Lewis Structures
Molecular Shape
VSEPR Model
(Valence Shell Electron Pair
Repulsion Theory)
PART 1: Ionic Compounds
Complete the table of Part 1 by writing:
• The Lewis dot structures for each
metallic and non-metallic atom
• The Lewis dot structures for the
common ions of the metal and nonmetal.
• The formula for the compound formed
between them.
• The Lewis structure of the compound
formed between them.
Here is an example of how the table should
be completed using barium and chlorine:
Element
Lewis Dot
Structure of
the Element
Lewis Dot
Structure of
its Ion
Ba
Ba :
Ba2+
Cl
..
: Cl .
..
.. _
: Cl :
..
Formula of
the
Compound
Lewis Dot Structure
of the Ionic
Compound
BaCl2
Ba2+
.. _
: Cl :
..
2
Ba:
Located in group 2 of the periodic table.
Therefore, the Lewis structure of the element has 2 dots.
It will lose its two valence electrons to form the +2 ion.
The Lewis structure of the ion of barium has no dots
and a +2 charge.
Cl:
Located in group 17 of the periodic table.
Therefore, the Lewis structure of the element has 7 dots.
It will gain one electron to give it a complete octet in the
valence shell to form the -1 ion. The Lewis structure of
the ion of chlorine has 8 dots and a -1 charge.
Since barium forms a +2 ion and chlorine
forms a -1 ion, they will combine in the ratio of
1 barium ion to 2 ions of chlorine: BaCl2
The Lewis structure of the ionic compound is
written by writing the Lewis structure of the
cation beside the Lewis structure of the anion.
If the formula of the compound requires more
than one of any ion, that ion is enclosed in
parentheses and the subscript is shown.
PART 2: Molecular Compounds
Complete the table of Part 2 by:
• Summing up the valence electrons in
molecular substance.
• Drawing the Lewis structure.
• Counting the total electron pairs on the
central atom.
• Counting the total non-bonding electron
pairs on the central atom.
• Writing the molecular geometry.
Here is how the table should be completed
showing an example of ClO3Total
Number
Lewis Dot
Molecule of
Structure
Valence
Electrons
ClO3-
26
.. .. .. : O – Cl – O :
.. |
..
:O:
..
Total #
of
Electron
Pairs on
the
Central
Atom
4
Total # of
Nonbonding
Electron
Pairs on the
Central
Atom
Molecular
Geometry
1
Trigonal
pyramid
A molecular model of each substance in
the table of part 2 is shown in the
following slides.
CH4
H2O
HF
H3O+
NH3
NH4+
OH-
F2
O2
N2
SO42CO2
HCN
NO3-
Molecular Shape
Overall molecular shape depends
upon bond angles.
Bond angles are arranged so that
electrons are spaced as far apart as
possible.
VSEPR Model
(Valence Shell Electron Pair Repulsion
Theory)
Electrons have like charges and repel
one another. Therefore, they space
themselves as far apart as possible.
1. Double and triple bonds are treated like
single bonds.
2. VSEPR can be applied to any resonance
structure.
Electron-Pair Geometry
¾ Orbital Geometry
¾ Geometrical arrangement of all
electron pairs about the central
atom.
2 Types of electron pairs:
1. Bonding or shared
2. Non-bonding or unshared
or lone pairs.
Molecular Geometry
¾ Geometrical arrangement of
atoms in space.
¾ Used to describe the shape of the
molecule.
Steps for Determining
Molecular Shape
1. Draw the Lewis Dot Structure.
2. Determine the orbital geometry by
counting up the total number of electron
pairs on the central atom, counting
double and triple bonds as one pair.
3. Determine the molecular geometry by
considering the number of lone pairs on
the central atom.
Linear
AB2
Trigonal Planar
AB3
Tetrahedral
AB4
Trigonal Bipyramidal
Octahedral
AB5
AB6
Bond Angles
Non-bonding electron pairs are not
attracted to the nucleus of another
atom.
Therefore, they exert a greater
repulsive force on adjacent electron
pairs and compress the bond angle.
Orbital Geometry
Molecular
Geometry
Bond Angle(s)
Linear
Linear
180o
Trigonal Planar
Trigonal Planar
120o
Bent
<120 o
Tetrahedral
109.5o
Trigonal Pyramidal
107o
Bent
105o
Trigonal Bipyramidal
90o, 120o, 180o
Seesaw
90o, <120o, 180o
T-Shaped
90o, 180o
Linear
180o
Octahedral
90o
Square Pyramid
90o
Square Planar
90o
Tetrahedral
Trigonal Bipyramidal
Octahedral
Polar Molecules
Molecules in which the centers of
positive and negative charge do not
coincide. (a dipole exists)
Dipole: Established when charges of
equal magnitude, but opposite
sign are separated by a distance.
Molecule Polarity
Depends upon:
1. Bond Polarity within the
molecule
2. Molecular Geometry
Symmetrical Geometries:
Linear
Trigonal Planar
Tetrahedral
Trigonal Bipyramid
Octahedral
Square Planar
Valence Bond Theory
When a valence orbital of 1 atom
overlaps with a valence orbital of
another and electrons are
concentrated (shared) between them
to form a covalent bond.
Formation of H2
Hybrid Orbitals
Formed when two or more atomic
orbitals combine in order to form
bonding orbitals.
The number of hybrid orbitals equals
the number of atomic orbitals
combined.
sp Hybridization
BeCl2
Be: 1s2 2s2
1s
..
..
:Cl – Be – Cl :
..
. . 1s
p orbital
2s
2p
2sp hybrid
sp hybrid
orbital
2p
p orbital
sp2 Hybridization
BF3
B: 1s2 2s22p1
1s
1s
..
:F:
.. | ..
: F – B – .F. :
..
2s
2sp2 hybrid
2p
2p
sp3 Hybridization
C: 1s2 2s22p2
CH4
H
|
H–C–H
|
H
1s
2s
2p
1s
2sp3 hybrid
NH3
H
|
:N–H
|
H
N: 1s2 2s22p3
1s
2s
1s
2sp3 hybrid
Lone pairs
2p
H2O
H
|
:O:
|
H
O: 1s2 2s22p4
1s
2s
1s
2sp3 hybrid
Lone pairs
2p
P: 1s2 2s2 2p6 3s2 3p3
PF5
3s
3p
3sp3d hybrid
F
| F
F–P
| F
F
3d
3d
sp3d Hybridization
S: 1s2 2s2 2p6 3s2 3p4
SF6
3s
3p
3sp3d2 hybrid
F
F | F
S
F | F
F
3d
3d
sp3d2 Hybridization
Sigma Bonds (σ)
ƒ First bond formed between atoms.
ƒ Head-to-head orbital overlap.
ƒ Internuclear axis passes through the
overlap region.
Pi Bonds (π)
ƒ Additional bonds formed between
atoms.
ƒ Side-to-side orbital overlap.
ƒ Overlap region is above and below
the internuclear axis.
ƒ Only forms when there are
unhybridized p orbitals on the
bonded atoms.
Ethene (ethylene)
C2H4
H
C: 1s2 2s2 2p2
H
C
H
C
H
1s
2sp2 hybrids
σ bond orbitals
2p
π bond orbital
Ethyne (acetylene)
C2H2
C: 1s2 2s2 2p2
H-C
1s
C-H
2sp hybrids
σ bond orbitals
2p
π bond orbitals