Brønsted-Lowry Theory

Chemistry: IB Diploma Higher Level, topic 6.1.1

Brønsted-Lowry Acids and Bases

The chemists Brønsted and Lowry defined what an acid-base reaction is.

A Brønsted-Lowry acid is a proton donor. This means that it gives up H⁺ ions in solutions. Examples of Brønsted-Lowry acids are:
Hydrochloric acid.
Sulfuric acid.
Phosphoric acid.

A Brønsted-Lowry base is a proton acceptor. This means that it accepts H⁺ ions in solutions. Examples of Brønsted-Lowry bases are:
Sodium hydroxide.
Potassium hydroxide.
Ammonia.

Brønsted-Lowry acids need something to give protons to - you don’t get free H⁺ ions.
When you add HCl to water, the water molecules act like a base and accept protons from the HCl. The equation is:
HCl + H₂O → H₃O⁺ + Cl⁻
The H₃O⁺ ion is called a ‘hydronium’ ion.
So, water plays an important role in aqueous reactions involving acids.
Water is able to accept protons and then donate them to other species.

The concept of conjugate pairs is useful in describing Brønsted-Lowry acid-base reactions.
When an acid donates H⁺, the species that remains is called the conjugate base of the acid.
This is because it reacts as a proton acceptor in the reverse reaction.
Likewise, when a base accepts H⁺, it is converted to its conjugate acid.

The reaction between water and ammonia illustrates the idea of conjugate pairs.
In the forward direction, water (H₂O) acts as an acid by donating a proton to ammonia.
Water subsequently becomes a hydroxide ion, OH⁻, the conjugate base of water.
The ammonia (NH₃ acts as a base in accepting this proton, becoming an ammonium ion, (NH₄⁺), the conjugate acid of ammonia.
In the reverse direction, a hydroxide ion acts as a base in accepting a proton from ammonium ion, which acts as an acid.

What are the conjugate acids in the below equation?

HSO₃⁻_(aq)+HI _(aq) ⇌ H₂SO_3(aq)+ I⁻_(aq)
When you add HCl to water, the water molecules act like a base and accept protons from the HCl. Finish the equation of this reaction
HCl + H₂O →