Reactivity 3.1.1 – Brønsted–Lowry Theory: Acids and Bases

• This theory was proposed by Johannes Brønsted and Thomas Lowry independently in 1923.
• It defines acid–base behavior in terms of proton (H⁺) transfer, rather than solution pH or hydroxide ions.

Definitions:

• Brønsted–Lowry Acid: A substance that donates a proton (H⁺) to another species.
• Brønsted–Lowry Base: A substance that accepts a proton (H⁺) from another species.

Representation of a Proton:

• In aqueous solution, a free proton (H⁺) does not exist alone — it forms a hydronium ion, \( \text{H}_3\text{O}^+ \), by bonding with a water molecule.
• Therefore, in IB Chemistry, both \( \text{H}^+ \) and \( \text{H}_3\text{O}^+ \) are acceptable to represent a proton in acid–base reactions. The hydronium ion is more accurate in describing aqueous proton behavior.

“Base” vs. “Alkali”

• Base: A broader term — any species that accepts a proton (Brønsted–Lowry base) or donates an OH⁻ (Arrhenius base).
• Alkali: A soluble base — typically a metal hydroxide that dissolves in water to release OH⁻ ions.
• All alkalis are bases, but not all bases are alkalis.

How to Deduce the Brønsted–Lowry Acid and Base in a Reaction

• To apply the Brønsted–Lowry theory to a reaction, look for the species that loses a proton (H⁺) and the one that gains it.
• The species that donates a proton is the acid.
• The species that accepts the proton is the base.
• After proton transfer, the acid becomes its conjugate base, and the base becomes its conjugate acid.

Key Steps to Identify Acid–Base Pairs:

1. Write the balanced chemical equation (can be reversible or one-way).
2. Identify the species that loses an H⁺ → this is the acid.
3. Identify the species that gains an H⁺ → this is the base.
4. Mark the conjugate pairs.

Equilibrium of a Brønsted–Lowry Acid and Base

Brønsted–Lowry acid–base reactions involve the transfer of a proton (H⁺) from an acid to a base. These reactions are typically reversible and establish a dynamic equilibrium in aqueous solution. At equilibrium, the rate of the forward reaction (acid donating H⁺) equals the rate of the reverse reaction (conjugate base accepting H⁺).

The reaction can be represented as:

\( \text{HA}(aq) + \text{B}(aq) \rightleftharpoons \text{A}^-(aq) + \text{BH}^+(aq) \)

In this equilibrium:

• HA is the Brønsted–Lowry acid (proton donor)
• B is the Brønsted–Lowry base (proton acceptor)
• A⁻ is the conjugate base
• BH⁺ is the conjugate acid

The position of equilibrium depends on the relative strength of acids and bases involved. If the acid and base are strong, the forward reaction is favored. If they are weak, the reverse reaction is more significant. In general, equilibrium lies toward the formation of the weaker acid–base pair.

Key Points:

• Proton donation and acceptance always occur together — acids and bases function in pairs.
• Thus, any acid–base reaction involves a proton transfer from the acid to the base.
• Acid → donates H⁺; Base → accepts H⁺.
• Conjugate pairs differ by one H⁺ ion.
• Use chemical equations to deduce acid and base roles based on the species gaining or losing a proton.