AQA GCSE FOUNDATION Chemistry: Types of Chemical Bonds

AQA GCSE Foundation Chemistry: Types of Chemical Bonds

This tutorial will explore the differences between ionic, covalent, and metallic bonding. These are the fundamental forces that hold atoms together to form compounds and materials.

1. Ionic Bonding

Formation: Ionic bonding occurs between metals and non-metals. Metals tend to lose electrons to achieve a stable electron configuration, becoming positively charged cations. Non-metals tend to gain electrons, becoming negatively charged anions. The electrostatic attraction between these oppositely charged ions forms the ionic bond.

Properties:

• High melting and boiling points: Due to strong electrostatic forces, a lot of energy is required to break the bonds.
• Solid at room temperature: The strong electrostatic attraction holds the ions in a rigid lattice structure.
• Conduct electricity when molten or dissolved: Free ions can move and carry charge.
• Brittle: Disruption of the lattice structure can cause ions of the same charge to come close, leading to repulsion and breakage.

Example: Sodium chloride (NaCl). Sodium (Na) loses one electron to become Na+, while chlorine (Cl) gains one electron to become Cl-. The attraction between these oppositely charged ions forms the ionic bond.

2. Covalent Bonding

Formation: Covalent bonding occurs between non-metals. Atoms share electrons to achieve a stable electron configuration, forming a shared pair of electrons between them.

Properties:

• Lower melting and boiling points than ionic compounds: The forces between molecules are weaker than the electrostatic forces in ionic compounds.
• Can be solid, liquid, or gas at room temperature: The strength of the forces between molecules determines the state of matter.
• Do not conduct electricity: No free electrons or ions to carry charge.
• Generally non-brittle: The molecules are able to slide past each other without breaking bonds.

Types of Covalent Bonding:

• Single covalent bond: One shared pair of electrons (e.g., H2).
• Double covalent bond: Two shared pairs of electrons (e.g., O2).
• Triple covalent bond: Three shared pairs of electrons (e.g., N2).

Example: Water (H2O). Each hydrogen atom shares one electron with the oxygen atom, forming two single covalent bonds.

3. Metallic Bonding

Formation: Metallic bonding occurs between metal atoms. The valence electrons of metal atoms are delocalized, forming a "sea of electrons" that surrounds the positively charged metal ions.

Properties:

• High melting and boiling points: The strong attraction between the delocalized electrons and the metal ions requires a lot of energy to overcome.
• Good conductors of heat and electricity: The delocalized electrons can move freely and carry charge.
• Malleable and ductile: The layers of metal ions can slide over each other without breaking the metallic bonds.
• Lustrous: The delocalized electrons absorb and re-emit light.

Example: Copper (Cu). Copper atoms contribute their valence electrons to the sea of electrons, allowing for excellent electrical conductivity and malleability.

Summary Table:

By understanding these fundamental types of chemical bonds, you can better understand the properties and behavior of different substances.