This document provides an overview of organic chemistry concepts including:
- The electronic structure of atoms and how this relates to bonding
- Different types of bonds (ionic, covalent, polar, nonpolar) and how they form
- Lewis structures and resonance forms
- Molecular shapes determined by hybridization of orbitals
- Isomerism including constitutional and geometric isomers
The document covers fundamental topics that provide context for understanding organic molecular structures and reactions.
The document discusses chemical bonding theories including the octet rule, ionic bonding, covalent bonding, limitations of the octet rule, and molecules with incomplete or expanded octets. It also covers topics like Lewis structures, formal charge, valence shell electron pair repulsion theory (VSEPR) and its application to molecules like NH3 and H2O. Further, it summarizes valence bond theory, hybridization, types of hybridization including sp3 and sp2, and their relation to the geometry of molecules.
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhDMaqsoodAhmadKhan5
applied chemistry lecture and slide,
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhD, lecturer in chemistry in pakistan institute of engineering and applied sciences
1. Organic reaction mechanisms involve the reaction of a substrate with a reagent, forming intermediates and ultimately products.
2. Bond cleavage can occur through either a heterolytic or homolytic process. Heterolytic cleavage leads to the formation of ions while homolytic cleavage leads to free radicals.
3. Electrophiles are electron seeking species that attack nucleophilic centers, while nucleophiles are electron pairing species that attack electrophilic centers. Common electrophiles include carbocations and carbonyl groups, while common nucleophiles include carbanions.
This document provides an overview of atomic structure, bonding, and electron distribution. It begins by defining the basic subatomic particles that make up atoms. It then discusses several historical atomic models including Thomson's plum pudding model, Rutherford's nuclear model, and Bohr's early quantum model. The document introduces concepts like electron orbitals and quantum numbers. It also covers bonding theories such as ionic and covalent bonding as well as localized and delocalized bonding. Hybridization of atomic orbitals is discussed through examples like sp, sp2, and sp3 hybridization. The summary concludes with an introduction to molecular orbital theory.
This document provides information about the periodic table and periodic trends. It discusses the organization of the periodic table into rows (periods) and columns (groups/families) and explains how elements in the same group have similar properties based on their electron configuration. Periodic trends are covered, including how atomic radius decreases across a period as effective nuclear charge increases, and how ionization energy and electron affinity vary periodically. Metals, nonmetals, and metalloids are defined based on their characteristic properties.
Class 11 Chapter 4 Chemical Bonding and Molecular Structure.pptxRajnishPrasadSarma
This document provides an overview of chemical bonding and molecular structure. It discusses topics such as octet rule, covalent bonds, limitations of the octet rule, ionic or electrovalent bonds, lattice enthalpy, bond parameters including bond length, bond angle, bond enthalpy and bond order. It also covers concepts of resonance, polar covalent bonds, dipole moment and covalent character in ionic bonds based on Fajans' rule. The document is presented as part of a Class XI chemistry curriculum on this unit.
The document discusses the properties of amino acids. It explains that amino acids can be hydrophobic, hydrophilic, or amphipathic based on their hydropathy index. Amphipathic amino acids are beneficial because they can tolerate changes in environment between hydrophilic and hydrophobic. The document also discusses the differences between ampholytes and amphoteric molecules, and explains concepts like electronegativity, dipole moments, and dielectric constants in the context of amino acid properties.
This document provides an overview of organic chemistry concepts including:
- The electronic structure of atoms and how this relates to bonding
- Different types of bonds (ionic, covalent, polar, nonpolar) and how they form
- Lewis structures and resonance forms
- Molecular shapes determined by hybridization of orbitals
- Isomerism including constitutional and geometric isomers
The document covers fundamental topics that provide context for understanding organic molecular structures and reactions.
The document discusses chemical bonding theories including the octet rule, ionic bonding, covalent bonding, limitations of the octet rule, and molecules with incomplete or expanded octets. It also covers topics like Lewis structures, formal charge, valence shell electron pair repulsion theory (VSEPR) and its application to molecules like NH3 and H2O. Further, it summarizes valence bond theory, hybridization, types of hybridization including sp3 and sp2, and their relation to the geometry of molecules.
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhDMaqsoodAhmadKhan5
applied chemistry lecture and slide,
Applied Chemistry, atomic and molecular structure, part 1, by Shiraz mahbob PhD, lecturer in chemistry in pakistan institute of engineering and applied sciences
1. Organic reaction mechanisms involve the reaction of a substrate with a reagent, forming intermediates and ultimately products.
2. Bond cleavage can occur through either a heterolytic or homolytic process. Heterolytic cleavage leads to the formation of ions while homolytic cleavage leads to free radicals.
3. Electrophiles are electron seeking species that attack nucleophilic centers, while nucleophiles are electron pairing species that attack electrophilic centers. Common electrophiles include carbocations and carbonyl groups, while common nucleophiles include carbanions.
This document provides an overview of atomic structure, bonding, and electron distribution. It begins by defining the basic subatomic particles that make up atoms. It then discusses several historical atomic models including Thomson's plum pudding model, Rutherford's nuclear model, and Bohr's early quantum model. The document introduces concepts like electron orbitals and quantum numbers. It also covers bonding theories such as ionic and covalent bonding as well as localized and delocalized bonding. Hybridization of atomic orbitals is discussed through examples like sp, sp2, and sp3 hybridization. The summary concludes with an introduction to molecular orbital theory.
This document provides information about the periodic table and periodic trends. It discusses the organization of the periodic table into rows (periods) and columns (groups/families) and explains how elements in the same group have similar properties based on their electron configuration. Periodic trends are covered, including how atomic radius decreases across a period as effective nuclear charge increases, and how ionization energy and electron affinity vary periodically. Metals, nonmetals, and metalloids are defined based on their characteristic properties.
Class 11 Chapter 4 Chemical Bonding and Molecular Structure.pptxRajnishPrasadSarma
This document provides an overview of chemical bonding and molecular structure. It discusses topics such as octet rule, covalent bonds, limitations of the octet rule, ionic or electrovalent bonds, lattice enthalpy, bond parameters including bond length, bond angle, bond enthalpy and bond order. It also covers concepts of resonance, polar covalent bonds, dipole moment and covalent character in ionic bonds based on Fajans' rule. The document is presented as part of a Class XI chemistry curriculum on this unit.
The document discusses the properties of amino acids. It explains that amino acids can be hydrophobic, hydrophilic, or amphipathic based on their hydropathy index. Amphipathic amino acids are beneficial because they can tolerate changes in environment between hydrophilic and hydrophobic. The document also discusses the differences between ampholytes and amphoteric molecules, and explains concepts like electronegativity, dipole moments, and dielectric constants in the context of amino acid properties.
The document discusses band theory of solids and how it explains the electrical conductivity of materials. It explains that when atoms are brought close together, the discrete energy levels of individual atoms overlap to form energy bands. Materials can be conductors, semiconductors, or insulators depending on whether their bands are fully overlapping, partially overlapping, or fully separated by a band gap. Conductors have overlapping bands allowing electrons to move freely. Semiconductors have a small band gap that electrons can cross with increased thermal energy. Insulators have a large band gap preventing electron flow.
This document discusses molecular polarity and how to determine if molecules are polar or nonpolar. It defines polar and nonpolar covalent bonds based on differences in electronegativity between atoms. A molecule is polar if it contains polar bonds in an asymmetrical arrangement, whereas a molecule is nonpolar if the polar bonds are symmetrical or if all bonds are nonpolar. Examples of polar molecules include HCl and H2O, while nonpolar examples include CO2 and CF4.
Lewis and Kossel were the first to successfully explain chemical bonding in terms of electrons. They proposed that atoms bond to achieve stable electron configurations like noble gases. Lewis pictured atoms as a positively charged nucleus surrounded by 8 electrons. Atoms bond by transferring or sharing electrons to achieve stable octets. Lewis symbols were introduced to represent valence electrons. Covalent bonds form when atoms share valence electrons in order to gain full outer shells. Resonance structures represent situations where no single Lewis structure adequately describes the molecule.
The document discusses atomic structure and chemical bonds. It explains that atoms are made up of protons and neutrons in the nucleus surrounded by electrons in energy levels. The number and arrangement of electrons determines an element's properties. Electrons closer to the nucleus have lower energy than those farther away. Elements in the same family on the periodic table have similar properties because they have the same number of electrons in their outer energy level, which determines how they bond with other elements.
CH 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE 3.pdfLUXMIKANTGIRI
Lewis and Kossel were the first to successfully explain chemical bonding in terms of electrons. They proposed that atoms bond to achieve stable electron configurations like noble gases. Lewis pictured atoms as a positively charged nucleus surrounded by 8 electrons. He proposed that atoms bond by transferring or sharing electrons to achieve stable octets. Lewis symbols were introduced to represent valence electrons. Covalent bonds form when atoms share valence electrons to achieve stable octets. Resonance structures represent situations where no single Lewis structure adequately describes the molecule.
UV spectroscopy is an analytical method used to detct the numbers of double and triple bonds present in dienes ,trienes and polyenes compounds.The energy corresponds to EM radiation in the ultraviolet (UV) region, 100-350 nm, and visible (VIS) regions 350-700 nm of the spectrum is known as UV spectrum.
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Resonance structures represent different arrangements of electrons in a molecule that have the same arrangement of nuclei. They are hypothetical structures that contribute to the true structure, which is a hybrid of all resonance forms. Benzene is an example where the true structure is an average of different Kekule and Dewar structures, giving each C-C bond equal length and bond order of 1.5. Resonance increases stability by delocalizing electrons over the whole molecule.
1. Semiconductors have electrical conductivity between conductors and insulators. Their small forbidden energy gap allows electrons to cross into the conduction band with increased temperature or an electric field.
2. Intrinsic semiconductors have few free electrons/holes at normal temperatures due to broken covalent bonds generating electron-hole pairs. Extrinsic semiconductors are doped with impurities to add more charge carriers.
3. N-type semiconductors are doped with pentavalent atoms whose extra electrons become donors in the conduction band. P-type are doped with trivalent atoms whose holes are the majority carriers.
1. Localized bonds involve electron density concentrated between two nuclei, while delocalized bonds involve electron density spread across multiple nuclei.
2. Conjugated systems have alternating single and multiple bonds, allowing p-orbitals to overlap and form delocalized molecular orbitals spread across multiple atoms. This delocalization increases stability.
3. Examples given include the allyl radical and cation, which have three p-orbitals overlapping to form bonding and antibonding molecular orbitals delocalized across all three carbons. 1,3-Butadiene also has four overlapping p-orbitals forming delocalized molecular orbitals.
This document provides an overview of UV spectroscopy. It discusses electronic transitions that can be observed via UV spectroscopy, including n→π*, π→π*, n→s*, and s→s* transitions. The energy required for different transitions is discussed, with n→π* requiring the lowest energy. Selection rules and factors that influence the observation of transitions are also covered. The document introduces concepts like chromophores, auxochromes, and how they can shift absorption bands.
This document discusses electronic spectroscopy techniques such as ultraviolet-visible (UV-vis) spectroscopy and chiroptical spectroscopy. It describes how UV-vis spectroscopy can be used to study electronic transitions in molecules, detect functional groups, and perform quantitative analysis. Chiroptical spectroscopy techniques like circular dichroism and optical rotary dispersion are used to determine stereochemistry by measuring differences in absorption of left and right circularly polarized light. Selection rules and origins of electronic transitions are also summarized.
This document discusses electronic spectroscopy techniques such as ultraviolet-visible (UV-vis) spectroscopy and chiroptical spectroscopy. It describes how UV-vis spectroscopy can be used to study electronic transitions in molecules, detect functional groups, and perform quantitative analysis. The fundamentals of UV-vis spectroscopy, including electronic transitions, selection rules, solvent effects, and absorption intensity are summarized. Chiroptical spectroscopy techniques like circular dichroism and optical rotary dispersion are also introduced.
The document discusses theories related to molecular structure and chemical bonding, including:
- Lewis theory which proposes that atoms bond by sharing valence electrons to achieve stable octet configurations.
- Limitations of the octet rule are discussed, including cases where the central atom has an incomplete or expanded octet.
- Sidgwick–Powell theory predicts molecular geometry based on the number of electron pairs around the central atom.
- VSEPR theory builds on this by accounting for differences in repulsion between bond pairs and lone pairs, allowing for more accurate prediction of molecular geometry. Lone pairs occupy more space and influence the shape.
1) Semiconductors have energy band gaps between 1-4 eV allowing electrons to move between bands with small energy inputs, unlike insulators. Metals have overlapping or no bands.
2) Intrinsic semiconductors have small numbers of electrons and holes as charge carriers from thermal excitation breaking covalent bonds. Extrinsic semiconductors are doped with donor or acceptor impurities to add more free electrons or holes.
3) Donor impurities introduce extra electrons in the conduction band of n-type semiconductors. Acceptor impurities introduce holes in the valence band of p-type semiconductors.
Ultraviolet spectroscopy involves the absorption of UV radiation by molecules, causing electronic transitions in valence electrons. The UV region is divided into near (2000-4000 Å) and far/vacuum (below 2000 Å) regions. UV absorption spectra arise from transitions of electrons between lower and higher electronic energy levels. Important terms include chromophores, which absorb UV radiation, and auxochromes, which shift absorption maxima. Woodward-Feiser rules can be used to calculate absorption maxima based on molecular structure. Common instrumentation includes hydrogen and deuterium lamps as well as mercury arcs as UV sources.
The document summarizes properties of the noble gas element xenon. It discusses xenon's electronic configuration, atomic structure, crystal lattice structure, isotopes, and reactivity including compounds formed with fluorine and oxygen such as XeF2, XeF4, XeF6, XeO3, XeOF4, and XeO2F2. Xenon has various applications in lighting, medicine, and nuclear technology due to its unique properties as a noble gas.
This document provides a summary of key concepts regarding transition elements covered in an inorganic chemistry lecture note. It discusses the definition of transition elements, their electronic configurations, atomic radii, ionization potentials, variable oxidation states, and ability to form metal complexes. It also introduces coordination chemistry concepts like metal complexes, ligands, and bonding theories. Properties of octahedral and tetrahedral complexes are examined along with nomenclature of coordination compounds.
How to Build a Module in Odoo 17 Using the Scaffold MethodCeline George
Odoo provides an option for creating a module by using a single line command. By using this command the user can make a whole structure of a module. It is very easy for a beginner to make a module. There is no need to make each file manually. This slide will show how to create a module using the scaffold method.
The document discusses band theory of solids and how it explains the electrical conductivity of materials. It explains that when atoms are brought close together, the discrete energy levels of individual atoms overlap to form energy bands. Materials can be conductors, semiconductors, or insulators depending on whether their bands are fully overlapping, partially overlapping, or fully separated by a band gap. Conductors have overlapping bands allowing electrons to move freely. Semiconductors have a small band gap that electrons can cross with increased thermal energy. Insulators have a large band gap preventing electron flow.
This document discusses molecular polarity and how to determine if molecules are polar or nonpolar. It defines polar and nonpolar covalent bonds based on differences in electronegativity between atoms. A molecule is polar if it contains polar bonds in an asymmetrical arrangement, whereas a molecule is nonpolar if the polar bonds are symmetrical or if all bonds are nonpolar. Examples of polar molecules include HCl and H2O, while nonpolar examples include CO2 and CF4.
Lewis and Kossel were the first to successfully explain chemical bonding in terms of electrons. They proposed that atoms bond to achieve stable electron configurations like noble gases. Lewis pictured atoms as a positively charged nucleus surrounded by 8 electrons. Atoms bond by transferring or sharing electrons to achieve stable octets. Lewis symbols were introduced to represent valence electrons. Covalent bonds form when atoms share valence electrons in order to gain full outer shells. Resonance structures represent situations where no single Lewis structure adequately describes the molecule.
The document discusses atomic structure and chemical bonds. It explains that atoms are made up of protons and neutrons in the nucleus surrounded by electrons in energy levels. The number and arrangement of electrons determines an element's properties. Electrons closer to the nucleus have lower energy than those farther away. Elements in the same family on the periodic table have similar properties because they have the same number of electrons in their outer energy level, which determines how they bond with other elements.
CH 4 CHEMICAL BONDING AND MOLECULAR STRUCTURE 3.pdfLUXMIKANTGIRI
Lewis and Kossel were the first to successfully explain chemical bonding in terms of electrons. They proposed that atoms bond to achieve stable electron configurations like noble gases. Lewis pictured atoms as a positively charged nucleus surrounded by 8 electrons. He proposed that atoms bond by transferring or sharing electrons to achieve stable octets. Lewis symbols were introduced to represent valence electrons. Covalent bonds form when atoms share valence electrons to achieve stable octets. Resonance structures represent situations where no single Lewis structure adequately describes the molecule.
UV spectroscopy is an analytical method used to detct the numbers of double and triple bonds present in dienes ,trienes and polyenes compounds.The energy corresponds to EM radiation in the ultraviolet (UV) region, 100-350 nm, and visible (VIS) regions 350-700 nm of the spectrum is known as UV spectrum.
Resonance structures represent different arrangements of electrons in a molecule that have the same positions of nuclei but different bonding patterns. Resonance contributes to the stability of molecules like benzene by delocalizing electrons across multiple equivalent structures. The actual structure of a molecule represented by resonance is a hybrid of the contributing structures, with bond lengths intermediate between single and double bonds. Delocalization of electrons is depicted using curved arrows between resonance structures.
Resonance structures represent different arrangements of electrons in a molecule that have the same arrangement of nuclei. They are hypothetical structures that contribute to the true structure, which is a hybrid of all resonance forms. Benzene is an example where the true structure is an average of different Kekule and Dewar structures, giving each C-C bond equal length and bond order of 1.5. Resonance increases stability by delocalizing electrons over the whole molecule.
1. Semiconductors have electrical conductivity between conductors and insulators. Their small forbidden energy gap allows electrons to cross into the conduction band with increased temperature or an electric field.
2. Intrinsic semiconductors have few free electrons/holes at normal temperatures due to broken covalent bonds generating electron-hole pairs. Extrinsic semiconductors are doped with impurities to add more charge carriers.
3. N-type semiconductors are doped with pentavalent atoms whose extra electrons become donors in the conduction band. P-type are doped with trivalent atoms whose holes are the majority carriers.
1. Localized bonds involve electron density concentrated between two nuclei, while delocalized bonds involve electron density spread across multiple nuclei.
2. Conjugated systems have alternating single and multiple bonds, allowing p-orbitals to overlap and form delocalized molecular orbitals spread across multiple atoms. This delocalization increases stability.
3. Examples given include the allyl radical and cation, which have three p-orbitals overlapping to form bonding and antibonding molecular orbitals delocalized across all three carbons. 1,3-Butadiene also has four overlapping p-orbitals forming delocalized molecular orbitals.
This document provides an overview of UV spectroscopy. It discusses electronic transitions that can be observed via UV spectroscopy, including n→π*, π→π*, n→s*, and s→s* transitions. The energy required for different transitions is discussed, with n→π* requiring the lowest energy. Selection rules and factors that influence the observation of transitions are also covered. The document introduces concepts like chromophores, auxochromes, and how they can shift absorption bands.
This document discusses electronic spectroscopy techniques such as ultraviolet-visible (UV-vis) spectroscopy and chiroptical spectroscopy. It describes how UV-vis spectroscopy can be used to study electronic transitions in molecules, detect functional groups, and perform quantitative analysis. Chiroptical spectroscopy techniques like circular dichroism and optical rotary dispersion are used to determine stereochemistry by measuring differences in absorption of left and right circularly polarized light. Selection rules and origins of electronic transitions are also summarized.
This document discusses electronic spectroscopy techniques such as ultraviolet-visible (UV-vis) spectroscopy and chiroptical spectroscopy. It describes how UV-vis spectroscopy can be used to study electronic transitions in molecules, detect functional groups, and perform quantitative analysis. The fundamentals of UV-vis spectroscopy, including electronic transitions, selection rules, solvent effects, and absorption intensity are summarized. Chiroptical spectroscopy techniques like circular dichroism and optical rotary dispersion are also introduced.
The document discusses theories related to molecular structure and chemical bonding, including:
- Lewis theory which proposes that atoms bond by sharing valence electrons to achieve stable octet configurations.
- Limitations of the octet rule are discussed, including cases where the central atom has an incomplete or expanded octet.
- Sidgwick–Powell theory predicts molecular geometry based on the number of electron pairs around the central atom.
- VSEPR theory builds on this by accounting for differences in repulsion between bond pairs and lone pairs, allowing for more accurate prediction of molecular geometry. Lone pairs occupy more space and influence the shape.
1) Semiconductors have energy band gaps between 1-4 eV allowing electrons to move between bands with small energy inputs, unlike insulators. Metals have overlapping or no bands.
2) Intrinsic semiconductors have small numbers of electrons and holes as charge carriers from thermal excitation breaking covalent bonds. Extrinsic semiconductors are doped with donor or acceptor impurities to add more free electrons or holes.
3) Donor impurities introduce extra electrons in the conduction band of n-type semiconductors. Acceptor impurities introduce holes in the valence band of p-type semiconductors.
Ultraviolet spectroscopy involves the absorption of UV radiation by molecules, causing electronic transitions in valence electrons. The UV region is divided into near (2000-4000 Å) and far/vacuum (below 2000 Å) regions. UV absorption spectra arise from transitions of electrons between lower and higher electronic energy levels. Important terms include chromophores, which absorb UV radiation, and auxochromes, which shift absorption maxima. Woodward-Feiser rules can be used to calculate absorption maxima based on molecular structure. Common instrumentation includes hydrogen and deuterium lamps as well as mercury arcs as UV sources.
The document summarizes properties of the noble gas element xenon. It discusses xenon's electronic configuration, atomic structure, crystal lattice structure, isotopes, and reactivity including compounds formed with fluorine and oxygen such as XeF2, XeF4, XeF6, XeO3, XeOF4, and XeO2F2. Xenon has various applications in lighting, medicine, and nuclear technology due to its unique properties as a noble gas.
This document provides a summary of key concepts regarding transition elements covered in an inorganic chemistry lecture note. It discusses the definition of transition elements, their electronic configurations, atomic radii, ionization potentials, variable oxidation states, and ability to form metal complexes. It also introduces coordination chemistry concepts like metal complexes, ligands, and bonding theories. Properties of octahedral and tetrahedral complexes are examined along with nomenclature of coordination compounds.
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3. INTRODUCTION
• Free radicals are formed in the organic reactions
when homolytic cleavage of bond takes place.
• In general, the free radicals are formed in the
reactions which occur in the presence of light or at
high temperature or in the presence of organic
peroxides such as benzoylperoxide.
• Free radicals are the species with a single electron
• They are electron deficient in nature (a species with
single electron always tends to pair up with another
electron and thus, looks for an electron rich site).
4. Characteristic features of free radicals
• A carbon free radical is sp2 hybridized with a p
orbital carrying single unpaired electron.
• A free radical has a planar geometry.
• The carbon in a radical is trivalent and has seven
electrons (septet).
5. • A radical, once formed, reacts immediately to
extract another radical from a bond and thus,
results in generation of new radical.
• For this reason, the reactions, which involve
free radicals as intermediates, are chain
reactions.
• Such reactions terminate only when two free
radicals combine with each other.
6. Stability of free radicals
• A free radical may be 1o, or 2o, or 3o, depending upon
the number of carbons attached to carbon carrying single
electron.
• Further, the free radicals may be categorised as alkyl,
allyl or benzyl radicals.
• For example
• The stability aspects of different categories of free radicals
may be explained through inductive, hyperconjugation and
resonance effect.
7. Stability of alkyl free radicals:
Explanation through inductive effect
• Free radicals are electron deficient species where
carbon carries a single unpaired electron.
• The alkyl groups release electrons through inductive
effect (+I effect).
• More the number of alkyl group attached to a carbon
radical, more is the availability of electrons and more
is the stabilization of free radical.
8. • This electron deficiency in carbon radicals is
compensated to maximum in 3o radicals because of
the presence of three electron releasing groups.
• Thus, the order of stability of methyl substituted free
radicals is as
• follows:
9. Radicals are stabilized by electron-
withdrawing groups
• when a radical centre finds itself next to an electron-
withdrawing group.
• Groups like C=O and C≡N are electron withdrawing because
they have a low-lying empty π* orbital.
• By overlapping with the (usually p) orbital containing the
radical (the SOMO), two new molecular orbitals are generated
• One electron (the one in the old SOMO) is available to fill the
two new orbitals.
• It enters the new SOMO, which is of lower energy than the old
one, and the radical experiences stabilization because this
electron drops in energy.
11. • Electron-rich groups, such as RO groups, in a similar way.
• Ether oxygen atoms have relatively high-energy filled n
orbitals,
• Their lone pairs interacting this with the SOMO again gives
two new molecular
• Three electrons are available to fill them.
• The SOMO is now higher in energy than it was to start with,
but the lone pair is lower.
• Because two electrons have dropped in energy and only one
has risen, there is an overall stabilization of the system, even
though the new SOMO is of higher energy than the old one.
Radicals are stabilized by electron-Donating
groups
13. Stability of alkyl free radicals:
Explanation through hyperconjugation
• The stability of free radicals may be explained not
only through inductive effect but also through
hyperconjugation.
• The C–H σ bond which is in conjugation with the p
orbital carrying a single electron, participates in
delocalization.
• For example, in case of ethyl radical (CH3 H2), the
three C–H σ bonds (of CH3 group) are in conjugation
with the p orbital on CH2 (carrying a single electron),
as shown below:
15. Hyperconjugation
• In (CH3)3C•, there are nine C–H σ bonds which participate in
delocalization with p orbital of the carbon carrying single
unpaired electron thus, nine contributing structures can be
obtained.
• In case of (CH3)2CH•,there are only six C–H σ bonds
available for participation in delocalization and only six
contributing structures are possible.
• There are three C–H σ bonds in CH3CH2•, available for
participation in delocalization and only three contributing
structures are possible.
16. • Thus, (CH3)3C• is more stable than (CH3)2CH• which in turn
is more stable than CH3CH2•. The CH3• is least stable as
there is no C–H σ bond available for participation in
delocalization with p orbital of the carbon carrying single
unpaired electron.
• Thus, the overall stability of free radicals can be given as
(CH3)3C• > (CH3)2C•H > CH3C•H2 > C•H3
17. Stability of allyl and benzyl radicals:
Explanation through Resonance
• The allyl and benzyl radicals are 1° free radicals but in
comparison to 1° alkyl radicals the allyl and benzyl radicals
are more stable.
• The stability of these radicals is attributed to resonance effect.
• In both the cases the p orbital, on carbon, carrying a single
electron is in conjugation with π bond and thus delocalization
occurs through p-π overlap.
18. • The benzyl radical is more stable compared to allyl
radical, because in radical, number of contributing
structures are more as compared to those in allylic
radical.
• Due to same reasons the diphenylmethyl radical (6
contributing structures) and triphenylmethyl radical
(9 contributing structures) are more stable compared
to benzyl radical (3 contributing structures).
19. • Due to resonance effect, the allyl and benzyl radicals are more
stable than 1o, 2o, or 3o alkyl free radicals.
• Thus, we can summarize the stability order of free radicals
discussed so far, as follows:
Benzylic > allylic > 3o > 2o > 1o > C•H3
20. Detection of free radicals
• The single electron have spinning motion and can be produced
magnetic field and magnetic momentum.
• While the single paired electrons can cancel the effects of each
other and cannot produced any magnetic field and magnetic
momentum.
• Therefore the free radicals can be detected by following
methods:
1. Magnetic susceptibility measurement
2. ESR Technique
3. Spin trapping technique
21. Magnetic susceptibility measurement
• The magnetic properties of free radical provide a
powerful tool for their detection and study.
• Detection of free radical associated with spin of
electron in magnetic field which can be expressed in
quantum number.
Quantum number of +1/2 and -1/2
• According to Pauli principle : any two electron
occupying the same orbital must have opposite spin
so that their magnetic momentum is zero for any
species in which all the electrons are paired.
22. Magnetic susceptibility
• Molecules with even numbers of paired electrons are
diamagnetic; i.e., they are slightly repelled by magnets
• Free radicals, however, are paramagnetic (attracted by
magnet) because of the spin of odd electron, the spins of
remaining paired electrons effectively cancelling each
other.
• Therefore in free radicals due to the presence of one or
more unpaired electrons there is net magnetic moment
and the species is paramagnetic.
• Therefore the free radicals can be detected by the
magnetic susceptibility.
23. Electron Spin Resonance spectroscopy
• The electron spin resonance spectra of free radicals
provide another technique for their detection and
study.
• A much more important technique is ESR this
method is similar to NMR technique except the
electron spin is involved rather than nuclear spin.
• The two electron spin states : +1/2 and -1/2
• ESR spectrum can detect both the presence of free
radical and their concentration.
24. • These two spin states are ordinarily of equal energy
but in magnetic field is applied the energies are
different and the electrons are excited from the lower
state to higher state by the absorption of appropriate
frequency signals.
• If two electrons are paired in orbital have opposite
spin cancel the effect of each other and hence ESR
spectrum arises only for a species that have one or
more unpaired electrons.
• Since the free radical only gives the ESR spectrum.
• ESR have been observed for the free radical with life
time considerably less than 1 second.
26. Spin Trapping Technique
• In case of failure of ESR technique another technique
is used called spin trapping technique for detection
and involvement of free radical.
• Spin trapping agent in this technique a compound is
added to reaction medium which combine to reactive
free radical to produced more persistent radical
which can be observed in ESR technique.
27. Spin trapping
• Free radicals are trapped by trapping agent and then
detected by ESR spectrum.
Trapping agent + free radical reaction ESR detector
• Example : azulenyl nitrone is used as spin trapping
reagent the most important spin trapping agents are
nitroso compounds, nitrosyl compounds which react
with to give stable nitroxide radicals which can be
determined by the ESR technique.