Basic Metallurgical Processes & Refining - Tenth Grade Science
The technological infrastructure of human society depends upon the extraction and refining of metallic elements from naturally occurring mineral deposits in the Earth's crust. Metallurgy is the applied science and technology of extracting metals economically from their ores and refining them for practical use. While low-reactivity metals like gold and platinum occur in the native metallic state, most metals exist chemically combined as oxides, sulfides, carbonates, or halides, intermixed with rocky, earthy impurities known as gangue. The chemical route chosen to extract a specific metal is dictated by its position in the reactivity series. In this chapter, you will investigate the sequential stages of metallurgy—concentration of ore, conversion to oxide, reduction to crude metal, and electrolytic refining—while exploring thermal roasting, calcination, aluminothermic reduction, and corrosion prevention.
Mineral Geological Abundance and Ore Terminology
Minerals are naturally occurring inorganic substances possessing a definite chemical composition and crystalline structure, distributed throughout the Earth's crust.
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| METALLURGICAL VOCABULARY FRAMEWORK |
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| MINERAL: Naturally occurring inorganic compound containing a metal. |
| ORE: A mineral deposit from which a metal can be extracted profitably. |
| Rule: "All ores are minerals, but not all minerals are ores." |
| GANGUE: Commercially worthless rocky, silicious, or clay impurities in ore. |
| FLUX: Chemical reagent added to combine with gangue to form fusible slag. |
| SLAG: Fusible, low-density waste product skimmed off molten metal. |
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Common Commercial Ores
Commercial ores typically occur in several major chemical classes:
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| Metal | Ore Name | Chemical Formula & Mineral Class |
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| Aluminium | Bauxite | Al2O3.2H2O (Hydrated Oxide) |
| Iron | Haematite | Fe2O3 (Oxide), Magnetite: Fe3O4 |
| Zinc | Zinc Blende | ZnS (Sulfide), Calamine: ZnCO3 (Carbonate) |
| Copper | Copper Pyrites| CuFeS2 (Sulfide), Cuprite: Cu2O |
| Mercury | Cinnabar | HgS (Sulfide) |
| Lead | Galena | PbS (Sulfide) |
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Beneficiation and Concentration of Ores
Crude ore extracted from mines contains high proportions of gangue that must be separated prior to chemical extraction to optimize thermal efficiency and reduce reagent consumption.
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| PRIMARY CONCENTRATION METHODS |
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| Method | Operational Mechanism | Applied Ore Classes |
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| Hydraulic Washing | Density differential under a | Heavy oxide ores |
| (Gravity Sep.) | stream of running water | (Haematite, tinstone) |
|-------------------+-------------------------------+-------------------------------+
| Magnetic Sep. | Differential magnetic | Magnetic ores |
| | susceptibility | (Magnetite Fe3O4 from silica) |
|-------------------+-------------------------------+-------------------------------+
| Froth Flotation | Differential wetting: pine oil| Sulfide ores |
| | for ore, water for gangue | (Zinc blende ZnS, Galena PbS) |
|-------------------+-------------------------------+-------------------------------+
| Chemical Leaching | Selective chemical dissolution| Bauxite (Bayer process), |
| | in hot aqueous reagent | Gold extraction using cyanide |
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The Froth Flotation Process for Sulfide Ores
Froth flotation is the standard concentration technique for sulfide ores, exploiting differences in surface wetting properties between mineral particles and rocky gangue:
Crushed ore is introduced into a large flotation tank containing water, pine oil (a frothing collector), and sodium ethyl xanthate (a collector surfactant). Air is injected through a rotating agitator, creating a vigorous froth. The hydrophobic sulfide ore particles adhere to pine oil bubbles and rise to the surface as a mineral-rich froth, which is skimmed off and dried. The hydrophilic gangue particles remain wetted by water and settle to the bottom of the tank.
Conversion of Concentrated Ore to Metallic Oxide
Metal oxides are chemically more straightforward to reduce than sulfides or carbonates. Consequently, concentrated sulfide and carbonate ores are converted to oxides via roasting or calcination.
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| ROASTING VERSUS CALCINATION |
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| ROASTING | CALCINATION |
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| Heated strongly in EXCESS of air | Heated strongly in the ABSENCE or LIMITED air |
| below the melting point. | below the melting point. |
| Applied to SULFIDE ores. | Applied to CARBONATE and HYDRATED ores. |
| Expels volatile sulfur as SO2 gas.| Expels CO2 gas and water of hydration. |
| Example: | Example: |
| 2ZnS + 3O2 -> 2ZnO + 2SO2 | ZnCO3 --[Heat]--> ZnO + CO2 |
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Roasting Sulfide Ores
Roasting involves heating concentrated sulfide ore strongly in a continuous excess supply of air below the metal's melting point. This oxidizes the metal to an oxide while converting sulfur into sulfur dioxide gas:
2ZnS(s) + 3O2(g) --[Heat]--> 2ZnO(s) + 2SO2(g)
2Cu2S(s) + 3O2(g) --[Heat]--> 2Cu2O(s) + 2SO2(g)
The evolved sulfur dioxide gas is scrubbed and routed to contact-process plants for sulfuric acid synthesis, preventing atmospheric pollution.
Calcination of Carbonate and Hydrated Ores
Calcination involves heating carbonate or hydrated ores strongly in the absence of air, or with limited air access, driving off carbon dioxide and moisture:
ZnCO3(s) --[Heat]--> ZnO(s) + CO2(g)
CaCO3.MgCO3(s) --[Heat]--> CaO(s) + MgO(s) + 2CO2(g) (Dolomite)
Al2O3.2H2O(s) --[Heat]--> Al2O3(s) + 2H2O(g) (Bauxite)
Calcination leaves the ore porous and dry, optimizing it for chemical reduction.
Reduction of Metallic Oxides to Crude Metals
The chemical method used to reduce a metal oxide to free elemental metal depends directly on the position of the metal in the reactivity series.
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| EXTRACTION STRATEGY BY REACTIVITY |
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| Reactivity Tier | Characteristic Metals | Extraction Methodology |
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| High Reactivity | K, Na, Ca, Mg, Al | Electrolytic reduction of |
| | | molten salts (Hall-Heroult|
| Moderate React. | Zn, Fe, Pb, Cu | Chemical reduction using |
| | | carbon (coke) or CO |
| Low Reactivity | Hg, Cu, Ag | Thermal self-reduction / |
| | | simple roasting |
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Self-Reduction of Low-Reactivity Metals
Low-reactivity metals can be reduced by thermal roasting alone. Cinnabar, mercury(II) sulfide ($\text{HgS}$), is converted to mercury(II) oxide upon heating in air, which spontaneously decomposes at roasting temperatures:
2HgS(s) + 3O2(g) --[Heat]--> 2HgO(s) + 2SO2(g)
2HgO(s) --[Heat]--> 2Hg(l) + O2(g)
Similarly, copper glance ($\text{Cu}_2\text{S}$) undergoes auto-reduction: partial roasting yields copper(I) oxide, which then reacts with the remaining copper sulfide:
2Cu2O(s) + Cu2S(s) --[Heat]--> 6Cu(s) + SO2(g)
Carbon Reduction of Moderately Reactive Metals
Moderately reactive metals like zinc, iron, and lead are reduced using carbon (coke) or carbon monoxide at elevated temperatures inside blast furnaces or retorts:
ZnO(s) + C(s) --[1400 deg C]--> Zn(g) + CO(g)
Fe2O3(s) + 3CO(g) --[Blast Furnace]--> 2Fe(l) + 3CO2(g)
PbO(s) + C(s) --[Heat]--> Pb(s) + CO(g)
Aluminothermic Reduction (The Thermite Process)
Certain metallic oxides, such as manganese dioxide ($\text{MnO}_2$) and chromium(III) oxide ($\text{Cr}_2\text{O}_3$), cannot be reduced cleanly by carbon because they form stable metallic carbides. Instead, they are reduced using aluminium powder as a reducing agent in an exothermic process known as the Goldschmidt thermite process:
3MnO2(s) + 4Al(s) -> 3Mn(l) + 2Al2O3(s) + Immense Heat
Fe2O3(s) + 2Al(s) -> 2Fe(l) + Al2O3(s) + Immense Heat
The reaction between iron(III) oxide and aluminium powder releases sufficient thermal energy to heat the produced elemental iron above its melting point ($1538^\circ\text{C}$). The resulting stream of molten iron is used in field thermite welding to join severed railway rails and repair cracked machinery castings.
Electrolytic Reduction of Highly Reactive Metals
Highly reactive metals located at the top of the reactivity series—such as sodium, magnesium, and aluminium—have a higher affinity for oxygen than carbon does. Consequently, carbon cannot reduce their oxides. These metals are extracted by electrolytic reduction of their molten anhydrous chloride or oxide salts.
In the extraction of sodium via the Down's cell, molten sodium chloride containing calcium chloride (added to depress the melting point from $801^\circ\text{C}$ to $600^\circ\text{C}$) is electrolyzed:
At the Cathode: Na+(l) + e- -----> Na(l) (Sodium metal collected)
At the Anode: 2Cl-(l) -----> Cl2(g) + 2e- (Chlorine gas collected)
Aqueous solutions of sodium chloride cannot be used for this extraction because water would reduce preferentially at the cathode, evolving hydrogen gas rather than depositing sodium metal.
In the Hall-Héroult process, purified alumina ($\text{Al}_2\text{O}_3$) is dissolved in molten cryolite ($\text{Na}_3\text{AlF}_6$) to lower the operating temperature and enhance electrical conductivity, yielding molten aluminium at the carbon cathode.
Refining and Purification of Crude Metals
Crude metals produced through smelting or chemical reduction contain residual impurities, slag inclusions, or unreacted metalloids that degrade mechanical and electrical performance. The primary method for high-purity metal refining is electrolytic refining.
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| ELECTROLYTIC REFINING OF COPPER |
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| [+] DC Power Supply [-] |
| | | |
| v v |
| +-----------+ +-----------+ |
| | ANODE | | CATHODE | |
| | Impure Cu | | Pure Cu | |
| | Block | | Strip | |
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| | | |
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| | Acidified Aqueous | |
| | CuSO4 Solution | |
| ~~~~~~~~~~~~~~~~~~~~~ |
| [ Anode Sludge (Au, Ag) ] |
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Electrolytic Refining of Copper
During the electrolytic refining of blister copper: The anode is cast from thick slabs of impure crude copper. The cathode consists of thin sheets of pure electrolytic copper. The electrolyte is an aqueous solution of copper(II) sulfate acidified with dilute sulfuric acid.
When current is applied: At the impure anode, copper atoms oxidize and dissolve into solution as cupric ions:
$$\text{Cu(s, impure)} \longrightarrow \text{Cu}^{2+}\text{(aq)} + 2\text{e}^-$$
At the pure cathode, cupric ions from the electrolyte accept electrons and deposit as pure metallic copper:
$$\text{Cu}^{2+}\text{(aq)} + 2\text{e}^- \longrightarrow \text{Cu(s, pure)}$$
More electropositive impurities in the crude copper (such as iron, nickel, and zinc) dissolve into the electrolyte as cations and remain in solution. Less electropositive noble impurities (such as silver, gold, and platinum) do not oxidize; they detach from the disintegrating anode and settle beneath it as valuable anode sludge (anode mud), which is collected for precious metal recovery.
Chapter Exercises
Exercise 1: An ore on treatment with dilute hydrochloric acid gives brisk effervescence with the evolution of a colorless, odorless gas that turns limewater milky. What type of ore is this? Outline the metallurgical steps required to convert this concentrated ore into the free metal, assuming the metal has moderate reactivity.
Exercise 2: Write balanced chemical equations for the following metallurgical transformations: first, the roasting of zinc blende; second, the thermite reaction between iron(III) oxide and aluminium powder; and third, the cathode and anode half-reactions during the electrolytic reduction of molten aluminium oxide.
Exercise 3: Explain why aluminium cannot be extracted from an aqueous solution of aluminium sulfate by chemical displacement using zinc metal, and explain why carbon reduction is ineffective for reducing aluminium oxide to aluminium metal.
Solutions and Explanations
Solution 1: The ore is a metal carbonate ore, such as calamine ($\text{ZnCO}_3$). The brisk effervescence is caused by the evolution of carbon dioxide gas ($\text{CO}_2$), which precipitates calcium carbonate when bubbled into limewater, turning it milky: $\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2$. To extract the free metal: first, concentrate the carbonate ore using gravity separation or hydraulic washing to remove rocky gangue; second, subject the concentrated carbonate ore to calcination by heating it strongly in the absence of air, driving off carbon dioxide to yield the metal oxide: $\text{ZnCO}_3\text{(s)} \rightarrow \text{ZnO(s)} + \text{CO}_2\text{(g)}$; third, reduce the resulting zinc oxide to free metallic zinc by heating it with carbon (coke) at approximately $1400^\circ\text{C}$ in a retort: $\text{ZnO(s)} + \text{C(s)} \rightarrow \text{Zn(g)} + \text{CO(g)}$, condensing the zinc vapor into molten metal; and fourth, refine the crude zinc by distillation or electrolytic refining to obtain pure zinc metal.
Solution 2: For the roasting of zinc blende: $2\text{ZnS(s)} + 3\text{O}_2\text{(g)} \rightarrow 2\text{ZnO(s)} + 2\text{SO}_2\text{(g)}$. For the aluminothermic reduction of iron(III) oxide: $\text{Fe}_2\text{O}_3\text{(s)} + 2\text{Al(s)} \rightarrow 2\text{Fe(l)} + \text{Al}_2\text{O}_3\text{(s)} + \text{Heat}$. For the electrolytic reduction of molten aluminium oxide in Hall-Héroult cells: at the carbon cathode, aluminium ions accept electrons to form liquid aluminium metal: $\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al(l)}$; at the carbon anode, oxide ions donate electrons to form oxygen gas, which reacts with the sacrificial carbon anode to evolve carbon dioxide: $2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-$, followed by $\text{C(s)} + \text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)}$.
Solution 3: Aluminium is situated higher in the electrochemical reactivity series than zinc, meaning aluminium has a higher oxidation potential and a greater thermodynamic tendency to remain as oxidized $\text{Al}^{3+}$ cations in aqueous solution. Because zinc is less electropositive than aluminium, zinc cannot displace aluminium from its aqueous salts. Furthermore, attempting chemical displacement in aqueous media would fail because water would reduce preferentially. Carbon reduction cannot be used to extract aluminium from aluminium oxide ($\text{Al}_2\text{O}_3$) because aluminium has an exceptionally high chemical affinity for oxygen, forming an extremely stable ionic lattice with a large negative Gibbs free energy of formation ($\Delta G^\circ_f \approx -1582\text{ kJ/mol}$). Carbon is a weaker reducing agent than aluminium at accessible temperatures; heating $\text{Al}_2\text{O}_3$ with carbon at extreme temperatures produces aluminium carbide ($\text{Al}_4\text{C}_3$) rather than elemental aluminium metal. Consequently, aluminium must be extracted by electrolytic reduction of molten alumina in cryolite.