1. Lecture 2 Water: The Medium of Life Mintel Office Hours Today 1:45 –3:00 Noyes 208

2. Fred Diehl – Univ. of Virginia Fred

3. Physical Properties of Compounds CompoundMP ( o C)BP ( o C)  H vap (cal/g)  H fus (cal/g) H2OH2O 010054080 H2SH2S -836013216.7 NH 3 -78-3332784 Explain the difference in these properties.

4. Fig. 2-1, p. 29 A Water Molecule

5. Structure of Ice – Hydrogen Bonding Each molecule of water can be hydrogen- bonded to up to four other water molecules

6. Fig. 2-2, p. 29 Structure of Ice

7. Liquid Water Lacks the lattice-like structure of ice. H-bonds are not colinear with a line joining the centers of the atoms involved. Therefore the H-bonds are weaker and water is fluid. H-bonds are dynamically formed and broken.

8. Fig. 2-3a, p. 30 Dynamic Formation of H-bonds in liquid water. Note the time scale.

9. Solvent Properties of Water Example – Solubility of sodium chloride Sodium and chloride ions are hydrated. Water molecules are oriented in an opposite direction about sodium and chloride ions, because the interaction is electrostatic.

10. Fig. 2-4, p. 31

11. Water’s Dielectric Constant SolventD Water78.5 Methanol32.6 F = e 1 e 2 /Dr 2 where F is the attractive force between oppositely charged ions e = charge on an ion r = distance between the ions D = dielectric constant

12. Table 2-1, p. 31

13. Hydrophobic Interactions Clathrate (“iceberg”) structure forms surrounding hydrocarbon tails in an aqueous environment, as shown on the next slide.

14. Fig. 2-5, p. 32 Iceberg structure

15. Fig. 2-6, p. 32 More iceberg cages Note disruption of cages when hydrocarbons come together.

16. Fig. 2-7a, p. 33 An Amphipathic Molecule

17. Fig. 2-7b, p. 33 Micelle Formation

18. Soaps and Detergents Grease is dissolved in the hydrocarbon tails of a soap or a detergent. Then, when our coated hands are placed into water, micelles form and disperse down the drain with the grease trapped inside.

19. Fig. 2-8, p. 34 P = iRTm, where i = number of ions, R=gas constant, T=absolute temperature, m = molality (Chemical purists: I know the units don’t work out. See me for an explanation.) Osmotic Pressure

20. Fig. 2-9, p. 34 Ionization of Water

21. p. 34 Formation of Hydronium Ions, H 3 O +

22. Fig. 2-10, p. 35 Hydration of a Hydronium Ion Itself

23. Ion Product of Water K W = 10 -14 = [H + ] [OH - ] In precisely neutral water, [H + ] = [OH - ] = 10 -7 M

24. Definition of pH pH = log [1/H + ] = -log [H + ] A logarithmic scale is more convenient for representing the large range of H + concentrations encountered in biochemistry, just as the Richter scale is more useful for representing the large range of energy values for earthquakes. By extension, pK = log (1/K) - log K

25. Table 2-2, p. 36

26. Table 2-3, p. 36

27. Dissociation of Strong Acids Example: HCl Completely dissociated in solution

28. Dissociation of Weak Acids Example: Acetic Acid Incompletely dissociated in solution

29. Table 2-4, p. 39

30. Fig. 2-11, p. 39 Titration of Acetic Acid – A Closed System (Matter is not ex- changed with the environment. Linear scale Logarithmic scale

31. Fig. 2-11a, p. 39

32. Henderson-Hasselbalch Equation pH = pK + log ([A - ]/[HA]) Derivation of the equation Describes the shape of a titration curve in the neighborhood of the pK. In a molecule with several ionizable groups, there is one H-H equation for each group that titrates. The pK of a weak acid is that pH where HA is half-titrated.

33. Fig. 2-11b, p. 39 Titration Curve for HAc

34. Fig. 2-12, p. 40 The Titration of Some Important Weak Acids

35. Phosphoric Acid Equilibria (1)H 3 PO 4 = H + + H 2 PO 4 - (pK = 2.15) (2)H 2 PO 4 - = H + + HPO 4 2- (pK = 7.20) (3)HPO 4 2- = H + + PO 4 3- (pK = 12.4)

36. Fig. 2-13, p. 41

37. Buffers Definition – A mixture of a weak acid and its conjugate base. Function – Maintains cellular pH, and that of bodily fluids like plasma. Important intracellular buffers are the phosphate system and the histidine system. Buffer capacity is generally best within 1 pH unit of the pK.

38. Fig. 2-14, p. 41

39. Bicarbonate Buffer System Most important buffer system in blood plasma. An open system – Exchanges matter with the environment. C0 2 + H 2 O = CO 2(d) (H 2 O) = H 2 CO 3 H 2 CO 3 = H + + HCO 3 - Henderson-Hasselbalch Equation pH = pK + log ([HCO 3 - ]/ [H 2 CO 3 ]

40. Bicarbonate Buffer System Normal values pH = 7.4 [HCO 3 - ] = 24 mM [H 2 CO 3 ] = 1.2 mM [HCO 3 - ]/ [H 2 CO 3 ] = 20/1 pCO 2 = 40 mmHg

41. Properties of Open System

42. Moral In a closed system a buffer is poorer as one moves away from the pK. In an open system a buffer is better as one moves away from the pK. This is why the bicarbonate buffer system, with a pK = 6.1, is effective at normal plasma pH = 7.4.

43. Respiratory Acidosis Caused, for example, by breathing in and out of a paper bag. The partial pressure of carbon dioxide in the blood increases, and plasma pH accordingly falls.

44. Respiratory Alkalosis Caused, for example, by hyperventilating. More carbon dioxide is blown off by the lungs, and the plasma pH accordingly rises.

45. Problems Do Problems 1 and 2 in Problem Set 1, which is posted on the course web site.

46. Learning Goals Know the physical and chemical properties of water important to biological system. Understand the dissociation of weak electrolytes, and their importance in maintaining pH in cells and tissues. Understand the bicarbonate buffer system and its importance