Water: Biomedical Importance PDF

Water: Biomedical Importance Water is the predominant chemical component of living organisms. Its unique physical properties, which include the ability to solvate a wide range of organic and inorganic molecules, derive from water's dipolar structure and exceptional capacity for forming hydrogen bond...

Water: Biomedical Importance Water is the predominant chemical component of living organisms. Its unique physical properties, which include the ability to solvate a wide range of organic and inorganic molecules, derive from water's dipolar structure and exceptional capacity for forming hydrogen bonds. The manner in which water interacts with a solvated biomolecule influences the structure of each. An excellent nucleophile, water is a reactant or product in many metabolic reactions. Water has a slight propensity to dissociate into hydroxide ions and protons. The acidity of aqueous solutions is generally reported using the logarithmic pH scale. Bicarbonate and other buffers normally maintain the pH of extracellular fluid between 7.35 and 7.45. Suspected disturbances of acid–base balance are verified by measuring the pH of arterial blood and the CO2 content of venous blood.  Water, an ideal Biological Solvent A water molecule is an irregular, slightly skewed tetrahedron with oxygen at its center. The two hydrogens and the unshared electrons of the remaining two sp 3-hybridized orbitals occupy the corners of the tetrahedron. The 105-degree angle between the hydrogens differs slightly from the ideal tetrahedral angle, 109.5 degrees. Ammonia is also tetrahedral, with a 107-degree angle between its hydrogens. Water is a dipole, a molecule with electrical charge distributed asymmetrically about its structure. The strongly electronegative oxygen atom pulls electrons away from the hydrogen nuclei, leaving them with a partial positive charge, while its two unshared electron pairs constitute a region of local negative charge. The ability of water to ionize, while slight, is of central importance for life. Since water can act both as an acid and as a base, its ionization may be represented as an intermolecular proton transfer that forms a hydronium ion (H3O+) and a hydroxide ion (OH – ) As its name suggests, the ion product Kw is numerically equal to the product of the molar concentrations of H+ and OH– Importance of pH Only a narrow range of pH change can be sustained by a living organism, any further change in pH can make the living difficult. For example: in the case of acid rain, the pH of water is less than 7. As it flows into a river, it lowers the pH of river water which makes the survival of aquatic life difficult. We know that our stomach contains hydrochloric acid which helps in the digestion of food. When the stomach produces too much of hydrochloric acid during indigestion, we feel a lot of pain and irritation. Hence, we generally use antacids or a mild base which increases the pH of the acidic stomach and thus decreases the pain. Bacteria present in our mouth sometimes lower the pH of our mouth by producing acids through degradation of the food particle. Hence, we are instructed to clean our mouths with toothpaste (which is generally basic) to prevent their decay by maintaining the pH. We experience a lot of pain in case of bee-sting as the bee injects the methanoic acid through its sting. Hence, we are generally advised to apply baking soda or other mild bases on the surface as it helps in maintaining the pH of the surface.  Strong Acids and Strong Bases The stronger the acid, the more readily it donates H+. For example, hydrochloric acid (HCl) is highly acidic and completely dissociates into hydrogen and chloride ions, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids Conversely, strong bases readily donate OH– and/or react with hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly basic and give up OH– rapidly when placed in water; the OH– ions react with H+ in solution, creating new water molecules and lowering the amount of free H+ in the system, thereby raising the overall pH. An example of a weak basic solution is seawater, which has a pH near 8.0, close enough to neutral that well-adapted marine organisms thrive in this alkaline environment.  BUFFERS How can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers usually consist of a weak acid and its conjugate base; this enables them to readily absorb excess H+ or OH–, keeping the system’s pH within a narrow range. Maintaining a constant blood pH is critical to a person’s well-being. The buffer that maintains the pH of human blood involves carbonic acid (H2CO3), bicarbonate ion (HCO3–), and carbon dioxide (CO2). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, excess carbonic acid can be converted into carbon dioxide gas and exhaled through the lungs; this prevents too many free hydrogen ions from building up in the blood and dangerously reducing its pH; likewise, if too much OH– is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to jeopardize survival. Key Points A basic solution will have a pH above 7.0, while an acidic solution will have a pH below 7.0. Buffers are solutions that contain a weak acid and its a conjugate base; as such, they can absorb excess H+ ions or OH–ions, thereby maintaining an overall steady pH in the solution. pH is equal to the negative logarithm of the concentration of H+ ions in solution: pH = – log[H+]. Key Terms alkaline: having a pH greater than 7; basic acidic: having a pH less than 7 buffer: a solution composed of a weak acid and its conjugate base that can be used to stabilize the pH of a solution