Questions: Acid and Base Strength: Ka, Kb, and Ionization

Question 1

The Ka of acetic acid (CH₃COOH) is 1.8 × 10⁻⁵. What is the Kb of its conjugate base, acetate ion (CH₃COO⁻)?

• A. 1.8 × 10⁻⁵
• B. 5.6 × 10⁻¹⁰
• C. 1.8 × 10⁻¹⁴
• D. 1.0 × 10⁻⁷

Answer: (B)

For a conjugate acid-base pair, Ka × Kb = Kw = 1.0 × 10⁻¹⁴. So Kb = Kw / Ka = (1.0 × 10⁻¹⁴) / (1.8 × 10⁻⁵) ≈ 5.6 × 10⁻¹⁰. This relationship shows that the stronger the acid, the weaker its conjugate base — acetic acid is a weak acid, so acetate is a moderately weak base.

Question 2

A solution of a weak acid with a larger Ka value will always have a lower pH than an equal-concentration solution of a weak acid with a smaller Ka.

• T. True
• F. False

Answer: True

For two weak acids at the same initial concentration, a larger Ka means a greater degree of ionization — more H⁺ ions in solution — which produces a lower pH. For example, a 0.1 M solution of an acid with Ka = 10⁻³ will be more acidic than a 0.1 M solution of an acid with Ka = 10⁻⁵. (This would be false if concentrations were different, which is a common source of confusion.)

Question 3

Why does a weak acid with Ka = 10⁻⁵ not fully ionize, even though the ionization reaction is spontaneous?

Model answer: The ionization establishes an equilibrium. As H⁺ and A⁻ accumulate, the reverse reaction (recombination) becomes more favorable, and the system reaches a balance where only a small fraction of HA has dissociated. Ka quantifies where that balance lies.

Ka is an equilibrium constant. A value much less than 1 means the equilibrium strongly favors the undissociated form HA. The reaction does proceed forward, but it also proceeds in reverse — and for weak acids, the reverse reaction is fast enough to keep most of the acid intact at equilibrium. This is fundamentally different from strong acids, where the reverse reaction is negligible.