Types of Chemical Bonds

Atoms combine to form molecules through chemical bonds, which hold them together. The three primary types of chemical bonds are ionic bonds, covalent bonds, and metallic bonds. Here's a detailed explanation of each:

Ionic Bond

Definition

An ionic bond forms when one atom transfers one or more electrons to another atom, resulting in the formation of oppositely charged ions that attract each other.

Formation

Usually occurs between a metal and a non-metal.

• Metals lose electrons to form positively charged ions (cations).
• Non-metals gain electrons to form negatively charged ions (anions).

Example

Sodium chloride (NaCl):

• Sodium (Na) loses one electron to become Na⁺.
• Chlorine (Cl) gains one electron to become Cl⁻.
• Na⁺ and Cl⁻ are held together by electrostatic attraction.

Properties

• High melting and boiling points due to strong electrostatic forces.
• Generally form crystalline solids.
• Soluble in water and conduct electricity when dissolved (as ions are free to move).
• Brittle in nature.

Covalent Bond

Definition

A covalent bond forms when two atoms share one or more pairs of electrons to achieve a stable electronic configuration.

Formation

Typically occurs between non-metal atoms. The bond can involve:

• Single bond: Sharing one pair of electrons (e.g., H₂, H-H).
• Double bond: Sharing two pairs of electrons (e.g., O₂, O=O).
• Triple bond: Sharing three pairs of electrons (e.g., N₂, N≡N).

Example

Water (H₂O):

Each hydrogen atom shares one electron with the oxygen atom, forming two single covalent bonds.

Properties

• Low melting and boiling points (weaker intermolecular forces in molecular compounds).
• Do not conduct electricity (no free ions or electrons).
• Can form polar or non-polar molecules depending on electronegativity differences:
  • Polar covalent bond: Unequal sharing of electrons (e.g., H₂O).
  • Non-polar covalent bond: Equal sharing of electrons (e.g., N₂).

Metallic Bond

Definition

A metallic bond is the force of attraction between positively charged metal ions and a "sea" of delocalized free electrons.

Formation

Found in pure metals and alloys. Metal atoms lose their outer electrons, which become delocalized and move freely throughout the lattice of positively charged metal ions.

Example

In a piece of copper (Cu):

• Copper atoms release their valence electrons to form a sea of electrons, which holds the metal ions together.

Properties

• High electrical and thermal conductivity due to free-moving electrons.
• Malleable (can be hammered into sheets) and ductile (can be drawn into wires).
• Lustrous (shiny) due to interaction of free electrons with light.
• High melting and boiling points, depending on the metal.

Summary Table