Chemical Bonding and Molecular Structure PDF

4 Chemical Bonding and Molecular Structure Chemical Bond The attractive force that binds the atoms together in a molecule is called a chemical bond. It is formed either by the transfer of electrons or by the sharing of electrons. Lewis symbols In the formation of...

4 Chemical Bonding and Molecular Structure Chemical Bond The attractive force that binds the atoms together in a molecule is called a chemical bond. It is formed either by the transfer of electrons or by the sharing of electrons. Lewis symbols In the formation of a chemical bond, only the outermost electrons participate and these electrons are called valence electrons. G.N Lewis introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols or Lewis notations. Here the nucleus and the inner electrons are represented by the symbol of the element and the valence electrons by dots or crosses. Octet rule This rule was proposed by Lewis and Kossel. According to this rule, atoms undergo chemical reaction in order to attain octet of electrons in the valence shell. Or, atoms containing 8 electrons in their valence shell are stable. To attain octet, they share electrons or exchange electrons. Ionic or Electrovalent Bond A bond formed by the transfer of electron between 2 atoms is called ionic or electrovalent bond. Or, it is the force of attraction between two oppositely charged ions. In Na atom, there is only one electron in the valence shell and in Cl, there are 7 electrons in the valence shell. In order to attain octet state, Na atom loses one electron and form Na+ ion and Cl atom gains one electron and form Cl- ion. These two ions attract each other by electrostatic force called ionic bond. Energy changes during the formation of an ionic bond An ionic bond formation involves the following steps: The formation of a positive ion (cation) from the electropositive atom. The amount of energy required for this process is called ionisation enthalpy (ΔiH). M(g) + ΔiH → M+(g) + e- Formation of negative ion from the electronegative atom. The energy involved in this process is called electron gain enthalpy (ΔegH). X(g) + e- → X-(g) + ΔegH The packing of the cation and anion to form an ionic compound. The energy change in this process is called lattice enthalpy (ΔlatticeH) M + + X- → MX(s) + ΔlatticeH Factors favouring the ionic bond formation are: 1  Low ionisation enthalpy of the electropositive atom (metal atom) High negative electron gain enthalpy of the electronegative atom (non-metal atom) High lattice enthalpy of the ionic compound formed. Covalent bond A bond formed by the mutual sharing of electrons between two or more atoms is called covalent bond. 1)The formation of Chlorine molecule (Cl2) Both the chlorine atoms have 7 electrons in their outermost shell. So here the exchange of electrons is not possible. For the formation of Cl2 molecule, each Cl atom shares one pair of electron. The bond thus formed is called a single covalent bond. 2) The formation of O2 molecule 8O + 8O → O2 2,6 2,6 To attain octet, each oxygen atom shares 2 pairs of electrons. Thus a double covalent bond is formed 3)The formation of N2 molecule 7N + 7N → N2 (2,5) (2,5) To attain the octet configuration, the nitrogen atoms share 3 pairs of electrons. Thus a triple bond is formed. In the formation of a covalent bond, each combining atoms contribute at least one electron to the shared pair. The number of electrons contributed by an atom to the shared pair (bonded pair) during the formation of a covalent bond is called covalency. In the formation of Cl2 molecule, the covalency of each Cl atoms is one, in O2 molecule, the covalency of O atom is 2 and in N2, it is 3. The double and triple bonds are together called multiple bonds. A covalent bond formed by the sharing of one pair of electron is called a single bond. A double bond is formed when 2 pairs of electrons are shared between 2 atoms and a triple bond is formed when 3 pairs of electrons are shared between two atoms. Limitations of Octet rule 2  It could not explain the stability of compounds containing less than 8 electrons around the central atom. E.g. LiCl, BeH2, BCl3 etc. It could not explain the stability of molecules containing odd number of electrons like NO, NO 2 etc. It could not explain the stability of molecules containing more than 8 electrons around the central atom (i.e. expanded octet). E.g. PF5, SF6, H2SO4, IF7 etc. Octet rule is based upon the chemical inertness of noble gases. But some noble gases like xenon and krypton form compounds with F and O. This theory does not account for the shape of molecules. It does not explain the relative stability of the molecules. Bond Parameters 1)Bond Length: It is defined as the average equilibrium distance between the nuclei of two bonded atoms in a molecule. It is measured by spectroscopic, X-ray diffraction and electron diffraction techniques. It is expressed in pico-metre (pm) or in angstrom unit (A0). 2) Bond Angle: It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule. It can be experimentally determined by spectroscopic methods. It is expressed in degree. 3) Bond Enthalpy: It is defined as the amount of energy required to break one mole of a particular bond between 2 atoms in gaseous state. Its unit is kJ/mol. For diatomic molecules, bond enthalpy is equal to bond dissociation enthalpy. Larger the bond dissociation enthalpy, stronger will be the bond formed. For poly atomic molecules, the bond enthalpy is the average of the bond dissociation enthalpies. 4) Bond order: It is defined as the number of bonds between 2 atoms in a molecule. For H2, bond order = 1, for O2, bond order = 2 and for N2, bond order = 3. Isoelectronic molecules and ions have identical bond orders. eg: F2 and O22- have bond order 1. & N2 , CO and NO+ have bond order 3. Generally with increase in bond order, bond enthalpy increases and bond length decreases. Resonance Structures 3  In the case of some compounds, all the observed properties cannot be explained by a single structure. Here we use more than one structures of the compound. These different structures are called resonance structures or canonical structures or contributing structures. The phenomenon is known as resonance. Resonating structures are separated by double headed arrows. 1)Ozone (O3) 2) Carbonate ion CO2 Characteristics of resonance Resonance stabilizes the molecule. As the number of resonating structures increases, the stability also increases. Resonance changes the bond length. The difference in energy between the actual structure and the most stable canonical structure is called resonance energy. Greater the resonance energy, more stable will be the molecule. Polarity of bonds – Dipole moment When a covalent bond is formed between 2 similar atoms, the shared pair of electrons is equally attracted by the two atoms. So the electron pair is exactly between the two nuclei. The bond so formed is called a non-polar covalent bond. When the covalent bond is formed between 2 dissimilar atoms, the shared electron pairs are more attracted by one of the atoms. So one atom gets a slight negative charge (δ-) and the other gets a slight positive charge (δ+). Such molecules are called polar molecules. E.g. HCl, HF, H2O, HI etc. The polarity of a molecule is expressed in terms of dipole moment (μ). Dipole moment is defined as the product of the magnitude of charge at one end (Q) and the distance between the charges (r). Mathematically, μ = Q x r. The unit of dipole moment is Coulomb metre (Cm). 4  Usually expressed in the unit Debye (D). 1D = 3.336 x 10-30Cm. Dipole moment is a vector quantity. i.e. it has both magnitude and direction. It is denoted by a small arrow with tail on the positive centre and head pointing towards the negative centre. e.g.: HCl In the case of poly atomic molecules, dipole moment depends on the individual bond dipoles and the spatial arrangement of bonds. The dipole moment of the molecule is the vector sum of the bond dipoles of various bonds. e.g. H2O For BeF2, the net dipole moment is zero, since the two equal bond dipoles are in opposite directions and cancel each other In BF3, the net dipole moment is zero. Here the resultant of any 2 bond dipoles is equal and opposite to the third. Both ammonia (NH3) and nitrogen fluoride (NF3) are pyramidal in shape. Even though F is more electro negative than H, the net dipole moment of NF3 is smaller than that of NH3. This is because in the case of NH3, the orbital dipole due to lone pair is in the same direction as the resultant dipole moment of the three N – H bonds. But in NF3, the orbital dipole is in the opposite direction to the resultant dipole moment of the three N-F bonds. So the dipole moments get partially cancelled. Covalent character in ionic bonds – Fajans Rules 5  The partial covalent character of ionic bonds was explained by Fajans in terms of the following rules: The smaller the size of the cation and the larger the size of the anion, the greater the covalent character of an ionic bond. The greater the charge on the cation, the greater the covalent character of the ionic bond. For cations of the same size and charge, the ion with electronic configuration (n-1)dnns0 is more polarising than the ion with a noble gas configuration (ns2 np6). The shapes of molecules Covalent bonds are directional in nature. i.e. they are directed to some specified positions in space. So covalent compounds have definite shapes. The Valence Shell Electron Pair Repulsion [VSEPR] Theory The important postulates of this theory are: The shape of the molecule depends on the number of valence shell electron pairs (VSEPRs) around the central atom. The valence shell electron pairs repel each other. In order to reduce the repulsion, the electron pairs stay at maximum distance. Presence of lone pairs of electron causes distortion in the expected geometry of the molecule. The repulsion between two lone pairs of electrons is different from those between two bond pairs or between a lone pair and bond pair. The repulsion decreases in the order lone pair - lone pair > lone pair - bond pair > bond pair - bond pair. As the angle between the electron pairs increases, the repulsion decreases. I)Molecules containing only bond pairs of electrons 1) AB2 type Here there are 2 VSEPs. In order to reduce the repulsion, these electron pairs are arranged at an angle of 1800. Thus the shape of the molecule is linear with bond angle 1800 e.g. BeCl2, HgCl2 etc. 2) AB3 type Here there are 3 VSEPs. In order to reduce the repulsion, these electron pairs are arranged at an angle of 1200. Thus the shape of the molecule is planar triangular (trigonal planar) with bond angle 1200. e.g. BF3, BCl3 etc. 3) AB4 type Here there are 4 VSEPs. These are arranged at the four corners of a tetrahedron and hence the shape of the molecule is tetrahedral with bond angle 109.50. e.g.: CH4, NH4+ etc 6 4) AB5 type Here there are 5 VSEPs. To reduce the repulsion, they are arranged at the five corners of a trigonal bipyramid with bond angles 1200 and 900. eg. PCl5 5) AB6 type Here there are 6 VSEPs. To reduce the repulsion, they are arranged at the six corners of an octahedron with bond angles 900. eg.: SF6 II) Molecules containing both bond pairs and lone pairs The presence of lone pairs of electron causes distortion in the shape of the molecules. Explanation of shapes of ammonia and water molecules by VSEPR theory 1. NH3 In ammonia, the central atom N has 5 valence electrons (7N – 2,5). Among these electrons, three are used for the formation of bonds with hydrogen atoms and the remaining 2 electrons stay as lone pairs. So there are 4 VSEPs. Hence the expected shape of the molecule is tetrahedral. But due to the repulsion between lp –bp, the shape is distorted to trigonal pyramidal and the bond angle changes from 109028l to 1070. 7 2. H2O In water, the central atom O has 6 valence electrons (8O – 2,6). Two of them are used for the formation of bonds with hydrogen atoms and the remaining 4 electrons stay as lone pairs. So there are 4 VSEPs. Hence the expected shape of the molecule is tetrahedral.But due to the presence of 2 lone pairs, the shape is distorted.ie,lone pair - lone pair repulsion > lone pair - bond pair repulsion > bond pair - bond pair repulsion The shape of the molecule is bent or angular or inverted v shape and the bond angle changes from 109028l to 104.50. Type of Total no.of No. of No. of Shape e.g. molecule VSEPs b.ps l.ps AB2E 3 2 1 Bent SO2, O3 AB3E 4 3 1 Trigonal Pyramid NH3 AB2E2 4 2 2 Bent H2O AB4E 5 4 1 See-saw SF4 AB3E2 5 3 2 T-Shape ClF3 AB5E 6 5 1 Square Pyramid BrF5 AB4E2 6 4 2 Square Planar XeF4 (Refer pg 112 & 113) Valence bond theory (VBT) This theory was introduced by Heitler and London and later developed by Linus Pauling in order to explain the shape of molecules theoretically. VBT can be explained by considering the formation of H2 molecule. Consider 2 hydrogen atoms A and B with nuclei NA and NB and electrons eA and eB respectively. When the two atoms are at large distance from each other, there is no interaction between them. So their potential energy is zero. When the two atoms approach each other, new attractive and repulsive forces begin to operate. Attractive forces arise between nucleus of one atom and electron of other atom ie NA– eB and NB– eA. Similarly repulsive forces arise between electrons of two atoms i.e., eA – eB, and the nuclei of two atoms NA – NB. Attractive forces bring the two atoms close to each other whereas repulsive forces push them away. Experimentally it has been found that the magnitude of new attractive force is more than the new repulsive forces. So the two atoms approach each other and potential energy decreases. At a particular stage, the net attractive force balances the net repulsive forces and the energy becomes minimum. At this stage the hydrogen atoms are said to be bonded together to form a stable molecule. 8  The distance between the two atoms at this stage is called bond length. The amount of energy liberated at this stage is called bond enthalpy. The potential energy diagram for the formation of H2 molecule is as shown below: Orbital overlap concept Orbital overlapping is the process of partial interpenetration of atomic orbitals. The important characteristics of orbital overlapping are: 1. A covalent bond is formed by the overlapping of half filled atomic orbitals present in the valence shell of atoms. 2. The overlapping orbitals should contain electron with opposite spin. 3. As a result of overlapping, the electrons get paired and a stable covalent bond is formed. 4. The strength of a covalent bond depends on the extent of overlapping. 5. The greater the extent of overlapping, the stronger will be the covalent bond formed. Overlapping of atomic orbitals When two atoms come close to each other ,there is overlapping of atomic orbitals.This overlap may be positive , negative or zero depending upon the properties of overlapping atomic orbitals. 9 Types of overlapping There are two types of orbital overlapping. 1. Axial overlapping: If the overlapping of atomic orbitals take place along inter nuclear axis, it is called axial overlapping or end to end overlapping. A bond formed by axial overlapping is called sigma (σ) bond. The electrons present in sigma bond are called sigma electrons. All single bonds are sigma bonds. A sigma bond can be formed by the following ways: i) s-s overlapping: Here the overlapping of two half filled s-orbitals take place along the inter nuclear axis. ii) s-p overlapping: It occurs with the overlapping of one half filled s-orbital and one half filled p-orbital. iii) p-p overlapping: Here two half filled p- orbitals of two atoms overlap. 2. Lateral overlapping: Here the overlapping take place perpendicular to the inter nuclear axis. The bond formed as a result of lateral overlapping is called pi (π) bond. The electrons in pi bond are called π electrons. A π bond is always present along with σ bonds. A double bond contains one σ bond and one π bond. A triple bond contains one sigma bond and two pi bonds. A sigma bond is stronger than a pi bond. This is because the extent of overlapping is greater in a sigma bond. Comparison of Sigma and Pi Bonds 10 Hybridisation It is the process of inter mixing atomic orbitals having slightly different energies to form new orbitals having equivalent energy and identical shape. The new orbitals formed are called hybrid orbitals. Characteristics of hybridisation The number of hybrid orbitals formed is equal to the number of atomic orbitals undergo hybridization. The hybrid orbitals are always equivalent in energy and in identical shape. The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals. The hybrid orbitals are directed to some fixed positions in space. So the type of hybridization gives the shape of the molecule. Types of hybridisation 1. sp3 hybridisation: It is the process of inter mixing of one s-orbital and three p-orbitals to form four new orbitals having equivalent energy and shape. The 4 new orbitals formed are called sp3 hybrid orbitals. They are directed to the four corners of a regular tetrahedron with bond angle 109028’. Each sp3 hybrid orbitals has 25% s-character and 75% p- character. 1)Formation of methane (CH4) In CH4, the central atom C has the electronic configuration 6C – 1s 2s 2p 2 2 2 One s-orbital and three p orbitals undergo sp3 hybridisation. These sp3 hybrid orbitals are directed to the four corners of a regular tetrahedron with bond angle 109.50. Each of these sp3 hybrid orbitals overlap with 1s orbital of H to form four C-H σ bonds. 11  2)Formation of ethane (C2H6) In ethane, each C atom undergoes sp3 hybridisation. Out of the 4 sp3 hybrid orbitals, one of each C overlaps axially to form a C-C σ bond. The remaining three sp3 hybrid orbitals of each C overlap with the 1s orbital of H to form 6 C-H σ bonds. 3)Formation of Ammonia(NH3 ) The one s-orbital and three p-orbitals of N undergo sp3 hybridisation. Out of the 4 sp3 hybrid orbitals one is occupied by a lone pair and the other three sp3 hybrid orbitals overlap with 1s orbital of H form 3 N-H bonds. Due to the greater repulsion between lone pair and bond pairs, the shape is distorted to pyramidal and the bond angle becomes 1070. 4)Formation of water( H2O) Due to the greater repulsion between lone pairs, the shape is distorted to angular shape or bent structure or inverted ‘v’ shape and the bond angle becomes 104.50 2.sp2 hybridisation: It is the process of inter mixing of one s and two p-orbitals to form three new orbitals having equivalent energy and shape. The 3 new orbitals formed are called sp 2 hybrid orbitals. 12  They are directed to the three corners of an equilateral triangle. So the shape of the molecule is planar triangular or trigonal planar with bond angle 1200. Each sp2 hybrid orbitals has 33% s character and 67% p-character 1) Formation of BCl3 One s-orbital and two p-orbitals undergo sp2 hybridisation. The 3 hybrid orbitals formed overlap with 2p orbitals of Cl to form 3 B-Cl σ bonds. Since the hybridisation is sp2 the shape of the molecule is planar triangular with bond angle 2)Formation of ethene or ethylene (C2H4) In ethene, each C atom undergoes sp2 hybridisation. Out of the 3 sp2 hybrid orbitals, one of each C overlaps axially to form a C-C σ bond. The remaining two sp2 hybrid orbitals of each C overlap with the 1s orbital of H to form 4 C-H σ bonds. Now each C atom has 1 unhybridized p- orbitals which overlap laterally to form a π bond. 3. sp hybridisation: It is the process of inter mixing of one s-orbital and one p-orbital to form two new orbitals having equivalent energy and shape. The 2 new orbitals formed are called sp hybrid orbitals. They are directed in a straight line. So the shape is linear with bond angle 1800. Each sp hybrid orbitals has 50% s-character and 50% p character. 13  1) Formation of BeCl2 In BeCl2, the central atom Be has the electronic configuration1s22s2. In the excited state, one of the 2s electrons is promoted to 2p level. So the configurationon is 2s 1 2p 1.One s-orbital and one p- orbital undergo sp hybridisation to form 2 new sp hybrid orbitals. Each of thes hybrid orbitals overlaps with the 2p orbitals of Cl to form 2 Be-Cl bonds. So the shape of the molecule is linear with bond angle 1800. 2) Formation of ethyne or acetylene(C2H2) In acetylene, each C atom undergoes sp hybridisation. Out of the 2 sp hybrid orbitals, one of each C overlaps axially to form a C-C σ bond. The remaining two sp hybrid orbitals of each C overlap with the 1s orbital of H to form 2 C-H σ bonds. Each C atom has 2 unhybridized p- orbitals which overlap laterally to form 2 π bonds. Thus the molecule has linear shape with bond angle 1800. Here there are 3 σ bonds and 2 π bonds in ethyne, 4. sp3d hybridisation: It is the process of inter mixing of one sorbital, three p-orbitals and one d orbital to form five new orbitals having equivalent energy and shape. The 5 new orbitals formed are called sp3d orbitals. These are directed to the five corners of a trigonal bipyramid with bond angle 900 and 1200 14  Formation of PCl5 In PCl5, three P-Cl bonds lie in one plane at an angle of 1200.These three bonds are called equatorial bonds. The other two P-Cl bonds lie one above and one below this plane. They are called axial bonds. The axial bond pairs suffer more repulsion from the equatorial bond pairs. So the axial bond length is greater than the equatorial bond length which makes PCl5 more reactive. 5. sp3d2 hybridisation : It is the process of inter mixing of one s-orbital, three p-orbitals and two d-orbitals to form six new orbitals having equivalent energy and shape. The 6 new orbitals formed are called sp3d2 hybrid orbitals. These are directed to the six corners of a regular octahedron with bond angle 900. Formation of SF6 In SF6 the central sulphur atom has the ground state outer electronic configuration 3s 23p4.In the exited state one electron each from 3s and 3p orbitals are promoted to 3d level. Now one s-orbital, three p-orbitals and two and two d-orbitals undergosp3d2 hybridisation.These hybrid orbitals 15  overlap with p orbitals of fluorine atoms to form 6 S–F sigma bonds.Thus SF6 molecule has a regular octahedral geometry with bond angle 900. Molecular Orbital Theory This theory was developed by F.Hund and R.S Mulliken. The important postulates of this theory are: In molecules, the electrons are present in some special type of orbitals called molecular orbitals(M.Os). The atomic orbitals (A.Os) of comparable energy and proper symmetry combine to form molecular orbitals. Atomic orbitals are monocentric, while molecular orbitals are polycentric. i.e. electrons present in atomic orbitals are attracted by only one nucleus. While the electrons present in molecular orbitals are attracted by more than one nuclei. The number of molecular orbitals formed is equal to the number of atomic orbitals combined. i.e. if 2 atomic orbitals combined, 2 molecular orbitals are formed. One is called bonding molecular orbital (BMO) and the other is called anti-bonding molecular orbitals (ABMO) The BMO has lower energy and greater stability than the corresponding ABMO. The molecular orbitals give the electron probability distribution around a group of nuclei. The molecular orbitals are filled according to 3 rules – Aufbau principle, Pauli’s exclusion principle and Hund’s rule. Formation of molecular orbitals – Linear Combination of Atomic Orbitals (LCAO) method Molecular orbitals are formed by the combination of atomic orbitals by an approximate method known as Linear Combination of Atomic Orbitals (LCAO). According to this theory, the combinations of atomic orbitals take place by addition and subtraction of wave functions of atomic orbitals. The M.O formed by the addition of A.Os is called the bonding molecular orbital (BMO) and by the subtraction of A.Os is called the anti-bonding molecular orbitals (ABMO). The electron density in a BMO is located between the nuclei of the bonded atoms. So the repulsion between the nuclei is very low. Therefore a BMO always possess lower energy than the combining atomic orbitals. In ABMO, most of the electron density is located away from the space between the nuclei. There is a nodal plane between the nuclei and hence the repulsion between the nuclei is high. So an ABMO possess higher energy than the combining atomic orbitals. ABMO is denoted by using an asteric (*) mark. 16 Conditions for the combination of atomic orbitals The combination of atomic orbitals to form M.Os takes place only if the following conditions are satisfied: The combining atomic orbitals must have the same or nearly the same energy. The combining atomic orbitals must have the same symmetry about the molecular axis. The combining atomic orbitals must overlap to the maximum extent. Types of molecular orbitals M.Os of diatomic molecules are designated as σ (sigma), π (Pi), δ (delta) etc. The sigma M.Os are symmetrical about the bond axis, while the pi M.Os are not symmetrical. Linear combination of two 1s atomic orbitals produces two M.Os – a BMO σ1s and an ABMO σ*1s. Linear combination of two 2s atomic orbitals produces two M.Os – σ2s and σ*2s. As z-axis is taken as the inter nuclear axis, linear combination of two 2pz orbitals produces two M.Os σ2pz and σ*2pz. Two 2px orbitals produces 2 M.Os – π2px and π*2px and two 2py orbitals produce 2 M.Os – π2py and π*2py. Energies of various M.Os The various M.Os are filled in the increasing order of their energies (Aufbau Principle). The increasing order of energy is: σ1s < σ*1s < σ2s < σ*2s < (π2px = π2py) Na). A molecule is unstable if Nb < Na For a single bond, B.O = 1, for a double bond B.O = 2 and so on. As the bond order increases, bond length decreases and bond enthalpy increases. Magnetic nature If all the M.Os in a molecule are doubly occupied, the substance is diamagnetic. i.e. that substance is repelled by an external magnetic field. If one or more M.Os are singly occupied, it is paramagnetic. i.e. it is attracted by an external magnetic field. 17 Molecular orbital Diagram The representation of various M.Os in the increasing order of energy is called M.O diagram. 1 ) H2 Bond order = ½ [Nb –Na] = ½ [ 2-0] = 1 Diamagnetic 2)He2 Bond order = ½ [Nb –Na] = ½ [ 2-2] = 0 He2 is unstable and doesnot exist because of zero bond order. 3) Li2 Bond order = ½ [Nb –Na] = ½ [ 4-2] = 1 Diamagnetic 18  4) N2 Bond order = ½ [Nb –Na] = ½ [ 10 - 4] = 3 Diamagnetic 5)O2 Bond order = ½ [Nb –Na] = ½ [ 10 - 6] = 2 Paramagnetic 19 Hydrogen Bonding The weak attractive force between hydrogen atom of one molecule and electronegative atom (like F, O or N) of the same or different molecule is termed as Hydrogen bond. It is weaker than a covalent bond but stronger than van der Waal’s force. It is represented by dotted line (……….). Cause of H bonding: When H is bonded to a strongly electronegative atom X, the shared electron pair is shifted more towards X. So H atom gets a slight positive charge (δ+) and the electronegative atom gets a slight negative charge (δ-). This results in the formation of a polar molecule. The electrostatic force of attraction between these polar molecules is termed as H-bonding. δ+ δ- δ+ δ- δ+ δ- δ+ δ- …….. H-X …….. H-X ……… H-X ………. H-X …… Types of H bond There are two types of H bonds- inter molecular H-bonding and intra molecular H-bonding. 1)Inter molecular H bonding: It is the H bond formed by H atom of one molecule and the electronegative atom of another molecule. e.g. i) HF ….H-F …….. H-F ……… H-F …… H-F …… ii) H2O Inter molecular H bonding influences the physical properties of the compounds. Water (H2O) is a liquid with high boiling point but hydrogen sulphide (H2S) is a gas. This is because in water inter molecular H bonding is possible which is not possible in H 2S. 2)Intra molecular H bonding: It is the H bond formed between H atom and the electronegative atom of the same molecule. e.g. Hydrogen bonding in ortho-nitrophenol 20 21