Acid-Base Equilibria — Flashcards & Quiz

Q1: Define acids and bases according to the Brønsted-Lowry theory.

A Brønsted-Lowry acid is a proton (H⁺) donor. A Brønsted-Lowry base is a proton (H⁺) acceptor. In any acid-base reaction, there are two conjugate acid-base pairs. Water is amphoteric — it can act as both an acid and a base depending on the reaction partner.

Q2: Distinguish between strong and weak acids in terms of equilibrium.

Strong acids fully dissociate in water — equilibrium lies completely to the right (Ka is very large). Examples: HCl, H₂SO₄, HNO₃. Weak acids partially dissociate — equilibrium lies to the left (Ka is small). A smaller Ka means a weaker acid. Weak acids have both undissociated acid molecules and ions present at equilibrium.

Q3: Define Ka and explain what it indicates.

Ka is the acid dissociation constant: Ka = [H₃O⁺][A⁻] / [HA]. It measures the strength of a weak acid. Larger Ka = stronger acid (more dissociation). Ka is constant at a given temperature. For strong acids, Ka is so large it is not usually quoted. Water is excluded from the expression.

Q1: A Brønsted-Lowry acid is defined as a proton acceptor.

Answer: FALSE

A Brønsted-Lowry acid is a proton (H⁺) DONOR. A proton acceptor is a Brønsted-Lowry BASE.

Q2: Strong acids fully dissociate in aqueous solution.

Answer: TRUE

Strong acids like HCl, H₂SO₄ and HNO₃ completely dissociate into ions in water. There are essentially no undissociated acid molecules remaining.

Q3: Water is amphoteric, meaning it can act as both an acid and a base.

Answer: TRUE

Water can donate a proton (acting as an acid) or accept a proton (acting as a base) depending on the reaction partner.