JEE Chemical Bonding Final PDF

Welcome to Chemical Bonding and Molecular Structure NaCl CaCO3 Iodine crystals Aluminium The substances are formed as a result of combination of atoms or molecules or ions. So Many Why? Why do some atoms combine while certain others do not? What is the nature of the force that exists between combini...

Welcome to Chemical Bonding and Molecular Structure NaCl CaCO3 Iodine crystals Aluminium The substances are formed as a result of combination of atoms or molecules or ions. So Many Why? Why do some atoms combine while certain others do not? What is the nature of the force that exists between combining atoms? Why does definite number of various atoms constitute a particular molecule? Why is it easy for some elements to lose electrons while it is harder for others? Chemical Bond The attractive force which holds various constituents (atoms, ions, etc.) together, in different chemical species. A chemical bond forms in order to reduce the energy of the chemical species involved in bonding, thereby increasing their stability. Potential energy (kJ/mol) Potential Energy Curve Internuclear distance, r (pm) r0 0 r = 300 r = 150 H - 435.8 r = 74 H H Observed bond distance in H2 H H H At 74 pm, two H atoms have together minimum potential energy (hence maximum stability in the form of a H2 molecule). Chemical Bonds Ionic Bond Covalent Bond Metallic Bond Electropositive & Electronegative atoms Two electronegative atoms Two electropositive atoms How Ionic Bonds are Formed? Ionic Bond Electrostatic force of attraction between oppositely charged ions. To attain stable electronic configuration Elements lose or gain electron(s) in order to have a stable electronic configuration in their valence shell Formation of Ionic Bond Elements involved in the ionic bond should possess (1) (2) The element losing electron should have LOW ionisation enthalpy. The element accepting electron should have HIGH electron gain enthalpy. (3) (4) Large difference in electronegativity of two elements High Lattice enthalpy Covalent Bond and Coordinate Bond A Covalent bond is formed by sharing electrons. A bond in which the shared pair of electrons originate from one atom and none from the other is called coordinate bond. Shared pair of electrons H H H Covalent Bond Coordinate Bond + O H H Metallic Bonds Electrostatic force of attraction between a metal kernel and valence electrons. Lattice Energy Energy required to completely separate one mole of solid ionic compound into gaseous constituent ions. NaCl (s) AB(s) A+(g) + B─(g) Na+ (g) + Cl─ (g) A+(g) + B─(g); Lattice Energy = Positive AB(s); Lattice Energy = Negative Factors Affecting Lattice Energy Lattice Energy (L.E.) ∝ 1 r r＋ + r─ = r r Interionic distance r＋ Radius of the cation rー Radius of the anion Lattice Energy (L.E.) Z+ Z─ ∝ Z+×Z Charge on the cation Charge on the anion ─ Lattice Energy Charge is the deciding factor Ionic compound r (Å) Z+ × Z ─ L.E. (kJ mol─1) LiF 2.01 1 1004 kJ mol─1 MgO 2.10 4 3933 kJ mol─1 Order of Lattice Energy BaO KI < KBr < KCl < KF < SrO < CaO < MgO < NaF < MgF2 < AlF3 BeO Properties of Ionic Compounds 01 Soluble in polar solvent High Melting point and Boiling point 03 02 Conduct Electricity in aqueous & molten state Exhibit isomorphism & polymorphism 05 04 Forms crystal in Solid state Highly Polar 06 Isomorphism Different ionic compounds having similar crystal structure are called isomorphs and this phenomenon is called isomorphism. Isomorphous compounds have the same type of formula. FeSO4.7H2O Green vitriol MgSO4.7H2O Epsom salt ZnSO4.7H2O White vitriol Polymorphism Occurence of a particular substance in more than one crystalline form is called polymorphism For example, ZnS exist as (1) (2) Sphalerite Wurtzite Hydration Energy Energy released when one mole of a gaseous ion is hydrated in large amount of water to form an aqueous ion. Size of ion Hydration energy Charge of ion Hydration energy As the dielectric constant of solvent increases, more energy is released on solvation. Solvation The interaction of the solute and the solvent molecules which stabilizes the solute in the solution If the solvent is water, then it is known as Hydration Dissolution of solute in water depends on Lattice energy and Hydration energy Types of Covalent Bond 01 Formed by the mutual sharing of electrons between two atoms. Number of shared pair of electrons Single Bond 0 2 0 3 Double Bond Triple Bond Types of Covalent Bond (1) One pair of electrons is shared Three pairs of (3) electrons are shared (2) Two pairs of electrons are shared Triple Covalent bond Single Covalent bond Double Covalent bond Theories of Covalent Bonding Kössel-Lewis Electronic theory MOT VBT VSEPR Kossel Lewis Electronic Theory Atoms can combine by the transfer of valence electrons from one atom to another or by sharing of electrons. Lewis Dot Structures Combining atoms contribute ≥ 1 electron(s) to the shared pair Covalent Bond Shared pair of electrons Shared pair of electrons (e─) O2 8e─ 8e─ Octet complete Attains noble gas configuration 01 Formal Charge It’s a theoretical charge over an individual atom of a molecule or an ion. 02 03 In polyatomic ions, the net charge is possessed by the whole ion. Feasible to assign a formal charge on each atom. Formal Charge 2 1 2 1 3 3 F.C. on O (1) = -1 F.C. on O (2) = +1 F.C. on O (3) = 0 Used to give the relative stability of possible Lewis structures Lowest energy structure: Smallest formal charge on the atoms Limitations of Octet Rule 1 Molecules with incomplete octet of the central atom Hypovalent compound Electrons around central atom < 8 4 valence electrons 6 valence electrons Limitations of Octet Rule 2 Molecules with odd electrons NO, NO2, ClO2, ClO3 3 Molecules with expanded octet Hypervalent compound 10 valence electrons Electrons around central atom > 8 Limitations of Octet Rule 4 Formation of Xe & Kr compounds Xe and Kr form compounds with F and O even though their octet is already complete. 5 Doesn’t account for the shape of the molecules 6 Doesn’t explain about the relative stability of the molecules Valence Bond Theory Linus Pauling J.C. Slater A covalent bond is formed by the overlap of half filled atomic orbitals that yield a pair of electrons shared between the two bonded atoms. Orbital Overlap Types of Overlap Positive Bond will be formed Negative Bond will not be formed All orbital overlappings do not result in bond formation Zero Directional Properties of Bonds Covalent Bond: Directional Ionic Bond: Non-Directional Coordinate or Dative Bond Bond formed by sharing of electrons between two atoms. Shared pair of electrons is contributed by only one of the two atoms. Co ordinate bond once formed cannot be distinguished from covalent bond. Covalent and coordinate bond are same with respect to bond properties. How to Identify Coordinate Bond? Different covalency than usual Presence of coordinate bond Example: [H3O+] Example: N2O Fourth bond formed by N is always coordinate N + N - O Third bond formed by O is always coordinate H + O H H Lewis Acid and Lewis Base Lone pair acceptors are called Lewis acids Lone pair donors are called Lewis bases H3N Donor + H+ Acceptor [NH4]+ Co-ordinate Bond or Dative Bond + H H O H + H+ H O H or H3O+ Sigma and pi-bonds Covalent Bond Sigma (σ) bond Sigma bond is formed when overlapping takes place along the internuclear axis of orbitals or when an axial overlap takes place. Pi (𝛑) bond Pi (𝛑) bond is formed when axes of combining orbitals are perpendicular to the internuclear axis i.e., lateral or sidewise overlapping takes place. Axial or Head-on Overlapping Cylindrically symmetrical about the internuclear axis Can undergo rotation about the internuclear axis Generally, 𝛑 bond between two atoms is formed in addition to a σ bond Types of π bond p𝛑-p𝛑 overlap p𝛑-d𝛑 overlap d𝛑-d𝛑 overlap Bond Strength In general, order of strength of bond 𝛑 < 𝞂 Greater the extent of overlapping, more will be the bond strength. For same value of n, s-s sigma overlap < s-p sigma overlap < p-p sigma overlap Strength of 𝜋 bonds: 3p-3p 𝜋 overlap < 2p-2p 𝜋 overlap Bonding in H2 Molecule 1 H 1s1 H ↿⇂ H 1 H 1s1 i.e. H 𝜎 1s H 1s Bonding in O2 Molecule O O ⥮ 1s2 2s2 ⥮ ⥮ 1s2 i.e. ⥮ O 2s2 𝜎 𝜋 O ⥮ ↿ ↿ 2p4 ⥮ ↿ 2p4 ↿ ↿ O ↿⇂ O ↿ 2p 2p Bonding in N2 Molecule N ⥮ 1s2 N ⥮ 1s2 i.e. N ⥮ ↿ 2s2 ↿ N ↿ 2p3 ↿ ↿ 2p3 ↿ 2s2 2𝜋 ↿ N ⥮ 𝜎 ↿ ↿ ↿ ↿⇂ N Limitations of VBT Paramagnetic nature of O2 could not be explained. Fails to account for the geometry and shapes of various molecules. Steps to Draw Structures 01 Largest size Select the central atom SO3 - Sulphur Central atom Least electronegative Less in number Can form maximum bonds Steps to Draw Structures Draw a symmetrical skeleton of atoms. 02 03 Calculate the total number of valence electrons O Valence electrons in SO3 = 6 +(3×6) = 24 S O O Electron pairs = 24 = 12 2 Steps to Draw Structures.. : O: S.. : O:.. : O: 04 Make a single bond using the electron pairs. Then complete the octet of the side atoms. If any electron pair is left, assign it to the central atom. Steps to Draw Structures 05 06 If the octet of central atom is not complete, use the lone pairs of side atoms to make the bonds and complete the octet. Assign formal charge on each atom VSEPR Theory Used to provide shape and electronic geometry of covalent compounds. 1. Shape of a molecule depends upon the number of valence shell electron pairs around the central atom 2. Valence shell is taken as a sphere with the electron pairs localising on the spherical surface 3. Electron pairs in the valence shell repel one another since, they are all negatively charged VSEPR Theory 4. Electron pairs occupy positions in space that tend to minimise repulsion. 5. Lone pair occupies more space on the sphere. So, the order of repulsion is: lp-lp > lp-bp > bp-bp (lp: Lone pair, bp: Bonding pair) VSEPR Theory 6. A multiple bond is treated as a single bonding pair. There is no effect of pi bond on geometry and shape Compound Shape 𝛑 bond(s) Cl - Be - Cl + O=C-H Linear 0 Linear 1 O=C=O Linear 2 H-C≡N Linear 2 Point to Remember!! For electronic geometry For shape Both bond pairs and lone pairs are considered Only bond pairs are considered VSEPR Theory (A: Central atom, B: Side atom) General Formula: AB2 Electron pairs Bonding pairs Lone pairs Electronic Geometry Shape 2 2 0 Linear Linear Bond Angle = 180o Example: CO2, BeCl2 VSEPR Theory (A: Central atom, B: Side atom) General Formula: AB3 Electron pairs Bonding pairs Lone pairs Electronic Geometry Shape 3 3 0 Trigonal Planar Trigonal Planar Bond Angle = Example: BF3, SO3 120o VSEPR Theory General Formula: AB2L (A: Central atom, B: Side atom, L: Lone pair) Electron Pairs Bonding Pairs Lone Pairs Electronic Geometry Shape 3 2 1 Trigonal Planar Bent/ V-Shape Bond Angle < 120o Example: SO2, SnCl2 VSEPR Theory Steric Number: 4 (sp3) (AB4) 109.50 Bond Angle = 109.5o Example: CH4, [NH4]+ , XeO4 VSEPR Theory General Formula: AB3L (A: Central atom, B: Side atom, L: Lone pair) Electron Pairs Bonding Pairs Lone Pairs Electronic Geometry Shape 4 3 1 Tetrahedral Pyramidal Bond Angle < 109.5o Example: NH3, XeO3, PCl3 VSEPR Theory General Formula: AB2L2 (A: Central atom, B: Side atom, L: Lone pair) Electron Pairs Bonding Pairs Lone Pairs Electronic Geometry Shape 4 2 2 Tetrahedral Bent or V-Shape Bond Angle < 109.5o Example: H2O, OF2 VSEPR Theory General Formula: AB5 Electron Pairs 5 Bonding Pairs 5 (A: Central atom, B: Side atom) Lone Pairs Electronic Geometry Shape 0 Trigonal Bipyramidal (T.B.P.) Trigonal Bipyramidal (T.B.P.) Bond Angle = 90o, 120o, 180o Example: PCl5, SOF4 Valence Shell Electron Pair Repulsion Theory (VSEPR) Steric Number: 5 (sp3d) (AB4L) Bond Angle < 90o , 120o Example: SF4, XeO2F2 VSEPR Theory General Formula: AB3L2 (A: Central atom, B: Side atom, L: Lone pairs) Electron Pairs Bonding Pairs Lone Pairs Electronic Geometry Shape 5 3 2 Trigonal Bipyramidal (T.B.P.) T-Shape Bond Angle < 90o, 180o Example: ClF3 VSEPR Theory General Formula: AB2L3 Electron Pairs 5 Bonding Pairs 2 (A: Central atom, B: Side atom, L: Lone pairs) Lone Pairs Electronic Geometry Shape 3 Trigonal Bipyramidal (T.B.P.) Linear AB2L3 Bond Angle = Example: XeF2, I3─ 180o VSEPR Theory Steric Number: 6 (sp3d2) (AB6) 900 Bond Angle = Example: SF6 90o VSEPR Theory General Formula Electron Pairs Bonding Pairs Lone Pairs Electronic Geometry Shape AB5L 6 5 1 Octahedral Square Pyramidal AB4L2 6 4 2 Octahedral Square Planar (A: Central atom, B: Side atom, L: Lone pairs) AB5L Bond Angle < 90o, 180o Example: BrF5, XeOF4 AB4L2 Bond Angle = 90o, 180o Example: XeF4 VSEPR Theory Steric Number: 7 (sp3d3) (AB7) 900 720 Bond Angle = Example: IF7 72o, 90o Need for Hybridisation Shape of CH4 Molecule Electronic configuration of carbon is: CGround State : [He] 2s2 2p2 ⥮ ⥮ 1s2 2s2 ↿ ↿ 2p2 Release of energy due to overlap between the orbitals of C and H On excitation, CGround State : [He] 2s2 2p2 CExcited state : [He] 2s1 2p3 ⥮ ⥮ 1s2 2s2 ⥮ ↿ 1s2 2s1 ↿ ↿ 2p2 ↿ ↿ 2p3 ↿ Shape of CH4 Molecule H C-H bonds formed by s (H)-p (C) overlapping H H Angle between them will be 90o H H H H 4th C-H Bond will be formed by s (C)-s (H) overlap Shape of CH4 Molecule s-orbital overlap can be in any direction Direction of fourth C–H bond cannot be determined Shape of CH4 Molecule Expected observations All H-C-H bond angles are not certain 3 C-H bonds formed by s-p overlap are stronger than 1 C-H bond formed by s-s overlap Experimental observations All H-C-H bond angles are identical with a value of 109.5° All C-H bond lengths and bond strengths are identical Limitations of VBT Formation of diatomic molecules are satisfactorily explained (except the paramagnetic nature of O2) VBT fails to explain the bond properties in polyatomic molecules Hybridization + VBT More complete theory to explain polyatomic molecules Hybridisation Pauling (1) Intermixing of atomic orbitals of equal or slightly different energies, results in the formation of new set of orbitals of equivalent energies and shape. (2) The orbitals present in the valence shell (and sometimes penultimate shell also) of the atom can hybridise. J.C. Slater Hybridisation and Shape of Hybrid Orbitals Number of hybrid orbitals (H.O.) = Number of atomic orbitals intermixing Larger lobe of H.O. takes part in bond formation (σ-bond) Did you Know? Hybrid Orbital Actual Shape Shape used for representation Naming of Hybrid Orbitals On the basis of atomic orbitals participating in hybridization: s-orbital + ‘n’ p-orbital spndm hybrid orbital + ‘m’ d-orbital Types of Hybridisation sp sp2 sp3 sp3d sp3d2 sp3d3 Important Conditions for Hybridisation Orbitals can Have a pair of electrons Have an unpaired e– All three types can undergo hybridisation Be vacant Salient Features of Hybridisation Promotion of electron is not an essential condition prior to hybridisation. Orbitals undergo hybridisation and not the electrons. Hybrid orbitals generally form 𝛔 bond. Participating atomic orbitals Number of hybridised orbitals Hybridisation 2 sp One s + One p + % s Character %s character sp = > Number of s orbitals Number of (s+p) orbitals sp2 > Decreasing order of s character ✕ 100 sp3 Percentage Character of Orbitals sp2 sp 50% s 50% p 33.33% s % p or % d character increases Orbital becomes bulkier and shorter Orbital becomes thinner and longer Energy of hybrid orbital decreases sp3 25% s 66.66% p % s character increases 75% p Electronegativity increases Energy of hybrid orbital increases Electronegativity decreases Features of Hybridisation % s character in hybrid orbital sp - p Stability of hybrid orbital > sp2 - p > sp3 - p Bond Strength > p- p Decreasing order of bond strength Hybrid orbitals are directed in space in a way to have minimum repulsion between the electron pairs in order to obtain a stable arrangement Steric Number Type of hybridisation is estimated by steric number Steric Number = Number of σ bonds of central atom Steric Number = + Number of lone pairs on central atom V+M-q 2 Type of hybridization indicates the geometry of the molecule Steric Number, Hybridization and Geometry Steric number Hybridization Geometry Involving Orbitals 2 sp Linear s, px/ pz/ py 3 sp2 Trigonal Planar s, px, pz/ py, pz/px, py 4 sp3 Tetrahedral s, px, pz, py Trigonal bipyramidal s, px, pz, py, dz2 Octahedral s, px, pz, py, dz2, dx2-y2 Pentagonal bipyramidal s, px, pz, py, dz2, dx2-y2 , dxy 3 5 sp d 6 3 2 3 3 7 sp d sp d Methods for Finding Hybridisation Determination of type of hybridisation Orbital Box Diagrams Finding Steric Number Formula for Steric Number sp Hybridisation Participating atomic orbitals Number of hybridised orbitals Hybridisation 2 sp One s + One p + 180° Linear +─ ─ + sp2 Hybridisation Atomic orbitals participating in hybridisation Number of hybridised orbitals Hybridisation 3 sp2 One s + two p 3 new sp2 hybridised orbitals Trigonal planar + ─ ─ + ─ 120° + sp2 Hybridisation sp3 Hybridisation Atomic orbitals participating in hybridisation Number of hybridised orbitals Hybridisation 4 sp3 One s + three p sp3 Hybridisation 4 new sp3 hybridised orbitals 109.5° + Tetrahedral + + + Bonding of CH4 Molecule Here, electronic configuration of carbon is Needs energy 2s2 CGround State : [He] 2s2 2p2 CExcited state : [He] 2s1 2px1 2py1 2pz1 ⥮ ↿ 2p2 ↿ ↿ ↿ ↿ ↿ Bent’s Rule In T.B.P. geometry, Equatorial position Axial position High % s character Occupied by lone pair, multiple bond Low % s character Occupied by more electronegative element Berry Pseudorotation Fluxional behaviour of PF5 due to fast exchange between axial and equatorial F atoms All P-F bonds are observed to be equivalent Energy difference b/w T.B.P. and square pyramidal geometry < Thermal energy at room temperature Examples of sp3d Hybridisation sp3d2 Hybridisation SF6 SGround state : [Ne] 3s2 3p4 On excitation: ⥮ ⥮ 3s ↿ ↿ 3p 3d Hybridized orbitals SExcited state : [Ne] 3s1 3p3 3d2 ↿ 3s ↿ ↿ 3p ↿ ↿ ↿ 3d sp3d2 Hybridisation No equatorial & no axial bonds SF6 XeOF4 All Bond lengths are observed to be identical XeF4 Hybridisation in Odd Electron Species F F C F Hybridisation: sp3 H C H H Hybridisation: sp2 Did you Know? PCl5(s) PBr5(s) PI5(s) [PCl4]+ [PCl6]─ sp3 sp3d2 [PBr4]+ [Br]─ sp3 Does not exist P cannot accomodate 6 large sized Br & I. So, it cannot form PX6─. Resonance Phenomenon of delocalisation of 𝞹 electrons If a single Lewis structure cannot represent a molecule Most important parameter to explain the stability of certain molecule Resonance structures describe the molecule accurately Need of the theory ＋ ＋ Resonance Structures of O3 ─ ─ ＋ Resonance Hybrid of O3 ─ 1/2 ─ 1/2 Resonance Molecules having Resonance Resonance Structures (R.S.) Resonance Hybrid (R.H.) Similar or degenerate Energy Actual structure of different possible structures Identical positions of nuclei Does not violate the rules of covalence maxima Same number of bonding & nonbonding electron pairs Most stable R.S. contributes maximum towards Resonance hybrid (R.H.) Resonance Structures (R.S.) and Resonance Hybrid (R.H.) + + + ─ ─ Resonance Structures of O3 _ 1 2 _ Resonance hybrid of O3 1 2 Characteristics of Resonance Resonance structures (R.S.) are hypothetical Resonance hybrid (R.H.) has its individual identity Conditions of Resonance Structure should be planar Occurs in adjacent parallel p-orbitals Bond lengths of R.H. are intermediate to those of R.S. Rules for Resonance Structures A Must have proper Lewis structures B Only π electrons or lone pairs can be moved C Overall charge of the system must remain the same D Bonding framework of a molecule must remain intact Benzene 1.5 4 C6H6 Å Hybridize 1. 3 4Å 2 ⥮ 1s2 2 2 CGround state : 1s 2s 2p C6H6 ⥮ 2 2s CExcited state : 1s2 2s1 2p3 ↿ ↿ 2p2 C (Excited and hybrid state) ⥮ ↿ 1s2 2s1 ↿ ↿ ↿ ↿ 2p3 ↿ ↿ 3 sp2-hybrid orbitals ↿ 2p Unhybridized p-orbital Benzene Each carbon of benzene has one unhybridised p-orbital Unhybridised p-orbitals of each carbon are parallel to each other Benzene Localised 𝛑 bonds Delocalised 𝛑 bonds Resonance Structures and Resonance Hybrid of Benzene 1.5 4 Å 1. 3 4 Å 1.3 9 Å Bond order of Resonance Structures For degenerate resonating structures: Bond Order between two atoms = Total no. of bonds between 2 atoms in all structures ÷ Total no. of resonance structures Less Contributing Less Contributing [ ] C = O bond length in CO2 is less than expected because of resonance Drago’s Rule ❖ Hybridization does not take place for compounds of elements of 3rd period onwards, bonded to a less electronegative element like hydrogen. ❖ It is because energy difference between participating orbitals is very high. Lewis Acid and Lewis Base 01 Lone pair donors are also called as Lewis base Lone pair acceptors are also called as Lewis acid H3N Donor + H+ Acceptor 02 NH4+ Lewis Acids BF3, AlCl3, BCl3 SiF4, PCl5, PF5 Mg2+, Al3+, Fe2+, Zn2+ CO2, SO2 Incomplete octet Central atom has vacant d-orbitals Metal cations Central atom is attached to a more E.N. atom with multiple bonds Lewis Bases NH3, H2O, ROH, RNH2 Cl–, F–, OH–, NH2– Central atom has at least one lone pair and is surrounded by less E.N. atom Anions Back Bonding Coordinate Bond 𝛔 - type Coordinate Bond 𝛑 - type 𝛔 - Coordinate Bond F H N H 𝝈-type F BF3 H3N B F H Filled ‘p’ F H F Vacant ‘p’ Lewis Base B ─ ＋ N Lewis Acid F H H 𝛑 - Coordinate Bond Back bonding Back bond forms between Kind of coordinate 𝛑 bonding Atom having non bonded electron pair & Partial double bond character Atom having vacant orbital 𝛑 - Coordinate Bond P H H Back Bonding Because Hydrogen has no lone pairs P H F F F Back Bonding Because F has lone pairs and can donate to one empty p-orbital of P CHCI3 ⇌ CHF3 ⇌ H+ H+ + + - CCI3 - CF3 F atom does not have vacant d -orbital Cl atom have vacant d -orbital to accommodate electron pair Vacant orbital Back bonding 2p𝛑-3d𝛑 - BF3 Species Bond Length (pm) BF4− 130.7 BF3 139.6 B-F bond in BF3 is found to be shorter and stronger than expected due to back bonding. Order of Accepting Tendency Lone pairs are present already Si > P > S > Cl When lone pairs are absent Si < P < S < Cl Factors for Coordinate 𝛑-Bonding Stability of 𝛑-bond Stability ∝ 1 Size of orbitals Size of orbital increases Weak 𝜋 bond Size of orbital decreases Strong 𝜋 bond Relative Stability of 𝛑-Bond BF3(2p𝜋-2p𝜋) > > 2p𝜋-3p𝜋 > PF3(3d𝜋-2p𝜋) 2p𝜋-3d𝜋 > 3p𝜋-3p𝜋 No. of lone pairs on the atom having vacant orbital ↑ Extent of back bonding ↑ 2p𝜋-2p𝜋 Backbonding Lewis acidic strength decreases. Lewis basic strength decreases. Bond length decreases. Bond angle may or may not change. Hybridisation may or may not change. Point to Remember!! One atom involved in back bonding must be of 2nd period Electron Deficient Compounds Insufficient no. of electrons to complete octet. Examples: BH3, BeCl2, BF3 Case 1: Electron Deficiency in BH3 ❏ Those bonds which has insufficient number of electrons and makes them stable are known as electron deficient bonds. ❏ BH3 is electron deficient compound. That’s why it undergoes dimerisation by means of 3c-2e─ bonds also known as banana bonds. ❏ Diborane (B2H6) is a dimer of BH3. ❏ In B2H6 there are two 3c-2e─ bond which are known as banana bonds. In B2H6 there are four 2c-2e─ bond which are known as terminal bonds. ❏ The hybridization of boron in B2H6 is sp3. Electron Deficient Bonds Those bonds which has insufficient number of electrons and makes them stable are known as electron deficient bonds. BH3 is electron deficient compound. That’s why it undergoes dimerisation by means of 3c-2e─ bonds also known as banana bonds. Case 2: Electron Deficiency in AlCl3 Cl Cl Al Cl Incomplete octet of Al Tendency to form dimer In Vapour Phase - Dimer of AlCl3 (Al2Cl6 ) sp3 - sp3 - sp3 overlap Two 3c - 4e─ bonds Case 3: Electron Deficiency in BeCl2 Incomplete octet of Be Forms dimer & polymer to get stabilized In Vapour Phase - Dimer of BeCl2 (Be2Cl4) sp2 - sp3 - sp2 overlap Cl Planar Cl Be Be Cl 3c - 4e─ bond Cl In Solid Phase Polymer of BeCl2 (BeCl2)n sp3 - sp3 - sp3 overlap Non planar 3c - 4e─ bond Case 4: Electron Deficiency in BeH2 H Incomplete octet of Be Be H Forms dimer & polymer to get stabilized In Vapour Phase - Dimer of BeH2(Be2H4) sp2- s - sp2 overlap 3c - 2ebond Planar In Solid Phase - Polymer of BeH2 (BeH2)n Non planar sp3 sp3 - s - sp3 overlap 3c - 2e─ bond sp3 Case 5: Electron Deficiency in ICl3 Forms dimer to minimize repulsion between lone pairs In Solid Phase - Dimer of ICl3 (I2Cl6) sp3d2 - sp3 - sp3d2 overlap Planar 3c - 4e─ bond Bond Parameters Bond Parameters Bond angle Bond length Bond energy Bond order Bond Angle Angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion 180° + - - Expressed in degrees & is spectroscopically determined + Gives ideas about distribution of orbitals around the central atom which helps in determination of shape. Factors Affecting Bond Angle (1) Hybridization 120° sp3 sp2 As % s character 109.5° Bond angle 180° sp Factors Affecting Bond Angle Steric Repulsions (2) Steric Repulsions Bond angle Same central atom (2nd period), same hybridisation (sp3) & side atoms are of 3rd period & onwards O F 103° O F Cl O Cl 110.13° B r B 112° r Factors Affecting Bond Angle (3) Number of lone pairs on the central atom As number of lone pairs Bond angle Same hybridisation of the central atom Factors Affecting Bond Angle Electronegativity of the central atom (4) E.N. of central atom Bond angle Same hybridization and number of lone pairs on central atom : : : : > : : : : : : : : : : : : : : : : > : : : : Factors Affecting Bond Angle (5) Electronegativity of the side atoms E.N. of side atom P F F 97.8° F Cl P Cl 100.3° Cl Bond angle Br Same central atom, same hybridization & same number of lone pairs P 101° Br Br I P 102° I I Point to Remember!! Regular geometry All the side atoms are identical and no lone pair on central atom Bond angle not affected by electronegativity Example 109.5o 109.5o 109.5o CCl4 = SiCl4 = 109.5o GeCl4 = SnCl4 Factors Affecting Bond Angle (6) Side atom Back Bonding Central atom Central atom Side atom Factors Affecting Bond Angle BF3 Due to back bonding Bond order But net effect in repulsion is zero Bond angle = 120o N(SiH3)3 Due to back bonding Hybridization changes from sp3 to sp2 Bond angle Bond Length Equilibrium distance between the nuclei of two bonded atoms in a molecule Factors Affecting Bond Length Size of the bonded atom Multiplicity of Bonds Number of lone pairs on bonded atoms % s-character Electronegativity difference Bond Energy Amount of energy required to break 1 mole of particular type of bonds between two atoms in gaseous state. Unit : kJ mol-1 Multiplicity of bond Bond Magnitude of Bond energy Energy (kJ mol-1) C C 347 C C 611 C C 837 Bond Energy Bond length (for same bonded atom) Bond Bond energy Bond length (pm) Energy (kJ mol-1) Cl Cl 199 243 Br Br 228 192 I I 267 151 Cl Cl > In group 15, 16 and 17 single bonds between 2nd period elements are exceptionally weaker due to l.p - l.p. repulsions Example : Br Br > F F > I I Bond Order The number of bonds between the two atoms in a molecule. Shared pairs of electrons Lone pair of electrons + Bond order = 2 What We Know? Types of bonds Covalent bond Sharing of electrons Ionic bond Transfer of electrons What Exists in Reality? There exists some covalent character in an ionic bond and some ionic character in a covalent bond! Covalent Bond Non-polar Covalent When 𝚫E.N. < 0.4 between the bonded atoms Covalent Bond < Polar Covalent When 𝚫E.N. 0.4 between the bonded atoms Non-polar & Polar Covalent Bond A A A B Symmetrical electron cloud Asymmetrical electron cloud H2, Cl2, N2, F2... HF, HCl, HBr, HI... E.N. of A < B Polarisation Due to polarisation Charged ends develop 𝛅+ H ─ 𝛅 Cl Act as an electric dipole Dipole Moment 1. 2. 3. 4. Dipole moment is a measure of the separation of charges(polarity) between the two ends of a dipole. It's magnitude is equal to the product of charge and the distance of separation. It a vector quantity. It is denoted by μ. Dipole Moment Formula of dipole moment is given as: 𝛍 = q ×d Unit = Debye d = Distance of separation (Å) q = Magnitude of charge (e.s.u.) ─q 𝛍 d +q Direction and Representation Represented by a small arrow with tail on the positive centre and head pointing towards the negative centre. Cl H ─ 𝛅 𝛍 d 𝛅+ Value of Dipole Moment (𝛍) 𝛍 Diatomic Molecule Difference in Electronegativities Polyatomic Molecule Bond dipole Spatial arrangement Dipole Moment (𝛍) In diatomic molecules, 𝛍 depends upon Difference in electronegativities & bond length In polyatomic molecules, 𝛍 depends upon Bond dipole and spatial arrangement Resultant Dipole Moment (R) 𝛍=P 𝛉 𝛍=Q R = (P2 + Q2 + 2PQ cos𝛉) Dipole Moment Generally, out of ‘q’ and ‘d’, ‘q’ is the dominant factor. q depends on 𝛥.E.N. q 𝛥.E.N. For a non-polar molecule, 𝛍net = 0 μ For a polar molecule, 𝛍net ≠ 0 Dipole Moment Diatomic molecules Homodiatomic (Non polar) Heterodiatomic (Polar) 𝛍net = 0 𝛍net ≠ 0 Polyatomic molecules Bond dipole Spatial arrangement Dipole Moment It can be zero as the two oppositely acting bond dipoles can cancel each other Net dipole moment( 𝛍net ) = 0 Regular Geometries B B B B A B B A B B B 𝛍net = 0 𝛍net = 0 Dipole Moment Lone pair contributes in dipole moment, but its contribution can’t be quantified as size of lone pair is not known. Dipole Moment = 0 Symmetrical Molecule Dipole Moment ≠ 0 Asymmetrical Molecule Predicting Geometry Using Dipole Moment CI H H C C C CI Cl CI H H H Cl Cl Cl CCl4 CH4 CHCl3 𝛍=0 𝛍=0 𝛍 = 1.04 D Some Important Order of Dipole Moment CD3F > CH3F F F C C D D D Δ E.N. in C - D H H H > Δ E.N. in C - H Dipole Moment of Dichlorobenzene Orthodichlorobenzene > Metadichlorobenzene > Paradichlorobenzene What are Cis and Trans? Cis: Similar groups on same side 𝛍net ≠ 0 Trans: Similar groups on opposite sides 𝛍net = 0 Effect of Dipole Moment on Boiling Point Generally, Dipole moment Boiling point B.P. of cis is greater than trans Dipole Moment and Percentage Ionic Character % Ionic character = 𝛍Observed 𝛍Theoretical × 100 𝛍Observed Experimental value of 𝛍 𝛍Theoretical Assuming 100% ionic compound Covalent Character in Ionic Compounds When an anion and a cation approach each other Valence shell of the anion is pulled towards the nucleus of the cation Polarisation Phenomenon of deformation of an anion by a cation The shape of the anion is deformed Polarising power of the cation The ability of a cation to polarise a nearby anion Polarisability of the anion Ability of an anion to get polarised Fajan’s Rule Greater is the polarisation of an anion in a molecule, more is the covalent character in the molecule. Charge on Cation 1 Size of Cation Polarization ∝ Charge on Anion Size on Anion Covalent character Example: BeCl2 > MgCl2 > CaCl2 > SrCl2 > BaCl2 (As size of cation increases from left to right, Polarisation decreases) Factors affecting polarisation For the cations of nearly the same size and charge, Order of polarizing power: Pseudo inert gas Inert gas > configuration configuration Eg: CuCl > NaCl (Covalent character) Cu+ [Ne] 3s2 3p6 3d10 Pseudo inert gas configuration Na+ 1s2 2s2 2p6 Inert gas configuration Fajan’s Rule Cations with pseudo inert gas configuration: (n-1)d10 ns0 More Zeff due to poor shielding effect of d and f electrons. 02 Polarising power increases 01 Polarisability of Anion Oxyanions are generally less polarisable because charge is present on O atom which is very small and we need to consider only the element which acquires the charge (and not the complete anion) Polarisability ∝ Charge on the anion S2─ Cl─ < > ClO2─ SO32─ < ClO3─ > SO42─ < ClO4─ Applications of Fajans’ Rule Determination of Covalent Character in an Ionic Compound Extent of polarisation LiF Covalent character < LiCl < LiBr < NaF < Na2O < Na3N LiI Variation in Melting Point Melting point of ionic solids depends upon 01 Lattice energy 02 Extent of polarisation Fluorides (F–) of s-block metals (except BeF2) & Al3+ and Cl–, Br–, I– of alkali metals (except Li+) are dominantly ionic. Melting point ∝ Lattice Energy Variation in Melting Point For Cl–, Br–, I– of Li+, all alkaline earth metals & Al3+, extent of polarisation is high. Melting point ∝ 1 Extent of polarisation As covalent character in an ionic compound, the melting point decreases. M.P. of covalent compound < M.P. of ionic compound Variation in Melting Point LiF > LiCl > LiBr > LiI NaCl > MgCl2 > AlCl3 > SiCl4 Intensity of Colour Colour of some compounds can be explained by Bigger anions are more polarised & hence their electrons get excited by Polarisation of their bigger negative ions. Partial absorption of visible light. PbCl2 SnCl2 PbBr2 SnI2 PbI2 Solubility in Water Extent of polarisation is high Solubility of p-block/d-block salts & halides of Be is low Solubility in Water ∝ Covalent character AgCl 1 Extent of polarisation Solubility in water > AgBr > Fe(OH)2 > Fe(OH)3 AgI Solubility in Water Ag2O > Ag2S Thermal Stability of Ionic Compounds ❖ For uniatomic anion, as interionic distance increases, lattice energy decreases, hence thermal stability decreases. ❖ Be2N2 > MgN2 > CaN2 > Sr2N2> Be3N2 ❖ For multiatomic anion (for compounds having the same anion) thermal stability increases down the group. Molecular Orbital Theory Features of MOT 01 Electrons in a molecule are present in the molecular orbitals (MO’s) Atomic orbitals (AO’s) of comparable energies & proper symmetry combine to form MO’s 03 02 AO is monocentric whereas a MO is polycentric Number of MO’s formed is equal to the number of combining AO’s 04 Features of MOT 05 There are two types of molecular orbitals: Bonding Molecular Orbitals (BMO) and antibonding Molecular Orbitals (ABMO). BMO has lower energy and hence greater stability than the corresponding ABMO 07 06 Electron probability distribution around a group of nuclei in a molecule is given by a MO MO’s are filled according to Aufbau principle, Pauli’s exclusion principle & Hund’s rule 08 Linear Combination of Atomic Orbitals (LCAO) BMO (Bonding Molecular Orbitals) ABMO (Anti Bonding Molecular Orbitals) Constructive interference Destructive interference ΨBMO = ΨA + ΨB ΨABMO Where, A and B are atoms = ΨA ─ ΨB Molecular Orbitals Bonding Molecular Orbital AntiBonding Molecular Orbital Electron Density in BMO Ψ Electron density increases in the internuclear region. 2 H H BMO Electron Density in ABMO Electron density decreases in the internuclear region. Ψ2 H H ABMO Linear Combination of Atomic Orbitals 01 Bonding Molecular Orbital (BMO) Types of MO’s 02 Antibonding Molecular Orbital (ABMO) Difference between BMO and ABMO Bonding Molecular Orbital (BMO) MO formed by the addition of Atomic orbitals Antibonding Molecular Orbital (ABMO) MO formed by the subtraction of atomic orbitals 𝛹BMO = 𝛹A + 𝛹B 𝛹ABMO = 𝛹A - 𝛹B Formed by constructive interference (Stabilized MO) Formed by destructive interference (Destabilized MO) Bonding Molecular Orbital (BMO) Antibonding Molecular Orbital (ABMO) Lower in energy as compared to atomic orbital Higher in energy as compared to atomic orbital Electron density increases in the internuclear region Electron density decreases in the internuclear region May or may not have a nodal plane Always has a nodal plane Represented by 𝞂1s, 𝞂2pz, 𝛑2px, 𝛑2py Represented by 𝞂*1s, 𝞂*2pz, 𝛑*2px, 𝛑*2py 𝞂 and 𝛑 Molecular Orbitals 𝞂 Head on Symmetrical around the bond axis 𝛑 Sideways Asymmetrical around the bond axis Shapes of MOs Formed by s-orbitals When two orbitals combine in same phase then constructive interference take place. When two orbitals combine out of the phase then destructive interference take place. (Constructive Interference) + s s (Destructive Interference) s σs s σ*s Molecular Orbital Energy Diagram Energy H2 molecule σ*1s 1s Atomic Orbital of hydrogen 1s σ1s Molecular Orbitals of H2 Atomic Orbital of hydrogen Shapes of MO’s: 𝞂2pz _ + + 2pz _ In Phase _ + _ 𝞂2pz + 2pz _ + 2pz + 2pz _ + Out of Phase _ + _ 𝞂*2pz + Shapes of MO’s Shapes of MO’s In Phase + _ 2px 2px 𝛑2px Out of Phase + 2py + 2py 𝛑*2py Gerade & Ungerade Molecular Orbitals On moving equal distance in the opposite direction from the centre of the MO, if the sign of 𝛹 is the same Gerade is not the same Ungerade Gerade orbital (𝛹g) 𝛔, 𝛑* Ungerade orbital (𝛹u) 𝛔*, 𝛑 Molecular Orbital Diagram for > 14 Electron System For molecules having > 14 electrons 𝛔*2pz 𝛑*2px 𝛑*2py 2p 2p Energy 𝛑2px 𝛑2py < (𝛔2pz) < [𝛑2px = 𝛑2py] < 𝛔2pz [𝛑*2px = 𝛑*2py] < (𝛔*2pz) 𝛔*2s 2s 2s 𝛔2s 1s 𝛔*1s 𝛔1s (𝛔1s) < (𝛔*1s) < (𝛔2s) < (𝛔*2s) 1s Molecular Orbital Diagram for ≤ 14 Electron System For molecules having ≤ 14 electrons 𝛔*2pz 𝛑*2px 𝛑*2py 2p 2p Energy 𝛔2pz < [𝛑2px = 𝛑2py] < (𝛔2pz) 𝛑2px 𝛑2py 2s 1s 𝛔*2s 𝛔2s 𝛔*1s 𝛔1s (𝛔1s) < (𝛔*1s) < (𝛔2s) Strength of H-bonding O > Decreasing tendency to donate lone pair F Point to Remember!! To compare strength of H-bond First check Δ E.N. and then tendency to donate lone pair Symmetrical Hydrogen Bonding Very strong H-bonding occurs in the alkali metal hydrogen fluorides of formula M[HF2] x y Bond lengths: x = y = 113 pm Bond energy of both H-F = 163 kJ/mol Types of Hydrogen Bonding Hydrogen Bonding Intermolecular Intramolecular Between two or more molecules Within a molecule Examples of Intermolecular Hydrogen Bonding 𝛅+ 𝛅- 𝛅+ Acetic Acid 𝛅- Intermolecular H - Bonding Homo Intermolecular Hetero Intermolecular 𝛅+ 𝛅 2 𝛅+ 𝛅- 𝛅+ 𝛅- Water Alcohol in Water Conditions for the Formation of Intramolecular Hydrogen Bond (1) Ring formed as a result of H bonding should be planar (2) 5 or 6 membered ring should be formed (3) Minimum strain should be there during ring closure Intramolecular H-Bonding Examples 𝛅+ 𝛅- o-Nitrophenol Point to Remember!! Cl usually doesn’t form H - bond due to their low charge density 𝛅+ 𝛅- 𝛅𝛅+ Chloral hydrate (CCl3CH(OH)2) H-Bonding Dependency on Physical State of Compounds Extent of H–bonding depends on the physical state of the compound. Gaseous state < Liquid state < Solid state Solubility Viscosity Effect of H-Bonding on Physical Properties Boiling Point Physical State Acidic and Basic Strength 01 Few organic compounds (Non-polar) are soluble in water (Polar solvent) due to H-bonding. Example: Alcohol in water. 02 Solubility 03 C2H2 is highly soluble in acetone due to H-bonding but not in water. Intramolecular hydrogen bonding leads to chelate formation, so the solubility of that species involved in intramolecular H-bonding in water decreases. Order of Boiling Point SbH3 > NH3 > AsH3 > PH3 H2O > H2Te > H2Se > H2S HF > Hl > HBr > HCl Why does Ice Floats over Water? Extensive network of H bonds Ice has cage like structure with vacant space H2O (s) is less dense than H2O (l) Did You Know? D2O (s) sinks in H2O (l) E.N. of D is less than H D forms stronger H-bond Density: D2O (s) > H2O (l) Clathrates Species formed by entrapment of appropriately sized gas molecules (e.g.: Xe, Kr etc.) into the voids of ice Ar.6H2O Noble Gas Clathrates Kr.6H2O Xe.6H2O Clathrates 𝛅+ 𝛅+ 𝛅+ 𝛅+ 𝛅- + 𝛅 𝛅- 𝛅+ 𝛅- 𝛅- 𝛅+ 𝛅+ 𝛅- 𝛅+ 𝛅𝛅+ 𝛅+ 𝛅+ Ar.6H2O