equilibrium which one acts as acids
In general chemistry, acids in an equilibrium are identified by the Brønsted–Lowry definition: an acid donates a proton, \( \mathrm{H^+} \). Reversibility matters because the product-side conjugate acid can donate a proton in the reverse direction.
Brønsted–Lowry roles in a reversible reaction
\[ \text{acid: proton donor} \qquad\qquad \text{base: proton acceptor} \]
\[ \mathrm{HA + B \rightleftharpoons A^- + HB^+} \]
In this equilibrium, \(\mathrm{HA}\) donates \(\mathrm{H^+}\) to \(\mathrm{B}\) in the forward direction, so \(\mathrm{HA}\) acts as an acid and \(\mathrm{B}\) acts as a base. In the reverse direction, \(\mathrm{HB^+}\) donates \(\mathrm{H^+}\) to \(\mathrm{A^-}\), so \(\mathrm{HB^+}\) acts as an acid and \(\mathrm{A^-}\) acts as a base.
Conjugate acid–base pairs
Conjugate pairs differ by exactly one proton. A conjugate acid is produced when a base gains \(\mathrm{H^+}\), and a conjugate base is produced when an acid loses \(\mathrm{H^+}\). The equilibrium \(\mathrm{HA + B \rightleftharpoons A^- + HB^+}\) contains two conjugate pairs:
- \(\mathrm{HA/A^-}\). \(\mathrm{HA}\) is the acid; \(\mathrm{A^-}\) is its conjugate base.
- \(\mathrm{HB^+/B}\). \(\mathrm{HB^+}\) is the conjugate acid; \(\mathrm{B}\) is its conjugate base.
Water as an acid or a base
Many acid–base equilibria in aqueous solution involve water. Water is amphiprotic: it can donate \(\mathrm{H^+}\) or accept \(\mathrm{H^+}\), depending on the reacting partner.
| Equilibrium | Acid in forward direction | Base in forward direction | Acid in reverse direction |
|---|---|---|---|
| \(\mathrm{HA + H_2O \rightleftharpoons H_3O^+ + A^-}\) | \(\mathrm{HA}\) | \(\mathrm{H_2O}\) | \(\mathrm{H_3O^+}\) |
| \(\mathrm{B + H_2O \rightleftharpoons BH^+ + OH^-}\) | \(\mathrm{H_2O}\) | \(\mathrm{B}\) | \(\mathrm{BH^+}\) |
| \(\mathrm{NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-}\) | \(\mathrm{H_2O}\) | \(\mathrm{NH_3}\) | \(\mathrm{NH_4^+}\) |
| \(\mathrm{CH_3COOH + H_2O \rightleftharpoons H_3O^+ + CH_3COO^-}\) | \(\mathrm{CH_3COOH}\) | \(\mathrm{H_2O}\) | \(\mathrm{H_3O^+}\) |
Strength of acids and the equilibrium position
For an acid in water, the acid ionization equilibrium and its equilibrium constant provide a quantitative view of acid strength:
\[ \mathrm{HA + H_2O \rightleftharpoons H_3O^+ + A^-} \qquad\qquad K_a=\frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]} \]
Larger \(K_a\) values correspond to greater product formation and stronger acid behavior in water. The logarithmic acidity measure satisfies \(pK_a=-\log(K_a)\), so smaller \(pK_a\) values correspond to stronger acids.
Visualization of acid roles in forward and reverse directions
Compact decision rule
A proton donor on the reactant side is an acid for the forward reaction, and the conjugate acid on the product side is an acid for the reverse reaction. The equilibrium contains an acid on each side because proton transfer is reversible.