Lecture 4 5

HBC1011 Biochemistry I
Trimester I, 2017/2018
Lecture 4-5: Biochemistry of Water
Ng Chong Han, PhD
MNAR1010, 06-2523751
chng@mmu.edu.my
Tan Chai Hong
MNBR0004 , 06-2523757
tan.chai.hong@mmu.edu.my

Overview
• Water, the Biological Solvent
• Hydrogen Bonding and Solubility
• Cellular Reactions of Water
• Buffer Systems
2

• Water covers about 70% of the Earth’s surface
and makes up 45-95% of living organisms
3

Water = molecule of life
One of NASA‘s guiding policies in the search for alien life is to “follow the
water”. Water is fairly common in the universe, but most of this water is in
the form of ice. Solid water can't act as a lubricant for the molecular
processes of life, so the search is for liquid water - a commodity that is far
more rare in the universe.
4

5
NASA Mars
Exploration
Programme
Curiosity and
Water on Mars

6
Water in our bodies
Approximately 55% of blood is plasma
Plasma is mainly composed of:
90% of WATER
10%: blood proteins,
inorganic
ions,
glucose,
lipids,
amino acids,
hormones,
metabolic end products….
Each of us has 5-6 liters of blood

7
Biological roles of water
Water is a biological solvent
eg. Biological fluids for delivery of nutrients and removal of
wastes
Water serves as an essential buffer
to regulate temperature (high specific heat capacity) and pH
(buffer solutions)
Water is a participant in many biochemical reactions
eg. the principal reactant in the photosynthesis process

Physical and chemical properties of
water
8
• Chemical formula H2O: one molecule of
water has two hydrogen atoms
covalently bonded to a single oxygen
atom.
• Water appears in nature in all three
common states of matter (solid, liquid,
and gas)
• Liquid at standard temperature and
pressure, tasteless and odorless.
• The intrinsic colour of water and ice is a
very slight blue hue, although both appear
colorless in small quantities.

• Polarity: a separation of electric charge leading to its chemical
groups having an electric dipole or multipole moment.
• Polar molecules interact through dipole–dipole intermolecular
forces and hydrogen bonds. Molecular polarity is dependent on
the difference in electronegativity (a pull on the electrons)
between atoms in a compound.
• Polarity affects physical properties eg. surface tension, solubility,
and melting and boiling-points.
9
Polar Molecules
The two charges are present with a negative
charge in the middle (red shade), and a
positive charge at the ends (blue shade).

• The polarity of chemical bonds: Non-polar bonds, Polar
bonds, Ionic bonds
• Bonds can fall between one of two extremes — being
completely nonpolar or completely polar.
– A completely nonpolar bond occurs when the
electronegativities are identical.
– A completely polar bond is more correctly called an ionic
bond.
• The terms "polar" and "nonpolar" are usually applied
to covalent bonds, that is, bonds where the polarity is
not complete.
• While the molecules can be described as "polar
covalent", "nonpolar covalent", or "ionic", this is often a
relative term.
10
Polar Molecules

11
Covalent and Noncovalent bonds
• Covalent bond: a chemical bond that involves the sharing of electron
pairs between atoms.
• Noncovalent bond: it does not involve the sharing of electrons, but
rather involves more dispersed variations of electromagnetic
interactions between molecules or within a molecule, eg ionic bond.

Covalent bonds
12
• If the two atoms have about equal attraction for electrons, a
nonpolar covalent bond (equal sharing) forms
• If the two atoms have different electronegativities, a polar
covalent bond (unequal sharing) forms
In polar bonds, one atom is slightly negative (δ-), while the
other is slightly positive (δ+)
• Electronegativity is a chemical property that describes the
tendency of an atom or a functional group to attract electrons
towards itself.

13
Noncovalent bonds
• Noncovalent bonds are important for
determining the structure of biomolecules.
• Types of noncovalent bonds or interactions
include:
– Hydrogen bonds
– Ionic bonds
– van der Waals forces
– Hydrophobic interactions

Noncovalent bonds
14
• Hydrogen bonds occur when a hydrogen atom (δ+) is
attracted by a negatively charged atom (δ-)
• Hydrogen bonds are weak (10% as strong as covalent).

Noncovalent bonds
15
• Ionic bonds are attractions between oppositely charged
atoms:
+  - for example, Na+ Cl-
• Positively charged atoms
are called cations,
always move to the
cathode.
• Negatively charged
atoms are called anions,
always move to the
anode.

Noncovalent bonds
17
• Hydrophilic molecules are ones that interact with water. These
molecules are polar: charged (+, -) or partially charged (δ+,δ-)
• Hydrophobic molecules do not interact with water, like oils and
fats. These molecules are nonpolar. These often have C’s and
H’s, but few or no O’s and N’s.

18
Noncovalent bonds
• Hydrophobic molecules tend to avoid water, hence
tend to bond together. This is called hydrophobic
interactions
• Very weak interactions between nonpolar molecules
that are tightly packed together are called van der
Waals forces
• Although these noncovalent interactions are
individually weak relative to covalent bonds, the
cumulative effect of many such interactions can be very
significant.

20
Water: Polar Molecules
• Since the water molecule is not
linear and the oxygen atom has a
higher electronegativity than
hydrogen atoms, the oxygen atom
carries a slight negative charge,
whereas the hydrogen atoms are
slightly positive.
• As a result, water is a polar
molecule with an electrical dipole
moment.

21
Water: Polar Molecules
• Water also can form an
unusually large number of
intermolecular hydrogen
bonds (four) for a molecule of
its size.
• These factors lead to strong
attractive forces between
molecules of water, giving rise
to water's high surface
tension and capillary forces.

POLES help to initiate…
HYDROGEN BONDING between
molecules!
THIS IS THE CHEMICAL BASIS FOR MOST OF
WATER’S ACTION IN LIFE PROCESSES
22

23
Hydrogen bonds in water
n Partial charges on atoms in water allow bonds to form
between molecules
• Hydrogen bonding
• results when H from one molecule is attracted to O of
a different molecule
d-
d+
d+
d-
O

1
4
3
2
Each molecule forms hydrogen bonds with 4
other molecules.
.
The H atom of one molecule of water interacts
with a lone pair of electrons in an orbital
of the O atom of another water molecule
24

These four hydrogen bonds increase the space the
water molecules take up, so water expands as it
freezes, and ice is less dense than liquid water.
For these reasons, ice floats in liquid water
25

27
Biological importance of hydrogen
bonds
a. Between an alcohol and
water or between alcohol
molecules.
b. Between a carbonyl group
and water
c. Between 2 peptide chains
d. Between 2 complementary
base pairs in DNA

Water and H-bonds
28
• Hydrogen bonding between water molecules gives
water its special properties
– Cohesion
– Adhesion
– Surface Tension
– Temperature Moderation

29
Water Properties - Cohesion
• High Cohesion
– Binding together of like molecules by H bonds
– High in water
– H-bonds constantly breaking and reforming
– most water molecules are bonded to neighboring
molecules at any instant
– Contributes to water transport in plants

30
Water Properties - Adhesion
• High Adhesion
– Clinging of one substance to another
– also involves H-bonds
– also contributes to water transport in plants
– water adheres to molecules of the walls of the
xylem vessels in plant stems (trunks) helps counter
the effects of gravity

31
Water Properties - Surface Tension
• High Surface Tension
– Measure of how difficult
it is to stretch or break
the surface of a liquid
– Higher in water than
other liquids

32
Water – Temperature moderation
High specific heat
Specific Heat: is the amount of heat that must be absorbed or lost
for one gram of a substance to change its temperature by 1°C.
Lots of heat is needed to break H-bonds and raise H2O temperature.
Therefore, H2O is a good insulator.
DAY or SUMMER NIGHT or WINTER
HEAT
HEAT

33
Water: solvent for life
Molecules dissolve in water,
which allows them to move
around more and interact.
Water facilitates all chemical
reactions in the body.
Since water is polar, the positive
and negative ends of a water
molecule will be attracted to
charged ions or other polar
molecules

34
Water and molecules
• Hydrophilic: hydros (water) and philia
(friendship)
– Ionic compounds dissolve in water
– Polar molecules (generally) are water soluble
• Hydrophobic: hydros (water) and phobos
(fear)
– Nonpolar compounds

35
Hydrophobicity
Solvent
Solute
Solution
• Polar molecules tend to be
hydrophilic.
• Substances that are ionic or
polar often dissolve in water
due to hydrogen and ionic
bonds.

How ionic solids dissolve
H
H
H
H
H
36

Hydrophobicity
37
• Hydrophobic compounds
and H2O don’t mix.
• Amphiphilic molecules are
part hydrophobic and part
hydrophilic
Amphiphile: amphis (both)
and philia (love, friendship)

38
Amphipathicity
• When an amphipathic compound is mixed with water,
the polar, hydrophilic region interacts favorably with
the solvent and tends to dissolve, but the nonpolar,
hydrophobic region tends to avoid contact with the
water.
It forms a stable structure,
called micelle in the water.

39
Cellular reactions of water
• Occasionally, a hydrogen atom shared by two water molecules shifts
from one molecule to the other.
– The hydrogen atom leaves its electron behind and is transferred
as a single proton - a hydrogen ion (H+).
– The water molecule that lost a proton is now a hydroxide ion
(OH-).
– The water molecule with the extra proton is a hydronium ion
(H3O+).

40
H2O H+ + OH-
Hydrogen
ion
Hydroxide
ion
• Reversible reaction
• At equilibrium the concentration of water molecules greatly
exceeds that of H+ and OH-.
• At equilibrium, the concentration of H+ or OH- is 10-7M (25°C)
• Hydroxide ions can accept a proton and be converted back into
water molecule
Dissociation of water molecules

41
• The ionization of water can be analyzed quantitatively.
• The concentrations of the reactants and the products at equilibrium:
The ratio of these concentrations defines the equilibrium constant
(Keq).
In case of water ionization: Keq= [H+] [OH-]
[H2O]
The concentration of water at equilibrium:
the mass of 1 liter of water is 1000g
And the mass of one mole of water is 18g
the pure water has a concentration of: 1000g/l = 55.5 mole/l or
18g/mole
= 55.5 M

42
Keq = [H+] [OH-]
[H2O]
Keq(55.5 M) = [H+] [OH-]
The Keq for the ionization of water has been determined under
standard conditions of pressure (1 atm) and temperature (25°C)
(1.8x10-16 M)(55.5 M) = [H+] [OH-]
1.0 x 10-14 M2 = [H+] [OH-]
Its value is: Keq= 1.8x10-16 M (the electrical conductivity
of pure water)

43
Because according to the chemical equation for
dissociation H+ and OH- must have equal concentrations in
pure water, then
Kw(ion product of water)=
[H+][OH-]=[H+]2=1.0 x 10-14 M2
[H+]=1.0 x 10-7 M
[H+]= 10-7 M =[OH-]
Hydrogen ion concentrations expressed in exponential form are
difficult to work with. A more useful terminology is pH, defined
as the negative logarithm of the [H+].

When [H+] = [OH-]
The solution is said
Neutral
When [H+] > [OH-] Acidic
When [H+] < [OH-] Basic
44
• The pH of a solution will depend little on the hydrogen
ions generated by the self-dissociation of water, but
rather on the presence of other substances (acids or
bases) that increase or decreases the H+ concentration.
• Acids and bases are chemical substances that change
the ionic properties of solutions.

45
Acids, bases and pH scale
• Some substances dissolve in water and release
hydrogen ions (H+); these are called acids. Their release
is called ionization.
• Acids release H+ ions in solution.
• If the reaction is complete, it is a strong acid, such as
HCl.
Hydrochloric acid in water
HCl H+ + Cl-

46
• Other substances dissolve in water and release
hydroxide ions (OH–); these are called bases.
• Bases accept H+ in solution.
• NaOH ionizes completely to Na+ and OH–. The OH–
absorbs H+ to form water. It is a strong base.
Sodium Hydroxide in water
NaOH Na+ + OH-

47
• Ionization of strong acids is virtually irreversible.
• Ionization of weak acids and bases is somewhat reversible.
• Many large molecules in biological systems contain weak acid or base
groups.
n Water is really a weak acid and has a slight tendency to ionize
into H+ and OH–.
n This ionization is very important for living creatures and the
chemical reactions they must perform because the H+ ion is so
reactive.

• Range from 0 to 14
• Basic pH > 7
• neutral pH = 7
• acidic pH < 7
The small p in pH stand for “potential” or “power”
48

pH is a negative
logarithmic
expression of
[H+]
49

50
• In pure H2O, [H+] and [OH-] = 10-7 M
pH = - log [H+]
So the pH of pure water is 7 Neutral

• Acids are compounds that donate protons, and bases are
compounds that accept protons.
• Strong acids, such as hydrochloric acid (HCl), dissociate
completely.
• Weak acids, such as acetic acid, dissociate only to a limited
extent:
where HA is the acid, and A- is its conjugate base.
• The dissociation constant for a weak acid is
51

52
Weak acids and the acid
dissociation constant (Ka)
• The stronger the acid, the lower the pKa the stronger the
base, the higher its pKa.
• The pKa can be determined experimentally; it is the pH at
the midpoint of the titration curve for the acid or base.
Ka = [H+][CH3COO-]
[CH3COOH]
pKa is a measure of acid strength

Some compound, such as acetic acid, is monoprotic; it can give up only one proton.
Others are diprotic (H2CO3 (carbonic acid) or triprotic (H3PO4 (phosphoric acid)). The
dissociation constant (Ka) and its negative logarithm, the pKa for each pair are shown
on a pH gradient.

Titration curve of weak acids
54
Titration is used to determine the
amount of an acid in a given solution.
A measured volume of the acid is titrated
with a solution of a strong base, usually
sodium hydroxide (NaOH), of known
concentration.
The NaOH is added in small increments
until the acid is consumed (neutralized),
as determined with a pH meter.
The concentration of the acid in the
original solution can be calculated from
the volume and concentration of NaOH
added.

Titration curve of weak acids
55
This value is plotted against the amount of
NaOH expressed as a fraction of the total
NaOH required to convert all the acetic acid
to its deprotonated form, acetate.
The points yield the titration curve. At the
midpoint of the titration, the concentrations
of the proton donor and proton acceptor
are equal, and the pH is numerically equal
to the pKa.
The shaded zone is the useful region of
buffering power, generally between 10%
and 90% titration of the weak acid.

Buffers
• Almost every biological process is pH dependent; a small change in
pH produces a large change in the rate of the process. The
enzymes that catalyze cellular reactions, and many of the
molecules on which they act, contain ionizable groups with
characteristic pKa values.
• In cells and tissues, phosphate and bicarbonate buffer systems
maintain intra-cellular and extra-cellular fluids at their optimum
(physiological) pH, which is usually close to pH 7. Enzymes
generally work optimally at this pH.
56

Optimal pH of some enzymes
• Pepsin is a digestive enzyme
secreted into gastric juice; trypsin,
a digestive enzyme that acts in the
small intestine; alkaline
phosphatase of bone tissue, a
hydrolytic enzyme thought to aid
in bone mineralization.
57

Buffers
58
Water has a very small [H+] (10-7).
Adding just a little bit of acid or base can change the pH
drastically.
Add 0.001 M HCl: pH goes from 7 to 3!
For many applications this sensitivity is undesirable.
One of the best ways to prevent pH swings is buffering:
the use of a mixture of a weak acid and its conjugate
base (which will be a weak base).

Buffers
59
How can pH changes be minimized?
• Buffers
– Substances that minimize changes in [H+] in
solution
– Present in all biological fluids
• Human blood maintained at pH 7.4
How do buffers work?
• accept H+ ions from the solution when in excess
• donate H+ ions to the solution when depleted

Buffers
• Whenever H+ 􏰒 or OH- 􏰒 is added
to a buffer, the result is a small
change in the ratio of the
relative concentrations of the
weak acid and its anion and
thus a small change in pH.
• The decrease in concentration
of one component of the
system is balanced exactly by an
increase in the other.
• The sum of the buffer
components does not change,
only their ratio.
60

• The shape of the titration curve of any weak acid is described by
the Henderson-Hasselbalch equation.
• For the dissociation of a weak acid HA into H􏰒 and A􏰒, the
Henderson-Hasselbalch equation can be derived as follows:
61
A Simple Expression Relates pH, pKa,
and Buffer Concentration

Buffer action
• The addition of the acid to 0.1 M sodium acetate solution results
in a much gradual change in pH until the pH drops below 3.5.
• When hydrogen ions are added to the solution, they react with
acetate ion to form acetic acid. This reaction consumes some of
the added hydrogen ion so that the pH does not drop. Hydrogen
ions continue reacting with acetate ions until essentially all of
the acetate ion is converted into acetic acid. After this point,
added protons remain free in solution and the pH begins to fall
sharply again.
63

Summary
1. Water, a nonlinear, polar molecule, serves at least three
functions in the cell: It is an effective solvent, it is a reactant
molecule, and it is a temperature buffer. As a solvent, water is
able to dissolve biomolecules that are ionic and polar.
2. The most important reaction of water is its reversible ionization
to generate proton (H+) and the hydroxide ion (OH-). The extent
of ionization is quantified by the pH scale (pH = -log [H+]).
3. The strength of an acid is defined its pKa, the negative log of its
dissociation constant.
4. Blood and other cellular fluids are maintained at a constant pH
by natural buffer systems.
64

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