Download Organic Reactions

An Overview of Organic Reactions
Reaction types: Classification by outcome
Most reactions produce changes in the functional group of the reactants:
1.
Addition (forward)
Gain of atoms across a bond
H
Example:
H3C
C
H
C
H
Cl
CH3
H3C
H
C
H
Cl
C
H
CH3
base, heat
2.
Elimination (reverse)
Loss of atoms across a bond
Dehydration and dehydrohalogenation are eliminations
Organic Oxidation & Reduction are variations of addition & elimination
An organic “oxidation”
HO
An organic “reduction”
H 3C
3.
CH3
C O
Na2Cr2O 7
O
NaBH4
H 3C
CH3
C CH3
H
C
H
OH
Substitution – a replacement reaction
H 3C
4.
CH3
C CH3
H
H
C
H
Br
-OH
H 3C
H
C
H
OH
Rearrangement - change in the alkyl group structure; may occur during
some substitution or elimination reactions
H3C
CH3
C C
H H
CH2
catalyst
H3C
CH3
C C
H
CH3
Mechanisms
Reactions can be classified according to the process by which bonds are made
and broken, or organic reaction mechanism. Two major types: hemolytic
(radical) and heterolytic (polar)
1. Radical reactions (Introduced in section 5.3 - more details in Chapter 10)
In homolytic bond cleavage, each atom gets one of the bonding electrons
Both are left with an odd electron:
A:B
A . + B . = “free radicals”
"Fishhook" arrow denotes movement of one eSince free radicals are unstable & reactive, further rxns occur (chain reaction)
Example from nature:
Steps in the mechanism of a radical substitution reaction:
Initiation:
Formation of the initial radical, usually by application of energy to
a compound with a weak or unstable σ bond:
hν or ∆
Example:
HO – OH
Propagation:
Ex:
Ex:
Reactions which produce more radicals, usually by substitution
Occurs when a radical collides with a stable molecule
Many different propagation steps are possible for each rxn.
Cl . +
.
CH3 +
Termination:
Cl .
2 . OH
H – CH3
Cl – Cl
HCl +
CH3Cl
.
CH3
+
Cl .
2 radicals combine to form a stable bond
+
.
CH3
CH3Cl
2. Polar Reactions (sections 5.4 – 5.6)
When bonds break unsymmetrically to leave both bonding electrons with one
atom (heterolytic cleavage) the products are charged species:
A : B
A+
+
B-
Polar reactions are much more common than radical reactions
Occur due to attraction between + and – charges of polar bonds in different
functional groups
Require the presence of a nucleophile (electron-rich species)
and an electrophile (electron-poor species)
Influences on bond polarity:
Atomic electronegativity
Solvents, acids or bases that can interact with bonded atoms
Polarizability of atoms
Flow of electrons is always from the nucleophile toward the electrophile
Arrows show the path of e-
A+ +
B: -
A–B
Typical nucleophiles:
Typical electrophiles:
Lewis bases
Electropositive atoms
Electron-rich atoms such as O, N
Electron-poor atoms
Lewis acids
Anions such as Cl-, BrH+ and other cations
Example: Identify the electrophile and the nucleophile in this reaction. Use arrows
to show the flow of electrons.
Classify each as one of the four reaction types
Follow the flow of electrons and predict the products:
Add curved arrows to indicate the flow of electrons
The Physical Chemistry of Reactions
Reaction mechanisms show the steps taking place but usually do not show us the
position of equilibrium, the rates of each step or energy changes that take place.
Thermodynamics
=
Study of the properties of a system at equilibrium
and its associated energy changes
The position of equilibrium of any reaction can be described by its Keq:
Reactants
Products
Keq = [products] / [reactants]
The associated energy change
= the Gibbs free energy change or ∆G
∆G = -RT ln Keq
For a reaction to favor product formation, the energy of the products must be
lower than the energy of the reactions (products more stable)
Favorable:
Unfavorable:
∆G is negative, energy is released, exergonic reaction
∆G is positive, energy absorbed, endergonic reaction
The free-energy change depends on two other changes: ∆G = ∆H – T∆S
A favorable reaction is one that yields products with strong bonds & less order
1. ∆S = entropy change:
degree of disorder
+ ∆S = reaction which produces greater number of particles
- ∆S = reaction which produces less particles, more order
2. ∆H = enthalpy change
“heat of reaction”
∆H can be obtained by using bond dissociation energies (D)
–Table 5.3 gives D values at standard temperature and pressure (∆Ho)
∆H = sum of D of bonds broken – sum of D of bonds formed
Kinetics: How fast does a reaction occur? Why is it fast or slow?
Any reaction can be described by a rate expression:
Ex:
A
B
rate = k [A]
Rate = Amount of product formed per unit time (measured experimentally)
Rate constant (k) depends on how easily energy barrier to rxn can be overcome
For any reaction to occur, there must be enough energy to attain the “transition
state” between reactants and products
Transition state =
∆G
Bonds partly broken and partly formed
Short-lived; not a stable species
= Energy of activation = energy difference between reactants and
Transition state
The smaller the ∆G, the faster reaction occurs
In a multi-step reaction, the slowest step (largest ∆G) limits the entire
process; often referred to as the “rate-limiting step”
Factors affecting reaction rates:
1. Temperature
2. Concentration of reactants
3. Number of collisions between reactant particles
4. Fraction of collisions with the correct orientation
5. Fraction of those collisions with enough energy to overcome activation energy!
What do “Reaction Energy Diagrams” tell us about a reaction?
H
H2C
Br
CH2
H3C
Br
CH2
1. Relative stabilities of reactants vs. products
2. Whether reaction is exergonic or endergonic
3. Whether equilibrium favors reactants
or products
4. How much activation energy is required to get rxn going
5. Structure of the transition state
6. Whether any intermediate species form
7. How many steps in the reaction
8. The relative rates of the steps in the reaction (which one is the rate-determining step?)
Problem 5.32 in McMurry
This reaction is an isomerization that occurs when the reactant, an alkene, is treated with acid:
At equilibrium, the reaction mixture contains about 30% reactant and 70% product.
• What is the approximate Keq for this reaction?
• The reaction occurs slowly at room temperature. What’s the approximate ∆G+?
• Draw an energy diagram for the reaction.