Chemical Bonding Definition
Chemical Bonding refers to the formation of a chemical bond between two or more atoms, molecules, or ions to give rise to a chemical compound.
The attractive force which holds various constituents (atoms, ions, etc.) together in different chemical species is called a chemical bond.
- Octet rule
- Formal charge
- Ionic bond
- Lattice enthalpy
- Covalent bond
- VSEPR Theory
- Fajans’ Rule
- VBT Theory
- Hybridization in chemical bonding : sp, sp2, sp3, sp3d, sp3d2 and sp3d3: their geometry and molecular structure.
- Atoms form chemical bonds in order to complete their octet i.e. eight electrons in their valence shell.
- Electrons present in the last shell of atoms are called valence electrons.
Exceptions to the Octet Rule:
- Incomplete octet for the central atom: LiCl, BeH2 and BCl3
- Species with odd number of electrons:NO2,NO
- The Expanded octet for the central atom: PF5, SF6 and H2SO4
- Difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.
- Formal charge = [Total number of valence electrons in the free atom ) – (Total number of lone pairs of electrons) -1/2(Total number of shared electrons i.e. bonding electrons)]
- Bond formed by transfer of electron
- Between metal and non-metal
The energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.
- Bond formed by sharing of electron pair
THE VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY:
- VSEPR Theory is very important for chemical bonding (in terms of exams as well).
- The shape of a molecule depends upon the number of valence shell electron pairs (bonded or non-bonded) around the central atom.
- Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
- These pairs of electrons tend to occupy such positions in space that minimise repulsion and thus maximise distance between them.
- The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at maximum distance from one another.
- A multiple bond is treated as if it is a single electron pair. And two or three electron pairs of a multiple bond are treated as a single super pair.
- Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure
The repulsive interaction of electron pairs decrease in the order:
Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp)
- It accounts for the covalent character in ionic compounds.
- The difference between the number of valence electrons in an isolated atom and number of electrons assigned to that atoms in Lewis structure.
Valence bond theory (VBT) of Chemical Bonding
- A covalent bond is formed by overlapping of valence shell atomic orbital of the two atoms having unpaired electron.
- Greater the overlapping of atomic orbital higher is the strength of chemical bond.
Types of bonds:
- σ-Bond : When covalent bond is formed by overlapping of atomic orbitals along the same axis, it is called s – bond. Such type of bond is symmetrical about the line joining the two nuclei.
- pi(π ) bond : The atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to sidewise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.
Hybridization in Chemical Bonding
Two or more than two orbitals having less energy difference are combined and form same number of new orbital having similar shape and energy but different orientation.
Types of hybridization in chemical bonding:
Direct questions comes on types of hybridization whenever it comes to chemical bonding. The following data is sufficient for the direct questions of boards, JEE and NEET as well.
sp hybridization is observed when one s and one p orbital in the same main shell of an atom mix to form two new equivalent orbitals. The new orbitals formed are sp hybrid orbitals.
- Geometry and shape: linear
- Bond Angle: 180 degree
- 2 bond pair & zero lone pair
- Examples: BeF2, BeH2, BeCl2 C2H2.
linear shape Co2
In sp2 hybridization, one ‘s’ orbital and 2 ‘p’ orbitals belonging to the same shell of an atom mix together. As a result, they form three new equivalent orbitals.
Different possible shapes in sp2 hybridization-
All bond pair(3bp + 0 lp)
- Geometry and shape : trigonal planer
- Bond Angle : 120 degree
- Examples : BF3, BH3
2 bond pair + 1 lone pair
- Geometry : trigonal planer
- Shape : V-shape
- Examples : SO2
In sp3 hybridization, one ‘s’ orbital and 3 ‘p’ orbitals belonging to the same shell of an atom mix together. As a result, they form four new equivalent orbitals.
- Geometry: tetrahedral
Different possible shapes in sp3 hybridization-
All bond pair(4bp + 0 lp)
- Shape : tetrahedral
- Bond angle : 109°28’
- Example : CH4
3 bond pair + 1 lone pair
- Shape : Pyramidal
- Bond angle : about 107°
- Example : NH3
2 bond pair + 2 lone Pair
- Shape : V-shape or bent
- Bond angle : 104.5
- Example : H2O
In sp3d hybridization, one ‘s’ orbital, 3 ‘p’ orbitals and 1’d’ orbitals belonging to the same shell of an atom mix together. As a result, they form five new equivalent orbital.
- Geometry- trigonal bi-pyramidal
- Bond angle -120 & 90
Different possible shape in sp3d hybridization-
All bond pair (5 bond pair+ 0 lone pair)
- Shape : trigonal bi-pyramidal
- Example : PCl5
4 bond pair + 1 lone pair
- Shape : see-saw
- Example : SF4
3 bond pair + 2 lone pair
- Shape : T-shape
- Example : ClF3
2 bond pair + 3 lone pair
- Shape : linear
- Example : XeF2
- 1S + 3P + 2D orbitals
- Geometry : octahedral
- Bond angle: 90
Different possible shapes in sp3d2 hybridization-
All bond pair(6 bond pair + 0lone pair)
- Shape : octahedral
- Example : SF6
5 bond pair + 1 lone pair
- Shape : square pyramidal
- Example : ClF5, BrF5
4 bond pair + 2 lone pair
- Shape : square planar
- Example : XeF4
- Geometry : Pentagonal bipyramidal
- Bond angle : 72 & 90
- Example : IF7