Worked example 1: Why does retaining extra CO₂ tend to lower blood pH?

More dissolved CO₂ drives CO₂ + H₂O ⇌ H₂CO₃ ⇌ HCO₃^- + H⁺ to the right.
The gas load changes the equilibrium composition, not just one number on a monitor.
The equilibrium shift produces more hydrogen ion.
More hydrogen ion means lower pH on the logarithmic scale.
Ventilation helps because removing CO₂ pulls the same equilibrium back left.
Blood buffering is therefore coupled to breathing, not isolated from it.

Worked example 2: What happens to the net charge of glycine as pH rises from strongly acidic to strongly basic conditions?

At low pH, both the amine and carboxyl group are protonated enough that glycine is net positive.
The acidic environment suppresses deprotonation.
Near intermediate pH, the carboxyl group is deprotonated but the ammonium group remains protonated, giving a zwitterion.
Opposite charges can coexist on the same molecule.
At high pH, the ammonium group also deprotonates, so the net charge becomes negative.
pH changes charge state because each ionizable group has its own acid-base equilibrium.

Worked example 3: Why can ocean acidification reduce shell formation even if the ocean remains basic overall?

Adding dissolved CO₂ raises hydrogen-ion concentration even if the pH stays above 7.
The ocean does not need to become acidic in the everyday sense for chemistry to shift meaningfully.
H⁺ converts some CO₃^2- into HCO₃^-.
That lowers the carbonate available to build calcium carbonate shells.
So a modest pH shift can still change material balance in a biologically important way.
On a logarithmic scale and inside equilibrium networks, small-looking pH changes can matter a great deal.

Common mistake: assuming an environment must cross pH 7 before acid-base chemistry becomes consequential. The relevant question is how proton activity shifts equilibria and charge states around the current operating point.

Acids & bases