Lesson 5: Summary

Chemical bonding refers to the attractive forces that hold atoms together in molecules, ions, or crystals. There are three main types of chemical bonds: covalent bonds, ionic bonds, and metallic bonds. the following are detailed summary of each type:

Covalent Bonds

• Covalent bonds form when atoms share one or more pairs of electrons to achieve a stable electron configuration.
• Electron Sharing: Each atom contributes one or more electrons to the shared pair, allowing both atoms to achieve a noble gas configuration.

Types of Covalent Bonds

• Single Covalent Bond: Involves the sharing of one pair of electrons (e.g., H₂).
• Double Covalent Bond: Involves the sharing of two pairs of electrons (e.g., O₂).
• Triple Covalent Bond: Involves the sharing of three pairs of electrons (e.g., N₂).

Properties:

• Covalent compounds can be solids, liquids, or gases.
• They generally have lower melting and boiling points compared to ionic compounds.
• They may be polar or non-polar depending on the electronegativity difference between bonded atoms.
• Examples: H₂O (water), CH₄ (methane), CO₂ (carbon dioxide).

Ionic Bonds

• Ionic bonds form between ions of opposite charges (cation and anion) through electrostatic attraction.
• Electron Transfer: One atom (typically a metal) loses electrons to become a cation, while another atom (typically a non-metal) gains those electrons to become an anion.
• Formation: The resulting ions attract each other to form a stable crystal lattice structure.

Properties:

• Ionic compounds are typically crystalline solids at room temperature.
• They have high melting and boiling points due to strong electrostatic attractions.
• They are brittle and often dissolve in water, where they conduct electricity.
• Examples: NaCl (sodium chloride), MgCl₂ (magnesium chloride), CaCO₃ (calcium carbonate).

Metallic Bonds

• Metallic bonds occur between atoms within metals and alloys.
• Electron Delocalization: Metal atoms release their valence electrons into a “sea” of delocalized electrons, which move freely throughout the metal lattice.

Properties:

• Metals are good conductors of electricity and heat due to the mobility of electrons.
• They are malleable and ductile, meaning they can be hammered into thin sheets or drawn into wires.
• Metals have a characteristic luster due to the reflection of light by delocalized electrons.
• Examples: Copper (Cu), Iron (Fe), Aluminum (Al).

Comparison:

• Bond Strength: Ionic bonds are typically stronger than covalent bonds, and metallic bonds are stronger than both.
• Nature of Attraction: Ionic bonds involve electrostatic attraction, covalent bonds involve electron sharing, and metallic bonds involve electron delocalization.
• Physical Properties: Each type of bond contributes distinct physical properties to substances, influencing their behavior in different environments.