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Molecular Geometry and
Polarity
Part A: Chemical Bonding
Review
Dr. Chin Chu
River Dell Regional High School
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Nature of Chemical Bond
• Atoms are held together by electrostatic attraction
between positively charged nuclei and negatively
charged electron clouds.
• Chemical Bond: a link between atoms that result from
mutual attraction of their nuclei for electrons.
• Bond energy: the energy required to break a bond.
• Forces in substances:
– Attractive: between electron clouds and respective bonding
nuclei (of the two atoms that bond).
– Repulsive: between all the electron clouds in the bonding
atoms; between positively charged nucleus.
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Formation of Bonds
• Bonding involves only the valence electrons (those in the
highest energy level).
• Use the periodic chart to guide determination of valence
electrons
• WHEN BONDING OCCURS:
– Atoms attain an OCTET: a stable Noble Gas configuration.
– the resulting system is at the lowest possible potential energy
level.
– The process of bonding is, therefore, exothermic: energy is being
released. If the energy released is Large we get a strong bond;
small ΔE bond is weak
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Types of Bonds
• Ionic bond: formed by transfer of
electrons from the valence energy
level of one atom to another’s
• Covalent bond: formed when atoms
share electrons.
• Metallic bond: ions of metals are
surrounded by sea of electrons that
bind all ions together.
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Lewis Theory: An Overview
• Valence e- play a
fundamental role in
chemical bonding.
• e- transfer leads to ionic
bonds.
• Sharing of e- leads to
covalent bonds.
• e- are transferred or shared
to give each atom a noble
gas configuration
– the octet.
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Lewis Symbols (Structures)
• A chemical symbol represents the
nucleus and the core e-.
• Dots around the symbol represent
valence e-.
Si
•• Al
•
• As •
•
P•
•
••
• Se
•
•
••
••
• Bi •
•
• Sb •
•
••
I
••
•
••
Ar
••
•
••
••
•N•
•
••
••
••
••
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Metallic Bond
• Metals consist of crystalline lattice in which
positive ions (kernels) are arranged in
fixed patterns.
• The valence electrons are free to move
and they belong to the entire crystal.
• “Electron Sea” model
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Metallic Bond
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Ionic and Molecular Bonds
• Formation of sodium chloride (ionic):

 Na+ [ Cl

]


Cl

Na  +



• Formation of hydrogen chloride (covalent):


H Cl



Cl

H +



A metal and a nonmetal transfer electrons to
form an ionic bond. Two nonmetals share
electrons to form a covalent bond.
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We know a
COVALENT bond
comes from sharing the
bonding pair of electrons.
FF
Shared pair
(bonding pair)
FF
The nucleus of each atom
pulls on the bonding pair.
FF
Both atoms have equal pull,
so the bonding pair is shared equally.
H Cl
If two different atoms share a bond,
one will pull more strongly
on the bonding electrons.
H Cl
H Cl
H Cl
H Cl
H Cl
H Cl
H Cl
H Cl
H Cl
The bonding electrons carry negative charge.
H Cl
The closer they get to the chlorine atom,
the more negative it gets.
The farther they get from the hydrogen,
the more positive it gets.
+
_
H Cl
But the charge is only partial.
Hydrogen has not lost the electrons
as in the formation of an ion.
H Cl
There is an unequal sharing of electrons.
d+
d–
H Cl
The partial charge is denoted by a + or –
and the Greek letter delta, d
d+
d–
The partial charge is denoted by a + or –
and the Greek letter delta, d
The degree of sharing (equal to unequal)
is determined by the electronegativity
difference between the two atoms.
FF
Two atoms of equal electronegativity
will share the bond equally
2.1
H Cl
3.0
Two atoms with a small difference in
electronegativity will share unequally,
resulting in partial charge.
d+
d–
H Cl
Two atoms with a small difference in
electronegativity will share unequally,
resulting in partial charge.
d+
d–
H Cl
This is a polar bond:
The bonding pair is, on average,
closer to one atom.
d+
d–
H Cl
Is a polar bond a covalent bond?
0.8
K
F
4.0
Two atoms with a large difference in
electronegativity will result in a loss
of an electron,
resulting in a full charge.
K
F
K
F
K
F
K
F
K
F
K
F
K
+
positive ion
_
F
negative ion
Electronegativity
• A measure of how strongly the atoms
attract electrons in a bond.
• The bigger the electronegativity difference
the more polar the bond.
• 0.0 - 0.3 Covalent nonpolar
• 0.3 - 1.0 Covalent moderately polar
• 1.0 -1.7 Covalent polar
• >1.7 Ionic
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Covalent:
DC = 0
Polar:
0.3 < DC
< 1.7
Table 8-1
Representative Electronegativity
Differences
Ionic
DC > 1.7
Pg 335
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