Next Steps: Resonance. Formal Charge When atoms do not exhibit ‘normal’ bonding patterns, they...

Next Steps:Resonance

Formal ChargeWhen atoms do not exhibit

‘normal’ bonding patterns, they will contain a ‘formal charge’.

Formal Charge does not indicate an actual ionic charge – it indicates the distribution of electrons

Dimethyl Sulfoxide (DMSO)

Normally, Sulfur owns 6 valence electrons, but in this structure, it only owns 5Therefore, Sulfur has formally lost 1 electron and has a + charge

Likewise, Oxygen normally owns 6 valence electrons – in this structure it owns 7, so it has a formal - charge

Calculating Formal ChargeFC = #valence e- - [(1/2 bonded e-) +

nonbonding e-]

Easier Calculation:FC = #valence e- - bonds – dots

You Try It: Calculate any fc’s for nonhydrogen atoms

H3C-C≡N-O

Note:From now on, lone pairs or formal

charges must be shown when needed.

You may show both, but it is not necessary.

Atoms that exhibit normal bonding patterns may assumed to have a formal charge of zero

Read pages 10-19 & try problems

ResonanceThis is why we study formal

charge:Consider Nitromethane:

Nitromethane EPMExperiments show that each N-O

bond is equivalent. Examine electron distribution:

Why?The true structure is a resonance

hybrid. The electrons are distributed evenly with both oxygen atoms bearing equal negative charge.

Remember:◦Resonance structures are not real.

They only help us to envision electron distribution. Only by knowing the contributing structures can we envision the real structure.

Benzene

2 Major Rules for Resonance1. Never break a single bond

2. Never exceed an octet for 2nd row elements

For more practice see handout problems 2.2 – 2.12 pgs 26-27

Drawing Arrows to Show Movement of Electrons: Pushing Electrons

Where the electrons come from

Where the electrons are moving to

Example:

You try itDraw arrows that show how one

structure becomes the other through resonance:

O-

-

O

More problems: pg 29; 2.14 – 2.19

Patterns for Drawing Resonance Structures:1. Lone pair next to pi bond2. Lone pair next to a positive

charge3. Pi bond next to a positive charge4. Pi bond between two atom

where one of those is electronegative

5. Pi bonds going all the way around a ring

6. Pi bond next to a free radical

1. Lone Pair Next to a Pi Bond“Next to” – a lone pair is

separated from a pi bond by exactly one single bond

2. Lone pair next to + chargeRemember a + charge means

that there is less electron density than usual, so there is an empty orbital available

Example:

N C+

CH3

CH3

H

H

3. Pi Bond next to + charge

+

4. Pi Bond between two atoms where one is electronegativeAn electronegative atom can

support an additional pair of electrons and a formal negative charge

CH3 CH3

O

5. Pi bonds going all the way around a ring

PhenanthreneHow many resonance structures

for this example?

6. Pi bond next to free radicalWhat is a free radical?o radical - (free radical) a neutral

substance that contains a single, unpaired electron in one of its orbitals, denoted by a dot (·) leaving it with an odd number of electrons.

o Radicals are highly reactive and unstable

o Radicals can form from stable molecules and can also react with each other.

Showing resonance of free radicalsUse half-arrows to represent the

movement of single electrons

You try itShow all of the resonance forms

for the following structure:

A look at Pyridine

The lone pair on the nitrogen does not participate in resonance due to its position in an sp2 hybrid orbital

N

Draw All Resonance Structures for Pyridine

N

Significant Resonance StructuresNot all resonance structures are

significant. Three rules help us choose structures that are significant◦1. Minimize Charges◦2. Electronegative atoms can bear

positive charge only if they have a full octet

◦3. Avoid resonance structures in which two carbon atoms bear opposite charges

1. Minimize Charges

2. Electronegative Atoms & positive charge

3. Avoid Carbons with opposite charges