Brønsted-Lowry Model Notes

Describe acids and bases in equilibrium systems using the Brønsted-Lowry model.

Brønsted–Lowry acid: A species that donates a proton (H⁺) in an equilibrium reaction.

Brønsted–Lowry base: A species that accepts a proton (H⁺) in an equilibrium reaction.

Explain the Brønsted-Lowry model using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs.

Acid donates a proton: HA + H2O⇌A− + H3O+

Base accepts a proton: NH3 + H2O⇌NH4+ + OH−

Identify that amphiprotic species can act as Brønsted-Lowry acid (or base).

Amphiprotic species: A species that can donate a proton (act as an acid) or accept a proton (act as a base).

Determine the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid).

Conjugate Acid Formula: Base + H+→Conjugate acid

Conjugate Base Formula: Acid − H+→Conjugate base

Identify that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added. (Buffer calculations are not required.)

Buffer solution: A solution containing a conjugate acid–base pair that resists pH change when small amounts of acid or base are added.

Apply Le Châtelier’s principle to explain how buffer solutions respond to the addition of hydrogen ions and hydroxide ions.

Buffer response (Le Châtelier’s principle) Added H⁺: The conjugate base in the buffer reacts with the extra H⁺ to form the conjugate acid. The equilibrium shifts left, removing most of the added H⁺ and minimising the pH change.

Buffer response (Le Châtelier’s principle) Added OH⁻: The conjugate acid donates a proton to neutralise OH⁻, forming water and the conjugate base. The equilibrium shifts right, removing most of the added OH⁻ and resisting pH increase.