Acid-Base Chemistry and Equilibrium

pH, Strong and Weak Acids

Water autoionizes with $K_w = [H^+][OH^-] = 10^{-14}$ at 25 degrees Celsius, so $pH + pOH = 14$ and $pH = -\log[H^+]$ The seven strong acids and strong bases dissociate completely; weak acids only partially, with strength set by $K_a$ (larger $K_a$, smaller $pK_a$, stronger acid). For a weak acid, approximate $[H^+] \approx \sqrt{K_aC}$.

Buffers and Henderson-Hasselbalch

A buffer pairs a weak acid with its conjugate base and resists pH change: $pH = pK_a + \log\frac{[A^-]}{[HA]}$ Buffering is best within one unit of $pK_a$. The bicarbonate system ($CO_2/HCO_3^-$) stabilizes blood at pH 7.4, linking ventilation to acid-base status.

Titrations and Equilibrium

Weak acid-strong base titrations show a buffer region, half-equivalence where $pH = pK_a$, and a basic equivalence point. For general equilibria, $Q$ vs. $K$ predicts reaction direction, Le Chatelier's principle predicts shifts from added species, pressure, or temperature, and $K_{sp}$ governs solubility, which drops with a common ion.

Key Takeaways

• Each pH unit is a tenfold change in $[H^+]$.
• At half-equivalence, $pH = pK_a$; at equivalence for a weak acid, pH is above 7.
• Only temperature changes the value of $K$; everything else shifts position only.
• Conjugates: $K_a \times K_b = K_w$.