pH in Biology 1. Definition and Formula of pH • pH is the measure of hydrogen ion concentration ([H + ]) in a solution, determining its acidity or basicity. • The mathematical formula for pH is: pH = − log10 [H + ] • Similarly, the pOH of a solution is defined as: pOH = − log10 [OH − ] • The relationship between pH and pOH is given by: pH + pOH = 14 where 14 represents the ion product constant of water (Kw ) at 25°C. • Key equation for pure water: Kw = [H + ][OH − ] = 1 × 10−14 mol2 /L2 In pure water, [H + ] = [OH − ] = 1 × 10−7 M, making the pH: pH = − log10 (1 × 10−7 ) = 7 2. Importance of pH in Biology • Biological systems are highly sensitive to pH changes as they influence molecular structure, chemical reactions, and biological functions. • Enzyme activity: – Enzymes typically have an optimal pH range where their catalytic activity is maximum. – Example: Pepsin (active in the stomach) has an optimal pH of 2, while trypsin (active in the small intestine) works best at 8. • Protein structure: – pH affects ionizable side chains (e.g., −COOH and −N H2 ) in amino acids, altering protein folding and stability. – The isoelectric point (pI) is the pH at which a protein carries no net charge. • Cellular processes: 1 – Cellular respiration and photosynthesis rely on specific pH gradients across membranes to drive ATP production. – Example: The proton gradient in mitochondria (pH ≈ 7.5 in the matrix vs. pH ≈ 7.0 in the intermembrane space). • Homeostasis: – Organisms regulate their internal pH tightly (e.g., human blood pH is maintained between 7.35 and 7.45) to sustain life. 3. Buffer Systems and pH Regulation • Buffers maintain pH stability by neutralizing added acids or bases. • Buffer mechanism: Buffers contain a weak acid (HA) and its conjugate base (A− ). The Henderson-Hasselbalch equation describes the relationship between pH and buffer components: − [A ] pH = pKa + log [HA] where: – Ka : Acid dissociation constant. – pKa = − log10 Ka . • Common biological buffers: – Bicarbonate buffer system: Regulates blood pH. H2 CO3 ↔ H + + HCO3− When blood becomes too acidic, bicarbonate (HCO3− ) neutralizes excess hydrogen ions. – Phosphate buffer system: Active in intracellular fluids. H2 P O4− ↔ H + + HP O42− – Protein buffers: Amino acids in proteins act as buffers by accepting or donating H + . 4. Effects of pH on Biological Systems • Acidic environments: – Stomach acid (pH ≈ 2) kills pathogens and aids digestion. – Acidophiles (e.g., Helicobacter pylori ) thrive in low-pH environments. • Basic environments: 2 – Alkaline soils (pH > 8) influence plant nutrient absorption. – Alkaliphiles thrive in high-pH environments (e.g., Bacillus alcalophilus). • Neutral environments: – Most cellular activities occur at pH 7, which is optimal for many biological macromolecules. 5. Experimental Measurement of pH • Indicators: – Substances like litmus and phenolphthalein change color in response to pH changes. • pH meters: – Electrodes measure H + concentration electronically, providing precise pH readings. • pH strips: – Strips of paper impregnated with indicators provide quick pH estimates. 6. Applications in Biology • Blood pH: – Maintained between 7.35 and 7.45 by the bicarbonate buffer system. Severe deviations can lead to acidosis or alkalosis. • Urine pH: – Ranges from 4.5 to 8.0 depending on diet and physiological state. • Soil pH: – Affects plant growth by influencing nutrient availability. – Example: Iron becomes less available at alkaline pH, leading to iron chlorosis in plants. • Enzyme activity: – Diagnostic tests (e.g., ELISA) rely on enzymes that function optimally at specific pH values. 3