Acid-base language and solution thinking

Brønsted–Lowry definition — An acid donates a proton (H$^+$); a base accepts one.

pH scale

$$\text{pH} = -\log_{10}[\text{H}_3\text{O}^+]$$

• pH $< 7$: acidic. pH $= 7$: neutral. pH $> 7$: basic.
• Autoprotolysis of water: $K_w = [\text{H}_3\text{O}^+][\text{OH}^-] = 10^{-14}$ at $25 ^{\circ}C$.

Weak acid equilibrium

$$\text{HA} + \text{H}_2\text{O} \rightleftharpoons \text{A}^- + \text{H}_3\text{O}^+ \qquad K_a = \frac{[\text{A}^-][\text{H}_3\text{O}^+]}{[\text{HA}]}$$

Henderson–Hasselbalch equation for buffer solutions:

$$\text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}$$

Titration — Adding a strong base to a weak acid traces a characteristic S-shaped curve:

• Half-equivalence point: $\text{pH} = \text{p}K_a$ (buffer capacity is maximum).
• Equivalence point: Moles of acid = moles of base; pH $> 7$ for a weak acid.
• Indicator choice: Select a dye whose colour change occurs near the equivalence pH.

Tip: The Henderson–Hasselbalch equation only works when the buffer ratio $[\text{A}^-]/[\text{HA}]$ is between about 0.1 and 10.

Common pitfall: pH is logarithmic: each unit is a factor of 10 in $H^{+}$. pH 3 is not "a bit" more acidic than pH 5 — it is one hundred times more.