Lecture 2: Bonding,
structure of water
and acids/bases
Dr. Nadeem Asad
BIOL 313
Properties of water
• WATER is the solvent of choice for biological systems
• Constitutes 70-85% of cell weight, typically
• Important as a solvent and a reactant in biochemical
reactions
• Helps regulate temperature since it is able to absorb large
amounts of heat
• Helps regulate intracellular pH
• Used for transport – delivers nutrients and removes waste
from cells
• Water is a unique solvent whose properties are
extremely important to biochemistry
Water and it’s importance
Water and it’s importance
Water is a unique and universal solvent whose
properties are extremely important to
biochemistry
Nature of bonding in water
• Angular geometry of water imp for living
organisms
• Water is polar- O2 with its unshared electrons
carries a partial -ve charge e -0.66 and the H
atoms a + ve charge e + 0.33
• Electrostatic interactions b/w dipoles of H2O imp
Nature of bonding in water
• Water is a dipole due to it’s geometry and the difference in
electronegativity between hydrogen and oxygen. Oxygen is more
electronegative than hydrogen
• Oxygen is sp3 hybridised, tetrahedral electron geometry, bent geometry
Polar nature of water
• The polar nature and geometry of the water molecule allows
water molecules to form hydrogen bonds with each other and
with dissolved hydrophilic substances.
• Hydrogen bonds between water molecules= electrostatic
attraction between the oxygen atom of one water and the
hydrogen of another
• Water can also form hydrogen bonds with functional groups of
hydrophilic (polar or ionic) biomolecules and organic
compounds.
- hydrogen bond donors
- hydrogen bond acceptors
Polar nature of water
Hydrogen bonding in water
Non-Covalent Interactions
• Relatively weak and
reversible
• 1-Hydrogen Bonds: Special
dipole-dipole interaction
(electronegative atom (e.g.
O or N) interacts with H
atom that is partially
positive (i.e. attached to N,
O, F)
• Very important for protein
and DNA structure.
Non-Covalent Interactions
• 2. Van Der Waals interactions
a. Dispersion Forces (London
Forces) (induced dipoles in
non-polar molecules)
b. Dipole-dipole forces
Non-Covalent Interactions
• 3. Ionic bondingElectrostatic
interactions between
two oppositely
charged ions
Non-Covalent Interactions
• 4. Hydrophobic Interactions/Hydrophobic Effect: Relations
between water and hydrophobic molecules (low water-soluble
molecules)
• Nonpolar substances tend to aggregate in aqueous solution
and exclude water molecules. Thermodynamically unfavorable
to dissolve hydrophobic substance in water
• Entropy-driven process. ΔG = ΔH – TΔS
Bond strength
Hydrophobic effect
• The hydrophobic effect is the observed tendency of nonpolar
substances to aggregate in aqueous solution and exclude
water molecules.
• Mostly entropy driven: Water more ordered at interface.
Entropy of water molecules increases upon exclusion of nonpolar species
Hydrophobic effect
• Water molecules align themselves around non-polar molecule
and lose freedom to form hydrogen bonds
• Entropy lost in system results in thermodynamic barrier
• Multiple molecules aggregate – increase the entropy of the
system because fewer water molecules needed to surround
the aggregate than to hydrate each dispersed molecule
Hydrophobic effect
• Exclusion of nonpolar
substances from aqueous
solution = HYDROPHOBIC
EFFECT
• Important concept in
biochemistry – governs
protein folding
(hydrophobic amino acids
are on the interior of the
protein), formation of
membranes (hydrophobic
lipid tails sequestered in
bilayer)
• Water soluble compounds
are those in which the
interactions between the
solute and water are greater
than those between solute
molecules. i.e. salts,
biological molecules that
have polar or ionic groups
(e.g. glucose, ethanol)
H bonding in drug design
• 25% of drugs contain Fluorine!
Why?
Takes the place of H in a chemical
structure. Very electronegative –
thus electron withdrawing. Gives
different properties to the group.
• CF3 – Electron withdrawing, can
decrease basicity of nearby
amino groups – fewer positive
charges – can penetrate cells
better.
• C – F bond can form hydrogen
bonds and other dipole-dipole
interactions, potentially
increasing binding to target
molecule in the body. May be
more effective at lower
concentrations.
Non-Covalent Interactions in proteins
• Non-covalent interactions also
form between two biomolecules
(e.g. proteins & DNA)
• The 3-dimensional structure of
many biological molecules (eg.
proteins) and macromolecular
structures (eg. membranes, DNA)
is determined by hydrogen
bonding, hydrophobic interactions,
ionic interactions and van der
Waals interactions.
• Hydrogen bonds are weak but
their high abundance makes them
important!
Key Concepts
•The ionization of water = equilibrium
expression in which the concentration of the
parent substance = denominator
Quantities in square brackets
And dissociated products =numerator
symbolize the molar concentrations
K is the dissociation constant;
•Because the concentration of the undissociated of the indicated substances,
which in many cases are only
H2O ([H2O])
is so much larger than the concentrations of its negligibly different from their
activities
component ions, it can be
considered constant and incorporated into K to
yield an expression for the ionization of water,
What is pH?
• pH is a measure of Hydrogen ion concentration
(acidity or alkalinity of a solution)
• The first important concept is that the acidity of
a solution reflects only the free hydrogen ions,
not those still bound to anions
Blood pH
• Blood is in contact with nearly every blood cell
• Regulation of it’s pH is particularly critical
• Normally, blood pH varies within a very narrow range (7.357.45
• If blood pH varies from these limits it maybe fatal . Living
range: pH 7.0-7.6
• Acidosis-normal range-alkalosis
pH scale
Where do acids in the body come from?
• Some acids can enter the body in food
• Most are generated by:
1. Breakdown of proteins
2. Incomplete oxidation of fats or glucose
3. The loading and transport of carbon dioxide in the blood
Acids can be classified
according to their relative
strengths, that is, their
abilities to transfer a proton
to water. Table 2-4 are
known as weak acids
because they are only
partially ionized in aqueous
solution
(K 1). Many of the so-called
mineral acids, such as HClO4,
HNO3, and
HCl, are strong acids (K 1).
Since strong acids rapidly
transfer all their
protons to H2O, the
strongest acid that can stably
exist in aqueous solutions is
H3O.
Likewise, there can be no
stronger base in aqueous
solutions than OH.
Acid-base balance
• It is regulated in the body by:
1. The lungs
2. The kidneys
3. Systems in the blood known as chemical buffers
Buffering
• Buffers resist abrupt and large swings in the pH of body fluids
by:
• releasing H+ (acting as acids) when the pH begins to rise
• binding H+ (acting as bases) when the pH drops
Dissociation of acids
•
•
•
•
•
Acids are proton donors
Bases are proton acceptors
Acids that dissociate completely in solution are strong acids
Acids that dissociate incompletely are weak acids
Strong bases are more effective proton acceptors than are
weak bases
pKa
•
•
•
•
pKa=-log Ka
Ka is the dissociation constant
It is the pH at which the acid is half dissociated
There are equal amounts of undissociated acid and it’s
conjugate base
• The lower the pKa, the stronger the acid
Henderson-Hasselbalch equation
•
•
•
•
pH=pKa+log [A-]/[HA]
pH=pKa+log [conjugate base]/[acid]
Buffers are mixture of weak acids and their conjugate bases
Buffering is the ability of a solution to resist a change in pH
when acid or alkali is added
• At the pKa buffering is the best