CHM 151- Module 6

Bonding pair

Bond shared between two atoms

Lone pair

Non-bonding electrons, bond only on one atom

Diatomics

Atoms that exist as pairs

7 diatomics

1. N

2. O

3. F

4. Cl

5. Br

6. I

7. H

Extra step of drawing polyatomic ion Lewis structure

Adding brackets and charge

VSEPR

Valence shell electron pair repulsion, electrons want to be as far from each other as possible, this allow us to predict the arrangement of valence electron pairs around a central atom

Electron geometry shapes

1. Linear (2 electron groups)

2. Trigonal planar (3 electron groups)

3. Tetrahedral (4 electron groups)

4. Trigonal bipyramidal (5 electron groups)

5. Octahedral (6 electron groups)

Molecular geometry shapes

1. Linear

2. Trigonal planar

3. Bent

4. Tetrahedral

5. Trigonal pyramidal

6. Trigonal bipyramidal

7. Seesaw

8. T-shaped

9. Octahedral

10. Square pyramidal

11. Square planar

Dipole moment

An atom’s ability to attract electrons

How to determine polarity

1. Calculate electronegativity difference

   1. 0-0.4 = nonpolair

   2. 0.4-2.0 = polar

   3. 2.0-3.3 = ionic bond

2. Evaluate lone pairs on central atom

   1. If yes, molecule is polar

   2. If no, molecule is not polar

Relationship between bond length (pm) and bond energy (kJ)

Inverse relationship, the longer the bond length, the lower the bond energy

Exceptions to the octet rule

1. Electron deficit molecules

2. Odd valence electron molecules (free radicals)

3. Expanded octet (any element in the 3rd energy level or higher)

Resonance

The ability to move around double bonds, when a compound has multiple Lewis structures.

How does resonance impact stabilization of bonds?

It increases stabilization

Formal charge

Theoretical perspective on which resonant structure would be the most stable

Calculating formal charge

1. Determine number of valence electrons

2. Count number of lone pairs

3. Count number of bonded electrons, divided by 2

4. Subtract sum of steps 2 and 3 from value in step 1