Metals and Non-metals

1. Physical Properties

Metals

• Metallic Lustre: In their pure state, metals have a shining surface.
• Hardness: Metals are generally hard, though hardness varies (e.g., Iron vs. Sodium).
• Malleability: Metals can be beaten into thin sheets. Gold and silver are the most malleable.
• Ductility: The ability to be drawn into thin wires. Gold is the most ductile metal.
• Conductivity: Metals are good conductors of heat and electricity. Silver and copper are the best thermal conductors; lead and mercury are poor.
• Sonority: Metals produce a ringing sound when striking a hard surface.

Non-metals

• State: Either solids or gases at room temperature (Exception: Bromine is a liquid).
• Properties: generally not lustrous, not sonorous, and are poor conductors of heat and electricity.
• Examples include carbon, sulphur, iodine, oxygen, and hydrogen.

2. Chemical Properties of Metals

Reaction with Air (Burnt in Oxygen)

• Almost all metals combine with oxygen to form Metal Oxides.
• Nature of Oxides: Metal oxides are generally basic. However, aluminium oxide and zinc oxide are Amphoteric Oxides (react with both acids and bases to produce salt and water).
• Reactivity varies:
  • Sodium and Potassium react so vigorously they catch fire in open air (stored in kerosene).
  • Magnesium, Aluminium, Zinc, and Lead form a protective oxide layer that prevents further oxidation.
  • Silver and Gold do not react with oxygen even at high temperatures.
• Anodising: A process to form a thick oxide layer on aluminium to improve corrosion resistance.

Reaction with Water

• Metal + Water → Metal Oxide + Hydrogen (Metal oxide may further form Metal Hydroxide).
• Cold Water: Sodium and Potassium react violently (exothermic, hydrogen catches fire). Calcium reacts less violently and floats.
• Hot Water: Magnesium reacts with hot water and floats.
• Steam: Aluminium, Iron, and Zinc do not react with hot/cold water but react with steam to form oxides.
• No Reaction: Lead, Copper, Silver, and Gold do not react with water at all.

Reaction with Acids

• Metal + Dilute Acid → Salt + Hydrogen.
• Nitric Acid (HNO₃): Hydrogen is generally not evolved because HNO₃ is a strong oxidising agent (oxidises H₂ to water). Exception: Magnesium and Manganese react with very dilute HNO₃ to evolve hydrogen.
• Aqua Regia: A mixture of conc. HCl and conc. HNO₃ (3:1) capable of dissolving gold and platinum.

Reaction with Other Metal Salts (Displacement)

• Reactive metals can displace less reactive metals from their compounds in solution or molten form.
• Reaction: Metal A + Salt Solution of B → Salt Solution of A + Metal B (if A is more reactive than B).

The Reactivity Series

Metals arranged in decreasing order of reactivity:

3. How Metals and Non-metals React

• Electronic Configuration: Elements react to attain a completely filled valence shell (stable octet).
• Ionic Compounds: Formed by the transfer of electrons from a metal (forming a positive cation) to a non-metal (forming a negative anion).
• Properties of Ionic Compounds:
  • Physical Nature: Solid and hard due to strong inter-ionic attraction; generally brittle.
  • Melting/Boiling Points: High, as considerable energy is needed to break the strong attraction.
  • Solubility: Soluble in water; insoluble in solvents like kerosene and petrol.
  • Conductivity: Conduct electricity in molten state or aqueous solution (ions move freely), but not in solid state.

4. Occurrence of Metals

• Minerals: Elements/compounds occurring naturally in the earth’s crust.
• Ores: Minerals from which metal can be profitably extracted.
• Gangue: Impurities (soil, sand) present in mined ores.

Extraction Based on Reactivity

• Metals low in activity series (e.g., Hg, Cu):
  • Oxides can be reduced to metals by heating alone.
  • Sulphide ores (like Cinnabar, HgS) are heated in air.
• Metals in the middle of activity series (e.g., Zn, Fe, Pb):
  • Usually found as sulphides or carbonates.
  • Roasting: Heating sulphide ores in excess air to convert to oxide.
  • Calcination: Heating carbonate ores in limited air to convert to oxide.
  • Reduction: Metal oxides are reduced using carbon (coke) or highly reactive metals (displacement).
  • Thermit Reaction: Reaction of Iron(III) oxide with Aluminium is highly exothermic and used to join railway tracks.
• Metals at the top of activity series (e.g., Na, Mg, Al, Ca):
  • Cannot be reduced by carbon due to high affinity for oxygen.
  • Obtained by Electrolytic Reduction of their molten chlorides or oxides. Metal deposits at the cathode.

Refining of Metals

• Electrolytic Refining: Used for Cu, Zn, Sn, Ni, Ag, Au.
• Anode: Impure metal. Cathode: Strip of pure metal. Electrolyte: Metal salt solution.
• Pure metal deposits on the cathode; insoluble impurities settle as Anode Mud.

5. Corrosion

• Silver: Turns black due to formation of silver sulphide (reaction with sulphur in air).
• Copper: Turns green due to formation of copper carbonate (reaction with moist CO₂).
• Iron: Forms brown flaky rust (hydrated iron oxide) when exposed to moist air. Both air (oxygen) and moisture (water) are required for rusting.

Prevention of Corrosion

• Methods include painting, oiling, greasing, chrome plating, anodising, or making alloys.
• Galvanisation: Coating steel and iron with a thin layer of Zinc.
• Alloying: A homogeneous mixture of two or more metals, or a metal and a non-metal.
  • Improves properties (e.g., pure iron is soft; Iron + Carbon is hard).
  • Stainless Steel: Iron + Nickel + Chromium (hard, does not rust).
  • Brass: Copper + Zinc.
  • Bronze: Copper + Tin.
  • Solder: Lead + Tin (low melting point, used for welding).
  • Amalgam: An alloy where one metal is Mercury.