Comprehensive Biochemistry

COMPItEHEHSI¥.E BIOCHEMISTlty Surya Prakash Dwivedi Neeraja Dwivedi M. Tech. (Biotechnology) M. Tech. (Biotechnology) . Assistant Professor Department of Biotechnology College of Engineering & Technology, I.F.T.M ., Moradabad (U.P.) Assistant Professor Department of Biotechnology College of Engineering & Technology, I.F.T.M., Moradabad (U'p.) P~G~TI \'\"',7. BUFFERS As shown above, the pH of a solution is dependent on the concentration of H+ ions. Addition or removal of H+ ions, then, can greatly affect the pH of a solution. In the body, the pH of cells and extracellular fluids can vary from pH 8 in pancreatic fluid to pH 1 in stomach acids. The average pH of blood is 7.4, and of cells is 7-7.3. Although there is great variation in pH between the fluids in the body, there is little variation within each system. For example, blood pH only varies between 7.35-7.45 in a healthy individual. Large changes in pH can be 16 Comprehensive Biochemistry life threatening. How does the body maintain a constant blood pH? The body uses a buffer system to withstand changes in pH. Buffers are made up of a mixture of a weak acid with its conjugate base or a weak base with its conjugate acid. Remember that an acid donates a H+. A weak acid does not donate its H+ as easily. Similarly, a weak base will not accept a H+ as well as a strong base. Buffers maintain pH by binding H+ or OH- ions. This stabilizes changes in pH. The bicarbonate buffer system maintains blood pH near pH 7.4. The carbonic acid, H2C03, in the blood is in equilibrium with the carbon dioxide (C0 2), in the air. Buffers are most effective in a pH range near its pK•. This is where the titration curve is most shallow, and where the pH is least affected by added acid or base. For example, in the titration curve of phosphoric acid (another blood buffering system): H 3P0 4 pKa 1 , H P0 4 2 + H+ pKa2 > HP04- 2+ H+ pKa3 > P0 4- 3 + H+ 14 12 10 8 6 4 2 0 2 0 3 Equivalents of 011 Added Figure: 2.4: Titration curve showing cell buffering region The region in the rectangle is the buffering region of biological importance. Note that the slope of the curve is shallow within one pH unit of the pK. value. This is where the buffering range is most effective, meaning that the buffer is able to resist changes in pH. NON-COVALENT BONDING Non-covalent bonds are not at strong as covalent bonds, but they are important in the stabilization of molecules. In contrast to covalent bonds, nOh-::ovalent bonds do not share electrons. Noncovalent bonds include: • • • • Electrostatic interactions Van der Waals forces Hydrophobic interactions Hydrogen bonds Water, pH and Buffers 17 Electrostatic interactions are fonned between positive and negative ions. The bond is non-directional, meaning that the pull of the electrons does not favor one atom over another. An example is NaCI, which is fonned between the positively charged Na+ ion and the negatively charged cr ion. The average strength of an electrostatic interaction is 15 kilojoules/mol (kJ/mol). The bond strength lessens when the distance between the two ions increases. Van der Waals forces (aka London forces) are weak forces between temporary dipoles. These forces may be attractive or repulsive. They are also non-directional. The average bond energy for van der Waals forces is on the order of 10 kJ/mol. Hydrophobic interactions result when non-polar molecules are in a polar solvent, e.g., H20. The non-polar molecules group together to exclude water (hydrophobic means water fearing). By doing so, they minimize the surface area in contact with the polar solvent. A hydrogen bond results when a hydrogen