1. ERT106 BIOCHEMISTRY WATER
Pn Syazni Zainul Kamal

2. Lecture contents Why water important to biochemistry Uses of water
Physics & chemistry of water
Unique physical properties of water
Molecular structure of water
Noncovalent Bonding in water
Ionic interactions
Hydrogen Bonds
van der Waals Forces
Thermal Properties of Water
Solvent Properties of Water
Hydrophilic, hydrophobic, and amphipathic molecules
Osmotic pressure

3. Ionization of Water
Acids, bases, and pH
Buffers
titration

4. Why water is important to biochemistry
More than 70% earth’s surface covered with water
The substance that make possible life on earth
Solvent & substrate for many cellular reaction
Transports chemicals from place to place
Helps to maintain constant body temperature
Cell components and molecules (protein, polysaccharides, nucleic acid, membranes) assume their shape in response to water

5. USES OF WATER?

6. Introduction Physic and chemistry of water
Water is the chemical substance with chemical formula H2O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom.
Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue.

7. Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom.
The presence of a charge on each of these atoms gives each water molecule a net dipole moment.

8. Introduction Unique physical properties of water
Exist in all three physical states of matter: solid, liquid, and gas.
Has high specific heat
Water conducts more easily than any liquid except mercury
Water has a high surface tension
Water is a universal solvent
Water in a pure state has a neutral pH

9. Molecular Structure of Water
Tetrahedral geometry
The oxygen in water is sp3 hybridized. Hydrogens are bonded to two of the orbitals. Consequently the water molecule is bent. The H-O-H angle is 104.5o.

10. The bent structure indicate water is polar coz linear structure is nonpolar.
Phenomenon where charge is separated to partial –ve charge and partial +ve charge is called dipoles.

11. Water is a polar molecule.
A polar molecule is one in which one end is partially positive and the other partially negative.
Oxygen is more electronegative than hydrogen, so oxygen atom bears a partial –ve charge, hydrogen atoms are partial +ve charge

13. Molecules eg water, in which charge is separated are called dipoles.
Molecular dipoles will orient themselves in the direction opposite to that of the field when subjected to an electric field.

14. Noncovalent Bonding Usually electrostatic
They occur between the positive nucleus of one atom and the negative electron clouds of another nearby atom
Relatively weak, easily disrupted
Large no. of noncovalent interactions stabilize macromolecules

15. Types of noncovalent bonding :
1)Ionic interactions
2)Hydrogen bonding
3)Van der Waals forces
-Dipole-dipole
-Dipole-induced dipole
-Induced dipole-induced dipole

16. Typical “Bond” Strengths
Type
kJ/mol
Covalent
>210
Noncovalent
Ionic interactions
4-80
Hydrogen bonds
12-30
van der Waals
0.3-9
Hydrophobic interactions
3-12

17. 1) Ionic Interactions
Interaction occur between charged atoms or group.
Oppositely charged ions are attracted to each other. (eg. NaCl)
ions with similar charges eg K+ and Na+ will repel each other

19. In proteins, certain amino acid side chains contain ionizable groups.
Glutamic acid ionized as –CH2CH2COO-
Lysine ionized as -CH2CH2CH2CH2NH3+
Attraction between +ve and –ve charged amino acid side chains forms a salt bridge (-COO-+H3N-)
Salt bridge

20. 2) Hydrogen bonding
Hydrogen bonding is a weak attraction between an electronegative atom (O,N,F) in one molecule and a hydrogen atom
in another molecule.
*Has both electrostatic
(ionic) and covalent
character.

21. Water molecule form hydrogen bond with one another
Four hydrogen bonding attraction are possible for each molecule:
*2 through the hydrogen
*2 through the
nonbonding electron
pairs

22. The resulting intermolecular hydrogen bond acts as bridge between water molecules.
Large no. of intermolecular bond (in liquid/solid states of water),the molecules become large, dynamic.
This explain why water have high boiling & melting point.

23. 3)Van Der Waals Forces Force between molecules
Occur between permanent and/or induced dipoles
3 types of van der waals forces :
- Dipole-dipole interactions
- Dipole-induced dipole interactions
- Induced dipole-induced dipole
interactions

24. a) Dipole-dipole interaction
Occur between molecules containing electronegative atoms, cause positive end of one molecule is directed toward negative end of another
eg. Hydrogen bonds are strong type of dipole-dipole interaction

25. b)Dipole-induced dipole interaction
A permanent dipole induces a transient dipole in a nearby molecule by distorting its electron distribution
eg. Carbonyl-containing molecule is weakly attracted to hydrocarbon
Weaker than dipole-dipole interaction

26. c)Induced dipole-induced dipole interactions
Forces between nonpolar molecules
Because of the constant motion of electron, an atom/molecule can develop a temporary dipole (induced dipole) when the electron are distributed unevenly around nucleus
Neighboring atom can be distorted by the appearance of the temporary dipole which lead to an electrostatic interaction between them
Also known as London dispersion forces
eg. Stacking of base ring in DNA molecule

28. Thermal Properties of Water
Hydrogen bonding keeps water in the liquid phase between 0oC and 100oC.
Liquid water has a high:
Heat of vaporization - energy to vaporize one mole of liquid at 1 atm
Heat capacity - energy to change the temperature by 1oC
Water plays an important role in thermal regulation in living organisms.

29. Relationship between temperature and hydrogen bond
Max number of hydrogen bonds form when water has frozen into ice.
Hydrogen bonds is approximately 15% break when ice is warmed.
Liquid water consists of continuously breaking and forming hydrogen bonds.
As the tempt rise, the broken of hydrogen bonds are accelerating.
When boiling point is reached, the water molecules break free from one another and vaporize.

30. Solvent properties of water
Water is an ideal biological solvent
Water easily dissolves a wide variety of the constituents of living organisms.
Water also unable to dissolve some substances
This behavior is called hydrophilic and hydrophobic properties of water.

31. Hydrophilic molecules
Ionic or polar substances that has an affinity for water
In Greek= Hydro, “water” philios, “loving”
Water dipole structure and its capacity to form hydrogen bond with electronegative atoms enable water to dissolve ionic and polar substance
These substances soluble in water due to 3 kinds of noncovalent bonding :
a) ion-dipole
b) dipole-dipole
c) hydrogen bonding

32. Salts (KCl,NaCl) held together by ionic interactions
When ionic compound eg. KCl,NaCl dissolved in water, its ions separate because the polar water molecules attract ions more than the ions attract each other. (ion-dipole interaction)
Shells of water mol. cluster around the ions = solvation spheres

33. Dipole-dipole Interactions
Organic molecules with ionize group
The polar water molecule interacts with carboxyl group of aldehyd & ketones (carbohyd) and hydroxyl group of alcohol
Dipole-dipole
interactions

34. Hydrogen Bonding hydrogen bond shown in yellow
A hydrogen attached to an O or N becomes very polarized and highly partial plus. This partial positive charge interacts with the nonbonding electrons on another O or N giving rise to the very powerful hydrogen bond.
hydrogen bond
shown in yellow

35. Hydrophobic molecules
Non ionic or nonpolar substance
These molecules do not form good attractions with the water molecule. They are insoluble and are said to be hydrophobic (water hating).
eg. Hydrocarbon : CH3CH2CH2CH2CH2CH3, hexane

36. Water forms hydrogen-bonded cagelike structures around hydrophobic molecules, forcing them out of solution. (droplet/into a separate layer)

37. Amphipathic Molecules
Amphipathic molecules contain both polar and nonpolar groups.
Ionized fatty acids are amphipathic. The carboxylate group is water soluble (hydrophilic) and the long carbon chain is not (hydrophobic).
Amphipathic molecules tend to form micelles when mixed with water.

38. polar head – orient themselves in contact with water molecules
Nonpolar tails – aggregate in the center, away from water

39. Osmotic Pressure
Osmosis is a spontaneous process in which solvent (eg water) molecules pass through a semi permeable membrane from a solution of lower solute concentration (dilute) to a solution of higher solute concentration (concentrated).
Osmotic pressure is the pressure required to stop osmosis (22.4 atm for 1M solution)

40. Over time, water diffuses from side B
(more dilute) to side A (concentrated) A A B B

41. Osmotic Pressure
Osmotic pressure (p) is measured using an osmometer.

42. Osmotic Pressure = iMRT
i = van’t Hoff factor (degree of ionization of solute)
M = molarity (concentration of solute in mole/L)
R = gas constant (0.082 L.atm/K.mole)
T = absolute temp (in Kelvin)
Osmolarity = iM (osmol/Liter)

43. i is the van't Hoff coefficient.
For non-electrolytes (non ionizable solute) i=1
For strong electrolytes i= the number of ions that are produced by the dissociation according to the molecular formula
e.g for NaCl you have 2 ions (1 Na+ and 1 Cl-) so i=2
for CaCl2, 3 ions (1 Ca+2 and 2 Cl-) so i=3
For weak electrolytes i=(1-a)+na
n = the number of ions coming from the 100% dissociation according to the molecular formula
a = the degree of dissociation
e.g the degree of ionization of 1M CH3COOH solution is 80%
a=80%/0.8 , n=2
so,
i=(1-0.8) + 2(0.8) =1.8

44. Question 1
1)Estimate the osmotic pressure of a solution 1M NaCl at 25°C. Assume 100% ionization of solute.
 = iMRT
i= 2 (1 Na+ and 1 Cl-)
M= 1 mol/L
R= L.atm/K.mol
T= 298K

45. Question 2
Estimate the osmotic pressure of a solution 0.2M Magnesium chloride at 25°C. Assume 70% ionization of solute.

46. Osmotic pressure is an important factor affecting cells
Cells contain high concentration of solutes – small organic mol., ionic salts, macromolecule
Cells may gain or lose water depend on concentration of solute in their environment.

47. Definitions of solutions
Isotonic – solutions of equal concentration on either side of the membrane
Cells placed in isotonic solution no net movement of water across the membrane
Volume of cells are unchanged bcoz water entering & leaving the cell at the same rate.

49. Hypotonic – solution with a lower solute concentration than the solution on the other side of the membrane
Cells placed in hypotonic solution water moves into the cells
Cause cells rupture
eg. Red blood cells swell & rupture when immersed in pure water (hemolysis)

50. Hypertonic – solution with higher concentration of solutes than the solution on the other side of the membrane
Cells placed in hypertonic solution water moves out the cells
Cause cells to shrink
eg. Red blood cells shrink when immersed in 3% NaCl solution. (crenation)

51. Water ionization, pH, titration, buffer
The self-ionization of water is the chemical reaction in which two water molecules react to produce a hydronium (H3O+) and a hydroxide (OH−) ion.
Water ionization occurs endothermically due to electric field fluctuations between molecules caused by nearby dipole librations resulting from thermal effects, and favorable localized hydrogen bonding.

52. Water dissociates. (self-ionizes)
H2O + H2O = H3O+ + OH-

53. Ions may separate but normally recombine within a few min. to seconds.
Rarely (about once every eleven hours per molecule at 25°C, or less than once a week at 0°C) the localized hydrogen bonding arrangement breaks before allowing the separated ions to return, and the pair of ions (H+, OH-) hydrate independently and continue their separate existence.

54. Ionization of water
may be expressed as
Keq = [H3O+][OH-]
[H2O]2
The conditions for the water dissociation equilibrium must hold under all situations at 25°C.
Kw= [H3O+][OH-] = 1 x 10-14M
Pure water ionize into equal amount of
[H3O+ ] = [OH-] = 1 x 10-7 M

55. Acids, Bases and pH
When external acids or bases are added to water, the ion product ([H3O+ ][OH-] ) must equal.
Kw= [H3O+][OH-] = 1 x 10-14
The effect of added acids or bases is best understood using the Bronsted-Lowry- theory of acids and bases.

56. Bronstead-Lowry theory
Bronsted-Lowry theory is an acid-base theory
Acid is a substance that can donate proton (ion H+ donor)
acid + base = conjugate base + conjugate acid
HCl + H2O = H3O+ + Cl-
Asid Base CA CB
C: conjugate (product) A/B

57. base is a substance that can accept proton
RNH2 + H2O = OH- + RNH3+
B A CB CA
C: conjugate (product) A/B

58. Measuring Acidity Added acids, increase concentration of hydronium ion
In acid solutions [H3O+] > 1 x 10-7 M
[OH-] < 1 x 10-7 M
Added bases, increase concentration of hydroxide ion.
In basic solutions [OH-] > 1 x 10-7 M
[H3O+] < 1 x 10-7 M
pH scale measures acidity without using exponential numbers.

59. pH Scale Define: pH = - log(10)[H3O+]
acidic basic
[H3O+]=1 x 10-7 M, pH = ?

61. pH Scale Questions 1) [H3O+]=1 x 10-5 M, pH = ?

62. pH Scale Questions 1) [H3O+]=2.6 x 10-5 M, pH = ?

63. pH Scale pH to [H3O+]? inverse log of negative pH
orange juice, pH [H3O+]=?
urine, pH 6.2. [H3O+]=?

64. Strength of Acids
Strength of an acid is measured by the percent which reacts with water to form hydronium ions.
Strong acids (and bases) ionize close to 100%.
eg. HCl, HBr, HNO3, H2SO4
eg. NaOH, KOH, CaOH

65. Strength of Acids
Weak acids (or bases) ionize typically in the 1-5% range
eg. Organic acid (contain carboxyl groups)
CH3COCOOH, pyruvic acid
CH3CHOHCOOH, lactic acid CH3COOH, acetic acid

66. Strength of Acids
Strength of an acid is also measured by its Ka or pKa values
Dissociation of weak acid :
HA + H2O = H3O+ + A-
Weak acid conjugate base of HA
Strength of weak acid may be determined :
pKa= -log Ka

67. Strength of Acids Ka pKa CH3COCOOH 3.2x10-3 2.5
CH3CHOHCOOH x
CH3COOH x
Larger Ka and smaller pKa values indicate stronger acids.

68. Monitoring acidity
The Henderson-Hasselbalch (HH) equation is derived from the equilibrium expression for a weak acid.

69. HH equation
The HH equation enables us to calculate the pH during a titration and to make predictions regarding buffer solutions.
What is a titration?
It is a process in which carefully measured volumes of a base are added to a solution of an acid in order to determine the acid concentration.

70. When chemically equal (equivalent) amounts of acid and base are present during a titration, the equivalence point is reached.
The equivalence point is detected by using an indicator chemical that changes color or by following the pH of the reaction versus added base, ie. a titration curve.

71. Titration Curve (HOAc with NaOH)
End point
Equivalence point pH
NaOH (equivalents added)

72. Titration Curve (HOAc with NaOH)
At the end point, only the salt (NaOAc) is present in solution.
At the equivalence point, equal moles of salt and acid are present in solution.
[HOAc] = [NaOAc]
pH = pKa

73. Questions
By using HH equation, calculate the pH of a mixture of 0.25M acetic acid and 0.1M sodium acetate. The pKa of acetic acid is 4.76
pH = pKa + log [A-]
[HA]
pH = log [acetate]
[acetic acid]
pH = log 0.1 = 4.36
0.25

74. 2) Calculate the ratio of lactic acid to lactate in a buffer at pH 5
2) Calculate the ratio of lactic acid to lactate in a buffer at pH The pKa for lactic acid is 3.86
= 13.8

75. 3) During the fermentation of wine, a buffer system consisting of tartaric acid and potassium hydrogen tartrate is produced by a chemical reaction. Assuming that at some time the concentration of potassium hydrogen tartrate is twice that of tartaric acid, calculate the pH of the wine. The pKa of tartaric acid is 2.96

76. pH = pKa + log [A-]
[HA]
= log [hydrogen tertrate]
[tartaric acid]
= log 2
= 3.26

77. Buffer solution
Buffer : a solution that resists change in pH when small amounts of strong acid or base are added.
A buffer consists of:
a weak acid and its conjugate base or
a weak base and its conjugate acid

78. How does buffer work?
Accepting hydrogen ions from the solution when they are in excess
Donating hydrogen ions from the solution when they have depleted

79. Buffer Solutions Maximum buffer effect occurs at the pKa for an acid.
Effective buffer range is at 1 pH unit above and below the pKa value for the acid or base.
eg. H2PO4-/HPO42-, pKa=7.20
buffer range pH

80. Buffer Solutions
High concentrations of acid and conjugate base give a high buffering capacity.
Buffer systems are chosen to match the pH of the physiological situation, usually around pH 7.

81. Physiological buffer 3 most important buffer in body:
Within cells the primary buffer is the phosphate buffer: H2PO4-/HPO42-
The primary blood buffer is the bicarbonate buffer: HCO3-/H2CO3.
Proteins also provide buffer capacity. Side chains can accept or donate protons. (eg. Hemoglobin, serum albumins)

82. A zwitterion is a compound with both positive and negative charges.
Zwitterionic buffers have become common because they are less likely to cause complications with biochemical reactions.

83. N-tris(hydroxymethyl)methyl-2-aminoethane sulfonate (TES) is a zwitterion buffer example.
(HOCH2)3CN+H2CH2CH2SO3-

84. Assignment (water)
Date of submission: 30/7/10
Explain how the changes in temperature give effect to hydrogen bonds in water molecule. Elaborate the situation with drawing of water molecules at every temperature level.
2. Explain how the acids produced in metabolism are transported to the liver without greatly affecting the pH of the blood.

85. Water : The Medium of Life
The End