Biochemistry I
Lecture 2
August 25, 2015
Lecture 2: Molecular Forces, Water & Hydrogen Bonds
Learning Goals:
 Distance and charge dependence of electrostatic forces
 Distance dependence of van der Waals interactions.
 Predict bond and molecular polarity
 Relate electronic structure of water to solution properties
 Identify hydrogen bond donors and acceptors
 Explain thermodynamics of salts dissolving in water.
 Explain thermodynamics of polar compounds dissolving in water
 Explain thermodynamics of non-polar compounds dissolving in water (hydrophobic effect).
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2A. Molecular interactions
i) Electrostatics: The force between two charged particles is:
qq
1 2
 o  8.854 10 12 C 2 / N  m 2
4 Dr 2
0
The force depends on the distance between the two charges
and the dielectric constant (D) of the media.
A high dielectric constant, such as that found in water, is
important because the forces between charges are
attenuated, reducing charge interactions, except at short
distance.
F
1
1
Th
Biochemistry I
Lecture 2
August 25, 2015
ii) van der Waals (induced dipole-induced dipole < induced
dipole-dipole < dipole-dipole) – an electrostatic interaction
that does not involve formal charges. Charges may be
temporary (induced dipole) or permanent (dipole).
Boiling points of hydrocarbons:
isobutane:
261 K
butane:
272 K
Same number of carbons, why
the difference?
2B. Polar bonds & Molecules
A bond is considered to be polar if there is a significant difference in the
electronegativities of the participating atoms:
1
2
H
He
2.1
3
4
5
6
7
8
9 10
Li Be B
C
N
O
F Ne
1.0 1.5 2.0 2.5 3.0 3.5 4.0
(Increases across the periodic table).
The dipole moment, μ, is defined by the following equation:  
 qr .
All atoms
Polar molecule: A molecule is considered polar if it is has a permanent dipole moment associated
with it.
Example: CO2
Does it have polar bonds?
Is it a polar molecule?
2
O=C=O
Biochemistry I
Lecture 2
2C. Structure of Water
i.
Oxygen has the following electronic configuration:
1s22s22p4.
ii.
In water, the 2s and the three 2p orbitals form four sp3
hybrid orbitals.
iii.
These orbitals are tetrahedral in their orientation; however,
the ideal bond angle of 109° is distorted to 104.5° by
electron repulsion between the full orbitals.
iv.
The orbitals are populated such that two orbitals are filled
and two contain one electron each.
v.
The filled orbitals cannot form bonds and are called lone
pairs of electrons.
vi.
The half-filled orbitals participate in the formation of a
sigma bond between oxygen and hydrogen.
vii.
"Bent" water molecule generates a permanent dipole
moment, making water a polar solvent with a high dielectric
constant.
2D. Hydrogen Bonds
i) Formation of H-bonds is primarily an electrostatic attraction
between:
 Electropositive hydrogen, attached to an electronegative
atom is the hydrogen bond donor (i.e. NH)
 Electronegative hydrogen bond acceptor (e.g. the lone pairs
of oxygen in the case of water, or C=O group of an amide).
ii) Typical length: 2.7-2.9 Å between electronegative atoms.
iii) Typical angle: 180° ± 20°
iv)Typical energy: 20 kJ/mole.
Biochemical Significance of Hydrogen Bonds in Water:
i). In ice, the hydrogen bonds cause
the formation of cavities in the ice,
lowering the density of the solid.
ii) In liquid water, the hydrogen bonds
persist,
and
are
transient,
generating small short-lived (nsec)
clusters of "ice" in liquid water.
iii) Hydrogen bonds are present over a
wide temperature range.
iv) The hydrogen bonds in water allow
water to absorb heat without a large
increase in temperature, giving
water a high heat capacity.
2E. Solvation – It is all about reaching the lowest energy.
ΔH: Enthalpy – A change in the electronic configuration of the system that
either releases heat (ΔH <0) or absorbs heat (ΔH >0).
A reaction that releases heat is favorable.
ΔS: Entropy – A change in the number of configurations of the system
(disorder). Either increasing the disorder (ΔS>0) or decreasing the
disorder (ΔS<0). An increase in disorder is favorable (ΔS>0)
3
August 25, 2015
Biochemistry I
ΔG=ΔH-TΔS
Lecture 2
ΔG <0 – favorable
August 25, 2015
ΔH<0 – favorable
ΔS>0 – favorable.
i) Solvation of ions:
 Energy is required to break the ionic bonds in the
crystal. H>>0. Heat added to system, unfavorable
 A large dipole moment on water means that the
solvent molecules can interact favorably with
charged solute molecules. This is energetically
favorable. H<0, releases heat.
H > O unfavorable
Na+ ClCl- Na+
Cl-
Na+
Cl-
Na+
H < O favorable
 Overall ΔH can be positive (unfavorable) or
negative (favorable), depending on the balance of
these two terms.
Na+
Cl-
Question: Why does NH4Cl readily dissolve in water, yet
the solution becomes cold, indicating that the reaction
consumes heat, and therefore should be unfavorable.
ClNa+
Release heat
H<0
ii) Hydrophilic (polar) compounds (e.g.
methanol):
ΔH<0 – favorable
T
H3C
O
Consume heat
H>0
H
ΔS>0 – favorable
Time
iii) Hydrophobic (apolar) compounds
(e.g. methane).
"oil"
"water"
http://chem.ps.uci.edu/~kcjanda/Group/Research_hydrates.html
iv) Amphipathic (or amphiphilic) compounds are both
polar (usually charged) and have a substantial
nonpolar section (e.g. fatty acids). These can form
micelles if the nonpolar part is sufficiently large.
Micelles are aggregates of amphipathic molecules that
sequester the nonpolar part on the inside.
4
CaCl2
O
O
NH4Cl