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Railway tracks are long steel rails that must be joined on site, often far from any power supply. One answer, still in use, is to pour molten iron made on the spot by a chemical reaction between aluminium and iron oxide. Hans Goldschmidt turned that violent reaction into a controlled tool. This post covers who he was, how he tamed the reaction, a worked thermochemistry example, and what students can learn from his approach.
Who was Hans Goldschmidt?
The Neue Deutsche Biographie, the standard German biographical dictionary, gives his full name as Johann (Hans) Wilhelm Goldschmidt, born 18 January 1861 in Berlin and died 21 May 1923 in Baden-Baden. His father had founded the chemical firm Th. Goldschmidt in Berlin in 1847. Hans studied at Berlin, Leipzig, Strasbourg and Heidelberg, where he took his doctorate in 1886. Among his teachers was Robert Bunsen. In 1888 he joined the family firm, which moved to Essen in 1889, and he led its science and technology. He also developed a process for recovering tin from tinplate scrap by electrolysis.
What problem did aluminothermy solve?
Most metals are made by reducing their oxides with carbon. The Neue Deutsche Biographie explains the drawback: carbon left in the metal harms the strength of high-grade steels and alloys. Aluminium can do the reducing instead, but earlier attempts had run out of control explosively.
Goldschmidt's idea, in the late 1890s, was to light the mixture of metal oxide and aluminium powder at one small spot with an ignition charge and let the reaction spread through the mix on its own, with no outside heating. The heat from burning aluminium gave a complete reaction and a clean separation of slag and metal. By this route, refractory metals such as chromium, manganese and molybdenum were first made pure. He published it in 1898, in Liebigs Annalen and in the Zeitschrift für Elektrochemie under the title (in translation) "A new method of producing high temperatures and preparing refractory carbon-free metals".
How did it become a way to weld rails?
The same biography says the thermite process was first used to weld the rails of electric tramways and then became a standard way to weld railway rails and steel structures. The NCERT Class 10 chapter on metals and non-metals gives the reaction and notes that it joins railway tracks and cracked machine parts:
Fe₂O₃(s) + 2Al(s) → 2Fe(l) + Al₂O₃(s) + heat
A 1908 manual from the Goldschmidt Thermit Company states that the mass reaches 5400 °F in less than 30 seconds, and that the liquid steel produced is one-half of the original thermite by weight.
Worked example: Hess's law for thermite
Standard enthalpies of formation from the NIST Chemistry WebBook: Al₂O₃ −1675.7 kJ/mol and Fe₂O₃ (hematite) −825.5 kJ/mol. Elements in their standard states are zero.
Reaction enthalpy. ΔH = ΣΔfH(products) − ΣΔfH(reactants) = (−1675.7 + 0) − (−825.5 + 0) = −850.2 kJ per mole of Fe₂O₃.
Mass of mixture. M(Fe₂O₃) = 2(55.85) + 3(16.00) = 159.69 g/mol; 2Al = 2 × 26.98 = 53.96 g. Total 213.65 g.
Energy per kilogram. 850.2 kJ / 0.21365 kg = 3,980 kJ/kg, about 4.0 MJ per kilogram of mixture.
Iron yield. 2Fe = 111.70 g out of 213.65 g, or 52 per cent, in line with the manual's "one-half by weight".
Mixture for a 10 kg pour. 10 / 0.523 = 19.1 kg of stoichiometric thermite.
Temperature in Celsius. (5400 − 32) × 5/9 = 2,982 °C, about 3,000 °C.
What mistakes do students make with thermite problems?
Reversing the subtraction and getting +850 kJ for an obviously exothermic reaction.
Forgetting that 2 mol of aluminium go with each mole of Fe₂O₃ when finding mass.
Treating ΔfH of Fe or Al as non-zero. Elements in their standard states have zero enthalpy of formation.
Hess's law and enthalpies of formation are in Unit 4 of chemistry in the JEE (Main) 2026 syllabus; check your year's syllabus. Our posts on Hess's law problems and metals and non-metals for Class 10 go further. If you want a step-by-step check of an enthalpy calculation, the Fermi Problem Library has worked problems with a step-by-step solver that explains the reasoning rather than printing an answer.
This reaction reaches about 3,000 °C and throws molten metal. It is for trained crews with proper equipment, not for home experiments.
What can a student take from how he worked?
Control beats raw power. Starting the reaction at one spot turned an explosion into a process.
Ask what the user needs. Carbon-free metals and on-site welding were real industrial problems.
Check your theory against practice. The stoichiometric iron yield matches what welders measured.
Next step: repeat the worked example for 3MnO₂ + 4Al → 3Mn + 2Al₂O₃ using enthalpies of formation you look up.
