HSSLIVE Plus One Chemistry Chapter 4: Chemical Bonding and Molecular Structure Notes – Hsslive | Hsslive.co.in | Hss Live

This chapter reveals how atoms join together to form molecules through different types of chemical bonds. Students explore ionic, covalent, and metallic bonding, along with intermolecular forces that determine physical properties. The concept of hybridization and molecular orbital theory are introduced to explain molecular geometries and bond characteristics. Through this knowledge, students begin to understand how molecular structure determines the properties and reactivity of substances.

Chapter 4: Chemical Bonding and Molecular Structure

Chemical bonding is the force that holds atoms together in a molecule or compound. The type of bond formed depends on the electronic configuration of the combining atoms and determines the physical and chemical properties of the resulting compound.

Types of Chemical Bonds

There are three primary types of chemical bonds:

Ionic bonds form when electrons are completely transferred from one atom to another, creating positively and negatively charged ions that attract each other. These typically form between metals (which readily lose electrons) and non-metals (which readily gain electrons). For example, in sodium chloride (NaCl), sodium donates its valence electron to chlorine.

Covalent bonds form when atoms share electrons to achieve stable electron configurations. These typically form between non-metals with similar electronegativities. For example, in a chlorine molecule (Cl₂), each chlorine atom shares one electron with the other.

Coordinate covalent bonds (or dative bonds) form when both shared electrons come from the same atom. The donor atom provides the electron pair, while the acceptor atom provides an empty orbital. For example, in the ammonium ion (NH₄⁺), the nitrogen atom donates a lone pair to the hydrogen ion (H⁺).

Ionic Bonding

Ionic bonds result from the electrostatic attraction between oppositely charged ions. The formation of an ionic bond involves several steps:

A metal atom loses one or more electrons to form a positive ion (cation).
A non-metal atom gains these electrons to form a negative ion (anion).
The oppositely charged ions attract each other, forming an ionic bond.

For example, in the formation of calcium fluoride (CaF₂), calcium loses two electrons to form Ca²⁺, while two fluorine atoms each gain one electron to form F⁻, resulting in an ionic compound.

Ionic compounds typically have high melting and boiling points due to the strong electrostatic forces between ions. They conduct electricity when dissolved in water or melted, as the ions become mobile and can carry current. They are generally soluble in polar solvents like water but insoluble in non-polar solvents like benzene.

Covalent Bonding

Covalent bonds involve the sharing of electron pairs between atoms. The number of shared pairs determines the bond type:

Single bond: One pair of shared electrons (e.g., H-H in hydrogen molecule) Double bond: Two pairs of shared electrons (e.g., O=O in oxygen molecule) Triple bond: Three pairs of shared electrons (e.g., N≡N in nitrogen molecule)

Covalent bonds can be polar or non-polar:

Polar covalent bonds form between atoms with different electronegativities, resulting in unequal sharing of electrons. The more electronegative atom pulls the shared electrons closer, creating a partial negative charge (δ-) on that atom and a partial positive charge (δ+) on the other. For example, in HCl, chlorine is more electronegative than hydrogen, making the H-Cl bond polar.

Non-polar covalent bonds form between atoms with the same or similar electronegativities, resulting in equal sharing of electrons. For example, in Cl₂, both chlorine atoms have the same electronegativity, making the Cl-Cl bond non-polar.

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory helps predict the three-dimensional shape of molecules. It is based on the principle that electron pairs (both bonding and non-bonding) around a central atom repel each other and arrange themselves to minimize repulsion.

The arrangement of electron pairs determines the molecular geometry. Some common arrangements include:

Linear: Two electron pairs (e.g., BeCl₂, CO₂)
Trigonal planar: Three electron pairs (e.g., BF₃)
Tetrahedral: Four electron pairs (e.g., CH₄)
Trigonal bipyramidal: Five electron pairs (e.g., PCl₅)
Octahedral: Six electron pairs (e.g., SF₆)

If some electron pairs are lone pairs (not involved in bonding), the molecular geometry differs from the electron pair geometry. For example, NH₃ has four electron pairs around nitrogen (tetrahedral electron pair geometry), but one is a lone pair, resulting in a pyramidal molecular geometry.

Hybridization

To explain the observed geometries of molecules, the concept of hybridization was introduced. Hybridization is the mixing of atomic orbitals of similar energies to form new hybrid orbitals of equal energy.

Different types of hybridization include:

sp hybridization: Mixing of one s and one p orbital to form two equivalent sp hybrid orbitals. This results in a linear geometry (180° bond angle), as seen in BeCl₂ and acetylene (C₂H₂).
sp² hybridization: Mixing of one s and two p orbitals to form three equivalent sp² hybrid orbitals. This results in a trigonal planar geometry (120° bond angles), as seen in BF₃ and ethylene (C₂H₄).
sp³ hybridization: Mixing of one s and three p orbitals to form four equivalent sp³ hybrid orbitals. This results in a tetrahedral geometry (109.5° bond angles), as seen in CH₄ and NH₃.
sp³d hybridization: Mixing of one s, three p, and one d orbital to form five equivalent sp³d hybrid orbitals. This results in a trigonal bipyramidal geometry, as seen in PCl₅.
sp³d² hybridization: Mixing of one s, three p, and two d orbitals to form six equivalent sp³d² hybrid orbitals. This results in an octahedral geometry, as seen in SF₆.

Molecular Orbital Theory

Molecular Orbital (MO) theory provides a more complete description of bonding by considering the wave-like properties of electrons. According to this theory, atomic orbitals combine to form molecular orbitals that extend over the entire molecule.

When atomic orbitals combine, they form two types of molecular orbitals:

Bonding molecular orbitals have lower energy than the original atomic orbitals. Electrons in these orbitals are concentrated in the region between nuclei, increasing the electron density and strengthening the bond.

Antibonding molecular orbitals have higher energy than the original atomic orbitals. Electrons in these orbitals are concentrated away from the region between nuclei, decreasing the electron density and weakening the bond.

The bond order, which indicates the stability of a bond, is calculated as: Bond order = (Number of electrons in bonding MOs – Number of electrons in antibonding MOs) / 2

A positive bond order indicates a stable molecule, while a bond order of zero indicates an unstable molecule.

For example, in O₂, the bond order is (8-4)/2 = 2, indicating a double bond, which matches experimental observations.

Complete Chapter-wise Hsslive Plus One Chemistry Notes

Our HSSLive Plus One Chemistry Notes cover all chapters with key focus areas to help you organize your study effectively: