Buffer solution

A solution that resists pH changes when small amounts of acid or base are added, usually through a weak acid and its conjugate base.

Henderson–Hasselbalch equation: An equation relating a buffer’s pH to its acid’s pKa and the ratio of conjugate base to acid. It estimates buffer pH from the relative amounts of the pair.
Acid–base equilibrium: The dynamic balance between proton-transfer reactions and the chemical species they produce. Buffer composition is governed by proton-transfer equilibrium.
Blood pH regulation: Physiological processes that keep blood pH within a narrow range. Bicarbonate and other buffering systems limit blood pH shifts.
Strong acid: An acid that ionizes essentially completely in water. Unlike a weak-acid buffer component, it does not establish a substantial acid–conjugate-base pair by partial ionization.
Acetic acid–sodium acetate buffer: An aqueous buffer made from acetic acid and sodium acetate. This familiar weak-acid system illustrates buffering near acetic acid’s pKa.
Conjugate acid–base pair: Two chemical species related by the gain or loss of a proton. The pair consumes added hydroxide and hydronium, limiting pH change.
Weak acid: An acid that only partially ionizes in water. Partial ionization leaves both acid and conjugate base available in a typical buffer.
Bicarbonate buffer system: The carbon dioxide, carbonic acid, and bicarbonate equilibrium that buffers aqueous systems, including blood. It is the principal extracellular buffer system in humans.
Strong base: A base that dissociates essentially completely in water. A strong base alone changes pH sharply rather than providing a conjugate pair to absorb additions.
Ammonia–ammonium buffer: An aqueous buffer made from ammonia and ammonium ions. This weak-base pair demonstrates buffering on the basic side of neutral pH.

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