Acid-Base Equilibria (DONE)

Question 1

What does the Lowry-Bronsted theory state

Answer 1

• Acids are defined as a substance which reacts by donating protons (H⁺) and hence bases are proton acceptors.

Question 2

Show the conjugate acid-base pair in the equation:
HIO₃ + H₂O <-> H₃O⁺ + IO₃⁻

Answer 2

HIO₃ is the acid, IO₃⁻ is its conjugate base
H₂O is the base, H₃O⁺ is its conjugate acid

Question 3

Describe what is meant by the terms strong and weak acid.

Answer 3

Strong acid is fully dissociated in water/almost completely
Weak acid is only partially dissociated in water

Question 4

How is the acid dissociation constant Ka calculated?
- What does a higher value of Ka suggest?

Answer 4

[Products]/[Reactants] in mol dm³
- The higher the Ka the stronger the acid

Question 5

How is pKa calculated?
- What does a higher value of pKa suggest?

Answer 5

pKa = -log₁₀Ka
- The higher the pKa the weaker the acid

Question 6

How is pH for STRONG acids calculated?

Answer 6

pH = -log₁₀[H⁺]

Question 7

Give detail of how pH of weak acids is calculated. Use weak acid HA
(3 marks)

Answer 7

• Depends on both its concentration of H⁺ and degree of dissociation (Ka)
• Since Ka - [H⁺][A⁻]/[HA] and [H⁺] = [A⁻]
  [H⁺] = √(Ka x [HA])
• Then pH = -log₁₀[H⁺]

Question 8

How can Kw be expressed?

Answer 8

Kw = [H⁺][OH⁻]

Question 9

How is the pH of a strong base calculated. Use MOH as the strong base

Answer 9

• Use Kw = [H⁺][OH⁻]
• [H⁺] = 1.0x10⁻¹⁴/[OH⁻]

Question 10

What is a buffer solution? What is its important in living things?

Answer 10

pH stays relatively constant as a small amount of acid or alkali is added
- pH of blood regulation, enzymes

Question 11

Describe what an acid buffer contains and how it acts as a buffer. Use sodium ethanoate as an example
- Give equations
(5 marks)

Answer 11

• Consist of a weak acid and a salt of the weak acid
• CH₃COONa -> CH₃COO⁻ + Na⁺
• CH₃COOH <-> CH₃COO⁻ + H⁺
• Adding an acid increases the amount of H⁺, equilibrium shifts left, removing H⁺ by reacting with CH₃COO⁻
• Adding an alkali increases the amount of OH⁻ and this removes some of the H⁺. Causes equilibrium to shift to right, producing H⁺ from CH₃COOH

Question 12

Describe what basic buffer solutions are.
Give an example of a basic buffer, showing equations and explaining the effects of adding OH ions and H ions
(5 marks)

Answer 12

• Consist of a weak base and a salt of the same base
• NH₄Cl -> NH₄⁺ + Cl⁻
• NH₃ + H₂O -> NH₄⁺ + OH⁻
• Adding a base will increase OH ions, shift eql left, removing OH ions by reacting with NH₄⁺
• Adding an acid increases H ions. This reacts with OH ions to form water. Shifts eql to right as OH ions are removed. More OH ions are produced

Question 13

Show equations of the salt hydrolysis for
- Salts of strong acid/weak base
- Salts of weak acid/strong base
Giving brief estimates of the pH

Answer 13

• NH₄⁺ + H₂O <-> NH₃ + H₃O⁺
  Hydrogen ions formed so pH < 7
• CH₃COO⁻ + H₂O <-> CH₃COOH + OH⁻
  Hydroxide ions formed so pH > 7

Question 14

What are the following terms:
- End point
- Equivalence point

Answer 14

End point when the indicator changes colour on addition of one drop of acid/base
Equivalence point is when there are equal amounts of H and OH ions. Neutralisation

Question 15

What are indicators? What is the most suitable indicator related to?

Answer 15

Very weak acids/bases that change colour as pH changes. They change colour of a range of pH values.
Suitable when the pH range lies within the vertical part of the titration curve