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Salivary amylase digests starches in your mouth at pH 7, but the exact second you swallow, your stomach's pH 2 acid turns it off permanently. Meanwhile, stomach pepsin thrives in that exact same acid.
Why does a flood of hydrogen ions supercharge one enzyme while instantly neutralizing another?
To answer this, we have to look at the sensitive electrical switches built into every amino acid.

The Chemistry of Side Chain Ionization

pH measures the concentration of free hydrogen ions (H+\text{H}^+) in solution, where lower pH values mean a higher concentration of protons. Several amino acid R-groups (side chains) can gain or lose these protons depending on the surrounding acidity.
Each ionizable group has a specific pKa\text{p}K_a, which is the pH at which exactly half of those groups hold a proton and half have released it. When the surrounding pH<pKa\text{pH} < \text{p}K_a, the excess protons force the group into its protonated state; when pH>pKa\text{pH} > \text{p}K_a, protons leave into solution, leaving the group deprotonated.
In 1909, Danish chemist Søren Sørensen invented the pH scale while studying brewing precisely because he noticed that tiny shifts in hydrogen ion concentration dramatically altered enzyme reaction rates.
So what happens when these shifting electrical charges sit directly inside the enzyme's catalytic pocket?

Active Site Charge Distribution

An enzyme's active site relies on precise electrostatic complementary charges to attract, position, and chemically attack its substrate. If the substrate carries a positive charge, the active site often uses a negatively charged aspartate (−COO−-\text{COO}^-) to form an attractive ionic bond.
If pH drops below the aspartate's pKa\text{p}K_a, that carboxylate group gains a proton and becomes neutral (−COOH-\text{COOH}). Without its negative charge, the active site loses its electrostatic attraction for the substrate, causing binding affinity to plummet.