CHAPTER 2

WATER

BOOK:LEHNINGER
water is the most abundant substance in living systems, making up 70% or more of
the weight of most organisms.

Physical and chemical properties of water, to which all aspects of cell structure and
function are adapted.

The attractive forces between water molecules and the slight tendency of
water to ionize are of crucial importance to the structure and function of
biomolecules.

We review the topic of ionization in terms of equilibrium constants, pH, and

titration curves, and consider how aqueous solutions of weak acids
or bases and their salts act as buffers against pH changes in biological systems.

The water molecule and its ionization products, H and OH, profoundly influence
the structure, self-assembly, and properties of all cellular components, including
proteins, nucleic acids, and lipids.

The noncovalent interactions responsible for the strength and specificityof

“recognition” among biomolecules are decisively influenced by the solvent
properties of water, including its ability to form hydrogen bonds with itself and
with solutes.
 Water

 Unusual properties of water:TM,

Boiling…..
 Structure of water:polar, tetrahedral….
 Ability of water molecules to form
hydrogen bonds
 Structure of ice
 1. water and its unusual properties

 Water has a higher melting point, boiling point, and heat of vaporization than
most other common solvents. attractions between adjacent water molecules
that give liquid water great internal cohesion (hydrogen bond = H bond).
Structure of a water molecule
2. Structure of the water molecule
1. Polar

2. Tetrahedral/ five molecules of water

form a tetrahedral structure by hydrogen
bonding.

3. bond angle is 104.5o and not 109.5o of

a perfect tetrahedron **due to electron
crowding of the nonbonding orbital of
the oxygen atom.

4. Sharing of electrons is unequal.The

oxygen nucleus attracts electron more
strongly than does the hydrogen nucleus
(PROTON)
the electrons are more often in the
vicinity of the oxygen atom (2d-) than the
hydrogen (d+), which generate an
electrostatic attraction between the
oxygen atom of one water molecule and
the hydrogen of another---Hydrogen
bond.
3. Ability to form Hydrogen bond
1. Length of hydrogen bond vs covalent
2. Hydrogen bonds are relatively weak.

bond dissociation energy (the energy

required to break a bond) of about 23
kJ/mol, compared with 470 kJ/mol for
the covalent O—H bond in water.

Bond dissociation
energy

23 kJ/mol Weak bond

470 kJ/mol Strong bond

 Life span hydrogen bond is just 1 to 20 picoseconds (where 1 ps= 10-12 s)

 On breaking of one hydrogen bond, another hydrogen bond is formed in 0.1

ps.

 Phrase for this phenomena is known “flickering clusters” applied to the

short-lived groups of water molecules interlinked by hydrogen bonds in
liquid water.

 Water molecules are disorganized and in continuous motion. At any

moment one (1) molecule of water forms 3.4 H-bonds in liquid state.
H2O in regular lattice structure 1. Density of ice is less than
water in liquid.

2. Ice form Hexagonal lattice by

forming H-bond. thus ice floats
on liquid water.

This crystal lattice of ice makes

it less dense than liquid and
regular.

3. Bond length of hydrogen

bond in ice is 0.274nm (whereas
in liquid state the H-bond length
is 0.177nm).

4. The maximum number of

hydrogen bonds that one water
molecule can have with
neighboring water molecules: 4
 Hydrogen bond not unique to water.

Common hydrogen bonds in biological systems. The hydrogen

acceptor is usually oxygen or nitrogen; the hydrogen donor is another
hydrogen that bonded with electronegative atom.
 Water forms Hydrogen Bonds with Polar solutes
1. Hydrogen bonds are not unique to water.
2. They readily form between an electronegative atom (the hydrogen
acceptor, usually O or N with a lone pair of electrons) and a hydrogen
atom covalently bonded to another electronegative atom (the hydrogen
donor) in the same or another molecule.
3. Hydrogen atoms covalently bonded to carbon atoms do not
participate in hydrogen bonding, because carbon is only slightly more
electronegative than hydrogen and thus the C-H bonds is only very
weakly polar.
4. Alcohols, aldehydes, ketones and compounds containing N-H bonds
all form hydrogen bonds with water molecules and tend to be soluble in
water.
Some biologically important hydrogen bonds
Directionality of the hydrogen bond

Acceptor atom O

 Hydrogen bonds are strongest when the bonded

molecules oriented to maximize electrostatic interaction,
which occurs when the hydrogen atom and the two
atoms that share it are in a straight line - that is, when
the acceptor atom is in line with the covalent bond
between the donor atom and H - holding two hydrogen
bonded molecules or groups in a specific geometric
arrangement.
 Charged or polar molecules are dissolved easily in water (Hydrophilic).
 Nonpolar solvents are poor solvents for polar molecule but easily dissolved
hydrophobic compound (lipid or waxes).
 Amphipathic compounds contain regions that are polar or (charged) and nonpolar
What is the difference between Enthalpy and entropy?
• Enthalpy is (number and kind of bond)the heat transfer
takes place in a constant pressure.
Entropy gives an idea of the randomness of a system.
Spontaneous reaction takes place in order to increase the
universal entropy.

• In a reaction, enthalpy change can be positive or negative.

• Enthalpy is the energy released or absorbed during a
reaction. Enthalpy is the energy surrounded by the system.

• Enthalpy is related with the first law of thermodynamics

where it says, “Energy can be neither created nor destroyed.”
But the entropy is directly related with the second law of
thermodynamics.
Entropy Increases as Crystalline Substances Dissolve

Water as solvent.
1. Water dissolves many crystalline salts by hydrating their component ions.
2. The NaCl crystal lattice is disrupted as water molecules cluster about the Cl and Na
ions.
3. The ionic charges are partially neutralized, and the electrostatic attractions necessary
for lattice formation are weakened.
Dissociation of molecules in water

 Electrolytes/non electrolytes.
 Sugar or acohol are not electrolytes / they
dissolve in water but not carry a charge or
dissociate into charged species.
Water dissolves salts such as NaCl by hydrating and stabilizing the Na+
and Cl- ions, weakening the electrostatic interactions between them.

The same factors apply to charged biomolecules, compounds with

functional groups such as ionized carboxylic acids (-COO-), protonated
amines (-NH4+) and phosphate esters or anhydrides.

Water readily dissolves such compounds by replacing solute-solute

hydrogen bonds with solute-water hydrogen bonds.

As a salt such as NaCl dissolves, the Na+ and Cl- ions acquire far greater
freedom of motion.

The resulting increase in entropy (randomness) of the system is largely

responsible for the ease of dissolving salts such as NaCl in water.
DG = DH - TDS, where DH has a small positive value and TD S a large positive value; thus DG is
negative.(H:enthalpy(heat), S:entropy).
Nonpolar Gases Are Poorly Soluble in Water

Why CO2

Polar molecules dissolves better at low T than do nonpolar molecules at high T

 Nonpolar compounds force energetically unfavorable
changes in the structure of water

Amphipathic compounds in
aqueous solution.
(a) Long-chain fatty acids have
very hydrophobic alkyl chains,
each of which is surrounded by a
layer of highly ordered water
molecules.

No compensation like NaCl

 The stable structure of Amphipathic compound in water=micelles.

Amphipathic compounds in
aqueous solution.
(b) By clustering together in micelles,
the fatty acid molecules expose the
smallest possible hydrophobic surface
area to the water, and fewer water
molecules are required in the shell of
ordered water. The energy gained by
freeing immobilized water molecules
stabilizes the micelle.
 The forces that hold the nonpolar regions of the molecules
together are called hydrophobic interactions

 The strength of hydrophobic interactions is not due to any

intrinsic attraction between nonpolar moieties. but from the
system's achieving greatest thermodynamic stability by
minimizing the number of ordered water molecules required
to surround hydrophobic portions of the solute molecules.

 Hydrophobic interactions among lipids, and between lipids

and proteins, are the most important determinants of
structure in biological membranes
What happen with enzyme
substarte complex??

Release of ordered water !

Release of ordered water favors formation

of an enzyme-substrate complex

Hydrophobic interaction between

nonpolar amino acids also stabilizes
the three-dimensional structures of
proteins.

Part of the driving force for

binding of a polar substrate (reactant)
to the complementary polar surface of
an enzyme is the entropy increase as
the enzyme displaces, ordered water
from the substrate
Weak Interactions Are Crucial to Macromolecular
Structure and Function
The folding of a single polypeptide or
polynucleotide chain into its three-
dimensional shape is determined by
this principle.

The binding of an antigen to a specific

antibody depends on the cumulative
effects of many weak interactions

The energy released when an enzyme

binds non covalently to its substrate is
the main source of the enzyme's
catalytic power.

The binding of a hormone or a

neurotransmitter to its cellular receptor
protein is the result of weak
interactions
 Solutes of all kinds alter the
colligative properties (tied
together) of aqueous
solutions.
Solutes affect the
the effect of solutes on all four colligative properties
properties has the same basis: the
concentration of water is lower in of aqueous
solutions than in pure water. solutions— the
concentration of water is
lower in solution than in
pure water (↓Vapor
effect of solute concentration pressure, ↑boiling point,
on the colligative properties of ↓melting point and
water is independent of the
chemical properties of the solute; ↑osmotic pressure)
it depends only on the number of
solute particles (molecules,ions) in a
given amount of water.
 Osmotic concentration,

 Is the measure of solute concentration.

 Osmolarity measures allows the measurement

of the osmotic pressure of a soln / the
determination of how the solvent will diffuse
across a semipermeable membrane (osmosis)
separating two solutions of different osmotic
concentration.
Osmosis and the measurement of osmotic
pressure How water move?

equilibrium

Osmotic pressure
Force that is needed to retur
soln in tube to the level
in that beaker
Osmosis
 How water move?
 higher water
concentration to one of
lower water concentration.
 Isotonic (equal
osmolarity); hypertonic
(higher osmolarity than
the cytosol); hypotonic
(lower osmolarity)
Naturally, cells generally contain higher concentration of
biomolecules and ions than water into the cell, this would
eventually cause bursting of the cell (osmotic lysis).

Several mechanisms have evolved to prevent this catastrophe.

1. Blood plasma and interstitial fluid (the extracellular fluid of

tissues) are maintained at an osmolarity close to that of the
cytosol. The high concentration of albumin and other proteins in
blood plasma contributes to its osmolarity.

2. Cells also actively pump out ions such as Na+ into the
interstitial fluid to stay in osmotic balance with their surroundings
Because the effect of solutes on osmolarity depends on the number of
dissolved particles, not their mass, macromolecules (proteins, nucleic
acids, polysaccharides) have for less effect on osmolarity of a solution
than would an equal mass of their monomeric components.

Example: a gram of a polysaccharide composed of 1,000 glucose units

has the same effect on osmolarity as a mg of glucose.

One effect of storage fuel as polysaccharides (starch or glycogen) rather

than as glucose or other simple sugars is prevention of an enormous
increase in osmotic pressure within the storage cell.
2.2 Ionization of Water, Weak Acids, and
Weak Bases
 Pure Water Is Slightly Ionized
 The Ionization of Water Is Expressed by an Equilibrium
Constant
 The pH Scale Designates the H+ and OH- Concentrations
 Weak Acids and Bases Have Characteristic Dissociation
Constants
 Titration Curves Reveal the pKa of Weak Acids
Hydronium ions and proton hopping
 1. Hydrogen ions formed
in water are immediately
hydrated to hydronium
ions (H3O+).
 2. No individual proton
moves very far through
the bulk solution, but a
series of proton hops
between hydrogen
bonded water molecules
causes the net movement
of a proton over a long
distance in a remarkably
short time.
The ionization of water is expressed by an
equilibrium constant (Keq)

ion-product of water=Kw
Kw is found by multiplying
Keq by the concentration of
water.

Application exercises
 The ion product of water, Kw, is the
basis for the pH scale

pH = -log [H+]

if concentration is 1.0×10-7M
pH = log (1.0×10-7) = 7

The pH of an aqueous solution reflects,

on a logarithmic scale, the
concentration of hydrogen ions
The pH Scale & some aqueous fluids

The pH scale is logarithmic, not

arithmetic.
Ex: two solutions differ in pH by 1 pH
unit means that one solution has ten
times the H concentration of the other.

A cola drink (pH 3.0) or red wine (pH 3.7)

has an H concentration approximately
10,000 times that of blood (pH 7.4).
Weak acids and bases have characteristic dissociation constants

HA H+ + A-
Keq 
H A 
 
K
HA a

1
pKa  log   log K a
Ka
pKa = _ log Ka (analogous to pH)

Equilibrium constants for ionization reactions are usually called

ionization or dissociation constants, often designated Ka.
Henderson-Hasselbalch equation

Ka 
H A 
  1. solve for H conc
HA 
H   Ka A 


HA 

 logH    log Ka  log

 HA 
2. take -log

A 
3. Substitute pH for –log H conc)

pH  pKa  log
A  

HA 4. Inver

pH  pKa  log
 proton _ acceptor  Conjugated base

 proton _ donor Conjugated acid

 The Henderson-Hasselbalch equation
can be used to calculate

1) the pH of a solution of an organic acid.

2) the amount of salt and acid to add to form a
specific buffer.
3) the pKa of a weak acid.
Acids: proton donors

Bases: proton acceptors.

A proton donor and its corresponding proton

acceptor make up a conjugate acid-base pair
Strong acids/bases (Completely ionized)

 Hydrochloric  NaOH
 Sulfuric  KOH
 Nitric acids

Weak acids/bases (not completely ionized)

 Important in biological systems.

 Metabolism regulation
The stronger the acid, the greater its tendency to lose its
proton.

The tendency of any acid (HA) to lose a proton and

form its conjugate base (A-) is defined by the equilibrium
constant (Keq) for the reversible reaction

Equilibrium constants for ionization reactions are usually

called ionization or dissociation constants, often
designated Ka.
pKa = _ log Ka (analogous to pH)

The stronger the tendency to dissociate a proton, the

stronger is the acid and the lower its pKa
Conjugate acid-base pairs consist of a proton donor
and a proton acceptor
Titration curves reveal the pKa of weak acids (acetic
acid)..acetate anion
 The conc. of the acid in a givin
solution can be calculated from
the volume and concentration of
NaOH added--- titration curve.
 See the axis: PH and NaOH
amount added
 At the midpoint of the titration, at
which exactly 0.5 equivalent of
NaOH has been added, [HA]=[A-]
and pH=pKa
 At this midpoint a very important
relationship holds: the pH of the
equimolar solution of acetic acid
and acetate is exactly equal to the
pKa of acetic acid (pKa = 4.76)

Keq 
H A 

K


HA a

Flat zone:is the buffering region of the acetic acid acetate buffer pair
By definition one equivalent (or equivalent weight) of a substance is the
amount of that substance which supplies or consumes one mol of
reactive species. In acid-base chemistry the reactive species is the
hydrogen ion (H1+) 5) At the midpoint of a titration curve
A) the concentration of a conjugate base is equal to the concentration
of a conjugate acid.
B) the pH equals the pKa.
C) the ability of the solution to buffer is best.
D) All of the above.
Titration Curves
Titration curves are used to determine the amount of an acids in a given
solution. A measured volume of the acids is titrated with a solution of a
strong base, usually NaOH of known concentration. A plot of pH against
the amount of NaOH added (a titration curve) reveals the pKa of the
weak acid. Consider the titration of 0.1M solution of acetic acid (HAc)
with 0.1M NaOH at 25oC

As Na OH is gradually introduced, the added OH- combines with the free

H+ in the solution to form H2O. As free H+ is removed, HAc dissociates
further to satisfy its own equilibrium constant, the net result as the
titration proceeds is that more and more HAc ionizes, forming Ac-, as the
NaOH is added.

At the midpoint of the titration at which exactly 0.5 equivalent of NaOH

has been added, one half of the original acetic acid has undergone
dissociation, so that the concentration of the proton donor (HAc) now
equals that of the proton acceptor, [Ac-]. At this midpoint a very important
relationship holds: the pH of the equimolar solution of acetic acid and
acetate is exactly equal to the pKa of acetic acid (pKa = 4.76)
Buffer are mixtures of weak acids and
their conjugate bases
 Buffers tend to resist changes in pH
when small amounts of acids or base
added.
 At 0.5 NAOH added, proton acceptor
eq.proton donor

Buffers Are Mixtures of Weak Acids and Their Conjugate Bases

By definition one equivalent (or equivalent weight) of a substance
is the amount of that substance which supplies or consumes one
mol of reactive species. In acid-base chemistry the reactive species
is the hydrogen ion (H1+)

Q1) At the midpoint of a titration curve

A) the concentration of a conjugate base is equal to the
concentration of a conjugate acid.
B) the pH equals the pKa.
C) the ability of the solution to buffer is best.
D) All of the above.
Comparison of the titration curves of three
weak acids
 The titration curves of these
acids have the same shape,
they are displaced along the
pH axis because the three
acids have different strengths.
1. Acetic acid, with the highest
Ka (lowest pKa, lowest pH) of
the three, is the strongest
(loses its proton most readily);
it is already half dissociated at
pH 4.76.
 2. Dihydrogen phosphate
loses a proton less readily,
being half dissociated at pH
6.86.
 3. Ammonium ion is the
weakest acid of the three and
does not become half
dissociated until pH 9.25.
2.3 Buffering against pH Changes in
Biological Systems

 Buffers Are Mixtures of Weak Acids and

Their Conjugate Bases
 A Simple Expression Relates pH, pKa,
and Buffer Concentration (Henderson-
Hasselbalch equation)
 Weak Acids or Bases Buffer Cells and
Tissues against pH Changes
 Buffering against pH Changes in
Biological Systems

Weak acids or bases buffer cells and tissues against

pH changes
Work example

 cytoplasm cells contain ionizable groups:

Protein (histidine pKa=6.0, buffer near
neutral PH), nucleotides such ATP, low
MW metabolites.
The organism’s first line of defense against
changes in
internal pH is provided by buffer systems
 Two Important biological buffer systems are the
phosphate and bicarbonate systems.

 The phosphate buffer system acts in the

cytoplasm of all cells.

 Maximally effective at a pH close to its pKa of

6.86 and thus tends to resist pH changes in
the range between about 5.9 and 7.9. It is
therefore an effective buffer in biological fluids;

 in mammals, extracellular fluids and most

cytoplasmic compartments have a pH in the
range of 6.9 to 7.4.
 The blood buffering system
 The buffering capacity of blood depends on 2 equilibria:

1. between gaseous CO2 dissolved in blood and carbonic acid formed by the
reaction.
CO2+H2O…………………………………..H2CO3

2. Between carbonic acid and bicarbonate formed by dissociation of H+.

H2CO3…………………………………..H+ + HCO3
 CO2 is a gas under normal conditions, and the concentration of dissolved CO2
is the result of equilibration with CO2 of the gas (g) phase:
 CO2 (g) …………………..CO2 (d)
 Hypoventelation (increased CO2): acidosis
 Hyperventilation ( decrease CO2): alkalosis

Acidosis: PH low as 7.1

Alkalosis: PH high as 7.6
 Blood plasma is buffered in part by the bicarbonate
system…………………….complex system
Carbonic acid bicarbonate
 Reaction 1: When H is added to
blood as it passes through the
tissues,- proceeds toward a new
equilibrium, in which the
concentration of H2CO3 is increased.
 Reaction 2 :This increases the
concentration of CO2(d) in the blood
plasma
 Reaction 3: thus increases the
Lactic acid formation pressure of CO2(g) in the air space
NH3 produced during protein catabolism of the lungs .The extra CO2 is
exhaled.
 Conversely, when the pH of blood
plasma is raised (i.e. by NH3
production during protein
catabolism),
 The H concentration of blood
plasma is lowered, causing more
H2CO3 to dissociate into H and
HCO3. This in turn causes more
CO2(g) from the lungs to dissolve
Enzymes typically show maximal catalytic activity at a
characteristic pH - pH optima
2.4 Water as a reactant

 Condensation
reaction: in which the
element of water are
eliminated.
 Hydrolysis reaction:
cleavage
accompanied by the
addition of the water
(reverse)