Chemical Bonding

Front 1

Octet Rule

Back 1

Atoms of different elements take part in bond formation to attain noble gas configuration

Front 2

Ionic Bond

Back 2

Complete transfer of electrons from one atom to another

Front 3

Ionization Enthalpy in Ionic Bond

Back 3

Metallic atom/Cation element to have less I.E for easy loss of electron

Front 4

Electron Gain Enthalpy in Ionic Bond

Back 4

Non-metallic atom/ anion element to have high e.g.e for easy acceptance of electron

Front 5

Lattice Energy

Back 5

Amount of energy required to separate 1 mole of ionic compd into separate oppositely charged ions

Front 6

Covalent Bond

Back 6

Bond is formed by mutual combination and sharing of electrons

Front 7

Homoatomic Molecule

Back 7

Same atoms combine. Eg- H2

Front 8

Heteroatomic

Back 8

Different atoms make up the molecule. Eg- C2H4

Front 9

Limitations of Octet Rule

Back 9

1. Incomplete octet of central atoms in some covalent structures. Eg- BCl3

2. Some compds have odd no of total electrons. Eg- NO2

3. Expanded octet of 3rd period elements

4. Noble gases make compounds

Front 10

Bond Length

Back 10

Equilibrium distance between the centres of the nuclei of the two bonded atoms.

Front 11

Bond Angle

Back 11

Angle b/w lines representing the orbitals containing the bonding

Front 12

Bond Enthalpy

Back 12

Amount of energy required to break one mole of bonds of a particular type to separate them into gaseous atoms.

Front 13

Bond Order

Back 13

No of bonds b/w a covalent structure. Eg-

H2 has b.o=1 (H-H)

Front 14

Non-Polar Covalent Bond

Back 14

When the atoms joined by covalent bonds are the same like H2, Cl2, etc the shared pair of electrons is equally attracted and is equidistant from each.

Front 15

Polar Covalent Bond

Back 15

When two dissimilar atoms make a covalent bond like HCl the more electronegative atom will pull the electron pair towards it so both atoms develop partial charge.

Front 16

Dipole Moment

Back 16

Product of magnitude of charge of + and - charge and the difference between them.

µ = Q x d

Front 17

VSEPR

Back 17

Valence Shell Electron Pair Repulsion Theory

Front 18

Postulates of VSEPR

Back 18

1.shape depends on no of e- pairs

2.e- pairs repel each other

3.e- pairs try to be in a position of min repulsion

4.valence shell is taken as a sphere w the electron pairs placed at min distance.

5.multiple bond treated as a single e- pair

Front 19

VBT

Back 19

Valence Bond Theory

Front 20

Postulates of VBT

Back 20

1.bond formed from overlapping of orbitals

2.overlapping orbital contain a pair of e-

3.e- density concentrates b/w bonded atoms

4.strength of the bond depends on degree of overlap

Front 21

Sigma Bonds

Back 21

involves s orbitals, strong bond, axial overlap

Front 22

Pi Bond

Back 22

does not include s orbitals, comparatively weak, sidewise overlap

Front 23

Format of VSEPR

Back 23

1.central atom

2.valence electron of c.a

3.no of b.p and l.p

4.total e- pairs

5.geometry

6.shape

Front 24

BeCl2

Back 24

Linear, 180

sp

Front 25

BF3

Back 25

Trigonal Planar, 120

sp2

Front 26

CH4

Back 26

Tetrahedral, 109 28'

sp3

Front 27

PCl5

Back 27

Trigonal Bipyramidal ,90 and 120

sp3d

Front 28

SF6

Back 28

Octahedral, 90

sp3d2

Front 29

IF7

Back 29

Pentagonal Bipyramidal

sp3d3

Front 30

AB2 type molecule

Back 30

B.P-2

L.P-1

Geometry- trigonal planar

Shape-Bent

eg- SO2

Front 31

AB3 type molecule

Back 31

B.P- 3

L.P-1

Geometry- tetrahedral

Shape- pyramidal

eg- NH3

Front 32

AB4 type molecule

Back 32

B.P- 4

L.P-1

Geometry- trigonal bypyramidal

Shape- see-saw

eg- SF4

Front 33

AB2 with 2 L.P

Back 33

B.P-2

L.P-2

Geometry-tetrahedral

Shape- bent

eg-H2O

Front 34

AB5 type molecule

Back 34

B.P- 5

L.P-1

Geometry- octahedral

Shape- square pyramid

eg- BrF5

Front 35

AB3 with 2 L.P

Back 35

B.P- 3

L.P-2

Geometry- TBP

Shape- T-shaped

eg- ClF3

Front 36

AB2 with 3 L.P

Back 36

B.P- 2

L.P-3

Geometry- TBP

Shape- linear

eg- XeF2

Front 37

AB4 with 2 L.P

Back 37

B.P-4

L.P-2

Geometry- octahedral

Shape-square planar

eg- XeF4

Front 38

AB6 type molecule

Back 38

B.P- 6

L.P-1

Geometry- pentagonal bipyramidal

Shape-distorted octahedron

eg- XeF6

Front 39

Bond Angle Comparison

Back 39

Depends on electronegativity of central atom. Higher the electronegativity higher the bond angle. Repulsion b/w surrounding atoms should be less for higher bond angle

Front 40

Hybridisation

Back 40

homogenisation of orbitals of valence shell w/ diff. energy & orientation to give new set of orbitals w/ similar shape and energy

Front 41

Resonance

Back 41

phenomenon of existence of two or more possible lewis dot structures of a molecule or ion

Front 42

Resonatng Structures

Back 42

diff structures for the same molecule/ion

Front 43

Molecular Orbital Theory

Back 43

1.e- of a molecule are present in various mol orbitals

2.e- are influenced by two or more nuclei

3.two atomic orbs combine to form 1 bonding orb and 1 antibonding orb

Front 44

Hydrogen Bond

Back 44

dipole-dipole interaction b/w between H of a molecule/ion carrying + charge and - charge on F,O,N,etc

Front 45

Dipole

Back 45

system of 2 equal and opposite charges separated by a small distance

Front 46

Dipole Moment

Back 46

Dip. Mom. (μ)= charge(Q) x distance(d)

vector quantity , SI unit- Debye (D)

1D= 3.33 x 10^-30 Cm

Front 47

LCAO

Back 47

Linear Combination of Atomic Orbitals

Front 48

Backbonding

Back 48

delocalisation of charge b/w 2 atoms involved in a sigma bond and one of them has a filled orbital having one l.p and the other has a vacant orbital.

Front 49

Fajan's Rule

Back 49

when a cation having high polarising power approaches an anion having high polarizability then the bond b/w them is predominantly covalent

Front 50

Polarizability

Back 50

measure of tendency to allow distortion of its e- cloud