Metals and Non-metals

Elements are broadly classified as metals and non-metals on the basis of their properties. Most of the metallic properties can be observed by simple physical tests.

Physical properties of metals:

• Lustre: Metals in their pure state have a shining surface, e.g. freshly cut sodium or silver.
• Malleability: Metals can be beaten into thin sheets. Gold and silver are the most malleable; aluminium foil is used to wrap food.
• Ductility: Metals can be drawn into thin wires, e.g. copper and aluminium wires.
• Conduction: Metals are good conductors of heat and electricity. Silver and copper are the best conductors; lead and mercury are comparatively poor.
• Sonority: Metals produce a ringing sound when struck, so they are sonorous (used in bells).
• Hardness, melting point, state: Most metals are hard solids with high melting points. Exceptions: sodium and potassium are soft (cut with a knife); mercury is a liquid metal.

Physical properties of non-metals: Non-metals exist as solids (carbon, sulphur), liquids (bromine) or gases (oxygen, chlorine). They are generally non-lustrous, brittle, non-sonorous and poor conductors of heat and electricity.

Important exceptions to remember: Iodine is a non-metal but is lustrous; graphite (a form of carbon) is a non-metal that conducts electricity; diamond is the hardest natural substance and has a very high melting point though it is a non-metal.

The chemical behaviour of a metal depends on its reactivity. Comparing four types of reactions lets us rank metals.

1. Reaction with oxygen (forming basic oxides): Metals combine with oxygen (many burn on heating) to form metal oxides which are generally basic. Note that some metals like copper do not burn, but their hot surface gets coated with a black oxide.

• 2Cu + O₂ → 2CuO (copper turns black)
• 4Al + 3O₂ → 2Al₂O₃

Some metal oxides such as Al₂O₃ and ZnO are amphoteric — they react with both acids and bases, e.g. Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O and Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O.

2. Reaction with water: Reactivity decreases down the activity series. Sodium and potassium react violently even with cold water; calcium reacts less vigorously (it floats because the hydrogen gas evolved sticks to its surface). Magnesium reacts with hot water; iron reacts only with steam; copper, silver and gold do not react at all.

• 2Na + 2H₂O → 2NaOH + H₂ + heat
• Ca + 2H₂O → Ca(OH)₂ + H₂
• 3Fe + 4H₂O (steam) → Fe₃O₄ + 4H₂

3. Reaction with dilute acids: Metals above hydrogen displace hydrogen gas from dilute acids; copper, silver and gold (below hydrogen) do not.

• Fe + 2HCl → FeCl₂ + H₂
• Zn + 2HCl → ZnCl₂ + H₂

4. Reaction with salt solutions (displacement): A more reactive metal displaces a less reactive metal from its salt solution. When an iron nail is dipped in blue copper sulphate solution, the solution turns pale green and a brown deposit of copper forms: Fe + CuSO₄ → FeSO₄ + Cu.

These observations together give the reactivity (activity) series, listing metals in decreasing order of reactivity.

Metals are electropositive — they tend to lose electrons to attain a stable (noble-gas) electron configuration. Non-metals are electronegative — they gain those electrons. The transfer creates positive and negative ions that attract each other strongly.

Example — formation of sodium chloride (NaCl):

• Sodium (2,8,1) loses 1 electron: Na → Na⁺ + e^- (now 2,8, like neon).
• Chlorine (2,8,7) gains that electron: Cl + e^- → Cl^- (now 2,8,8, like argon).
• The oppositely charged ions attract: Na⁺ + Cl^- → NaCl. Overall: 2Na + Cl₂ → 2NaCl.

Compounds formed by such transfer of electrons are called ionic (electrovalent) compounds. Another example is magnesium chloride, where Mg loses 2 electrons (Mg → Mg²⁺ + 2e^-) and each of two chlorine atoms gains one, giving MgCl₂.

Properties of ionic compounds:

• Physical state: Solid and generally hard due to strong force of attraction between ions.
• Melting and boiling points: Very high, because a large amount of energy is needed to break the strong inter-ionic bonds.
• Solubility: Generally soluble in water but insoluble in solvents like kerosene and petrol.
• Conduction of electricity: They conduct electricity in the molten state or in aqueous solution (ions become free to move) but not in the solid state (ions are fixed in the lattice).

Metals occur in nature as free elements or, more often, as ores (minerals from which a metal can be profitably extracted). The extraction route depends on reactivity.

Step 1 — Enrichment/Concentration: The ore is first freed from earthy impurities (gangue).

Step 2 — Conversion of concentrated ore to oxide: Oxides are easier to reduce than sulphides or carbonates, so ores are converted to oxides.

• Roasting: A sulphide ore is heated strongly in the presence of excess air to form the oxide, e.g. 2ZnS + 3O₂ → 2ZnO + 2SO₂.
• Calcination: A carbonate ore is heated strongly in the absence or limited supply of air to form the oxide, e.g. ZnCO₃ → ZnO + CO₂.

Step 3 — Reduction of oxide to metal:

• Metals of medium reactivity (Zn, Fe, Pb) are reduced by heating their oxides with carbon: ZnO + C → Zn + CO.
• Highly reactive metals like aluminium can reduce the oxides of less reactive metals in the thermit reaction (used to join railway tracks): Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + heat.
• Highly reactive metals (K, Na, Ca, Mg, Al) are obtained by electrolytic reduction of their molten chlorides/oxides, e.g. at the cathode: Na⁺ + e^- → Na; at the anode: 2Cl^- → Cl₂ + 2e^-.
• Low reactivity metals like mercury and copper are obtained just by heating their ores; e.g. cinnabar: 2HgS + 3O₂ → 2HgO + 2SO₂, then 2HgO → 2Hg + O₂.

Step 4 — Refining (electrolytic refining): The impure metal is made the anode and a strip of pure metal the cathode, dipped in a solution of a salt of the metal. On passing current, pure metal dissolves from the anode and deposits on the cathode; impurities settle below the anode as anode mud.

Corrosion is the gradual deterioration of a metal by the action of air, moisture and chemicals in its surroundings.

Common examples:

• Iron forms a reddish-brown flaky coating of rust (hydrated iron(III) oxide, Fe₂O₃·xH₂O). Both air (oxygen) and water/moisture are essential for rusting.
• Silver develops a black coating of silver sulphide (Ag₂S) on reacting with H₂S in air.
• Copper develops a green coating of basic copper carbonate (CuCO₃·Cu(OH)₂) in moist air.

Corrosion causes huge economic loss because it damages car bodies, bridges, ships, and iron railings, and weakens structures.

Methods of prevention:

• Painting, oiling, greasing: forms a barrier that keeps out air and moisture.
• Galvanisation: coating iron/steel with a thin layer of zinc. The zinc protects the iron even if the coating is scratched, because zinc is more reactive.
• Electroplating / chrome plating / tin plating: coating with a less reactive metal.
• Anodising: forming a thick protective oxide layer on aluminium.
• Alloying: e.g. making stainless steel (iron with chromium and nickel), which does not rust.

Alloys in general (a homogeneous mixture of a metal with other metals or non-metals) are made to improve properties: brass (Cu + Zn) and bronze (Cu + Sn) are harder and more useful than pure copper; solder (Pb + Sn) has a low melting point; adding carbon to iron gives hard steel.