Menu

Tools Submit Materials Pricing About Contact Login Register

Unit 10: Chemistry of Metals

Chemistry - Class 11

This chapter delves into the fundamental principles of metallurgy, exploring various techniques for extracting and refining metals from their ores. It then provides a detailed examination of the alkali and alkaline earth metals, covering their characteristic properties, important compounds, and industrial applications.

No MCQ questions available for this chapter.

Unit 10: Chemistry of Metals

10.1 Metals and Metallurgical Principles

Metallurgy is the science and technology concerned with the extraction of metals from their ores and their subsequent purification, alloying, and fabrication for use. It encompasses a wide range of processes designed to transform raw mineral resources into usable metallic products.

Types of Metallurgical Processes

  • Hydrometallurgy: This involves the use of aqueous solutions to extract metals from ores. The metal is dissolved in a solvent, and then recovered from the solution, often by precipitation or electrolysis.
  • Pyrometallurgy: These are high-temperature processes involving chemical reactions to extract metals. Examples include roasting, calcination, and smelting, which typically occur in furnaces.
  • Electrometallurgy: This method uses electrochemical processes, primarily electrolysis, to extract or refine metals. It is particularly effective for highly reactive metals or for achieving high purity.

Key Terms in Metallurgy

  • Ores: A naturally occurring mineral or rock from which a metal or valuable mineral can be economically extracted. Examples include bauxite (for aluminium), haematite (for iron), and galena (for lead).
  • Gangue/Matrix: These are the unwanted, commercially worthless materials (impurities) that are present in an ore. They are typically rocky or earthy substances.
  • Flux: A substance added during smelting to combine with the gangue, forming a more easily fusible product called slag. Common fluxes include limestone (CaCO3) and silica (SiO2).
  • Slag: The fusible product formed when a flux combines with gangue during the smelting process. Slag is lighter than the molten metal and floats on top, allowing for easy separation. It often consists of silicates.
  • Alloy: A mixture of two or more metals, or a mixture of a metal and a non-metal, designed to enhance specific properties like strength, hardness, or corrosion resistance. For example, brass is an alloy of copper and zinc. An amalgam is a specific type of alloy where one of the constituent metals is mercury.

General Principles of Metal Extraction

The extraction of a metal from its ore typically involves several sequential steps:

  1. Concentration/Beneficiation: This initial step aims to remove a significant portion of the gangue from the ore, increasing the concentration of the desired metal-bearing mineral.
    • Gravity Separation: Used for ores where the metal compound and gangue have significant density differences. The ore is washed with water, and the heavier ore particles settle faster than the lighter gangue particles.
    • Magnetic Separation: Applicable to ores where either the ore particles or the gangue particles are magnetic. The crushed ore is passed over a magnetic roller, separating magnetic from non-magnetic components. For example, chromite ore (FeCr2O4) can be separated from non-magnetic impurities.
    • Froth Flotation: Primarily used for sulfide ores. The finely crushed ore is mixed with water, a frothing agent (e.g., pine oil), and a collector (e.g., potassium ethyl xanthate). Air is blown through the mixture, creating froth. The sulfide ore particles preferentially attach to the oil and rise with the froth, while the gangue settles at the bottom.
    • Leaching: A chemical concentration method where the ore is treated with a suitable chemical reagent that selectively dissolves the metal-bearing compound, leaving the gangue undissolved. For example, bauxite (Al2O3.xH2O) is leached with concentrated NaOH solution in the Bayer process to dissolve alumina.
  2. Calcination: This process involves heating the concentrated ore in a limited supply of air (or in the absence of air) to a high temperature below its melting point.
    • Purpose: To remove volatile impurities (like moisture, organic matter) and to convert carbonate ores into their respective metal oxides.
    • Example: Decomposition of limestone to quicklime: CaCO3 (s) → CaO (s) + CO2 (g)
    • Example: Decomposition of hydrated bauxite: Al2O3.xH2O (s) → Al2O3 (s) + xH2O (g)
  3. Roasting: This process involves heating the concentrated ore strongly in the presence of excess air (or oxygen) to a high temperature below its melting point.
    • Purpose: To remove volatile impurities like sulfur, arsenic, and antimony as their gaseous oxides (e.g., SO2) and to convert sulfide ores into their corresponding metal oxides.
    • Example: Roasting of zinc sulfide ore: 2ZnS (s) + 3O2 (g) → 2ZnO (s) + 2SO2 (g)
    • Example: Roasting of galena: 2PbS (s) + 3O2 (g) → 2PbO (s) + 2SO2 (g)
  4. Smelting: This is a pyrometallurgical process where the metal oxide (obtained from calcination or roasting) is reduced to its molten metallic state. It typically occurs at very high temperatures in a furnace, often with a reducing agent and a flux.
  5. Carbon Reduction: A common smelting method where coke (carbon) or carbon monoxide (CO) acts as the reducing agent to convert metal oxides into free metals. This method is suitable for metals that are less reactive than carbon.
    • Principle: Carbon reduces the metal oxide by removing oxygen, forming carbon monoxide or carbon dioxide.
    • Example: Reduction of zinc oxide: ZnO (s) + C (s) → Zn (l) + CO (g)
    • Example: Reduction of iron oxide in a blast furnace: Fe2O3 (s) + 3CO (g) → 2Fe (l) + 3CO2 (g)
  6. Thermite Reduction: This is an exothermic reduction process where aluminium powder is used as a strong reducing agent for metal oxides of metals that have a high affinity for oxygen (e.g., Fe2O3, Cr2O3, MnO2).
    • Principle: Aluminium is more reactive than many metals and displaces them from their oxides, releasing a large amount of heat which melts the product metal.
    • Example: Reduction of ferric oxide: Fe2O3 (s) + 2Al (s) → 2Fe (l) + Al2O3 (s) + Heat
  7. Electrochemical Reduction: This method uses electrolysis to extract highly reactive metals (e.g., sodium, potassium, calcium, magnesium, aluminium) that cannot be easily reduced by carbon or other chemical reducing agents.
    • Principle: The molten ore (or a molten mixture containing the metal compound) is electrolyzed. The metal ions gain electrons at the cathode and are deposited as pure metal, while non-metal ions lose electrons at the anode.
    • Example: Extraction of aluminium from molten alumina (Al2O3) in the Hall-Héroult process.
    • Example: Extraction of sodium from molten sodium chloride (Down's process).

Refining of Metals

After extraction, most metals contain impurities and require further purification, known as refining, to achieve the desired purity level.

  • Poling: This is a method used for refining metals that contain oxides as impurities, particularly those like copper or tin that dissolve their own oxides. The molten impure metal is stirred with green wood logs. The hydrocarbons from the green wood act as reducing agents, reducing the dissolved metal oxides (e.g., Cu2O) back to the metal and removing dissolved gases like O2 and SO2.
  • Electro-refining: This is a highly effective method for obtaining very pure metals, especially copper.
    • Setup:
      • Anode: Made of the impure metal.
      • Cathode: A thin sheet of pure metal.
      • Electrolyte: A solution of a salt of the same metal (e.g., CuSO4 solution for copper refining).
    • Process: When an electric current is passed, the impure metal at the anode oxidizes (M → M^n+ + ne-) and dissolves into the electrolyte. More reactive impurities also oxidize and dissolve. Less reactive impurities simply fall to the bottom as "anode mud." At the cathode, only the pure metal ions from the electrolyte are preferentially reduced and deposited as pure metal (M^n+ + ne- → M).
    • Example: Refining of copper, where impure copper is the anode, pure copper is the cathode, and CuSO4 solution is the electrolyte.

10.2 Alkali and Alkaline Earth Metals

Alkali Metals (Group 1)

The alkali metals include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). They are highly reactive elements with distinct properties.

  • General characteristics:
    • They readily form +1 ions by losing their single valence electron.
    • They are highly reactive, increasing in reactivity down the group.
    • They are soft, silvery-white metals that can be cut with a knife.
    • They have low densities and low melting points, which decrease down the group.
    • They are strong reducing agents, meaning they are easily oxidized.
    • They react vigorously with water, producing hydrogen gas and forming strong bases (metal hydroxides).

Sodium (Na)

Sodium is a very important alkali metal with numerous applications.

  • Extraction from Down's process: Sodium metal is extracted by the electrolytic reduction of molten sodium chloride (NaCl).
    • Principle: Molten NaCl is electrolyzed in a Down's cell. CaCl2 (calcium chloride) is added to the NaCl to lower its melting point from 801°C to about 600°C, making the process more energy-efficient.
    • Reactions:
      • At Cathode (reduction): Na+ (l) + e- → Na (l)
      • At Anode (oxidation): 2Cl- (l) → Cl2 (g) + 2e-
  • Properties:
    • Action with Oxygen: Sodium reacts with oxygen to form different oxides depending on the conditions.
      • In limited oxygen: 4Na (s) + O2 (g) → 2Na2O (s) (sodium oxide)
      • In excess oxygen: 2Na (s) + O2 (g) → Na2O2 (s) (sodium peroxide)
    • Action with Water: Sodium reacts vigorously and exothermically with water, producing sodium hydroxide and hydrogen gas. The reaction is so exothermic that the hydrogen gas often ignites.
      • 2Na (s) + 2H2O (l) → 2NaOH (aq) + H2 (g)
    • Action with Acids: Sodium reacts explosively with dilute acids, producing a salt and hydrogen gas.
      • 2Na (s) + 2HCl (aq) → 2NaCl (aq) + H2 (g)
    • Action with Nonmetals: Sodium reacts directly with many nonmetals upon heating to form ionic compounds.
      • With Chlorine: 2Na (s) + Cl2 (g) → 2NaCl (s)
      • With Sulfur: 2Na (s) + S (s) → Na2S (s)
    • Action with Ammonia: Sodium dissolves in liquid ammonia to form a deep blue solution, which is a strong reducing agent.
      • Na (s) + (x+y)NH3 (l) → Na+ (am) + e- (am) (where 'am' denotes ammoniated species)

Sodium Hydroxide (NaOH)

Sodium hydroxide, also known as caustic soda, is a strong base.

  • Precipitation reactions: NaOH reacts with solutions of many metal salts to precipitate insoluble metal hydroxides.
    • CuSO4 (aq) + 2NaOH (aq) → Cu(OH)2 (s) + Na2SO4 (aq) (blue precipitate)
    • FeCl3 (aq) + 3NaOH (aq) → Fe(OH)3 (s) + 3NaCl (aq) (reddish-brown precipitate)
  • Action with CO: Sodium hydroxide can react with carbon monoxide under specific conditions (high temperature and pressure) to form sodium formate.
    • NaOH (s) + CO (g) → HCOONa (s) (sodium formate)
  • Uses: Used in the manufacture of soap, paper, textiles, detergents, and as a strong base in various chemical processes.

Sodium Carbonate (Na2CO3)

Sodium carbonate, or soda ash, is an important industrial chemical.

  • Solvay process preparation: The Solvay process is the primary industrial method for producing sodium carbonate.
    • Key steps: Ammonia (NH3) is dissolved in saturated brine (NaCl solution), and then carbon dioxide (CO2) is passed through the ammoniated brine. This leads to the formation and precipitation of sodium bicarbonate (NaHCO3).
      • NH3 (g) + H2O (l) + CO2 (g) → NH4HCO3 (aq)
      • NH4HCO3 (aq) + NaCl (aq) → NaHCO3 (s) + NH4Cl (aq)
    • The precipitated NaHCO3 is then filtered and heated (calcined) to produce sodium carbonate.
      • 2NaHCO3 (s) → Na2CO3 (s) + H2O (g) + CO2 (g)
  • Properties:
    • Action with CO2: Sodium carbonate solution can react with excess carbon dioxide to form sodium bicarbonate.
      • Na2CO3 (aq) + H2O (l) + CO2 (g) → 2NaHCO3 (aq)
    • Action with SO2: Sodium carbonate solution can absorb sulfur dioxide, forming sodium sulfite.
      • Na2CO3 (aq) + SO2 (g) → Na2SO3 (aq) + CO2 (g)
    • Action with Water: Sodium carbonate dissolves in water to form a basic solution due to hydrolysis of the carbonate ion.
      • CO3^2- (aq) + H2O (l) ↔ HCO3- (aq) + OH- (aq)
    • Precipitation: Sodium carbonate can be used to precipitate insoluble metal carbonates from solutions of their salts.
      • CaCl2 (aq) + Na2CO3 (aq) → CaCO3 (s) + 2NaCl (aq)
  • Uses: Used in glass manufacturing, paper industry, detergents, water softening, and as a raw material for other sodium compounds.

Alkaline Earth Metals (Group 2)

The alkaline earth metals include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). They are less reactive than alkali metals but still quite reactive.

  • General characteristics:
    • They readily form +2 ions by losing their two valence electrons.
    • They are reactive, though less so than the corresponding alkali metals in the same period. Reactivity increases down the group.
    • They are harder, denser, and have higher melting points than alkali metals.
    • They are strong reducing agents.
    • They react with water to form hydroxides and hydrogen gas, though less vigorously than alkali metals (e.g., magnesium reacts slowly with cold water, calcium reacts more readily).

Important Compounds of Alkaline Earth Metals

  • Quick lime (CaO): Calcium oxide, produced by heating limestone. Used in cement, steelmaking, and water treatment.
  • Bleaching powder (CaOCl2): Calcium oxychloride, produced by passing chlorine gas over dry slaked lime (Ca(OH)2). Used as a bleaching agent and disinfectant.
  • Magnesia (MgO): Magnesium oxide, a refractory material with a high melting point. Used in furnace linings, antacids, and laxatives.
  • Plaster of Paris (CaSO4.1/2H2O): Calcium sulfate hemihydrate, formed by heating gypsum (CaSO4.2H2O). Used in casts for broken bones, dental molds, and decorative work.
  • Epsom salt (MgSO4.7H2O): Magnesium sulfate heptahydrate. Used as a laxative, bath salt, and in agriculture as a fertilizer.

Solubility Trends of Alkaline Earth Metal Compounds

  • Hydroxides (M(OH)2): Solubility generally increases down the group (e.g., Be(OH)2 is amphoteric, Mg(OH)2 is sparingly soluble, Ca(OH)2, Sr(OH)2, Ba(OH)2 are increasingly soluble).
  • Carbonates (MCO3): Solubility generally decreases down the group (e.g., MgCO3 is more soluble than CaCO3, which is more soluble than BaCO3).
  • Sulphates (MSO4): Solubility generally decreases down the group (e.g., MgSO4 is highly soluble, CaSO4 is sparingly soluble, while SrSO4 and BaSO4 are largely insoluble).

Stability Trends of Alkaline Earth Metal Compounds

The thermal stability of carbonates and nitrates of alkaline earth metals is influenced by the polarizing power of the cation. Following the provided instructions:

  • Carbonates (MCO3): The thermal stability of alkaline earth metal carbonates is stated to become less stable down the group. This means they decompose at lower temperatures as you go from Be to Ba.
  • Nitrates (M(NO3)2): Similarly, the thermal stability of alkaline earth metal nitrates is stated to become less stable down the group. They decompose more readily at lower temperatures from Be to Ba.
Previous Next