1
MB ChB PHASE I
WATER
LECTURE 1
AIM: To review:
the role of water as a solvent.
http://www.abdn.ac.uk/~bch118/index.htm
2
Life evolved in and is adapted to the
properties of water.
Water is the most abundant substance
in living organisms.
[The Molecules of Life Lecture 2]
3
WATER AS A SOLVENT
1 Water is a polar molecule.
O is more electronegative than H, and
more attractive to the electrons of the
two bonds.
So, O and H have partial charges:
Thus, the two ‘ends’ of the molecule
have opposite partial charges,
and hence water is said to be polar.
4
2 Polarity gives rise to hydrogen
bonding.
Electrostatic interaction can
between the partial charges:
occur
5
6
The hydrogen bonding explains why
water has unexpectedly high melting
and boiling points
(compared to other common liquids.)
ICE
LIQUID
WATER
WATER
VAPOUR
Regular
lattice of
H-bonds
Many,
short-lived
H-bonds
(‘flickering
clusters’)
Few
H-bonds
considerable
energy needed
to break bonds
and progress
7
3 Hydrogen bonds are not unique to
water.
They occur between any electronegative
atom (usually O or N)
and a H atom covalently bound to
another electronegative atom;
for example,
polymers:
between
two
protein
8
[Learning Guide, p.
]
[See later in Proteins lectures.]
9
and between bases in a DNA doublehelix:
thymine
adenine
[Learning Guide, p.
]
[See later in Genetic Information lectures.]
10
4 Hydrogen bonds link groups and
molecules in precise patterns in
space.
H-bonds are strongest when the three
atoms involved lie in a straight line.
So,
11
is stronger
than
This confers very precise 3-D structures
on molecules like proteins and DNA, in
which much H-bonding occurs.
12
5 Molecules that form hydrogen bonds
(i.e. are polar) are water-soluble
(hydrophilic).
sugars
alcohols,
aldehydes,
ketones,
compounds with N-H groups.
When they dissolve,
the water-water H-bonding, and
the solute-solute H-bonding
are replaced by even more energetically
favourable
solute-water H-bonding.
13
6 O2 and CO2 are not polar.
O=O has no polarity.
In O=C=O,
the three atoms lie in a straight line,
So the two ‘ends’ of the molecule don’t
have opposite partial charges,
(unlike
).
Thus O2 and CO2 are relatively poorly
water-soluble,
and are transported in blood
bound to the water-soluble protein
haemoglobin,
and as H2CO3 (which is polar)
respectively.
14
7 Charged molecules are also watersoluble (hydrophilic).
Thus,
Na+Clsolid
in water
Na+
Cl-
This ‘charge screening’ stops Na+Cl- reforming, and hence keeps the salt in
solution.
Similar screening makes more complex
charged biomolecules water-soluble.
15
8 Non-polar
(and
uncharged)
molecules are hydrophobic.
They arrange themselves in water so as
to minimise disruption of hydrogen
bonding among surrounding water
molecules,
e.g. forming separate layers,
hydrophobic
liquid
water
or droplets:
At the molecular level,
with
H-bonded
water
16
is more
energetically
favourable than
with the
H-bonding
disrupted
17
9 Amphipathic molecules contain both
hydrophilic and hydrophobic parts.
E.g. a phospholipid:
hydrophilic ‘head’
(charged/polar)
hydrophobic ‘tail’
(non-polar, hydrocarbon
chains)
In water, disruption of water-water Hbonding is again minimised, by
formation of:
a micelle
(in cross
-section)
or
bilayer
(in cross
-section)
Such layers must have been the origin of
the first cell membranes.
[See later in Lipids lectures.]
18
Some proteins fold, so as to place
hydrophilic regions on the outside, and
hydrophobic regions on the inside:
this makes them water-soluble.
cross
-section
[See later in Proteins lectures.]
Hydrophobic lipids are transported
through watery blood plasma
as particles,
in which hydrophilic parts of proteins
and phospholipids are on the outside,
and the lipid on the inside.
cross
-section
[See later in Lipids lectures.]
19
Even the overall structure of the DNA
double-helix obeys this ‘rule’:
‘cross
-section’
hydrophilic
sugar
-phosphate
chain
(polar/charged)
hydrophobic
bases
[The Molecules of Life Lecture 2]
20
MB ChB PHASE I
WATER
LECTURE 2
AIM: To review:
the role of water as a weak acid.
21
1 Water dissociation
So far, water has been considered as an
uncharged molecule (or, at least, one
with only partial charges),
but it does dissociate slightly:
H+ + OH-
H2O
Equilibrium between these three species
is described thus:
Keq
=
[H+] [OH-]
[H2O]
(where the
concentrations
are those at
equilibrium.)
22
[H2O], the concentration of undissociated
water in pure water,
is very large, and essentially constant.
It can be moved across the equation:
Keq [H2O]
=
[H+] [OH-]
=
[H+] [OH-].
or
Kw
Kw is the ‘ion product of water’.
It has been experimentally determined,
and, (at 250C), is
1 x 10-14 (mol/L)2.
In pure water, of course,
there must be 1 H+ for every 1 OH(as dissociation produces 1 of each).
So, in pure water,
[H+] = [OH-] = 1 x 10-7 mol/L.
23
When dealing with very small numbers
like this (or very large numbers),
logarithms can be useful.
A brief guide to logarithms, for reference, is in
the Learning Guide, p.
.
24
2 The meaning of pH
The log of 1 x 10-7
(that is, the concentration in mol/L of H+
and of OH- in pure water)
is -7.
The minus sign is inconvenient (easily
mislaid), so it’s better to express the
small number as a negative log,
i.e.
= 7.
So, in pure water,
- log[H+]
=
- log[OH-]
This is still very cumbersome, so
- log[H+]
are abbreviated to
pH
respectively.
and - log[OH-]
and
pOH
= 7.
25
So, in pure water,
pH = pOH = 7.
26
3 The pH scale
So far, only pure water has been considered,
but the water dissociation equilibrium
occurs in all aqueous solutions,
in which there may be solutes that
produce or remove H+ and OH-.
In such solutions,
Kw still
= 1 x 10-14 (mol/L)2,
so, if [H+] increases,
[OH-] must correspondingly decrease,
and vice-versa,
so that [H+] [OH-] always = 1 x 10-14 (mol/L)2
and
pH + pOH always = 14.
27
We can imagine a series of aqueous
solutions.
(Learning Guide, p. .)
[H+]
[OH-]
(mol/L)
pH pOH
100
10-14
0
14
10-1
10-13
1
13
10-2
10-12
2
12
10-3
10-11
3
11
10-4
10-10
4
10
10-5
10-9
5
9
10-6
10-8
6
8
10-7
10-7
7
7
10-8
10-6
8
6
10-9
10-5
9
5
10-10
10-4
10 4
10-11
10-3
11 3
10-12
10-2
12 2
10-13
10-1
13 1
10-14
100
14 0
28
So, the pH scale is simply a convenient
way of expressing [H+] in aqueous
solutions over a very wide (10-14 to 100
mol/L) range of concentrations.
There is no (theoretical) reason why a
solution should not have a pH less than
0, or greater than 14.
(It’s just extremely unlikely in
Chemistry, and never seen in Biology).
Remember,
A change of just 1 pH unit means a 10fold change in [H+].
29
4 Buffers
Many biomolecules are ionisable
(i.e. may bear a charge).
Their charge depends on the pH of their
aqueous environment.
[See this later for amino-acids in Proteins
lectures.]
Their shape depends on their charge.
Their activity depends on their shape.
So,
for
optimal
activity,
such
biomolecules need to be at a particular,
optimal pH.
Organisms maintain a particular pH
(usually about 7) in biological aqueous
solutions using buffers.
30
A buffer consists of a mixture of a weak
acid and its conjugate base.
An acid is a molecule able to donate H+.
A base is a molecule able to accept H+.
A strong acid donates readily.
A weak acid donates reluctantly.
(Similar definitions for bases.)
Acid dissociation occurs thus:
H+ +
HA
acid
A-
conjugate
base
Equilibrium between these three species
is described thus:
Keq
=
[H+] [A-]
[HA]
(where the
concentrations
are those at
equilibrium.)
31
Keq (the dissociation constant for the
acid) is often referred to as Ka.
The stronger the acid, the larger the Ka.
Water, as we saw earlier, is a very weak
acid, with a very small Ka.
32
Buffer action depends on the occurrence
of the two equilibria we have already
seen.
They are linked, because both involve
H+.
(a)
H2O
H+ + OH-
(b)
HA
H+ +
A-
33
What happens if the biological solution
contains a molecule that generates OH(or removes H+)?
(a) would no longer be at equilibrium.
(There’s too much OH-.)
A little HA dissociates to give H+ (and A-).
(It’s the HA rather than H2O, because, although
both are at high concentration, HA is the stronger
acid.
Because [HA] and [A-] are very large, equilibrium
(b) is hardly affected.)
The H+ generated combines with the
extraneous OH- to bring (a) back to
equilibrium.
34
What happens if the solution contains a
molecule that generates H+ (or removes
OH-)?
(a) would no longer be at equilibrium.
(There’s too much H+.)
A little A- associates with the H+.
(It’s the A- rather than OH- because, although Ais a weaker base than OH-, it’s present at a much
higher concentration.
Again, because [HA] and [A-] are very large,
equilibrium (b) is hardly affected.)
Association of the A- with the extraneous
H+ brings (a) back to equilibrium.
35
A simple equation
relationship between
expresses
the
the buffering action of the weak acid/
conjugate base;
the Ka of the weak acid;
the maintained pH.
It’s just a re-working of the equilibrium
equation for dissociation of the weak acid
seen earlier:
=
[H+] [A-]
[HA]
=
Ka [HA]
[A-]
Ka
Rearranging:
[H+]
Taking negative logs:
- log[H+]
= - logKa - log [HA]
[A-]
Defining pKa as - logKa:
pH
=
pKa + log [A-]
[HA]
36
This is the
equation.
Henderson-Hasselbalch
It’s useful when considering setting up
artificial buffers.
Natural buffers occur when large,
essentially unchanging concentrations of
A- and HA occur naturally in biological
aqueous solutions.
Thus,
inside cells, a phosphate buffer is
important:
H2PO4(HA)
H+ + HPO42(A-)
and, in plasma, bicarbonate buffer
operates:
H2CO3
(HA)
H+ + HCO3(A-)
(Tutorial 1 Question 6 refers to this buffer and
to the Henderson-Hasselbalch equation.)
37
You make artificial buffers by mixing, in
aqueous solution, large, essentially
unchanging concentrations of A- and
HA, for example, acetate and acetic
acid.
Looking back at the
Hasselbalch equation,
pH
=
Henderson-
pKa + log [A-]
[HA]
because pKa is a constant,
you can set the pH by choosing the ratio
[A-]
[HA] .
In fact, all buffers work best when the
ratio is 1, i.e. when
log [A-]
= 0,
[HA]
and
pH
= pKa
that is, they work best at a pH equal to
the pKa of the buffer’s weak acid.
38
However, over a small pH range, (+/- 1
unit), as indicated above, you can set a
desired pH by choosing the [A-]/[HA]
ratio.
Approximate pKa values for the weak
acids in the buffers mentioned are:
Acetate
4.8
H2PO4-
6.9
H2CO3
6.1 (effectively)
[See later four types of acid-base imbalance:
respiratory acidosis, alkalosis;
metabolic acidosis, alkalosis,
in Systems 1 lectures.]