Acid-base language and solution thinking

Acid-Base Foundations

Chemistry changes completely when you track protons. The Brønsted-Lowry definition cuts through the noise: an acid donates a proton ($H^+$), and a base accepts one.

Water itself participates through autoprotolysis, splitting into ions with a constant product at $25 ^{\circ}C$:

$K_w = [H_3O^+][OH^-] = 10^{-14}$

Acidity is measured on a logarithmic scale. The pH scale formula is:

$\text{pH} = -\log_{10}[H_3O^+]$

Common pitfall: Because pH is logarithmic, each unit is a factor of 10 in $[H^+]$. pH 3 is one hundred times more acidic than pH 5, not just a bit more.

Equilibria and Titrations

For a weak acid equilibrium, molecules partially dissociate in water. The acid dissociation constant is:

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

To calculate buffer pH, use the Henderson-Hasselbalch equation:

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

Tip: This equation only works when the buffer ratio $[A^-]/[HA]$ stays between 0.1 and 10.

Adding strong base to a weak acid traces an S-shaped curve during a titration with two critical milestones:

Indicator choice: Always select a dye whose colour change matches the equivalence pH.