1. VESPR Theory

2. VSEPR Theory
VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)
A model for describing the shapes of molecules whose main postulate is that the structure around a given atom is determined by minimizing the electron pair repulsion
Therefore, the electrons and elements bonded to the central atom want to be as far apart as possible

3. VSEPR Steps Draw the Lewis structure for the molecule
Count the total number of things that are around the central atom to determine the electron pair geometry
Imagine that the lone pairs of electrons are invisible and describe the molecular shape

4. Summary VSEPR and Hybridization Table
Electron Domains
Electron Domain
Geometry
Predicted
Bond Angle(s)
Hybridization of Central Atom
Molecular Geometry
0 Lone Pair
1 Lone Pair
2 Lone Pair 2
Linear
180º sp 3
Trigonal Planar
120º
sp2
Bent 4
Tetrahedral
109.5º
sp3
Trigonal Pyramidal 5
Trigonal Bipyramidal
90º, 120º
sp3d
Seesaw
T-shaped 6
Octahedral
90º
sp3d2
Square Pyramidal
Square Planar

5. 2 Electron Pairs
If there are 2 things attached to the central atom, the shape is linear
Bond angle = 180°
Hybridization = sp

6. 3 Electron Pairs
If there are 3 electron pairs the shape will be trigonal planar
Bond angle = 120°
Hybridization =sp2

7. 3 electron pairs
Now imagine that you have 3 electron pairs, but one is just a lone pair (invisible) what would it look like then?

8. 4 electron pairs
If there are 4 electron pairs, the shape will be tetrahedral
Bond angle = 109.5°
Hybridization = sp3

9. 4 electron pairs
What if 1 of the electron pairs is a lone pair (invisible)? What would it look like then?
Trigonal Pyramidal

10. 4 electron pairs
What if there are 2 lone pairs (invisible)? What would it look like then?
bent

11. 5 electron pairs
If there are 5 electron pairs the shape will be Trigonal Bipyramidal
Bond angles = 90º & 120º
Hybridization = sp3d

12. 5 electron pairs
What is there is 1 lone pair (invisible)
Seesaw

13. 5 electron pairs
What is there are 2 lone pairs (invisible)
T-shaped

14. 6 electron pairs
If there are 6 electron pairs the shape will be octahedral
Bond angle = 90°
Hybridization = sp3d2

15. 6 electron pairs What if there is 1 lone pair (invisible)?
Square pyramidal

16. 6 electron pairs What if there are 2 lone pairs (invisible)
Square planar

17. Formal Charge Formal charges can be used in 1 of 2 ways…
Suggest where the charges are
Help select the most plausible structure from a set of resonance structures

18. 1 - Suggest where the charges are
Formal charge =

19. Example
Calculate the formal charge on each element in the carbonate ion
CO3 2-

20. Example

21. Example

22. Example
The sum of the formal charges of the individual charges equals the formal charge on the molecule or ion
The formal charge for carbonate =
= -2

23. 2 - Help select the most plausible structure from a set of resonance structures
When choosing the most likely resonance structure
Most likely – All formal charges are zero
Next likely – All formal charges add up to zero
Next likely – Formal charges add up to the lowest possible charge
Next likely – Negative charge is on most electronegative atom

24. Example
Which of the following resonance structures is most likely for CH2O and why?

25. Example

26. Another Example
Which is the most likely structure for N2O?

27. Another example

28. Polar bonds & polar molecules (Dipole or non dipole)
In order for a substance to be polar, the bonds within the molecule must carry different charges and cannot cancel out due to symmetry

29. Polar or non polar
CHF3
CO2
BCl3
CH4
H2O

30. Polar or non polar

31. Polar or non polar

32. Rule for solubility Like dissolves like Polar will dissolve in polar
Non polar will dissolve in non polar

33. Bonding
Intramolecular forces – bonding within molecule (ionic or covalent)
Intermolecular forces – bonding between molecules

34. Intermolecular Bonding
2 factors determine if a substance is a solid, liquid, or a gas:
Kinetic energy
Intermolecular forces holding the particles together

35. Intermolecular Bonding

36. Hydrogen bonding H is special when it bonds with another element.
The electron is on one side leaving an exposed nucleus
An approaching charged group can get very close the H nucleus creating a lrge electrostatic attraction
These attractions are especially large when H is bonded to a highly electronegative atom like F, O, or N

37. Hydrogen bonding These bonds are called Hydrogen bonds
They are VERY strong leading to
High boiling points
Viscous

38. Van der Waals Forces Dipole – Dipole
Remember, dipoles mean that the molecule has a partial positive & a partial negative charges at one end
This has a significant effect only when the molecules are close together
The partial positive and partial negative will attract
These attractions are called dipole dipole attractions
These come from polar molecules ONLY!!!

39. London Dispersion forces
Small electrostatic forces caused by the movement of the electron in molecules that n=have no permanent dipole
In all molecules – polar & non polar
Heavier atoms  stronger LDF because valence electrons are further apart in larger molecules & held less tightly so they can more easily form temporary dipoles

40. What type of intermolecular forces are present?
HCl HF
CaCl2
CH4 CO
NaNO3
LDF
DD, LDF
HB, LDF
Ionic

41. Which will have the … Highest boiling point… HBr, Kr, Cl2
HBr (DD, LDF)
Highest freezing point…H2O, NaCl, HF
NaCl (ionic)
Lowest freezing point…N2, CO, CO2
N2 (smallest non polar present LDF only)
Lowest boiling point…CH4, CH3CH3, CH3CH2CH3
CH4 (smallest nonpolar, LDF only)
Highest boiling point…HF, HCl, HBr
HF (Hydrogen bonding)

42. More examples At 25C ONF is a gas where H2O is a liquid. Why?
H2O forms H bonds which are stronger than the dipole dipole forces in ONF
At 25C Br2 is a liquid when Cl2 is a liquid. Why?
Both have only LDF, but since Br2 is heavier, the LDF are greater