1. Molecular Geometry Cocaine
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Review of Chemical Bonds
There are 3 forms of bonding:
1) _________—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another
2) _________—some valence electrons shared between atoms
3) _________ – holds atoms of a metal together
Most bonds are somewhere in between ionic and covalent.

3. The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

4. Electronegativity Difference
If the difference in electronegativities is between:
1.7 to 4.0: Ionic
0.3 to 1.7: Polar Covalent
0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!

5. Bonding and Lone Pairs
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl
shared or
bond pair
lone pair (LP)
This is called a LEWIS structure.

6. Bond Formation H Cl H – Cl Note that each atom has a
A bond can result from an overlap of atomic orbitals on neighboring atoms. Cl H •• • + ••
H – Cl • ••
Overlap of H (1s) and Cl (2p)
Note that each atom has a
single, unpaired electron.

7. Review of Valence Electrons
Remember valence electrons are the electrons in the OUTERMOST energy level…

8. Steps for Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H. Why?
If there is a choice, the central atom is atom of lowest affinity for electrons. (Most of the time, this is the least electronegative atom… Therefore, N is central on this one
Add up the number of valence electrons that can be used. Ammonia, NH3
H = 1 and N = 5
Total = (3 x 1) = 8 electrons ••
= 8 electrons = 4 pairs

9. H N H N 3 single bonds= 3 pairs
3. Form a single bond between the central atom and each surrounding atom (each bond takes 2 electrons!) H N
3 single bonds= 3 pairs H N
4. Remaining electrons form LONE PAIRS to complete the octet as needed (or duet in the case of H).
3 BOND PAIRS + 1 LONE PAIR = 4 Pairs
=8 electrons = 4 pairs

10. Check to make sure there are 8 electrons around each atom except H
Check to make sure there are 8 electrons around each atom except H. H should only have 2 electrons. This includes SHARED pairs. H •• N
6. Also, check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. If you have less electrons in the drawing than in step 2, you made a mistake!
= 8 electrons = 4 pairs

11. Carbon Dioxide, CO2
1. Central atom =
2. Valence electrons=
3. Form bonds.
C 4 e- O 6 e- (x2) O’s = 12 e- Total: 16 electrons = 8 pairs
2 single bonds = 2 pairs
4. Place lone pairs on outer atoms.
6 Loan pairs = 6 pairs
5. Check to see that all atoms have 8 electrons
(except for H, which can have 2.)
What is wrong with our drawing?

12. What is wrong with our drawing?
Each single bond = 1 pair (2)
Each loan pair = 1 pair (2) 8 4 8
These loan pairs should actually be bonding pairs
4 single bonds = 4 pairs
4 loan pairs = 4 pairs 8 8 8
C 4 e- O 6 e- (x2) O’s = 12 e- Total: 16 electrons = 8 pairs

13. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
C2F4
SO3

14. Violations of the Octet Rule (Honors only)
Common exceptions are: Be, B, P, S, and Xe.
Be: 4 (not 8)
B: 6
P: 8 OR 10
S: 8, 10, OR 12
Xe: 8, 10, OR 12
SF4
BF3 :
Loan Pair not shown

15. VSEPR MOLECULAR GEOMETRY Valence Shell Electron Pair Repulsion theory.
Most important factor in determining geometry is relative repulsion between electron pairs.
Molecule adopts the shape that minimizes the electron pair repulsions.
(get as far away as possible)
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16. Arrangement of the number of electron pairs around the central atom
4 charge clouds, tetrahedral
3 charge clouds, trigonal planar
2 charge clouds, linear
6 charge clouds, Octahedral
5 charge clouds, trigonal bipyramidal
electron pairs (bonding & loan pair)

17. Some Common Geometries
Linear
Trigonal Planar
Tetrahedral
(click me)
(click me)

19. VSEPR charts 1 region = electron pairs (bonding & loan pair)
Use the Lewis structure to determine the geometry of the molecule
Electron arrangement establishes the bond angles
Molecule takes the shape of that portion of the electron arrangement
Charts look at the CENTRAL atom for all data!
Think REGIONS OF ELECTRON DENSITY rather than bonds (for instance, a double bond would only be 1 region – CHARGE CLOUD)
1 region = electron pairs (bonding & loan pair)

22. VSEPR Molecular Geometry bond length and angles
determined experimentally
where Lewis Structures
 bonding geometry
VSEPR
Valence
Shell
Electron
Pair
Repulsion Theory
octahedron
90o bond angles
6 electron pairs (bonding & loan pair)
trigonal bipyramid
equatorial
120o
axial
180o
5 electron pairs (bonding & loan pair)

23. tetrahedron
109.5o
4 electron pairs (bonding & loan pair)
trigonal planar
120o
3 electron pairs (bonding & loan pair)
linear
180o
2 electron pairs (bonding & loan pair)

24. Be is an exception to the octet rule with 4 Ve-
BeCl2
valence e- =
2 +
(2 x 7)
= 16e- Cl .. Be
Be is an exception to the octet rule with 4 Ve- 8 4 8
2 valence pairs on Be =
2 bonding pairs
+ 0 lone pair
Electron pair arrangement = linear
Molecular geometry = linear

25. CO2 .. .. valence e- = 4 + (2 x 6) = 16e- C O 2 valence pairs on C =
2 bonding pairs
+ 0 lone pair C O .. 8 8 8
single and double bonds same
Electron pair arrangement =
linear
Molecular geometry =
linear

26. B = exception to the octet rule
BF3
valence e- = 3+
(3 x 7)
= 24 e- B F ..
B = exception to the octet rule
3 valence pairs on B =
3 bonding pairs
+ 0 lone pair
120o
Electron arrangement = trigonal planer
Molecular geometry = trigonal planer

27. SO2 .. : .. : .. : valence e- = 6+ (2 x 6) = 18e- S O S O S O
3 valence pairs on S =
2 bonding pairs
+ 1 lone pair
< 120o
Electron pair arrangement= trigonal planer
Molecular geometry= bent (angular)

28. CH4 valence e- = 4+ (4 x 1) = 8e- = 4 pairs C H 109.5o
4 valence pairs on C =
4 bonding pairs
+ 0 lone pair
Electron pair arrangement = tetrahedral
Molecular geometry = tetrahedral

29. NH3 : valence e- = 5+ (3 x 1) = 8e- = 4 pairs N H
4 valence pairs on N =
3 bonding pairs
+ 1 lone pair
< 109.5o
Electron pair arrangement = tetrahedral
Molecular geometry = trigonal pyramidal

30. H2O : valence e- = 6+ (2 x 1) = 8e- = 4 pairs O H < 109.5o
4 valence pairs on O =
2 bonding pairs
+ 2 lone pair
Electron pair arrangement = tetrahedral
Molecular geometry = Bent (angular)

31. P = exception to the octet rule
PCl5
valence e- = 5+
(5 x 7)
= 40e- P Cl ..
P = exception to the octet rule
90o
120o
180o
5 valence pairs on P =
5 bonding pairs
+ 0 lone pair
Electron pair arrangement = trigonal bipyramidal
Molecular geometry= trigonal bipyramidal

32. S = exception to the octet rule
SF4
valence e- = 6+
(4 x 7)
= 34e-
S = exception to the octet rule S .. F :
< 180o
5 valence pairs on P =
4 bonding pairs
+ 1 lone pair
Electron pair arrangement = trigonal bipyramidal
Molecular geometry= Seesaw (distorted tetrahedron)

33. S = exception to the octet rule
SF6
valence e- = 6+
(6 x 7)
= 48e-
S = exception to the octet rule S F ..
90o
6 valence pairs on P =
6 bonding pairs
+ 0 lone pair
Electron pair arrangement = Octahedral
Molecular geometry= Octahedral