Campbell6e lecture ch2

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Campbell6e lecture ch2

  1. 1. Chapter Two Water: The Solvent for Biochemical Reactions Paul D. Adams University of Arkansas
  2. 2. What makes water polar <ul><li>What is a polar bond: </li></ul><ul><li>• Electrons are unequally shared,more negative charge found closer to one atom. </li></ul><ul><li>• Due to difference in electronegativity of atoms involved in bond. </li></ul>
  3. 3. Electronegativity <ul><li>Electronegativity: a measure of the force of an atom’s attraction for electrons it shares in a chemical bond with another atom </li></ul><ul><ul><li>Oxygen and Nitrogen, more electronegative than carbon and hydrogen </li></ul></ul><ul><ul><li>Fluorine is most electronegative (4) </li></ul></ul>
  4. 4. Polar Bonds & Molecules <ul><li>Molecules such as CO 2 have polar bonds but, given their geometry, are nonpolar molecules; that is, they have a zero dipole moments </li></ul>
  5. 5. Solvent Properties of H 2 O <ul><li>Ionic compounds (e.g.,KCl) and low-molecular- weight polar covalent compounds (e.g., C 2 H 5 OH and CH 3 COCH 3 ) tend to dissolve in water </li></ul><ul><li>The underlying principle is electrostatic attraction of unlike charges; the positive dipole of water for the negative dipole of another molecule, etc. </li></ul><ul><ul><li>ion-dipole interaction: e.g., KCl dissolved in H 2 O </li></ul></ul><ul><ul><li>dipole-dipole interactions: e.g., ethanol or acetone dissolved in H 2 O </li></ul></ul><ul><ul><li>dipole induced-dipole interactions: weak and generally do not lead to solubility in water </li></ul></ul>
  6. 6. Hydration Shells Surrounding Ions in Water
  7. 7. • Ion-dipole and dipole-dipole interactions help ionic and polar compounds dissolve in water Ion-dipole and Dipole-dipole Interactions
  8. 8. Solvent Properties of H 2 O <ul><li>Hydrophilic : water-loving </li></ul><ul><ul><li>tend to dissolve in water </li></ul></ul><ul><li>Hydrophobic : water-fearing </li></ul><ul><ul><li>tend not to dissolve in water </li></ul></ul><ul><li>Amphipathic : has characteristics of both properties </li></ul><ul><ul><li>molecules that contain one or more hydrophobic and one or more hydrophilic regions, e.g., sodium palmitate </li></ul></ul>
  9. 9. Amphipathic molecules <ul><li>• both polar and nonpolar character </li></ul><ul><li>Interaction between nonpolar molecules is very weak </li></ul><ul><ul><li>called van der Waals interactions </li></ul></ul>
  10. 10. Micelle formation by amphipathic molecules <ul><li>Micelle: a spherical arrangement of organic molecules in water solution clustered so that </li></ul><ul><ul><li>their hydrophobic parts are buried inside the sphere </li></ul></ul><ul><ul><li>their hydrophilic parts are on the surface of the sphere and in contact with the water environment </li></ul></ul><ul><ul><li>formation depends on the attraction between temporary induced dipoles </li></ul></ul>
  11. 11. Examples of Hydrophobic and Hydrophilic Substances
  12. 12. Hydrogen Bonds <ul><li>Hydrogen bond: the attractive interaction between dipoles when: </li></ul><ul><ul><li>positive end of one dipole is a hydrogen atom bonded to an atom of high electronegativity, most commonly O or N, and </li></ul></ul><ul><ul><li>the negative end of the other dipole is an atom with a lone pair of electrons, most commonly O or N </li></ul></ul><ul><li>Hydrogen bond is non-covalent </li></ul>
  13. 13. Interesting and Unique Properties of Water <ul><li>• Each water molecule can be involved in 4 hydrogen bonds: 2 as donor, and 2 as acceptor </li></ul><ul><li>• Due to the tetrahedral arrangement of the water molecule (Refer to Figure 2.1). </li></ul>
  14. 14. Hydrogen Bonding <ul><li>Even though hydrogen bonds are weaker than covalent bonds, they have a significant effect on the physical properties of hydrogen-bonded compounds </li></ul>
  15. 15. Other Biologically Important Hydrogen bonds <ul><li>• Hydrogen bonding is important in stabilization of 3-D structures of biological molecules such as: DNA, RNA, proteins. </li></ul>
  16. 16. Acids, Bases and pH <ul><li>Acid: a molecule that behaves as a proton donor </li></ul><ul><li>Strong base: a molecule that behaves as a proton acceptor </li></ul>
  17. 17. <ul><li>One can derive a numerical value for the strength of an acid (amount of hydrogen ion released when a given amount of acid is dissolved in water). </li></ul><ul><li>Describe by K a : </li></ul><ul><li>Written correctly, </li></ul>Acid Strength
  18. 18. Ionization of H 2 O and pH <ul><li>Lets quantitatively examine the dissociation of water: </li></ul><ul><li>• Molar concentration of water (55M) </li></ul><ul><li>• K w is called the ion product constant for water. </li></ul><ul><li>• Must define a quantity to express hydrogen ion concentrations…pH </li></ul>
  19. 19. Henderson-Hasselbalch <ul><li>Equation to connect K a to pH of solution containing both acid and base. </li></ul><ul><li>We can calculate the ratio of weak acid, HA, to its conjugate base, A - , in the following way </li></ul>
  20. 20. Henderson-Hasselbalch (Cont’d) <ul><li>Henderson-Hasselbalch equation </li></ul><ul><li>From this equation, we see that </li></ul><ul><ul><li>when the concentrations of weak acid and its conjugate base are equal, the pH of the solution equals the pK a of the weak acid </li></ul></ul><ul><ul><li>when pH < pK a , the weak acid predominates </li></ul></ul><ul><ul><li>when pH > pK a , the conjugate base predominates </li></ul></ul>
  21. 21. Titration Curves <ul><li>Titration: an experiment in which measured amounts of acid (or base) are added to measured amounts of base (or acid) </li></ul><ul><li>Equivalence point: the point in an acid-base titration at which enough acid has been added to exactly neutralize the base (or vice versa) </li></ul><ul><ul><li>a monoprotic acid releases one H + per mole </li></ul></ul><ul><ul><li>a diprotic acid releases two H + per mole </li></ul></ul><ul><ul><li>a triprotic acid releases three H + per mole </li></ul></ul>
  22. 22. Buffers <ul><li>buffer: a solution whose pH resists change upon addition of either more acid or more base </li></ul><ul><ul><li>consists of a weak acid and its conjugate base </li></ul></ul><ul><li>Examples of acid-base buffers are solutions containing </li></ul><ul><ul><li>CH 3 COOH and CH 3 COONa </li></ul></ul><ul><ul><li>H 2 CO 3 and NaHCO 3 </li></ul></ul><ul><ul><li>NaH 2 PO 4 and Na 2 HPO 4 </li></ul></ul>
  23. 23. Buffer Range <ul><li>A buffer is effective in a range of about +/- 1 pH unit of the pK a of the weak acid </li></ul>
  24. 24. Buffer Capacity <ul><li>Buffer capacity is related to the concentrations of the weak acid and its conjugate base </li></ul><ul><ul><li>the greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity </li></ul></ul>
  25. 25. Naturally Occurring Buffers <ul><li>H 2 PO 4 - /HPO 4 2- is the principal buffer in cells </li></ul><ul><li>H 2 CO 3 /HCO 3 - is an important (but not the only) buffer in blood </li></ul><ul><ul><li>hyperventilation can result in increased blood pH </li></ul></ul><ul><ul><li>hypoventilation can result in decreased blood pH </li></ul></ul><ul><ul><li>(Biochemical Connections p. 60) </li></ul></ul>
  26. 26. Selecting a Buffer <ul><li>The following criteria are typical </li></ul><ul><ul><li>suitable pK a </li></ul></ul><ul><ul><li>no interference with the reaction or detection of the assay </li></ul></ul><ul><ul><li>suitable ionic strength </li></ul></ul><ul><ul><li>suitable solubility </li></ul></ul><ul><ul><li>its non-biological nature </li></ul></ul>
  27. 27. Laboratory Buffers

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