Transcript Ch. 13 - Molecular Structure

Molecular Geometry
(p. 183 – 187)
Ch. 6 – Molecular Structure
Molecular Geometry
Chemical Formulas and Lewis Structures
reveal very little information about the
shapes of molecules
Molecular geometry refers to the three
dimensional arrangement of a molecule’s
atoms in space
There are two theories – VSEPR Theory
and hybridization
VSEPR Theory
Valence Shell Electron Pair Repulsion
Theory
States that the repulsion between the sets
of valence electrons surrounding an atom
causes these sets to be oriented as far
apart as possible, resulting in various bond
angles and shapes.
VSEPR Theory
Lone pairs repel electrons more strongly
than bonded pairs. Lone pairs reduce
the bond angle between atoms.
Bond Angle
Predicting Molecular Geometry
Use the letter ‘A’ to represent the central
atom
Use the letter ‘B” to represent atoms bonded
to the central atom. Count them.
Use the letter ‘E’ to represent lone pairs on
the central atom. Count them.
Draw the Lewis Structure
Write the molecular geometry formula based
on the Lewis structure to determine the
shape
Common Molecular Shapes
AB2
180°
LINEAR
BeF2
Be is an exception to the octet rule and needs
only two pairs of electrons
Common Molecular Shapes
AB3
BF3
120°
TRIGONAL PLANAR
Common Molecular Shapes
AB4
CH4
109.5°
TETRAHEDRAL
Common Molecular Shapes
AB5
PCl5
TRIGONAL
BIPYRAMIDAL
120°/90°
Common Molecular Shapes
90°
AB6
SF6
OCTAHEDRAL
Common Molecular Shapes
AB2E
SO2
<120°
BENT
Common Molecular Shapes
AB3E
NH3
107°
TRIGONAL PYRAMIDAL
Common Molecular Shapes
AB2E2
H2O
105°
BENT
Examples
CO2
AB2
O C O
LINEAR
180°
Examples
PF3
AB3E
F P F
F
TRIGONAL
PYRAMIDAL
107°
Hybridization
The hybridization model is the mixing of
two or more atomic orbitals of similar
energies on the same atom to produce
new orbitals of equal energy
It is used to explain the geometry of
molecular orbitals.
sp hybridization
1 s and 1 p orbitals mix
to form 2 hybrid
orbitals
end up with two lobes
180º apart.
linear
sp2 hybridization
1 s and 2 p orbitals mix to form 3
hybrid orbitals
trigonal planar
120º
3
Sp
hybridization
1 s and 3 p orbitals
mix to form 4 sp3
orbitals.
This leads to
tetrahedral shape.
(Gives us trigonal
pyramidal and bent
shapes also.)
109.5º
Intermolecular Forces
What holds molecules to each
other
I
II
III
Intermolecular Forces
They are what make solid and liquid
molecular compounds possible.
The weakest are called van der Waal’s
forces - there are two kinds
– Dispersion forces
– Dipole Interactions
Dispersion Force
Temporary Instantaneous dipoles induce
dipoles in adjacent molecules
Depends only on the number of electrons in the
molecule
Bigger molecules more electrons
More electrons stronger forces
F2 is a gas
Br2 is a liquid
I2 is a solid
Dispersion force
d+
d-
d+
d-
H
H
H
H
Dipole interactions
Occur when polar molecules are attracted
to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely
hooked like in ionic solids.
Dipole interactions
Occur when polar molecules are attracted
to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely
hooked like in ionic solids.
+
d
d
H F
+
d
d
H F
+
-
+
+
-
-
Hydrogen bonding
Are the attractive force caused by
hydrogen bonded to F, O, or N.
F, O, and N are very electronegative so it
is a very strong dipole.
They are small, so molecules can get
close together
The hydrogen partially share with the lone
pair in the molecule next to it.
The strongest of the intermolecular forces.
Hydrogen Bonding
d+ dH O
+
Hd
Hydrogen bonding
H O
H