Transcript Organic Chemistry Introduction

Organic Chemistry I
Organic Reactions
Unit 5
Dr. Ralph C. Gatrone
Department of Chemistry and Physics
Virginia State University
Fall, 2009
1
Objectives
• Classification of Organic Reactions
• Reaction Mechanisms
– Radical Reactions
– Polar Reactions
• Describing Reactions
– Thermodynamics
– Kinetics
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Classifying Organic Reactions
• Addition
• Elimination
• Substitution
• Rearrangements
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Addition Reactions
• A+B
A-B
• Hybridization change occurs
• sp2 to sp3 or sp to sp2
sp2
Br
Br2
sp3
Br
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Elimination Reactions
• A-B
A + B
• Hybridization change occurs
• sp3 to sp2 or sp2 to sp
sp3
KOH/ethanol
+ KBr + H2O
Br
sp2
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Substitution Reactions
• A-B + C
A-C + B
• Groups exchange
• No hybridization change occurs
sp2
Br2/FeBr3
Br
sp2
Carbon has four bonds, Hydrogens are understood to be present
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Rearrangements
• Relatively uncommon
• Groups migrate
• Different atom connections result
OH
O
heat
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Examples
1. NaNH2
H
CH3
2. CH3Br
acid
OH
NO2
O2N-NO2
NO2
NO2
light
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Examples
1. NaNH2
H
CH3
no hybridization change
substitution
2. CH3Br
hybridization change
sp3 to sp2
elimination
acid
OH
NO2
O2N-NO2
light
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NO2
NO2
hybridization change
sp2 to sp3
addition
no hybridization change
substitution
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Mechanism
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•
Definition:
detailed description of how a reaction takes place
Accounts for
bond breaking and forming
Order
relative rates
all reactants
all products
catalysts needed
all amounts used and produced
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Steps in Mechanisms
• Steps in a reaction sequence are classified
• A step involves either the formation or breaking
•
•
of a covalent bond
Steps can occur individually or in combination
with other steps
When several steps occur at the same time they
are said to be concerted
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Types of Steps in Reaction
Mechanisms
• Formation of a covalent bond
– Homogenic (symmetrical)
– Heterogenic (unsymmetrical)
• Breaking of a covalent bond
– Homolytic (symmetrical)
– Heterolytic (unsymmetrical)
• Oxidation of a functional group
• Reduction of a functional group
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Homogenic Bond Formation
• One electron comes from each fragment
• No electronic charges are involved
• Not common in organic chemistry
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Heterogenic Bond Formation
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•
•
•
One fragment supplies two electrons
One fragment supplies no electrons
Involves electronic charges
Common in organic chemistry
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Homolytic Breaking of Covalent Bonds
• Each product gets one electron from the bond
• Not common in organic chemistry
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Heterolytic Breaking of Covalent Bonds
• Both electrons from the bond that is broken
•
become associated with one resulting
fragment
A common pattern in reaction mechanisms
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Indicating Steps in Mechanisms
• Curved arrows indicate
•
•
breaking and forming of bonds
Arrowheads with a “half” head
(“fish-hook”) indicate homolytic
and homogenic steps
Arrowheads with a complete
head indicate heterolytic and
heterogenic steps
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Radical Reactions
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•
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Homolytic cleavage leads to a radical
Neutral, highly energetic species
React with many things to complete octet
Radicals may abstract an atom
Radicals may add to a pi bond
Radicals may react with another radical
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Radical Reactions
abstraction of an atom
.
H
addition to a pi bond
.
new sigma bond
H
.
new sigma bond
.
new sigma bond
radical adding to radical
. .
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Mechanism of Radical Reactions
• Initiation
– Radicals are produced
• Propagation
– Radicals react with substrate
– New Radicals are produced
• Termination
– Radicals react to terminate reaction
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Example of Radical Reaction
The Bromination of Methane
CH4 + Br2
No Reaction
light
CH3Br + multibrominated products
CH4 + Br2
Initiation:
light
Br
Br2
.
Propagation:
CH4 + Br
.
HBr + CH3
+ CH3.
Br2
.
CH3Br + Br
.
Termination:
Br
.
+
Br
.
Br2
CH3
.
+
Br
.
Br2
CH3 .
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CH3
.
CH3CH3
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Radical Halogenation
• Alkane + Cl2 or Br2 + light
• Not very useful
• Produces poly-halogenated products
• Mixtures result
• Illustrative or radical processes
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Polar Processes
• Electrical attraction between positive and
negative centers of the functional groups
• Review polar covalent bonding
• Review electronegativity
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Polar Bonds
• Result from unsymmetrical electron distribution
in a bond
+
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
X
-

-
+

Li
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Polar Bonds
• Interactions with acids or bases can alter bond
polarity

+

H-A
OH

+
H
O
+ H
Electrons are more strongly attracted to the O atom
Makes C more positive
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Polar Bonds
• Polarizability
• Electric field around an atom reacts to
•
changing electric environment
Large atoms respond stronger than smaller
atoms, for example I
+
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
I
-
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Polar Reactions
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•
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Unlike charges attract
Electron rich sites react with electron poor sites
Electron rich sites are: nucleophiles
Electron poor sites are: electrophiles
Nucleophiles react with electrophiles
nucleophile
Nu:
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electrophile
E+
Nu
E
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Polar Reactions: An Example
• Consider the following reaction:

H
Br
Is there an acid or base present?
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Polar Reaction
• Is there an acid or base present?

H
Br
HBr is an acid. No base present.
Do we have an electrophile and
nucleophile present?
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Polar Reaction

H
Br
• Pi electrons lie above plane of carbons – Nu
• H+ in HBr is electrophile
• Draw curved arrow Nu to E+
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Draw curved arrows
H
+
H Br
+
Br-
Intermediate is a carbocation
It is an electrophile (+ charge)
Br- is a nucleophile.
Opposite charges attract.
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Draw curved arrows
H
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+
+
Br-
H
Br
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Describing a Reaction
• Microscopic level – all reax – reversable
A
Keq =
B
B
A
If Keq > 1, formation of products is favored
If Keq < 1, formation of starting materials is favored
If Keq = 103 or more, describe reaction as at completion
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Free Energy of Reaction
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For Keq > 1
energy of products < energy of reactants
For Keq < 1
Energy of reactants < energy of products
Energy change is given by Gibbs Free Energy
• Energy is released for a favorable reaction
• Energy is absorbed for an unfavorable reaction
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Gibbs Free Energy
Keq > 1,
Go is negative - exergonic
Keq < 1,
Go is positive - endergonic
Gibbs Free Energy and the Equilibrium Constant:
ΔE = -RTlnK
ΔE = lnK
-RT
K = e-ΔE/RT
where ΔE = difference in energy of the conformers in J/mole;
R = gas constant (8.315 J/Kmole)
T = temperature in K
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Changes in Energy at
Equilibrium
• Relationship: DGº = DHº - TDSº
– where DGº change in Gibbs Free Energy
– where DHº = change in enthalpy
– heat given off (exothermic)
– bonds forming
– heat absorbed (endothermic)
– bonds breaking
– where DSº = change in entropy
– the amount of disorder in the system
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Limitations on the Keq
• Tells us the position of a reaction
• How much product is theoretically possible
• No information is available for
• How fast do we reach equilibrium?
• What is the rate of the reaction?
• How do reactions occur?
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Molecular Collisions
• Molecules collide – electrons rearrange
Molecules surrounded by electrons
As atoms approach each other – electron – electron
repulsion must occur
Energy available must overcome this repulsive term
The activation energy
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Graphical Representation
• The energy needed to go from reactant to
transition state is the activation energy
(DG‡)
• The highest energy point in a reaction step
is called the transition state
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Energy of Activation
• If (DG‡) is large = reaction is slow
• If (DG‡) is small = reaction is fast
• Adding heat raises the ground state energy of
the reactants – speeding the reaction
• Cooling the reactants lowers the GS energy of
the reactants – slowing the reaction
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The Transition State
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The Transition State
• In the addition of HBr the (conceptual)
transition-state structure for the first step
• The  bond between carbons begins to
break
• The C–H bond begins to form
• The H–Br bond begins to break
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Transition State
• From the [TS] ‡ the reaction may
– reverse
– To starting materials
– Releases E (amount equals DG‡)
• Or
– proceed forward
– To products
– Energy change = net difference between energy of the starting materials
and products (DGo)
• If DGo is + : endergonic (endothermic)
• If DGo is – : exergonic (exothermic)
– In organic chemistry we may generalize
DH = DG
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allows us to use endothermic and exothermic terms
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Intermediates
• If a reaction occurs in more than one step, it
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•
•
must involve species that are neither the
reactant nor the final product
These are called reaction intermediates or
simply “intermediates”
Each step has its own free energy of activation
The complete diagram for the reaction shows
the free energy changes associated with an
intermediate
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Carbocation Intermediate Reactions
with Anion
• Bromide ion adds an
•
•
electron pair to the
carbocation
An alkyl halide
produced
The carbocation is a
reactive intermediate
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Reaction Diagrams
• Four Types of Reaction Diagrams
– a.
– b.
– c.
– d.
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Fast, exothermic
Slow, exothermic
Fast, endothermic
Slow, endothermic
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Laboratory and Biological Reactions
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Organic solvents
Variable temperature
Catalysts are rare
Simple reagents
Non-specific
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Water
Body temperature
Enzyme catalysts
Complex reagents
Specific
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Enzymes
• Globular proteins
• Active site
• Reaction occurs at active site
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