IntroductionOrganic chemistry is a branch of chemistry which involves scientific study of different aspects of organic compounds such as structure, properties, and their reactions with other inorganic and organic compounds. Initially, those compounds were considered organic which were obtained from living sources i.e. plants and animals. They were considered to be formed by some vital force and it was believed that these could not be prepared by man using synthetic methods. Then, some compounds were prepared in laboratory and the theory lost its significance. Then, scientists coined another definition of organic compounds according to which compounds of carbon are organic compounds. This also did not serve the purpose as many compounds like carbides, carbonates, bicarbonates in addition to carbon dioxide, carbon monoxide, carbon tetrachloride, etc. are established as inorganic compounds. Finally, after analysing a large number of organic compounds, the most widely accepted definition came and according to it organic compounds are either hydrocarbons or their derivatives. Organic compounds contain two main elements which are carbon and hydrogen. Compounds formed by combination of these two elements are known as hydrocarbons. We have a large number of families of hydrocarbons. Some elements like oxygen, nitrogen, halogens and sulphur (also called heteroatoms) are commonly present in organic compounds. Organic compounds containing these elements are called derivatives of hydrocarbons (they are obtained by replacing one or more hydrogens from a hydrocarbon). In most of the organic compounds carbon, hydrogen and heteroatoms are linked to each other with the help of covalent bond (a bond formed by sharing of electrons by the combining atoms). There are three main types of covalent bond which we come across in organic compounds. These are single, double or triple covalent bonds. Single covalent bond is formed when combining atoms share one electron each. Similarly double and triple covalent bonds are formed when combining atoms share two and three electrons each respectively.Carbon is tetravalent (it forms four bonds), carbon-carbon bond is very strong and due to its ability to form all types of covalent bonds (single, double and triple), organic compounds may be straight chain, branched or cyclic which may contain chains and rings of carbon atoms linked to each other or some heteroatom may also be present.For studying different aspects of organic compounds, it is important to study about their structures. This is due to the reason that physical and chemical properties of the organic compounds are directly related to their structures. In simpler words if we know the structure of a molecule of an organic compound, we can easily understand its properties. Structure of a molecule means relative position of the atoms with respect to each other as nature of bonding between them.So, covalent bond is the main type of bonding that, results in formation of vey large number of various types of organic compounds. Formation of covalent bond is explained by many theories. Most widely accepted concept is orbital overlap concept. According to this concept, covalent bond is formed by overlapping of orbitals of combining atoms. In simple molecules like hydrogen it is believed that pure atomic orbitals of combining atoms overlap. However, if we apply the same concept to higher atoms, we fail to explain many observations related with structure and properties of the covalent molecules (compounds). To justify the structure and properties of covalent molecules or compounds, it was suggested by Linus Pauling that combining atoms do not form covalent bond by using pure atomic orbitals. They rather use orbitals obtained by mixing of atomic orbitals and these orbitals were called hybrid orbitals and the phenomenon was called hybridisation.Here, we will study hybridisation in detail along with various types of covalent bonds and their effect on the properties of covalent molecules (compounds). Our focus will be on organic compounds. The concept of hybridisation can also be applied on inorganic covalent compounds which we will discuss in the book of inorganic chemistry.Hybridisation.The simplest hydrocarbon is Methane (CH4), which has a tetrahedral structure. The four hydrogens attached to carbon atom are arranged as far apart in space as possible and are symmetrically arranged around the carbon atom. All the four bonds are equivalent. This structure cannot be explained using concept of overlapping of pure atomic orbitals. Similar are the cases of CCl4, ethane, CBr4 etc.

The electronic configuration of carbon (atomic number = 6) is 1s2 2s2 2px1 2py1 2pz0. According to overlapping theory it can make two single covalent bonds with two hydrogen atoms. But, in methane carbon forms four bonds. Formation of four bonds can be explained by considereing involvment of carbon in bond in excited state. Excited state can be achieved by pomoting one electron from 2s orbitral to 2pz orbital. The configuration of carbon in excited state is 1s2 2s1 2px1 2py1 2pz1. This configuration justifies formation of four covalent bonds or tetracovalency of carbon.

If simple and pure atomic orbitals of carbon overlap with half filled atomic orbital of hydrogen, we will get two types of bonds. First type of bond is formed by overlap of 2s orbital of carbon with 1s orbital of hydrogen. Second type of bond is formed by the overlap of 2p orbital with 1s orbital of carbon. these twotypes of bonds are different from each other. However, in reality we have all the four bonds equivalent. Equivalence in all the bonds in methane can be explained on the basis of hybridisation. It was staed that before participating in bonding the pure atomic orbital of carbon undergo intermixing and redistribution of energy to form four mixed type of orbitals having same shape and equal energy. This intermixing is called hybridisation. The orbitals obtained as a result of hybridisation are called hybrid orbitals. Thus, orbitals of carbon in excited state overlap with orbitals of hydrogen only after undergoing hybridisation. In simple words methane is formed by the overlap of four hybrid orbitals of carbon with four pure obitals of four hydrogen atoms (one from each hydrogen).Hybridisation is, thus defined as, a hyprothetical process of intermixing of orbitals, which either possess same or slightly different energy, so as to form, same number of orbitals (which is equal to the number of orbitals mixed and constitute set of equivalent orbitals) by redistribution, having same energy, same shape and symmetrical arrangement in space. Such orbitals are called hybrid orbitals. From the definition it is clear that this is a hypothetical process which means there is no proof of hybridisation. You may ask a very valid question that if there is no proof of hybridisation then how can we say that bonding in most of cases occur after the obitals have undergone hybridisation. The answer to this question is simple. The concept of occurence of Hybridisation helps us to justify almost all the problems related with the geometry of the covalent molecules both simple and complex. So the comfort brought to us forces us to believe that hybridisation occurs even if there is no proof till date.Types of hybridisation. Depending upon the number of types of orbitals involved we may have a very large number of different types of hybridisation. However, in organic chemistry carbon is the central atom and it can involve only two types of orbitals viz. 2s and 2p orbitals. Depending upon the number of p-orbitals involved (s-orbital is always involved and there is only one 2s orbital) we have three types of hybridisation.These are sp3-hybridisation, sp2-hybridisation and sp-hybridisation. Let us discuss these in detail.1.1.1. sp3-hybridisation. If all the three 2p-orbitals of carbon participate along with 2s-orbital in hybridisation resulting in formation of four equivalent orbitals, the hybridisation is called sp3-hybridisation. These four hybrid orbitals are directed towards the corners of a regular tetrahedron. Thus, any molecule possing this type of hybridisation will possess tetrahedral geometry. That is why this hybridisation is also known as tetrahedral hybridisation. Examples of molecules possessing this type of hybridisation include all saturated hydrocarbons and their derivatives like alkanes, cycloalkanes, alkyl halides, etc.Let us explain this with the help of two examples i.e. methane (CH4) and ethane (C2H6).(a) Structure of methane (CH4) molecule. One carbon atom when combines with four hydrogen atoms through single bonds result in formation of methane. Carbon in methane undergoes sp3-hybridisation. Four bonds are formed by the overlap of each sp3-hybridised orbital of carbon with 1s orbtial of each hydrogen. Thus, four sp3s, CH, -bonds are formed which are diagramatically represented in figure-1.

From the figure it is clear that methane has tetrahedral geometry. All HCH bond angles are of 10928. All the CH bond lengths are the same and are equal to 110 pm. Each bond has bond dissocitation energy of 416 kJ/mol.(b) Structure of ethane (CH3CH3) molecule. In ethane two carbons and six hydrogens are present. Each carbon is linked to three hydrogens so there are six CH bonds. Both carbons are also bonded to each other through a single bond (CC bond). Each carbon atom when combines with three hydrogen atoms and one carbon atom through single bonds result in formation of four bonds. Both carbons in ethane undergo sp3-hybridisation. There are two types of bonds; three bonds are formed by the overlap of each sp3-hybridised orbital of carbon with 1s orbtial of each hydrogen. Fourth bond is formed by overlap of sp3-hybridised orbital of each carbon. Thus, three sp3s, CH, -bonds and one sp3sp3, CC, -bonds are formed, which are diagramatically represented in figure-2.

From the figure, this is clear that each carbon of ethane has tetrahedral geometry. All bond angles are of 10928(HCH and HCC). All the CH bond lengths are the same and are equal to 110 pm. CC bond length is 154 pm. Each CH bond has bond dissocitation energy of 409.6 kJ/mol and CC bond dissociation energy is 347 kJ/mol.1.1.2. sp2-hybridisation. If two out of three 2p-orbitals of carbon participate along with 2s-orbital in hybridisation resulting in formation of three equivalent orbitals, the hybridisation is called sp2-hybridisation. These three hybrid orbitals are directed towards the corners of an isosceles triangle. Thus, any molecule possing this type of hybridisation will possess triangular or trigonal planar geometry. That is why this hybridisation is also known as triangular or trigonal planar hybridisation. The 2p-orbital which do not participate in hybridisation lie perpendicular to the plane containing hybrid orbitals. If these unhybridised or pure orbitals form a bond they do so by sidewise or lateral overlap and the bond formed between them is called -bond. Examples of molecules possessing this type of hybridisation include all unsaturated hydrocarbons possessing double bond and their derivatives like alkenes, cycloalkenes, aldehydes, ketones, carboxylic acids etc.Let us explain this with the help of following example i.e. Ethene (C2H4).Structure of ethene (CH2=CH2) molecule. In ethene two carbons and four hydrogens are present. Each carbon is linked to two hydrogens so there are four CH bonds. Both carbons are also bonded to each other through a single bond (CC bond). Each carbon atom when combines with two hydrogen atoms and one carbon atom through single bonds result in formation of three -bonds. Both carbons in ethene undergo sp2-hybridisation. There are two types of -bonds; two bonds are formed by the overlap of each sp2-hybridised orbital of carbon with 1s orbtial of each hydrogen. Third bond is formed by overlap of sp2-hybridised orbital of each carbon. Thus, two sp2s, CH, -bonds and one sp2sp2, CC, -bonds are formed, which are diagramatically represented in figure-2.Each carbon is left with one unhybridised 2p orbital which is perpendicular to the plane of carbon and hydrogen atoms. These p-orbitals are half filled and parallel to each other. Thus, they overlap from sides which is called sideways overlap or lateral overlap and results in formation of -bond. We can easily infer that -bond and -bond are perpendicular to each other. -bond will have two overlap areas; one lying above and one lying below the plane containing the carbon and hydrogen atoms.It is also important to note that though their are two overlapping areas but the bond is counted as one. More over number of areas where the overlap is occurring has nothing to do with strength of the bond. The strength of the bond depends upon the total area involved in the overlap. The total area of overlap is less in case of -bond. Thus, -bond is weaker than the -bond. In addition we can see -bond can be formed in the absence of -bond, whereas -bond can not be formed in the absence of -bond.

Thus, where ever you see a double bond you can easily infer that it has one -bond and one -bond. The -bond has a symmetry axis along the CC bond and -bond does not. -bond is bisected by a plane containing two carbon and four hydrogen atoms. In simple words the nature of the structure where the hybridisation is sp2 is planar or two dimensional.From the figure, this is clear that each carbon of ethene has trianular geometry or trigonal planar geometry. All HCH bond angles are of 120. All the CH bond lengths are the same and are equal to 108.7 pm. CC bond length is 134 pm. The shortening of bond length for C=C is due the fact that size of sp2-hybrid orbital is smaller than size of sp3-orbital. This is due to increased percentage of s-character in sp2-hybrid orbital (33%). Each CH bond has bond dissocitation energy of 443 kJ/mol. Increase in bond dissociation energy is due to the fact that the bond length is shorter and overlapping is more when a bond is formed between sp2-hybrid orbital and 1s-orbital of hydrogen.The bond dissociation energies for carbon-carbon double bond (composed of one sigma and one pi bond) is 607 kJ/mol. On the other hand bond dissociation energy of carbon-carbon single bond (sigma bond) is 347 kJ/mol. The difference in two bond dissociation enrgies will give value for bond dissociation energy of pi-bond. Thus, bond dissocaition energy is = 607 347 = 260 kJ/mol. From this value it is proved that pi-bond is weaker than sigma bond.As, pi-bond can be cleaved easily, the alkenes are more reactive than alkanes. More over since only pi-bond is broken, alkenes will undergo addition reaction (involves conversion of 1 pi-bond nto two sigma bonds).Structure of ethyne (CHCH) molecule. In ethyne two carbons and two hydrogens are present. Each carbon is linked to one hydrogen so there are two CH bonds. Both carbons are also bonded to each other through a single sigma bond (CC bond). Each carbon atom when combines with one hydrogen atom and one carbon atom through double bond result in formation of two -bonds. Both carbons in ethyne undergo sp-hybridisation. There are two types of -bonds; one bond is formed by the overlap of each sp-hybridised orbital of carbon with 1s orbtial of hydrogen. Second bond is formed by overlap of sp-hybridised orbital of each carbon. Thus, one sps, CH, -bond and one spsp, CC, -bond are formed, which are diagramatically represented in figure-3.Each carbon is left with two unhybridised 2p orbitals (2py and 2pz) which are perpendicular to the plane of carbon and hydrogen atoms and also perpendicular to each other. These p-orbitals are half filled and parallel to each other. Thus, they overlap from sides which is called sideways overlap or lateral overlap and results in formation of two -bonds. All the -bond and -bonds are perpendicular to each other. Each -bond will have two overlap areas. So, overlap areas of one -bond lie above and below the plane containing carbon and hydrogen. Other two overlap areas lie on the front and back side of the plane containing carbon and hydrogen. However, the four lobes or halves of the electron clouds of two -bonds do not stay as such but merge together to form a single electron cloud which which possess cylindrical symmetry about the internuclear axis.Thus, where ever you see a triple bond you can easily infer that it has one -bond and two -bonds. The -bond has a symmetry axis along the CC bond and -bonds do not. -bond is bisected by a plane containing two carbon and two hydrogen atoms. In simple words the nature of the structure where the hybridisation is sp is linear or diagonal.

From the figure, this is clear that each carbon of ethyne has linear geometry. All HCC bond angles are of 180. All the CH bond lengths are the same and are equal to 106 pm. CC bond length is 120 pm. The shortening of bond length for CC is due the fact that size of sp-hybrid orbital is smaller than size of sp3- and sp2-orbital. This is due to increased percentage of s-character in sp-hybrid orbital (50%). Each CH bond has bond dissocitation energy of 507 kJ/mol. Increase in bond dissociation energy is due to the fact that the bond length is shorter and overlapping is more when a bond is formed between sp-hybrid orbital and 1s-orbital of hydrogen. For CC bond, the same fact is applicable which has a bond dissociation energy of 803 kJ/mol.The bond dissociation energies for carbon-carbon double bond (composed of one sigma and one pi bond) is 607 kJ/mol. On the other hand bond dissociation energy of carbon-carbon single bond (sigma bond) is 347 kJ/mol. The difference in two bond dissociation enrgies will give value for bond dissociation energy of pi-bond. Thus, bond dissocaition energy is = 607 347 = 260 kJ/mol. From this value it is proved that pi-bond is weaker than sigma bond.Due to symmetrical nature of two mutually perpendicualr -bonds the delocalisation of -electron cloud decreases and this becomes the reason for lesser reactivity of alkynes as compared to alkenes towards addition reaction. 1.2. Bond length. From our knowledge of bond formation process, we know that atoms willing to make bond approach each other. This lowering of distance between the atoms lowers their potential energy and brings stability oto the system. The lowering of potential energy is directly proportional to the decrease in distance between combining atoms upto a certain limit. This distance up to which the potential energy of the system decreases with decrease in distance is called distance of closest approach or bond length. IF the atoms are brought further closer, then the potential energy of the system starts increasing with decrease in potential energy. At distance of closest approach repulsive and attractive forces are adjusted in such a way that the system possess minimum potential energy or maximum stability.Thus, bond length is defined as, the minimum distance between the centres of nuclei of the two atoms combined or bonded with the help of a covalent bond at which the attractive and repuslsive forces adjust in such a way that the system is in a state of miniumum energy or maximum stability is called bond length.Now, at advance level we can understand that atoms are not in stationary state but keeps on viberating and the bond may undergo inward or outward stretching reulsting in chainge of value of bond length. Then, the defination of the bond length was modified by taking the average of maxmum possible bond length and minimum possible bond length between the two atoms. Thus, improved definiation is, the average distance between the centres of nuclei of the two atoms combined or bonded with the help of a covalent bond at which the attractive and repuslsive forces adjust in such a way that the system is in a state of miniumum energy or maximum stability is called bond length. The units in which bond distance or bond lengths are measured are; nanometer (nm), Angstrom () or picometer (pm). It should be remembered that 1 nanometer (nm) = 109 m or 107 cm, 1Angstrom () = 1010 m or 108 cm1picometer (pm = 1012 m or 1010 cm.The bond length can be found wth the help of various methods. Some of the common methods include X-ray diffraction method, neutron diffraction method and different spectroscopic techniques. The magnitude of bond length or bond distance between two atoms say A and B remains the same if different molecules of same class are considered for finding the bond length. In case of alkanes (methane, ethane, propane, etc.) the CH bond length is same i.e. 110 pm or 1.10 . Similarly OH bond length in different alcohols is same and is equalt to 97 pm or 0.97.Bond lengths for some commonly interacted bonds during study of organic chemistry are given below in the table 1.1.Factors affecting bond length. From the table we see that bond lengths are different for different types of bonds. Some of the factors that effect bond-length are size of the combining atoms, multiplicity of the bond and type of the hybridisation. Let us study in brief, how these factors effect the bond length of the bond between different atoms.1. Size of the atoms. Bond length is directly related to the size of the two atoms bonded together. Larger the size of the atoms, larger will be the bond length. This is clear from the HH, ClCl bond length values which are 74 pm and 199 pm respectively. This is because hydrogen atom is smaller and chlorine atom is larger. Similarly if we see bond lengths of HF, HCl, HBr and HI we find that these are 91.7 pm, 127.46, 141.45 pm and 160.9 pm repectively. The increase in bond length is due to increase in size of halogen atom.2. Multiplicity of bond or bond order. The multiplicity of bond or the bond order is defined as the number of covalent bonds formed between two atoms. It is one, two or three depending upon wether single, double or triple bonds are formed. The bond length is inversely related to bond order. HIgher the bond order lower is the bond length. This can be understood if we see bond length values for CC, C=C and CC bonds which are 154 pm, 134 pm and 120 pm repectively. OO bond length and O=O bond length are 148 pm and 121 pm repectively. Similarly, NN, N=N and NN bond lengths are 147 pm, 124 pm and 110 pm respectively.We know that a double bond comprises of 1 and 1 bond. On the other hand a triple bond comprises of 1 and 2 bonds. -bond is formed by the overlapping along internuclear axis and is formed by the sideways or lateral overlp of unhybridised p-orbitals which occurs after the formation of -bond. Additional bonding will bring two atoms closer and shorten the bond length value. The same explanation can be extende to explain further lowering of bond length when a triple bond is formed between two carbon atoms. 3. Type of hybridisation. Hybridsation also effects the bond length value. In case of organic molecules we keep carbon as central atom and it is capable of showing three types of hybridation i.e. sp3, sp2 and sp. IF we see bond length values for CH bond having carbon in different hybridisation state, we find that CH bond lengths decrease in the order (sp3-hybridised)CH > (sp2-hybridised)CH > (sp-hybridised)CHThe values are repectively 110 pm, 108.7 pm and 106 pm. Same has been observed in case of CC bond lengths.(sp3)CC(sp3) > (sp2)CC(sp2) > (sp)CC(sp)The values are repectively 154 pm, 148 pm and 138 pm.Thus to conclude we can summerise:The bond length of the homonuclear diatomic molecules are twice the covalent radii.Bond length of polar bond is smaller than the theoretical non-polar bond length.Bond length increases with the increase in size of the bonding atoms.The bond length will decrease with increase in bond order.The bond length will decrease with increase in s-character in the hybridised orbitals.Average Bond Lengths for Some Single, Double and Triple BondsBondBond Length ()BondBond Length ()CC1.54NN1.47C=C1.34N=N1.24C=C1.20NN1.10CN1.43NO1.36C=N1.38N=O1.22CN1.16OO 1.48CO1.43O=O 1.21C=O1.23HH0.74CO1.13FF1.42CH1.10OH0.97NH1.03CS1.82HF0.92HCl1.27HBr1.41HI1.61

1.3. Bond angles.We have studied many polyatomic molecules (molecules with three or more atoms of same or different types). They possess more than one type of covalent bonds. Covalent bond has directional nature. This means they are present at definite angle with repect to each in space. In simple words polyatomic molecules have atoms at definite angles called bond angles. Bond angle can be difined as, the angle between any two adjacent covalnet bonds at an atom in a polyatomic molecule. Bond angles are very important feature of structure of a covalent molecule and can be measured with the help of various methods. Some of the common methods with the help of which we calcluate bond angles are X-ray diffraction and various spectroscopies. Like bond lengths we can not calculate exact bond angles due to viberations among the various bonds as well as various atoms. So, we calculate average bond angles and bond angles given in most of the books are average bond angles.Factors affecting bond angles. From the table we see that bond angles are different for different types of molecules. Some of the factors that effect bond-length are hybridisation of the central atom, types of electron pairs present around the central atom, electronegativity of the central atom. Let us study in brief, how these factors effect the bond angle.1. Hybridization: Bond angle depends on the state of hybridization of the central atom. Generally increase in s- character in the hybrid orbital increases the bond angle.HybridizationBond angleExamplesp310928Methane (CH4)sp2120Ethene (C2H4)sp180Ethyne (C2H2)2. Types of electron pairs present around the central atom. Bond angle is affected by the presence of lone pair of electrons at the central atom. A lone pair of electrons at the central atom always tries to repel the shared pair (bonded pair) of electrons. Due to this, the bonds are displaced slightly inside resulting in a decrease of bond angle.

Therefore, the regular geometry gets distorted with the presence of lone pairs and bond angles decrease from their expected values. We have already seen this in the case of CH4, NH3 and H2O, where the central atom is sp3-hybridized in each case but they have bond angles 109 28, 107 and 104, respectively. This is due to the presence of one lone pair in NH3 and two lone pairs in H2O. Similarly, variations are found from ideal values of 120 for sp2 carbon. e.g., in ethylene HCC and HCH bond angles are 121.7 and 116.6 respectively, due to the presence of electrons which repel CH bonds away.3. Electronegativity of the central atom. If the electronegativity of the central atom decreases, bond angle decreases. This fact is best explained when we study hydrides of group-13-16. In case the central atom remains the same, bond angle increases with the decrease in electronegativity of the surrounding atom.4. Vander Waal Interactions. Repulsion between atoms or groups attached to the central atom may increase or decrease the bond angle.1.4. Bond Energy. During the formation of a bond, certain amount of energy is released. This is called bond formation energy. The same amount of energy will be needed to break this bond. This will be known as bond dissociation energy. Bond formation energy and bond dissociation energies are equal but have opposite signs. For a diatomic molecule having one bond, bond energy (considered equal to bond dissociation or bond formation energy) may thus be defined as, ther energy required to break or released during the formation of one mole of a particular bond between two atoms. Bond dissociation energies are symbolized by D. These are expressed in units of kJ mol1. For example, 435 kJ mol1 of heat is needed to break a mole of hydrogen molecules into individual atoms. Therefore, the bond energy of hydrogen is 435 kJ mol1.

This definition is applicable to polyatomic molecules also provide they possess one bond of a particular type. For example, bond formation or bond dissociation energies in case of hypochlorous acid for OH and ClO bonds are equal to their bond energies. The values of bond energies for certain diatomic molecules are given in Table 1.2.Table 1.2: Bond energies of certain diatomic moleculesBondEnergy (kJ mol1)BondEnergy (kJ mol1)H H 435O O 146N N 297HI296.8Cl Cl 247HBr367.8Br Br 192HF568.5F F 155HCl430.5I I 150However, for a polyatomic molecule containing more than one covalent bond of same type, the term bond energy may be equal to bond dissociation energy (D) or different from the value of bond dissociation energy (D). It should be noted that bond dissociation energy values change with the change in environment of the bond which is to be broken. For example bond dissociation energies of two OH bonds in water are not same. First bond dissociation energy in water is more (498.7 kJ/mol) where as the second bond dissociation energy is less (428 kJ/mol).Thus, bond energy value will be different from bond dissociation energy in molecules possessing more than one bond of the same type. The water is stable molecule and OH radical is unstable. Breaking of bond from stable water molecule to form unstable OH radical is thermodynamically unfavourable and thus rquires a little more energy than when we convert unstable OH radical into repective atoms.From the above example we can also infer that i is better if we use bond energy which is average of bond dissociation energies of many bonds of one type present in the molecule. Otherwise we have to remember too many values along with the conditions which will not be easy.Thus, for a polyatomic molecule containing more han one bond of a particular type bond dissociation energiy and bond energy are different from each other in magnitude.We can illustrate the difference between these two values by considering another molecule such as methane. If we begin to remove the four hydrogen atoms one by one by splitting of carbon-hydrogen bonds, we get four different D values, These are shown below

Why do we get four different values when we are breaking a CH bond in each step? Before going for the answer of this question, we should not that the hybridisation of carbon changes after the loss of each hydrogen or rupturing of each CH bond. This process also gain or give energy. The hybridisation change which increases the percentage of s-character in the hybridised orbitals will be more stable and thus we need more energy in such cases.Now, we got the reason which justifies that the C H bond dissociation energy not only includes the energy needed for rupturing the bond, but it also includes the energy changes accompanying the rehybridization at the carbon atom in each step. Direct atomization of methane (conversion of methane into C and 4 H) requires 1665 kJ mol1, which is the sum of all the four D values listed above. Thus, the bond energy per CH bond in methane is thus one fourth of this value, i.e., 1665/4 = 416.25 kJ mol1.Since, a diatomic molecule has only one bond, D is equal to E in such cases as told earlier. The same is applicable for polyatomic molecules having one bond of each type. We normally use bond energy values rather than bond dissociation energy values due to the above explanantion.Some bond energies obtained by taking the average from measurements on several polyatomic molecules are listed in Table 1.4.Table 1.4: Average bond energies (E values in kJ mol1)BondAverage bond energy (E)BondAverage bond energy (E)O H 448C N285C H 414C = N615C F 448C N866C Cl 326C (sp3) H435C Br 284C (sp2) H443C I 213C(sp) H507C C 347C (sp3) C (sp3)347C = C 609C (sp2) C (sp2)383C C 804C (sp) C (sp)433C O 360N = N418C = O 740N N941Factors Effecting Bond Energies. The bond energy gives an approximate idea about the strength of a particular bond and depends upon the following factors:1. Size of atom. Bond energy is related with bond length which is indirectly related with size of the bonding atoms. Longer bonds are weaker and shorter bonds are stronger. Longer bonds are formed by lager atoms, thus, larger the size of atom forming covalent bond smaller will be the bond energy, e.g., as the size of halogen atom (X) increases from F to I, the bond energy of CX bond decreases i.e., CF (448) > CCl (326) > CBr (284) > CI (213) kJ mol1.2. Bond order or multiplicity of the bond. Bond energy increases with a increase in bond order. This means if two atoms are capable of forming more than one type of covalent bonds (single, double or/and triple) then the bond with higher order or higher multiplicity will have higher value of bond energy, e.g., for carbon-carbon bond where bond length decreases due to increase in multiplicity of the bond from single to triple bond the order of bond energy is as given below.C C (347 kJ mol1) < C = C (609 kJ mol1) < C C (804) kJ mol1 Similarly for oxygen-oxygen bond and nitrogen-bond the order of bond energy is as given below.N N (160 kJ mol1) < N = N (418 kJ mol1) < N N (941) kJ mol1O O (146 kJ mol1) < O = O (495 kJ mol1) 3. Type of Hybridization. We know that, the size of hybrid orbitals decreases in the order: sp3 > sp2 > sp and their electronegativity increases in the order. sp3 < sp2 < sp. Both these factors cause a decrease in bond length (from sp3 sp2 sp) and hence bond energy increases, e.g., for CH bond. C (sp3) H (416 kJ mol1) < C (sp2) H (443 kJ mol1) < C (sp) H (50 kJ mol1) and for C C bond C (sp3) C (sp3) (347 kJ mol1) < C (sp2) C (sp3) (383 kJ mol1) < C (sp) C (sp3)(433 kJ mol1).4. Stability of the product formed as a result of dissociation. Let us take example of cycloalkane molecules where the product formed is open chain free radical. Straight chain radicals formed from three and four membered cyclic compounds are more stable and thus they are easily formed. Bond dissociation energies leading to such processes will be quite less than the normal bond dissociation energy of a particular bond.5. Attainment of aromaticity. If a bond cleavage results in attainment of aromaticity, then the bond being more prone to cleavage will have smaller bond dissociation energy than expected. For example, cycloheptatriene will cleave CH bond by absorbing much less energy than required for normal CH bond due to aromatic nature of cycloheptatriene cation.1.6. LOCALIZED AND DELOCALIZED BONDSCovalent bond is formed by sharing of electrons by the bonding atoms. According to orbital overlap concept, atomic or hybrid orbitals overlap to form covalent bonds. We find two ways in which these orbitals can overlap viz along the intermolecular axis or above and below the intermolecular axis. The first type of overlapping results in formation of -bond and the latter gives -bond. IF we take example of hydrogen, we find that there is only one type of overlap possible between the two hydrogens which is along the intermolecular axis. The shared pair of electron for most of time stay in the space which is under the control of two hydrogen nuclei. They can not move freely. On the other hand in benzene, the -electrons move over the entire ring. Thus, we come across two types of bonded pairs of electrons or bonds on this basis. These are called localised electron pairs or bonds and delocalised electron pairs or bonds. Let us define these two types of bonds, first.Localised electron pair or localised bond. When the electron pair forming a bond spends most of their time in the space between the two bonded atoms controlled by their nuclei, it is called localized electron pair and the bond formed by this pair is called localized bond. All -bonds are localized bonds. e.g., -bonds formed between carbon and hydrogen (irrespective of hybridisation of carbon), -bonds formed between two carbons (in case of alkanes, alkenes, alkynes, cycloalkanes, benzene, etc.), etc. are all localised. Delocalised electron pair or delocalised bond. On the other hand, when the electron pair moves in and out of the space between the two bonded atoms, it is called delocalized electron pair and the bond formed by this pair is called delocalized bond. Some of the -bonds show localisation and some others show delocalisation. -bonds showing localisation are called localised -bonds and -bonds showing delocalisation are called delocalised -bonds.This is explained below:-bond or electron pair of oxygen is localized. -bond or electron pair of ethylene is localized because the electrons forming the -bond in ethylene stay in the space between two carbon atoms in such a way that these electrons are distributed equally in the space above and below the plane of CC, -bond.

Acetylene forms two -bonds which are not only perpendicular to plane containing carbons but are perpendicular to each other. These two -bonds of acetylene which should be localised due to their mutually perpendicular orientation are actually delocalized. This is because the electrons forming the -bond in the plane of paper do not remain confined to the space above and below the plane of CC, -bond and similarly the electrons of other -bond formed in perpendicular direction, do not remain confined to that space. Actually, all these electrons merge together to form a cylindrical electron cloud around CC -bond.Other examples of delocalized bonds which we come commonly come across are 1, 3-butadiene and benzene. In these cases, two or more than two -bonds are in conjugation (present at alternate position). The electrons of one -bond do not remain confined to the atoms responsible for their formation but move into the space of other -bond and vice-versa. This movement of electrons into and out of their alloted space is called delocalisation. This delocalization occurs through the overlap of unhybrid p-orbitals present on each sp2 carbon. Let us explain these two examples in little detail.(a) 1,2- and 1, 3-Butadiene. Butadiene (C4H6) possess two double bonds. They may be formed in two manners resulting in formation of two isomeric butadienes. These are called 1,2- or 1,3-butdienes. Three of the four carbons in 1,2-butadiene possess double bonds. The first and third cabon are sp2 hybridised where as the second carbon is sp-hybridised. The fourth carbon is sp3-hybridised. Two -bonds are formed between 1,2 and 2,3 carbons. These -electrons can move over the three carbons rather than staying between the carbons forming thembecause they lie in the same plane and possess close proximity. Thus they show delocalisation. Similarly, all the four carbon atoms (having sp2 hybridisation) of 1, 3-butadiene lie in the same plane due to which all the p-orbitals at four carbon atoms overlap with each other and the -electrons can move to a limited extent over all the four carbon atoms i.e., -electrons of 1, 3-butadiene are delocalized as shown. Hence, the p-bonds of butadiene are delocalized bonds.

In contrast, the -electrons and hence the -bonds are localised in isolated dienes because of larger distance between the p-orbtals as they are intervened by one carbon (sp3-hybridised) possessing no p-ornbital. One example of such diene is 1,4-pentadiene, where each pair of -electrons is confined to the space between two carbon atoms. (b) Benzene. Molecular formula of benzene is C6H6. It possesses 12 -bonds (6 CC -bonds and 6 CH -bonds) and three -bonds. Benzene has six sp2-hybridized carbon atoms and each sp2-carbon has one unhybridized p-orbital containing one electron. These p-orbitals are so close to each other that they can overlap sideways to form a -bond. The structure of benzene shows that -bonds are alternate in position. We can also say that there is high degree of conjugation. There are two ways in which overlapping of adjacent p-orbitals can take place. The structures formed due to these two eays are also called resonating structures. These are shown as:

Actually, each p-orbital overlaps equally well with the p-orbitals on adjacent two carbon atoms on both sides to form a doughnut shaped -electron cloud above and below the plane of carbon and hydrogen atoms which simply means that these three -bonds of benzene are delocalized.

A very commonly asked question is about representing its structure. You might have seen that at place benzene is represented with single and double bonds and at some places it is shown as a hexagonal ring with a dotted or solid circle inside it. This circle denotes six completely delocalised -electrons.

The later structure of benzene is becasue the three -bonds of benzene are completely delocalized. So, it is not proper to represent benzene with a hexagonal ring with three double bonds at alternate positions since the position of -bonds is not fixed. Therefore, benzene is written as a hexagonal ring with a dotted or solid circle inside it. This representation was given by Armit and Robinson.Consequences of delocalisation. There are many effects which we can observe in organic chemistry arisen because of delocalisation. Many features of organic compounds are also effected by this phenomenon of delocalisation. Some of the characteristics of delocalised bonds are given below:1. Stability. The delocalisation of the electrons effectively spreads their charge over a larger region of space lowering the energy of the system making the molecule more stable. This situation can easily be recognised in conjugated systems i.e. systems that have alternating double and single bonds. Conjugated dienes are more stable as compared to their isolated counterparts and this is due to the delocalisation. The extent of stability is directly related to the extent of delocalisation. Greater the delocalisation greater will be the stability of the molecule.Similarly, benzene without delocalisation would have 151 kJ/mol more energy or it would have been more unstable and more reactive.The difference in energy that is between actual benzene (with delocalisation) and hypothatical benzene (with no delocalisation) is also termed as resonance energy.2. Effect on carbon-carbon bond lengths. The bond lengths of carbon carbon would have been of two types corresponding to CC (154 pm) and C=C bonds (134 pm). But these bond lengths in benzene are of one type and their value is equal to 139 pm which is inbetween the value of the two bond lengths. This is easily understood with the idea of delocalisation.Similarly due to overlap between p-orbitals on second and third carbon atoms in 1,3-butadiene molecule, the bond length decreases from expected value of 154 pm to 149 pm. This is becasue partial double bond character is acquired by these two carbons due to delocalisation.3. Effect on acidic strength. Delocalisation in the anions of protic acids decide their acidic strengths. If the anion has higher degree of delocalisation (which maes t more stable), the compound will be more acidic. This explains the order of acidic strength of different oxyacids of halogens which is as follows:Perchloric acid > Chloric acid > Chlorous acid > Hypochlorous acid.4. Stability of intermediates. Some intermediate become more stable due to delocalisation. For example stability of tropyllium carbocation and allyl radical and carbocation show stability which is more than expected.5. Acidic nature of -Hydrogen. Whenever, we have unsaturated polar bonds like carbonyl, nitrile etc. the -H becomes more acidic. In simple words, it is easily removed due to formation of carbonaion which has higher stabiliy due to delocalisation.6. Formation of unexpected products. In many reactions we do not get the desired product. We rather get unexpected products. Their formation can be easily understood with the help of concept of delocalisation in the intermediates to result in formation of more stable unexpected intermediates. We will study many examples in the later portions of this book.1.6. VAN DER WAALS INTERACTIONSWe have learnt that atoms combine to become more stable. For this they do many manupulations with their electrons like transferring electrons or sharing electrons, etc. Most easily understandable and widely accepted concept for reason of bonding is octet rule. According to this we believe that atoms combine to acquire octet of electrons in their valence shells. In different manner we can say different atoms combine to form molecules in their desire to acquire configuaration of the nearest noble gas. If this is true then noble gas should show complete inertness. Their atoms should show no type of interactions with each other or with atoms of other elements. They should not be converted into liquid or solid form. But this is not so. The same should be applicable to gases and non-polar molecules. Gases must obey gas law under all the conditions of temperature. But these also are not observed. We know that (i) Gases do not obey the ideal gas equation, PV = nRT, and (ii) A non-polar gaseous compound can be condensed to the liquid state which in turn can be converted into solid by increasing the pressure and decreasing the temperature.How can these facts be accounted for? In order to account for these observations, van der Waals, a Dutch chemist, proposed the existence of some sort of attractive forces between the molecules of non-polar compounds. These attractive forces between the molecules of non-polar compounds are called van der Waals forces. These are short range forces and are responsible for bringing the molecules close together.Origin of van der Waals Attractive Forces. We know in a non-polar neutral molecule, the centres of positive and negative charge densities are not separated and consequently these molecules do not possess any dipole moment. In simple words we can say that negative charge density is uniformly distibuted over the entire space of the molecule. The movement of electrons around the nucleus is thus, uniform. These molecules keep on moving randomly (the movement is called Brownian movement) and colliding with each other. Because of collisons amongst the molecules or some other factors, movement of electrons around the nucleus becomes irregular or non-uniform. Because of this, the centres of positive and negative charge densities get separated at that instant of time and this leads to the generation of small instantaneous local dipole commonly known as instantaneous induced dipole. This instantaneous dipole moment induces some polarity in the neighbouring molecules. The dipole moment corresponding to such induced polarity is known as induced dipole moment. When such polar molecules come close enough, their dipoles attract each other and forces them to stay together in the vacinity of each other. Because of this, such attractions are also referred to as instantaneous dipole-induced dipole attractions. These instantaneous and induced dipoles are constantly changing but the overall result is the attraction between such molecules. Such attractions are also called London forces or dispersive forces.London forces vary inversely as the sixth power of the distance between them i.e. F 1/r6These forces are always attractive forces and are at maximum in the solid and minimum in the gases. This is simple to understand on the basis of distnace between the molecules which is very less in solid state and quite large in gaseous state. Moreover, these attractive forces are appreciable only between those portions of the molecules over which they can touch each other, i.e., between the surfaces of the molecules. For this reason it is understandable that surface area plays significant role in deciding the magnitude of van der Waal forces. Greater the surface area of the molecule, the greater will be overall van der Waals forces.Factors influencing Vander-Waals forces. Vander-Waals forces are very weak and their strength depends upon following factors :1. Molecular and atomic size. The atoms or molecules with larger size possess large surface area. As a result of this, their electron clouds are more diffused and distorted. due to greater distortion of electron clouds, the strength of Vander-Waal forces will be more.2. Number of valence electrons. The magnitude of Vander-Waal forces increases with increase in the number of electrons in the molecule. Due to increase in the number of electrons, the chances of distortion of the electron cloud by the movement of electrons also increases, thus, the magnitude of Vander-Waal forces increases.3. Atomic or molecular size. The magnitude of Vander-Waal forces depends upon the shape and structure of the molecule. The straight chain isomers have large surface area as compared to the branched chain isomer. Consequently, straight chain isomers show greater magnitude of Vander-Waal forces and have boiling point than branched chain isomers.For example, if we compare the boiling points of n-pentane (309 K), iso-pentane (301 K) and neo-pentane (283 K), it has been observed that n-pentane possess highest boiling point as compared to iso-pentane which in turn shows more boiling point than neo-pentane. The explanantion depends upon the fact that the n-pentane possess straight carbon chain where as iso-pentane and neo-pentane show branching. So, the magnitude of van der Waal forces will be maximum in n-pentane and minimum in neo-pentane. Thus, the boiling point increases accordingly.4. Molecular mass. Vander-Waal forces tend to increase with molecular mass. This is because molecules with larger molecular mass usually have more electrons and London forces increase in strength with the number of electrons. Also, larger molecular mass often means larger atoms, which are more polarisable. Example boiling point of n-octane (C8H18, molecular mass = 114) is greater than n-hexane (C6H14, molecular mass = 86).5. Temperature. The decrease in temperature decreases the kinetic energy of molecules so that molecules come close to each other resulting in increase of intermolecular forces. In other words, low temperature favours Vander-Waal forces.6. High pressure. On increasing the pressure, the molecules come close to each other resulting in the increase of the magnitude of Vander-Waal forces.van der Waals forces and physical properties of substaces. These forces are directly associated with many properties of the liquids and gases. The strength of intermolecular forces influence the magnitude of these properties. Some of the properties are discussed below:1. Vapour pressure. The ease or difficulty with which a molecule leaves the liquid depends upon the strength of its attraction to other molecules. The vapour pressure is the pressure extended by the vapours on the surface of the liquid in equilibrium with the liquid. The vapour pressure depends upon the intermoleculer forces. when the intermolecular forces are strong, the vapour pressure will be low. In liquids, intermolecular attractions are usually observed. Many liquids particularly show Vander-Waal forces. Out of different types of Vander-Waal forces, the London forces are usually dominant. The dipole-dipole force is appreciable only in smaller polar molecules or in large molecules having very large dipole moments. The London forces tend to increase with molecular mass, therefore, it is expected that liquids with high molecular mass possess low vapour pressure as due to strong interactions liquid molecules are unable to escape into vapour state. The vapour pressure of various liquids are shown in the table and it is clear from the table that the vapour pressure decreases with increas in the molecular mass of the liquid.2. Boiling point. The boiling point of a liquid depends upon inter molecular forces. The boiling point decreases with increase in the magnitude of intermolecular forces as due to strong interations between the molecules of liquid, higher energy will be required to overcome these forces and to form vapours. Normal boiling points are approximately proportional to the energy of intermolecular attractions. Thus, they are lowest for liquids with the weakest intermolecular forces.3. Surface tension. Surface tension also depends on intermolecular forces . Surface tension is the energy needed to increase the surface area of a liquid. To increase the surface area, it is necessary to pull molecules apart against the intermolecular forces of attraction. Thus, surface tension would be expected to increase with the strength of attractive force. If London forces are the dominant attractive forces in the liquid, surface tension whould increase with molecular mass.4. Viscosity of a liquid. The viscosity of a liquid depends upon the intermolecular forces because increasing the attractive forces between the molecules increase the resistance to flow.5. Physical state of substances. The difference in physical state of substances can be explained on the basis of van der Waals forces. This can be explained by considering the examples of physical state of halogens. F2 and Cl2 are gases due to weak van der Waal forces of attraction. Iodine is a solid due to maximum magnitude of Vander-Waal forces of attraction while bromine is a liquid due to intermediate van der Waal forces of attraction. van der Waals Attraction in Polar Molecules. Apart from non-polar molecules the van der Waals forces of attraction also occur in polar molecules. They may originate from dipole-dipole and dipole-induced dipole interactions.(i)Dipole - dipole interaction: Polar molecules having permanent dipole moment are held together by dipole-dipole interactions. For instance, in gases such as NH3, HCl, HF, SO2, etc., there are significant dipole-dipole attractions between the molecules of these gases due to the presence of permanent dipole moment in the molecules. The extent of this type of interaction depends upon the magnitude of the dipole-moment of the molecule. Thus, greater the dipole moment, stronger is the dipole-dipole interaction.(ii)Dipole-induced dipole interaction: They are attractive interactions between polar and non- polar molecules. A polar molecule may induce polarization in a non-polar molecule present in its vicinity. This induced dipole then interacts with the dipole moment of the first molecule and thereby the two molecules are attracted to each other [Fig. 1.25 B]. The extent of such interaction depends upon the magnitude of the dipole moment of the polar molecule and the polarizability of the non-polar molecule. For instance, the increasing order of water solubility of noble gases from He to Rn is attributed to increase in the magnitude of dipole induced dipole intraction due to increase in their polarizability in the same order.van der Waals Repulsions. We have learnt that the van der Waals forces of attraction are inversely related to the sixth power of the distance between the interacting moieties (increase in magnitude with decrease in the distance between non-bonded atoms or molecules). Due to short range nature of these forces, this attraction is effective only to a certain point (range is 500 pm). If such molecules are brought further closer (distance between them is less than the sum of van der Waal radii of the interacting species), the attractive forces turn into great repulsive forces, called van der Waals repulsions. This looks like definitiation of chemical bond and thus, it can be said that all molecules in gaseous state possess an effective size which can be named as van der Waals radius. It may be mentioned that the van der Waals radii of various atoms are greater than their covalent radii by about 80 pm. The minimum distance beyond which the forces of attraction turn into forces of repulsion is equal to the sum of the van der Waals radii of two atoms or molecules. These repulsive forces are very strong in nature. They bring strain in the molecule which is called van der Waals strain.In the following table van der Waals radii of some atoms and molecules are listed. Table: van der Waals radii (pm) of some atoms and groupsSpeciesVan der Waals Radii (pm)SpeciesVan der Waals Radii (pm)H120Cl180N150Br195O140I220S190CH3200F135Benzene 170Effects of van der Waals repulsive forces. These repuslive forces play very important role in determining shapes, conformations and geometries of the molecules. 1. Stability of a particular isomer. Covalent molecules have definite geometries due to directional nature of the covalent bond. Some time the atoms or groups in a molecule can exist in two or more geometries having different energies. The geometry corresponding to a particular arrangment will have fixed energy and different geometries will have different geometries. The molecule will have different geometries at different temperatures. However if the atoms or groups are situated at such positions that the distance between them is smaller than the sum of their van der Waals radii, they are said to be crowding together and such a geometry will not be preferred. In order to achieve maximum stabilitya situation where such repulsive forces could be avoidedthe molecule orients itself in such a way that the distance between these atoms increases.For instance, cis-1, 3-dimethylcyclohexane is more stable when the substituents are equatorial rather than axial because in the latter case, 1, 3-diaxial interactions (because of van der Waals repulsions) make the molecule less stable. In 1, 3-diaxial conformation the distance between the two axial methyl groups and one axial hydrogen is much less than the sum of their van der Waals radii which results in repulsive forces commonly known as1, 3-diaxial interactions. However, in the diequatorial conformation, the two methyl groups are held at a distance which is more than the sum of their van der Waals radii and hence such repulsive forces are absent. Therefore, cis-1, 3-dimethylcyclohexane is more stable in its diequatorial conformation than in diaxial.Similarly if we look at the fully eclipsed conformation of butane we find that it is less stable and more energetic than partly eclipsed, skew and anti coformations for the same reason. You will study these in stereochemistry in detail. 2. These repulsions are important during asymmetric synthesis where a particular geometrical or optical isomer is obtained in large proportions.3. These repulsive forces are very significant in case of biochemical reactions.1.7. INCLUSION COMPOUNDS AND CLATHERATESA bird in the cage is common sight in our surrounding. Similarly, we can see some molecules trapped in cage like formations of some ther molecules. Organic solids such as urea, thiourea, hydroquinone, cyclodextrins, cholic acid, deoxycholic acid, etc. form cage like structures which are capable of trapping molecules. These are called host molecules. They have a particular sized cavity in them. They may have have channel like spaces (like space inside a spring). The molecules which are trapped inside these cavities or channels are called guest molecules. The guest and host molecules do not form any chemical bond. They remain together due to weak van der Waals forces of attraction. These compounds are called inclusion compounds (because guest is included into the space or cavity of the host). They are sometimes also known as caged compounds. There is no stoichiometric molar ratio between host and guest molecule. There is no electronic or chemical interaction criterion for the guest molecule and only the shape and size decides the guest molecule. The size of the space or cavity is the host molecule and size of the guest must be comparable. Guest molecules which are too big or too small are not suitable for forming inclusion compounds because larger molecules will not fit into the cavity and smaller molecules will excape easily because the forces of attraction between the guest and host are weak and they can not stop the smaller molecule from leaving the guest molecule.Types of Host-Guest Addition Compounds.Depending upon the type of the space available within the crystal lattice of the host molecule the host- guest addition compounds or inclusion compounds are broadly classified into following two categories:1. Inclusion compounds. The host-guest addition compounds are known as inclusion compounds when the space available within the crystal lattice of host molecule is in the form of long channels.2. Clatherates. The host-guest addition compounds are known as clatherates, when the space available within the crystal lattice of the host molecule is in the form of a cage.Both these types of addition compounds are well defined crystalline solids but they are not useful for derivatization since they decompose at the melting point of the host compounds. Although the two types of addition compounds differ from each other yet they have one thing in common that is the guest molecules of only right size can be trapped into the crystal lattice of the host molecules.The inclusion compounds are generally solids which melt at the melting point of the host with decomposition and therefore are not useful as derivatives. We will now discuss both types of addition compounds in more details.1. Inclusion Compounds. (i) Urea as host: The most widely studied host molecule for inclusion compounds is urea. Ordinarily urea crystallizes in tetragonal shape. However, when a guest is present, urea crystallizes in the hexagonal lattice trapping the guest molecules (n-alkanes) in the space. Since hexagonal type of lattice is formed only in presence of some guest molecule. The diameter of the urea channel is about 500 pm and the type of guest molecules will depend upon their shape and size only. It may be reasonable to believe that the van der Waals forces of attraction between the host and the guest molecule are key forces for the stability of inclusion compounds. Some examples are described here.(a) Sizes and shapes of n-Octane, 2-chlorooctane and 1-bromooctane are just appropriate and thus they act as guest for urea. But sizes and shapes of 2-bromooctane, 2-methylheptane and 2-methyloctane are not appropriate and thus they do not fit into the channels made by urea. Hence, they do not act as guests for urea lattice. (b)Both dibutyl maleate and dibutyl fumarate are guests but neither diethyl maleate nor diethyl fumarate is a guest.(c)Dipropyl fumarate is a guest but dipropyl maleate is not.Thus, mixture of two compounds (one of which can act as guest and one which can not act as guest are spearated by forming these compounds)In these complexes there is no fixed whole number stochiometry (i.e., whole number molar ratio) of host and guest. If we see an example where it is simple whole number ration, it is by chance. For example, the host-guest compound of n-octane and urea has ratio of 1 : 6.73. (ii)Thiourea as host: Thiourea also forms inclusion compounds with channels of larger diameter. Sizes and shapes of n-Octane, 2-chlorooctane and 1-bromooctane are not appropriate and thus they do not act as guest for thiourea. But sizes and shapes of 2-bromooctane, 2-methylheptane and 2-methyloctane are just appropriate for becoming guest and thus they fit into the channels made by urea. Hence, they act as guests for thiourea lattice. (iii) Amylose as host: Iodine gives deep blue clour with starch which is a test for the presence of starch in any mixture. Starch (which is a mixture of two components i.e., amylose which is water soluble and amylopectin which is water insoluble) gives a deep blue colour with iodine. Amylose has helical structure (spring like). It has channel like space inside the helix which is of the appropriate size and polarity to accommodate an iodine molecule. Thus the blue colour is due to the formation of an inclusion complex between amylose and iodine.(iv) Cholic acid as a host: Scientists have reported the formation of a large number of inclusion compounds where cholic acid acts as a host to wide variety of organic substances such as aliphatic, alicyclic and aromatic alcohols, ketones, aldehydes, ethers and carboxylic acids, esters and nitrite. Most of these compounds were formed by recrystallization of cholic acid from liquid guest component. It can be seen that cholic acid tends to form stable inclusion compounds with organic substances having polar functional groups suggesting that hydrogen bonds between host and guest molecules play some role in stabilizing these inclusion compounds.Uses of inclusion compounds: Some of the more important uses of inclusion compounds are as under:(i)These complexes are quite useful in separating certain isomers that would be otherwise difficult to separate. For example, n-octane can be separated from its branched chain isomers, because only n-octane can form inclusion compound with urea.(ii)Urea inclusion compounds have also been used for resolving racemic mixtures. For example, () racemic mixture of 2-chlorooctane forms two different inclusion compounds (diastereomeric) which can be separated by fractional crystallization.(iii) Amongst the isomeric threo- and erythro-9,10-dihydroxystearic acid (formed by hydroxylation of oleic acid) only threo isomer forms an inclusion compound with host urea while erythroisomer having the two OH groups on opposite side of the chain does not.(iv) Dipropyl fumarate is accepted as a guest by host urea but its isomeric dipropyl maleate is not.CLATHRATE OR CAGE COMPOUNDSLike inclusion compounds the clathrate or cage compounds also require a host as well as a guest. The difference between the two types of compound lies in the nature of crystal lattice spaces in the host compound. While these spaces are tunnel type in the inclusion compounds, in clathrate compounds these spaces arc completely enclosed in which the guest molecules are trapped. Again there is no formal bonding between the host and guest molecules and only the shape and size decides the guest molecule. Also in contrast to inclusion compounds, some of the crystal lattice spaces in the clathrate compounds may remain vacant. Hence, these compounds are also non-stoichiometric.Most of these clathrate compounds are solids and are stable at low temperature but decompose when kept at room temperature.Important hosts forming clathrate compounds include sodium chloride, some other halides of alkali metals, hydroquinone and water. Hydroquinone and water which are capable of forming hydrogen bonds.(i) Hydroquinone as host or hydroquinone clatherates. Three molecules of hydroquinone combine through hydrogen bonding to enclose a cage like space into which guests of proper size can fit in. There are many cavaties having a diameter of 400 pm. The common guests for host hydroquinone are SO2, CO2, Argon and methanol. It is important to note that methanol forms a clathrate whereas ethanol does not as it is not a suitable guest. Similarly neon does not form clathrate with hydroquinone host as it is not a suitable guest. Normally all the cavities are not filled by guests. Hence, they do not show simple and fixed molar ratio.The clatherates of noble gases and hydroquinone are quite stable and can be kept for years. However they decompose on heating or on adding some liquid capable of forming hydrogen bonding with hydroquinone results in escape of gases entrapped in hydroquinone cage. This is because the hydroquinone cage is broken due to formation of H-bonding with the added liquid.(ii) water as host or water clatherates. Water can act as host for many guests like noble gases (argon, krypton and xenon), Cl2, propane, methyl iodide etc. Six molecules of water combine through hydrogen bonding to form a cage. When mixture of water and noble gases is cooled below 273 K under high pressure we get clatherate of noble gas and water having water noble gas ratio of 575 : 1. These clatherates are stable at low temperature and thus stored below 273 K. These clatherates decompose when the clatherate is heated to room temperature.Methane hydrate, a clathrate compound of host water which forms an ice-like cage at low temperature and high pressure to encapsulate methane as a guest, is the clean fuel of the future. Amongst ionic compounds sodium chloride and some other alkali metal halides are good hosts for guest organic molecules like benzene and naphthalene forming clathrate compounds.In addition there is one type of host which can form both inclusion as well as clathrate compounds. These are known as Cyclodextrins. Here the host molecule is made up of six to eight glucose units connected to form a large ring with primary hydroxyl group projecting outside the ring while the secondary hydroxyl groups on the inner side of the ring. When cyclodextrins are dissolved in water the spaces inside the ring gets filled with water molecules held in place through hydrogen bonding. However the less polar inner ring side can accommodate nonpolar organic molecules displacing water. Accordingly cyclodextrins can form 1 : 1 stoichiometric clathrate complexes with guests like acetic acid, propionic acid and butyric acid. Here again the proper size of the guest is all too important for the formation of clathrate compounds.Cyclodextrins are also able to form inclusion compounds where guest molecules are stacked, like coins, in row over each other in a tunnel like space. Thus cyclodextrins form inclusion compounds with valeric and other higher acids.The spaces inside the different kind of cyclodextrins are of different sizes and a large variety of guests can be accommodated either as inclusions or clathrate complex. Hence these nontoxic cyclodextrins are used industrially as host molecules to encapsulc food or drugs as guest.Uses of Clathrates(a)Krypton-85 clathrates provide a safe and useful source of beta radiation which are useful for measuring thickness of gauges.(b)Xenon-133 clathrates provide a compact source for gamma radiation.(c)In the separation of noble gases, for example, neon can be separated from argon, krypton and xenon because neon is the only gas that does not form clathrates with quinol.(d)Clathrates play an important role in some physiological actions, e.g. it is thought that the anaesthetic action of xenon and many other anaesthetics is due to aqueous clathrate in physiologically strategic spots. When the anaesthetic is no longer administered, the clathrate equilibrium is destroyed, the clathrate decomposes and the consciousness returns.1.8. Charge-transfer ComplexesA charge-transfer complex (or CT complex, electron-donor-acceptor- complex) is a chemical association of two or more molecules, or of different parts of one very large molecule, in which attraction between the molecules (or parts) occurs due to an electronic transition into an excited electronic state, such that a fraction of electron charge is transferred between the molecules. The resulting electrostatic attraction stabilizes the molecular complex. These can be understood as, molecular complexes formed by mixing an electron rich molecule (called the donar molecule, represented as D) to the electron deficient or poor molecule (also known as acceptor, represented as A) usually in 1 : 1 molar ratio are called charge transfer complexes.The source molecule from which the charge is transferred is called the electron donor and the receiving molecule is called the electron acceptor. Hence the alternate name, electron-donor-acceptor-complex (EDA). Species which have lone pairs, delocalised -electron system, etc. can act as donar and species which have empty orbitals or positive charge or are electron deficient can act as acceptors.During formation of these complexes charge is transferred from donar to acceptor and this is represented as:It is interesting to note that the attraction in a charge-transfer complex is not equivalent of a stable chemical bond and is much weaker than covalent forces. These complexes are formed due to weak coloumbic forces rather than normal bonds.The formation of electron donor acceptor (EDA) complex is confirmed from their spectrum (called charge transfer spectrum) which is different from the spectrum of both donor as well as the acceptor molecules. The difference between the energies of first excited state and ground state of these complexes is relatively small, as a result of which most of the EDA complexes absorb in visible or near UV region and are often coloured. Optical spectroscopy is a powerful technique which is used to characterise charge-transfer bands. These optical absorption bands are referred to as charge-transfer bands or CT bands.Charge-transfer complexes exist in organic as well as inorganic molecules and in all states of matter, i.e. in solids, liquids and even gases.The charge transfer or EDA complexes may be categorised in two groups:(i)EDA complexes where donor is an olefin or an aromatic compound and acceptor is a metal ion, and(ii)EDA complexes where donor may be -donor and acceptor is an organic molecule.In inorganic chemistry, most charge-transfer complexes involve electron transfer between metal atoms and ligands (electron rich species capable of donation of electrons). The charge-transfer band in transition-metal complexes result from movement of electrons between Molecular Orbitals (MO) that are predominantly metal in character and those that are predominantly ligand in character.The classic example of charge-transfer complex is that between iodine and starch forming an intense purple colour. Some electron-rich organic molecules combine to give crystalline CT-complexes like picrate and trinitrobenzoates. Most thoroughly studied CT-complexes or EDA complexes in organic chemistry are pictrates in which picric acid acts as acceptor. Let us describe picrates in detail to understand charge transfer complexes.Picrates. 2,4,6-trinitrophenol is known as picric acid. Presence of three electron withdrawing groups (nitro groups) make it highly electron deficient. When it comes close to electron rich molecules like amines (both alphatic and aromatic), phenols, naphthalene, anthracene, mesitylene, phenanthrene and other polynuclear hydrocarbons which are electron rich species, it gets bonded to them though weak coloumbic interactions. As a result we get addition compounds or EDA-adducts which are commonly known as picrates. These picrates are intensely coloured and different donors give different colours with picric acid. This sometimes help us to identify the donar species.The formation of picrates by different donors are shown asThe distance between donor and acceptor molecules is always more than that of normal bonds which proves that these type of compounds are formed only due to weak coloumbic interactions. The extent of electron richness of the donor and electron deficiency or acceptor decides the extent of the charge transfer and thus, colour of the complex. If the donor has high electron density available for donation, the compound will be more intensely coloured. Picrate of xylene are less intensely coloured as compared to picrates of mesitylenes.The transfer of energy can take place only when the difference in energies of HOMO of donor matches with the energy of the LUMO of the acceptor. This means their energies should either be same or there should be very small difference in their energies. HOMO means Highest Occupied Molecular Orbital and LUMO means Lowest Unoccupied Molecular Orbital. Determination of crystal structure of picrates has revealed that the aromatic rings are held in parallel planes. If donor and acceptors are not in parallel planes no charge complexes are formed. This is proved by taking into example of 3-methyl-2,2,4,4,6-pentanitrodiphenyl which is highly electron deficient does not form the CT-complexes like picric acid. The only reason is non-coplanarity of the molecule.What ever be the nature of attractive forces, the bonding in charge transfer complex is usually weak and hence, the complex readily dissocaites in solution. The extent of charge separation in a charge transfer complex is measured by its dipole moment which is often higher than the dipole moment of acceptor molecule before formation of charge transfer complex. For example, a charge transfer complex of pyridine and iodine shows that dipole moment of pyridine (228 D) is raised to 490 D after the formation of complex.Similarly, dipole moment of 1,3,5-trinitrobenzene is raised from zero (due to symmetrical nature of the molecular structure) to significant values when it forms charge transfer complexes with various donors. Electron Displacement in Organic CompoundsThe behaviour of an organic compound (derivative of hydrocarbon) is influenced by the electron displacement taking place in its covalent bonds. These displacements may be permanent or temporary, which take place in presence of another species in the molecule. The acidity and basicity of organic compounds, their stability, their reactivity towards other substances, etc., can be predicted based on such electronic displacements. The electron displacement in an organic molecule may take place either in the ground state under the influence of an atom or a substituent group or in the presence of an appropriate attacking reagent. The electron displacements due to the influence of an atom or a substituent group present in the molecule cause permanent polarlisation of the bond. Inductive effect and resonance effects are examples of this type of electron displacements. Temporary electron displacement effects are seen in a molecule when a reagent approaches to attack it. This type of electron displacement is called electromeric effect or polarisability effect. In the following sections we will learn about these types of electronic displacements.Inductive Effect. Organic compounds are either hydrocarbons (saturated) or their derivatives (unsturated hydrocabons and functional groups contain derivatives). In saturated hydrocarbons we have two types of bonds. These are (i) bonds between carbon and hydrogen, (ii) bonds between two carbons. When a carbon atom is bonded to a hydrogen (CH) or another carbon (CC) atom by a covalent bond as in alkanes, the sharing of electron pair is symmetrical between them. Thus, no charges are induced on the atoms. However, in case of functional group derivatives of alkanes, one or more hydrogens are replaced with some atom or group of atoms. The atom or the group of atoms attached to carbon may have different electronegativity than that of carbon. This results in polarisation of bond between the carbon and the group attached to it. The electron pair no longer stays in the middle. It is shifted either away from carbon or twowards the carbon depending upon the electronegativity of the group or atom attached with the carbon.The polarization of a bond due to electron withdrawing or electron donating effect of adjacent groups or atoms is called inductive effect or simply I-effect.Let us try to understand it with the help of following example.The CCl bond in the butyl chloride, CH3CH2CH2CH2Cl is polarized due to electronegativity difference of carbon and chlorine. The electrons are shifted or displaced towards the chlorine atom due to its higher electronegativity. Thus the first carbon atom (C1) gets partial +ve charge (+) and chlorine gets partial ve charge ().

In turn, this partially positively charged carbon atom withdraws electron density from the next carbon (C2). As a result, C2 will also acquite small positive charge (+) which will be less than that at the first carbon (C1). Similalrly, C3 carbon will acquire small positive charge which is smaller than charge on C2 (+). Thus, the inductive effect is transmitted through the carbon chain with decreasing magnitude of positive charge on carbons (as we move away from halogen). So, sometime it is also called transmission effect or T-effect.The inductive effect is not significant beyond 3rd carbon atom. It diminishes as we move away from the heteroatom or the group causing it. This means that inductive effect is most prominant at the carbon which is attached to the heteroatom.Types of inductive effects. The inductive effect is divided into two types depending on their strength of electron withdrawing or electron releasing nature with respect to carbon. These are:(i) Negative inductive effect (I)(ii) Positive inductive effect (+I)(i) Negative inductive effect (I): If the heteroatom attached to the end of carbon chain is more electronegative than carbon, the electrons are diplaced towards the heteroatom or more electronegative group and the effect is called negative inductive effect or I effect. In simple words the electron withdrawing nature of groups or atoms is called as negative inductive effect. It is indicated by I.This can be diagramatically represented as

The extent of I effect is directly proportional to the electronegativity of the heteroatom or the group. Higher the electronegativity, higher will be the I effect. Following are the examples of groups in the decreasing order of their I effect: NR3+ > SR2+ > NH3+ > NO2 > SO2R > CN > SO2Ar > SO3H > CHO > CO > COOH > COCl > CONH2 > F > Cl > Br > I > OC6H5 > COOR > OR > SR > SH > OH > NH2 > Alkynyl > C6H5 > Alkenyl > H where R means alkyl and Ar means aryl.(ii) Positive inductive effect (+I): If the heteroatom attached to the end of carbon chain is less electronegative than carbon, the electrons are diplaced towards the carbon or away from heteroatom or group and the effect is called positive inductive effect or +I effect. In simple words the electron releasing nature of groups or atoms is called as negative inductive effect. It is indicated by +I.This can be diagramatically represented as

The extent of +I effect is inversely proportional to the electronegativity of the heteroatom or the group. Higher the electronegativity, lesser will be the +I effect. Positive inductive effect shown by metals are observed when we study organometallic compounds like alkyl lithium, grignard reagents or tetramethyl tin, tetra alkyl lead, tetra methyl silane, etc. Following are the examples of groups in the decreasing order of their +I effect: C(CH3)3 > CH(CH3)2 > CH2CH3 > CH3 > HDeuterium (D) causes greater +I effect than hydrogen, perhaps, because of larger atomic radii which make it more electropositive than hydrogenAll alkyl groups show +I effect. The magnitude of +I effect increases with the increase in number of alkyl groups attached with the carbon under consideration. In simple words more substituted alkyl groups have more +I effect. This can be easily explained on the basis of hyperconjugation which we will discuss later in this chapter. Why alkyl groups are showing positive inductive effect?Though the CH 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. Salient features of inductive effect.* It involves displacement of -electrons along the carbon chain to which the heteroatom or group is attached. * It arises due to electronegativity difference between two atoms forming a sigma bond. * It is transmitted through the sigma bonds. * The magnitude of inductive effect decreases while moving away from the groups causing it. The inductive effect is not significant beyond 3rd carbon atom.* It is a permanent effect and can not be reversed. * It influences the chemical and physical properties of compounds.APPLICATIONS OF INDUCTIVE EFFECT. A large number of phenomena have now been explained on the basis of inductive effect. Some of these like decrease in bond lengths, origin of dipole moment, strength of acids and bases, are considered as applications of inductive effect.Effect on bond lengths. Usually the bond length decreases with increase in inductive effect, perhaps, because the covalent bond assumes increased ionic character. The bond length of CF, C Cl, C Br and CI bonds are in the order given below: C F < CCl < CBr < CI .This is due to increase in the radius of halogen from F to I but the C C bond lengths in following skeletal structures should be same.CC F = C C Cl = CC Br = CCI.However, the inductive effect of halogen is in the order F > Cl > Br > I which results in the following order of bond lengths for CC bondsCCF < CCCl < CC Br < CCIStability of reaction intermediates. We will study about many reaction intermediates like carbocations, carbanions, free radicals, etc. They all are unstable but different types of reaction intermediates have different order of stbility. Their order of stability can be explained on the basis of inductive effect. We will study about reactive intermediates in detail in second chapter of this book. Their orders of stability are given here in brief. Stability of carbocations: The stability of carbocations 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 stabilizes the ion. 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: Polarity of covalent bond and dipole moment of organic compounds.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: