Rucete ✏ Lehninger Principles of Biochemistry In a Nutshell
2.3 Buffering against pH Changes in Biological Systems
This chapter introduces how biological systems resist changes in pH through buffering, why pH constancy is crucial for life, and explains the chemical principles and physiological roles of biological buffers.
The Importance of pH Stability in Biology
• Almost every biological process is pH-dependent; small pH changes can greatly affect reaction rates.
• Many biomolecules (like enzymes, amino acids, nucleotides) have ionizable groups with characteristic pKa values, influencing their function.
• Cells maintain a specific cytosolic pH (near 7), and extracellular fluids are tightly regulated to support optimal biomolecular activity.
What Is a Buffer?
• Buffers are aqueous mixtures of a weak acid and its conjugate base that resist pH changes when small amounts of acid or base are added.
• Example: Acetic acid and acetate ion form a buffer with a relatively flat titration curve around pH 4.76, where buffering is most effective.
• Buffering occurs because addition of acid or base causes only small changes in the ratio of weak acid to its conjugate base, minimizing pH change.
How Buffering Works Chemically
• Buffering is due to two reversible equilibrium reactions between the proton donor (acid) and proton acceptor (base).
• When acid (H+) is added, the conjugate base soaks it up; when base (OH-) is added, the weak acid donates a proton.
• Each conjugate acid-base pair has a specific pH range where it is most effective as a buffer (about ±1 pH unit around its pKa).
The Henderson-Hasselbalch Equation
• The Henderson-Hasselbalch equation relates pH, pKa, and the ratio of proton donor to acceptor: pH = pKa + log ([A-]/[HA])
• It explains buffer action and allows calculation of pH, pKa, or the acid/base ratio as needed.
• At the midpoint of the buffer region (when [A-] = [HA]), pH equals pKa.
Biological Buffers in Cells and Tissues
• Cells and tissues use weak acids and bases (including proteins and metabolites) to buffer pH changes.
• Histidine, with a pKa of 6.0, is a common amino acid in proteins that buffers near neutral pH.
• Other buffer sources: ATP, metabolic intermediates, organic acids (in plant vacuoles), ammonia (in urine).
The Phosphate Buffer System
• Phosphate buffer acts in the cytoplasm of all cells: H2PO4- (acid) ↔ HPO42- (base) + H+
• It is most effective at pH ≈ 6.86, buffering between pH 5.9 and 7.9.
• Example: Addition of strong base to buffered vs. unbuffered water shows that buffer maintains pH near neutral, while pure water’s pH jumps drastically.
The Bicarbonate Buffer System
• Bicarbonate buffers blood plasma: CO2 (gas) dissolves in water to form H2CO3 (acid), which dissociates into HCO3- (base) and H+.
• The system is complex due to rapid equilibrium between CO2 in lungs, dissolved CO2 in plasma, and bicarbonate.
• Blood pH is tightly maintained at ~7.4; breathing rate adjusts CO2 levels and thus pH.
• Hyperventilation (rapid breathing) leads to excessive CO2 loss, causing alkalosis (pH rises above normal), which can be relieved by re-breathing CO2-rich air.
• At normal blood pH, bicarbonate’s effectiveness relies on the large reserve of dissolved CO2 and rapid equilibration in the lungs.
• The apparent pKa for the combined system is about 6.1; with a normal HCO3- : CO2 ratio of ~20:1, pH is ≈ 7.4.
Acidosis, Alkalosis, and Clinical Relevance
• Enzymes have optimal activity at specific pH values, and even slight deviations can sharply reduce their function.
• Untreated diabetes causes acidosis: lack of insulin forces fat breakdown, producing acidic ketone bodies, lowering blood pH below 7.35.
• Symptoms of acidosis include headache, drowsiness, nausea, and in severe cases, coma or death.
• Other causes: starvation, heavy exercise (lactic acid buildup), kidney failure, and lung diseases.
• Acidosis is treated by correcting the underlying problem (e.g., insulin for diabetes, or IV bicarbonate for severe acidosis).
• Intravenous bicarbonate raises plasma pH by increasing the HCO3- concentration, shifting buffer equilibrium.
In a Nutshell
• Biological systems resist pH changes using buffers, which are mixtures of weak acids and their conjugate bases.
• The Henderson-Hasselbalch equation explains how buffers work and allows calculation of pH under varying conditions.
• Phosphate and bicarbonate systems are the most important buffers in cells and blood, maintaining pH near 7.4 for optimal enzyme activity.
• Disturbances in buffer systems can cause life-threatening acidosis or alkalosis, underlining the critical importance of pH regulation in physiology.