Chapter 8: Acids and Bases

Question 1

Define acids and bases using the Bronsted-Lowry definition

Answer 1

Acid: Proton (H+) donor
Base: Proton (H+) acceptor

Question 2

Define conjugate acid-base pairs and conjugate acids/bases

Answer 2

Conjugate acid-base pairs: Substances that differ by one proton (H+)

Conjugate bases: Formed when acids donate a proton (H+), allowing it to act as a base in the reverse reaction

Conjugate acids: Formed when bases accept a proton (H+), allowing it to act as an acid in the reverse reaction

Question 3

Define amphiprotic and amphoteric, and explain how these two terms differ

Answer 3

Amphiprotic: Substance that can both accept and donate a proton, and can hence act as both a base and an acid

Amphoteric: Substance that can react with both acids and bases, not necessarily involving proton transfer (all amphiprotic substances are amphoteric, but not all amphoteric substances are amphiprotic)

Question 4

Define a salt

Answer 4

Salt: Neutral substance made up of a metal cation (from the base) and non-metal anion (from the acid)

Question 5

Explain how diprotic acids can form 2 different salts, depending on which reactant is in excess

Answer 5

If base is in excess, the diprotic acid is able to undergo two-step dissociation:
Eg. H₂SO₄ + 2 NaOH –> Na₂SO₄+ 2H₂O

If base is limiting, the diprotic acid can only undergo its first dissociation:
Eg. H₂SO₄+ 1 NaOH –> NaHSO₄+ H₂O

Question 6

Describe the reactions of:
- Acids + Reactive metals
- Acids + carbonates/hydrogencarbonates
- Acids + bases

Answer 6

• Acids + Reactive metal –> Salt + H₂ (g)
  (*Link to Redox Processes: Reactive metals are metals higher than H on the activity series, reducing H+ to H₂)
• Acids + Carbonates –> Salt + H₂O (l) + CO₂ (g)
• Acids + Bases –> Salt + H₂O (l)
  [Neutralisation –> exothermic reaction]

Question 7

Explain how to interpret the pH scale

Answer 7

Acidic: pH<7 at 25°C
Neutral: pH=7 at 25°C
Alkaline: pH>7 at 25°C

Question 8

Define and state the formula for pH, and the inverse formula for [H⁺] given the pH

Answer 8

pH: simplified expression of the concentration of H+ ions in a solution (not in mol dm⁻³)

pH = - log₁₀[H⁺]
[H⁺] = 10⁻ᵖᴴ

Question 9

Understand that pH operates on a log scale, and does not increase linearly

Answer 9

pH 0 = 1mol dm⁻³ [H⁺]
pH 1 = 0.1 mol dm⁻³ [H⁺]
pH 2 = 0.01 mol dm⁻³ [H⁺]

0.2 mol dm⁻³ [H⁺] ≠ pH 2

Question 10

Example question:
0.1 mol dm⁻³ of H⁺ ions in 10cm³ has a pH of 1. If another 10cm³ of water is added, how will the pH change?

Answer 10

[H⁺] in 10cm³ = 0.1 mol dm⁻³
+ 10cm³ of water –> 2x dilution
[H⁺] in 20cm³ = 0.05 mol dm⁻³

pH will increase

Question 11

Write an expression for the dissociation of water

Answer 11

H₂O (l) ⇌ H⁺ (aq) + OH⁻ (aq)

Question 12

State the expression for the ionic product constant (Kᵥᵥ) and use Kᵥᵥ to find [H⁺] or [OH⁻] in a strong acid/base, and use the answer to justify why a substance is acidic/basic

Answer 12

Kᵥᵥ = 1.0 x 10⁻¹⁴

[H⁺] [OH⁻] = 1.0 x 10⁻¹⁴
[H⁺] = 1.0 x 10⁻¹⁴/[OH⁻]
[OH⁻] = 1.0 x 10⁻¹⁴/[H⁺]

Acidic: [H⁺] > [OH⁻]
Neutral: [H⁺] = [OH⁻]
Basic: [H⁺] < [OH⁻]

Question 13

Example question:
[HCI] = 0.10 mol dm⁻³
Find [OH⁻]

Answer 13

Since HCI is a strong acid, assume complete dissociation:
[HCI] = [H⁺] = 0.10 mol dm⁻³

Since [H⁺] [OH⁻] = 1.0 x 10⁻¹⁴,
[OH⁻] = 1.0 x 10⁻¹⁴/0.10
= 1.0 x 10⁻¹³

Hence, HCI is acidic, as [H⁺] > [OH⁻]

Question 14

Define a strong/weak acid/base and list examples of strong/weak acids/bases

Answer 14

Acids: Dissociate in water by donating a proton (H⁺) to water
HA (aq) + H₂O (l) ⇌ A (aq) + H₃O (aq)

Strong acids: Dissociate completely in aqueous solution (equilibrium lies long way to the right, allowing us to assume [HA] = [H⁺]
Eg. HCI, H₂SO₄, HNO₃

Weak acids: Dissociate partially in aqueous solution (equilibrium lies to the left: [HA] ≠ [H⁺] )
Eg. H₂CO₃, CH₃COOH

Bases: Ionise in water by receiving a proton (H⁺) from water
B (aq) + H₂O (l) ⇌ BH⁺ (aq) + OH⁻ (aq)

Strong base: Ionise completely in aqueous solution; [B]=[OH⁻]
Eg. Group 1 metal hydroxides and Ba(OH)₂

Weak base: Ionise partially in aqueous solution [B] ≠ [OH⁻]
Eg. NH₃

Question 15

Explain why the strength of an acid/base does not correlate to proticity

Answer 15

• Strength: Degree of dissociation
• Proticity: Number of H atoms that can be donated

Eg.
- Both HCI (monoprotic) and H₂SO₄ (diprotic) are strong acids
- H₂SO₄ and H₂CO₃ are both diprotic, but H₂SO₄ is strong, and H₂CO₃ is weak

Question 16

Describe the relationship between acid/base strength and conjugate base/acid strength

Answer 16

The stronger the acid/base, the weaker its conjugate base/acid
- Strong acids/bases: Dissociate/ionise completely in aqueous solution –> equilibrium lies long way to the right –> conjugate base/acid unlikely to reform parent acid/base

The weaker the acid/base, the stronger its conjugate base/acid
- Weak acids/bases: Dissociate/ionise partially in aqueous solution –> equilibrium lies to the left –> high tendency for conjugate base/acid to reform parent acid/base

Question 17

Describe experiments to distinguish strong and weak acids/bases

Answer 17

*Only acids/bases of equal concentrations can be compared (concentrated weak acid can have a higher [H⁺] than a diluted strong acid):

• Strong acids are better conductors of electricity than weak acids (dissociates completely in aqueous solution = greater concentration of ions = more charge carriers): Conductivity meter/lightbulb in an electrolytic cell set-up
• Strong acids will have a lower pH than weak acids (dissociate completely = higher [H⁺] = higher pH): measure using pH probe/universal indicator
• Strong acids will react faster and more vigorously with metals to produce H₂: More rapid bubbling will be observed

Question 18

Explain why the strength of an acid does not correlate to its concentration

Answer 18

Strength: Degree of dissociation
Concentration: Number of moles of acid before dissociation in mol dm⁻³

Question 19

Explain why the strength of an acid does not correlate with the quantity of its neutralising substance

Answer 19

Equal concentrations of acids, regardless of their strength, will react with equal number of moles of base

Question 20

Explain why rain is naturally acidic, and state its pH range

Answer 20

H₂O (l) + CO₂ (g) –> H₂CO₃ (aq) [carbonic acid: weak acid]
pH range: 5.6 to 7

Question 21

Define acid deposition and state the acidic pollutants that cause it

Answer 21

Acid deposition: when acidic substances leave the earth’s atmosphere and are deposited onto the earth’s surface (wet and dry deposition)

Acidic pollutants: sulfur and nitrogen oxides (from coal power plants and combustion engines)

S (s) + O₂ (g) –> SO₂ (aq)
2 SO₂ (aq) + O₂ –> SO₃ (aq)
SO₃ (aq) + H₂O (l) –> H₂SO₄ (aq) [sulfuric acid]

N₂ (g) + O₂ (g) –> 2 NO (g)
2 NO (g) + O₂ (g) –> 2 NO₂ (g)
2 NO₂ (g) + H₂O (l) –> HNO₂ + HNO₃ [nitric acid]

Question 22

Understand effects of acid deposition

Answer 22

Vegetation: H⁺ ions in acid rain displace the metal ions in soil, which are crucial for the survival of vegetation (Eg. Mg²⁺ ions are crucial for plants to produce chlorophyll, causing them to be unable to photosynthesise properly if displaced)

Aquatic life: cannot survive below certain pH levels

Marble/limestone structures: Eroded by acid rain (acid-carbonate reaction)
CaCO₃ (s) + H₂SO₄ –> CaSO₄ + H₂O + CO₂

Human health: mucosal membranes are irritated by acids, causing respiratory illnesses. acid rain can also react with heavy metal compounds and release poisonous ions

Question 23

Understand pre- and post-combustion methods of reducing sulfur dioxide emissions

Answer 23

Pre-combustion desulfurisation: Remove sulfur before fuel is burned
- Heat crude oil fractions with hydrogen in the presence of a catalyst, to produce H₂S, which can be removed by bubbling gas through an alkaline or converting it back to S, which can be sold to companies to manufacture sulfuric acid

Post-combustion desulfurisation: Remove sulfur after fuel is burned
- Pass gases through a vessel, where SO₂ can react with calcium oxide/carbonate/hydroxide
- Liming of lakes (using Ca(OH)₂ )