Lecture 6: Stereochemistry

Mechanism of organic reactions:

Electronic effects
Dr Mohammed Jasim Mohammed Al Yasiri

1
Mechanism of organic reactions: Electronic effects
Electronic Effects
Partial Polarity
Distribution of electron density
Examples of Electronic Effects
◦ Inductive effect
◦ Types of Inductive Effect
◦ Mesomeric or Resonance Effects
◦ Rules of Resonance
◦ Hyperconjugation

Dr Mohammed Jasim Al Yasiri 2

 Electronic Effects
Electronic factors that influence organic reactions include the inductive effect, electromeric
effect, resonance effects, and hyperconjugation. These electronic factors involve organic
molecules, most of which are made from a combination of the following six elements: carbon,
hydrogen, nitrogen, oxygen, phosphorus, and sulfur (known collectively as CHNOPS). Yet, the
limited number of building blocks does not prevent organic compounds from taking on diverse
properties in their physical characteristics and chemical reactivity. The subtle differentiation of
various compounds in organic chemistry is essential for the biological functions of the molecules
and creates a wide variety of reactions.

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 Electronic Effects

Part of this variety in organic chemistry stems from differences in electron behaviour when
elements other than carbon and hydrogen participate in molecular bonds. For example, the
three compounds pictured above have similar formula units and structures, but react very
differently from one another because of these electronic factors. Varying electronegativity can
cause delocalization effects, where the electron cloud for a given bond expands to more than
two atoms within the molecule.

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 Partial Polarity
Partial polarity within a molecule leads to electron transfer among the atoms in a molecule,
leading to different behaviour than what would be expected in a non-polar version of the
compound, where no sections were electron-rich or electron-deficient.
Saturated hydrocarbons are nonreactive because there is no polarity in C-C bond and practically
no polarity in C-H bonds. Carbon and hydrogen are almost identical in electronegativity, so the
electrons involved in a bond between the two atoms are equally attracted to each nucleus and
spend roughly the same amount of time orbiting one as the other.
Electron density is evenly distributed between the two atoms in a non-polar bond, which
prevents charged species from attacking or altering the bond. In contrast, charged species
(electrophiles and nucleophiles) react with polar organic molecules because there is an
imbalance in electron density or polarity.

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 H H H C H
H H
H
RACTICE? Try Problem 1.41 Acetylcholine

1.5 Induction and Polar Covalent Bonds

Distribution of electron density

Chemists classify bonds into three categories: (1) covalent, (2) polar covalent, and (3) ionic. T ese
categories emerge from the electronegativity values of the atoms sharing a bond. Electronegativity is
a measure of the ability of an atom to attract electrons. Table 1.1 gives the electronegativity values for
elements commonly encounter ed in organic chemistr y.
Electronegativity (Pauling, 1932) – the tendency of an atom to attract electron density
TABLE 1.1 ELECTRONEGATIVITY VALUES OF SOME COMMON ELEMENTS

Increasing electronegativity

H
2.1
Li Be B C N O F
1.0 1.5 2.0 2.5 3.0 3.5 4.0
Increasing
Na Mg Al Si P S Cl electronegativity
0.9 1.2 1.5 1.8 2.1 2.5 3.0
K Br
0.8 2.8

When two atoms form a bond, one critical consideration allows us to classify the bond:
What is the dif erence in the electronegativity values of the two atoms? Below are some rough
guidelines:
If the dif erence in electronegativity is less than 0.5, the electrons are considered to be
equally shared between the two atoms, resulting in a covalent bond. Examples include C− C and
C− H :

C C C H

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 Examples of Electronic Effects
Examples of electronic effects: 1) The Inductive Effect 2) Resonance 3) The Mesomeric Effect 4)
Electromeric Effect 5) Hyperconjugation

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 Inductive Effect
Inductive effect may be defined as the permanent displacement of electrons forming a covalent
bond towards the more electronegative element or group.
The inductive effect is represented by the symbol, arrow pointing towards the more
electronegative element or group of elements. In case of 1-Chloro Butane inductive effect may be
represented as below.

The C-Cl bond is a Polar Covalent Bond. This polarization results is partial positive charge (d+) on
Carbon-1 and Partial Negative Charge (d-) on Chlorine. The shift of electron density is shown by
an arrow that points from d+ to d- end of the polar bond

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 Inductive Effect
In turn Carbon-1 draws some Positive charge (d+) towards it from the adjacent C-C bond . In this
way the polar C-Cl bond induces polarity in the adjacent bond. Such polarization of s-bond
caused by the polarisation of the adjacent s-bond is called Inductive Effect. The effect rapidly
dies out and is usually not significant after the 2nd carbon atom, or at most the 3rd.

The inductive effect is permanent, but relatively weak, and can be easily overshadowed by the
electronic effects discussed later.

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 Types of Inductive Effect
There are two categories of inductive effects: the electron-withdrawing (-I effect) and the electron-releasing (+I
effect). The latter is also called the electron-donating effect. In the image above, X is electron-withdrawing
and Y is electron-donating. These relative inductive effects are measured with reference to hydrogen:

NO2 > COOH > F > Cl > Br > I > OR > OH > C6H5 (benzene) > H > Me3C- > Me2CH- > MeCH2- > CH3-

-I effect:
The -I effect is seen around a more electronegative atom or group, and electron density is higher there than
elsewhere in the molecule. Electron-withdrawing groups include halogen, nitro cyano carboxy ester and aryloxy

+I effect:
The +I effect is observed among the less electronegative atoms of the molecule by electron-releasing (or
electron-donating) groups. The alkyl groups are usually considered electron-releasing (or electron-donating)
groups.

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 Mesomeric or Resonance Effects
Sometimes, there are several correct Lewis structures for a given molecule. Ozone is one
example. The compound is a chain of three oxygen atoms, and minimizing the charges while
giving each atom an octet of electrons requires that the central oxygen atom form a single bond
with one terminal oxygen and a double bond with the other terminal oxygen.
When drawing the Lewis structure, the choice of placement for the double bond is arbitrary, and
either choice is equally correct. The multiple correct ways of drawing the Lewis structure are
called the resonance forms.

The resonance hybrid for ozone is found by identifying the multiple resonance structures for the molecule.

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 Mesomeric or Resonance Effects
Based on the resonance forms, a beginning chemistry student might wonder if ozone has bonds
of two different lengths, since single bonds are generally longer than double bonds. However,
the ozone molecule is perfectly symmetrical, with bonds that are the same length. None of the
resonance forms represent the true structure of the molecule. Rather, the negative charge of
the electrons that would form a double bond are delocalized, or distributed evenly across the
three oxygen atoms. The true structure is a composite, with bonds shorter than what would be
expected for single bonds, but longer than the expected double bonds.
Thus, for the two structures (I and II) shown above constitute the canonical structures or
resonance structures and their hybrid (i.e. the III structure) represents the structure of more
accurately. Resonance is represented by a double-headed arrow between the resonance
structures, as illustrated above.

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 Mesomeric or Resonance Effects
The resonance hybrid is more stable than its canonical Examples
forms, i.e. the actual compound (hybrid) is at a lower
energy state than its canonical forms. Resonance stability
increases with increased number of resonance structures.

The difference in the experimental and calculated

energies is the amount of energy by which the compound
is stable. This difference is known as resonance energy or
delocalization energy.

All resonance structures are not equivalent. The following

rules help determine whether or not a resonance structure
will contribute significantly to the hybrid structure.

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 Rules of Resonance
All resonance structures are not equivalent. The following rules help determine whether or not a
resonance structure will contribute significantly to the hybrid structure.
Rule 1: The most significant resonance contributor has the greatest number of full octets (or if
applicable, expanded octets).

Open octet on carbon: All atoms have full octets

Less significant resonance contributor More significant resonance contributor
Rule 2: The most significant resonance contributor has the least number of atoms with formal charges.

Two formal charges No formal charges

Less significant resonance contributor More significant resonance contributor

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 Rules of Resonance
Rule 3: If formal charges cannot be avoided, the most significant resonance contributor has the
negative formal charges on the most electronegative atoms, and the positive formal charges on the
least electronegative atoms.

Negative formal charge on carbon (EN = 2.5) Negative formal charge on oxygen (EN = 3.5)
Less significant resonance contributor More significant resonance contributor

Rule 4: The most significant resonance contributor has the greatest number of covalent bonds.

Three covalent bonds Four covalent bonds

Less significant resonance contributor More significant resonance contributor

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 Rules of Resonance
Rule 5: If a p bond is present, the most significant resonance contributor has this p bond between atoms of
the same row of the periodic table (usually carbon p bonded to boron, carbon, nitrogen, oxygen, or
fluorine).

Carbon-chlorine double bond Carbon-fluorine double bond

Less significant resonance contributor More significant resonance contributor

Rule 6: Aromatic resonance contributors are more significant than resonance contributors that are
not aromatic.

Nonaromatic Aromatic
Less significant resonance contributor More significant resonance contributor

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 Rules of Resonance
Rule 5: If a p bond is present, the most significant resonance contributor has this p bond between atoms of
the same row of the periodic table (usually carbon p bonded to boron, carbon, nitrogen, oxygen, or
fluorine).

Carbon-chlorine double bond Carbon-fluorine double bond

Less significant resonance contributor More significant resonance contributor

Rule 6: Aromatic resonance contributors are more significant than resonance contributors that are
not aromatic.

Nonaromatic Aromatic
Less significant resonance contributor More significant resonance contributor

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 acrylic aldehyde -M

p - electron-donating groups (+M-effect) p - electron-withdrawing groups (-M-effect)

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 Hyperconjugation
According to classical resonance theory, electron delocalization could occur only via parallel
overlap of p orbitals. According to hyperconjugation, also known as no-bond resonance, and a
variant of resonance theory, electron delocalization could also occur via parallel overlap of p
orbitals with hybridized orbitals participating in sigma bonds. For example, consider the ethyl
carbocation (1), which is shown in a specific conformation (2) below.

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 Hyperconjugation
In 2, the empty p orbital on C1 and the sp3-hybridized orbital on C2 participating in C2—H1 bond are more or
less parallel, allowing parallel overlap, which lowers the electron deficiency at C1 but makes the H1 electron
deficient.

This overlap is not strong enough to completely prevent the free rotation around the C1—C2 bond.
Consequently, C2—H2 bond and C2—H3 bond could also share electrons with the empty p orbital on C1.

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 Hyperconjugation
The structure of the ethyl carbocation,
according to the theory of
hyperconjugation, can be shown
conveniently using a series of resonance
forms.

Based on the above resonance forms,

the structure of the ethyl carbocation
can be shown roughly as follows.

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 Hyperconjugation
The isopropyl carbocation is more stable than methyl carbocation by hyperconjugation, explain ?

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