Metals: Extraction, Reactivity, and Corrosion and Quantitative Analysis and Corrosion

Characteristics and Industrial Applications of Common Metals

• Iron

  • Common Ore: Haematite ($Fe_2O_3$)
  • Main Uses: Construction work, ship bodies.
  • Properties and Reasons for Use: Hard, strong, malleable, and relatively cheap.

• Copper

  • Common Ore: Copper Pyrite ($CuFeS_2$)
  • Main Uses and Reasoning:
    1. Electric wires: Good electrical conductor and ductile.
    2. Water pipes: Corrosion-resistant and non-toxic.
    3. Cooking utensils: Good heat conductor, non-toxic, and corrosion-resistant.

• Aluminium

  • Common Ore: Bauxite ($Al_2O_3$)
  • Main Uses and Reasoning:
    1. Overhead power cables: Good electrical conductor and ductile.
    2. Aircraft bodies: Strong, corrosion-resistant, and low density ($2.7\,g/cm^3$).
    3. Kitchen foil and soft drink cans: Non-toxic, corrosion-resistant, and malleable.
    4. Window frames: Strong and corrosion-resistant.

• Gold

  • Common Ore: Gold ore ($Au$), often found in a native state.
  • Main Use: Jewellery.
  • Reasons for Use: Highly corrosion-resistant and has an attractive appearance.

• Mercury

  • Common Ore: Cinnabar ($HgS$)
  • Main Use: Thermometers.
  • Reasons for Use: Expands uniformly on heating and remains liquid at room conditions.

• Silver

  • Common Ore: Argentite ($Ag_2S$)
  • Main Uses and Reasoning:
    1. Jewellery: Corrosion-resistant and attractive appearance.
    2. Electronic components: Best electrical conductor, ductile, and corrosion-resistant.

• Lead

  • Common Ore: Galena ($PbS$)
  • Main Use: Fishing sinkers.
  • Reasons for Use: High density.

• Zinc

  • Common Ore: Zinc sulphide ($ZnS$)
  • Main Use: Galvanizing iron.
  • Reasons for Use: Forms a protective layer that prevents corrosion of the underlying iron.

• Titanium

  • Main Uses and Reasoning:
    1. Supersonic aircraft bodies: Low density, strong, and has a high melting point.
    2. Medical implants (in the human body): Low density, corrosion-resistant, and biocompatible.

Extraction Methodologies Based on Metal Reactivity

• Extraction of Very Unreactive Metals (e.g., Gold, Platinum)

  • These metals are often found as free elements in nature.
  • The process involves mining the native metal from the earth.

• Extraction of Unreactive Metals (e.g., Silver, Mercury)

  • From Metal Oxide: Heat the metal oxide alone ($\text{Metal oxide} \rightarrow \text{Metal} + \text{Oxygen}$).
  • From Metal Sulphide: Heat the metal sulphide in air ($\text{Metal sulphide} + \text{Oxygen} \rightarrow \text{Metal} + \text{Sulphur dioxide}$).

• Extraction of Moderately Reactive Metals (e.g., Copper, Lead, Iron, Zinc)

  • From Metal Oxide: Heat the metal oxide with a reducing agent like carbon or carbon monoxide ($\text{Metal oxide} + C/CO \rightarrow \text{Metal} + CO_2$).
  • From Metal Sulphide: First, convert the sulphide to an oxide by heating in air, then reduce the oxide with carbon ($\text{Metal sulphide} + O_2 \rightarrow \text{Metal oxide} + SO_2$ followed by $\text{Metal oxide} + C \rightarrow \text{Metal} + CO_2$).

• Extraction of Very Reactive Metals (e.g., Aluminium, Magnesium, Calcium, Sodium, Potassium)

  • Process: Electrolysis of the molten ore.
  • From Metal Oxide: Electrolysis of the molten oxide ($\text{Metal oxide} \xrightarrow{Electrolysis} \text{Metal} + O_2$).
  • From Metal Chloride: Electrolysis of the molten chloride ($\text{Metal chloride} \xrightarrow{Electrolysis} \text{Metal} + Cl_2$).

The Blast Furnace Process for Iron Extraction

• Step 1: Production of Carbon Dioxide

  • Hot air is blasted near the bottom of the furnace where coke (carbon) reacts with oxygen.
  • Word Equation: $\text{Carbon} + \text{Oxygen} \rightarrow \text{Carbon dioxide}$
  • Chemical Equation: $C + O_2 \rightarrow CO_2$

• Step 2: Production of Carbon Monoxide

  • Coke reacts with the carbon dioxide produced in the first step to form carbon monoxide, the primary reducing agent.
  • Word Equation: $\text{Carbon} + \text{Carbon dioxide} \rightarrow \text{Carbon monoxide}$
  • Chemical Equation: $C + CO_2 \rightarrow 2CO$

• Step 3: Reduction of haematite to Iron

  • Carbon monoxide reacts with haematite ($Fe_2O_3$). The iron(III) oxide is reduced to iron metal.
  • Word Equation: $\text{Carbon monoxide} + \text{Iron(III) oxide} \rightarrow \text{Carbon dioxide} + \text{Iron}$
  • Chemical Equation: $3CO + Fe_2O_3 \rightarrow 3CO_2 + 2Fe$

• Step 4: Slag Formation and Impurity Removal

  • Limestone ($CaCO_3$) is added to the furnace. It decomposes upon heating to form calcium oxide ($CaO$).
  • $CaCO_3 \rightarrow CaO + CO_2$
  • Calcium oxide reacts with impurities such as silicon dioxide ($SiO_2$) to form molten slag (calcium silicate), which is then removed.
  • Word Equation: $\text{Calcium oxide} + \text{Silicon dioxide} \rightarrow \text{Calcium silicate}$
  • Chemical Equation: $CaO + SiO_2 \rightarrow CaSiO_3$

Electrolytic Extraction of Aluminium

• Raw Materials

  1. Bauxite: The source of the aluminium ore, containing mainly Al_2_O_3.
  2. Cryolite: Used to reduce the melting point of the aluminium oxide mixture from above $2000\,^\circ\text{C}$ to below that temperature, facilitating electrolysis.

• Electrode Setup

  • Anode (+ve electrode): Graphite electrodes.
  • Cathode (-ve electrode): Graphite lining of the electrolytic cell.

• Chemical Reactions

  • The aluminium oxide must be in molten form.
  • Half equation at the Anode (+ve): $2O^{2-} \rightarrow O_2 + 4e^-$
  • Half equation at the Cathode (-ve): $Al^{3+} + 3e^- \rightarrow Al$
  • Due to high operative temperatures, the aluminium metal is produced in a liquid (molten) state.

Comparative Reactivity and Historical Discovery of Metals

• Reactivity Series and Discovery

  • The discovery order of metals in history is directly related to the ease of their extraction.
  • Reactive metals form very stable ores (compounds), making them harder to extract.
  • Less reactive metals form less stable ores, which can be extracted with less energy or simpler methods.

• Metal Appearances

  • Copper: Reddish-brown solid.
  • Gold: Golden-yellow solid.
  • Mercury: Silvery-grey liquid.
  • Others (Iron, Aluminium, etc.): Silvery-grey solids.

Availability, Conservation, and Resource Management

• Factors Affecting Availability

  1. Abundance of the metal in the Earth's crust.
  2. The ease and financial cost of mining the ore.
  3. The ease and financial cost of extracting the metal from its ore.

• Metal Conservation Strategies

  • Metals are non-renewable resources (limited quantity).
  • Primary Methods: Reusing metal articles, reducing the use of metals, and recycling used metals.
  • Benefits: Conserves ore resources, saves energy, reduces metal waste/land use, and minimizes pollution.
  • Problems: Difficulty and time required for sorting; high costs associated with collection, transportation, melting, and purification.

Chemical Properties and Reaction Observations of Metals

• Reaction with Oxygen ($\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$)

  • Potassium: Gentle heating; burns with a lilac flame; forms a white powder ($4K(s) + O_2(g) \rightarrow 2K_2O(s)$).
  • Sodium: Gentle heating; burns with a golden yellow flame; forms a white powder ($4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$).
  • Calcium: Strong heating; burns with a brick-red flame; forms a white powder ($2Ca(s) + O_2(g) \rightarrow 2CaO(s)$).
  • Magnesium: Strong heating; burns with a bright white flame; forms a white powder ($2Mg(s) + O_2(g) \rightarrow 2MgO(s)$).
  • Aluminium: Strong heating; burns to give heat and a white powder ($4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(s)$).
  • Zinc: Strong heating; powder is yellow when hot and white when cold ($2Zn(s) + O_2(g) \rightarrow 2ZnO(s)$).
  • Iron: Strong heating; burns with yellow sparks; forms a black powder ($3Fe(s) + 2O_2(g) \rightarrow Fe_3O_4(s)$). Iron(III) oxide is otherwise brown.
  • Lead: Strong heating; melts and yields a powder that is orange when hot and yellow when cold ($2Pb(s) + O_2(g) \rightarrow 2PbO(s)$).
  • Copper: Very strong heating; forms a black powder ($2Cu(s) + O_2(g) \rightarrow 2CuO(s)$).
  • Mercury: Very strong heating; forms a red powder ($2Hg(l) + O_2(g) \rightarrow 2HgO(s)$).
  • Silver, Platinum, Gold: No reaction.

• Reaction with Water and Steam

  • Potassium and Sodium: React violently with cold water to form metal hydroxide and hydrogen gas. They melt into silvery balls and move on the surface with a hissing sound. Potassium burns with a lilac flame; Sodium burns with a golden yellow flame.
  • Calcium: Reacts readily with cold water; sinks; forms a milky suspension of calcium hydroxide.
  • Magnesium: Reacts very slowly with hot water. Reacts vigorously with steam to produce magnesium oxide and hydrogen gas, giving off an intense white light.
  • Aluminium: Reaction with steam is usually prevented by a thin, tough layer of aluminium oxide.
  • Zinc and Iron: React with steam to produce metal oxides. Iron forms a black powder ($Fe_3O_4$).
  • Lead, Copper, Mercury, Silver, Platinum, Gold: No reaction with water or steam.

• Reaction with Dilute Acids ($\text{Metal} + \text{Acid} \rightarrow \text{Metal salt} + \text{Hydrogen}$)

  • Potassium and Sodium: Reaction is explosive and not attempted.
  • Calcium: Reacts readily. Note: with dilute sulphuric acid, the reaction stops quickly because insoluble calcium sulphate coats the metal surface.
  • Magnesium, Aluminium, Zinc: React readily to give colorless hydrogen gas.
  • Iron: Reacts readily; the solution turns from colorless to green ($Fe^{2+}$).
  • Lead: Reacts very slowly. The reaction stops quickly with dilute hydrochloric or sulphuric acid due to the formation of insoluble lead(II) chloride or lead(II) sulphate layers.
  • Copper, Mercury, Silver, Platinum, Gold: No reaction.

Advanced Displacement Reactions and Ionic Equations

• Definition: A reaction where a more reactive element displaces a less reactive element from its compound ($A + B^{n+} \rightarrow A^{n+} + B$).

• Example: Zinc and Copper(II) Sulphate

  • Chemical Equation: $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$
  • Ionic Equation: $Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$
  • Observations: Zinc metal dissolves ($Zn \rightarrow Zn^{2+}$); blue color of the solution pales ($Cu^{2+}$ used up); a reddish-brown solid (copper) forms; heat is evolved (exothermic).

• Example: Sodium and Copper(II) Sulphate

  • Sodium reacts with the water in the solution first: $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$.
  • Secondary reaction: The $OH^-$ ions formed react with the $Cu^{2+}$ ions to form a blue precipitate of copper(II) hydroxide ($Cu(OH)_2$).

• Steps for Writing Ionic Equations

  1. Write the full balanced chemical equation.
  2. Expand soluble aqueous compounds into individual ions.
  3. Cancel "spectator ions" (ions that appear unchanged on both sides).
  4. If all ions cancel, there is no reaction.

Quantitative Chemical Analysis and Stoichiometry

• Mole and Particle Calculations

  • $\text{Mole} = \frac{\text{Mass}}{\text{Molar Mass}}$
  • $\text{Number of particles} = \text{Mole} \times 6.02 \times 10^{23}$
  • Example: In $60.1\,g$ of quartz ($SiO_2$), there are $1.204 \times 10^{24}$ oxygen atoms (since $1\,mol$ of $SiO_2$ has $2\,mol$ of $O$ atoms).

• Percentage Yield

  • $\text{Percentage yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$
  • Reasons for

• Empirical and Molecular Formulas

  • Empirical formula is the simplest whole-number ratio of atoms in a compound.
  • Molecular formula is a multiple ($n$) of the empirical formula where $n = \frac{\text{Molecular mass}}{\text{Empirical formula mass}}$.
  • Example: An oxide of lead with $90.6\%$ lead by mass yields an empirical formula of $Pb_3O_4$ (Mole ratio $Pb:O = 1:1.34 \approx 3:4$).

Corrosion, Rusting Mechanisms, and Prevention Strategies

• Essential Conditions for Rusting

  • Rusting specifically refers to the corrosion of Iron.
  • Requirements: Water ($H_2O$) and Oxygen ($O_2$).

• The Chemistry of Rusting

  • Step 1: Iron loses electrons to form Iron(II) ions ($Fe \rightarrow Fe^{2+} + 2e^-$). Water and oxygen gain electrons to form hydroxide ions ($O_2 + 2H_2O + 4e^- \rightarrow 4OH^-$).
  • Step 2: Formation of Iron(II) hydroxide ($Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2$).
  • Step 3: Further oxidation to Iron(III) hydroxide ($4Fe(OH)_2 + O_2 + 2H_2O \rightarrow 4Fe(OH)_3$).
  • Step 4: Dehydration to form hydrated iron(III) oxide (Rust): $Fe_2O_3 \cdot nH_2O$.

• Rust Indicators

  • Potassium hexacyanoferrate(III): Turns blue in the presence of $Fe^{2+}$ ions.
  • Phenolphthalein: Turns pink in the presence of $OH^-$ ions.

• Factors Speeding Up Rusting

  1. Presence of soluble ionic compounds (salts) or acids (increase conductivity).
  2. Attachment of less reactive metals (e.g., Tin, Copper).
  3. High temperatures.
  4. Physical damage (scratches, sharp bends).

• Methods of Rust Prevention

  • Painting/Plastic Coating: Physical barrier; cheap, but can be scratched off (e.g., Bridges, car bodies, coat hangers).
  • Greasing/Oiling: For moving parts; prevents contact with air/water and lubricates (e.g., Tools, engines).
  • Galvanizing (Zinc-plating): Zinc is more reactive than iron, providing sacrificial protection even if scratched (e.g., Buckets, construction iron).
  • Tin-plating: Barrier method. If the tin layer is scratched, the iron underneath rusts faster because iron is more reactive than tin (e.g., Food cans).
  • Cathodic Protection: Connecting the iron to the negative terminal of a DC source to supply it with electrons.
  • Sacrificial Protection: Attaching a more reactive metal (e.g., Magnesium or Zinc) which loses electrons in place of the iron.
  • Alloying: Mixing iron with other elements (e.g., Stainless steel containing Cr and Ni) to form a self-protecting oxide layer.

Aluminium Anodization and Surface Protection

• Natural Resistance: Aluminium forms a thin, impermeable layer of aluminium oxide ($Al_2O_3$) immediately upon contact with air, which adheres firmly and prevents further corrosion.

• Anodization Process: An electrolytic process used to thicken this oxide layer.

  • Electrolyte: Dilute sulphuric acid.
  • Cathode (-ve): Aluminium sheet.
  • Anode (+ve): The aluminium object to be anodized.
  • Reaction at Anode: $2Al(s) + 3H_2O(l) \rightarrow Al_2O_3(s) + 6H^+(aq) + 6e^-$.
  • Outcome: Anodized aluminium is harder, more corrosion-resistant, and easier to dye for decorative purposes.

Questions & Discussion

• Question: Explain why iron was discovered and used much earlier than aluminium.

• Answer: Iron is less reactive than aluminium. The ores of iron are less stable and can be reduced by carbon in a furnace, whereas aluminium requires electricity for extraction, which was not available in early history.

• Question: Explain whether carbon can be used to extract magnesium from magnesium oxide.

• Answer: No. Magnesium oxide is extremely stable because magnesium is a very reactive metal. Carbon is not a strong enough reducing agent to remove the oxygen from MgO. Magnesium can only be extracted through the electrolysis of molten magnesium ore.

• Question: Why does the reaction between lead and dilute sulphuric acid stop quickly?

• Answer: Lead reacts with the acid to form lead(II) sulphate, which is insoluble in water. This insoluble salt forms a layer over the remaining metal, preventing further contact between the lead and the acid.