Unit 4: Chemical Bonding-II

Weak Chemical Forces (Intermolecular Forces)

These are the attractive forces between molecules (or between ions and molecules). They are much weaker than covalent or ionic bonds (intramolecular forces) but are responsible for the physical properties of matter (e.g., boiling point, melting point, solubility).

(i) van der Waals Forces

This is a general term for all intermolecular forces *except* hydrogen bonds and ion-dipole forces. It is often used to specifically refer to the three types of dipole-based interactions:

1. Ion-Dipole Forces

This is the force between an ion (e.g., Na⁺) and a polar molecule (e.g., H₂O). This is the strongest of the weak forces. It is the primary force responsible for the dissolution of ionic salts in polar solvents like water.

2. Dipole-Dipole Interactions

This is the electrostatic attraction between two permanent polar molecules. The δ⁺ end of one molecule attracts the δ^- end of another.
Example: HCl ⋯ HCl.
These forces are significant only when molecules are close. They are responsible for the higher boiling points of polar molecules compared to non-polar ones of similar mass (e.g., HCl vs F₂).

3. Induced Dipole Interactions

This category involves at least one non-polar molecule.

• Dipole-Induced Dipole: A permanent dipole (like H₂O) can distort the electron cloud of a non-polar molecule (like O₂), inducing a temporary, weak dipole. This force explains why non-polar gases like O₂ can dissolve slightly in water.
• Instantaneous Dipole-Induced Dipole (London Dispersion Forces):
  • This is the weakest intermolecular force, but it is universal (exists between *all* molecules, polar or non-polar).
  • It arises from the random motion of electrons. At any given instant, a non-polar molecule (like He or CH₄) can have a temporary, instantaneous dipole.
  • This temporary dipole can then induce a temporary dipole in a neighboring molecule, leading to a weak, fleeting attraction.
  • Key Point: The strength of London forces increases with:
    1. Increased molar mass / number of electrons (more polarizable cloud).
    2. Increased surface area (more points of contact). (e.g., n-pentane has a higher boiling point than neopentane).

Summary of Intermolecular Forces (Strongest to Weakest)

Force	Interaction	Example
Ion-Dipole	Ion + Polar Molecule	Na+ in H2O
Hydrogen Bond (see below)	H bonded to N/O/F	H2O ⋯ H2O
Dipole-Dipole	Permanent Dipole + Permanent Dipole	HCl ⋯ HCl
Dipole-Induced Dipole	Permanent Dipole + Non-polar	H2O ⋯ O2
London Dispersion	Instantaneous Dipole + Induced Dipole	CH4 ⋯ CH4

Hydrogen Bonding

Definition: A special, strong type of dipole-dipole interaction that occurs when a hydrogen atom, which is covalently bonded to a highly electronegative atom (N, O, or F), is attracted to another nearby electronegative atom.

The H-N, H-O, or H-F bond is extremely polar (δ⁺H - X^δ-). The tiny, highly positive H is then strongly attracted to a lone pair on a N, O, or F atom of another molecule.

Theories of Hydrogen Bonding

1. Electrostatic Theory: This is the simplest view. The H-bond is just a very strong dipole-dipole interaction between the H^δ+ and the X^δ-.
2. Valence Bond / MO Theory: This more advanced view suggests a degree of covalent character. The lone pair orbital (X₂) and the H-X₁ antibonding orbital (σ^*) can overlap, forming a weak 3-center-4-electron bond. This helps explain the short, directional nature of H-bonds.

Types of Hydrogen Bonds

• Intermolecular H-bonding: Occurs between two or more different molecules.
  Examples: H₂O, HF, alcohols (R-OH), NH₃.
  Consequence: Drastically increases boiling points and melting points. (e.g., H₂O boils at 100°C, while H₂S boils at -60°C).
• Intramolecular H-bonding: Occurs within a single molecule.
  Examples: o-nitrophenol, salicylaldehyde.
  Consequence: "Chela-tion" or ring formation. This *prevents* the molecule from H-bonding with other molecules, leading to a lower boiling point and lower water solubility compared to its para-isomer (which can only form intermolecular H-bonds).

Metallic Bond

This is the force of attraction that holds atoms together in a metallic solid. It is responsible for properties like high electrical/thermal conductivity, malleability, ductility, and luster.

(ii) Qualitative Idea of Valence Bond and Band Theories

1. Valence Bond Theory (Electron Sea Model)

This is the simplest model. It visualizes the metal as a rigid lattice of positive ions (kernels or cations) sitting in a "sea" of delocalized valence electrons. The metallic bond is the electrostatic attraction between the positive cations and the mobile, negative electron sea. This mobile sea easily explains conductivity.

A VBT extension (by Pauling) describes the metallic bond as a resonance of a large number of covalent bonds between all adjacent atoms.

2. Band Theory (Molecular Orbital Theory for Solids)

This is the most successful model. It applies MOT to an entire crystal.

• When N atoms come together (where N is huge, ≈ 10²³), their N atomic orbitals (e.g., 3s) overlap to form N molecular orbitals.
• These N MOs are so numerous and so closely spaced in energy that they merge into a continuous energy band.
• Valence Band: The energy band formed from the valence atomic orbitals. In a metal at 0K, this band is either partially filled (like Na, 3s¹) or is full but overlaps with an empty band (like Mg, 3s² overlaps 3p⁰).
• Conduction Band: The next higher, empty (or partially filled) energy band.
• Band Gap (E_g): The energy difference between the top of the valence band and the bottom of the conduction band.

Semiconductors, Insulators, and Defects in Solids

Semiconductors and Insulators (explained by Band Theory)

The electrical conductivity of a solid is determined by the size of its band gap (E_g).

• Conductors (Metals): E_g = 0. The valence band is partially filled, or it overlaps with the conduction band. Electrons can move freely into higher energy states with minimal energy input, allowing current to flow.
• Insulators: E_g is very large (e.g., > 3 eV). The valence band is full, and the conduction band is empty. A huge amount of energy is needed to promote an electron, so no current flows under normal conditions (e.g., diamond, quartz).
• Semiconductors: E_g is small (e.g., 0.5 - 2 eV). The valence band is full, but the band gap is small enough that thermal energy (at room temp) can promote some electrons to the conduction band, allowing a small current to flow (e.g., Silicon, Germanium).
  • Conductivity of semiconductors increases with temperature (more e- promoted).
  • Conductivity of metals decreases with temperature (lattice vibrations hinder e- flow).

Defects in Solids

An ideal crystal has a perfect, repeating arrangement of atoms. Any deviation from this is a defect or imperfection. Defects are crucial for the properties of many materials.

The syllabus mentions "defects in solids" in the context of metallic bonding, which most likely refers to point defects in ionic/metallic crystals.

Point Defects (Stoichiometric)

These defects do not disturb the overall stoichiometry (ratio of atoms) of the solid.

1. Schottky Defect:
   • What it is: A pair of vacancies (one cation and one anion) are missing from the lattice to maintain electrical neutrality.
   • Effect: Decreases the density of the crystal.
   • Found in: Highly ionic compounds where cation and anion are of similar size (e.g., NaCl, KCl, CsCl).
2. Frenkel Defect:
   • What it is: A smaller ion (usually the cation) is dislocated from its normal lattice site and moves into an interstitial site (a small gap).
   • Effect: Does not change the density of the crystal.
   • Found in: Compounds with a large difference in ion size (e.g., AgCl, AgBr, ZnS).

Point Defects (Non-Stoichiometric)

• Metal Excess Defect (F-centers): An anion is missing from its site, and the hole is filled by an electron to maintain neutrality. This "trapped electron" is an F-center (from German Farbe, "color") and is responsible for the color of compounds like NaCl (yellow) when heated in Na vapor.