The Electrolytic Crown

How Two Twenty-Two-Year-Olds Unlocked the Crust's Most Abundant Metal

Aluminum makes up $8.2\text{ percent}$ of the Earth's continental crust by weight, making it the most abundant metallic element on the planet. Yet throughout the mid-nineteenth century, metallic aluminum was rarer and costlier than silver. At the state banquets of French Emperor Napoleon III in the 1850s, lesser dukes ate with solid gold forks, while the emperor and his honored guests dined with lightweight spoons of aluminum. When engineers capped the Washington Monument in December 1884 with a 2.8-kilogram pyramid of cast aluminum, Tiffany's jewelers displayed the apex in their New York window so crowds could step over it.

Why did the crust's most common metal resist the blast furnaces that poured out millions of tons of iron? The bottleneck was aluminum's fierce affinity for oxygen. Every aluminum atom in bauxite ore is locked to oxygen inside a crystal lattice of alumina ($\text{Al}_2\text{O}_3$), with a Gibbs free energy of formation of $-1,582\text{ kilojoules per mole}$. That bond energy is nearly double the stability of iron oxide. Heating alumina with carbon charcoal in a blast furnace fails because carbon cannot strip oxygen from aluminum below $2,000^\circ\text{C}$. At that extreme heat, reduced aluminum vaporizes and forms useless aluminum carbide.

Chemists could isolate small nuggets of aluminum only by reducing volatile aluminum chloride ($\text{AlCl}_3$) with metallic potassium or sodium. Friedrich Wöhler pioneered that expensive batch reaction in 1827, and Henri Sainte-Claire Deville scaled it up in 1854. Direct electrical splitting appeared equally blocked. Passing an electric current through an aqueous solution of aluminum salts never deposits aluminum at the cathode. Instead, water molecules split first and bubble away as hydrogen gas. Conversely, electrolyzing pure molten alumina without water was impossible because solid $\text{Al}_2\text{O}_3$ melts at $2,072^\circ\text{C}$, a temperature that destroyed crucibles and electrodes.

To split alumina with electricity, metallurgists needed a molten salt bath that could dissolve $\text{Al}_2\text{O}_3$ at a moderate temperature the way warm water dissolves table sugar. In 1886, working independently on opposite sides of the Atlantic, Charles Martin Hall in a woodshed in Oberlin, Ohio, and Paul Héroult in a tannery shed in Gentilly outside Paris discovered the identical chemical key within two months of one another.

Both inventors turned to cryolite ($\text{Na}_3\text{AlF}_6$, sodium hexafluoroaluminate), a white fluoride mineral mined at Ivigtut on the coast of Greenland. Pure cryolite melts at $1,012^\circ\text{C}$. Dissolving five to ten percent refined alumina powder into the molten mineral lowers the eutectic melting point of the bath to $960^\circ\text{C}$. Because fluoride ions in cryolite have a higher decomposition voltage than aluminum-oxygen bonds, an electric current passing through the $960^\circ\text{C}$ bath ignores the cryolite solvent and splits only the dissolved alumina.

On February 23, 1886, using a carbon-lined crucible connected to Bunsen battery cells, Hall cooled his cryolite melt and hammered out shiny aluminum globules. Two months later, on April 23, 1886, Héroult received a French patent for the identical carbon-anode cryolite cell. Inside a Hall-Héroult smelting pot, direct current at four to five volts flows from consumable carbon anodes through the cryolite bath to the carbon floor. Oxygen ions migrate to the top carbon anodes and burn off as carbon dioxide ($\text{2Al}_2\text{O}_3 + \text{3C} \to \text{4Al} + \text{3CO}_2$). Meanwhile, $\text{Al}^{3+}$ cations migrate to the bottom cathode to form liquid aluminum. Because molten aluminum at $960^\circ\text{C}$ is denser ($2.3\text{ g/cm}^3$) than the cryolite bath above it ($2.1\text{ g/cm}^3$), the liquid metal sinks to the cell floor, shielded from air by the floating salt blanket.

The mass $m$ of metallic aluminum deposited on the pot floor obeys Michael Faraday's first law of electrolysis:

Here $Q = I \cdot t$ is the total electrical charge in coulombs delivered by current $I$ over time $t$, $M$ is the molar mass of aluminum ($26.98\text{ g/mol}$), $F$ is the Faraday constant ($96,485\text{ C/mol}$), and $z = 3$ is the valence electrons required to reduce each $\text{Al}^{3+}$ ion. Because each light aluminum atom demands three electrons, smelting one metric ton of aluminum consumes roughly $14,000\text{ kilowatt-hours}$ of electricity. Within fifty years of Hall and Héroult's 1886 molten-cryolite cells, the price of commercial aluminum fell from twelve dollars per pound to twenty cents per pound while global production rose from a few metric tons a year to hundreds of thousands of tons.