body fluids

1. ERT106 BIOCHEMISTRY
WATER
Pn Syazni Zainul Kamal

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

3. 9. Ionization of Water
1. Acids, bases, and pH
2. Buffers
3. 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-)
CH2
CH2
COO
-
CH2
CH2
N
H3
+
Salt bridge

20. H
O
H


 H
O
H



H
O
H


 Hydrogen bonding is a
weak attraction
between an
electronegative atom
(O,N,F) in one
molecule and a
hydrogen atom
in another molecule.
2) Hydrogen bonding
*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
H
O
H
H
O
H
H
O
H
H
O
H
H
O
H

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
C O C O
+
-
+
-

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
C O H
H
H
H
+
- +
-

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.  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.
Relationship between temperature and
hydrogen bond

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
K+
Cl-
H
O
H
H
O
H
H
O
H
H
O
H
H
O
H
H
O
H
H
O
H H
O
H

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
H
O
H
H
O H
C
H3
C
CH3
O
H
O
H
+
-
Dipole-dipole
interactions

34. Hydrogen Bonding
 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.
R1 O H
H
O
H
H
O
H
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. B
A B
A
•Over time, water diffuses from side B
(more dilute) to side A (concentrated)

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

42. Osmotic Pressure
p = 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.
p = iMRT
i= 2 (1 Na+ and 1 Cl-)
M= 1 mol/L
R= 0.0821 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.  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.
Definitions of solutions

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.  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
Ionization of water

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.  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
Bronstead-Lowry theory

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+]
0---------------7---------------14
acidic basic
[H3O+]=1 x 10-7 M, pH = ?

61. pH Scale
Questions
1) [H3O+]=1 x 10-5 M, pH = ?
2) [H3O+]=1 x 10-10 M, pH = ?
3) [H3O+]=1 x 10-8 M, pH = ?

62. pH Scale
Questions
1) [H3O+]=2.6 x 10-5 M, pH = ?
2)[H3O+]=6.3 x 10-9 M, pH = ?
3)[H3O+]=7.8 x 10-3 M, pH = ?

63. pH Scale
 pH to [H3O+]?
inverse log of negative pH
 orange juice, pH 3.5. [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 : Ka = [H3O+][A-]
[HA]
pKa= -log Ka

67. Strength of Acids
Ka pKa
CH3COCOOH 3.2x10-3 2.5
CH3CHOHCOOH 1.4x10-4 3.9
CH3COOH 1.8x10-5 4.8
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.
pH = pKa + log [A-]
[HA]

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)
Equivalence point
End point
NaOH (equivalents added)
pH

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
1) 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 = 4.76 + log [acetate]
[acetic acid]
pH = 4.76 + log 0.1 = 4.36
0.25

74. 2) Calculate the ratio of lactic acid to
lactate in a buffer at pH 5.00. The pKa for
lactic acid is 3.86
5.00 = 3.86 + log [lactate]
[lactic acid]
5.00-3.86 = log [lactate]
[lactic acid]
antilog 1.14 = [lactate]
[lactic acid]
= 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]
= 2.96 + log [hydrogen tertrate]
[tartaric acid]
= 2.96 + 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
-/HPO4
2-, pKa=7.20
buffer range 6.20-8.20 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
-/HPO4
2-
 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
1. 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. The End
Water : The Medium of Life