1
Summary
• Hopefully you are starting to build up a picture of molecular structure
• You should know how to draw molecules and both their Lewis
structure and simple molecular orbitals
O
H
C
≡
H
O
H
H
O
≡
H
C
H
• You
should also be aware that IR spectroscopy provides information
about functional groups in the molecule
• For
the molecule to absorb IR there must be a change in dipole
moment - so electrons cannot be evenly distributed in a bond
• Today we look at this effect...
2
Formal charges
• MO theory tells us that electrons shared out over whole molecule
• But it is useful to "localise" electrons and charge in order to understand
reactivity & properties
• Localise electrons in bonds
• Localise charge as formal charge (fc)
• NOTE: this is just "bookkeeping" charge distributed over molecule
=
fc
C
+
old group
number
–
number of
unshared
electrons
4H
carbon; fc = 4-0-½(8)=0
N
+ 3H
+ H
nitrogen; fc = 5-0-½(8)=+1
–
H
H C H
H
H
H N H
H
1/2 number
of shared
electrons
H
≡
H
C
H
H
H
≡
H
N
H
cation
H
3
Formal charges II
B
+ 3H
+ H
boron; fc = 3-0-½(8)=-1
H
H
H B H
H
≡
H
B
H
H
anion
• Atoms in neutral molecules can have formal charges
• This can frequently be used to understand their chemistry
O
+
+
O
O O
O
central oxygen; fc = 6-2-½(6)=+1
terminal oxygen; fc = 6-6-½(2)=-1
C
+ 3H
+
N
+ 2 O
nitrogen; fc = 5-0-½(8)=+1
oxygen; fc = 6-6-½(2)=-1
O
O3
ozone
≡
O
O
O
neutral
H
O
H C N
O
H
neutral
≡
≡O
O
O
atom's formal
charges
O
H3C N
O
atom's formal
charges
4
Carbanions and carbocations
• Learn to count to four!
• Carbon has four bonds in stable molecules
• If there are only three bonds present carbon has either a formal positive
charge or a formal negative charge
Carbanion
3 covalent bonds and a formal negative charge
carbon; fc = 4-2-½(6)=-1
R
C
R
≡
R
R
C
R
R
sp3
≡
R
C
R
Carbocation
3 covalent bonds and a formal positive charge
carbon; fc = 4-0-½(6)=+1
R
C
R
R
≡
R C
R
R
• Here the & are not the same as formal charge!
• Negative charge shows the two non-bonding electrons
• Positive charge shows the absence of 2 electrons
R
empty p
orbital
5
Polarisation
• IR has shown us that electrons not evenly distributed in a bond
• Different atoms have different electronegativities
• The more electronegative an atom the closer the electrons will be
σ*
energy
•
held to it
This effects the shape of the bonding & antibonding MOs
s orbital of less
electronegative
atom
σ
bond polarised
towards
electronegative
atom
s orbital of more
electronegative
atom
6
Polarisation II
•
•
•
•
These effects occur in σ-bonds
Below highlights the effect of the more common functional groups
Power of effect diminishes with distance from X or Z
NOTE: Formal charges effect bond polarisation
δ+ δ–
C X
X = Br, Cl, NO2, OH, OR,
SH, SR, NH2, NHR, NR2,
CN, CO2H, CHO, C(O)R
δ– δ+
C
Z
Z = R (alkyl or aryl),
metals (Li, Mg etc)
• If X is more electronegative than carbon
• Electrons are attracted to X
• The more electronegative X the stronger the effect
• This effects orbitals & reactivity
• If Z is less electronegative than carbon
• Electrons are attracted to C
• The more electropositive Z the stronger the effect
• This effects orbitals & reactivity
7
Polarisation III
•
•
•
•
Electrons shared evenly in alkene
Electrons not shared evenly in carbonyl group
Polarisation of bond effects reactivity
Carbonyl C=O is a stronger bond than C=C - but more reactive!
C
C
C
CCπ*
COπ*
energy
energy
p AO
carbon
CCπ
O
p AO
carbon
p AO
carbon
COπ
p AO
oxygen
C
C
C
O
8
Resonance
•
•
•
•
•
Electrons can also move in π-bonds
π-bonds can stabilise , & lone pairs
Electrons are delocalised or spread out
We draw the extremes or resonance forms - electrons "inbetween"
"Curly arrow" shows the movement of 2 electrons only
Positive charges
can be stabilised
2 p orbitals with 1
electron each
≡
R1
R3
R3
R2
R2
empty p orbital
+
R1
≡
≡
R1
R3
R2
Antibonding
Non-bonding
LUMO
Bonding
HOMO
p orbitals
May look complicated but it is
just the same as benzene!
9
Resonance II
Negative charges can be stabilised
R1
R3
R1
R2
R3
R2
≡
R1
R3
R2
MO diagram identical to previous slide except an addition 2 electrons
Neutral compounds can also have delocalised electrons
R
O
R2
R
R2
O
Delocalisation
This effect is general - any 2 p-orbitals separated by 1 bond can overlap!
O
R
O
O
R
O
≡
O
R
O
MO diagram is again identical to previous slide except it will be polarised
10
Nucleophiles
• Electron rich - form a covalent bond by donating two electrons
• Nucleophiles are either negatively charged
• Or they are neutral with a lone pair
Anions
H
O
≡
H O
≡
Br
3 lone pairs
≡
Br
4 lone pairs
lone pairs as
nucleophile
O
H
water
H
N
H
H
ammonia
Pr
C
H
1 lone pair
Lone pairs on heteroatoms
H
H
Me
S
Me
dimethyl sulfide
(DMS)
11
Nucleophiles II
High energy or reactive σ bonds are nucleophiles
δ– δ+
C Y
O
H
H
H
H
O
+
H
Et
R
R
nucleophilic
site
B
HO
+ Et MgBr
Y=Li, Mg
H
R
B
H
HO
H
H
R
π bonds can also act as nucleophiles
H
H
H
C C
H
+
C C
H
high electron density
– nucleophile
H
H
Br Br
Br
Br
H
Br
R
R = electron
donating group
+ Br Br
R
R
12
Strength of nucleophiles
Anions are normally more nucleophilic than lone pairs
HO
>
H2O
• More electronegative an atom the less nucleophilic it is
• Reason is the atom holds on to the electrons more readily
• This parallels BASICITY (roughly)
most
nucleophilic
R3C
>
RNH2
R2N
>
RO
>
ROH
>>
>>
F
least
nucleophilic
R–X
• Basicity = donation of 2 electrons to H or H
• Basicity measured in pKa (higher the value the more basic)
• Nucleophilicity = donation of 2 electrons to anything else
• Nucleophilicity similar to basicity but effected by sterics (size)
13
Electrophiles
• Electron deficient - form a covalent bond by accepting two electrons
• Nucleophiles are often positively charged
• But can be neutral
Electrophiles with an empty atomic orbital
H
H
≡
H
H
O
H
O
B
F
F
empty 2p
orbital
Cl
Cl
Al Cl
Cl
empty 3p
orbital
H
H
empty 1s
orbital
F
H
H
N
Al
Cl
Et
Cl
Et
Et
Cl
Cl Al
Cl
O
Et
N
Et
Et
H
14
Electrophiles II
Electrophiles with an empty molecular orbital
Cl
δ+ δ– Y = Cl,
C Y
Br, N, O
H
H
O
O
H
Cl
Me
filling σ* orbital
so break bond
σ*
δ– δ+
Cl H
electrophilic
site
Me
O
Me
Cl
H
O
Me
• Y must be be able to accept and stabilise electrons of σ bond
• If Y can it is a leaving group
• The more stable or electronegative Y is the better leaving group it is
δ+ δ– Y = O,
C Y
O
NR
electrophilic
site
+ Et MgBr
R
HO
R
Et
15
Strength of electrophiles
• Charged species are always more powerful electrophiles than
•
neutral compounds
The more stable the cation the less reactive the electrophile
H
H
O
C
O
C
>
O
• Electrophilic
strength of the electrophilic site of a neutral species
depends on the partial positive charge (δ+)
• H or C atoms are electrophilic if attached to electronegative atoms
• More electronegative the atom the more electrophilic H or C
δ+ δ–
>
δ+ δ–
>
δ+ δ–
H NH2
H Cl
H OH
H more electrophilic as chloride
more electronegative than nitrogen
O δ–
O δ–
>
C δ+ δ–
C δ+
R
Cl
R
R
C in acid chloride is more electrophilic
than C in ketone as both O & Cl are
electronegative