Comprehensive Study Notes on Metallurgy: Metals, Non-metals, and Extraction Processes

Definitions and Classification of Elements

• Metals: Elements with a tendency to lose electrons and form positively charged ions. They are found in almost all groups of the Periodic Table.
  • The first three groups consist solely of metals, with the exceptions of hydrogen and boron.
  • Groups IA and IIA contain the most active metals.
  • Group IA elements (excluding hydrogen) are classified as alkali metals.
  • Group IIA elements are classified as alkaline earth metals.
• Non-metals: Elements with a tendency to accept electrons and form negatively charged ions. They are found in Groups IIIA, IVA, VA, VIA, VIIA, and Zero.
  • Group VIIA is known as the halogen family.
  • Group Zero elements are known as noble gases.

Physical Properties of Metals

• Physical State: All metals are solids at room temperature, except for mercury and gallium.
• Lustre: They possess a characteristic brilliant lustre known as metallic lustre.
• Hardness: Metals are generally hard.
• Density: Metals typically have high density, though alkali metals are an exception.
• Melting and Boiling Points: They generally possess high melting points (except for sodium and potassium) and high boiling points (except for mercury and gallium).
• Conductivity: They are good conductors of both heat and electricity, with the exception of tungsten.
• Malleability: Metals are malleable, meaning they can be beaten into thin sheets. Zinc and mercury are exceptions.
• Ductility: They are usually ductile, meaning they can be drawn into wires. Exceptions include zinc, mercury, and gallium.
• Brittleness: Metals are generally not brittle, though zinc is an exception.
• Sonorousness: Metals are sonorous, meaning they produce a musical note when struck. Sodium and potassium are exceptions.
• Alloy Formation: They form alloys and amalgams.
• Solubility: They are generally insoluble in water or other solvents.

Chemical Properties of Metals

• Valence Shells: Metals generally have 1, 2, or 3 electrons in their valence shells.
• Electropositive Nature: Metals form cations by losing electrons:
  • $Na - e^- \rightarrow Na^+$
  • $Ca - 2e^- \rightarrow Ca^{2+}$
  • $Al - 3e^- \rightarrow Al^{3+}$
• Oxide Formation: Metals generally form basic oxides. These oxides react with acids to produce salts and water:
  • $Na_2O + 2HCl \rightarrow 2NaCl + H_2O$
  • $CaO + 2HCl \rightarrow CaCl_2 + H_2O$
  • Note: $Cr_2O_3$ is acidic. Oxides of $Al$, $Zn$, and $Pb$ are amphoteric.
• Reaction with Acids: Metals placed above hydrogen in the activity series can replace hydrogen ions from acids to form salts:
  • $Mg + 2HCl(dil) \rightarrow MgCl_2 + H_2 \uparrow$
  • $Fe + H_2SO_4(dil) \rightarrow FeSO_4 + H_2 \uparrow$
• Hydride Formation: Generally, metals do not form hydrides. Only highly reactive metals like $Na$, $K$, $Ca$, and $Mg$ react with hydrogen to form unstable hydrides:
  • $2Na + H_2 \rightarrow 2NaH$
• Chloride Formation: Metals generally form non-volatile solid chlorides that are not hydrolysed by water (exceptions include $BiCl_2$ and $SbCl_3$).
  • Metallic chlorides are electrovalent and act as electrolytes.
  • Example ionization: $NaCl \rightarrow Na^+ + Cl^-$
  • They conduct electricity in a fused state or in solution. For example, passing electric current through molten calcium chloride liberates calcium at the cathode and chlorine at the anode:
    • $CaCl_2 \rightarrow Ca^{2+} + 2Cl^- \text{ (molten)}$
    • Reaction at cathode: $Ca^{2+} + 2e^- \rightarrow Ca$
    • Reaction at anode: $Cl^- - e^- \rightarrow Cl$
    • Final anode product: $Cl + Cl \rightarrow Cl_2 \uparrow$
• Reducing Nature: Metals act as reducing agents because they donate electrons:
  • $Na - e^- \rightarrow Na^+$

Physical Properties of Non-metals

• Physical State: Non-metals can be gases, liquids, or brittle solids at room temperature.
• Lustre: They do not possess lustre, except for iodine and graphite.
• Hardness: They are generally soft, with the exception of diamond.
• Density: They generally have low density, with the exception of diamond.
• Melting and Boiling Points: Typically low, except for carbon, boron, and silicon.
• Conductivity: They are bad conductors of heat and electricity, with the exception of graphite and gas-carbon.
• Malleability and Ductility: They are non-malleable and do not possess ductility (carbon fibre is an exception).
• Brittleness: Non-metals are brittle.
• Alloys: They generally do not form alloys and amalgams, except for carbon and phosphorus.
• Sonorousness: Generally non-sonorous.
• Solubility: They dissolve in many liquid solvents.

Chemical Properties of Non-metals

• Valence Shells: Non-metals generally have 5, 6, or 7 electrons in their valence shells.
• Electronegative Nature: They form anions by accepting electrons (except hydrogen):
  • $Cl + e^- \rightarrow Cl^-$
  • $O + 2e^- \rightarrow O^{2-}$
• Oxide Formation: They generally form acidic oxides that react with bases to form salt and water:
  • $CO_2 + 2NaOH \rightarrow Na_2CO_3 + H_2O$
• Reaction with Dilute Acids: Non-metals do not react with dilute acids.
• Chloride Formation: They generally form volatile chlorides which are liquids or gases. Their chlorides are hydrolysed by water:
  • $PCl_3 + 3H_2O \rightarrow H_3PO_3 + 3HCl$
• Hydride Formation: They react with hydrogen to form stable, covalent hydrides:
  • $S + H_2 \rightarrow H_2S$
• Oxidizing Nature: They act as oxidizing agents by accepting electrons:
  • $S + 2e^- \rightarrow S^{2-}$

Electrochemical (Metal Activity) Series

The activity series is a list of metals arranged in order of decreasing chemical reactivity. Hydrogen is included to indicate its relative position.

• Potassium ($K$): Most reactive metal.
• Sodium ($Na$): Very reactive metal.
• Calcium ($Ca$): Very reactive metal.
• Magnesium ($Mg$): Very reactive metal.
• Aluminium ($Al$): Very reactive metal.
  • Note on extraction: These first five metals (K to Al) are never found as free elements and must be extracted by electrolytic reduction.
• Zinc ($Zn$): Moderately reactive metal.
• Iron ($Fe$): Moderately reactive metal.
• Lead ($Pb$): Moderately reactive metal.
  • Note on occurrence and extraction: Moderately reactive metals are found as oxides, carbonates, or sulphites. They are extracted with carbon or carbon monoxide.
• Hydrogen ($H^*$): Reference point.
• Copper ($Cu$): Less reactive metal.
• Mercury ($Hg$): Less reactive metal.
• Silver ($Ag$): Less reactive metal.
• Gold ($Au$): Least reactive metal.
  • Note on occurrence: Less reactive metals may be found as free elements.

Occurrence of Metals and Basic Terminology

• Free State: Metals like $Au$, $Pt$, $Hg$, and $Cu$ are found in nature in an elementary or native form.
• Combined State: Metals found in the form of compounds in the earth's crust along with impurities.
• Minerals: Naturally occurring compounds of metals found in the earth's crust along with impurities (oxides, sulphides, halides, silicates, sulphates, phosphates, chlorides, and carbonates).
• Ore: A mineral from which a metal can be extracted easily and profitably.
• Gangue: The collective term for various types of impurities present in minerals.
• Metallurgy: The process of extracting metals in pure form from their ores.

The Process of Metallurgy

1. Concentration (or Dressing) of Ore: Removal of impurities (gangue).
   • Methods: Electromagnetic separation, Gravity separation (hydraulic washing), and Froth floatation process.
2. Treatment of Concentrated Ore:
   • Conversion: Converting the ore into its oxide through calcination or roasting.
   • Reduction: The process where metallic ions gain electrons to become atoms (e.g., using carbon reduction or electrolysis).
3. Refining of Metals: Purification via Liquation, Distillation, Oxidation, or Electrorefining.

Iron ($Fe$)

• Ores of Iron:
  • Haematite: Ferric oxide ($Fe_2O_3$)
  • Magnetite: Tri-iron tetroxide ($Fe_3O_4$) [Note: Transcript says $Fe_2O_4$, likely a typo for Magnetite which is traditionally $Fe_3O_4$]
  • Iron pyrites: Iron sulphide ($FeS_2$)
  • Limonite: Hydrated ferric oxide ($2Fe_2O_3 \cdot 3H_2O$)
  • Spathic iron ore: Ferrous carbonate ($FeCO_3$)
• Uses of Iron:
  • Cast Iron: Used for articles not requiring great strength (drain pipes, gutter covers, railings, stoves, steam radiators).
  • Wrought Iron: Used for making wires, cables, chains, and forged articles.
• Corrosion (Rusting): Iron exposed to moist air forms rust, a brown, flaky, soft, and porous substance.
• Prevention of Rusting: Surface coating and galvanization.
• Galvanization: Coating iron or steel with a thin layer of zinc to prevent corrosion.

Aluminium ($Al$)

• Ores of Aluminium:
  • Bauxite: Hydrated aluminium oxide ($Al_2O_3 \cdot 2H_2O$)
  • Cryolite: Sodium aluminium fluoride ($Na_3AlF_6$)
  • Corundum or Alumina: Aluminium oxide ($Al_2O_3$)
• Extraction: Extracted from bauxite in two steps:
  • Baeyer’s Process: Conversion of impure bauxite into pure alumina.
  • Hall-Heroult’s Process: Recovery of aluminium from fused alumina via electrolytic reduction.
• Uses: Electric wires/cables, alumino-thermic (thermite) welding, extraction of metals like chromium and barium, deoxidiser in the steel industry, light/strong alloys, cooking utensils, and frames for doors/windows.

Zinc ($Zn$)

• Ores of Zinc:
  • Zinc blende: Zinc sulphide ($ZnS$)
  • Calamine: Zinc carbonate ($ZnCO_3$)
  • Zincite: Zinc oxide ($ZnO$)
• Uses: Galvanizing steel, amalgamated rods for electrodes, manufacturing alloys (brass, German silver, bronze), negative electrodes in dry Leclanche cells, extraction of silver and gold, reducing agent, and zinc dust for drugs/dyes.

Alloys

• Definition: A mixture of two or more metals, or a metal and a non-metal, fused together in a molten state in a fixed proportion.
• Purpose: To enhance the properties of constituents to achieve desired characteristics.
• Examples:
  • Brass: Alloy of $Cu$ and $Zn$.
  • Bronze: Alloy of $Cu$ and $Zn$.