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MHT-CET : Chemistry Entrance Exam

MHT - CET : Chemistry - Nature of Chemical Bond Page 1

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Introduction
The stability of a molecule arises due to chemical bonding. The stability of atoms lies in their ability to form molecules. Only electrons in the outermost orbital take part in bond formation. The study of nature of bond involves - the study of nature of molecules, their stability, arrangement of atoms in the outermost shell, etc. All these factors play an important role in determining the nature of a chemical bond.

What is bond?
In a molecule, the atoms are held together by a strong force of attraction to form a bond. The force of attraction may be due to oppositely charged ions or due to orbital overlap.

Types of bond:
Three different types of bonds are formed depending on the electropositive or electronegative character of atoms involved.

  1. Electropositive element + electronegative element = Ionic Bond
  2. Electropositive element + Electropositive element = Metallic bond
  3. Electronegative element + electronegative element = Covalent bond.

Octet Rule
During bond formation atoms gain, loose or share electrons so that the outermost or valence shell of an atom has eight electrons as in inert gases.

Electronic Theory
Put forth by Kossel and Lewis in 1916. The main postulates are

  • Valence shell electrons take part in bond formation.
  • Inert gases have stable outermost configuration.
  • Elements tend to acquire inert gas configuration by gaining or losing electrons. On this basis, Ionic and Covalent bonds are explained.
    1. Ionic Bond: Bonds formed by gaining or losing electrons, in which the ions formed are held together by electrostatic force of attraction.
      Formation of Ionic Bond:
      11Na
      Na+ + e-
      1s2, 2s2, 2p6, 3s1
      1s2, 2s2, 2p6 + e-
                                      
                                          octet
      Cl + e
      - Cl-
      1s2, 2s2, 2p6, 3s2, 3p5 +e
      - 1s2, 2s2, 2p6, 3s2, 3p6
                                                                        
                                                                            octet
      Na+ + Cl
      - NaCl
    2. Covalent Bond: Bond formed between a pair of atoms by sharing one or more electrons in the outermost shell.

Formation of Covalent Bond:

xx

+

xx

 xx

x
x

xx

xx Cl x

x Cl xx

  xx Cl

Cl xx

xx

xx

 xx

xx

(2, 8, 7)

(2, 8, 7)

(2, 8, 7)

 

(2, 8, 7)

Limitations of Octet Rule

  • Fails to explain formation of compounds with incomplete and expanded octets.
  • Fails to explain about nature of forces responsible for the combination of atoms.
  • Does not explain energy, stability and reactivity of molecule.
  • Does not explain geometry and shape of different molecules.

Postulates of the Valence Bond Theory

  1. Covalent bond is formed by overlapping of atomic orbitals and hence energy of the system decreases.
  2. Atomic orbitals of two atoms having unpaired electrons overlap to form a covalent bond.
  3. Electrons in overlapping orbitals should have opposite spins and in the process spins are neutralised.
  4. Overlapping of orbitals causes increase in electron density in the region where overlapping occurs.
  5. Overlapping orbitals should have comparable energies.
  6. The bond formed has directional character and the strength of the bond is directly proportional to the extent of overlap.
  7. Number of unpaired electrons an atom possesses determines number of bonds formed, and hence its valency.

Formation of H2 Molecule

  • When two hydrogen atoms with opposite spin approach each other from infinity initial energy of the system is arbitrarily taken as zero.
  • When they come sufficiently close, following forces act:
    1. Attractive forces between nuclei and electrons of other atom.
    2. Repulsive forces between electrons and nucleus of the two hydrogen atoms.
  • Potential energy of the system goes on decreasing till some point.
  • At bond length,distance is minimum
  • At minimum energy, atomic orbitals overlap and the H-H bond is formed. This represents the lowest point in potential energy curve.

Structure of Helium

  1. Helium electronic configuration is 1s2 - hence there are no unpaired electrons in valence shell.
  2. 1s - orbitals repel - potential energy increases, hence no bond is formed. Helium is monoatomic in nature.


Polar Bond: Bonds between elements with electronegativity difference are polar bonds
example: HCl, HF, H2O
Non-polar Bond: Bond between similar atoms (no electronegativity difference).
example: H2, Cl2, Br2, etc.

Sigma Bond

Pi( p) bond

Linear overlap along inter-nuclear axis

Lateral overlap perpendicular
to inter-nuclear axis.

Maximum overlap occurs.

Extent of overlap is less.

Bond is rotationally symmetrical
along inter-nuclear axis.

Not rotationally symmetrical.

Stronger than p - bond.

Weaker than a sigma bond.

Overlap of Atomic Orbitals

  • s-s Overlap: Overlap between two half filled 's' orbitals. example: H2 molecule
  • s-p Overlap: Overlap between half-filled 's' orbital of one atom and 'p' orbital of another. example:H - F molecule
  • p-p Overlap: Overlap between half-filled 'p' orbitals of two atoms is called p - p overlap. example: F2 molecule.

Hybridisation
Definition:The
process of mixing and recasting to form same number of equivalent orbitals with maximum symmetry and definite orientation in space is called hybridisation.

Hybridisation involves the following steps:

Formation of
Excited State: Paired electrons jump to higher energy levels to create if necessary more number of half-filled orbitals.
Examples:

Ground state configuration of carbon

 

1s2       2s2         2p2

Excited state

 

1s2        2s1         2p3

sp3 - hybridisation

 

1s2       Four sp3 - hybrid orbitals

  • Mixing and Recasting of Atomic Orbitals: Orbitals of valence shell mix to form new set of atomic orbitals having same energy. The new orbitals then formed are called hybrid orbitals.
  • Orientation of Hybrid Orbitals in Space: Hybrid orbitals are then arranged symmetrically in available space.

Need for the Concept of Hybridisation

  • To explain valencies of element.
  • To explain equivalence of bonds.
  • To explain geometry of molecule.

Types of Hybridisation:

sp3-Hybridisation:
Mixing and recasting of 's' orbitals with three 'p' orbital of same atom forming four identical orbitals tetrahedrally arranged in space.

sp2-Hybridisation: One 's' and two 'p' orbitals of the same atom mix and form three identical orbitals trigonally arranged in space.

sp-Hybridisation: One 's' and one 'p' orbital of the same atom mix and form two identical orbitals diagonally arranged in space.

Formation of Methane (CH4) Molecule:-

  • Electronic configuration of carbon is 1s2, 2s2, 2p2.
  • Excited state configuration is 1s2, 2s1, 2p1x, 2p1y, 2p1z
  • 2s and 2p orbitals mix to form a set of four equivalent orbitals.
  • The hybrid orbitals are arranged tetrahedrally and the bond angle is 109, 28.

Diagrammatic Representation

Ground state configuration of carbon

 

1s2       2s2         2px',2py',2pzo

Excited state

 

1s2       2s1         2px',2py',2pz1

sp3 - hybridisation

 

1s2       sp3 - hybrid orbitals

Methane Molecule


Ammonia Molecule
This also undergoes sp3 - hybridisation.
Nitrogen has an electronic configuration of 1s2, 2s2, 2p3.
2s orbitals with three 2p orbitals form a set of four equivalent orbitals.
The lone pair of electron in nitrogen causes repulsion between two bonding pairs, hence bond angle is not 10928 but 10728.
Geometry is a trigonal pyramid.
The three N - H bonds in ammonia are sp3 - s - bonds.

Water Molecule
Oxygen has an electronic configuration 1s2, 2s2, 2p4 (no space for excitation)
2s and 3p orbitals mix and form 4 equivalent hybrid orbitals.
The four O - H bonds in water are sp3 - ss - bonds.
Two hybrid orbitals contain one lone pair of electron while two others contain one unpaired electron.
Geometry of water molecule is V - shaped or angular.

sp2 - Hybridisation
example: Formation of ethane:

Ground state electronic configuration is 1s2, 2s2, 2p2
Excited state electronic configuration is 1s2, 2s1, 2p1x, 2p1y, 2p1z
In sp2 - hybridisation, the 2s - orbital mixes with two of the three 2p - orbitals to form three equivalent hybrid orbitals.
The three C - H bonds in ethane are sp2 - s s - bonds.
Unhybridised 2p - orbital remains.
Bond angle is 120 and geometry is trigonal planar.

Diagrammatic Representation:

Ground state configuration of carbon

 

1s2       2s2      2px',2py',2pzo

Excited state

 

             2s1     2px',2py',2pz1

sp2 - hybridisation

         

 

1s2      

sp3 - hybrid orbitals   

Unhybridised

 

 

 

2pz - orbital

Ethane Molecule

            

 

s - bonds     p - bonds        



Boron Trifluoride Molecule:

Boron has ground state electronic configuration 1s2, 2s2, 2p1
2s and 2p orbitals mix to form three equivalent hybrid orbitals
The three B - F bond are sp2 - s, s - bonds in BF3
Boron has three B - F bonds and the resulting boron trifluoride molecule is with triangular planar geometry.
Boron trifluoride bond angle is 120.

sp-hybridisation:
Acetylene:

C2H2 (CH CH), carbon atoms are sp - hybridised.
Singly occupied sp - hybridised orbital overlap with singly occupied 1s - orbital of hydrogen atom to form C - H sigma bond.
Remaining sp - hybrid orbital overlap on other carbon atom to form C - C sigma bond.
Unhybridised half-filled py-orbital overlaps laterally with py of other carbon atom to form p bonds. Thus two p bonds are present.
Two p - bonds overlap to give a cylinder of electrons coaxial 'X' with bond axis.
Acetylene molecule is linear, C - C bond has one sigma and 2 - p bonds.
The two C - H bonds in acetylene are sp - s s - bonds.

Diagrammatic representation:

Ground state configuration of carbon

 

1S2       2S2     2px',2py',2pzo

Excited state

 

1S2       2S2 2px',2py',2pz1

Sp - hybridisation

 

1S2   

   sp- hybridised 

 2py1, 2pz1 Unhybridised

 

         

orbitals  

orbitals 

Acetylene molecule

            

 

s - bonds     p - bonds


sp-hybridisation in Beryllium Difluoride
Electronic configuration of Be is 1s2, 2s2

  • Excited state configuration is 1s2, 2s2, 2px1
  • In BeF2, beryllium atom is sp-hybridised.
  • Fluorine has the electronic configuration 1s1, 2s2, 2p5
  • Half-filled sp hybrid orbital of Be overlaps with 2p orbital of fluorine forming Be - F sigma bond.
  • Fe - Be - F bond angle is 180 and geometry is linear.
  • The two Be - F bonds in BeF2 are sp - s s- bonds.
  • Bond Energy: Average amount of energy per mole required to break a particular bond to produce free atoms or radicals.
  • Bond Dissociation Energy: Amount of energy required to break the bond in a molecule producing free atoms or molecules.

Factors Affecting Bond Energy

  • Bond Length:
    • Distance between nuclei of covalently bonded atoms.
    • Bond length is more, the bond energy is less and vice-versa.
  • Bond Polarity:
    • Two atoms with large electronegativity differences, bond is polar.
    • Acquires ionic character.
    • Bond polarity is more - bond energy increases.
  • Type of Bond orUnsaturation Character:
    • More number of bonds between 2 atoms in a molecule - more unsaturation.
    • Bond energy a unsaturation character of bond.
  • Percentage of 's' character:
    • Greater the percentage of s - orbital character in hybrid bond, stronger the bond.
    • Bond energy a percentage of 's' character.

 

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