Chapter 8: Acids and Bases Flashcards
Define acids and bases using the Bronsted-Lowry definition
Acid: Proton (H+) donor
Base: Proton (H+) acceptor
Define conjugate acid-base pairs and conjugate acids/bases
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
Define amphiprotic and amphoteric, and explain how these two terms differ
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)
Define a salt
Salt: Neutral substance made up of a metal cation (from the base) and non-metal anion (from the acid)
Explain how diprotic acids can form 2 different salts, depending on which reactant is in excess
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
Describe the reactions of:
- Acids + Reactive metals
- Acids + carbonates/hydrogencarbonates
- Acids + bases
- 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]
Explain how to interpret the pH scale
Acidic: pH<7 at 25°C
Neutral: pH=7 at 25°C
Alkaline: pH>7 at 25°C
Define and state the formula for pH, and the inverse formula for [H⁺] given the pH
pH: simplified expression of the concentration of H+ ions in a solution (not in mol dm⁻³)
pH = - log₁₀[H⁺]
[H⁺] = 10⁻ᵖᴴ
Understand that pH operates on a log scale, and does not increase linearly
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
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?
[H⁺] in 10cm³ = 0.1 mol dm⁻³
+ 10cm³ of water –> 2x dilution
[H⁺] in 20cm³ = 0.05 mol dm⁻³
pH will increase
Write an expression for the dissociation of water
H₂O (l) ⇌ H⁺ (aq) + OH⁻ (aq)
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
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⁻]
Example question:
[HCI] = 0.10 mol dm⁻³
Find [OH⁻]
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⁻]
Define a strong/weak acid/base and list examples of strong/weak acids/bases
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
Explain why the strength of an acid/base does not correlate to proticity
- 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
Describe the relationship between acid/base strength and conjugate base/acid strength
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
Describe experiments to distinguish strong and weak acids/bases
*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
Explain why the strength of an acid does not correlate to its concentration
Strength: Degree of dissociation
Concentration: Number of moles of acid before dissociation in mol dm⁻³
Explain why the strength of an acid does not correlate with the quantity of its neutralising substance
Equal concentrations of acids, regardless of their strength, will react with equal number of moles of base
Explain why rain is naturally acidic, and state its pH range
H₂O (l) + CO₂ (g) –> H₂CO₃ (aq) [carbonic acid: weak acid]
pH range: 5.6 to 7
Define acid deposition and state the acidic pollutants that cause it
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]
Understand effects of acid deposition
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
Understand pre- and post-combustion methods of reducing sulfur dioxide emissions
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 )
