Water and Life Lecture 8 PDF

Water & Life Module 2 Dr. Priyanka Singh IITJ Polar Covalent bonds in water result in H-bonding Four emergent properties of water contribute to Earth’s suitability for life Cohesion of Water Molecules Ability to moderate temperature Floating of ice in water Ve...

Water & Life Module 2 Dr. Priyanka Singh IITJ Polar Covalent bonds in water result in H-bonding Four emergent properties of water contribute to Earth’s suitability for life Cohesion of Water Molecules Ability to moderate temperature Floating of ice in water Versatility as a solvent  Hydrogen bond Cohesion Behavior Water molecules stay close to each other due to hydrogen bonding. These linkages make water more structured than most other liquids. Liquid water Adhesion of water by hydrogen bonds to the molecules of cell walls helps counter the downward pull of gravity. Related to cohesion is surface tension, a measure of how difficult it is to stretch or break the surface of a liquid Moderation of Temperature by water Water’s High Specific Heat The specific heat of a substance is defined as the amount of heat that must be absorbed or lost for 1 g of that substance to change its temperature by 1°C. The specific heat of water is 1 calorie per gram and per degree Celsius A calorie of heat causes a relatively small change in the temperature of water because much of the heat is used to disrupt hydrogen bonds before the water molecules can begin moving faster. And when the temperature of water drops slightly, many additional hydrogen bonds form, releasing a considerable amount of energy in the form of heat. Floating of Water is one of the few substances that are less dense as a solid than as a liquid. In other words, ice ice in water floats on liquid water Water: The solvent of life  The Hydrophobic Effect The hydrophobic effect refers to the entropy-driven aggregation of nonpolar molecules in aqueous solution that occurs to minimize the ordering of water molecules with which they are in contact. This is not an attractive force, but rather a thermodynamically driven process. The hydrophobic effect drives the formation of membranes and contributes to the folding of proteins and the formation of double helical DNA. The position of equilibrium of any chemical reaction is given by its equilibrium constant In pure water at 25°C, the concentration of water is 55.5 M the ion product of water at 25°C. The value for Keq is 1.8 × 10−16 M at 25°C as calculated from electrical conductivity measurements When there are exactly equal concentrations of both H+ and OH−, as in pure water, the solution is said to be at neutral pH pH scale: logarithmic scale for expressing the H+ concentration  Strong and Weak Acids A strong acid is defined as a substance that has a greater tendency to lose its proton and therefore completely dissociates (or ionizes) in water, such as HCl and H2SO4. A weak acid, on the other hand, is a molecule that has a lesser tendency to lose its proton (or, in other words, displays a high affinity for its proton) and, therefore, does not readily dissociate in water, such as CH3COOH. The dissociation of the weak organic compound, acetic acid, is written as : CH3COOH  H+ + CH3COO− At a given temperature, the extent of ionization at equilibrium can be calculated by the following equation : Ka = [H+ ] [CH3 COO−] / [CH3 COOH]  pH & pKa The strength of an acid depends on its degree of dissociation in water and can be determined from the equilibrium expression Hendersson-Hasselbach equation relates a relationship between the concentration of hydrogen ions in solutions and acid strength  Buffer A buffer solution is one which resists changes in pH when small quantities of an acid or an alkali are added to it. Most commonly, the buffer solution consists of a mixture of a weak acid and its conjugate base; for example, mixtures of acetic acid and sodium acetate or of ammonium hydroxide and ammonium chloride are buffer solutions Cells maintain a constant pH of 7.2-7.4 in cytoplasm (physiological pH). This is because it has reservoirs of weak acids and bases, called buffers, that ensures that pH remains constant. Buffers soak up protons or hydroxy ions when added to cells. This ability of a buffer to minimize changes in pH, its buffering capacity, depends on the relationship between its pKa value and the pH  Buffering range Maximum buffering capacity occurs at a pH equal to the pKa, but a conjugate acid/base pair can still serve as an effective buffer when the pH of a solution is within approximately ± 1pH unit of the pKa 1. A buffer is prepared containing 1.00 M acetic acid and 1.00 M sodium acetate. What is its pH? (The Ka of acetic acid is 1.77 x 10¯5) (Ans: pH = 4.752) 2. A buffer is prepared containing 0.800 molar acetic acid and 1.00 molar sodium acetate. What is its pH? (Ans: pH = 4.85) 3. A buffer is prepared containing 1.00 molar acetic acid and 0.800 molar sodium acetate. What is its pH? (Ans: pH = 4.65) 4. (a) Calculate the pH of a 0.500 L buffer solution composed of 0.700 M formic acid (HCOOH, Ka = 1.77 x 10¯4) and 0.500 M sodium formate (HCOONa). (Ans: pH = 3.6) (b) Calculate the pH after adding 50.0 mL of a 1.00 M NaOH solution. (Ans: pH = 3.77)