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