By David Stewart | Updated Aug 30, 2022
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A buffer is a chemical substance that stabilizes the pH of a solution, even when acids or bases are added. In living systems, buffering is crucial for maintaining a stable internal environment—homeostasis. Small molecules such as bicarbonate and phosphate, as well as macromolecules like hemoglobin and proteins, all provide this buffering capacity.
The bicarbonate buffer system regulates blood pH through a dynamic equilibrium between carbonic acid (H₂CO₃) and bicarbonate ions (HCO₃⁻). When blood becomes too acidic, the buffer converts excess hydrogen ions into carbon dioxide, which the lungs then expel. Conversely, during alkalosis, bicarbonate is excreted in the urine, restoring neutral pH.
Within cells, the phosphate buffer—comprising hydrogen phosphate (HPO₄²⁻) and dihydrogen phosphate (H₂PO₄⁻)—provides a stronger buffering capacity than the bicarbonate system. It neutralizes excess hydrogen ions by forming the less reactive phosphate species, and in alkaline conditions it accepts hydroxide ions, thereby returning the intracellular environment to neutrality.
Proteins, built from amino acids linked by peptide bonds, have side chains that can donate or accept protons. At physiological pH, carboxyl groups exist as negatively charged carboxylate ions (COO⁻), while amino groups are protonated as NH₃⁺. In acidic conditions, carboxylates capture protons to form COOH, and in alkaline conditions, NH₃⁺ releases protons, becoming NH₂—both actions dampen pH swings.
Hemoglobin, the oxygen‑carrying pigment in red blood cells, also buffers tissue pH. Its globin subunits can bind protons while the heme iron binds oxygen. During exercise, excess protons are taken up by hemoglobin, which simultaneously releases oxygen, thereby mitigating acid buildup in muscles.
