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.