The Cryolite Bath

Aluminum, Napoleon III's State Cutlery, and the Hall-Héroult Cell

At the Paris Exposition Universelle of 1855, visitors crowded around a velvet-lined glass display case guarding a dozen small silvery ingots resting beside the French Crown Jewels. The metal was aluminum, newly isolated by French chemist Henri Sainte-Claire Deville, who reduced aluminum chloride with metallic sodium. Emperor Napoleon III was so captivated by the metal's lightness—barely one-third the density of steel or silver ($2.70\text{ g/cm}^3$)—that he commissioned a state dinner service of aluminum spoons and forks for his most honored guests at the Tuileries Palace, relegating lesser nobility to eat with gold and silver utensils. Three decades later, in December 1884, American engineers capped the tip of the Washington Monument with a one-hundred-ounce cast pyramid of solid aluminum. Priced at $1.10 per ounce on par with silver, the pyramid was displayed beforehand in the window of Tiffany's jewelry store in New York.

The economic absurdity of treating aluminum as a precious jewelry metal was that aluminum ($Z = 13$) is the single most abundant metallic element in the Earth's crust. Making up 8.2 percent of continental rock by weight, aluminum is more abundant than iron, copper, lead, and tin combined. Every clay bank and bauxite deposit held millions of tons of the metal. Yet unlike iron or copper, which ancient metallurgists smelted in charcoal furnaces, no blast furnace on Earth could tear aluminum free from oxygen.

The barrier is thermodynamic. An aluminum atom possesses three outer valence electrons ($3s^2 3p^1$) that it surrenders completely to oxygen to form a compact $\text{Al}^{3+}$ cation with a small ionic radius ($0.0535\text{ nanometers}$). Packed together with $\text{O}^{2-}$ anions inside the corundum crystal lattice of alumina ($\text{Al}_2\text{O}_3$), those small, triply charged ions generate an electrostatic lattice energy exceeding 15,000 kilojoules per mole and a standard Gibbs free energy of formation of $-1,582\text{ kJ/mol}$. Carbon monoxide cannot reduce alumina at blast-furnace temperatures, and heating alumina with carbon above $2,000^\circ\text{C}$ produces useless yellow aluminum carbide ($\text{Al}_4\text{C}_3$). Nor could nineteenth-century electrochemists simply electrolyze an aqueous aluminum salt solution: water reduces to hydrogen gas at $-0.83\text{ volts}$, long before the cathode reaches the $-1.66\text{ volts}$ required to plate metallic aluminum. Melting pure alumina to electrolyze it without water required $2,072^\circ\text{C}$, a temperature that melted iron retorts and destroyed electrodes.

The breakthrough required finding a non-aqueous, molten inorganic mineral that could dissolve solid alumina like sugar in hot tea at a manageable temperature. Working independently in 1886—Charles Martin Hall inside a woodshed behind his family home in Oberlin, Ohio, and Paul Héroult in a tannery workshop in Gentilly outside Paris—two researchers discovered the exact same chemical key within weeks of each other. Both men dissolved alumina in molten cryolite, sodium hexafluoroaluminate ($\text{Na}_3\text{AlF}_6$), a rare snow-white fluoride mineral mined from a coastal fjord at Ivigtut, Greenland.

Cryolite melts at roughly $1,000^\circ\text{C}$, and adding ten percent alumina depresses the eutectic bath temperature to $960^\circ\text{C}$. Inside the molten fluoride bath, dynamic fluoride ions attack the rigid $\text{Al}-\text{O}$ lattice bridges of alumina, breaking it apart into mobile oxyfluoride complex ions ($\text{Al}_2\text{OF}_6^{2-}$ and $\text{Al}_2\text{O}_2\text{F}_4^{2-}$). When Hall passed direct current from homemade Bunsen battery jars through a carbon-lined crucible on February 23, 1886, $\text{Al}^{3+}$ ions reduced to liquid metal at the cathode while oxide ions reacted with the consumable carbon anode to bubble off as carbon dioxide ($\text{CO}_2$). Crucially, at $960^\circ\text{C}$, liquid aluminum ($2.30\text{ g/cm}^3$) is slightly denser than the molten cryolite bath ($2.10\text{ g/cm}^3$). The newly formed silver droplets sank to the floor of the cell, shielded beneath the fluoride blanket from atmospheric re-oxidation. Within ten years, aluminum plunged from silver's price to thirty cents a pound, turning smelters into permanent geographic companions of massive hydroelectric dams from Niagara Falls to Quebec and Iceland, where fourteen kilowatt-hours of electricity are still poured into every kilogram of metal to sever the aluminum-oxygen bond.