Lecture 2 Water: The Medium of Life

1. Lecture 2Water: 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 Explain the difference in these properties.

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

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

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

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

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 F = e1e2/Dr2 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. Iceberg structure Fig. 2-5, p. 32

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

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

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

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. Osmotic Pressure 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.) Fig. 2-8, p. 34

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

21. Formation of Hydronium Ions, H3O+ p. 34

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

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

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

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

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

35. Phosphoric Acid E quilibria • H3PO4 = H+ + H2PO4- (pK = 2.15) • H2PO4- = H+ + HPO42- (pK = 7.20) • HPO42- = H+ + PO43- (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. • C02 + H2O = CO2(d)(H2O) = H2CO3 • H2CO3 = H+ + HCO3- • Henderson-Hasselbalch Equation • pH = pK + log ([HCO3-]/ [H2CO3]

40. Bicarbonate Buffer System • Normal values pH = 7.4 [HCO3-] = 24 mM [H2CO3] = 1.2 mM [HCO3-]/ [H2CO3] = 20/1 pCO2 = 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