K o s s e l L e w i sA p p r o a c h T oC h e m i c a l B o n d i n g• In 1916, Kossel and Lewis were the first tobecome independently successful in givinga satisfactory explanation about theformation chemical bond in terms ofelectrons.• They were the first to provide some logicalexplanation of valence which was based onthe inertness of noble gases.
L e w i s a p p r o a c h• He pictured the atom in terms of a positively charged‘Kernel’ and the outer shell which could accommodatea maximum of eight electrons.• He assumed that these eight electrons occupy thecorners of a cube which surround the ‘Kernel’.• Lewis postulated that atoms achieve the stable octetwhen they are linked by chemical bonds.• In sodium and chlorine, this can happen by transferof electron from sodium to chlorine.• In case of other molecules like CL2 the bond is formedby sharing of a pair of electrons between atoms.
L e w i s s y m b o lIn formation of a molecule, only theouter shell electron take part inchemical combination and theyare called valence electrons.The inner shells do not get involvedin the combination process.G.N. Lewis, an American chemistintroduced simple notations torepresent valence shell electronsin an atom and these are knownas Lewis symbols.
S i g n i f i c a n c e o fL e w i s s y m b o l• The number of dots represent the number ofvalence electrons.• This number of valence electrons help tocalculate the common or group valence ofthat element.• The group valence of the element is generallyeither equal to the number of dots in Lewissymbol or 8 minus the number of dots.
• Other important facts given by Kossel are-• In periodic table, the highly electronegative halogens and thehighly electropositive alkali metals are separated by noblegases.• The formation of a negative ion from a halogen atom and apositive ion from an alkali metal atom is associated with thegain and loss of electron by the respective atoms.• The negative and positive ions thus formed attain stablenoble gas electronic configurations.• The negative and positive ions are stabilized by electrostaticattraction.
Example• Na Na+ + e-[N e ] 3s 1 [N e ]• Cl + e- Cl-[N e } 3s 2 3p 5 [N e ] 3s 2 3p 6 o r [A r ]=> Na+ + Cl- NaCl
O c t e t r u l e• According to electronic theory of chemicalbonding, atoms can combine either bytransfer of valence electrons from one atomto another or by sharing of valence in orderto have an octet in their valence shells. this isknown as octet rule.
R e c a l l : E l e c t r o n s a r ed i v i d e d b e t w e e n c o r ea n d v a l e n c ee l e c t r o n s .A T O M c o r ev a l e n c eN a 1s 2 2s 2 2p 6 3s 1 [N e ]3s 1B r [A r ] 3d 10 4s 2 4p 5 [A r ] 3d 104s 2 4p 5C o v a l e n tB o n d i n gCovalent bond is the sharing of the VALENCEELECTRONS of each atom in a bondBr Br
C o v a l e n t b o n d• The bond formed by atoms when they share theirvalence electrons to gain octet is known as covalentbond.• When two atoms share one electron pair they aresaid to be joined by a single covalent bond.• If two atoms share two pairs of electrons, thecovalent bond between them is called a double bond.• When combining atoms share three electron pairs, a triple bond is formed.
C o v a l e n t b o n dd i a g r a m sH 2ON H 3C H 4
L e w i s r e p r e s e n t a t i o no f s i m p l em o l e c u l e s (t h e L e w i ss t r u c t u r e s )• The Lewis dot structures provide a picture of bonding inmolecules and ions in terms of shared pairs of electrons andthe octet rule.• The Lewis dot structures can be written by adopting thefollowing steps.– 1)the total number of electrons required for writing thestructures are obtained by adding the valence electrons ofthe combining atoms.– 2) for anions, each negative charge would mean additionof one electron. For cations, each positive charge wouldresult in subtraction of one electron from the totalnumber of valence electrons.
– 3)one must know the chemical symbols of thecombining atoms and their skeletal structures.– 4)in general, the least electronegative atomoccupies the central position.– 5)after accounting for shared pairs of electronsfor single bonds, the remaining electron pairs areeither utilized for multiple bonding or remain aslone pairs.
F o r m a l c h a r g e• A formal charge (FC) is the charge assigned toan atom in a molecule, assuming that electrons ina chemical bond are shared equally betweenatoms, regardless of relative electro negativity.• The formal charge of an atom in a polyatomicmolecule or ion may be defined as the differencebetween the number of valence electrons of thatatom in an isolated or free state and the number ofelectrons assigned to that atom in Lewis structure.
• Formal charge = (total no. of valenceelectrons in the freeatom)-(total no. of nonbonding electrons)-(1/2(total no. of bondingelectrons)
• Formal charge on ‘1’ =6-2-1/2(6)=+1• Formal charge on ‘2’ =6-4-1/2(4)=0• Formal charge on ‘3’ =6-6-1/2(2)=-1
• Formal charges do not indicate real charge separationwithin the molecule. These charges help in keepingtrack of the valence electrons in the molecule.• Formal charges help in the selection of the lowestenergy structure from a number of possible Lewisstructures for a given species.• Generally the lowest energy structure is the one withthe smallest formal charges on the atoms. The formalcharge is a factor based on pure covalent view ofbonding in which electron pairs are shared equallyby neighboring atoms.
L i m i t a t i o n s o ft h e o c t e t r u l e• It is not universal. There areexceptions to it.1)the incomplete octet of thecentral atomIn some compounds, thenumber of electronssurrounding the centralatom is less than eight.Examples- LiCl, BeH2 andBCl3
2) Odd-electron moleculein molecules with an odd number ofelectrons like nitric oxide, NO and nitrogendioxide, the octet rule is not satisfied for allatoms.
3)The Expanded OctetElements in and beyond the third period ofthe periodic table have, apart from 3s and 3porbitals, 3d orbitals also available forbonding. In a number of compounds of theseelements there are more than eight valenceelectrons around the central atom. This istermed as expanded octet.Examples-PF5, SF6, H2SO4
4) Octet rule is based on the chemical inertnessof noble gases. However, some noble gasesalso combine with oxygen and fluorine toform a number of compounds likeXeF2, KrF2, XeOF2 etc.5) This theory does not account for the shape ofthe molecules.6) It does not explain the relative stability ofthe molecules being totally silent about theenergy of a molecule.
B o n dp a r a m e t e r s• Bond lengthbond length is defined as the equilibriumdistance between the nuclei of two bondedatoms in a molecule.Bond lengths are measured by spectroscopic, X-ray diffraction techniques etc.Each atom of the bonded pair contribute to thebond length.
• In case of a covalent bond, the contributionfrom each atom is called the covalent radiusof that atom.• The covalent radius is measuredapproximately as the radius of an atom’s corewhich is in contact with the core of anadjacent atom in a bonded situation.• The van der Waals radius represents theoverall size of the atom which includes itsvalence shell in a non bonded situation.
• Bond angleIt is defined as the angle between the orbitalscontaining bonding electron pairs around thecentral atom in a molecule or complex ion.Bond angle is expressed in degree.
• Bond enthalpyIt is defined as the amount of energy required to break one mole of bondsof a particular type between two atoms in a gaseous state.The unit of bond enthalpy is kJ per mol.Example-The H-H bond enthalpy in hydrogen molecule is 435.8 kJ/mol.H2 -------> H(g)+H(g) ΔaH=435.8 kJ/ molO2 (O=O) (g) ------> O(g)+O(g) ΔaH=498 kJ/mol
• The larger the bond dissociation enthalpy, stronger will bethe bond in the molecule.HCl---H + Cl• In case of polyatomic molecules, the measurement of bondstrength in the following way-let us take the example of H2O molecule.H2O---H+OH; ΔaH1=502 kJ/molOH---H+O; ΔaH2=427 kJ/molIn this case, the average bond enthalpy is taken.average bond enthalpy=502+427=464.5 kJ/mol2
B o n d o r d e r• In Lewis description of covalent bond, thebond order is given by the number of bondsbetween the two atoms in a molecule.• For example, the bond order of h2 which has asingle shared pair of electrons is 1• The bond order of O2 with 2 shared pairs ofelectrons is 2• The bond order of N2 with 3 shared pairs ofelectron is 3.
• For N2, the bond order is 3 and its bondenthalpy is 946 kJ/mol and it is one of thehighest for diatomic molecule.• Isoelectronic molecules and ions have identicalbond orders.example- N2 and CO have bond order 3.A general correlation useful for understandingthe stabilities of molecules is that: with increasein bond order, bond enthalpy increases andbond length decreases.
R e s o n a n c es t r u c t u r e s• A single Lewis structure may be inadequate forrepresentation of a molecule.• According to the concept of resonance, whenevera single Lewis structure cannot describe amolecule accurately, a number of structureswith similar energy, position of nuclei, bondingand non-bonding pairs of electrons are taken asthe canonical structures of the hybrid whichdescribes the molecule accurately.
• For example, O3 molecule canbe represented by twostructures as shown in thefigure.• For O3 these two structuresconstitute the canonicalstructures or resonancestructures and their hybridshown by the third figurerepresents the structure of o3more accurately. this is calledresonance hybrid.
• In general it may be stated that– Resonance stabilizes the molecule as the energyof the resonance hybrid is less than the energy ofany single canonical structure– Resonance averages the bond characteristics as awhole.Resonance EnergyIt is the difference between the actual bond energy ofthe molecule and that of the most stable of theresonating structures (having least energy).
Polarity o f B o n d sThe existence of a hundred percent ionic or covalent bond represents an ideal situation. Inreality no bond or a compound is either completely covalent or ionic. Even in case ofcovalent bond between two hydrogen atoms, there is some ionic character.When covalent bond is formed between two similar atoms, for example in H2, O2, Cl2, N2 or F2,the shared pair of electrons is equally attracted by the two atoms. As a result electron pairis situated exactly between the two identical nuclei. The bond so formed is called non polarcovalent bond. Contrary to this in case of a heteronuclear molecule like HF, the sharedelectron pair between the two atoms gets displaced more towards fluorine since theelectronegativity of fluorine (Unit 3) is far greater than that of hydrogen. The resultantcovalent bond is a polar covalent bond.As a result of polarisation, the molecule possesses the dipole moment (depicted below) which canbe defined as the product of the magnitude of the charge and the distance between thecentres of positive and negative charge. It is usually designated by a Greek letter ‘μ;’.Mathematically, it is expressed as follows :Dipole moment (μ) = charge (Q) x distance of separation (r)Dipole moment is usually expressed in Debye units (D).The conversion factor is1 D = 3.33564 x 10– -30 C mwhere C is coulomb and m is meter.
• Further dipole moment is a vector quantity and is depicted by asmall arrow with tail on the positive centre and head pointingtowards the negative centre. For example the dipole moment of HFmay be represented as :• The shift in electron density is symbolised by crossed arrow ( )above the Lewis structure to indicate the direction of the shift.• In case of polyatomic molecules the dipole moment not only dependupon the individual dipole moments of bonds known as bonddipoles but also on the spatial arrangement of various bonds in themolecule. In such case, the dipole moment of a molecule is the vectorsum of the dipole moments of various bonds. For example in H2Omolecule, which has a bent structure, the two O-H bonds are orientedat an angle of 104.50. Net dipole moment of 6.17 x 10– -30 C m(1D = 3.33564 x 10– -30 C m) is the resultant of the dipolemoments of two O-H bonds.• Peter Debye, the Dutch chemist received Nobel prize in 1936 for hiswork on X-ray diffraction and dipole moments. The magnitude ofthe dipole moment is given in Debye units in order to honour him.• Net Dipole moment, μ = 1.85 D = 1.85 x 3.33564 x 10– -30 C m =6.17 10– 30 C m
• The dipole moment in case of BeF2 is zero.This is because the two equal bond dipolespoint in opposite directions and cancel theeffect of each other.• In tetra-atomic molecule, for example in BF3,the dipole moment is zero although the B – Fbonds are oriented at an angle of 120° to oneanother, the three bond moments give a netsum of zero as the resultant of any two isequal and opposite to the third.
• In case of NH3 and NF3 molecule. Both the moleculeshave pyramidal shape with a lone pair of electronson nitrogen atom. Although fluorine is moreelectronegative than nitrogen, the resultant dipolemoment of NH3 ( 4.90 x 10–3 the orbital dipole due tolone pair is in the same direction as the resultantdipole moment of the N – H bonds, whereas in NF3theorbital dipole is in the direction opposite to theresultant dipole moment of the three N – F bonds. Theorbital dipole because of lone pair decreases the effect ofthe resultant N – F bond moments, which results inthe low dipole moment of NF3.
The partial covalent character of ionic bonds was discussed byFajan’s in terms of the following rules:• The smaller the size of the cation and the larger the sizeof the anion, the greater the covalent character of an ionicbond.• The greater the charge on the cation, the greater thecovalent character of the ionic bond.• For cations of the same size and charge, the one, withelectronic configuration (n-1)dnnso, typical of transitionmetals, is more polarising than the one with a noble gasconfiguration, ns2 np6, typical of alkali and alkalineearth metal cations. The cation polarises theanion, pulling the electronic charge toward itself andthereby increasing the electronic charge between the two.This is precisely what happens in a covalentbond, i.e., build-up of electron charge density between thenuclei. The polarising power of the cation, thepolarisability of the anion and the extent of distortion(polarisation) of anion are the factors, which determine theper cent covalent character of the ionic bond.
V a l e n c e S h e l lE l e c t r o n P a i rR e p u l s i o nT h e o r y• Sidgwick and Powell in 1940, proposed a simple theorybased on the repulsive interactions of the electron pairs inthe valence shell of the atoms. It was further developed andredefined by Nyholm and Gillespie (1957).• The main postulates of VSEPR theory are as follows:• • The shape of a molecule depends upon the number ofvalence shell electron pairs (bonded or non-bonded) aroundthe central atom.• • Pairs of electrons in the valence shell repel one anothersince their electron clouds are negatively charged.• • These pairs of electrons tend to occupy such positions inspace that minimise repulsion and thus maximise distancebetween them.• • The valence shell is taken as a sphere with the electron pairslocalising on the spherical surface at maximum distance
• • A multiple bond is treated as if it is a single electron pairand the two or three electron pairs of a multiple bond aretreated as a single super pair.• • Where two or more resonance structures can represent amolecule, the VSEPR model is applicable to any suchstructure.• The repulsive interaction of electron pairs decrease in theorder:• Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair(bp) > Bond pair (bp) – Bond pair (bp)• Nyholm and Gillespie (1957) refined the VSEPR model byexplaining the important difference between the lone pairsand bonding pairs of electrons. While the lone pairs arelocalised on the central atom, each bonded pair is sharedbetween two atoms. As a result, the lone pair electrons in amolecule occupy more space as compared to the bondingpairs of electrons. This results in greater repulsion betweenlone pairs of electrons as compared to the lone pair – bondpair and bond pair – bond pair repulsions. These repulsioneffects result in deviations from idealised shapes andalterations in bond angles in molecules.
• For the prediction of geometrical shapes of moleculeswith the help of VSEPR theory, it is convenient todivide molecules into two categories as (i) moleculesin which the central atom has no lone pair and (ii)molecules in which the central atom has one or morelone pairs.• The VSEPR Theory is able to predict geometry of alarge number of molecules, especially the compoundsof p-block elements accurately. It is also quitesuccessful in determining the geometry quite-accurately even when the energy difference betweenpossible structures is very small. The theoretical basisof the VSEPR theory regarding the effects of electronpair repulsions on molecular shapes is not clear andcontinues to be a subject of doubt and discussion.
• Shape (g e o me t r y ) o f S o meS i mp l e M o l e c u l e s /I o n s w i t hC e n t r a l I o n s h a v i n g O n e o rmo r e L o n e P a i r s o fE l e c t r o n (E ).
• S h a p e s o f M o l e c u l e sc o n t a i n i n g B o n d P a i r a n dL o n e P a i r
O R B I T A LO V E R L A PC O N C E P T• If orbitals of 2 atoms are mixed with each other partially during bondformation, then the phenomenon is called as overlapping of orbitals.• TYPES OF ORBITAL OVERLAPS:• 1.Positive overlap: If the symmetry of both the atomic orbitals is the same, then itis called as positive overlap i.e. if the symmetry of the overlapping orbitals iseither positive, or negative, then it is a type of positive overlap.• 2.Negative overlap: If the symmetry of the atomic orbitals is not thesame, i.e., one is positive and the other negative, then it is a type of negativeoverlap.• 3.Zero overlap: If overlap of orbitals present in 2 different planes takes place, then it is called as zero overlap. E.g.. Px overlaps with Py (in realsituation, overlapping does not takes place.)
T Y P E S O F O V E R L A P SR E S U L T I N G I N S I G M AB O N D F O R M A T I O N
T Y P E S O FO V E R L A P SR E S U L T I N GI N P I B O N DF O R M A T I O NDue to lateral/sidewaysoverlap of P-P orbitals presentin the same plane, Pi bond isformed.
C O M P A R I S I O NSIGMA BONDS• It is a strong bond.• Electron cloud issymmetrical along the inter-nuclear axis.• There can be free rotations ofatoms around this bond.• These are less reactive.• Shape of the molecule isdetermined by these bonds.• Sigma bonds haveindependent existence.PI BONDS• It is a weak bond.• Electron cloud isasymmetrical.• Free rotation is not possiblearound this bond.• These are more reactive.• These bonds do not affect theshape of the molecule.• Pi bond always exists with asigma bond.
H Y B R I D I Z A T I O NThe intermixing of different atomic orbitals of approximately equal energy levels toproduce hybrid orbitals before bond formation is called as HYBRIDIZATION.Here arrangement of hybrid orbitals are such that there is minimum repulsion inbetween the hybrid orbitals.No. of orbitals mixed=No. of hybrid orbitals produced.DIFFERENT TYPES OF HYBRIDISATIONS:Sp HybridizationNO. of hybrid orbitals produced=2Structure = LINEARBond angle = 180 degree …..e.g. BeF₂
• SP2 Hybridization:• No. of hybrid orbitals produced = 3• Arrangement of these orbitals = TRIGONAL PLANAR• Bond angle=120 degree……..ex. BF₃, etc.• SP3 Hybridization:• No. of hybrid orbitals = 4• Arrangement= TETRAHEDRAL• BOND ANGLE…..• IN CH₄ = 109.28 degree• IN NH₃ = 107.3 degree.• Here, in methane, sigma bond is formed between H and C atom due tooverlapping of sp3 orbital of H atom with S orbital of H atom.• Structure of NH₃ – TRIANGULR PYRAMIDAL.
• SP3d Hybridization:• NO. of hybrid orbitals produced = 5• Structure = TRIGONAL BIPYRAMIDAL• BOND ANGLES:• Equatorial = 120 degree.• Axial = 90 degree.• If lone pair of electron is present at central atom, its position is alwaysequatorial.Ex.PCl5• Length of axial bonds is longer than that of equatorial bonds because ofminimum repulsion.• IT may have different shapes according to no. of lone pairs it has:• 1 lone pair – seesaw shape………e.g. SF₄• 2 lone pairs – bent T shape………E.g. BrF₃• 3 Lone pairs – Linear shape
• SP3d2 Hybridization:• No. of hybrid orbitals produced = 6• Arrangement = OCTAHEDRAL……e.g. SF₆ ( lone pairs = 0)• Shape may be:• Square pyramidal…..e.g. BrF₅ ( lone pairs = 1)• Square planar…….e.g. XeF4₄ (lone pairs = 2)• SP3d3 Hybridization:• No. of hybrid orbitals = 7• Arrangement = PENTAGONAL BIPYRAMIDAL• BOND ANGLES:• Axial with equatorial = 90 degree.• Equatorial to equatorial = 72 degree.
R U L E S R E G A R D I N GH Y B R I D I Z A T I O N• Only orbitals of approximately same energy levels can take part.• No. of orbitals mixed = No. of hybrid orbitals produced.• Most hybrid orbitals are similar but not always identical in shape. Theymay differ from one another in their orientation in space.• The electron waves in hybrid orbitals repel each other and this tend to thefarthest apart.• Hybrid orbitals can form only sigma bonds.• Depending on the number and the nature of the orbitals undergoinghybridization, various types of hybrid orbitals directing towards thecorners of specified geometrical figures come into existence.
• C O N D I T I O N S F O RC O M B I N A T I O N O FA T O M I C O R B I T A L S• For atomic orbitals to combine, resulting in theformation of molecular orbitals , the main conditionsare :• The combining atomic orbitals should have almost thesame energies. For example, in the case of diatomicmolecules, 1s-orbital of one atom can combine with 1s-orbital of the other atom, but 1s-orbital of one atomcannot combine with 2s-orbital of the other atom.• The extent of overlap between the atomic orbitals of thetwo atoms should be large.• The combining atomic orbitals should have the samesymmetry about the molecular axis. Forexample, 2Pxorbital of one atom can combine with2Px orbital of the other atom but not with 2Pz orbital .• Note : It may be noted that Z-axis is taken as theinter-nuclear axis according to modern conventions.
D e s i g n a t i o n s o fM o l e c u l a r O r b i t a l sJust as atomic orbitals are designated ass, p, d, f etc molecular orbitals of diatomic molecules arenamed σ (sigma) ,π (pi) , δ (delta) etc.M O L E C U L A R O R B I T A L SThe molecular orbitals which are cylindrically symmetricalaround inter-nuclear axis are called σ - molecular orbitals.The molecular orbital formed by the addition of 1s orbitalsis designated as σ 1s and the molecular orbital formed bysubtraction of 1s orbitals is designated as σ * 1s .Similarly combination of 2s orbital results in theformation of two2 s - molecular orbitals designated as σ 2s and σ * 2s
1.Determine the number of electrons in the molecule. We get thenumber of electrons per atom from their atomic number on theperiodic table. (Remember to determine the total number of electrons,not just the valence electrons.)2.Fill the molecular orbitals from bottom to top until all the electronsare added. Describe the electrons with arrows. Put two arrows in eachmolecular orbital, with the first arrow pointing up and the secondpointing down.3. Orbitals of equal energy are half filled with parallel spin beforethey begin to pair up.
Stability of the molecule with bond order.Bond order = 1/2 (#e- in bonding MOs - #e- inantibonding MOs)We use bond orders to predict the stabilityof molecules :-• If the bond order for a molecule is equal to zero, themolecule is unstable.• A bond order of greater than zero suggests a stablemolecule.• The higher the bond order is, the more stable the bond.We can use the molecular orbital diagram to predict whetherthe molecule is paramagnetic or diamagnetic. If all theelectrons are paired, the molecule is diamagnetic. If one ormore electrons are unpaired, the molecule is paramagnetic.
1. The molecular orbital diagram for a diatomic hydrogen molecule, H2, is• The bond order is 1. Bond Order = 1/2(2 - 0) = 1• The bond order above zero suggests that H2is stable.• Because there are no unpaired electrons, H2 is diamagnetic.2. The molecular orbital diagram for a diatomic helium molecule, He2, shows the following.• The bond order is 0 for He2. Bond Order = 1/2(2 - 2) = 0• The zero bond order for He2suggests that He2is unstable.• If He2did form, it would be diamagnetic
3. The molecular orbital diagram for a diatomic oxygen molecule, O2, is• O2has a bond order of 2. Bond Order = 1/2(10 - 6) = 2• The bond order of two suggests that the oxygen molecule is stable.• The two unpaired electrons show that O2is paramagnetic
Diatomic molecules are molecules composed only of two atoms,of either the same or different chemical elements. The prefix di-is of Greek origin, meaning two. Common diatomic moleculesare hydrogen (H2), nitrogen (N2), oxygen (O2), and carbonmonoxide (CO). Seven elements exist as homonuclear diatomicmolecules at room temperature: H2, N2, O2, F2, Cl2, Br2, and I2.Many elements and chemical compounds aside from theseform diatomic molecules when evaporated. The noble gases donot form diatomic molecules: this can be explained usingmolecular orbital theory (see molecular orbital diagram).
INTRODUCTION1.Two H atoms in their ground state configurationcome together and form a single bond. The bondformation stabilizes both atoms and, therefore, is lowerin energy than the atomic orbitals. This is alsoobserved in Valence Bond Theory, which implies thateach H atom in H2shares its electron with one another,so that both can achieve the stable configuration of He.2.On top of that, MO Theory allows one to compute theamount of energy released from a bond formation anda distance between two bonded atoms as well as predictthe magnetic property of a molecule (or a substance).For H2, the bond strength is -432 kJ/Mol, and thebond length is 74 angstrom (or 74 pm). H2 is adiamagnetic molecule because the electrons paired up;therefore, it is not attracted by a magnetic field.
B o n d i n g a n d A n t i -b o n d i n gm o l e c u l a r o r b i t a l s i n H 21.Each H atom has a 1s atomic orbital. When two H atoms come to a proper proximity, their 1sorbitals interact and produce two molecular orbitals: a bonding MO and an anti-bonding MO.2.If the electrons are in phase, they have a constructive interference. This results in a bondingsigma MO (σ1s). This MO has an increased probability of finding electrons in the bondingregion.Figure 2: Schematic representation of the bonding molecular orbital σ(1s)If the electrons are out of phase, they have a destructive interference. This results in an anti-bonding sigma MO (σ*1s). This MO has a decreased probability of finding electrons in thebonding region. (Valence Bond Theory does not explain this phenomenon.)Figure 3:Schematic representation of antibonding molecular orbital σ*(1s)Note that there is a nodal plane in the anti-bonding
B o n d o r d e r i n H 2Bond order = 1/2 (#e- in bonding MO - #e- in antibonding MO)For H2, bond order = 1/2 (2-0) = 1, which means H2 has only onebond. The antibonding orbital is empty. Thus, H2 is a stablemolecule.Again, in the MO, there is no unpaired electron, so H2 isdiamagnetic
H Y D R O G E N B O N DIn compounds of hydrogen with strongly electronegativeelements, such as fluorine, oxygen andnitrogen, electron pair shared between the two atomslie far away from the hydrogen atom. As a result, thehydrogen atom becomes highly electropositive withrespect to the other atom. This phenomenon of chargeseparation in the case of hydrogen fluoride isrepresented as . Such a molecule is said to be polar .The molecule behaves as a dipole because one endcarries a positive charge and the other end a negativecharge. The electrostatic force of attraction betweensuch molecules should be very strong. This isbecause the positive end of one molecule is attracted bythe negative end of the other molecule . Thus, two ormolecules may associate together to form largercluster of molecules. This is illustrated below for theassociation of several molecules of hydrogen fluoride.
• The cluster of HF molecules may be described as(HF)n.• It may be noted that hydrogen atom is bonded tofluorine atom by a covalent bond in one molecule andby electrostatic force or by hydrogen bond to thefluorine atom in the adjacent molecule . Hydrogenatom is thus seen to act as a bridge between the twofluorine atoms.
• The hydrogen bond is represented by a dotted line. Thesolid lines represent the original(covalent ) bondpresent in the molecule.• Chlorine, bromine and iodine are not as highlyelectronegative as fluorine and therefore, the sharedpair of electrons in the case of HCl , HBr and HI donot lie as far away from hydrogen as in the case ofHF. The tendency to form hydrogen bond in thesecases is therefore less.• Water molecule, because of its bent structure, is also adipole, oxygen end carrying a negative charge andhydrogen end carrying a positive charge. Hydrogenbond taking place in this case as well, as representedbelow:• The cluster of water molecules may be described as(H2O)n
• The nature of hydrogen bond• The hydrogen bond is a class in itself. It arises fromelectrostatic forces between positive end (pole) of onemolecule and the negative end(pole) of the othermolecule generally of the same substance. Thestrength of hydrogen bond has been has been found tovary between 10 - 40 kJ mol−1 (i.e., 6.02 x1023 bonds) while that of a covalent bond has beenfound to be of the order of 400 kJ mol−1 . Thus ahydrogen bond is very much weaker than a covalentbond. Consequently, the length of hydrogen bond isbigger than the length of a covalent bond.• In the case of hydrogen fluoride, for instance, whilethe length of the covalent bond between F and H atomsis 100 pm, the length of hydrogen bond between F andH atoms of neighbouring molecules is 155 pm.
• T y p e s o f h y d r o g e nb o n d i n g• Hydrogen bonding may be classified intotwo types :• I n t e r m o l e c u l a rh y d r o g e n b o n d i n gThis type of hydrogen bonding involves electrostaticforces of attraction between hydrogen andelectronegative element of two different molecules ofthe substance. Hydrogen bonding in molecules ofHF, NH3 , H2O etc. are examples of intermolecularhydrogen bonding.
• I n t r a m o l e c u l a rh y d r o g e n b o n d i n gThis type of bonding involves electrostatic forces ofattraction between hydrogen and electronegativeelement both present in the same molecule of thesubstance. Examples o-nitrophenol andsalicylaldehyde.• p-Nitrophenol , on account of large distance betweentwo groups , does not show any intramolecularhydrogen bonding. On the other hand, it shows theusual inter molecular hydrogen bonding , asillustrated below:
• As a result of intermolecular hydrogen bonding, thepara derivative undergoes association, resulting in anincrease in molar mass and hence an increase inboiling point. In ortho derivative, on accountof intramolecular hydrogen bonding , no suchassociation is possible. Consequently, the orthoderivative is more volatile than the para derivative.Thus, while ortho nitrophenol is readily volatile insteam , para nitrophenol is completely non-volatile.The two derivatives can thus be separated from eachother by steam distillation.
• Density in solid state(ice) is less than that in liquidstate . This is some what unusual because in mostsubstances density in solid is more than that inliquid state.• Water contracts when heated between 0°C and 4°C .This is again unusual because most substancesexpand when heated in all temperature ranges.• Both these peculiar features are due to hydrogenbonding, as discussed below :
• In ice, hydrogen bonding between H2Omolecules is more extensive than in liquidwater. A substance in solid state has adefinite structure and the molecules aremore rigidly fixed relative to one anotherthan in the liquid state. In ice, the H2Omolecules are tetrahedrally oriented withrespect to one another.At the same time, each oxygen atom is surroundedtetrahedrally by four hydrogen atoms, twoof these are bonded covalently and theother two by hydrogen bonds.Thetetrahedral open cage-like crystal structureof ice. The central oxygen atom A issurrounded tetrahedrally by the oxygenatoms marked1,2, 3 and 4.The hydrogen bonds areweaker and therefore, longer than covalentbonds. This arrangement gives rise to anopen cage-like structure , as shown in theFig. There are evidently a number of ‘holesor open spaces.
• These holes are formed because the hydrogen bondsholding the H2O molecules in ice are directed incertain definite angles . In liquid water suchhydrogen bonds are fewer in number. Therefore, as icemelts, a large number of hydrogen bonds are broken.The molecules, therefore, move into the ‘hole or openspaces and come closer to one another than they werein the solid state. This results in a sharp increase indensity . The density of liquid water is, thereforehigher than that of ice.• As liquid water is heated from 0°C to 4°C, hydrogenbonds continue to be broken and the molecules comecloser and closer together. This leadsto contraction. However, there is some expansion ofwater also due to rise in temperature as in otherliquids. It appears that up to 4°C, the former effectpredominates and hence the volume increases as thetemperature rises.
• It can be easily realised that withouthydrogen bonding , water would have existedas a gas like hydrogen sulphide. In that caseno life would have been possible on this globe.• Hydrogen bonding also exists in all livingorganisms, whether of animal or of vegetablekingdom. Thus, it exists in varioustissues, organs, blood, skin and bones inanimal life. It plays an important role indetermining structure of proteins which areso essential for life.
• Hydrogen bonding plays an important role inmaking wood fibres more rigid and thus makes it anarticle of great utility. The cotton, silk or syntheticfibres owe their rigidity and tensile strength tohydrogen bonding. Thus hydrogen bonding is ofvital importance for our clothing as well. Most of ourfood materials also consists of hydrogen bondedmolecules. Sugars and carbohydrates , for example,have many -OH groups. The oxygen of one such groupin one molecule is bonded with -OH group of anothermolecule through hydrogen bonding. Hydrogenbonding is thus a phenomenon of great importance inevery day life.