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Chemical Bonding

1LECTURE

2LECTURE

3LECTURE

4LECTURE

5LECTURE

6LECTURE

7LECTURE

8LECTURE

9LECTURE

10LECTURE

11LECTURE

12LECTURE

13LECTURE

14LECTURE

Chemical Bonding 2

Chemical Bonding

Chemical Bonding 3

Chemical Bond is force of attraction between elements.

They are formed so that elements/atoms become stable.

It is explained by following theories:

I. Kossel-Lewis Concept

II. VSEPR Theory

III. Valence Bond Theory

IV. Molecular Orbital Theory

Chemical Bonding 4

I. Kossel-Lewis Concept

According to them, atoms form bonds to become stable i.e., to attain nearest noble gas configuration i.e. Octet (and Duplet in some cases.)

For this, atoms of elements need to lose or gain or share e-.

Therefore they form bonds which are of 3 types:

(b) Covalent bond

(a) Ionic or Electrovalent bond

(c) Dative/Co-ordinate bond

Chemical Bonding 5

(a) Covalent Bond

Bond formed by sharing of e–.

It is of further 3 types

¨ Single Covalent Bond

¨ Double Covalent Bond

¨ Triple Covalent Bond

Chemical Bonding 6

¨ Single Covalent Bond : Sharing of 1 pair/2e–s

Formation of Hydrogen moleculeH + H H H or H—

HFormation of H2O molecule

+ orO H H—O—HH O + H H

Formation of NH3 molecule

+ orN H H—N—HH N + H H

H H H

Formation of Methane (CH4) molecule

+ orC H H—C—HH C + H H

+ H HH

+H

H H

Chemical Bonding 7

¨ Double Covalent Bond : Sharing of 4e– /2 pairs

Formation of Oxygen molecule

O + orO O O O=O

Formation of Ethylene (C2H4) molecule

+

or C = C

H

C

+ HH H

+H

C

+ HC C

H H

H

H

H

H

Chemical Bonding 8

¨ Triple Covalent Bond : Sharing of 6e– /3 pairs

Formation of Nitrogen molecule

N + orN N N N º N

Formation of Ethyne

or H—CºC—H+H C + C + H C C HH

Chemical Bonding 9

(b) Dative Bond:

Bond formed by sharing of 2e– which belongs to one of the atoms.

O = S ® OO S O S OO

Formation of sulphur dioxide molecule

H N + H+

H

HN H+H

H

H

Formation of ammonium (NH4)+ ion

H—N—>H

H

H

+

H—O—S—O—H

O

O

Chemical Bonding 10

Exceptions of Octet Theory:

1. Molecules with less than 8e– in outermost shell like BeCl2, BF3, AlCl3 etc.

Be

ClCl

(Be has 4 electrons around it)

B FF

(B and Al have 6 electrons around them)

, Al ClCl

F Cl

2. Expanded Octet

P

Cl

Cl Cl

Cl Cl

(10e– around P)

S

FF F

(12e– around S)

F FF

I

FF F

(14e– around I)

F FF

F

H—O—S—O—H

O

O

(12e– around S)

3. Odd number of e–

O—N = O

N º O

Chemical Bonding 11

Drawbacks of Octet Theory

¨ According to theory atoms form molecules to attain stability like noble gases but then why noble gases like Xenon forms molecules.

e.g.XeO3, XeOF4, XeF2

¨ No explanation about stablity of molecules.

¨ No explanation about shapes of molecules.

Chemical Bonding 12

Formal Charge

Charge possessed by an atom in a molecule/radical.

F.C. = No. of v. e– – Number of unpaired e– or non bonded–1/2 Number of shared e–

Example:

O = O — O

F.C. of O–1 Þ 6 – (4 × ½) – 4 = 0

F.C. of O–2 Þ 6 – (6 × ½) – 2 = + 1

F.C. of O–3 Þ 6 – (2 × ½) – 6 = – 1

Chemical Bonding 13

NH4+

F.C. on H = 1 – (2 × ½) – 0 = 0

F.C. on N =

5 – (8 × ½) – 0 = + 1

H — N — H

H

H

(CO)

C º O

F.C. of C = 4 – (6 × ½) – 2 = – 1

F.C. of O = 6 – (6 × ½) – 2 = + 1

Chemical Bonding 14

(NO3)–

O = N — O

O

–

F.C. of O–1 Þ 6 – (4 x ½) – 4 = 0

F.C. of O–2 Þ 6 – (2 x ½) – 6 = – 1

F.C. of O–3 Þ 6 – (2 x ½) – 6 = – 1

F.C. of N Þ 5 – (8 x ½) – 0 = + 1

Chemical Bonding 15

Resonance

When a single structure cannot explain all properties, then two or more structures are drawn (canonical/resonating structure) for it.But the single structure which can explain all its properties is called resonating hybrid or resonance hybrid.

Out of all structures Resonating hybrid is most stable.

For example:

Ozone (O3) molecule can be represented by structures 1 and 2.

The O—O is 140 pm and O = O is 121 pm, but the bond length which exist in O3 molecule, are both 128 pm i.e. it shows partial double bond character which is much better and stable than other 2 structures.

O

O OI

O

O OII

O

O O

128

pm

128

pm

Chemical Bonding 16

Partial Double Bond character is actually delocalisation of e–.

O = C = O I

OºC— OII

+ – O—C º OIII

+–

C

O O

O

––

I

C

O O

O

–

II

–

C

O O

O

–

III

–

I II

Kekule structures

Chemical Bonding 17

S

O O

O

S

O O

O

S

O O

O

N

O O

O–

N

O O

O

–

N

O O

O

–

Chemical Bonding 18

¨ Out of all canonical forms it is the resonance hybrid which exist in nature.

¨ The total number of Lone Pairs (L.P.) and Bond Pairs (B.P.) remains the same.

¨ Total charge on molecules or ion remains same.

¨ None of the canonical forms exist.

Chemical Bonding 19

(c) Ionic Bond

Bond formed by complete transfer of e– is called ionic bond.

It means it is formed between Cations and Anions i.e., between metal and non-metal.

.. ..

.. ..(2, 8,1) (2, 8)(2, 8, 7) (2, 8, 8)

.. : ::Na Cl Na Cl

2.. ..2

.. ..(2, 8)(2, 8, 2) (2, 6) (2, 8)

: :Mg: O Mg :O or MgO

2 2(2, 8, 2) (2, 8)

Mg: Mg or MgBr

. .

. .::Br

. .

. .. :Br

. .

. .

(2, 8,18, 8)

::Br

. .

. .

(2, 8,18, 7)

. :Br

Chemical Bonding 20

Factors Governing the Formation of Ionic Bonds

(i) Low Ionisation Enthalpy.

(ii) High Electron Gain Enthalpy.

(iii) High Lattice Enthalpy.

General Characteristics of Ionic Compounds

(i) All are solid in nature.

(ii) Crystal Structure ions are arranged in a regular pattern in the three dimensional space to form a lattice.

(iii) High melting and boiling points.

(iv) Solubility: Soluble in water (as ions combine with the solvent to liberate energy called the hydration enthalpy which is sufficient to overcome the attractive forces between the ions).(v) Ionic compounds are good conductors of electricity in solution or in molten state, as their ions are free to move.

Chemical Bonding 21

Bond Parameters

(i) Bond LengthIt is the distance between the centre of 2 nuclei of atoms closely bonded.

From bond length we can get atomic radii of respective atoms.Different types of atomic radii are:

(a) Covalent radii

Cl—Cl

Bond Length = 198pm

Covalent radii of Cl-atom = 198/2 = 99 pm

(b) Metallic radiiIn metallic CuCu—Cu = 256pmM.R. of Cu-atom = 256/2

= 128 pm

(c) Vander Waal radius

V.W. of radii of H2 = 360/2

= 180 pm

Chemical Bonding 22

(ii) Bond Angle

It is the angle between two bonds.

It is formed due to repulsion between bonded and non bonded e– to get max. stability.

H

C

H

109°28¢

H

H

H

N

H107

°H H

O

104.5°

H F

B

F

F

120°

(iii) Bond EnthalpyEnergy required to break all the bonds present in 1 mole of a molecule. OR

Energy released when all the bonds are formed in 1 mole molecule.

Bond Breaking ® + DH (Endo)

Bond Formation ® - DH (Exo)

Chemical Bonding 23

DH = + 496 kJ / molO2 (O = O) + BE ® 2O1 mol atom

DH = + 946 kJ / molN2 (N º N) + BE ® 2N1 mol atom

DH = + 431 kJ / molH—Cl + BE ® H + Cl1 mol

BE (DH) = 435 kJ/molH2 + Heat ® 2H(H—H)1 mol

(Stable)(Unstable)

E

Chemical Bonding 24

(iv) Average Bond Enthalpy

In a molecule when same bond is 2 or more in number then each bond energy is different, in fact successive B.E. is less, hence average bond enthalpy is preferred.

H2O

H—O—H + BE1 ® H—O + H DH1 = 502 kJ / mol

H—O + BE2 ® H + O DH2 = 427 kJ / mol

Average Bond Enthalpy =

DH1 + DH2

2

=502 + 4272

= 464.5 kJ/mol

Chemical Bonding 25

(v) Bond Order

Number of bonds present between 2 atoms in a molecule.

H2 (H—H) Þ B.O. = 1

O2 (O = O) Þ B.O. = 2

N2 (N º N) Þ B.O. = 3

Iso-electronic species have same B.O.

O22– and F2 Þ B.O. = 1

N2, CO and NO+ Þ B.O. = 3

B.O. , B.S. , B.L. ¯, Stability .

(vi) Lattice EnthalpyEnergy required to separate one mole of ionic solid into its corresponding gaseous ions is called Lattice Enthalpy. S.I. unit is kJ/mole.

Na+Cl–(s) + L.E. ® Na+(g) + Cl–(g) L.E. Þ + 788 kJ/mol

Na+(g) + Cl–(g) ® NaCl(s) + 788 kJ/mol

Chemical Bonding 26

Polarity of Bond

Due to high electronegativity of atom in a bond, the shared pair of e– is pulled towards itself and attains d– (small/partial) charge whereas other atom attains d+, therefore it results in formation of 2 poles within a molecule hence is called dipole.

Dipole MomentIt is the existence of partial charges d– and d+ within a molecule.Mathematically it is defined as the product of magnitude of charge on atom and distance between them.

Cl—Cl

Cl : Cl

H—Cl

Hd+ : Cld–

Hd+—Cld–

Dipole moment (µ) = q × r

S.I. unit of µ is Debye (D) or Cm

1D = 3.33 × 10–30 Cm

Chemical Bonding 27

1. Polar Bond/Molecule: Molecules having value of dipole moment. i.e., µR ¹ 0

On the basis of Dipole Moment, covalent bonds are two types:

O–2d

H+d H+d

N–3d

H+d

H+d

H+d

N

F

Resultant of 3N—F bonds

F

F

µ = 0.24 D

Chemical Bonding 28

2. Non-Polar bonds/molecule: Molecules having zero dipole moment. i.e., µR = 0.

–d

+2dO C

–dO

–d

+2d

–d

–d

+2d

–d

S C S F Be F

F–d B+3d

F–d

F–d

C–4d

H+

dH+d

H+d

H+

d

C+4d

Cl–d

Cl–d

Cl–d

Cl–d

Chemical Bonding 29

Fajan’s Rule

If covalent molecules show partial ionic character then ionic compounds also possess some covalent character, which depends on following factors:

¨ Greater charge on cation more is covalent character.

¨ Smaller cation or bigger anion shows greater covalent character.

¨ d-elements forms more of covalent character compounds.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Theory which gives and explains shapes of molecules.

¨ Electrons of valence shell of central atom participates to form bond.

Main Postulates

¨ Electrons of surrounding atoms form bond with e– of central atom.

Chemical Bonding 30

¨ The central atom is linked to similar or different atoms to get geometry.

¨ Bonded e– and lone pair e– repel each other to attain a shape having minimum repulsion and maximum stability.

Repulsion : lp– lp > lp – bp > bp – bp

[lp: Ione pair and bp : bond pair]

(a) Shape of Molecules Without Lone Pair:

B BA

Linear

or O = C = O

(BeF2, BeCl2, HgCl2)

Chemical Bonding 31

or

B B

A

Triangular planar

B

B

(AlCl3, SO3)

F

F F

or

B B

A

Tetrahedral

B

B

C

(SiF4, NH4+)

H

HHH

Chemical Bonding 32

or

B

BA

Trigonal bipyramidal

B

B

B

PCl5

(PF5, SbCl5, AsF5)

or

B

A

Octahedral

B

BB

B

B

SF6

Chemical Bonding 33

or

(SO2, O3, SnCl2)

N

O O

or

B B

A

Bent (V-shape)

B B

A

Pyramidal

B

N

(PCl3, NCl3, NF3, PH3, H3O+)

H HH

(b) Shape of Molecules With Lone Pair:

Chemical Bonding 34

orB

A

Bent

B

O

(H2S, NH2–)

H H

orA

See saw

B

B

B

B

S

(SF4)

F

F

F

F

Chemical Bonding 35

orBA

T-shaped

B

B

FCl

(ClF3)

F

F

orA

Linear

B

B

XeF2

(I3–, Br3–, ICl2–)

Chemical Bonding 36

or ClF5

(IF5, BrF5)

or XeF4

B

A

Square pyramidal

B

B

B

B

B

A

Square planar

B

B

B

Chemical Bonding 37

Valence Bond Theory (V.B.T.)

According to VBT when atoms are brought close to each other they form stable molecule with a bond for e.g.

Case I : When two unstable H-atoms are far away then force of repulsion and force of attractions are negligible.

Case II : When they come closer, at some minimum distance (B.L.), force of attraction is high and force of repulsion is minimum.

Case III : On further decreasing distance between them, attraction is too high but repulsion also becomes so high which makes the molecule unstable.

P.E.

B.E

.B.L.

Internuclear Distance

Chemical Bonding 38

E = -ve (Released-Exothermic) E = +ve (Absorbed-Endothermic)

¨ e– present in valence shell of central and surrounding atoms participate.

Postulates/Salient Features of V.B.T.:

¨ Only unpaired e– of valence shell participates but not lone pair.¨ Unpaired valence e- overlap or forms bond with unpaired e– of central atom but with opp. spin.

¨ These bonds formed are of two types.

Chemical Bonding 39

¨ Extent of overlapping is more, therefore stronger bond.

¨ less hence weaker

¨ can exist independently ¨ only with s bond

¨ Head on overlapping ¨ Sideways

¨ Exist between s—s, s—p, p—p

¨ p—p

Sigma bond (s) Pi bond (p)

¨ Atomic orbitals overlap on inter nucleus axis (INA)

¨ Overlapping ^ to INA

Chemical Bonding 40

Drawbacks of Valence Bond Theory (V.B.T.)

¨ Acc. to VBT only unpaired e– are paired, therefore the simplest hydrocarbon which exist in nature should be CH2. But CH2 does not exist and is highly unstable and the hydrocarbon which exists is CH4.

¨ Similarly according to VBT instead of BF3 or BH3, only BF and BH should exist as it has one unpaired e- whereas Be should not form any compound as it doesn’t have any unpaired e-.

Chemical Bonding 41

Hybridisation:

¨ Atomic orbitals of similar energies can participate.

It is the mixing of atomic orbitals of similar energies to get new stable hybridised orbital of same energy and shape.

¨ Number of atomic orbitals(AO) combining gives same number of hybrid orbitals (HO).

Salient features:

¨ Paired, unpaired and vacant orbitals can participate.

¨ These HO’s form stronger and stable bonds.

¨ These HO’s repel to give proper geometry.

Chemical Bonding 42

¨ AO with similar energy will participate.

¨ Not necessary for excitation of e–.

Important Condition for Hybridisation

¨ Not necessary only unpaired e– participate.

¨ Electrons in last and second last shell also can participate.

Chemical Bonding 43

Types of Hybridisation:

(i) sp3 Hybridisation

When one ‘s’ and three ‘p’ (px, py, pz) A.O. of similar energies combine, they form four new stable H.O.’s each having same energy and shape.

These four HO form s bond and each having 25% ‘s’ character and 75% ‘p’ character.

These four HO’s repel each other to give tetrahedral geometry with 109.28° angle.

(i) sp3 (ii) sp2

(iii) sp (iv) sp3d

(v) sp3d2

Examples: Alkanes (CH4), NH3, NH4+, H2O, H3O+, etc.

Chemical Bonding 44

Formation of sp3 hybrid orbitals

Chemical Bonding 45

(ii) sp2 Hybridisation

One ‘s’ and two ‘p’ orbitals participate to give three new sp2 hybridised orbitals.

Each possess 33% ‘s’ character and 67% ‘p’ character. They repel each other to give trigonal planar geometry with 120° bond angle. Example: Alkenes (C2H4), BH3, BF3, etc

Chemical Bonding 46

(iii) sp Hybridisation

One ‘s’ and one ‘p’ orbital participate to form two new hybridised orbitals each having 50% ‘s’ character 50% ‘p’ character. Both hybridised orbitals repel each other to give linear shape with bond angle 180°. Since it has max. s-character hence has max. electronegativity.

Example:

Alkynes (C2H2), BeCl2, BeF2, BeH2 etc.

Chemical Bonding 47

(iv) sp3d Hybridisation

One ‘s’, three ‘p’ and one ‘d’ orbitals of same shell participates to give sp3d hybd. with trigonal bipyramidal shape with bond angles 90o &120o. Example: PF5 & PCl5,

P(G.S.) = 1s22s22p63s23p3

P(E.S.) = 1s22s22p63s1 3p3 3d1

sp3d

Chemical Bonding 48

So, now have 5 hybridised orbitals which form 5s bonds.

3P—Cl form trigonal planar shape with 120° bond angle and are called equitorial bonds which are short and strong.

2P—Cl are placed perpendicular which experience greater repulsion hence are long and weak called axial bonds.

So the overall shape is trigonal bipyramidal.

PCl5 is thermally unstable and gives PCl3 and Cl2 gas on heating because its two axial P—Cl bonds are weak.

PCl5 (s)D

PCl3 (s) + Cl2 (g)

Chemical Bonding 49

(v) sp3d2 Hybridisation

One ‘s’, three ‘p’ and two ‘d’ orbitals of same shell participate and give sp3d2 with octahedral shape having all bond angles 90o.Example: SF6

S(G.S.) = 1s22s22p63s23p4

S(E.S.) = 1s22s22p63s1 3p3 3d2

sp3d2

Chemical Bonding 50

Molecular Orbital Theory

2. When two A.O. of equal/comparable energies combine or overlap, they lose their identity and forms two new M.O.

3. The two new M.O.’s are Bonding M.O. and Antibonding M.O.

1. M.O. gives the electron probability distribution around a group of nuclei, just as an A.O. gives the electron probability distribution around the single nucleus.

Molecular Orbital (M.O.): It’s the space around the nuclei of molecule where the probability of finding electron is maximum.

Salient Features of M.O.T.

4. Bonding M.O. has less energy (more stable and denoted by , , ) whereas Antibonding M.O. has more energy (less stable and denoted by *, *,*).

5. Follows the same rules of filling electrons as that of A.O.6. A.O. with proper orientation can overlap/combine. For example, 1s

can combine with 1s and not with 2s. Similarly, s can combine with pz but not with px or py, or pz can not with px or py (taking Z-axis as the INA).

Chemical Bonding 51

Condition for Linear Combination of Atomic Orbitals (LCAO)

1. The combining A.O’s should have comparable energies. i.e., 1s combines with 1s, 2s with 2s and 2p with 2p.

2. The combining A.O’s should have proper orientation. i.e., same symmetry. e.g., 2pz can combine with 2pz or 2s but not with 2px or 2py.

3. The extent of overlapping should be max. to gives stronger bond.

Chemical Bonding 52

Case I.

When the two waves are in phase (constructive interference) so that they add up and the amplitude of the new wave is

Formation of Molecular Orbitals

f = yA + yB

yA

yB

Addition

yA + yB

Additive effect of the electron waves in phase

Chemical Bonding 53

Case II.

When the two waves are out of phase (destructive interference), the waves are substracted from each other so that the amplitude of the new wave is

¢f = yA – yB

yA

yB

Subtraction

yA – yB

Subtractive effect of the electron waves out of phase

Chemical Bonding 54

Molecular Orbital (MO) Theory

BONDING

ANTIBONDINGThese two new orbitals have different energies. 

Chemical Bonding 55

MO’s of 2p A.O’s

sigma

pi

Chemical Bonding 56

1. Electronic Configuration

Characteristic of Molecular Orbitals

s (1s) s* (1s) < s (2s) < s* (2s) < [p (2px) = p (2py)] < s (2pz) < [p* (2px) = p* (2py)] < s* (2pz)

2. Stability of molecules in terms of bond orderBond order is defined as half of the difference between the number of electrons present in the bonding and the antibonding orbitals, i.e.,

Bond order (B.O.) =

12

(Nb – Na).

3. Greater the bond order, shorter is the bond length.

4. Diamagnetic and paramagnetic nature of the molecules. If all the electrons in the molecule are paired, it is diamagnetic in nature. On the other hand, if the molecule has some unpaired electrons, it is paramagnetic in nature.

(B.M. = Bohr Magneton)

Magnetic moment =

n (n + 2)

B.M.

Chemical Bonding 57

2px 2py 2pz 2px 2py 2pz

s*(2pz)

p*(2px

)p*(2py

)s(2pz)

p(2px) p(2py)

s*(2s)

s(2s)

2s 2s

s*(1s)

s(1s)

1s 1sAtomic orbitals

Atomic orbitals

Molecular orbitals

Incre

asin

g E

nerg

y

Chemical Bonding 58

2px 2py 2pz 2px 2py 2pz

s*(2pz)

p*(2px

)p*(2py

)

s(2pz)

p(2px) p(2py)

s*(2s)

s(2s)

2s 2s

s*(1s)

s(1s)

1s 1s

Incre

asin

g E

nerg

y

Electronic Configuration (from O2 and above)

s(1s) s* (1s) < s (2s) < s* (2s) < s (2pz) < [p (2px) = p (2py)] < [p* (2px) = p* (2py)] < s* (2pz)

Chemical Bonding 59

Hydrogen Bonding:

When a hydrogen atom linked to a highly electronegative atom (like F, O or N) comes under the influence of another strongly electronegative atom, then a weak bond is developed between them.

This type of bond is called hydrogen bond and is represented by dotted lines as shown below:

¨ The greater is the electronegativity difference, more is the polarisartion of the bond of the molecule.

¨ Smaller the size of electronegative atom more is the electronegative attraction.

Chemical Bonding 60

Types of hydrogen bond:

(i) Intermolecular hydrogen bonding

(ii) Intramolecular hydrogen bonding

(i) Intermolecular hydrogen bonding:

When hydrogen bonding takes place between various molecules of the same or different compounds, it is called intermolecular hydrogen bonding e.g.

¨ They show abnormally high melting and boiling points.

(ii) Intramolecular H.B.:

When hydrogen bonding occurs within the same molecule is termed as intramolecular hydrogen bonding.

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