SEMESTER III Course Code: CHE3C03 Complementary Course III: ORGANIC CHEMISTRY

Module I: Organic Chemistry – Some Basic Concepts (9 hrs)

Introduction: Origin of organic chemistry – Uniqueness of carbon – Homologous series – Nomenclature of alkyl halides, alcohols, aldehydes, ketones, carboxylic acids and amines. Structural isomerism: Chain isomerism, position isomerism, functional isomerism and metamerism. Hybridisation in organic molecules (a brief study) - Curved arrow formalism - Homolysis and heterolysis of bonds – Electrophiles and nucleophiles.

Electron Displacement Effects: Inductive effect: Definition - Characteristics - +I and -I groups. Applications: Explanation of substituent effect on the acidity of aliphatic carboxylic acids. Mesomeric effect: Definition – Characteristics - +M and -M groups. Applications: Comparison of electron density in benzene, nitrobenzene and aniline. Hyperconjugation: Definition – Characteristics. Example: Propene. Applications: Comparison of stability of 1-butene & 2-butene. Electromeric effect: Definition - Characteristics - +E effect (addition of H+ to ethene) and -E effect (addition of CN- to acetaldehyde). Steric effect (causes and simple examples).

Reaction Intermediates: Carbocations, carbanions and free radicals (types, hybridization and stability).

Carbon is the only element that can form so many different compounds because each carbon atom can form four chemical bonds to other atoms, and because the carbon atom is just the right, small size to fit in comfortably as parts of very large molecules.

Having the atomic number 6, every carbon atom has a total of six electrons. Two are in a completed inner orbit, while the other four are valence electrons — outer electrons that are available for forming bonds with other atoms. The carbon atom's four valence electrons can be shared by other atoms that have electrons to share, thus forming covalent (shared-electron) bonds. They can even be shared by other carbon atoms, which in turn can share electrons with other carbon atoms and so on, forming long strings of carbon atoms, bonded to each other.

 Carbon's ability to form long carbon-to-carbon chains is the first reason that there can be so many different carbon compounds; a molecule that differs by even one atom is, of course, a molecule of a different compound.  The second reason for carbon's astounding compound-forming ability is that carbon atoms can bind to each other not only in straight chains, but in complex branchings. They can even join "head-to-tail" to make rings of carbon atoms. There is practically no limit to the number or complexity of the branches or the number of rings that can be attached to them, and hence no limit to the number of different molecules that can be formed.  The third reason is that carbon atoms can share not only a single electron with another atom to form a single bond, but it can also share two or three electrons, forming a double or triple bond. This makes for a huge number of possible bond combinations at different places, making a huge number of different possible molecules

 The fourth reason is that the same collection of atoms and bonds, but in a different geometrical arrangement within the molecule, makes a molecule with a different shape and hence different properties. These different molecules are called isomers.  The fifth reason is that all of the electrons that are not being used to bond carbon atoms together into chains and rings can be used to form bonds with atoms of several other elements. The most common other element is hydrogen, which makes the family of compounds known as hydrocarbons. But nitrogen, oxygen, phosphorus, sulphur, halogens, and several other kinds of atoms can also be attached as part of an organic molecule. There is a huge number of ways in which they can be attached to the carbon-atom branches, and each variation makes a molecule of a different compound. These atoms when attached to the rest of the carbon frame work are known as a functional groups and it defines the chemical and physical properties of that particular molecule.  The sixth reason is that carbon has the ideal atomic size to form chemical bonds with even larger sized atoms to be part of large molecules yet avoid considerable steric effect in it. This enables carbon to make large but stable complex structured molecules.

Catenation: Catenation is the binding of an element to itself through covalent bonds to form chain or ring molecules. Carbon is the most common element that exhibits catenation. It can form long hydrocarbon chains through the formation large number of C-C bonds and even cyclic structures possible if a head to tail bond formation has taken place inside the molecule. carbon is not the only one with catenation property. Silicon shows it to quite a good extent, sulphur and boron has also been shown to catenate. Carbon has highest degree of catenation because:

 high C-C bond energy  tetravalency (large number of bonds)  small atomic size hence less diffused orbital

Broad Classification of organic compounds

There are a large number of organic compounds and therefore a proper systematic classification was required. Organic compounds can be broadly classified as acyclic (open chain) or cyclic (closed chain). Moving on to their classification in detail:

Acyclic or open chain compounds: Organic compounds in which all the carbon atoms are linked to one another to form open chains (straight or branched) are called acyclic or open chain compounds. These may be either saturated or unsaturated. These compounds are also called as aliphatic compounds.

Alicyclic or closed chain or ring compounds: These are cyclic compounds which contain carbon atoms connected to each other in a ring (homocyclic). When atoms other than carbon are also present then it is called as heterocyclic. They exhibit some properties similar to aliphatic compounds. Examples of this type are as follows:

Aromatic compounds: These compounds consist of at least one benzene ring, i.e., a six-membered carbocyclic ring having alternate single and double bonds. Generally, these compounds have some

fragrant odour and hence, named as aromatic (Greek word aroma meaning sweet smell).Similar to alicyclic, they can also have hetero atoms in the ring. Such compounds are called as heterocyclic aromatic compounds. Some of the examples are as follows:

Benzenoid aromatic compounds

Non-benzenoid aromatic compounds: There are aromatic compounds, which have structural units different from benzenoid type and are known as Non-benzenoid aromatics e.g. Tropolone,

Tropolone

Heterocyclic aromatic compounds: When atoms of more than one kind make up the ring in the compounds, they are known as heterocyclic compounds or heterocycles. In these compounds generally one or more atoms of elements such as nitrogen 'N', oxygen 'O', or sulphur 'S' are present. The atom other than that of carbon viz., N, O or S, present in the ring is called hetero atom. Heterocyclic compounds with five and six atoms in the ring are termed as five-membered, and six- membered heterocycles respectively.

Hydrocarbons can be further classified into four types on the basis of their structures. These are:

 Alkanes: Hydrocarbons that contain only C-C single bonds in their molecules are called alkanes. These include open chain as well as closed chain (cyclic) hydrocarbons. For example, ethane, propane cyclopentane.Alkanes are further divided into:  Open chain or acyclic (simple alkanes not having any closed chains). They have the general formula CnH2n+2. Examples are methane(CH4), propane(C3H8) and butane(C4H10).  Cycloalkanes or cyclic alkanes (having a closed chain or rings in their molecules). They have the general formula CnH2n. Examples are cyclopropane(C3H6) and cyclobutane(C4H8).  Alkenes: These are hydrocarbons that contain at least one carbon-carbon double bond. For example, ethene, but-2-ene, but-1-ene.  Alkynes: These hydrocarbons contain at least one carbon-carbon triple bond. For example, ethyne, propyne.  Arenes: These are hydrocarbons that contain at least one special type of hexagonal ring of carbon atoms with three double bonds in their alternate positions. The ring is called aromatic or benzene ring. For example, benzene, toluene, o-xylene. They also contain more than one benzene rings. For example, naphthalene (2 rings) and anthracene (3 rings).

Hydrocarbons can also be classified into:  Saturated hydrocarbons: Those that contain carbon-carbon single bonds e.g. alkanes  Unsaturated hydrocarbons; Those that contain carbon-carbon double or triple bonds e.g. alkenes(C=C), alkynes(C=C).

Classification of organic compounds based on functional groups

Functional group: A specific grouping of elements that is characteristic of a class of compounds, that give a compound certain physical and chemical properties. A functional group is a specific group of atoms or bonds within a compound that is responsible for the characteristic chemical reactions of that compound. The same functional group will behave in a similar fashion, by undergoing similar reactions, regardless of the compound of which it is a part. Functional groups also play an important part in organic compound nomenclature; combining the names of the functional groups with the names of the parent alkanes provides a way to distinguish compounds.

The atoms of a functional group are linked together and to the rest of the compound by covalent bonds. The first carbon atom that attach to the functional group is referred to as the alpha carbon; the second, the beta carbon; the third, the gamma carbon, etc. Similarly, a functional group can be referred to as primary, secondary, or tertiary, depending on if it is attached to one, two, or three carbon atoms.

Homologous Series

A homologous series is a group of organic compounds (compounds that contain C atoms) that differ from each other by one methylene (CH2) group. For example, methane, ethane, and propane are part of a homologous series. The only difference among these molecules is that they have different numbers of CH2 groups.Each member of a homologous series is called a homologue. For example, methane and ethane are homologues and belong to the same homologous series. They differ from each other by one CH2 group. The formula of methane is CH4 and the formula of ethane is C2H6.

Characteristics of Homologous Series

 Members of the series can be represented by a general formula. Ethanol, propanol and butanal have the same general formula, CnH2n+1OH  Successive members differ from each other by –CH2. The difference between propanol and ethanol is CH2 that has a relative molecular mass of 14.  Physical properties change regularly with increasing number of carbon atoms. The boiling points of alcohols increase from ethanol, propanol to butanol.  Members have similar chemical properties because they have same functional group. Ethanol, propanol and butanol undergo oxidatoin produce carboxylic acids.  Members of the homologous series can be prepared using the same method. Ethanol, propanol & butanol can be prepared by hydration of alkene.  All the members of homologous series contain same functional group.

The simplest example of a homologous is that of alkanes. Alkanes consist of carbon and hydrogen atoms only, in proportions according to the general formula: CnH2n+2 where the letter n represents the number of carbon atoms in each molecule of the compound. Hence the first 10 molecules in the homologous series of linear alkanes may be listed as follows:

Example of the Homologous Series of Alkanes, Structure: CnH2n+2 Number Name of Carbon Chemical Simple Structure Alkane atoms Formula (Molecular Diagram)

Methane 1 C H4

Ethane 2 C2H6

Propane 3 C3H8

Butane 4 C4H10

Pentane 5 C5H12

Hexane 6 C6H14

Heptane 7 C7H16

Octane 8 C8H18

Nonane 9 C9H20

Decane 10 C10H22

As can be seen in the case of the example of the homologous series of alkanes (right), the basic molecular structure of all members of the series takes the same form. The difference between

members of the homologous series depends on the value of n, which represents the number of carbon atoms in the chain.

Name of Series General chemical formula*

Alkanes CnH2n+2

Alkenes CnH2n

Alkynes CnH2n-2

1-bromoalkanes CnH2n+1Br (e.g. of Haloalkanes)

1-alcohols CnH2n+1OH (e.g. of Alcohols) = CnH2n+2O

Aldehydes CnH2nO

Carboxylic Acids CnH2nO2

Acid Chlorides CnH2n-1OCl

Amines CnH2n+1NH2 = CnH2n+3N

Amides CnH2n-1ONH2 = CnH2n+1ON

Nitriles CnH2n-3N

Nomenclature Of Organic Compounds

In chemical nomenclature, the IUPAC nomenclature of organic chemistry is a systematic method of naming organic chemical compounds as recommended by the International Union of Pure and Applied Chemistry (IUPAC). In order to name organic compounds you must first memorize a few basic names. These names are listed within the discussion of naming alkanes. In general, the base part of the name reflects the number of carbons in what you have assigned to be the parent chain. The suffix of the name reflects the type(s) of functional group(s) present on (or within) the parent chain. Other groups which are attached to the parent chain are called substituents.

Alkanes - saturated hydrocarbons: The names of the straight chain saturated hydrocarbons for up to a 12 carbon chain are shown below. The names of the substituents formed by the removal of one hydrogen from the end of the chain are obtained by changing the suffix -ane to -yl.

Number of Carbons Name 1 methane 2 ethane 3 propane 4 butane 5 pentane 6 hexane 7 heptane

8 octane 9 nonane 10 decane 11 undecane 12 dodecane

There are a few common branched substituents which you should memorize. These are shown below.

Here is a simple list of rules to follow. Some examples are given at the end of the list.

 Identify the longest carbon chain. This chain is called the parent chain.  Identify all of the substituents (groups appending from the parent chain).  Number the carbons of the parent chain from the end that gives the substituents the lowest numbers. When compairing a series of numbers, the series that is the "lowest" is the one which contains the lowest number at the occasion of the first difference. If two or more side chains are in equivalent positions, assign the lowest number to the one which will come first in the name.  If the same substituent occurs more than once, the location of each point on which the substituent occurs is given. In addition, the number of times the substituent group occurs is indicated by a prefix (di, tri, tetra, etc.).  If there are two or more different substituents they are listed in alphabetical order using the base name (ignore the prefixes). The only prefix which is used when putting the substituents in alphabetical order is iso as in isopropyl or isobutyl. The prefixes sec- and tert- are not used in determining alphabetical order except when compared with each other.  If chains of equal length are competing for selection as the parent chain, then the choice goes in series to:

a) the chain which has the greatest number of side chains. b) the chain whose substituents have the lowest- numbers. c) the chain having the greatest number of carbon atoms in the smaller side chain. d) the chain having the least branched side chains.  A cyclic (ring) hydrocarbon is designated by the prefix cyclo- which appears directly in front of the base name.

In summary, the name of the compound is written out with the substituents in alphabetical order followed by the base name (derived from the number of carbons in the parent chain). Commas are

used between numbers and dashes are used between letters and numbers. There are no spaces in the name.

Structural isomerism

Isomerism in organic chemistry is a phenomenon shown by two or more organic compounds having the same molecular formula but different properties due to difference in arrangement of atoms along the carbon skeleton (structural isomerism) or in space (Stereo isomerism). Structural isomerism can be six types; chain isomerism, positional isomerism, functional group isomerism, tautomerism, metamerism and ring-chain isomerism. All these types of isomers have same molecular formulae but different structural formulae.

Chain isomers can be defined as the structural isomers which have different number of carbon atoms in the parent chain. In other words we can say that they have different carbon skeleton but same molecular formulae. Such type of structural isomerism is common in organic compounds which have parent carbon atoms. In such molecules, chain can be straight, branched or chain with multiple side groups.

Position isomerism arises when there is a difference in the positions occupied by the substituent atoms or a group of atoms or due to the unsaturation occurring in the chain. When the position of the functional groups with respect to main chain atom changes, the phenomenon is called as position isomerism arises. For example; 2-cholropentane and 3-chloropentane are positional isomers. Different position of multiple bonds in the parent chain also results the formation of positional isomers such as 2-pentene and 3-pentene are positional isomers of each other.

Functional isomers are the structural isomers with same molecular formulae but different functional groups such as alcohols and ethers are functional isomers of each other. Similarly carboxylic acid and esters are functional isomers. Ethanol and Ethoxyether have same molecular formula but different structural formulae. Due to different bonding and functional group in these isomers, they show different chemical and physical properties. Straight-chain alkanes are functional isomers of cycloalkane.

Metamerism arises due to the unequal distribution of carbon atoms on either side of functional group. Metamers belongs to same homologous series.

Tautomers are isomers of a compound which differ only in the position of the protons and electrons. The carbon skeleton of the compound is unchanged. A reaction which involves simple proton transfer in an intramolecular fashion is called a tautomerism. Keto-enol tautomerism is a very common process, and is acid or base catalysed. Typically the 'keto' form of the compound is more stable, but in some instances the 'enol' form can be the more stable.

Ring chain isomerism is a special case of functional isomerism where a particular linear compound has an isomer with a cyclic structure with same molecular formula.

Hybridization in organic compounds

Hybridization is defined as the concept of mixing two atomic orbitals of similar energy to give equal number of hybrid degenerated orbitals. The atomic orbitals of the same energy level can only take part in hybridization and both full filled and half-filled orbitals can also take part in this process provided they have equal energy. During the process, the atomic orbitals of similar energy are mixed together such as the mixing of two ‘s’ orbitals or two ‘p’ orbital’s or mixing of an ‘s’ orbital with a ‘p’ orbital or ‘s’ orbital with a ‘d’ orbital.

From the ground state electron configuration, one can see that carbon has four valence electrons, two in the 2s subshell and two in the 2p subshell. The 1s electrons are considered to be core electrons and are not available for bonding. There are two unpaired electrons in the 2p subshell, so if carbon were to hybridize from this ground state, it would be able to form at most two bonds. Recall that energy is released when bonds form, so it would be to carbon's benefit to try to maximize the number of bonds it can form. For this reason, carbon will form an excited state by promoting one of its 2s electrons into its empty 2p orbital and hybridize from the excited state. By forming this excited state, carbon will be able to form four bonds. The excited state configuration is shown below:

There are three ways in which this mixing process can take place:

 the 2sorbital is mixed with all three 2porbitals. This is known as sp3 hybridization;  the 2s orbital is mixed with two of the 2p orbitals. This is known as sp2 hybridization;  the 2s orbital is mixed with one of the 2p orbitals. This is known as sp hybridization. sp3 hybridization- structure of methane

The valence electrons for carbon can now be fitted into the sp3 hybridized orbitals configuration. There were a total of four electrons in the original 2s and 2p orbitals. The s orbital was filled and two of the p orbitals were half filled. After hybridization, there is a total of four hybridized sp3 orbitals all of equal energy. By Hund’s rule, they are all half filled with an electron which means that there are four unpaired electrons. Four bonds are now possible.

Each of the sp3 hybridized orbital has the same shape – a rather deformed looking dumbbell. This deformed dumbbell looks more like a p orbital than an s orbital since more p orbitals were involved in the mixing process.

Each sp3 orbital will occupy a space as far apart from each other as possible by pointing to the corners of a tetrahedron. Here, only the major lobe of each hybridized orbital has been shown and the angle between each of these lobes is 109.5. This is what is meant by the expression tetrahedral carbon. The three dimensional shape of the tetrahedral carbon can be represented by drawing a normal line for bonds in the plane of the page. Bonds going behind the page are represented by a hatched wedge, and bonds coming out the page are represented by a solid wedge.

A half-filled sp3 hybridized orbital from one carbon atom can be used to form a bond with a half- filled sp3 hybridized orbital from another carbon atom. The major lobes of the two sp3 orbitals overlap directly leading to a strong σ bond. It is the ability of hybridized orbitals to form strong σ bonds that explains why hybridization takes place in the first place. The deformed dumbbell shapes allow a much better orbital overlap than would be obtained from a pure s orbital or a pure p orbital. A σ bond between an sp3 hybridized carbon atom and a hydrogen atom involves the carbon atom using one of its half-filled sp3 orbitals and the hydrogen atom using its half-filled 1s orbital.

Notes: Sigma bond (σ bond): A covalent bond formed by overlap of atomic orbitals and/or hybrid orbitals along the bond axis (i.e., along a line connected the two bonded atoms).

Hund's Rule: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied and each of the electrons in singly occupied orbitals has the same spin (to maximize total spin). sp2 hybridization- structure of ethene

2 In sp hybridization, the s orbital is mixed with two of the 2p orbitals (e.g. 2px and 2pz) to give three 2 sp hybridized orbitals of equal energy. The remaining 2py orbital is unaffected. The energy of each hybridized orbital is greater than the original s orbital but less than the original p orbitals. The remaining 2p orbital (in this case the 2py orbital) remains at its original energy level.

For carbon, there are four valence electrons to fit into the three hybridized sp2 configuration orbitals and the remaining 2p orbital. The first three electrons are fitted into each of the hybridized orbitals

according to Hund’s rule such that they are all half- filled. This leaves one electron still to place. 2 There is a choice between pairing it up in a half-filled sp orbital or placing it into the vacant 2py orbital. The usual principle is to fill up orbitals of equal energy before moving to an orbital of higher energy. However, if the energy difference between orbitals is small (as here) it is easier for the 2 electron to fit into the higher energy 2py orbital resulting in three half-filled sp orbitals and one half- filled p orbital. Four bonds are possible.

2 Geometry The 2py orbital has the usual dumbbell shape. Each of the sp hybridized orbitals has a deformed dumbbell shape similar to an sp3 hybridized orbital. However, the difference between the sizes of the major and minor lobes is larger for the sp2 hybridized orbital. The hybridized orbitals and the 2py orbital occupy spaces as far apart from each other as possible. The lobes of the 2py orbital occupy the space above and below the plane of the x and z axes. The three sp2 orbitals (major lobes shown only) will then occupy the remaining space such that they are as far apart from the 2py orbital and from each other as possible. As a result, they are all placed in the x–z plane pointing toward the corner of a triangle (trigonal planar shape). The angle between each of these lobes is 120.

sp2 Hybridization results in three half-filled sp2 hybridized orbitals which form a trigonal planar shape. The use of these three orbitals in bonding explains the shape of an alkene, for example ethene (H2C==CH2). As far as the C–H bonds are concerned, the hydrogen atom uses a half-filled 1s orbital to form a strong σ bond with a half filled sp2 orbital from carbon. A strong σ bond is also possible between the two carbon atoms of ethene due to the overlap of sp2 hybridized orbitals from each carbon.

Ethene is a flat, rigid molecule where each carbon is trigonal planar. We have seen how sp2 hybridization explains the trigonal planar carbons but we have not explained why the molecule is rigid and planar. If the σ bonds were the only bonds present in ethene, the molecule would not remain planar since rotation could occur round the C–C σ bond. Therefore, there must be further bonding which ‘locks’ the alkene into this planar shape. This bond involves the remaining half-filled 2py orbitals on each carbon which overlap side-on to produce a pi (p) bond), with one lobe above and one lobe below the plane of the molecule. This π bond prevents rotation round the C–C bond since the π

bond would have to be broken to allow rotation. A π bond is weaker than a σ bond since the 2py orbitals overlap side-on, resulting in a weaker overlap. The presence of a π bond also explains why alkenes are more reactive than alkanes, since a π bond is more easily broken and is more likely to take part in reactions.

sp hybridization- structure of ethyne

......

“Curved Arrow Formalism” or Pushing Electrons

Carbon and other second row elements such as B, N, O, and F follow the octet rule, i.e. they try to have the sum of bonding electrons and electrons in lone pairs around them equal to 8. For the first row, hydrogen tries to have 2 electrons.

Sometimes molecules have atoms that are short of an octet by one or more electron pairs – they tend to be very reactive. For example: 1. H+ has 0 electrons and it needs 2, thus it is deficient by 2.

BF3 is an electron deficient compound. The boron atom in boron tri-fluoride has 6 electrons, and it needs 8. Thus it is deficient by 2 electrons. One additional lone pair is needed to fill its octet.

Methyl cation has 6 electrons, and it needs 8, thus is deficient by 2

 Electron deficient compounds, which can behave a electron pair acceptors are Lewis acids. A species that donates an electron pair is a Lewis base. The reaction above is called Lewis acid/ Lewis base association reaction.

Each curved arrow with two barbs on the head represents the shift of one electron pair. The curved arrows shows the direction of electron flow. The tail shows the electron origin, and always come from an electron source, usually a lone pair or bonding pair from a sigma or pi bond. The head of the arrow indicates the electron pair destination, either as a new lone pair or a new bond. If the arrowhead points to another atom, that atom must either have an open octet and thus be able to accept the electron pair, or have an electron pair that can be displaced by the incoming electron pair. In general, the arrow starts at the Lewis base end and ends at the electron deficient species (the Lewis acid). Electrons never flow from atoms which are electron-poor to atoms which are electron-rich, so a curved arrow will never point from an atom with a positive charge to an atom with a negative charge.

 Curved arrows can help one draw resonance structures.

 Conjugate Acid-Base Pairs

Fission of Covalent Bond: This bond can be broken either by :

1. Homolytic fission (homolysis) 2. Heterolytic fission (heterolysis)

Heterolysis : When a bond breaks such that one fragment takes away both electrons of the bond results in the formation of charged species, leaving the other fragment with an empty orbital, this kind of cleavage is called heterolysis (Greek: hetero, different, + lysis, loosening or cleavage). Heterolysis produces charged fragments or ions and is termed an ionic reaction. The broken bond is said to have cleaved heterolytically:

Heterolysis of a bond normally requires that the bond be polarized:

Polarization of a bond usually results from differing electronegativities of the atoms joined by the bond. The greater the difference in electronegativity, the greater is the polarization. In the instance just given, atom B is more electronegative than A.

Even with a highly polarized bond, heterolysis rarely occurs without assistance. The reason: Heterolysis requires separation of oppositely charged ions. Because oppositely charged ions attract each other, their separation requires considerable energy. Often, heterolysis is assisted by a molecule with an unshared pair that can form a bond to one of the atoms:

Formation of the new bond furnishes some of the energy required for the heterolysis.

Carbocation : A positively charged carbon ion is called carbocation. Carbocations are highly unstable and reactive species. Carbocation can be primary, secondary or tertiary depending upon how many alkyl groups are attached to the positively charged carbon. The stability of carbocations follows the order : tertiary > secondary > primary due to +I effect and hyperconjugation.

+ + + + Stability of carbocation : (CH3)3C > (CH3)2CH > CH3CH2 > CH3

Carbanion: If its carbon that gets away with shared pair of electrons, then the carbon is termed as carbanion. The order of stability of carbanion is exact opposite of that of carbocation due to inductive effect.

− − − − Stability of carbanion : (CH3)3C < (CH3)2CH < CH3CH2 < CH3

Homolysis : When a bond breaks so that each fragment takes away one of the electrons of the bond results in the formation of neutral species, this process is called homolysis (Greek: homo, the same, + lysis). Homolysis produces fragments with unpaired electrons called radicals.

Reactions that involve homolytic cleavage are called free radical reactions. Homolytic fission usually occurs in non-polar bonds. Conditions that generally favour homolytic fission are :

 High temperature  Ultraviolet light  Presence of peroxides

. Alkyl Free Radicals: Free radicals of carbon such as CH3 are known as alkyl free radicals. They can be primary, secondary or tertiary.The stability of alkyl free radicals increases in the order : primary < secondary < tertiary. The main reason for this order is hyperconjugation.

Nucleophiles and Electrophiles

The term is broken down into the word “nucleo” which refers to the nucleus and the Latin word “phile” which means loving. It simply means nucleus loving. Nucleophiles are rich in electrons and, as thus, donate electron pairs to electrophiles to form covalent bonds in chemical reactions. These substances are best noticed with lone pairs, pi bonds and negative charges. Ammonia, iodide and hydroxide ions are examples of nucleophile substances. The nucleophilic center in a compound is detected with the most electronegative atom. Consider ammonia NH3; the nitrogen is more electronegative and thus draws electrons to the center. The compound has high electron density and, when reacting with an electrophile, say water, it donates electrons.

A nucleophile is a reactant that provides a pair of electrons to form a new covalent bond. This is the exact definition of a Lewis base. In other words, nucleophiles are Lewis bases. Nucleophiles all have pairs of electrons to donate, and tend to be rich in electrons.

The word “electro” is from electrons and the Latin word “phile” refers to “loving”. In simple terms, it means electrons-loving. It is a reagent that is characterized with a low density of electrons in its valance shell, and, therefore, reacts with a high-density molecule, ion or atom to form a covalent bond. Hydrogen ion in acids and methyl-carbocation are examples of electrophilic substances. They are electron deficient.

An electrophile is easily detected by a positive charge or neutral charge with empty orbitals (not satisfying the octet rule). Electrons move from an area of high density to the one with low density, and unlike charges attract each other. This theory explains the attraction of electrons by the electron- deficient electrophile atoms, molecules or ions. By definition, an electrophile is interchangeably called a Lewis acid as it accepts electrons in line with the definition of the acid. In general, an electrophile is identified by a partial positive charge as in hydrogen chloride, a formal positive charge as in methyl carbocation or vacant orbitals. Polarized neutral molecules such as acyl halides, carbonyl compounds, and alkyl halides are typical examples of electrophiles.

Summary of Electrophile Verses Nucleophile

 An electrophile is an electron-deficient atom,, ion or molecule while the nucleophile is an electron-rich atom, molecule or ion  An electrophile can be positively or neutrally charged while the nucleophile can be negatively or neutrally charged  An electrophile is called the Lewis acid and the nucleophile is called the Lewis base  An electrophile accepts electrons and donates protons while a nucleophile donates electrons and accepts protons.

Difference between electrophile and nucleophile

In order to distinguish between good or poor nucleophiles when SN reactions are carried out the better form of nucleophile forms the main product.

 Nucleophilicity measures the ability of the nucleophile to make an electron pair available to the electrophile.  Electrophile and nucleophile differ in mechanism of a reaction and depends upon the nature of double bond and attacking species and also on reaction conditions. In Electrophilic mechanism the electrophile adds to double bond first, while nucleophile adds to double bonds first in nucleophilic mechanism.  Nucleophile is considered as lewis base while electrophile is best compared to a lewis acid.  A nucleophile donates electron pair and an electrophile accepts an electron pair.  A nucleophiles are normally measured on a kinetic scale, while the electrophiles are measured on an equilibrium basis.  Charged nucleophiles are more nucleophilic than uncharged species with same central atom, the effect is with halogens I- > Br > Cl- and is due to polarizability of larger ions. Electrophiles act exactly the opposite to these.  Electrophile is involved in rate determining step of SN1 reaction, where the rate order is first order. The nucleophile is involved in the fast product forming step which has no relevance to kinetic part.

Inductive effect

The polarization of a σ bond due to electron withdrawing or electron donating effect of adjacent groups or atoms is called inductive effect. Inductive effect occurs due to the polarisation of σ bonds within a molecule or ion. This is typically due to an electronegatvity difference between the atoms at either end of the bond.The more electronegative atom pulls the electrons in the bond towards itself creating some bond polarity. The inductive effect is a distance-dependent phenomenon. It is transmitted through the sigma bonds. The magnitude of inductive effect decreases while moving away from the groups causing it. For example the O-H and C-Cl bonds in the following:

If the electronegative atom is connected to a chain of carbon atoms, then the positive charge is relayed to the other carbon atoms. With its positive charge, exerts a pull on the electrons, but the pull is weaker than it is between on. The effect rapidly dies out and is usually not significant after the 4th carbon atom.

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.

1) Negative inductive effect (-I): The electron withdrawing nature of groups or atoms is called as negative inductive effect. It is indicated by -I. Following are the examples of groups in the decreasing

+ NH3 > NO2 > CN > SO3H > CHO > CO > COOH > COCl > CONH2 > F > Cl > Br > I > OH > OR > NH2 > C6H5 > H

2) Positive inductive effect (+I): It refers to the electron releasing nature of the groups or atoms and is denoted by +I. Following are the examples of groups in the decreasing order of their +I effect.

-C(CH3)3 > -CH(CH3)2 > -CH2CH3 > -CH3 > -H

Why alkyl groups are showing positive inductive effect?Though the C-H bond is practically considered as non-polar, there is partial positive charge on hydrogen atom and partial negative charge on carbon atom. Therefore each hydrogen atom acts as electron donating group. This cumulative donation turns the alkyl moiety into an electron donating group.

Stability of carbonium ions: The stability of carbonium ions increases with increase in number of alkyl groups due to their +I effect. The alkyl groups release electrons to carbon, bearing positive charge and thus stabilize the ion.

In a simple alkyl carbocation, the positive C attracts the electrons in the σ bonds connected to that center towards itself and therefore away from the atom at the other end of the σ bond. Electrons in C- C bonds are more readily polarised than those in a C-H bond. Therefore, alkyl groups are better at stabilising C+ than H atoms.

The order of stability of carbonium ions is :

Stability of free radicals: In the same way the stability of free radicals increases with increase in the number of alkyl groups.

Thus the stability of different free radicals is:

Stability of carbanions: However the stability of carbanions decreases with increase in the number of alkyl groups since the electron donating alkyl groups destabilize the carbanions by increasing the electron density.

Thus the order of stability of carbanions is:

Acidic strength of carboxylic acids and phenols: The electron withdrawing groups (-I) decrease the negative charge on the carboxylate ion and thus by stabilizing it. Hence the acidic strength increases when -I groups are present. However the +I groups decrease the acidic strength.

E.g.

i) The acidic strength increases with increase in the number of electron withdrawing Fluorine atoms as shown below.

CH3COOH < CH2FCOOH < CHF2COOH < CF3COOH

ii) Formic acid is stronger acid than acetic acid since the –CH3 group destabilizes the carboxylate ion.

On the same lines, the acidic strength of phenols increases when -I groups are present on the ring. e.g. The p-nitrophenol is stronger acid than phenol since the -NO2 group is a -I group and withdraws electron density whereas the para-cresol is weaker acid than phenol since the -CH3 group shows positive (+I) inductive effect. Therefore the decreasing order of acidic strength is:

Basic strength of amines: The electron donating groups like alkyl groups increase the basic strength of amines whereas the electron withdrawing groups like aryl groups decrease the basic nature. Therefore alkyl amines are stronger Lewi bases than ammonia, whereas aryl amines are weaker than ammonia.

Thus the order of basic strength of alkyl and aryl amines with respect to ammonia is:

CH3NH2 > NH3 > C6H5NH2

Reactivity of carbonyl compounds: The +I groups increase the electron density at carbonyl carbon. Hence their reactivity towards nucleophiles decreases. Thus formaldehyde is more reactive than acetaldehyde and acetone towards nucleophilic addition reactions.

Thus the order of reactivity follows:

Mesomeric Effect

The withdrawal effect or releasing effect of electrons attributed to a particular substituent through the delocalization of π or pi-electrons that can be seen by drawing various canonical structures is called as resonance effect or mesomeric effect.

Types of Resonance Effects: There are two types of Resonance effects namely positive resonance effect and negative resonance effect.

 Positive Resonance Effect- Positive resonance effect occurs when the groups release electrons to the other molecules by the process of delocalization. The groups are usually denoted by +R or +M. In this process, the molecular electron density increases. For example- -OH, -SH, -OR,-SR.

 Negative Resonance Effect- Negative resonance effect occurs when the groups withdraw the electrons from other molecules by the process of delocalization. The groups are usually denoted by -R or -M. In this process, the molecular electron density is said to decrease. For example- -NO2, C=O, -COOH, -C≡N.

The para- and ortho- isomers each have a significant resonance form where the nitrogen lone pair donates into the ring and a pi bond breaks in the nitro group.The resulting NH2+ group is non-basic (no free lone pair). The meta isomer lacks this resonance form, which explains why it is the most basic of the three isomers. The ortho has the nitro group closer to the amine than the para, which explains why it is less basic. (inductive effect)

Hyperconjugation or no-bond resonance or Baker-Nathan effect

Hyperconjugation is the stabilising interaction that results from the interaction of the electrons in a σ- bond (usually C-H or C-C) with an adjacent empty or partially filled p-orbital or a π-orbital to give an extended molecular orbital that increases the stability of the system.

For example, consider the ethyl carbocation (1), which is shown in a specific conformation (2) 3 below. In 2, the empty p orbital on C1 and the sp -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. 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.

The stabilisation arises because the orbital interaction leads to the electrons being in a lower energy orbital.

The ethyl cation has 3 C-H σ-bonds that can be involved in hyperconjugation. The more hyperconjuagtion there is, the greater the stabilisation of the system. For example, the t-butyl cation + has 9 C-H σ-bonds that can be involved in hyperconjugation. Hence (CH3)3C is more stable than + CH3CH2 .

In propene molecule, Hyper conjugation arise due to partial overlap of sp3-s sigma bond orbital and the empty p-orbital or pi-bond orbital of an adjacent carbon atom. Here one of the carbon-hydrogen bonds of methyl group can lie in the plane of pi-bond orbital, hence partial overlap with pi-bond orbital. This results the delocalization of pi-electrons and increase the stability of molecule.

In resonating structures of propene, there is no bond between carbon and hydrogen ion, therefore hyper conjugation is also called as no bond resonance. As the number of methyl group bonded in double bonded carbon atom increases, the possibility of Hyper conjugation increases which results more stability. That is the reason, more substituted alkene are more stable than less substituted alkene. The increasing order of stability of some alkenes is as follow.

(CH3)2C=C(CH3)2 > (CH3)2C=CHCH3 > CH3CH=CHCH3 > CH3CH=CH2

Similarly free radicals get stabilized through hyper conjugation. Like carbocation, the sigma electrons of carbon-hydrogen bonds of methyl group next to carbon atom contain odd electron interact with p-orbital having odd electron. As the number of alpha carbon-hydrogen bond increases, contributing structures increases results more stability.

In nitromethane, the nitrogen-oxygen pi-bond interacts with alpha carbon-hydrogen bond and show hyper conjugation.

The presence of triple bond between carbon-carbon or carbon-nitrogen also show hyper conjugation with alpha carbon-hydrogen bond like in acetonitrile and propyne molecule.

In Toluene, the carbon - hydrogen sigma bond interacts with pi-bond of aromatic ring to form four contributing structures of toluene.

Stability of unsaturated hydrocarbons :

 The stability of unsaturated hydrocarbons like nitriles, alkenes effects with hyper conjugation. The more possible contributing structures in hyper conjugation increase the stability of molecule.  Since hyper conjugation mainly involves alpha carbon-hydrogen sigma bond and pi-electrons, therefore as the number of alpha sigma bonds increases, hyper conjugation increases.  For example, 2-butene consists of six alpha carbon-hydrogen sigma bonds while there are only 2 carbon-hydrogen bonds next to double bonded carbon atom in 1-butene.  Hence 2-butene shows six contributing structures while 1-butene shows only two which make 2-butene more stable compare to 1-butene.  This rule is applicable on other alkenes also. Hence as the number of alkyl group on double bonded carbon atoms increases, hyper conjugation increases which stabilized the molecule.

Or more substituted alkenes are more stable than less substituted alkenes.

Stability of reaction intermediates: Intermediates like carbocations and free radicals also show hyper conjugation due to the presence of empty p-orbitals or p-orbital with odd electrons. The stability of carbocation and free radical get affected with the number of contributing structures. For example; in primary carbocation like ethyl carbocation there are three alpha carbon-hydrogen bonds which delocalized in to the empty p-orbital of C+. While in secondary carbocation like iso-propyl carbocation, there are six alpha carbon-hydrogen bonds and in tertiary carbocation like tert-butyl carbocation, there are nine alpha carbon-hydrogen bonds. Hence the increasing order of stability of carbocations can be given as:

Primary < Secondary < Tertiary carbocation

Similarly the increasing order of stability of free radical is primary < secondary < tertiary free radical.

Dipole moment & bond length: Hyper conjugation induces polarity in molecule which affects the dipole moment and bond length of molecule. As the polarity increases, dipole moment increase and bond length decreases.

Orientation effect of electrophilic substitution reactions on benzene ring: The methyl group of toluene shows positive inductive effect and hyper conjugation which releases electrons towards aromatic ring. These two effects increases the electron density on aromatic ring atortho and para- positions, therefore coming electrophile effectively attack on these two positions to give ortho and para-substituent products.

Anomeric effect: Anomeric Effect may define as the tendency of anomeric substituents to prefer an axial configuration over equatorial position. The high stability of αα-methyl glucoside compare to ββ- anomer can be explained by using hyper conjugation. As αα-anomer can show hyper conjugation of lone pair of oxygen atom with axial methyl group which is not possible in ββ-anomer due to equatorial position.

Electromeric effect

Electromeric effect is a temporary effect and may be defined as the complete transfer of shared pair of pi electrons of multiple bonds to one of the atoms in presence of an attacking reagent.

For example, consider the carbonyl group, >C=O, present in aldehydes and ketones. When a negatively charged reagent say approaches the molecule seeking positive site, it causes instantaneous shift of electron pair of carbonyl group to oxygen (more electronegative than carbon). The carbon thus becomes deprived of its share in this transferred-pair of electrons and acquires positive charge. In the meanwhile oxygen takes complete control of the electron pair and becomes negatively charged. Therefore, in the presence of attacking reagent, one bond is lost and this negatively charged attacking reagent links to the carbon having positive charge. This phenomenon of movement of electrons from one atom to another at the demand of attacking reagent in multibonded atoms is called electromeric effect, denoted as E effect. The electromeric shift of electrons takes place only at the moment of reaction.

+E effect: If the attacking species is an electrophile, the π electrons are transferred towards the positively charged atom. This is the +E effect. An example is the protonation of ethene. When the H⁺ comes near the double bond, the bond is polarized towards the proton.

–E Effect : If the attacking reagent is a nucleophile, the electrons are transferred away from the attacking reagent and into the π system. This is the –E Effect.

Steric effect

The word steric is derived from ‘stereos’ meaning space. So this effect is manifested when two or more groups or atoms come in close proximity to each other precisely within each other’s van der Waals radii and result in a mutual repulsion. This makes the molecule unstable. The situation can be compared to a crowded bus or train where each passenger stands touching the other one and there is collision, one steps on the other’s feet, hits one another with elbows and so on and so forth. It’s clearly not a very pleasant scenario! It is the same things with molecules. Sheer bulk of the atoms or groups and their proximity can have serious implications in deciding the reactivity of an entity. The usual physical clash between groups, almost always is accompanied by an electronic component as well. This is called stereoelectronic effect, which is not the same as the electronic effects discussed above and does not carry have an effect on some other part of the molecule like inductive and resonance effects. When the two atoms get to close, into each other’s van der Waal’s radii, the electron cloud surrounding each atom repel each other leading to a lot of destabilization. Steric effect affects different properties of molecules, like acidity, basicity and general reactivity.

A substitution reaction on a halide by a hydroxide does not work in this case because of steric hindrance. The second figure represents the same reaction, with spheres replacing the alkyl groups to show the spatial perspective.

Reaction intermediates

A reaction intermediate is transient species within a multi-step reaction mechanism that is produced in the preceding step and consumed in a subsequent step to ultimately generate the final reaction product. Synthetic intermediate are stable products which are prepared, isolated and purified and subsequently used as starting materials in a synthetic sequence. Reactive intermediate, on the other hand, are short lived and their importance lies in the assignment of reaction mechanisms on the pathway from the starting substrate to stable products. These reactive intermediates are not isolated, but are detected by spectroscopic methods, or trapped chemically or their presence is confirmed by indirect evidence.

1. Carbocation

A carbocation is molecule having a carbon atom bearing three bonds and a positive formal charge. Carbocations are generally unstable because they do not have eight electrons to satisfy the octet rule.

A carbocation forms bonds with three other groups and is positively charged. With three bond pairs and zero lone pairs it has trigonal planar geometry by VSEPR theory, and trigonal planar is characteristic of sp2 hybridisation. If you look at it another way, carbon has 4 electrons in the second shell (electronic configuration 1s22s22p2). For a carbocation to be formed, the carbon atom must lose one electron. A carbocation has only 3 electrons and it has to form 3 sigma bonds with its 2 substituents. Therefore, it needs 3 hybridised orbitals, so it is sp hybridised. The pz-orbital that is not utilised in the hybrid set is empty and is often shown bearing the positive charge since it represents the orbital available to accept electrons.

A carbocation in which the open valence shell carbon is not bonded to any carbon groups is termed a methyl carbocation. A primary carbocation (1o carbocation) is one in which there is one carbon group attached to the carbon bearing the positive charge. A secondary (2o) carbocation is one in which there are two carbons attached to the carbon bearing the positive charge. Likewise, a tertiary (3o) carbocation is one in which there are three carbons attached to the carbon bearing the positive charge.

When the carbon bearing the positive charge is immediately adjacent to a carbon-carbon double bond, the carbocation is termed an allylic carbocation. The simplest case (all R = H) is called the allyl carbocation.

When the carbon bearing the positive charge is immediately adjacent to a benzene ring, the carbocation is termed a benzylic carbocation. The simplest case is called the benzyl carbocation.

When the carbon bearing the positive charge is part of an alkene, the carbocation is termed a vinylic carbocation. The simplest case is called the vinyl carbocation. Note that the carbon bearing the positive charge has two attachments and thus adopts sp hybridization and linear geometry

. When the carbon bearing the positive charge is part of a benzene ring, the carbocation is termed an aryl carbocation. The simplest case is called the phenyl carbocation.

Carbocations are stabilized by neighboring carbon atoms: The number of carbon groups attached to the carbon carrying the positive formal charge is a major factor which decides the stability of a carbocation. Bonding electrons in sigma bonds adjacent to the open valence shell carbon can delocalize the positive charge to some extent by overlapping with the unoccupied p orbital of the carbocation.

This phenomenon is termed hyperconjugation. Since the overlap supplies electron density to the electrondeficient carbocation carbon, we predict that increasing the number of hyperconjugative increases carbocation stability. Extending this idea, we predict that increasing the number of bonds adjacent to the carbocation by increasing the number of alkyl groups attached to the carbocation carbon results in an increase in carbocation stability. For example, a tertiary carbocation should be more stable than a secondary carbocation. The order of stability of carbocation is 3° > 2° > 1°.

Carbocations are stabilized by neighboring carbon-carbon multiple bonds: Carbocations adjacent to another carbon-carbon double or triple bond have special stability because overlap between the empty p orbital of the carbocation with the p orbitals of the π bond allows for charge to be shared between multiple atoms. This effect, called “delocalization” is illustrated by drawing resonance structures where the charge “moves” from atom to atom. This is such a stabilizing influence that even primary carbocations – normally very unstable – are remarkably easy to form when adjacent to a double bond.

Carbocations are stabilized by adjacent lone pairs: The key stabilizing influence is a neighboring atom that donates a pair of electrons to the electron-poor carbocation. Note here that this invariably results in forming a double bond (π bond) and the charge will move to the atom donating the electron pair. Hence this often goes by the name of “π donation”.The strength of this effect varies with basicity, so nitrogen and oxygen are the most powerful π donors. Strangely enough, even halogens can help to stabilize carbocations through donation of a lone pair.

2. Carbanions

A carbanion is an anion in which carbon has an unshared pair of electrons and bears a negative charge Heterolytic cleavage of a bond, where carbon retains both the shared pair of electrons results into the formation of a carbanion (i.e, carbon atom having negative charge).In these species, carbon atom carrying negative charge has eight electrons in the valence shell- six from three covalent bonds and two from lone pair of electrons.

The basic form of carbanion is methide ion (CH3-) also commonly known as methyl carbanion. It is carbanion of methane (CH4) formed by loss of a proton.

Hybridisation of carbanion is sp3 with a lone pair of electrons being carried in one of the sp3 hybrid orbital with other 3 are responsible for the three sigma bonds a canbanions carries. Geometry is trigonal pyramidal but shape of the carbanion is pyramidal with a lone pair of electrons on carbon towards upward direction.

Factors which can stabilize or disperse the negative charge on carbon will stabilize a carbanion. The stability of carbanion depends on the following factors:  Inductive effect  Extent of conjugation of the anion  Hybridization of the charge-bearing atom  Aromaticity

Inductive effect: If the groups attached to carbanion are electron releasing in nature they will increase the negative charge on carbon and thus destabilize it. However, electronegative atoms or electron withdrawing groups adjacent to the negatively charged carbon will stabilize the carbanion. The alkyl groups are electron releasing in nature due to inductive effect (+I). More the number of alkyl groups attached lesser will be the stability. Carbanions prefer a lesser degree of alkyl substitution. Therefore the order of stability order of alkyl carbanion is methy l> 1o > 2o > 3o.

Presence of electronegative atoms (F, Cl, Br) or electron withdrawing groups (NO2, CN, COOH, CO) close to the negatively charged carbon will stabilize the charge. Thus more the number of such groups in a carbanion greater will be the stability.

Extent of conjugation of the anion: If negatively charged carbon is in conjugation with a double bond the resonance effects will stabilize the anion by spreading out the charge by rearranging the electron pairs.

Hybridization of the charge bearing atom: Stability of anion will depend upon the s character of carbanion i.e. more the s character, higher will be the stability of anion. The percentage s character in the hybrid orbitals is as follows: sp(50%)> sp2 (33%)>sp3 (25%).

Orbital with greater s character is more close to the nucleus and feels more nuclear charge. The sp hybridized atoms (50% s character) are more electronegative than sp2 (33% s character) and sp3 (25% s character). The distance of lone pair and nucleus is less if the lone pair is sp hybridized than in a sp2 hybrid orbital. Since, it is more favourable for the negative charge of an anion to be in an orbital close to the positively charged nucleus. Therefore sp hybrid anion is more stable than sp2

Aromaticity: In some carbanions, the lone pair of electrons of the negative charge is involved in delocalization to add on to the aromatic character of the molecule which gives them extra stability. For example, in cyclopentadienyl anion there are 6 π electron and thus it obeys Huckel rule, (4n+2) π electron. This anion is stabilized by aromatization. Cyclooctatetraene on reaction with potassium gets converted to cyclooctatetraenyldianion potassium salt. This is 10 π electron system which is stable due to aromaticity.

3. Free Radical

A free radical is a species containing one or more unpaired electrons. Free radicals are electron- deficient species, but they are neutral species, so their chemistry is very different from the chemistry of even-electron electron-deficient species such as carbocations and carbenes. The alkyl radical (·CR 3) is a seven-electron, electron-deficient species. The geometry of the alkyl radical is considered to be a shallow pyramid, somewhere between sp2 and sp3 hybridization,

In carbon free radical,there are total three bond pairs. This means carbon satisfies its three valency with other atoms and a single electron is left there. So, here in carbon free radical, only three bonds are present and there is no lone pair. Instead of lone pair,there is a single electron. so it is sp2 hybridised. It has a planar structure and bond angle is 120°. The unpaired electron goes in a unhybridized pz-orbital.

Both alkyl radicals and carbocations are electron-deficient species, and structural features that stabilize carbocations also stabilize radicals. Alkyl radicals are stabilized by adjacent lone-pair- bearing heteroatoms and by π bonds, just as carbocations are, and the order of stability of alkyl radicals is 3° > 2° > 1°.

A C atom surrounded by seven electrons is not as electron-deficient as a C atom surrounded by six electrons, so alkyl radicals are generally not as high in energy as the corresponding carbocations. Thus, the very unstable aryl and 1° alkyl carbocations are almost never seen, whereas aryl and 1° alkyl radicals are reasonably common.

 Free radicals decrease in stability as the % of s-character in the orbital increases [i.e. as the half-empty orbital becomes closer to the nucleus]. For that reason, free radical stability decreases as the atom goes from sp3 > sp2 > sp.

 Across a row of the periodic table, free radicals decrease in stability as the electronegativity increases

 Free radicals increase in stability going down a column of the periodic table,

F• < Cl• < Br• < I• since the electron-deficient orbital is spread out over a greater volume.

 Free radicals adjacent to an electron-withdrawing group are less stable, since in effect, electron-density is being taken away from what is already an electron deficient species. [Watch out, however – this only applies to electron withdrawing groups that cannot donate a pair of electrons, like CF3 or CN.]