Shape of molecules

1. CHAPTER 10: The Shape of
Molecules

2. ELECTRON-DOT STRUCTURES
• Bonding involves only the valence electrons
• Electron sharing can be represented by electron-
dot structures
N + 3 H N
H H
H
H O H
lone pair
bonding
pair

3. • Writing electron-dot structures
1. Select a reasonable (symmetrical) “skeleton”
(a) Central element is least electronegative
(b) Oxygen atoms do not bond to each other
2. Find total number of electrons in molecule
(a) Add one e- for each negative charge
(b) Subtract one e- for each positive charge

4. 3. Form single bonds between the atoms
- assign remaining electrons to form octets
on all atoms
4. If any electron remain, place them the central
atom as lone pairs
5. If central atom does not have a filled octet, form
multiple bonds with a neighboring atom

5. RESONANCE
• Consider the Lewis structure for CO3
2-
O
C
O
O O
C
O O O
C
O
O
2- 2- 2-
• The three individual forms are called resonance
structures
• The correct structure is a blend of all three – the
resonance hybrid

6. FORMAL CHARGES
• It may be possible to draw more than one Lewis
• The likely structure can be determined from the
concept of formal charges
• Formal charge: the charge on an atom in a
molecule or polyatomic ion
• Most favorable structure has atoms with zero or
near-zero formal charges on each atom

7. • Guidelines for assigning formal charges:
(a) In a molecule the sum of the formal charges
is zero
(b) In a polyatomic ion, the sum of the formal
charges is equal to the charge on the ion
FC = Group number – [(# of bonds) + (# of
unshared electrons)]

8. • Not all compounds obey the octet rule
e.g. BF3, BeCl2
:- these are examples of electron deficient species
e.g. NO2
:- this is an example of an odd-electron species
e.g. SF6
:- an example of a compound with the central
atom having an expanded valence shell

9. MOLECULAR SHAPES:VSEPR MODEL
• Lewis structures are not true representations of
the molecular shape
• Molecular shape can be predicted from the
valence-shell electron-pair repulsion model
• In this model
1. Count number of electron clouds around the
central atom

10. 2. Arrange the electron clouds around the central
atom so that they are as far apart as possible
• This arrangement is called the atom’s electronic
geometry
• Molecular geometry: arrangement of the bonded
atoms
• Double and triple bonds count as one

12. Two charge clouds
e.g. CO2
-: electron clouds 180° apart – give rise to
LINEAR molecules
O C O
Three charge clouds
e.g. BF3
• Most stable if electron clouds are 120° apart

13. • Give TRIGONAL PLANAR electronic geometry
F
B
F
F
S
O
O
e.g. SO2
e-
s at corners of an
equilateral triangle
BENT molecular geometry
Four charge clouds
e.g. CH4, NH3, H2O
• All have TETRAHEDRAL electronic geometry –
all angles are 109°

14. • Molecular geometry is different
N
H H
H
O
H
H
bent trigonal
pyrimidal
molecular
geometry
• Bond angles are 107° and 104.5° in ammonia and
water respectively

15. Five charge clouds
e.g. PCl5
• Have TRIGONAL BIPYRAMIDAL electronic
geometry
Cl P
Cl
Cl
Cl
Cl
equatorial
positions
axial
position
trigonal bipyramidal

16. • In general (for molecular geometry and L = lone
pair):
A
X
X
X
X
see-saw
AX4L e.g. SF4
X
A
X X
T-shaped
AX3L2 e.g . ClF3
A
X X
linear
AX2L3 e.g. I3
-

17. Six charge clouds
• OCTAHEDRAL electronic
geometry; bond angles of
90° and 180°
S
F
F F
F
F
F
octahedral
A
X
X X
X
X
square pyramidal
AX5L e.g. BrF5
A
X
X X
X
square planar
AX4L2 e.g. XeF4

21. BOND POLARITY
 Recall the nature of bond in HCl
H Cl
 
 The separation of charge in a polar covalent
bond creates an electric dipole
 A molecule typically has a number of polar
bonds
H Cl H I
EN = 0.9 EN = 0.9

22.  The arrangement of the bonds determines if
the molecule is polar
 For a molecule to be polar
1. There must be at least one polar bond or one
lone pair on the central atom
and
2. The polar bonds, if more than one, must not be
symmetrically arranged
or

23.  If there are two or more lone pairs on the
central atom, they must not be symmetrically
arranged
C O
O
no net dipole
F
C
F F
F
no net dipole
H
N
H
H net
dipole
polar
nonpolar
nonpolar