1. Valence Shell Electron Pair Repulsion Theory
Molecular Geometries

2. Molecular Shape
VSEPR theory assumes that the shape of a molecule is determined by the repulsion of electron pairs.

3. Chemical Bonds Bond Type Single Double Triple # of e’s 2 4 6 Notation— = 
Bond order 1 3
Bond strength
Increases from Single to Triple
Bond length
Decreases from Single to Triple

4. Bond Polarity and Electronegativity
Electronegativity: The ability of one atom in a molecule to attract other electrons.
Electronegativities range from 0.7 (Cs) to 4.0 (F).
Trend in Electronegativity increases…
From left to right across a period
From bottom to top in a group.

5. Electronegativities of Elements

6. Bond Polarity and Electronegativity
There are two types of covalent bonds.
Nonpolar Covalent Bonds – Equal sharing of electrons
Polar Covalent Bonds – Unequal sharing of electrons
The positive end (or pole) in a polar bond is represented + and the negative pole -.
HyperChe m

8. Polar Covalent Bond
A Polar Covalent Bond is unequal sharing of electrons between two atoms (H-Cl)
In a polar covalent bond, one atom typically has a negative charge, and the other atom has a positive charge

9. Nonpolar Covalent Bond
A Nonpolar Covalent Bond is an equal sharing of electrons between two atoms (Cl-Cl, N-N, O-O)

11. Classification of Bonds
You can determine the type of bond between two atoms by calculating the difference in electronegativity values between the elements.
The bigger the electronegativity difference the more polar the bond.
Type of Bond
Electronegativity Difference
Nonpolar Covalent
0  0.4
Polar Covalent
0.5  1.9
Ionic
2.0  4.0

12. Your Turn To Practice Practice
What type of bond is HCl? (H = 2.1, Cl = 3.1)
Difference = 3.1 – 2.1 = 1.0
Therefore it is polar covalent bond.
Your Turn To Practice
N(3.0) and H(2.1)
H(2.1) and H(2.1)
Ca(1.0) and Cl(3.0)
Al(1.5) and Cl(3.0)
Mg(1.2) and O(3.5)
H(2.1) and F(4.0)

13. How to show a bond is polar
d+ means a partially positive (less electronegative) d- means a partially negative (more electronegative) The Cl pulls harder on the electrons The electrons spend more time near the Cl d+ d- H Cl

14. Polar Molecules Molecules with a positive and a negative end
Requires two things to be true
The molecule must contain polar bonds
This can be determined from differences in electronegativity.
Asymmetric molecule.

15. Symmetrical Molecules
Because of symmetry, molecules that have polar bonds are overall a nonpolar molecule (+ and – charges cancel out or balance out) Examples: CO2 BF3 CCl4

16. Asymmetrical Molecules
If a molecule has polar bonds (and there is no symmetry to cancel out + and – charges), the molecule is polar. Examples: H2O HCl NH3

17. Dipole When there is unequal sharing of electrons a dipole exists
Dipole is a molecule that has two poles or regions with opposite charges
A dipole is represented by a dipole arrow pointing towards the more negative end

19. Is it Polar?
HF H2O NH3 CF4 CO2

20. Exceptions to the Octet Rule
Central Atoms Having Less than an Octet
Relatively rare.
Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A.
Most typical example is BF3.

21. Exceptions to the Octet Rule
Central Atoms Having More than an Octet
This is the largest class of exceptions.
Atoms from the 3rd period onwards can accommodate more than an octet.
Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.
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22. Valence Shell Electron Pair Repulsion
VSEPR Theory
Valence Shell Electron Pair Repulsion

23. VSEPR Theory Predicts the molecular shape of a bonded molecule.
Electron pairs in the valence shell of an atom tend to orient themselves so that the total energy is minimized.
Electrons around the central atom arrange themselves as far apart from each other as possible.
Unshared pairs of electrons (lone pairs) on the central atom repel the most.

24. VSEPR Theory
Because lone pairs of electrons are spread out more broadly than bond pairs, repulsions are greatest between two lone pairs, intermediate between a lone pair and a bond pair, and weakest between two bonding pairs of electrons.
Repulsive forces decrease rapidly with increasing angle between the central atom and the outside atoms.
Greatest at 90o, much weaker at 120o, and very weak at 180o.

25. 1) The central atom is called A.
2) All the outer atoms are designated with an X.
3) Any lone pair electrons are designated with an E.

26. Compounds with paired electrons around the central atom.

27. Linear 2 atoms attached to center atom 0 unshared pairs (lone pairs)
Bond angle = 180o
Type: AX2
Ex. : CO2 BeF2,

28. Trigonal Planar 3 atoms attached to center atom 0 lone pairs
Bond angle = 120o
Type: AX3
Ex. : BF3

29. Tetrahedral 4 atoms attached to center atom 0 lone pairs
Bond angle = 109.5o
Type: AX4
Ex. : CH4, CCl4

30. Compounds with unshared (lone) pairs of electrons around the central atom.

31. Bent 2 atoms attached to center atom 1 lone pair Bond angle = 119o
Type: AX2E
Ex. : NO2-

32. Bent 2 atoms attached to center atom 2 lone pairs Bond angle = 104.5o
Type: AX2E2
Ex. : H2O

33. Trigonal Pyramidal 3 atoms attached to center atom 1 lone pair
Bond angle = 107o
Type: AX3E
Ex. : NH3

35. Number of electron pairs
SHAPE
Molecule
Lewis Structure
Number of electron pairs
CH4
NH3 4
Tetrahedral
Trigonal Pyramidal 4
(3 shared
1 lone pair)

36. Number of electron pairs
Molecule
Lewis Structure
Number of electron pairs
H2O
SHAPE
Bent or V 4
(2 shared
2 lone pairs)
Linear
CO2 2

37. Number of electron pairs
SHAPE
Molecule
Lewis Structure
Number of electron pairs
BeCl2
BF3
Linear 2
Trigonal Planar 3