Chemical
Bonding
• Bonding within
a molecule is
called
intramolecular
attraction
– Ionic bonds
– Covalent
bonds
– Polar covalent
bonds
Electronegativity
-
+
0 :
H
-
+
N :
H
• Electronegativity is the
tendency of an atom to
attract electrons.
• Group 1 and 2 lose
electron easily, these
elements have a low
electonegativity
• Group 17 and 16 can
attain electrons easily,
these elements have high
electronegativities
• Delta represents a partial
charge or fractional charge
Ionic Bonds
• Steals an e-, this causes
one atom to have a
positive charge and the
other a negative charge.
This is the reason for the
attraction
• Opposites attract!!
• Group 1 and 2 lose
electrons to group 16 and
17 (usually)
• When the difference in
electonegativities between
atoms is greater than 1.7
the molecules is ionic
•
•
•
•
•
•
Example NaCl
Na = .9
Cl = 3.0
3.0 - .9 = 2.1
2.1 > 1.7 = ionic bond
Cl would be the more
negative atom because it
has a higher
electronegativity (it steals
the e- from Na)
Covalent Bonds
• Atoms share electrons because
they have similar
electronegativities
• Atoms share electron to fill
octet rule
• Hydrogen form stable
molecules were it shares two
electrons, this is the duet rule
• Single bond formed when
atoms share one pair of e• Double bond formed when
atoms share two pair of e• Triple bond formed when atoms
share three pair of e• The electonegativites
differences are 0 - .3
•
•
•
•
Example O2
O = 3.5
3.5 – 3.5 = 0
0 falls between 0 and
.3 therefore O2 is a
covalent bond
More examples of
covalent bonds
Ø
Polar covalent bonds
• Electrons not always shared
equally.
• The atom with the higher
electronegativity attract the shared
electron pair more strongly
pulling it away from the other
atom.
• The shared pair is shifted from the
center between the two
participating atoms making one
end of the molecule positive and
the other end negative. The bond
is polarized. (Dipole – one side of
molecule is slightly negative and
one part slightly positive)
• The difference in electronegativity
among the atoms is .4 to 1.7
•
•
•
•
•
Example PO3
P = 2.1
O = 3.5
3.5 – 2.1 = 1.4
1.4 falls between .4 and
1.7 therefore PO3 is a
polar covalent bond
• Oxygen would be the
more negative atom
because of the
greaterelectonegativity
Now practice with worksheet #55
Lewis Structures
Represents individual
valence electrons
Represents a
nonbonding pair or
lone pair of electrons
Represents a pair of e-
B. Lewis Structures
• Electron Dot Diagrams
X
– show valence e- as dots
– distribute dots like arrows
in an orbital diagram
– Show a single line for a single bond, double line for a
double bond, triple line for a triple bond
– 4 sides = 1 s-orbital, 3 p-orbitals
– EX: oxygen
2s
2p
O
Lewis Structures
• 1) Count all valence electrons
in all the atoms in the molecule;
it doesn’t matter which atoms
they come from.
• 2) If the compound has more
than 2 atoms, the least
electronegative atom is the
central atom, often a single
atom. If carbon is present, it is
almost always the central atom.
Hydrogen is never a central
atom since it can only form one
bond.)
• 3) Make a bond between the
central atom and the other
atoms using a dash (-) to show a
pair of shared electrons.
• 4) Place the remaining valence
electrons around each of the
atoms so that they have an
octet.
Lewis structure
• When more than one
Lewis structure can be
drawn for a particular
molecule this
molecule exhibits
resonance
• Example of an electron
dot diagram for CCl4
• First step determine the
number of valence
electrons
Draw the e- dot diagram
– C has 4 (remember 2 from 2s
O
CL
O
and 2 from 2p)
– Cl has 7 (4) (remember the 7
comes from 2e- from the 3s
and 5ee from 3p) the 4 comes
from Cl4
– 4 + 7 (4) = 32 valence
electrons
O
Now count to see if you have 32 valence e-
• Your turn, draw the e- dot diagram for CH3I
• Step 1 – count valence e• Step 2 – draw e- dot diagram remember (If the compound has more
than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it
is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.)
• Step 3 count to see if you have the correct number of
Valence e-
Did you remember to:
• Check to see if all of
the valence electrons
got used?
• Check to make sure
every atom has an
octet of electrons (or 2
for Hydrogen)
• There are 14 Valence e- (4 +3 + 7)
• H only needs two e- to satisfy the duet rule
• Iodine and Carbon needs 8 e- to satisfy the octet
rule
I
H
C
H
H
How did you do? Now practice with worksheet #56
and #57
Metallic Bonding
• Metallic bonding is the
attraction between metal
atoms and a sea of
surrounding valence
electrons
• A result of this bonding is
mobile electons which
gives rise to the excellent
electrical conductivity of
most metals.
Intermolecular Forces
• Intermolecular
forces – bond that
holds molecules
together.
– Effect boiling and
freezing points
– http://www.bcpl.net/~k
drews/interactions/inter
actions.html
Intermolecular Forces
• Dipole – Dipole
– Molecules are attracted to
each other as a result of
partial charges of dipole
molecules this is called a
dipole-dipole
intermolecular force
– In general, intermolecular
forces are about 1% as
strong as intramolecular
forces.
Intermolecular Forces
• Hydrogen Bonding
– Especially strong dipoledipole are called hydrogen
bonding, example = water
– When F,O and N are
attached to hydrogen they
will form hydrogen bonds
with other molecules
– Common molecules that
form hydrogen bonds are
HF, H2O and NH3.
– Molecules with an OHbond, like alcohols also
exhibit hydrogen bonding
Hydrogen Bonding – type of
attraction that holds two water
molecules together.
• Cohesion – attractive force between
particles of the same kind
• Adhesion – attractive force between unlike
substances (meniscus)
Intermolecular Forces
• London Dispersion
Forces (LDF)
– The force that holds noble
gas atoms and nonpolar
molecules together
– No permanent dipoles
however, at any given time,
the e- can be mostly on one
side of the molecule
creating a slightly negative
charge on one side
compared to the other
– Weak and easily broken
force call LDF
– Larger the molecule the
more e- it has therefore the
stronger LDF
Bonding forces and boiling and
freezing points
• Strongest bonding force is
a hydrogen bond, followed
by a dipole-dipole force,
and then LDF.
• H-Bonds have highest
boiling, melting and
freezing points
• Dipole-dipole force have
intermediate boiling,
freezing and melting
points
• LDF have lowest boiling,
freezing and melting
points
Now practice with worksheet #58
VSEPR model
•
Valance Shell Electron
Pair Repulsion (VSEPR)
is a 3D model of a
molecule
1.
Draw the Lewis structure for the
molecule
Count the electron pairs and
arrange them in the way that
minimizes repulsion(that is, put
the pairs as far apart as possible)
Determine the positions of the
atoms from the way the electron
pairs are shared
Determine the name of the
molecular structure from the
positions of the atoms.
http://wunmr.wustl.edu/EduDev/
Vsepr/table1.html
2.
3.
4.
5.
Linear
• Number of bonding
pairs around central
atom = 2
• Number of lone pairs
around central atom =
0
• 180 bond
• Example- BeCl2, CO2
Triangle (Trigonal planar
structure)
• Number of bonding
pairs around central
atom = 3
• Number of lone pairs
around central atom =
0
• Examples- BF3, SO3
tetrahedral
• Number of bonding
pairs around central
atom = 4
• Number of lone pairs
around central atom =
0
• 109.5 bond angle
• Example- CH4
Trigonal pyramid
• Number of bonding
pairs around central
atom = 3
• Number of lone pairs
around central atom =
1
• 120 bond angle
• Example-NH3
Bent or v-shaped
• Number of bonding
pairs around central
atom = 2
• Number of lone pairs
around central atom =
1 or 2
• 106 bond angle
• Example H2O, SO2