Metallurgy and the Chemistry of Metals Study Notes

Metallurgy and the Chemistry of Metals

Introduction to Metallurgy

• Mineral: A naturally occurring substance with a range of chemical compositions.
• Ore: A mineral deposit concentrated enough for economical recovery of a desired metal.
• Metallurgy: The science and technology of separating metals from their ores and compounding alloys.
• Alloy: A solid solution of two or more metals, or of a metal(s) with one or more nonmetals.

Recovery of Metals from Ores

The recovery of a metal from its ore involves three main steps:

1. Preparation of the ore
2. Production of the metal
3. Purification of the metal

Principal Types of Minerals

• Uncombined Metals: Examples include Ag, Au, Bi, Cu, Pd, Pt
• Carbonates: Examples include $BaCO<em>3$ (witherite), $CaCO</em>3$ (calcite, limestone), $MgCO<em>3$ (magnesite), $CaCO</em>3 \cdot MgCO<em>3$ (dolomite), $PbCO</em>3$ (cerussite), $ZnCO_3$ (smithsonite).
• Halides: Examples include $CaF<em>2$ (fluorite), $NaCl$ (halite), $KCl$ (sylvite), $Na</em>3AlF_6$ (cryolite).
• Oxides: Examples include $Al<em>2O</em>3 \cdot 2H<em>2O$ (bauxite), $Al</em>2O<em>3$ (corundum), $Fe</em>2O<em>3$ (hematite), $Fe</em>3O<em>4$ (magnetite), $Cu</em>2O$ (cuprite), $MnO<em>2$ (pyrolusite), $SnO</em>2$ (cassiterite), $TiO_2$ (rutile), $ZnO$ (zincite).
• Phosphates: Examples include $Ca<em>3(PO</em>4)<em>2$ (phosphate rock), $Ca</em>5(PO<em>4)</em>3OH$ (hydroxyapatite).
• Silicates: Examples include $Be<em>3Al</em>2Si<em>6O</em>{18}$ (beryl), $ZrSiO<em>4$ (zircon), $NaAlSi</em>3O<em>8$ (albite), $Mg</em>3(Si<em>4O</em>{10})(OH)_2$ (talc).
• Sulfides: Examples include $Ag<em>2S$ (argentite), $CdS$ (greenockite), $Cu</em>2S$ (chalcocite), $FeS_2$ (pyrite), $HgS$ (cinnabar), $PbS$ (galena), $ZnS$ (sphalerite).
• Sulfates: Examples include $BaSO<em>4$ (barite), $CaSO</em>4$ (anhydrite), $PbSO<em>4$ (anglesite), $SrSO</em>4$ (celestite), $MgSO<em>4 \cdot 7H</em>2O$ (epsomite).

Metals and Their Best-Known Minerals

• The periodic table highlights where different metals are found and the types of minerals (sulfides, oxides, chlorides, and uncombined) in which they are commonly found.

Production of Metals

• Roasting:
  • $CaCO<em>3(s) \rightarrow CaO(s) + CO</em>2(g)$
  • $2PbS(s) + 3O<em>2(g) \rightarrow 2PbO(s) + 2SO</em>2(g)$
• Chemical Reduction:
  • $TiCl<em>4(g) + 2Mg(l) \rightarrow Ti(s) + 2MgCl</em>2(l)$
  • $Cr<em>2O</em>3(s) + 2Al(s) \rightarrow 2Cr(l) + Al<em>2O</em>3(s)$
  • $WO<em>3(s) + 3H</em>2(g) \rightarrow W(s) + 3H_2O(g)$
• Electrolytic Reduction:
  • $2MO(l) \rightarrow 2M \text{ (at cathode)} + O_2 \text{ (at anode)}$
  • $2MCl(l) \rightarrow 2M \text{ (at cathode)} + Cl_2 \text{ (at anode)}$

Reduction Processes for Common Metals

Blast Furnace for Producing Iron

• Process Overview:
  • A blast furnace is used to reduce iron ore to molten iron.
  • The charge (ore, limestone, coke) descends through the furnace as hot gases rise.
• Reactions at Different Temperatures:
  • 200°C: $3Fe<em>2O</em>3 + CO \rightarrow 2Fe<em>3O</em>4 + CO_2$
  • 700°C: $CaCO<em>3 \rightarrow CaO + CO</em>2$
  • $Fe<em>3O</em>4 + CO \rightarrow 3FeO + CO_2$
  • $C + CO_2 \rightarrow 2CO$
  • 1200°C - 1500°C: $FeO + CO \rightarrow Fe + CO_2$
  • 2000°C: $2C + O_2 \rightarrow 2CO$
• Outputs:
  • Molten iron is collected at the bottom.
  • Molten slag forms.
  • Hot air is blasted into the furnace.

Steel Manufacturing via the Oxygen Process

• Process Overview:
  • The oxygen process is used to manufacture steel from molten iron.
  • Oxygen is blown into the molten iron to oxidize impurities.
  • CaO or $SiO_2$ is added to form slag.

Types of Steel

• A single number indicates the maximum amount of the substance present.

Purification of Metals

• Distillation:
  • $Ni(s) + 4CO(g) \rightleftharpoons Ni(CO)_4(g)$
  • @70°C: $Ni(s) + 4CO(g) \rightarrow Ni(CO)_4(g)$
  • @200°C: $Ni(CO)_4(g) \rightarrow Ni(s) + 4CO(g)$
• Electrolysis:
  • $Cu(s) \text{ (impure)} \rightarrow Cu^{2+}(aq) + 2e^-$
  • $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s) \text{ (pure)}$
• Zone Refining:

Zone Refining of Metals

• A heating coil is used to melt a small zone of a metal rod.
• Impurities concentrate in the molten zone, which is moved along the rod to purify the metal.

Band Theory of Conductivity

• Delocalized electrons move freely through “bands” formed by overlapping molecular orbitals.
• Example: Mg has the electronic configuration $1s^22s^22p^63s^2$ or $[Ne]3s^2$

Energy Gaps Between Valence and Conduction Bands

• Metals: Valence and conduction bands overlap, allowing electrons to move freely.
• Semiconductors: Small energy gap between valence and conduction bands.
• Insulators: Large energy gap between valence and conduction bands.

Semiconductors

• Silicon (Si): $[Ne]3s^23p^2$
• n-type semiconductor:
  • Doped with donor impurities like Phosphorus (P) which has the electronic configuration $[Ne]3s^23p^3$
• p-type semiconductor:
  • Doped with acceptor impurities like Boron (B) which has the electronic configuration $[Ne]3s^23p^1$

Periodic Trends in Metallic Character

• Increasing Metallic Character: Metallic character generally increases down a group and from right to left across a period.
• The alkali metals (Li, Na, K, Rb, Cs) and alkaline earth metals (Be, Mg, Ca, Sr, Ba) are good examples of this trend.

Properties of Alkali Metals

*Refers to the cation $M^+$, where M denotes an alkali metal atom.
The half-reaction is $M^+(aq) + e^- \rightarrow M(s)$.

Reactivity of Alkali Metals

• General reaction: $M \rightarrow M^{+1} + 1e^-$
• Reaction with water: $2M(s) + 2H<em>2O(l) \rightarrow 2MOH(aq) + H</em>2(g)$
• Reaction with oxygen: $2M(s) + O<em>2(g) \rightarrow M</em>2O(s)$
• Reactivity increases down the group.

Properties of Alkaline Earth Metals

*Refers to the cation $M^{2+}$, where M denotes an alkaline earth metal atom.
The half-reaction is $M^{2+}(aq) + 2e^- \rightarrow M(s)$.

Reactivity of Alkaline Earth Metals

• General reaction: $M \rightarrow M^{+2} + 2e^-$
• Reaction with water:
  • Be(s) + 2H2O(l) → No Reaction
  • $Mg(s) + H<em>2O(g) \rightarrow MgO(s) + H</em>2(g)$
  • $M(s) + 2H<em>2O(l) \rightarrow M(OH)</em>2(aq) + H_2(g)$, where M = Ca, Sr, or Ba
• Reactivity increases down the group (excluding Be).

Aluminum Production

• Purification:
  • $Al<em>2O</em>3(s) + 2OH^-(aq) \rightarrow 2AlO<em>2^-(aq) + H</em>2O(l)$
  • $AlO<em>2^-(s) + H</em>3O^+(aq) \rightarrow Al(OH)_3(s)$
  • $2Al(OH)<em>3(s) \rightarrow Al</em>2O<em>3(s) + 3H</em>2O(g)$
• Electrolytic Process:
  • Anode: $3[2O^{2-} \rightarrow O_2(g) + 4e^-]$
  • Cathode: $4[Al^{3+} + 3e^- \rightarrow Al(l)]$
  • Overall: $2Al<em>2O</em>3 \rightarrow 4Al(l) + 3O_2(g)$

Chemistry in Action: Recycling Aluminum

• Recycling aluminum is an important process.