Chemical Bonding and Molecular Structure is an important NEET chapter because concepts like hybridisation, molecular geometry, polarity, and bonding are used in many Organic and Inorganic Chemistry topics. You may sometimes find the chapter confusing while identifying shapes, bond angles, orbital overlap, or electron arrangement in molecules. Regular revision of structures, formulas, and bonding theories along with practice questions helps you understand concepts more clearly and improve accuracy in NEET Chemistry. With Physics Wallah, you can strengthen these concepts through mind maps, sample papers, MCQs, PYQs, and formula sheets designed for quick revision and exam-focused practice.

Ionic Bond

An ionic bond is formed when one atom completely transfers electrons to another atom. The atom losing electrons becomes positively charged, while the atom gaining electrons becomes negatively charged. These oppositely charged ions attract each other through electrostatic force and form an ionic bond.

Ionic bonds are generally formed between metals and non-metals because metals tend to lose electrons easily, whereas non-metals tend to gain electrons.

For example, in sodium chloride (NaCl), sodium transfers one electron to chlorine. Sodium becomes Na⁺ , and chlorine becomes Cl⁻. The strong attraction between these ions forms the ionic compound.

Important characteristics of ionic compounds include:

• High melting and boiling points because strong electrostatic forces hold the ions together.

• Solubility in water due to the polar nature of water molecules.

• Ability to conduct electricity in a molten or aqueous state because ions become free to move.

Ionic compounds are usually hard and brittle solids.

Covalent Bond

A covalent bond is formed when atoms share electrons to achieve stability. Covalent bonding generally occurs between non-metal atoms.

The shared electrons remain between the bonded atoms and help hold them together. Depending on the number of electron pairs shared, covalent bonds can be single, double, or triple bonds.

For example:

• H₂ molecule contains a single covalent bond.

• O₂ molecule contains a double bond.

• N₂ molecule contains a triple bond.

Covalent compounds generally have lower melting and boiling points compared to ionic compounds because intermolecular forces are weaker.

Covalent molecules may be polar or non-polar depending on the distribution of electrons. Water (H₂O) is polar, whereas methane (CH₄) is non-polar.

Lewis Dot Structures

Lewis Dot Structures represent valence electrons around atoms using dots. These structures help explain how atoms bond with each other and how electrons are arranged in molecules.

Gilbert N. Lewis introduced this method to explain chemical bonding. In Lewis structures, the symbol of the element is written at the centre, and dots around the symbol represent valence electrons.

Shared electron pairs between atoms represent covalent bonds. Lone pairs represent non-bonding electrons that remain on atoms.

Lewis structures help in:

• Predicting bond formation

• Determining lone pairs

• Understanding molecular geometry

• Calculating formal charge

• Explaining resonance structures

Bond representation in Lewis structures:

• Single bond = one shared pair

• Double bond = two shared pairs

• Triple bond = three shared pairs

Formula for Formal Charge

Formal Charge = Valence Electrons − Non-bonding Electrons − (Bonding Electrons / 2)

Octet Rule and Exceptions

According to the octet rule, atoms tend to gain, lose, or share electrons until they achieve eight electrons in their valence shell. This configuration provides stability similar to noble gases.

The octet rule successfully explains bonding in many molecules, but several important exceptions exist.

Some molecules contain incomplete octets. For example, BF₃ and BeCl₂ do not have eight electrons around the central atom.

Some molecules contain expanded octets where the central atom has more than eight electrons. Examples include PCl₅ and SF₆.

Certain molecules, such as NO and NO₂ contain an odd number of electrons and cannot satisfy the octet rule completely.

Understanding these exceptions is important because NEET often includes conceptual questions related to the octet rule limitations.

VSEPR Theory

VSEPR stands for Valence Shell Electron Pair Repulsion Theory. According to this theory, electron pairs around the central atom repel one another and arrange themselves as far apart as possible to minimise repulsion.

This theory helps predict molecular geometry and bond angles. Both bonding pairs and lone pairs affect molecular shape, but lone pairs create greater repulsion because they occupy more space.

The order of repulsion is:

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

Some important molecular shapes are given below:

The decrease in bond angle from CH₄ to NH₃ and H₂O occurs because lone pairs increase repulsion and compress bond angles.

Hybridisation

Hybridisation refers to the mixing of atomic orbitals to form new hybrid orbitals having similar energy and shape. Hybrid orbitals form stronger and more stable covalent bonds.

Hybridisation helps explain molecular geometry, bond angle, and orientation of bonds in molecules.

The different types of hybridisation are:

In sp hybridisation, one s orbital mixes with one p orbital to form two hybrid orbitals. In sp² hybridisation, one s orbital mixes with two p orbitals to form three hybrid orbitals.

Sigma bonds are generally formed through hybrid orbitals, while pi bonds are formed by unhybridised p orbitals.

Hybridisation is frequently asked in NEET through molecular geometry and bond angle questions.

Molecular Orbital Theory

Molecular Orbital Theory explains chemical bonding through the combination of atomic orbitals to form molecular orbitals spread over the entire molecule.

When atomic orbitals combine, two molecular orbitals are formed:

• Bonding molecular orbitals, which are lower in energy and increase stability.

• Antibonding molecular orbitals, which are higher in energy and decrease stability.

Electrons first fill bonding orbitals before entering antibonding orbitals.

Bond Order Formula

Bond Order = 1/2 × (Number of electrons in Bonding Orbitals − Number of electrons in Antibonding Orbitals)

Bond order helps determine:

• Stability of molecules

• Strength of bonds

• Bond length

A higher bond order indicates a stronger and shorter bond. If the bond order is zero, the molecule cannot exist.

Molecular Orbital Theory also explains magnetic behaviour. Oxygen (O₂) is paramagnetic because it contains unpaired electrons in antibonding orbitals.

Bond Parameters

Bond parameters describe different characteristics of chemical bonds and molecular structure.

Bond length refers to the distance between the nuclei of two bonded atoms. Stronger bonds usually have shorter bond lengths.

Bond angle refers to the angle between two covalent bonds around the central atom. Bond angles depend on molecular geometry and electron pair repulsion.

Bond enthalpy is the energy required to break one mole of bonds in the gaseous state. Stronger bonds possess higher bond enthalpy.

These parameters help explain molecular stability and reactivity.

Dipole Moment and Polarity

Dipole moment measures the separation of positive and negative charges in a molecule. It helps determine whether a molecule is polar or non-polar.

When atoms with different electronegativities form bonds, electrons are shared unequally. This creates partial positive and negative charges.

Dipole Moment Formula

μ = Q × d

Where:

• μ = dipole moment

• Q = magnitude of charge

• d = distance between charges

Polar molecules possess a non-zero dipole moment, whereas non-polar molecules possess a zero dipole moment.

For example:

• H₂O is polar because of its bent geometry.

• CO₂ is non-polar because dipole moments cancel each other due to linear geometry.

Dipole moment affects solubility, boiling point, and intermolecular interactions.

Hydrogen Bonding

Hydrogen Bonding is a weak attractive force between a hydrogen atom attached to a highly electronegative atom and another electronegative atom nearby.

Hydrogen Bonding occurs only when hydrogen is bonded to fluorine, oxygen, or nitrogen.

There are two types of Hydrogen Bonding:

Hydrogen Bonding is responsible for:

• High boiling point of water

• Solubility of alcohols in water

• Molecular association in biological compounds

Hydrogen Bonding is important because it strongly affects the physical properties of compounds.

Resonance

Resonance occurs when a molecule can be represented by two or more valid Lewis structures differing only in electron arrangement. The actual molecule exists as a resonance hybrid of all contributing structures.

Resonance stabilises molecules because electrons become delocalised over multiple atoms.

Examples of resonance include:

• Benzene

• Ozone (O₃)

• Carbonate ion (CO₃²⁻)

Resonance is important in understanding the stability, bond length, and reactivity of molecules.

Chemical Bonding and Molecular Structure Complete Study Resources

Physics Wallah provides multiple study and revision resources for chapter-wise NEET preparation. These resources help improve conceptual understanding, formula revision, and problem-solving skills.