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Quick Review of Lecture 9 Feb. 4, 2019

Lewis Structures

Concepts and Vocabulary: FormalCharges, Valences, Oxidation States

Resonance Structures; Hyper andHypovalent sites

VSEPR

Lewis Structures

A bond between two atoms is formed by means of sharing of a pair ofelectrons

Each atom shares electrons with neighbors to achieve a total of eightvalence electrons

Determine connectivity of the atoms in the molecule

Sum up the total number of valence electrons in the molecule

Distribute the electrons so that each atom acquires an octet (duet for H!)in either

a) bonding pairs (denoted : or – ) shared between a pair of atoms, orb) lone pairs (denoted : ) that

belong

to a single atom (i.e., “unused”

in

making bonds and occupy more space than bonded pairs).

Examples:

HF, CF

4

, NH

3

, COCl

2

, CO, CO

2

, N

2

O, H

2

CN

2

, N

3

‐

, N

5

Resonance Structures:

A Way to Delocalize

Electrons in Valence Bond Descriptions

Resonance structures represent different Lewis structures for the same molecule

Resonance structures must have the same connectivity and the same total no. of

electrons

Resonance structures are

not isomers

, but are various limiting descriptions of the same

molecule.

If a molecule can be represented by more than one viable resonance structure, itselectronic structure should be thought of as a “blend” of the resonance structures. Such“blending” generally lowers the energy of the system.

Resonance structure may be of equal or unequal importance (“weight”).

Guiding principles:

Favor

octets;

minimize

formal charges

disfavor

formal positive

charge on more electronegative atoms

Pauling’s Electroneutrality Principle

Hypervalence



Hypervalence may be a misnomer by some definitions, but it wouldprobably be the most commonly used misnomer in chemistry



Aka “octet expansion”



Can be represented either by an increased number of atoms boundto the “hypervalent” atom or by an increased number of (multiple)bonds to the “hypervalent” atom.

Applies to 3

rd

row and higher

non‐metals bonds to highly electronegative elements.

Lecture 10 February 6, 2019

Positions of Electron Pairs in Molecules

Hybrid Orbitals

Shapes of Molecules

(Common

Stereochemistries in Main Group Compounds)

VSEPR

Symmetry Operations and Elements

Can Atomic Orbital Overlap Predict Molecular Shapes?

Consider: HF, H

2

S, PH

3

Bonds of all can be

ascribed to overlap of 3p valence orbital onF, S, or P with 1s H orbital

But: H

O, NH

, CH

Need Hybrid orbitals: sp

Makeup of hybrid orbitals

s + p

z

s + p

x

• p

y

s + p

x

• p

y

• p

z

(s + p

x

• p

y

) + ( p

z

• d

z

2

(s + p

x

• p

y

• p

z

• d

z

2

d

x

2 –y

Descriptions of geometries in mono-centric molecules:

It is where

The atoms are...

Geometries may be predicted via VSEPR

:

Valence Shell Electron Pair Repulsion



Electron pairs of bonds and lone pairs repel each other



The geometry around any atom is a consequence ofminimizing these repulsive interactions



Lone pairs are considered to be larger than bonding pairs



Multiple bonds are considered to be in the same space as,but larger than, single bonds



Several common geometries depending on the number of“occupants” around the atom in question

The “Steric number” of the molecule determines the hybrid

orbitals used to account for sigma bonds and lone pairs

Remember the following things:

https://www.youtube.com/watch?v=1ZlnzyHahvohttps://www.youtube.com/watch?v=xNYiB_2u8J

In TBP, lone pairs go inEquatorial positions ratherThan axial positions.This minimizes repulsions.

VSEPR rationalizes bond angles and geometry of molecules

For H

2

O and NH

3

, lone pairs are

in sp

3

hybrid orbitals; for H

2

S and PH

3

No hybrid orbitals needed.

For SF

4

the lone pair is more stable in the

sp

2

subset of the sp

3

d hybrid orbitals

VSEPR and the AXE Description of Electron Arrangements