Introduction of Biochemistry with H2O role in Biochemistry

1. Introduction of Biochemistry and
importance of H2O
By: Asheesh Kumar Pandey

2. Pancreatic cell section

3. Universal features of living cells

4. Phylogeny of the three domains of life
Introduce by Carl Woese in 1977 on the basis of
16SrRNA

5. Classification of Organisms

10. Several functional groups in a single biomolecule

12. Complementary fit between a
macromolecule and a small
molecule

13. Stereoisomers have different effects in humans

14. Many drugs are racemic mixtures

15. Energy
Transformations in
Living Organisms

16. The major carrier of chemical
energy in all cells is adenosine
triphosphate (ATP)

17. Very polar
Oxygen is highly electronegative
H-bond donor and acceptor
High b.p., m.p., heat of vaporization, surface
tension
Properties of water

19. Water dissolves polar
compounds
solvation shell
or
hydration shell

20. Water

21. Hydrogen
bonding in
ice

22. 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

23. Biologically important hydrogen bonds

24. 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

26. non-covalent interactions

27. Directionality of the hydrogen bond

29. Water as a solvent

30. Non-polar substances are insoluble in water
Many lipids are amphipathic

31. How detergents work?

33. Amphipathic compounds in aqueous solution

34. Dispersion of lipids
in water

35. Release of ordered water
is energetically favorable

37. Ionization of Water

38. Ionization of
water

39. 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

41. Conjugate acid-base pairs consist of a proton donor and a proton acceptor

42. Titration curve of acetic acid

44. [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

45. 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

46. 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

47. pKa values determined by titration

48. Phosphate has three
ionizable H+ and three pKas

49. 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

50. 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

51. 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

52. 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

54. Case where 10% acetate ion 90%
acetic acid
pH = pKa + log10 [0.1 ]
¯¯¯¯¯¯¯¯¯¯
[0.9]
pH = 4.76 + (-0.95)
pH = 3.81

55. pH = pKa + log10 [0.5 ]
¯¯¯¯¯¯¯¯¯¯
[0.5]
pH = 4.76 + 0
pH = 4.76 = pKa
Case where 50% acetate ion 50% acetic
acid

56. pH = pKa + log10 [0.9 ]
¯¯¯¯¯¯¯¯¯¯
[0.1]
pH = 4.76 + 0.95
pH = 5.71
Case where 90% acetate ion 10% acetic
acid

57. 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

58. 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

60. 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: