pH scales from 0 to 14, measuring hydrogen ion concentration
An acid donates protons (H+); a base accepts them. In water, the autoionization reaction creates H+ and OH- in equal amounts: [H+] = [OH-] = 1e-7 M at 25 C, giving pH = 7 (neutral). pH is defined as -log[H+], so pH 7 corresponds to [H+] = 1e-7 M. Lower pH (0-7) is acidic (more H+). Higher pH (7-14) is basic (more OH-, less H+).
Strong acids (HCl, H2SO4, HNO3) dissociate completely in water, so a 1 M solution gives [H+] = 1 M and pH = 0. Strong bases (NaOH, KOH) dissociate completely, so 1 M NaOH gives [OH-] = 1 M, [H+] = 1e-14 M, and pH = 14.
Weak acids establish equilibrium
Weak acids (acetic acid, lactic acid, carbonic acid) only partially dissociate. The equilibrium HA <-> H+ + A- is described by the equilibrium constant Ka. If Ka is small, only a small fraction of HA dissociates, so pH is higher than it would be for a strong acid at the same concentration. The Henderson-Hasselbalch equation pH = pKa + log([A-]/[HA]) describes the pH of a weak acid or buffer.
Buffers are mixtures of a weak acid and its conjugate base (e.g., acetic acid and acetate). They resist pH change because added H+ is consumed by the base form, and added OH- is neutralized by the acid form. Buffers are essential in biology: blood pH is buffered by the carbonic acid / bicarbonate system, keeping it near 7.4. Deviations damage proteins and disrupt enzyme function. The pH scale is logarithmic, so a pH change of 1 unit represents a tenfold change in [H+], a massive difference for the cell.