The Bronsted-Lowry Theory of acids and bases
An acid is a proton donor.
A base is a proton acceptor.
Note: A proton is the same as a H+ ion.
A monobasic acid is able to donate one proton, e.g. HCl.
A bibasic acid is able to donate two protons, e.g. H2SO4.
A tribasic acid is able to donate three protons, e.g. H3PO4.
Acid-base reactions involve the transfer of protons
When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride molecule gives a proton to a water molecule, forming a hydroxonium ion.
\(H_2O_{(l)} + HCl_{(g)} ⟶ H_3O^+_{(aq)} + Cl^-_{(aq)} \)
This is usually written in its simplified form: \(HCl ⟶ H^+ + Cl^- \)
When an acid reacts with a base, the hydroxonium ion reacts with hydroxide ion. A proton is transferred from a hydroxonium ion (the proton donor) to a hydroxide ion (the proton acceptor) to make water.
\(H_3O^+_{(aq)} + OH^-_{(aq)} ⟶ 2H_2O_{(l)} \)
This is commonly written as \(H^+_{(aq)} + OH^-_{(aq)} ⟶ H_2O_{(l)} \), but it is important to realise that whenever you talk about H+ in the context of acids, you are actually referring to the hydroxonium ion H3O+.
Conjugate pairs
The general equation for any acid dissolving in water is:
\(HA + H_2O ⇌ H_3O^+ + A^- \)
Consider the forwards reaction:
The HA is an acid because it is donating a proton to the water.
The water is a base because it is accepting a proton from the HA.
Consider the reverse reaction:
The H3O+ is an acid because it is donating a proton to the A- ion.
The A- ion is a base because it is accepting a proton from the H3O+.
When the acid, HA, loses a proton, it forms a base, A-. When the base, A-, accepts a proton, it reforms the acid, HA. These are called conjugate pairs.
If HA is considered as the acid, then A- is its conjugate base.
If water is considered as the base, then H3O+ is its conjugate acid.
An amphoteric substance can act as either an acid or a base.
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