KEKULE STRUCTURE
Problems with the Kekule strcuture
The Kekulé structure of benzene, proposed by August Kekulé in 1865, was a
groundbreaking step in understanding aromatic compounds. However, it has
several limitations and problems that conflict with experimental observations. Here’s
a breakdown of the key issues:
1. Stability of Benzene (Resonance Energy)
•
Kekulé’s Model: Benzene was depicted as a six-membered ring with alternating
single and double bonds (cyclohexatriene).
o
•
Expected reactivity: Should resemble conjugated dienes (e.g., undergoing
addition reactions like alkenes).
Reality: Benzene is unusually stable and resists addition reactions. Instead, it
undergoes substitution reactions (e.g., nitration, sulfonation).
o
Resonance energy: Benzene’s stability (measured as ~36 kcal/mol
resonance energy) arises from delocalized π-electrons, which Kekulé’s
static structure fails to explain.
2. Bond Lengths and Bond Order
•
Kekulé’s Model: Alternating single (1.54 Å) and double (1.34 Å) bonds.
•
Experimental Evidence:
o
X-ray crystallography shows all C–C bonds in benzene are
identical (1.39 Å), intermediate between single and double bonds.
o
This uniformity arises from electron delocalization, not fixed alternating
bonds.
3. Lack of Isomerism
•
Kekulé’s Claim: Two distinct structures (ortho-dibenzenes) with double bonds
in different positions.
o
•
Expected: Rapid interconversion between these structures (dynamic
equilibrium).
Reality:
o
No isomers of benzene are observed.
o
If Kekulé’s structures existed, substituted derivatives (e.g., odichlorobenzene) should show multiple isomers, but they do not.
o
Modern Explanation: Delocalized π-electrons create a single, symmetric
structure.
4. Hydrogenation Data
•
Kekulé’s Prediction: Hydrogenating three isolated double bonds should release
3 × the heat of hydrogenation of cyclohexene (~85.8 kcal/mol).
•
Experimental Result:
o
Benzene releases only ~49.8 kcal/mol upon hydrogenation (much less
than predicted).
o
The difference (~36 kcal/mol) is the resonance stabilization energy,
confirming benzene’s extra stability.
What is Hydrogenation?
Hydrogenation is a reaction where hydrogen (H₂) is added to a molecule, typically
converting double bonds (C=C) into single bonds (C–C). For benzene (C₆H₆),
hydrogenation adds 3 H₂ molecules to form cyclohexane (C₆H₁₂):
C₆H₆+3H₂→C₆H₁₂(ΔH=−49.8 kcal/mol)C₆H₆+3H₂→C₆H₁₂(ΔH=−49.8kcal/mol)
Expected vs. Observed Energy Release
•
Expected ΔH (if benzene had three isolated double bonds):
Hydrogenating one C=C bond (e.g., in cyclohexene) releases -28.6 kcal/mol:
Cyclohexene+H₂→Cyclohexane(ΔH=−28.6 kcal/mol)Cyclohexene+H₂→Cyclohexane(ΔH=
−28.6kcal/mol)
For three isolated double bonds, total energy release would be:
3×(−28.6)=−85.8 kcal/mol.3×(−28.6)=−85.8kcal/mol.
•
Observed ΔH (actual benzene):
Benzene releases only -49.8 kcal/mol when hydrogenated.
This is 36 kcal/mol less exothermic than expected.
The smaller energy release (-49.8 vs. -85.8 kcal/mol) means benzene starts at a lower
energy (is more stable) than a molecule with three isolated double bonds.
o
Benzene is already "relaxed," so releasing it (hydrogenation) gives less
energy.
o
Kekulé structure is "tightly wound," so releasing it gives much more energy.
1. Lower ΔH = Higher Stability:
Benzene’s hydrogenation is less exothermic because it’s already stabilized
by resonance.
2. Resonance Energy:
The 36 kcal/mol difference proves benzene is 36 kcal/mol more stable than
the hypothetical Kekulé structure.
3. Aromaticity:
Delocalized electrons (aromaticity) are why benzene resists reactions that
break its π-system (like addition reactions).
5. Magnetic and Spectroscopic Properties
•
Kekulé’s Model: Fails to explain benzene’s unique magnetic behavior.
o
Diamagnetic Ring Current: Benzene exhibits a ring current in a magnetic
field (observed via NMR), a hallmark of aromaticity.
o
Delocalized electrons generate this current, which is impossible in a
structure with localized double bonds.
6. Failure to Explain Aromaticity
•
Kekulé’s Structure: Does not account for Hückel’s rule (4n+2 π-electrons) or
the concept of aromaticity.
o
Benzene’s planar structure, conjugation, and delocalized π-cloud define
its aromatic behavior, not alternating bonds.
Resolution: Modern Theories
1. Resonance Theory:
o
Benzene is a hybrid of two equivalent Kekulé structures, with delocalized
π-electrons.
o
Explains bond uniformity, stability, and lack of isomers.
2. Molecular Orbital (MO) Theory:
o
π-electrons occupy delocalized molecular orbitals spanning the entire
ring.
o
Accounts for aromaticity, bond lengths, and magnetic properties.
1. What is Resonance?
•
Definition: Resonance describes the delocalization of π-electrons or lone
pairs over multiple atoms, resulting in a hybrid structure that is more stable
than any individual contributing structure.
•
Key Idea: No single Lewis structure fully represents the molecule. Instead, the
true structure is a weighted average (resonance hybrid) of all valid resonance
structures.
2. Why is Resonance Necessary?
•
Limitations of Lewis Structures:
o
Lewis structures assume localized electrons, but many molecules (e.g.,
benzene, ozone) exhibit delocalized bonding that cannot be captured by
a single structure.
o
Example: Benzene cannot be represented by alternating single/double
bonds (Kekulé structures) because all C–C bonds are experimentally
identical.
3. Key Postulates of Resonance Theory
1. Multiple Valid Lewis Structures:
o
Resonance structures must be valid Lewis structures (obey octet rule,
proper formal charges).
o
Example: Ozone (O₃) has two resonance forms with delocalized double
bonds.
2. Same Atomic Positions:
o
Only electrons (π-electrons or lone pairs) move; atoms remain in fixed
positions.
3. Resonance Hybrid:
o
The actual molecule is a hybrid of all resonance structures, not a rapid
interconversion between them.
4. Increased Stability:
o
The hybrid is more stable than any individual resonance structure
(resonance stabilization energy).
4. Rules for Drawing Resonance Structures
1. Only π-electrons and lone pairs move (σ-bond framework remains intact).
2. Avoid breaking single bonds or violating the octet rule (except for hypervalent
species like SO₄²⁻).
3. Follow electron movement conventions:
o
Curved arrows show electron flow:
▪
A double-headed arrow (↔) connects resonance structures.
▪
A curved arrow (↷) indicates electron movement.
4. Maximize Bonding:
o
Structures with more bonds and fewer formal charges are better
contributors.
5. Examples of Resonance
a) Benzene (C₆H₆)
•
Two Kekulé structures with alternating double bonds.
•
Hybrid: All six C–C bonds are identical (1.39 Å), intermediate between single and
double bonds.
Resonance hybrid=Average of all contributing structuresResonance hybrid=Average of a
ll contributing structures
b) Ozone (O₃)
•
Two resonance structures:
O=O–O↔O–O=OO=O–O↔O–O=O
•
Hybrid: Delocalized π-electrons create two equivalent O–O bonds (bond order =
1.5).
c) Carbonate Ion (CO₃²⁻)
•
Three resonance structures with double bonds rotating among the three oxygen
atoms.
•
Hybrid: All C–O bonds are equivalent (bond order = 1.33).
6. Resonance Effects on Molecular Properties
1. Stability:
o
Resonance lowers energy (e.g., benzene is stabilized by ~36 kcal/mol).
2. Bond Lengths:
o
Bond lengths average between single and double bonds (e.g., 1.39 Å in
benzene).
3. Acidity/Basicity:
o
Resonance stabilizes conjugate bases (e.g., carboxylic acids are acidic
because their conjugate base is resonance-stabilized).
4. Reactivity:
o
Delocalization directs electrophilic substitution in aromatic compounds
(e.g., benzene nitration).
7. Resonance vs. Other Concepts
•
•
Resonance vs. Tautomerism:
o
Tautomers are real, interconvertible structures (e.g., keto-enol
tautomerism).
o
Resonance structures are hypothetical and do not interconvert.
Resonance vs. Hyperconjugation:
o
Hyperconjugation involves σ-electrons delocalizing (e.g., stabilizing
carbocations).
8. Common Misconceptions
•
Myth: "Resonance structures are in equilibrium."
o
•
Reality: The molecule exists as a single hybrid, not a mix of structures.
Myth: "Resonance creates new bonds."
o
Reality: It redistributes electron density without altering the σ-bond
framework.
9. Limitations of Resonance Theory
•
Qualitative Only: Does not quantify stabilization energy.
•
Fails for Extended Systems: Molecular Orbital (MO) Theory is better for large
conjugated systems (e.g., graphene).
Understanding the structure of Benzene.
Hybridization, in Chemistry, is defined as the concept of mixing two atomic orbitals to
give rise to a new type of hybridized orbitals. This intermixing usually results in the
formation of hybrid orbitals having entirely different energy, shapes, etc. The atomic
orbitals of the same energy level mainly take part in hybridization. However, both fullyfilled and half-filled orbitals can also take part in this process, provided they have equal
energy.
sp Hybridization
sp hybridization is observed when one s and one p orbital in the same main shell of an
atom mix to form two new equivalent orbitals. The new orbitals formed are called sp
hybridized orbitals. It forms linear molecules with an angle of 180°.
•
This type of hybridization involves the mixing of one ‘s’ orbital and one ‘p’ orbital
of equal energy to give a new hybrid orbital known as an sp hybridized orbital.
•
The sp hybridization is also called diagonal hybridization.
•
Each sp hybridized orbital has an equal amount of s and p characters – 50% s
and 50% p characters.
sp3 Hybridization
When one ‘s’ orbital and 3 ‘p’ orbitals belonging to the same shell of an atom mix
together to form four new equivalent orbitals, the type of hybridization is called
a tetrahedral hybridization or sp3. The new orbitals formed are called sp3 hybrid
orbitals.
•
These are directed towards the four corners of a regular tetrahedron and make
an angle of 109°28’ with one another.
•
The angle between the sp3 hybrid orbitals is 109.280
•
Each sp3 hybrid orbital has 25% s character and 75% p character.
•
Examples of sp3 hybridization are ethane (C2H6) and methane.
Benzene (C₆H₆) is a planar, hexagonal aromatic hydrocarbon with a unique electronic
structure due to sp² hybridization of its carbon atoms. Here's a breakdown:
1. Hybridization of Carbon in Benzene
Each carbon atom in benzene undergoes sp² hybridization:
•
•
Atomic Orbitals Involved:
o
One 2s orbital + two 2p orbitals (e.g., 2pₓ, 2pᵧ) → Three sp² hybrid orbitals.
o
One unhybridized 2p_z orbital remains perpendicular to the plane.
Geometry:
o
Trigonal planar arrangement (bond angles = 120°).
o
Explains the hexagonal symmetry of benzene.
2. Bond Formation in Benzene
•
Sigma (σ) Bonds:
o
•
Each carbon uses three sp² orbitals to form σ bonds:
▪
Two with adjacent carbons (C–C σ bonds).
▪
One with a hydrogen atom (C–H σ bond).
Pi (π) Bonds:
o
The unhybridized 2p_z orbitals of all six carbons overlap sideways
(laterally) to form a delocalized π-electron cloud above and below the
ring.
o
This creates a resonance-stabilized conjugated system (not alternating
single/double bonds).
3. Delocalization and Resonance
•
•
Resonance Hybrid:
o
The six π-electrons are delocalized over all six carbon atoms.
o
Results in equal bond lengths (1.39 Å), intermediate between single
(1.54 Å) and double bonds (1.34 Å).
Resonance Stabilization Energy:
o
Benzene is ~36 kcal/mol more stable than hypothetical cyclohexatriene
(Kekulé structure).
sp² Hybridization
1. Planar Structure:
o
All atoms lie in the same plane due to sp² hybridization.
2. Reactivity:
o
Delocalized π-electrons make benzene resistant to addition reactions
(unlike alkenes).
o
Prefers electrophilic substitution to preserve aromaticity.
3. Magnetic Properties:
o
The ring current from delocalized electrons causes diamagnetic
anisotropy (observed in NMR).
5. Comparison with Other Hybridization States
Feature
Benzene (sp²)
Alkane (sp³)
Alkene (sp²)
Bond
Angles
120°
~109.5°
120°
Bond Type
Delocalized πsystem
Single σ bonds
Localized π bond
Reactivity
Aromatic
substitution
Free radical
substitution
Addition
reactions
6. Why Benzene Doesn’t Behave Like an Alkene
•
Localized vs. Delocalized π Bonds:
o
In alkenes (e.g., ethylene), π-electrons are localized between two
carbons, making them reactive.
o
In benzene, delocalization distributes electron density evenly, reducing
reactivity.
Summary
The sp² hybridization in benzene:
•
Creates a planar hexagonal structure with 120° bond angles.
•
Enables delocalization of π-electrons, leading to resonance stabilization.
•
Explains benzene’s unique stability, equal bond lengths, and preference for
substitution over addition.
This hybridization is central to the concept of aromaticity, defining benzene’s chemical
behaviour and physical properties.