1. IB Chemistry Power Points
Topic 4
Bonding
www.pedagogics.ca
LECTURE
Shapes of Molecules

2. Much taken from
AN INTRODUCTION TO
BONDING
and
SHAPES OF MOLECULES
Great thanks to
JONATHAN HOPTON & KNOCKHARDY PUBLISHING
www.knockhardy.org.uk/sci.htm

3. VALENCE SHELL ELECTRON PAIR
REPULSION (VSPER) THEORY
“THE SHAPE ADOPTED BY A SIMPLE MOLECULE OR ION IS
THAT WHICH KEEPS REPULSIVE FORCES TO A MINIMUM”
Bonds are closer
Molecules contain covalent together so repulsive
bonds. As covalent bonds forces are greater
consist of a pair of Al
electrons, each bond will Bonds are further
repel other bonds. apart so repulsive
forces are less
Bonds will therefore push each
other as far apart as possible All bonds are
to reduce the repulsive forces. equally spaced
out as far apart
Because the repulsions are Al as possible
equal, the bonds will also be
Note – you must
equally spaced. think of spacing
in 3D not just 2D
A lone pair of electrons will
exert even greater repulsive
forces on bonds.

4. ADDING ANOTHER ATOM - ANIMATION

5. ADDING ANOTHER ATOM - ANIMATION

6. 2 regions of electron density

7. 2 Bonding Pairs- BERYLLIUM CHLORIDE
Be Cl Cl Be Cl
Beryllium - has two electrons to pair up Two covalent bonds are formed
Chlorine - needs 1 electron for „octet‟ Beryllium still has an incomplete shell
BOND PAIRS 2 180°
LONE PAIRS 0 Cl Be Cl
BOND ANGLE... 180°
Geometry ... LINEAR

8. MOLECULES WITH DOUBLE BONDS
Treat as a single region of electron density
The shape of a compound with a double bond is calculated in the same way.
A double bond repels other bonds as if it was single e.g. carbon dioxide
C O O C O
Carbon - needs four electrons to complete its shell The atoms share two electrons
Oxygen - needs two electron to complete its shell each to form two double bonds
DOUBLE BOND PAIRS 2 180°
LONE PAIRS 0 O C O
Double bonds behave exactly as single
BOND ANGLE... 180°
bonds for repulsion purposes so the
shape will be the same as a molecule with
Geometry ... LINEAR
two single bonds and no lone pairs.

9. 3 regions of electron density

10. 3 Bonding Pairs - ALUMINIUM CHLORIDE
Al Cl Cl Cl
Al
Aluminium - has three electrons to pair up
Cl
Chlorine - needs 1 electron to complete „octet‟
Three covalent bonds are formed; aluminium
still has an incomplete outer shell.
BOND PAIRS 3
LONE PAIRS 0 Cl
120°
Cl Al
BOND ANGLE... 120°
Cl
Geometry ... TRIGONAL PLANAR

11. 4 regions of electron density

12. 4 Bonding Pairs- METHANE
H
H C H
C H
H
Carbon - has four electrons to pair up Four covalent bonds are formed
Hydrogen - 1 electron to complete shell C and H now have complete shells
BOND PAIRS 4
H
LONE PAIRS 0
109.5°
C H
BOND ANGLE... 109.5°
H H
Geometry ... TETRAHEDRAL

13. 4 regions of electron density - AMMONIA
H
BOND PAIRS 3
N H H N H LONE PAIRS 1
TOTAL PAIRS 4
• The shape is based on a tetrahedron but not all the repulsions are the same
• LP-BP REPULSIONS > BP-BP REPULSIONS
• The N-H bonds are pushed closer together
• Lone pairs are not included in the shape! N
H
107°
H
H
ANGLE... 107°
SHAPE... PYRAMIDAL

14. 4 regions of electron density - WATER
H
BOND PAIRS 2
O H H O LONE PAIRS 2
TOTAL PAIRS 4
• The shape is based on a tetrahedron but not all the repulsions are the same
• LP-LP REPULSIONS > LP-BP REPULSIONS > BP-BP REPULSIONS
• The O-H bonds are pushed even closer together
• Lone pairs are not included in the shape
O
104.5°
H
O O H
H H ANGLE... 104.5°
H H
SHAPE... BENT

15. HOW TO DETERMINESHAPES OF IONS
H
BOND PAIRS 3 PYRAMIDAL
NH3 N H N
LONE PAIRS 1 H-N-H 107°
H
H
+
NH4 N
+
H N
+
H
BOND PAIRS 4 TETRAHEDRAL
LONE PAIRS 0 H-N-H 109.5°
H
H
- BOND PAIRS 2 BENT
NH2 N
H N
LONE PAIRS 2 H-N-H 104.5°

16. 5 Bonding Pairs (HL only) - PHOSPHORUS(V) FLUORIDE
F
P F F
F P
Phosphorus - has five electrons to pair up
F
Fluorine - needs one electron to complete „octet‟
F
Five covalent bonds are formed; phosphorus can
make use of d orbitals to expand its „octet‟
F
BOND PAIRS 5
90°
LONE PAIRS 0 F
120° P F
F
BOND ANGLE... 120° & 90°
F
Geometry ... TRIGONAL BIPYRAMIDAL

17. 6 Bonding Pairs (HL only) - SULPHUR(VI) FLUORIDE
F
S F
F F
S
Sulphur - has six electrons to pair up
F F
Fluorine - needs one electron to complete „octet‟
F
Six covalent bonds are formed; sulphur can
make use of d orbitals to expand its „octet‟
F
BOND PAIRS 6
90°
LONE PAIRS 0 F
F
S
BOND ANGLE... 90° F F
Geometry ... OCTAHEDRAL F

18. HL only - XENON TETRAFLUORIDE
F F BOND PAIRS 4
F Xe LONE PAIRS 2
Xe
TOTAL PAIRS 6
F F
• As the total number of electron pairs is 6, the shape is BASED on an octahedron
• There are two possible spatial arrangements for the lone pairs
• The preferred shape has the two lone pairs opposite each other
F
F
Xe
F
F F
Xe F
F F ANGLE... 90°
SHAPE ... SQUARE PLANAR

19. SUMMARY
FUNDAMENTAL SHAPES – no lone pairs
Molecules, or ions, possessing ONLY BONDING
PAIRS of electrons fit into a set of standard
shapes. All the bond pair-bond pair C
repulsions are equal.
All you need to do is to count up the number A covalent bond will repel
of bond pairs and chose one of the following another covalent bond
examples...
BOND ELECTRON BOND
PAIRS GEOMETRY ANGLE(S) EXAMPLE
2 LINEAR 180º BeCl2
3 TRIGONAL PLANAR 120º AlCl3
4 TETRAHEDRAL 109.5º CH4
HL ONLY
5 TRIGONAL BIPYRAMIDAL 90º & 120º PCl5
6 OCTAHEDRAL 90º SF6

20. Effect of Lone Pairs on Molecular Shape
If a molecule, or ion, has lone pairs on the central atom, the shapes are slightly
distorted away from the regular shapes. This is because of the extra repulsion
caused by the lone pairs.
BOND PAIR - BOND PAIR < LONE PAIR - BOND PAIR < LONE PAIR - LONE PAIR
O O O
As a result of the extra repulsion, bond angles tend to
be slightly less as the bonds are squeezed together.

21. SUMMARY - CALCULATING THE SHAPE OF IONS
The shape of an ion or molecule is determined by...
• calculating the number of electrons in the outer shell of the central species *
• pairing up electrons, making sure the outer shell maximum is not exceeded
-
• calculating the number of bond pairs and lone pairs (regions of e density)
• using ELECTRON PAIR REPULSION THEORY to calculate shape and bond
angle(s)
Note for ions
* the number of electrons depends on the charge on the ion
* if the ion is positive you remove as many electrons as there are positive
charges
* if the ion is negative you add as many electrons as there are negative charges
-
e..g. for PF6 add one electron to the outer shell of P
+
for PCl4 remove one electron from the outer shell of P

22. OTHER EXAMPLES TO TRY
2-
SO4 O
BOND PAIRS
LONE PAIRS
O S O- SHAPE
ANGLE
O-
BrF3 F
BOND PAIRS
LONE PAIRS
F Br SHAPE
ANGLE
F
BrF5 BOND PAIRS
F F LONE PAIRS
Br SHAPE
F F ANGLE
F

23. OTHER EXAMPLES TO TRY
2- O
SO4 O
BOND PAIRS 4
LONE PAIRS 0
O S O- TETRAHEDRAL S
O-
ANGLE 109.5°
O O-
O-
BrF3 F F
BOND PAIRS 3
LONE PAIRS 2
F Br ‟T‟ SHAPED F Br
ANGLE <90°
F F
BrF5
BOND PAIRS 5
F F Br
LONE PAIRS 1 F F
Br
SQUARE PYRAMID F F
F F ANGLES 90° <90° F
F

24. TEST QUESTIONS
For each of the following ions/molecules, state the number of bond pairs
state the number of lone pairs
state the bond angle(s)
state, or draw, the shape
BF3
SiCl4
+
PCl4
-
PCl6
2-
SiCl6
H2S

25. TEST QUESTIONS
For each of the following ions/molecules, state the number of bond pairs
state the number of lone pairs
state the bond angle(s)
state, or draw, the shape
BF3 3 bp 0 lp 120º trigonal planar boron pairs up all 3 electrons in
its outer shell
SiCl4 4 bp 0 lp 109.5º tetrahedral silicon pairs up all 4 electrons in
its outer shell
+
PCl4 4 bp 0 lp 109.5º tetrahedral as ion is +, remove an electron
in the outer shell then pair up
-
PCl6 6 bp 0 lp 90º octahedral as the ion is - , add one electron to
the 5 in the outer shell then pair up
2-
SiCl6 6 bp 0 lp 90º octahedral as the ion is 2-, add two electrons
to the outer shell then pair up
H2S 2 bp 2 lp 92º bent planar sulphur pairs up 2 of its 6
electrons in its outer shell -
2 lone pairs are left