AP Chemistry · Summary

Unit 8 summary

1. pH = $-\log[\text{H}_3\text{O}^+]$; at 25 °C, $K_w = 1.0 \times 10^{-14}$ and pH + pOH = 14.00. Neutral means equal ion concentrations, which is pH 7.00 only at 25 °C.
2. Strong acids and bases ionize completely: count every hydroxide in a formula like $\text{Ba(OH)}_2$.
3. For a weak acid, $[\text{H}_3\text{O}^+] \approx \sqrt{K_a C}$ when under 5 percent ionizes; for a weak base, $[\text{OH}^-] \approx \sqrt{K_b C}$. For a conjugate pair, $K_a K_b = K_w$.
4. A buffer holds a weak acid and its conjugate base; pH = $\text{p}K_a + \log([\text{A}^-]/[\text{HA}])$, and it works best within one unit of the $\text{p}K_a$.
5. When strong acid or base is added to a buffer, react it first in moles, then use the buffer equation.
6. At the equivalence point, moles of titrant match moles of analyte; for a weak acid with a strong base the pH there is above 7.
7. At the half-equivalence point, pH = $\text{p}K_a$; choose an indicator whose color change includes the equivalence-point pH.