Chemical Bonds

1. Meaning of Chemical Bonding

Chemical bonding is the force that holds atoms or ions together in a molecule or compound to achieve stability. Atoms bond by sharing, transferring, or pooling electrons to attain stable electron configurations, often resembling noble gases.

2. Interatomic Bonding

2.1 Types of Interatomic Bonds

1. Ionic Bonding: Involves the transfer of electrons from a metal to a non-metal.
2. Covalent Bonding: Involves the sharing of electron pairs between non-metal atoms.
3. Metallic Bonding: Delocalized electrons move freely in a lattice of metal ions.

3. Ionic Bonding

3.1 Formation of Ionic Bonds and Compounds

• Occurs between atoms with large differences in electronegativity.
• Process:
  • A metal atom loses electrons to form a cation.
  • A non-metal atom gains electrons to form an anion.
  • Electrostatic attraction between oppositely charged ions creates the ionic bond.

Example: Sodium Chloride (NaCl):

3.2 Properties of Ionic Compounds

• Melting and Boiling Points: High, due to strong ionic bonds.
• Electrical Conductivity: Conduct in molten or aqueous states.
• Solubility: Soluble in polar solvents (e.g., water).
• Hardness: Brittle and crystalline structure.

3.3 Naming Ionic Compounds

• Name the cation first, followed by the anion.
• Use Roman numerals for transition metals.
  • Example: $\text{FeCl}_3$: Iron (III) chloride.

4. Covalent Bonding

4.1 Formation of Covalent Bonds and Compounds

• Formed when two non-metal atoms share electrons.
• Electron pairs are localized between nuclei, stabilizing the molecule.

Example: Formation of $\text{H}_2$:

4.2 Properties of Covalent Compounds

• Melting and Boiling Points: Low, due to weak intermolecular forces.
• Electrical Conductivity: Poor, as no free ions or electrons exist.
• Solubility: Soluble in non-polar solvents (e.g., benzene).

4.3 Coordinate (Dative) Covalent Bonding

• A covalent bond where both electrons come from one atom. Example: Formation of $\text{NH}_4^+$:

5. Shapes of Molecular Compounds

• Linear: $\text{BeCl}_2$, bond angle = $180^\circ$.
• Planar Triangular: $\text{BF}_3$, bond angle = $120^\circ$.
• Tetrahedral: $\text{CH}_4$, bond angle = $109.5^\circ$.
• Trigonal Pyramidal: $\text{NH}_3$, bond angle = $107^\circ$.
• Bent (Angular): $\text{H}_2\text{O}$, bond angle = $104.5^\circ$.

6. Metallic Bonding

6.1 Nature of Metallic Bonding

• Positive metal ions in a sea of delocalized electrons.
• Factors Influencing Formation:
  • Atomic Radius: Smaller size strengthens bonds.
  • Ionization Energy: Lower values favor bonding.
  • Number of Valence Electrons: More delocalized electrons increase bond strength.

6.2 Properties of Metals

• High Conductivity: Free electrons allow electrical and thermal conduction.
• Malleability and Ductility: Layers of atoms slide easily.
• Luster: Free electrons reflect light.
• High Melting Points: Strong metallic bonds.

7. Intermolecular Bonding

7.1 Types of Intermolecular Forces

1. Dipole-Dipole Interactions: Occur between polar molecules.
2. Van der Waals Forces:
   • Induced Dipole-Induced Dipole: Weak, e.g., noble gases.
   • Dipole-Induced Dipole: Between a polar molecule and a non-polar molecule.
3. Hydrogen Bonding: Strong attraction between $\text{H}$ and $\text{F, O, or N}$.

7.2 Properties of Compounds with Intermolecular Bonding

• Boiling/Melting Points:
  • Increase with stronger forces.
  • Example: $\text{H}_2\text{O}$ (high due to hydrogen bonding) vs $\text{H}_2\text{S}$ (low, weak dipole-dipole forces).
• Solubility: Polar compounds dissolve in polar solvents.

8. Comparison of Bond Types

9. Factors Influencing Bond Formation

9.1 Ionic Bonds

• Ionization Energy: Lower favors cation formation.
• Electron Affinity: High values favor anion formation.
• Electronegativity Difference: Larger differences favor ionic bonds.

9.2 Covalent Bonds

• Atomic Size: Smaller atoms form stronger bonds.
• Electronegativity: Similar values favor covalent bonding.

10. Real-World Applications

• Ionic Compounds: Salt (NaCl) for food and industry.
• Covalent Compounds: Plastics and water.
• Metallic Bonding: Conductors like copper wires.

11. Common Misconceptions

• Misconception: Covalent compounds are always non-conductors.
  • Reality: Graphite (a covalent network) conducts electricity due to delocalized electrons.
• Misconception: Ionic compounds are soluble in all solvents.
  • Reality: Insoluble in non-polar solvents.

This structured format offers clarity on chemical bonds, ensuring understanding of their formation, properties, and significance.