First Dissociation Constant of Carbonic Acid

This is the step that sets the pH of almost every natural water on earth. Dissolved CO₂ hydrates to carbonic acid and then loses a proton to give bicarbonate, with K_a1 = 4.45 × 10⁻⁷ at 25 °C. Rainwater in equilibrium with the atmosphere ends up near pH 5.6 from this reaction alone, and a river or aquifer that has picked up limestone climbs to pH 7–8.5 as the bicarbonate builds — which is exactly the alkalinity that water treatment measures.

The constant quoted here is the apparent one, which folds in the fact that most dissolved CO₂ never actually hydrates: only about 1 molecule in 600 exists as true H₂CO₃, and the true acid is roughly 600 times stronger than the tabulated figure suggests. The same equilibrium runs your bloodstream, where the bicarbonate buffer holds arterial pH at 7.40 and the lungs adjust it in seconds by changing how fast CO₂ is exhaled — a buffer with an open valve on one side, which is why it is far more powerful than its pK_a alone would predict.