Metals and Non-Metals

1. Physical Properties of Metals and Non-Metals

Physical Properties of Metals

• Metallic lustre: Metals generally have a shiny surface in their pure state.
• Hardness: Most metals are hard, although their hardness varies. Lithium, sodium and potassium are soft enough to be cut with a knife.
• Malleability: The ability to be beaten into thin sheets. Gold and silver are among the most malleable metals.
• Ductility: The ability to be drawn into thin wires. Gold is the most ductile metal; approximately 2 km of wire can be drawn from 1 gram of gold.
• Conductivity: Metals are generally good conductors of heat and electricity. Silver and copper are excellent conductors, while lead and mercury are relatively poor conductors of heat.
• Melting point: Most metals have high melting points. Gallium and caesium are exceptions because of their very low melting points.
• Sonority: Metals produce a ringing sound when struck, so they are called sonorous.
• Physical state: All metals are solids at room temperature except mercury, which is liquid.

Physical Properties of Non-Metals

• Non-metals are fewer in number than metals and include carbon, sulphur, iodine, oxygen and hydrogen.
• They may exist as solids or gases. Bromine is the only liquid non-metal at room temperature.
• They are generally poor conductors of heat and electricity and are neither malleable nor ductile.
• Iodine is a lustrous non-metal.
• Diamond, an allotrope of carbon, is the hardest natural substance and has a very high melting point.
• Graphite, another allotrope of carbon, conducts electricity despite being a non-metal.

2. Chemical Properties of Metals

2.1 Reaction with Air or Oxygen

Metals react with oxygen to form metal oxides.

When copper is heated, it develops a black coating of copper(II) oxide.

Most metal oxides are basic in nature. Some soluble metal oxides react with water to form hydroxides.

Amphoteric Oxides

Variation in Reactivity with Oxygen

• Sodium and potassium react so vigorously with oxygen that they are stored under kerosene oil.
• Magnesium, aluminium, zinc and lead develop thin protective oxide layers at room temperature.
• Iron filings burn vigorously in a flame, although a solid iron piece does not burn easily.
• Silver and gold do not react with oxygen even at high temperatures.

Anodising

Anodising is an electrochemical process used to form a thick protective oxide layer on aluminium. The layer improves corrosion resistance and may also be dyed for decorative purposes.

2.2 Reaction with Water

Metals may react with water to form metal oxides or hydroxides and hydrogen gas.

• Sodium and potassium: React violently with cold water and release heat.
• Calcium: Reacts less violently with cold water and floats because hydrogen bubbles stick to its surface.
• Magnesium: Reacts with hot water and floats due to hydrogen bubbles.
• Aluminium, zinc and iron: React with steam rather than cold or hot water.
• Lead, copper, silver and gold: Do not react with water.

2.3 Reaction with Acids

Metals generally react with dilute acids to form a salt and hydrogen gas.

The approximate order of reaction with dilute hydrochloric acid is:

Copper does not react with dilute hydrochloric acid. Nitric acid generally does not release hydrogen because it is a strong oxidising agent.

Aqua Regia

Aqua regia is a freshly prepared, highly corrosive mixture of concentrated hydrochloric acid and concentrated nitric acid in a 3:1 ratio. It can dissolve noble metals such as gold and platinum.

2.4 Displacement Reactions

A more reactive metal displaces a less reactive metal from its salt solution.

For example, iron displaces copper from copper sulphate solution:

3. Reactivity Series

4. Ionic Bonding

Electronic Configuration and Reactivity

• Atoms react to achieve a stable electronic configuration, usually a complete outer shell.
• Metals generally lose 1, 2 or 3 valence electrons to form positively charged ions called cations.
• Non-metals generally gain electrons to form negatively charged ions called anions.

Formation of Sodium Chloride

• Sodium has electronic configuration (2, 8, 1) and loses one electron to form Na⁺ with configuration (2, 8).
• Chlorine has configuration (2, 8, 7) and gains one electron to form Cl⁻ with configuration (2, 8, 8).
• The oppositely charged ions attract each other through strong electrostatic forces to form sodium chloride.

Formation of Magnesium Chloride

• Magnesium has configuration (2, 8, 2) and loses two electrons to form Mg²⁺.
• Two chlorine atoms each gain one electron to form two Cl⁻ ions.
• The resulting formula is MgCl₂.

Properties of Ionic Compounds

• They are hard, crystalline and brittle solids.
• They have high melting and boiling points due to strong electrostatic forces.
• They are generally soluble in water and insoluble in non-polar solvents such as kerosene and petrol.
• They do not conduct electricity in the solid state because ions cannot move freely.
• They conduct electricity in molten form or aqueous solution because ions become mobile.

5. Occurrence and Extraction of Metals

Important Definitions

• Mineral: A naturally occurring element or compound found in the Earth's crust.
• Ore: A mineral containing enough metal to make extraction economically profitable.
• Gangue: Unwanted earthy impurities such as soil, sand and clay present in an ore.

General Extraction Flow

Ore
 ↓
Concentration of ore
 ↓
Selection of method according to reactivity
 ↓
Conversion into oxide, if required
 ↓
Reduction to metal
 ↓
Refining

5.1 Metals Low in the Reactivity Series

Low-reactivity metals can often be obtained by heating their oxides.

Mercury from cinnabar:

Copper from copper glance:

5.2 Metals in the Middle of the Reactivity Series

• Roasting: Sulphide ores are heated strongly in excess air to form oxides.
• Calcination: Carbonate ores are heated strongly in limited air to form oxides.

Highly reactive metals such as aluminium may reduce oxides of less reactive metals.

This is the thermit reaction, used for welding railway tracks and repairing cracked machine parts.

5.3 Metals High in the Reactivity Series

Potassium, sodium, calcium, magnesium and aluminium cannot be extracted by carbon reduction. They are obtained by electrolytic reduction of molten salts or oxides.

Aluminium is obtained through electrolytic reduction of molten aluminium oxide.

5.4 Electrolytic Refining

• Anode: Thick block of impure metal.
• Cathode: Thin strip of pure metal.
• Electrolyte: Solution containing a salt of the metal, such as acidified copper sulphate.
• Pure metal dissolves from the anode and deposits on the cathode.
• Insoluble impurities settle below the anode as anode mud.

6. Corrosion and Its Prevention

Examples of Corrosion

• Silver: Turns black due to formation of silver sulphide (Ag₂S).
• Copper: Develops a green coating of basic copper carbonate in moist air containing carbon dioxide.
• Iron: Develops a brown, flaky coating of rust.

Conditions Required for Rusting

Rusting requires the simultaneous presence of oxygen and water.

• Air and water together: rusting occurs.
• Boiled water covered with oil: rusting does not occur because oxygen is absent.
• Dry air with anhydrous calcium chloride: rusting does not occur because moisture is absent.

Methods of Prevention

Corrosion can be reduced by painting, oiling, greasing, galvanising, chrome plating, anodising and alloying.

Alloys

• Galvanisation: Coating iron or steel with zinc.
• Steel: Iron mixed with a small amount of carbon becomes harder and stronger.
• Stainless steel: Alloy of iron, nickel and chromium; hard and corrosion-resistant.
• 24-carat gold: Pure gold, which is too soft for most jewellery.
• 22-carat gold: 22 parts gold mixed with 2 parts copper or silver.
• Amalgam: An alloy containing mercury.
• Brass: Copper + zinc.
• Bronze: Copper + tin.
• Solder: Lead + tin; it has a low melting point and is used for joining wires.

7. Ancient Indian Metallurgy

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