Revision Notes: Metals and Non-metals

Exploring the distinct properties and behaviors of elements

1. Introduction and Classification

Elements are classified as metals or non-metals based on their properties.

2. Physical Properties

2.1 Metals

  • Metallic Lustre: Metals have a shining surface in their pure state.
  • Hardness: Generally hard, though hardness varies from metal to metal.
    • Exception: Alkali metals (lithium, sodium, potassium) are so soft they can be cut with a knife.
  • Malleability: Can be beaten into thin sheets. Gold and silver are the most malleable metals.
  • Ductility: Can be drawn into thin wires. Gold is the most ductile metal (1 gram can make a 2 km wire).
  • Conductivity (Heat): Good conductors of heat and have high melting points.
    • Best conductors of heat: Silver and copper.
    • Poor conductors of heat: Lead and mercury.
    • Exceptions: Gallium and caesium have very low melting points and will melt on your palm. All metals except mercury exist as solids at room temperature.
  • Conductivity (Electricity): Good conductors of electricity. Electric wires are coated with PVC or rubber-like material because these are insulators.
  • Sonorousness: Produce a sound when striking a hard surface. School bells are made of metals for this reason.

2.2 Non-metals

  • General State: Mostly solids or gases at room temperature, except bromine which is a liquid.
  • Physical Properties: Generally have properties opposite to metals.
  • Lustre: Generally non-lustrous. Exception: Iodine is a non-metal but is lustrous.
  • Hardness: Not generally hard. Exception: Diamond (an allotrope of carbon) is the hardest natural substance known.
  • Malleability/Ductility: Neither malleable nor ductile.
  • Conductivity (Heat/Electricity): Bad conductors of heat and electricity. Exception: Graphite (an allotrope of carbon) is a conductor of electricity.

Allotropes: Carbon can exist in different forms called allotropes (e.g., diamond, graphite).

3. Chemical Properties of Metals

Elements can be more clearly classified as metals and non-metals based on their chemical properties.

Noble gases have completely filled valence shells and show little chemical activity. The reactivity of elements is a tendency to attain a completely filled valence shell.

3.1 Reaction with Air (Oxygen)

Almost all metals combine with oxygen to form metal oxides.

Metal + Oxygen → Metal oxide

(e.g., Copper forms black copper(II) oxide; Aluminium forms aluminium oxide).

Nature of Metal Oxides:

  • Most metal oxides are basic in nature.
  • Some metal oxides (e.g., aluminium oxide, zinc oxide) show both acidic and basic behaviour and are known as amphoteric oxides. They react with both acids and bases to produce salts and water.

Solubility of Metal Oxides:

Most are insoluble in water, but some dissolve to form alkalis (e.g., sodium oxide, potassium oxide).

Reactivity towards Oxygen:

Different metals show different reactivities.

  • Highly Reactive: Potassium and sodium react so vigorously that they catch fire in the open; kept immersed in kerosene oil to prevent accidental fires.
  • Moderately Reactive: Magnesium, aluminium, zinc, lead form a thin protective oxide layer at ordinary temperature, preventing further oxidation. Iron does not burn on heating but iron filings burn vigorously.
  • Less Reactive: Copper does not burn but gets a black layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.

Anodising: Process of forming a thick oxide layer of aluminium to improve corrosion resistance. Aluminium is made the anode and electrolysed in dilute sulphuric acid, and oxygen gas forms a thicker protective oxide layer.

3.2 Reaction with Water

Metal + Water → Metal oxide + Hydrogen

Metal oxide + Water → Metal hydroxide (if soluble)

Reactivity towards Water:

Not all metals react with water.

  • React with cold water: Potassium and sodium react violently, the reaction is exothermic, and evolved hydrogen catches fire. Calcium reacts less violently, and hydrogen bubbles make it float.
  • React with hot water: Magnesium reacts with hot water to form magnesium hydroxide and hydrogen; it also floats.
  • React with steam: Aluminium, iron, and zinc do not react with cold or hot water but react with steam to form metal oxides and hydrogen.
  • Do not react with water: Lead, copper, silver, and gold.

3.3 Reaction with Acids

Metal + Dilute acid → Salt + Hydrogen gas

Reactivity towards Dilute Acids:

Not all metals react in the same manner.

  • Decreasing Reactivity: Magnesium > Aluminium > Zinc > Iron.
  • No Reaction: Copper does not react with dilute HCl.

Exception (Nitric Acid): Hydrogen gas is generally not evolved when metals react with nitric acid (HNO₃) because it is a strong oxidising agent, converting H₂ to water and getting reduced to nitrogen oxides (N₂O, NO, NO₂). Exceptions: Magnesium (Mg) and Manganese (Mn) react with very dilute HNO₃ to evolve H₂ gas.

Aqua Regia: A freshly prepared mixture of concentrated hydrochloric acid and concentrated nitric acid (3:1 ratio). It can dissolve gold and platinum, which neither acid can do alone. It is highly corrosive and fuming.

3.4 Reaction with Solutions of other Metal Salts (Displacement Reactions)

Reactive metals can displace less reactive metals from their compounds in solution or molten form.

Metal A + Salt solution of B → Salt solution of A + Metal B
(if Metal A is more reactive than Metal B)

Displacement reactions provide better evidence for reactivity than reactions with oxygen, water, or acids, especially for metals that do not react with those reagents.

3.5 The Reactivity Series (Activity Series)

A list of metals arranged in the order of their decreasing activities.

  • Most Reactive (Top): Potassium (K), Sodium (Na), Calcium (Ca), Magnesium (Mg), Aluminium (Al).
  • Medium Reactive: Zinc (Zn), Iron (Fe), Lead (Pb).
  • Hydrogen [H] is included as a reference point; metals above hydrogen can displace it from dilute acids.
  • Least Reactive (Bottom): Copper (Cu), Mercury (Hg), Silver (Ag), Gold (Au).

4. How Metals and Non-metals React

Elements react to attain a completely filled valence shell (like noble gases).

Formation of Ionic Compounds (Electrovalent Compounds):

Formed by the transfer of electrons from a metal to a non-metal.

  • Metals lose electrons to form positively charged cations (e.g., Na → Na⁺).
  • Non-metals gain electrons to form negatively charged anions (e.g., Cl → Cl⁻).
  • Oppositely charged ions attract each other via strong electrostatic forces of attraction (e.g., NaCl, MgCl₂).

Ionic compounds exist as aggregates of oppositely charged ions, not as molecules.

4.1 Properties of Ionic Compounds

  • Physical Nature: Solids, somewhat hard due to strong inter-ionic forces. Generally brittle and break into pieces under pressure.
  • Melting and Boiling Points: Have high melting and boiling points. This is because considerable energy is needed to break the strong inter-ionic attraction.
  • Solubility: Generally soluble in water and insoluble in solvents like kerosene, petrol.
  • Conduction of Electricity:
    • Solid state: Do not conduct electricity because ions are in a rigid structure and cannot move.
    • Molten state or in solution: Conduct electricity because the electrostatic forces are overcome by heat (molten) or water (solution), allowing ions to move freely.

5. Occurrence and Extraction of Metals (Metallurgy)

5.1 Key Definitions

  • Minerals: Elements or compounds occurring naturally in the earth's crust.
  • Ores: Minerals from which a metal can be profitably extracted due to a high percentage of that metal.
  • Gangue: Impurities (soil, sand, etc.) contaminating ores.

5.2 Occurrence of Metals

  • Earth's crust is the major source; seawater also contains soluble salts.
  • Free state (Native state): Metals low in the activity series (least reactive) like gold, silver, platinum, and sometimes copper.
  • Combined state: Metals high and middle in the activity series (more reactive) are found as compounds (sulphides, oxides, carbonates).

Oxides are common ores because oxygen is very reactive and abundant.

5.3 Extraction of Metals

Different techniques are used based on reactivity. Impurities (gangue) must be removed first.

5.3.1 Extracting Metals Low in the Activity Series

Very unreactive. Their oxides can be reduced to metals by heating alone.

Example: Cinnabar (HgS) heated in air forms mercuric oxide (HgO), which then reduces to mercury upon further heating. Copper sulphide (Cu₂S) also yields copper upon heating.

5.3.2 Extracting Metals in the Middle of the Activity Series

Moderately reactive (iron, zinc, lead, copper). Usually found as sulphides or carbonates.

Conversion to Oxides:

Easier to obtain metal from oxide.

  • Roasting: Sulphide ores converted to oxides by strong heating in excess air. (e.g., 2ZnS + 3O₂ → 2ZnO + 2SO₂).
  • Calcination: Carbonate ores converted to oxides by strong heating in limited air. (e.g., ZnCO₃ → ZnO + CO₂).
Reduction of Metal Oxides:

Reduced to metals using suitable reducing agents.

  • Using Carbon: Common method (e.g., ZnO + C → Zn + CO).
  • Using more reactive metals (Displacement): Highly reactive metals (Na, Ca, Al) act as reducing agents, displacing less reactive metals from their compounds.

These reactions are highly exothermic, producing metals in molten state.

Thermit Reaction: Reaction of iron(III) oxide (Fe₂O₃) with aluminium, used to join railway tracks and cracked machine parts (e.g., Fe₂O₃ + 2Al → 2Fe + Al₂O₃ + Heat).

5.3.3 Extracting Metals towards the Top of the Activity Series

Very reactive (Na, Mg, Ca, Al). Cannot be reduced by carbon because they have a higher affinity for oxygen than carbon.

Electrolytic Reduction: Obtained by electrolysis of their molten chlorides or oxides.

  • At cathode (negative electrode): Metal ions gain electrons and are deposited as pure metal (e.g., Na⁺ + e⁻ → Na).
  • At anode (positive electrode): Non-metal ions lose electrons and are liberated (e.g., 2Cl⁻ → Cl₂ + 2e⁻).

Aluminium is obtained by electrolytic reduction of aluminium oxide.

5.4 Refining of Metals

Metals obtained from extraction processes are not pure and contain impurities.

Electrolytic Refining: Most widely used method for purifying many metals (copper, zinc, tin, nickel, silver, gold).

  • Anode: Impure metal.
  • Cathode: Thin strip of pure metal.
  • Electrolyte: A solution of the metal salt.

Process: Pure metal from the anode dissolves into the electrolyte, and an equivalent amount of pure metal from the electrolyte is deposited on the cathode.

Anode Mud: Insoluble impurities settle at the bottom of the anode.

6. Corrosion and Prevention

6.1 Corrosion

The surface of some metals (e.g., iron, silver, copper) gets corroded when exposed to moist air for a long time.

  • Silver: Becomes black due to reaction with sulphur in air, forming silver sulphide.
  • Copper: Loses shiny brown surface and gains a green coat (basic copper carbonate) from reacting with moist carbon dioxide in air.
  • Iron (Rusting): Acquires a coating of a brown flaky substance called rust when exposed to moist air (both air and water are needed).

6.2 Prevention of Corrosion (Rusting of Iron)

Rusting of iron can be prevented by:

  • Painting
  • Oiling
  • Greasing
  • Galvanisation: Coating steel and iron with a thin layer of zinc. The galvanised article is protected even if the zinc coating is broken.
  • Chrome plating
  • Anodising (for aluminium)
  • Making Alloys

6.3 Alloys

Definition: A homogeneous mixture of two or more metals, or a metal and a non-metal.

Preparation: Primary metal is melted, then other elements are dissolved in definite proportions, and the mixture is cooled.

Purpose: A very good method for improving the properties of a metal and getting desired properties. Pure iron is soft and stretches when hot, but alloying changes its properties.

Properties of Alloys vs. Pure Metals:

  • Electrical conductivity and melting point of an alloy are less than that of pure metals.

Examples:

  • Iron with Carbon: Small amount (0.05%) makes iron hard and strong.
  • Stainless Steel: Iron mixed with nickel and chromium; hard and does not rust.
  • Brass: Alloy of copper and zinc (Cu + Zn).
  • Bronze: Alloy of copper and tin (Cu + Sn).
  • Solder: Alloy of lead and tin (Pb + Sn); has a low melting point, used for welding electrical wires.
  • Gold: Pure gold (24 carat) is very soft and not suitable for jewellery. It is alloyed with silver or copper to make it hard (e.g., 22 carat gold is 22 parts pure gold + 2 parts copper/silver).
  • Amalgam: An alloy where one of the metals is mercury.

Iron Pillar at Delhi: An ancient example of Indian metallurgy that prevented iron from rusting for over 1600 years.