Joshua Yeo
Ong Han Wee
Danny Li
ELECTRONEGATIVITY
POLAR BONDS
MOLECULAR POLARITY
Some terms we will be explaining
Electronegativity
Polarity
Hydrogen
bonding
Electron density
Dipole moment
Etc…
ELECTRONEGATIVITY IN
COVALENT AND IONIC
BONDING
Electronegativity
Definition
 A measure of the tendency of an atom to
attract electrons towards itself.
e-
What happens when two atoms of
equal electronegativity bond together?




Equally electronegative
Same tendency to attract the bonding pair of electrons
Electrons average half way between the two atoms
A non-polar bond is formed
(To get a bond like this, A and B would usually have to be the same
atom.)
A.K.A - a "pure" covalent bond - where the electrons are shared evenly
between the two atoms.
What happens if B is slightly more
electronegative than A?


B end of the bond has more than its fair share of
electrons and so becomes slightly negative.
A end, short of electrons, becomes slightly
positive.
In the diagram, - (read as "delta") means “slightly
negative”, while + means “slightly positive”.
Polar bonds


This is described as a polar bond.
A covalent bond in which there is a separation of
charge between one end and the other
◦ One end is slightly positive and the other slightly negative.

Examples: most covalent bonds. The hydrogenchlorine bond in HCl or the hydrogen-oxygen
bonds in water.
What happens if B is a lot more
electronegative than A?
Electron pair is dragged right over to B's end of the bond.
A has lost control of its electron, and B has complete control over both
electrons.
Ions have been formed.
Electronegativity
Attraction that a bonding pair of electrons feels
for a particular nucleus depends on:
Number of
protons in
the nucleus
Distance
from the
nucleus
Amount of
screening
by inner
electrons
Pauling’s Scale
Electronegativity cannot be directly measured
and must be calculated from other atomic or
molecular properties
 Most commonly used method of calculation is
that originally proposed by Pauling
 Commonly referred to as the Pauling scale,
on a relative scale running from 0.7 to 4.0


Electronegativity in Pauling units
Pauling’s Scale
Explaining the trends
Number of protons in the nucleus
1.
◦
Proton number increases, charge increases
Distance from the nucleus
2.
◦
Equal distance since bonding electrons are all in the same
valence shell
Amount of screening by inner electrons
3.
◦
Same valence shell, equal screening effect
Explaining the trends
Number of protons in the nucleus
1.
◦
Proton number increases, charge increases
Distance from the nucleus
2.
◦
Increase since number of electron shells and quantum number
increase
Amount of screening by inner electrons
3.
◦
Increase since number of electrons in inner shells increase
ELECTRON DENSITY
Electron Density

Electron density is the measure of
the probability of an electron being
present at a specific location. (i.e. how
likely you are to find an electron at a
particular place)
Electron Density

Heisenberg
Uncertainty
Principle : you
can't know with
certainty where an
electron is and
where it's going
next
Electron Density
a 2p orbital

A region of space is
called an orbital is
where the electron
will be found 95% of
the time

Higher electron
density (where the
dots are thicker)
nearer the nucleus
POLARITY
Dipole Moment
Separation of positive and
negative charges
 Formed when the electron
density of one side of a
molecule is higher than the
other
 Due to a higher
electronegativity
 A polar bond must be
present

Polar Molecules
A molecule would be polar when:
1. It has dipoles
2. It does not have rotational symmetry /
dipoles do not cancel one another
Polar Molecules?
Is this a polar molecule?
1. It has dipoles
2. It does not
have rotational
symmetry
Is this a polar molecule?
1. It has dipoles
2. It does not
have rotational
symmetry
Is this a polar molecule?
1. It has dipoles
2. It does not
have rotational
symmetry
Polar Molecules?
Is this a polar molecule?
1. It has dipoles
2. Dipoles do
not cancel one
another
Is this a polar molecule?
Acetic Acid
1. It has dipoles
2. Dipoles do
not cancel one
another
Is this a polar molecule?
1. It has dipoles
Acetone
2. Dipoles do
not cancel one
another
Physical Properties

Cl-
Na+
◦ Non-polar solutes are soluble in non-polar solvents
(eg. Hexane)
ClNa+
Na◦+ Most organic
molecules
are relatively non-polar
◦ Polar solutes are soluble in polar solvents (eg. Water
the universal solvent)
Na+
ClCl◦- Mineral
salts and most sugars are highly polar

Cl-
Solvent
Applications
◦+ To dissolve
certain
+ materials for usage
Cl
Na
Na
◦ Liquid-liquid separation
 Purification and separation of solutes
Na+
Cl-
Na+
Cl-
Cl-
Na+
Cl-
Na+
Na+
Cl-
Na+
Cl-
Cl-
Na+
Cl-
Na+
Na+
Cl-
Na+
Cl-
Solvent
Boiling
point[7
Chemical formula
]
Dielectric
constant[8]
Density
Dipole
mome
nt
Non-polar solvents
Pentane
CH3-CH2-CH2-CH2CH3
36 °C
1.84
0.626
g/m
l
0.00 D
0.00 D
Cyclopentane
C5H10
40 °C
1.97
0.751
g/m
l
Hexane
CH3-CH2-CH2-CH2CH2-CH3
69 °C
1.88
0.655
g/m
l
0.00 D
2.02
0.779
g/m
l
0.00 D
0.00 D
Cyclohexane
C6H12
81 °C
Benzene
C6H6
80 °C
2.3
0.879
g/m
l
Toluene
C6H5-CH3
111 °C
2.38
0.867
g/m
l
0.36 D
1,4-Dioxane
/-CH2-CH2-O-CH2CH2-O-\
101 °C
2.3
1.033
g/m
l
0.45 D
Chloroform
CHCl3
61 °C
4.81
1.498
g/m
l
1.04 D
4.3
0.713
g/m
l
1.15 D
Diethyl ether
CH3CH2-O-CH2-CH3
35 °C
Chloroform
CHCl3
61 °C
4.81
Diethyl ether
CH3CH2-O-CH2-CH3
35 °C
4.3
0.713
g/m
l
1.15 D
9.1
1.3266
g/m
l
1.60 D
1.75 D
g/m
l
1.04 D
Polar aprotic solvents
Dichloromethane
(DCM)
CH2Cl2
Tetrahydrofuran
(THF)
/-CH2-CH2-O-CH2CH2-\
66 °C
7.5
0.886
g/m
l
Ethyl acetate
CH3-C(=O)-O-CH2CH3
77 °C
6.02
0.894
g/m
l
1.78 D
21
0.786
g/m
l
2.88 D
3.82 D
Acetone
CH3-C(=O)-CH3
40 °C
56 °C
Dimethylformamide
(DMF)
H-C(=O)N(CH3)2
153 °C
38
0.944
g/m
l
Acetonitrile (MeCN)
CH3-C≡N
82 °C
37.5
0.786
g/m
l
3.92 D
46.7
1.092
g/m
l
3.96 D
Dimethyl sulfoxide
(DMSO)
CH3-S(=O)-CH3
189 °C
Polar protic solvents
1.21
highly
g/m
l
These
bind
charged solutes
Formic
acid to positively
H-C(=O)OH
1.41 D
101 °C well due58to the
electronegative atom at one side of the solvent molecule (usually O)
0.810
Dimethyl sulfoxide
(DMSO)
189 °C
CH3-S(=O)-CH3
46.7
1.092
g/m
l
3.96 D
58
1.21
g/m
l
1.41 D
1.63 D
Polar protic solvents
101 °C
Formic acid
H-C(=O)OH
n-Butanol
CH3-CH2-CH2-CH2OH
118 °C
18
0.810
g/m
l
Isopropanol (IPA)
CH3-CH(-OH)-CH3
82 °C
18
0.785
g/m
l
1.66 D
20
0.803
g/m
l
1.68 D
1.69 D
n-Propanol
CH3-CH2-CH2-OH
97 °C
Ethanol
CH3-CH2-OH
79 °C
30
0.789
g/m
l
Methanol
CH3-OH
65 °C
33
0.791
g/m
l
1.70 D
1.74 D
1.85 D
Acetic acid
CH3-C(=O)OH
118 °C
6.2
1.049
g/m
l
Water
H-O-H
100 °C
80
1.000
g/m
l
These bind to negatively charged solutes well using hydrogen bonding
from the singular outward H atom(s)
Physical Properties

Hydrophilic VS Hydrophobic
◦
◦
◦
◦
◦
Hydrophilic  likes water
Hydrophobic  dislike water
Polar molecules are hydrophilic
Non-polar molecules are hydrophobic
Certain molecules have non-polar and polar ends of
the molecule, displaying both non-polar and polar
characteristics
◦ This would result in a hydrophobic end and a
hydrophilic end
Applications - Soap





Soap contains a hydrophilic head and a
hydrophobic hydrocarbon tail
Hydrophobic tail is attracted to dirt particles
or soap surfaces
Hydrophilic head is attracted to water
This forms a link between water and the dirt
molecules
When water is run through a soaped dirt layer
the soap will “pull” the dirt off the surface
Dirt
Dirt
Dirt
Common Examples

Polar
◦ Water
◦ Ammonia

Non-polar
◦ Carbon dioxide
◦ Methane
Intermolecular Bonding

Van der Waal’s Force
1. Hydrogen bonds: formed between molecules which have
a strongly electronegative atom and a hydrogen atom, with
the hydrogen gaining a partial positive charge.
2. Permanent Dipole (PD) - PD: one atom of a molecule is
distinctly more electronegative than the other. This results
in one side having a permanent partial positive charge and
the other side having a permanent negative charge
3. Induced Dipole (ID) - ID: random movement of the
electrons in the molecule. At any point in time, the electron
cloud at one part of the molecule may be more dense than
another side of the atom
Thank you