The acid dissociation constant ($K_a$) of acetic acid (ethanoic acid) at 298 K (25 °C) is $1.75 \times 10^{-5} \text{ mol L}^{-1}$.

• Mathematical Formula: The $\text{p}K_a$ is defined as the negative logarithm (base 10) of the dissociation constant:
  $$\text{p}K_a = -\log_{10}(K_a)$$
  $$\text{p}K_a = -\log_{10}(1.75 \times 10^{-5}) = 5 - \log_{10}(1.75) \approx 4.76$$
• Buffer Applications: An equimolar buffer solution of acetic acid and sodium acetate ($ ext{CH}_3 ext{COOH} / ext{CH}_3 ext{COONa}$) maintains a pH equal to its $\text{p}K_a$ (4.76) by the Henderson-Hasselbalch equation.
• Acid Strength: Formic acid ($\text{HCOOH}$) is stronger with a lower $\text{p}K_a$ of 3.75.
• Food Use: Acetic acid (4-8% aqueous solution) is commonly used as vinegar for food preservation and pickling.

At 298 K (25 °C), the autoionization of pure liquid water produces hydronium and hydroxide ions according to the equilibrium:

• Expression & Calculation: The ionic product ($K_w$) is defined as:
  $$K_w = [\text{H}^+][\text{OH}^-]$$
  In pure neutral water at 298 K:
  $$[\text{H}^+] = [\text{OH}^-] = 1.0 \times 10^{-7} \text{ M}$$
  $$K_w = (1.0 \times 10^{-7} \text{ M}) \times (1.0 \times 10^{-7} \text{ M}) = 1.0 \times 10^{-14} \text{ M}^2$$
• Dimension / Unit: The concentration terms are in molarity ($\text{M}$), making the unit $\text{M}^2$ (or $\text{mol}^2 \text{L}^{-2}$).
• Temperature Dependence: Autoionization of water is an endothermic process; therefore, $K_w$ increases with temperature (e.g., at 373 K, $K_w \approx 5.13 \times 10^{-13} \text{ M}^2$).
• Relation to pH: $\text{pH} + \text{pOH} = \text{p}K_w = 14$ at 298 K.