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Introduction of Biochemistry

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Introduction of Biochemistry with H2O role in Biochemistry

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Introduction of Biochemistry

  1. 1. Introduction of Biochemistry and importance of H2O By: Asheesh Kumar Pandey
  2. 2. Pancreatic cell section
  3. 3. Universal features of living cells
  4. 4. Phylogeny of the three domains of life Introduce by Carl Woese in 1977 on the basis of 16SrRNA
  5. 5. Classification of Organisms
  6. 6. Several functional groups in a single biomolecule
  7. 7. Complementary fit between a macromolecule and a small molecule
  8. 8. Stereoisomers have different effects in humans
  9. 9. Many drugs are racemic mixtures
  10. 10. Energy Transformations in Living Organisms
  11. 11. The major carrier of chemical energy in all cells is adenosine triphosphate (ATP)
  12. 12. Very polar Oxygen is highly electronegative H-bond donor and acceptor High b.p., m.p., heat of vaporization, surface tension Properties of water
  13. 13. Water dissolves polar compounds solvation shell or hydration shell
  14. 14. Water
  15. 15. Hydrogen bonding in ice
  16. 16. Hydrogen Bonding of Water Crystal lattice of ice One H2O molecule can associate with 4 other H20 molecules •Ice: 4 H-bonds per water molecule •Water: 2.3 H-bonds per water molecule
  17. 17. Biologically important hydrogen bonds
  18. 18. Relative Bond Strengths Bond type kJ/mole H3C-CH3 88 H-H 104 Ionic 40 to 200 H-bond 2 - 20 Hydrophobic interaction 3 -10 van der Waals 0.4 - 4
  19. 19. non-covalent interactions
  20. 20. Directionality of the hydrogen bond
  21. 21. Water as a solvent
  22. 22. Non-polar substances are insoluble in water Many lipids are amphipathic
  23. 23. How detergents work?
  24. 24. Amphipathic compounds in aqueous solution
  25. 25. Dispersion of lipids in water
  26. 26. Release of ordered water is energetically favorable
  27. 27. Ionization of Water
  28. 28. Ionization of water
  29. 29. pH Scale  Devised by Sorenson (1902)  [H+] can range from 1M and 1 X 10-14M  using a log scale simplifies notation  pH = -log [H+] Neutral pH = 7.0
  30. 30. Conjugate acid-base pairs consist of a proton donor and a proton acceptor
  31. 31. Titration curve of acetic acid
  32. 32. [OH-] [H+] Keq = [H2O] Kw = [OH-] [H+] = 10-14 M2 Pure H2O : [H+] = [OH-] = 10-7 M pH = - log [H+] = -log (10-7) = 7 If [H+] < 10-7 M then pH < 7 (acidic) If [H+] < 10-7 M then pH < 7 (basic) Blood: [H+] = 4 x 10-8 M Blood pH = 7.4 H2O OH- + H+ Equilibrium constant = = 1.8 x 10-16 M Ion product of water = H2O
  33. 33. Ionization of Water H20 + H20 H3O+ + OH- Keq= [H+] [OH-] [H2O] H20 H+ + OH- Keq=1.8 X 10-16M [H2O] = 55.5 M [H2O] Keq = [H+] [OH-] (1.8 X 10-16M)(55.5 M ) = [H+] [OH-] 1.0 X 10-14 M2 = [H+] [OH-] = Kw If [H+]=[OH-] then [H+] = 1.0 X 10-7
  34. 34. Acid/conjugate base pairs HA + H2O A- + H3O+ HA A- + H+ HA = acid ( donates H+)(Bronstad Acid) A- = Conjugate base (accepts H+)(Bronstad Base Ka = [H+][A-] [HA] Ka & pKa value describe tendency to loose H+ large Ka = stronger acid small Ka = weaker acid pKa = - log Ka
  35. 35. pKa values determined by titration
  36. 36. Phosphate has three ionizable H+ and three pKas
  37. 37. Buffers Buffers are aqueous systems that resist changes in pH when small amounts of a strong acid or base are added. A buffered system consist of a weak acid and its conjugate base. The most effective buffering occurs at the region of minimum slope on a titration curve (i.e. around the pKa). Buffers are effective at pHs that are within +/-1 pH unit of the pKa
  38. 38. Henderson-Hasselbach Equation 1) Ka = [H+][A-] [HA] 2) [H+] = Ka [HA] [A-] 3) -log[H+] = -log Ka -log [HA] [A-] 4) -log[H+] = -log Ka +log [A-] [HA] 5) pH = pKa +log [A-] [HA] HA = weak acid A- = Conjugate base * H-H equation describes the relationship between pH, pKa and buffer concentration
  39. 39. Relationship between pH and pKa pH = pKa when: The molar concentration of acid and conjugate base are equal [H2PO4-] = [HPO42-] pH = pKa = 6.8 Henderson – Hasselbalch equation
  40. 40. Physiological pH The pH in the human body needs to remain ~7. Enzyme catalysis, protein-protein interactions, receptor binding, and other biological processes are very sensitive to pH. pH balance of the blood is maintained using a CO2 - bicarbonate buffer. CO2 + H2O H2CO3 H+ + HCO3- pKa = 6.1 (acid) (hydrated (bicarbonate CO2) base) There is more than 10-fold more base (HCO3-) than acid (CO2) so pH < pK (pH= 7.4) CO2 is exhaled by the lungs H+ + HCO3- CO2 + H2O Breathing rate controls CO2 CO2 balance is controlled by the lungs, HCO3- by the kidneys
  41. 41. Case where 10% acetate ion 90% acetic acid pH = pKa + log10 [0.1 ] ¯¯¯¯¯¯¯¯¯¯ [0.9] pH = 4.76 + (-0.95) pH = 3.81
  42. 42. pH = pKa + log10 [0.5 ] ¯¯¯¯¯¯¯¯¯¯ [0.5] pH = 4.76 + 0 pH = 4.76 = pKa Case where 50% acetate ion 50% acetic acid
  43. 43. pH = pKa + log10 [0.9 ] ¯¯¯¯¯¯¯¯¯¯ [0.1] pH = 4.76 + 0.95 pH = 5.71 Case where 90% acetate ion 10% acetic acid
  44. 44. pH = pKa + log10 [0.99 ] ¯¯¯¯¯¯¯¯¯¯ [0.01] pH = 4.76 + 2.00 pH = 6.76 pH = pKa + log10 [0.01 ] ¯¯¯¯¯¯¯¯¯ [0.99] pH = 4.76 - 2.00 pH = 2.76 Cases when buffering fails
  45. 45. Weak Acids and Bases Equilibria •Strong acids / bases – disassociate completely •Weak acids / bases – disassociate only partially •Enzyme activity sensitive to pH • weak acid/bases play important role in protein structure/function
  46. 46. LEHNINGER PRINCIPLES OF BIOCHEMISTRY Sixth Edition David L. Nelson and Michael M. Cox © 2013 W. H. Freeman and Company CHAPTER 1 The Foundations of Biochemistry Reference:

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