weak acid
/ WEEK AS-id /
Imagine a crowd at a doorway where only a few people step out while most stay inside, and the in-and-out keeps balancing. A weak acid is like this: drop it in water and only a small fraction of its molecules actually let go of their hydrogen ion, while the rest stay whole.
Formally, a weak acid is one that only partially dissociates in water, settling into an equilibrium between the intact acid and its ions. Because some acid stays undissociated, the reaction has a meaningful, modest acid dissociation constant Ka — small enough that you must solve the equilibrium to find the pH, rather than reading it off the concentration. Most organic acids, like acetic and citric acid, and many biological acids fall here.
Weak acids matter because they are the workhorses of buffering and the reason titration curves of real-world acids have gentle, informative slopes. The honest caveat is that 'weak' is a comparison, not a flaw: a weak acid can still be plenty acidic at high concentration, and pairing it with its conjugate base is precisely what gives buffers their steadying power.
In a 0.1 mole-per-litre acetic acid solution only about one percent of the molecules dissociate, so the pH is near 2.9 — far less acidic than a strong acid at the same concentration, which would sit at pH 1.
Only a small fraction dissociates, so pH must come from the equilibrium.
The fraction that dissociates rises as the acid is diluted, so a weak acid does not have a single fixed percent dissociation — it depends on concentration. This is why pH does not simply track concentration the way it does for strong acids.