Chemical Bonding: Full Chapter Revision For NEET

Chemical Bonding is a major chapter for NEET because it connects electronic configuration with molecular structure and properties. It explains why atoms combine and how the shape of a molecule decides its behaviour, which is why questions from this chapter appear regularly and often test several ideas at once.

The key areas include the Octet Rule, formal charge, Valence Bond Theory, hybridisation, VSEPR theory, bond parameters, dipole moment, Molecular Orbital Theory, resonance, and hydrogen bonding. Molecular structure questions deserve special attention, since recent NEET papers have asked about the geometry and lone pairs of molecules like ClF3 and about formal charge in ozone. Hydrogen bonding is also useful beyond this chapter, as it links to Organic Chemistry and p-Block Elements.

Major Concepts In Chemical Bonding

Chemical Bonding is built step by step. It starts with why atoms combine (the Octet Rule and formal charge), moves to how bonds form (orbital overlap, Valence Bond Theory and hybridisation), and then to the shapes and properties that result (VSEPR theory, bond parameters, dipole moment, Molecular Orbital Theory and hydrogen bonding). Learning the topics in this order makes each one easier to follow.

Octet Rule And Its Limitations

The Octet Rule states that atoms form chemical bonds to achieve a stable electronic configuration, generally containing eight electrons in the valence shell. The rule is based on the stability of noble gases.

However, it does not explain every stable species.

Electron-Deficient Species

Some central atoms have fewer than eight electrons but remain stable. These are called electron-deficient or hypovalent species. They may also act as Lewis acids.

Examples include:

• BeCl2

• BCl3

Hypervalent Species

Some atoms have more than eight electrons around them. These are called hypervalent or electron-rich species.

Examples include:

• PCl5

• SF4

• SF6

• ICl3

• IF7

Odd-Electron Species

Species with an odd total number of electrons cannot complete a complete octet.

Examples include:

• NO

• NO2

Noble-Gas Compounds

Noble gases can form compounds, which is another limitation of the rule. Examples include XeF2, XeF4, XeF6, XeO3, KrF2, and RnF2.

Transition-metal ions and their compounds also generally do not follow the Octet Rule. Therefore, the rule is useful as an introductory model but is not universal.

Formal Charge

Formal charge is a hypothetical charge assigned to an atom in a bonded structure. It compares the valence electrons of the free atom with the electrons assigned to it in the structure.

Formal Charge = Valence Electrons − Non-Bonded Electrons − Half Of Bonded Electrons

The valence-electron count can be obtained from the group number:

The entire calculation is based on the electrons present in the structure. The sum of all formal charges must equal the overall charge of the species.

The RBI Rule For Formal Charge

The memory aid RBI helps recall important sources of formal charge:

• R: Resonance Or Coordinate Bond
  In a coordinate bond, the donor atom receives a formal charge of plus one, while the acceptor receives a formal charge of minus one.

• B: Background Or Pre-Existing Charge
  Any charge already present on an atom contributes to its formal charge.

• I: Ion Charge
  The sum of formal charges must equal the charge on the complete ion.

In ozone, the three oxygen atoms have formal charges of zero, plus one, and minus one in a given representation. Their sum is zero, matching neutral ozone.

Valence Bond Theory And Orbital Overlap

Valence Bond Theory explains bond formation through orbital overlap and reduction in system energy.

Bond formation occurs because the energy of the system is minimised. As two atoms approach, attractive forces lower the energy. At the equilibrium bond distance, attractive and repulsive forces balance. If the atoms move closer, repulsion increases and the energy rises.

The energy curve provides two important bond parameters:

• The x-coordinate gives bond length.

• The y-coordinate gives bond energy.

According to VBT:

• Valence-shell electrons participate in bonding.

• A covalent bond contains two electrons.

• The bonding electrons must have opposite spins.

• Atomic orbitals overlap to form bonds.

Types Of Orbital Overlap

An s orbital can form only sigma overlap. A p orbital can form sigma and pi overlap. A d orbital can form sigma, pi, and delta overlap.

A sigma bond is stronger than a pi bond in the basic overlap comparison. When the principal quantum numbers of two orbitals are represented by n1 and n2, overlap strength decreases as n1 plus n2 increases. If the sum is the same, the stated comparison is s less than p for sigma overlap and p less than d for pi overlap.

Hybridisation

Hybridisation is the mixing of atomic orbitals of nearly similar energy to form new hybrid orbitals. The number of hybrid orbitals equals the number of atomic orbitals mixed.

Hybrid orbitals generally form sigma bonds. Pi bonds are always formed by pure, unhybridized orbitals.

The percentage of s-character is:

Percentage s-character = Number Of s Orbitals Mixed divided by Total Orbitals Mixed, multiplied by 100

Thus:

• sp has 50 percent s-character.

• sp2 has about 33 percent s-character.

• sp3 has 25 percent s-character.

VSEPR Theory And Molecular Geometry

Valence Shell Electron Pair Repulsion Theory states that electron pairs around a central atom arrange themselves to reduce repulsion.

The repulsion order is:

Lone Pair–Lone Pair > Lone Pair–Bond Pair > Bond Pair–Bond Pair

VSEPR count equals the number of bond pairs plus the number of lone pairs.

In trigonal bipyramidal geometry, lone pairs occupy equatorial positions because this reduces repulsion. Therefore:

• SF4 has see-saw geometry.

• ClF3 has T-shaped geometry.

• XeF2 and I3 minus have linear geometry.

In octahedral arrangements:

• Five bond pairs and one lone pair give square pyramidal geometry.

• Four bond pairs and two opposite lone pairs give square planar geometry.

For XeOF4, six electron domains are present: five bonding regions and one lone pair. Its molecular shape is square pyramidal.

Resonance And Bond Parameters

Resonance occurs when more than one valid structure can be drawn for a species. Resonance has an averaging effect. It produces average bond order and equalises equivalent bonds.

In ozone, both oxygen–oxygen bonds have equal lengths. Their length is greater than a pure double bond but less than a pure single bond because each bond has partial double-bond character.

Important Bond Parameters

• Bond Order: Average number of bonds between two atoms.

• Bond Length: Distance between the nuclei of bonded atoms.

• Bond Energy: Energy required to break a bond.

• Bond Angle: Angle between two bonds around a central atom.

Higher bond order gives shorter bond length. Bond length also depends on atomic size and electronegativity difference.

The important halogen bond-enthalpy order is:

Cl2 > Br2 > F2 > I2

F2 is an exception because strong electron–electron repulsion occurs between the two small fluorine atoms.

Bond-Angle Comparison

Use these rules in order

1. Different hybridisation: sp > sp2 > sp3

2. Same hybridisation: more lone pairs give a smaller bond angle.

3. If hybridisation and lone-pair count are the same, compare electronegativity.

The common order is:

CH4 > NH3 > H2O

Dipole Moment And Molecular Polarity

Dipole moment is a vector quantity. Bond polarity depends on electronegativity difference, but molecular polarity depends on the vector sum of all bond dipoles.

• Zero resultant dipole moment means a nonpolar molecule.

• Nonzero resultant dipole moment means a polar molecule.

CO2 is nonpolar because it is linear and its identical bond dipoles cancel. BF3 is nonpolar because of its symmetrical trigonal-planar structure. NH3, H2O, and NF3 are polar.

A molecule with one lone pair is generally polar because its lone-pair dipole cannot be cancelled by another lone-pair dipole.

NH3 is more polar than NF3 because the bond dipoles in NH3 support the lone-pair dipole, while those in NF3 oppose it.

Molecular Orbital Theory

For species containing up to 14 electrons, the molecular-orbital order includes sigma and pi bonding orbitals followed by their antibonding orbitals. Above 14 electrons, the order of the relevant 2p orbitals changes. The lecture uses the memory aid 12, 21 for this interchange.

Bond order is calculated as:

Bond Order = Half Of Bonding Electrons Minus Antibonding Electron

Molecular species are paramagnetic when unpaired electrons are present. O2 is paramagnetic because it has two unpaired electrons. CO and NO plus, with 14 electrons in the stated examples, are diamagnetic.

Hydrogen Bonding

Hydrogen bonding occurs when hydrogen is directly bonded to fluorine, oxygen, or nitrogen.

Intermolecular Hydrogen Bonding

This occurs between different molecules and generally increases:

• Boiling point

• Melting point

• Viscosity

Intramolecular Hydrogen Bonding

This occurs within the same molecule. Compared with intermolecular bonding, it generally produces a lower boiling point, melting point, and viscosity.

Chemical Bonding links electronic configuration to molecular shape, polarity and properties. This revision covers the Octet Rule, hybridisation, VSEPR, bond parameters, dipole moment, MOT and hydrogen bonding. It also highlights key exceptions like F2 bond enthalpy and exam-favourite structures like ClF3.

Related Sources

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