1. Introduction to Chemical Bonding

Atoms combine with each other to form molecules and compounds in order to attain a stable electronic configuration, usually that of the nearest noble gas (octet rule). This process is called chemical bonding.

There are three main types of chemical bonds:

• Ionic bond (Electrovalent bond)
• Covalent bond
• Coordinate bond (Dative bond)

2. Ionic Bond (Electrovalent Bond)

Definition:

An ionic bond is formed when one atom transfers electrons to another atom, resulting in the formation of positively charged and negatively charged ions. These oppositely charged ions are held together by electrostatic forces of attraction.

Formation Process:

• Atoms of metals tend to lose electrons to achieve a stable configuration, becoming positively charged ions (cations).
• Atoms of non-metals tend to gain electrons to achieve a stable configuration, becoming negatively charged ions (anions).
• The ionic bond is the strong attraction between these oppositely charged ions.

Example 1: Formation of Sodium Chloride (NaCl)

• Sodium atom (Na) has 1 electron in its outermost shell. It loses this electron to achieve the stable configuration of neon (2,8). It becomes a Na⁺ ion.
• Chlorine atom (Cl) has 7 electrons in its outermost shell. It gains 1 electron to complete its octet (2,8,8). It becomes a Cl⁻ ion.
• The Na⁺ and Cl⁻ ions attract each other to form an ionic bond, resulting in the compound sodium chloride (NaCl).

Example 2: Formation of Magnesium Oxide (MgO)

• Magnesium atom (Mg) loses 2 electrons to become Mg²⁺.
• Oxygen atom (O) gains 2 electrons to become O²⁻.
• The Mg²⁺ and O²⁻ ions form an ionic bond, resulting in magnesium oxide (MgO).

Properties of Ionic Compounds:

• High melting and boiling points due to strong electrostatic forces between ions.
• Conduct electricity in molten or aqueous solutions but not in solid state.
• Soluble in water but insoluble in organic solvents (e.g., alcohol, ether).

3. Covalent Bond

Definition:

A covalent bond is formed when two atoms share one or more pairs of electrons in order to achieve a stable configuration.

Types of Covalent Bonds:

• Single Covalent Bond: A bond formed by sharing one pair of electrons.
  • Example: Hydrogen (H₂) – H : H
• Double Covalent Bond: A bond formed by sharing two pairs of electrons.
  • Example: Oxygen (O₂) – O = O
• Triple Covalent Bond: A bond formed by sharing three pairs of electrons.
  • Example: Nitrogen (N₂) – N ≡ N

Example 1: Formation of Water (H₂O)

• Oxygen atom (O) needs 2 electrons to complete its octet.
• Each Hydrogen atom (H) needs 1 electron to complete its duplet.
• Oxygen shares one electron with each hydrogen atom, forming two single covalent bonds.

Example 2: Formation of Carbon Dioxide (CO₂)

• Carbon atom (C) needs 4 electrons to complete its octet.
• Each Oxygen atom (O) needs 2 electrons to complete its octet.
• Carbon shares two electrons with each oxygen atom, forming two double covalent bonds.

Properties of Covalent Compounds:

• Low melting and boiling points because of weak intermolecular forces.
• Generally, do not conduct electricity as they don’t have free ions.
• Insoluble in water but soluble in organic solvents.

4. Coordinate Bond (Dative Bond)

Definition:

A coordinate bond is a type of covalent bond where both electrons in the shared pair come from the same atom. One atom donates a lone pair of electrons to form the bond.

Example: Formation of Ammonium Ion (NH₄⁺)

• Ammonia (NH₃) has a lone pair of electrons on nitrogen.
• Hydrogen ion (H⁺) has no electrons.
• Nitrogen donates its lone pair to bond with H⁺, forming a coordinate bond.

Properties of Coordinate Compounds:

• Similar properties to covalent compounds.
• Can conduct electricity if they ionize in solution.

5. Comparison Between Ionic and Covalent Compounds

6. Important Concepts to Remember

• Octet Rule: Atoms bond to complete their valence shell with 8 electrons (or 2 electrons for H and He).
• Lewis Dot Structures: A way to represent atoms, ions, and molecules showing the valence electrons and how they bond.
• Electronegativity: The tendency of an atom to attract electrons in a bond. A large difference in electronegativity leads to ionic bonds, while a small difference leads to covalent bonds.

7. Practice Questions

1. Define ionic and covalent bonds with examples.
2. Why do ionic compounds have high melting points?
3. Draw the Lewis structure of water (H₂O) and explain the type of bond formed.
4. Compare the properties of ionic and covalent compounds.
5. Explain the formation of NaCl and CO₂ using the octet rule.