Chemical Bonding

Introduction: The Drive for Stability

Atoms form chemical bonds to achieve a more stable, lower-energy state. This stability is typically achieved by obtaining a full outer electron shell, resembling the electron configuration of the noble gases (Group 18). This principle, often called the octet rule, is the fundamental reason chemical reactions occur. There are three main ways atoms can achieve this: by transferring, sharing, or delocalising electrons, leading to ionic, covalent, and metallic bonding, respectively.

1. Ionic Bonding: The Transfer of Electrons

Ionic bonding occurs between a metal and a non-metal. Metals, having few valence electrons, tend to lose them to form positively charged ions (cations). Non-metals, with nearly full outer shells, tend to gain electrons to form negatively charged ions (anions). The ionic bond is the strong electrostatic force of attraction between these oppositely charged ions.

Example: Sodium Chloride (NaCl)

1. A sodium atom (Na) has the electron configuration 2,8,1. It loses its single valence electron to achieve a stable configuration of 2,8. It becomes a sodium ion, Na⁺.
2. A chlorine atom (Cl) has the configuration 2,8,7. It gains one electron to achieve a stable configuration of 2,8,8. It becomes a chloride ion, Cl⁻.
3. The resulting Na⁺ and Cl⁻ ions are strongly attracted to each other.

Structure and Properties of Ionic Compounds:

• Giant Ionic Lattice: Ions are not found in discrete pairs but are arranged in a vast, three-dimensional, repeating pattern called a **giant ionic lattice**. Each cation is surrounded by anions, and each anion by cations.
• High Melting and Boiling Points: A large amount of thermal energy is required to overcome the strong electrostatic forces holding the lattice together. For example, NaCl melts at 801 °C.
• Electrical Conductivity:
• Solid: Do not conduct electricity because the ions are held in fixed positions and are not free to move.
• Molten or Aqueous (dissolved): Conduct electricity because the ions are mobile and can move to carry an electric current. This is a crucial distinction often tested in exams.
• Solubility: Many are soluble in polar solvents like water. The polar water molecules can surround the ions and weaken the electrostatic forces, causing the lattice to break down.

Drawing Dot-and-Cross Diagrams for Ionic Bonding:

• Show the valence electrons of the metal and non-metal atoms before bonding.
• Use arrows to show the transfer of electrons from the metal to the non-metal.
• Draw the resulting ions in square brackets, with the new full outer shell of electrons and the charge shown outside the bracket.

2. Covalent Bonding: The Sharing of Electrons

Covalent bonding occurs between non-metal atoms. To achieve a stable outer shell, these atoms share one or more pairs of electrons. A shared pair of electrons is attracted to the positive nuclei of both atoms, holding them together in a strong covalent bond.

• Single bond: One shared pair of electrons (e.g., H₂, Cl₂, CH₄).
• Double bond: Two shared pairs of electrons (e.g., O₂, CO₂).
• Triple bond: Three shared pairs of electrons (e.g., N₂).

Structures and Properties of Covalent Substances:

A) Simple Molecular Structures

• Structure: Consist of discrete, small molecules (e.g., water H₂O, methane CH₄, carbon dioxide CO₂). The atoms within a molecule are joined by strong covalent bonds, but the forces **between** the molecules (**intermolecular forces**) are very weak.
• Properties:
• Low Melting and Boiling Points: Only a small amount of energy is needed to overcome the weak intermolecular forces. The strong covalent bonds *within* the molecules do not break.
• Do not conduct electricity: There are no free-moving charged particles (no ions or delocalised electrons).

B) Giant Covalent Structures (Macromolecules)

• Structure: A vast number of atoms are joined by a network of strong covalent bonds, forming a single giant molecule or lattice.
• Properties:
• Very High Melting and Boiling Points: A large amount of energy is needed to break the numerous strong covalent bonds throughout the structure.
• Variable Conductivity: Most do not conduct electricity. An important exception is graphite.

Allotropes of Carbon:

• Diamond: Each carbon atom is covalently bonded to four others in a rigid **tetrahedral** arrangement. This makes diamond extremely hard (used in cutting tools) and gives it a very high melting point. It does not conduct electricity as all electrons are localised in bonds.
• Graphite: Each carbon atom is bonded to three others, forming **hexagonal layers**. The layers are held by weak intermolecular forces, allowing them to slide over each other, making graphite soft and useful as a lubricant. Each carbon atom has one **delocalised electron**, which is free to move along the layers, allowing graphite to conduct electricity.

3. Metallic Bonding: A Sea of Electrons

Metallic bonding is found in metals and alloys. It consists of a regular lattice of positive metal ions surrounded by a 'sea' of delocalised electrons. These delocalised electrons are the valence electrons that are no longer associated with any single atom.

The metallic bond is the electrostatic attraction between the positive ions and the delocalised electrons.

Properties of Metals:

• Good Electrical and Thermal Conductors: The delocalised electrons are mobile and free to move throughout the structure, carrying charge (electricity) or kinetic energy (heat).
• Malleable and Ductile: When a force is applied, the layers of positive ions can slide over one another without breaking the metallic bond. The delocalised electrons adjust to the new arrangement, maintaining the attraction. This allows metals to be hammered into shape (**malleable**) or drawn into wires (**ductile**).
• High Melting and Boiling Points: The electrostatic forces between the positive ions and the delocalised electron sea are very strong and require significant energy to overcome.

Common Exam Trap: When explaining the low boiling points of simple molecules like methane, students often incorrectly state that 'weak covalent bonds are broken'. The correct answer is that weak **intermolecular forces** between the molecules are overcome.