Metallic Bonding and Properties of Metals

Characteristics of Metals

• Metals are located on the left side of the periodic table.
• Tend to lose valence electrons to form cations (positive ions).
• Form lattice structures with delocalized valence electrons.
• Metallic Bonding: Delocalized electrons are not attached to any single cation, allowing them to move freely within the lattice.

Physical Properties of Metals

• Electrical and Thermal Conductivity: Excellent due to mobile electrons.
• High Melting and Boiling Points: Result from strong metallic bonds.
• Ductility and Malleability: Space between cations allows movement without disrupting bonding.
• Luster: The interaction of delocalized electrons with light.

Factors Affecting the Strength of Metallic Bonds

• Number of valence electrons.
• Size of metallic cation and packing within lattice.
• Example: Calcium (Ca) vs. Potassium (K)
  • Calcium has stronger metallic bonds than potassium, reflected in higher melting point (Ca: 842°C vs. K: 63.5°C).

Group Trends

• Metals with fewer energy levels exhibit stronger metallic bonds due to tighter packing and stronger forces.
• Group 1 Alkaline Metals: Smaller metals like lithium form the strongest metallic bonds in the group.

Historical Use and Alloy Formation

• Metals have been used since the Bronze Age, over 5,000 years ago.
• Alloys: Mixtures of metallic elements to enhance properties.
  • Bronze: Copper + Tin, enhanced hardness, durability, corrosion resistance.
  • Integration of elements into metallic structures, either interstitially or substitutionally.
  • Examples:
    • Steel: Carbon atoms between iron cations.
    • Nichrome: Chromium substitutes for nickel.

Summary

• Metallic bonding allows for conductivity, malleability, and workability.
• Alloys enhance metal characteristics, impacting technological and industrial advancements.
  • Examples of alloys include Bronze, Brass, Steel, Pukor, Nichrome, Sterling Silver, and varieties of Gold (White and Rose).