 mL of 0.5000 mol L⁻¹ hydrochloric acid, an excess. When the fizzing stops, she transfers the mixture to a 250.0 mL volumetric flask and makes it up with distilled water. 25.00 mL aliquots of this solution need a mean titre of 15.10 mL of 0.1000 mol L⁻¹ sodium hydroxide. The reactions are CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) and HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l). M(CaCO₃) = 100.09 g mol⁻¹. Assume that only the calcium carbonate in the tablet reacts with the acid.

(a) Calculate the amount of hydrochloric acid left over after the reaction with the tablet, in the whole 250.0 mL flask. [2]

mol

(b) Calculate the percentage by mass of calcium carbonate in the tablet. [3]

(c) Explain why a back titration is used for this analysis instead of titrating the calcium carbonate directly with hydrochloric acid. [2]

0 words · a sentence or two

Neutralising acid wasteExtended answer · 7 marks

A technician in a Pilbara mine-site laboratory treats acid waste. She mixes 150.0 mL of 0.200 mol L⁻¹ nitric acid with 100.0 mL of 0.250 mol L⁻¹ potassium hydroxide solution at 25 °C. Kw = 1.00 × 10⁻¹⁴.

(a) Calculate the pH of the mixture. [3]

(b) Calculate the volume of 0.100 mol L⁻¹ barium hydroxide solution, Ba(OH)₂, that would exactly neutralise the acid left in the mixture. [2]

(c) A 0.0200 mol L⁻¹ solution of ethanoic acid has the same concentration of acid as the mixture in (a). Explain why its pH is higher. [2]

0 words · a sentence or two

Why salt solutions differ in pHShort answer · 6 marks

Students at a Geraldton school use a pH meter to test 0.10 mol L⁻¹ solutions of three salts at 25 °C. Their results are in the table.

pH of 0.10 mol L⁻¹ salt solutions at 25 °C
Salt solution	pH
sodium carbonate, Na₂CO₃	11.6
ammonium chloride, NH₄Cl	5.1
potassium nitrate, KNO₃	7.0

(a) Write an equation that shows why the sodium carbonate solution is basic. [2]

0 words · a sentence or two

(b) Write an equation that shows why the ammonium chloride solution is acidic. [2]

0 words · a sentence or two

How an ethanoate buffer worksShort answer · 5 marks

A student makes a buffer by dissolving 0.050 mol of sodium ethanoate in 500.0 mL of 0.10 mol L⁻¹ ethanoic acid. The buffer contains the equilibrium CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq).

(a) Explain, using Le Châtelier's principle, how this buffer keeps its pH nearly constant when a few drops of hydrochloric acid are added. [3]

0 words · about 50 to 80 words

(b) The student then adds 0.10 mol of sodium hydroxide to the buffer. Explain why the pH now rises sharply. [2]

0 words · a sentence or two

Standardising hydrochloric acidExtended answer · 9 marks

A student at a Kalgoorlie school dissolves 1.325 g of anhydrous sodium carbonate (M = 105.99 g mol⁻¹) in distilled water and makes the solution up to 250.0 mL in a volumetric flask. She titrates 20.00 mL aliquots of it with a hydrochloric acid solution, using methyl orange; the mean concordant titre is 22.85 mL. The reaction is Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g).

(a) Calculate the concentration of the sodium carbonate solution. [2]

mol L⁻¹

(b) Calculate the concentration of the hydrochloric acid. [2]

mol L⁻¹

0 words · about 50 to 80 words

(d) Before filling the burette with the acid, the student rinsed it with distilled water only. Explain how this affects the value she calculates for the concentration of the hydrochloric acid in its stock bottle. [2]

0 words · a sentence or two

Phosphoric acid in a soft drinkShort answer · 6 marks

Phosphoric acid, H₃PO₄, is added to some cola soft drinks. It is a triprotic weak acid. The table gives the Ka value for each of its three ionisation steps at 25 °C.

Acidity constants for the three ionisation steps of phosphoric acid at 25 °C
Ionisation step	Acid that donates the proton	Ka at 25 °C
first	H₃PO₄	0.0069
second	H₂PO₄⁻	6.2 × 10⁻⁸
third	HPO₄²⁻	4.8 × 10⁻¹³

(a) Write equations for the first and second ionisations of phosphoric acid in water. [2]

0 words · a sentence or two

(b) Use the Ka values to explain why the second ionisation adds very little H₃O⁺ to the solution. [2]

0 words · a sentence or two

(c) A 0.10 mol L⁻¹ solution of phosphoric acid has a higher pH than 0.10 mol L⁻¹ hydrochloric acid, even though phosphoric acid can donate three protons. Explain why. [2]

0 words · a sentence or two