Strong and weak acids and bases Knowledge-Based Study

Study notes

1. In acid–base chemistry, it is essential to distinguish between the strength of an acid or base and its concentration:

   • Strength: Refers to the extent to which an acid or base dissociates or ionizes into ions in aqueous solution.
   • Concentration: Refers to the amount of solute present in a given volume of solution, typically expressed in molarity (mol/L).

   For example, a concentrated solution of a weak acid contains many acid molecules per liter, yet only a small fraction of them ionize. Conversely, a dilute solution of a strong acid contains few total acid molecules per liter, but virtually every molecule dissociates into ions.

   Strong acids and strong bases are considered strong electrolytes because they dissociate essentially completely in dilute aqueous solutions.

   Strong Acids

   Strong acids transfer their acidic protons completely to water molecules, producing hydronium ions:

   $$\mathrm{HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)}$$

   Common strong acids include:

   • Hydrochloric acid ($\mathrm{HCl}$)
   • Hydrobromic acid ($\mathrm{HBr}$)
   • Hydroiodic acid ($\mathrm{HI}$)
   • Nitric acid ($\mathrm{HNO_3}$)
   • Sulfuric acid ($\mathrm{H_2SO_4}$, first dissociation step)
   • Perchloric acid ($\mathrm{HClO_4}$)

   Strong Bases

   Common strong bases are soluble metal hydroxides that dissociate completely in water to yield hydroxide ions ($\mathrm{OH^-}$):

   • Group 1 hydroxides, such as sodium hydroxide ($\mathrm{NaOH}$) and potassium hydroxide ($\mathrm{KOH}$)
   • Heavier Group 2 hydroxides, such as calcium hydroxide ($\mathrm{Ca(OH)_2}$) and barium hydroxide ($\mathrm{Ba(OH)_2}$)

   Weak acids and weak bases ionize only partially in aqueous solutions, setting up a dynamic chemical equilibrium between the intact neutral molecules and the formed ions.

   Weak Acids

   When a weak monoprotic acid ($\mathrm{HA}$) dissolves in water, only a small percentage of its molecules donate a proton:

   $$\mathrm{HA(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + A^-(aq)}$$

   Because the reverse reaction is significant, an equilibrium is established. Common weak acids include acetic acid ($\mathrm{CH_3COOH}$) and hydrofluoric acid ($\mathrm{HF}$).

   Weak Bases

   A weak base accepts a proton from water molecules in a reversible process:

   $$\mathrm{B(aq) + H_2O(l) \rightleftharpoons BH^+(aq) + OH^-(aq)}$$

   Ammonia ($\mathrm{NH_3}$) is a classic weak base:

   $$\mathrm{NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)}$$

   The extent of ionization of a weak acid or base is quantitatively described by its equilibrium constant.

   Acid Dissociation Constant ($K_a$)

   For the ionization equilibrium of a weak acid $\mathrm{HA}$, the equilibrium constant expression is:

   $$K_a = \frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$$

   The value of $K_a$ indicates the relative strength of the acid: a larger $K_a$ corresponds to a higher equilibrium concentration of hydronium ions and thus a stronger acid.

   Base Dissociation Constant ($K_b$)

   Similarly, for a weak base $\mathrm{B}$, the equilibrium constant expression is:

   $$K_b = \frac{[\mathrm{BH^+}][\mathrm{OH^-}]}{[\mathrm{B}]}$$

   A higher $K_b$ indicates a greater degree of protonation and a stronger weak base.