1. VSEPR Theory and Polar Molecules
Molecular Geometry
VSEPR Theory and Polar Molecules

2. Valence-Shell Electron-Pair Repulsion (VSEPR) Theory
VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.
We’ll start with a simple molecule that has NO unshared electron pairs, BeCl2:

3. VSEPR Theory
As you can see from the ball and stick model below, the distance between the electron pairs is maximized if the bonds to chlorine are on opposite sides of the beryllium atom, 180 degrees apart. This is a LINEAR molecule.
If we represent the central atom by the letter A and the atoms bonded to the central atom by the letters B, then according to VSEPR, BeCl2 is an example of an AB2 molecule which is linear.

4. VSEPR Theory
In the general formula, B can represent a single atom, a group of identical atoms, or a group of different atoms. However, different sizes of B molecules will change the bond angles.

5. VSEPR Theory What type of geometry would a BF3 molecule have?
The three B-F bonds stay farthest apart by pointing at the vertices of an equilateral triangle, giving what angle between the bonds?

6. Trigonal Planar Domain
This is called trigonal planar geometry.
The general formula for such a molecule is AB3.
The bond angle is 120 degrees.

7. VSEPR Theory
But what about molecular geometries when there are UNBONDED (unshared) pairs of electrons?
VSEPR theory states that these unbonded pairs occupy space around the central atom just like bonded pairs do.
However, observation of real atoms demonstrates that unshared electron pairs repel electrons more strongly than bonded pairs do, so molecular geometries are different when there are unbonded electron pairs.

8. Another Type of Trigonal Planar Molecule
A v-shaped or bent molecule with two atoms bonded to the central atom with one lone pair of electrons is in the Trigonal Planar Electron Domain.
The general form of this geometry is AB2E.
An example is SnCl2
PHET Molecular Shapes

9. VSEPR Theory
What would an AB4 molecule look like? An example would be methane, CH4. What would the bond angle be between H-C-H bonds?

10. Tetrahedral Domain
An AB4 molecule is a tetrahedral or tetragonal molecule. The bond angle between B-A-B bonds is degrees.
PHET simulation

11. VSEPR Theory
Let’s examine ammonia, NH3. What molecular geometry does it have? Would it be tetrahedral?

12. Tetrahedral Domain
Well, its electron pair geometry is tetrahedral, but the bond angle between the unbonded pair and the other N-H bonds is larger than degrees, making its molecular geometry a trigonal pyramid. The angle between H-N-H bonds is 107.8°.
The general formula for this molecular geometry is AB3E, where E represents the unshared electron pair.

13. What’s the difference?
Why is the model on the left “trigonal planar” while the one on the right is “trigonal pyramidal”?
There is an unbonded electron pair on the trigonal pyramidal model, decreasing the bond angle between the bonded pairs.

14. Tetrahedral Domain
What if there is more than one unshared pair of electrons? Consider this molecule:
Its electron pair geometry is tetrahedral.
Its molecular geometry is ‘bent’ or ‘angular’.
This an example of a AB2E2 molecular geometry.
The bond angle will be <109.5° (water is 104.5°).

15. VSEPR Theory So, what geometry do we get for 5 bonds?
The electron pair geometry and molecular geometry are both trigonal bipyramidal.
Bond angles are 90⁰ and 120⁰.
An example is phosphorus pentachloride.

16. Trigonal Bipyramidal Domain
This same electron geometry applies to situations with some combination of 5 bonded and lone pairs:
4 bonded pairs and 1 lone pair is called a see-saw
Bond angles are 90, 120, and 180.
An example is sulfur tetrafluoride (SF4).

17. Trigonal Bipyramidal Domain
This same electron geometry applies to situations with some combination of 5 bonded and lone pairs:
3 bonded pairs and 2 lone pairs is called a T-structure.
Bond angles are 90 and 180.
An example is ClF3

18. Trigonal Bipyramidal Domain
This same electron geometry applies to situations with some combination of 5 bonded and lone pairs:
2 bonded pairs and 3 lone pairs are a linear structure.
Bond angle is 180.
An example is XeF2

19. VSEPR Theory So, what geometry do we get for 6 bonds?
The electron pair geometry and molecular geometry are both octahedral.
All bond angles are 90⁰.
An example is sulfur hexachloride (SCl6).

20. Octahedral Domain
This same electron geometry applies to situations with some combination of 6 bonded and lone pairs:
5 bonded pairs and 1 lone pair form a square pyramid.
Bond angles are 90.
An example is ClF5

21. Octahedral Domain
This same electron geometry applies to situations with some combination of 6 bonded and lone pairs:
4 bonded pairs and 2 lone pairs have square planar geometry.
Bond angles are 90.
An example is XeF4

22. VSEPR Theory
Lastly, double and triple bonds are treated in the same way as single bonds.
Polyatomic ions are treated the same way that molecules are.

23. Lewis Structures of Polyatomic Ions
With a partner, draw the Lewis structures of these polyatomic ions on your white board, then write them on a sheet of paper and turn in:
PO43-
ClO4-
SO42-
(d) Also draw the Lewis structure for xenon tetroxide

24. Intermolecular Forces
The forces of attraction between molecules are known as intermolecular forces. While they vary in strength, they are generally weaker than the bonds that join atoms in molecules, ions in ionic compounds, or metal atoms in solid metals.
A compound’s boiling point is a good measure of the force of attraction between particles of a substance when it is a liquid. Substances with ionic and metallic bonds have much higher boiling points than substances with covalent bonds.

25. Intermolecular Forces
However, the strongest intermolecular forces exist between covalently bonded molecules that are POLAR.
Polar molecules act as tiny dipoles because of their uneven charge distribution. A dipole is created by equal but opposite charges that are separated by a short distance. The direction of a dipole is from the dipole’s positive pole to its negative pole.

26. Intermolecular Forces
The negative region in one polar molecule attracts the positive region in adjacent molecules, and these forces of attraction are called dipole-dipole forces. They are short-range forces, acting only between nearby molecules.

27. Molecular Polarity
If a molecule has more than one polar bond, you need to ‘add up’ the polarity and orientation of each bond to determine the molecule’s overall ‘dipole moment.’

28. Molecular Polarity
Of course, not all molecules are polar. Some are symmetric and their dipole moments cancel each other so that there is NO net molecular dipole.
Which molecular geometries have NO net dipole moment?

29. Molecular Polarity
Now let’s go back through your VSEPR and Molecular Geometry chart and label each type of molecule as ‘polar’ (P) or ‘non-polar’ (NP) by its geometry.
The polar geometries are:
Linear: v-shaped
Tetrahedral: trigonal pyramidal, v-shaped
Trigonal bipyramidal: see-saw, T-structure (shaped) and linear
Octahedral: Square pyramid

30. Review: Electronegativity
The tendency of an atom to attract pairs of electrons towards itself for bonding. An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus.

31. Electronegativity and Bonding
The difference in electronegativity between two elements determines the type of bonding that occurs between the elements.
If the difference is large, the bond will be IONIC.
If the difference is small, the bond will be COVALENT.

32. Electronegativity
Moving left to right across a period, electronegativity increases due to the stronger attraction for electrons as the nuclear charge increases.
Moving down a group, electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, reducing the attraction for electrons.

33. Bond Energies, Bond Lengths
Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.
Remember, the bond length of a covalent bond is the average distance between the atoms at their minimum potential energy.

34. How does bond length relate to bond strength?
Average Bond Length (pm)
Average Bond Energy (kJ/mol)
H-H 75
436
F-F
142
159
Cl-Cl
199
243
Br-Br
229
193
I-I
266
151
H-F 92
569
H-Cl
127
432
H-Br
141
366
C-H
109
418

35. Bond Energies, Bond Lengths
What is the general trend? As bond length increases, bond energy ____________.

36. Demonstrations
We need two boys to demonstrate types and lengths of bonds. Who will volunteer?
Now we need a girl to help the boys demonstrate bond strength.

37. Bond Energies and Lengths for Multiple Covalent Bonds
Aver. Bond length (pm)
Aver. Bond energy (kJ)
Aver. Bond length (pm)
C-C
154
346
C-O
143
358
C=C
134
612
C=O
120
732
C≡C
835
C≡O
113
1072
C-N
147
305
N-N
145
163
C=N
132
615
N=N
125
418
C≡N
116
887
N≡N
110
945

38. Bond Energies and Lengths for Multiple Covalent Bonds
What’s the trend? As the number of bonds increases, the bond length ____________ and the bond energy ___________.

39. Hydrogen Bonding
Some hydrogen-containing compounds, such as hydrogen fluoride (HF), water (H2O), and ammonia (NH3) have unusually high boiling points.
This is explained by a particularly strong type of dipole-dipole force. The large electronegativity differences between hydrogen and fluorine, oxygen, and nitrogen atoms makes their bonds highly polar.

40. How does hydrogen bonding affect water’s boiling point?
In these compounds, the hydrogen atom’s large relative positive charge and small size allows it to come very close to an unshared electron pair on an adjacent molecule.
The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule is known as hydrogen bonding.
How does hydrogen bonding affect water’s boiling point?

41. Resonance Structures
Following the octet rule, draw the Lewis structure for ozone, O3.
Some molecules and ions cannot be represented adequately by a single Lewis structure. Ozone, O3, is such a molecule, as it can be represented by either of these Lewis structures:

42. Resonance Structures
Scientists used to believe that ozone spent its time constantly alternating or ‘resonating’ between these two structures.
Experiments have shown, however, that the oxygen-oxygen bonds in ozone are identical, not sometimes a single and sometimes a double bond. Chemists now say that ozone has a structure that is the average of these two. Together the structures are called resonance structures or resonance hybrids.
Resonance refers to bonding in molecules that cannot be correctly represented by a single Lewis structure.

43. Resonance Structures
To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures:

44. Resonance Structures
Write Lewis structures for each of the following molecules. Remember to (1) sum up the valence electrons, (2) put in single bonds, and (3) satisfy the octet rule.
CO2
CO32-
NO2-
NO3-

45. Notice that two are polar while one is not (which one?).
Resonance Structures
CO2
Notice that two are polar while one is not (which one?).

46. Resonance Structures
CO32-

47. NO2-

48. Resonance Structures
NO3-