Pyrophoric Iron
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The following article was originally published in the journal for educators Chemia w Szkole (eng. Chemistry in School) (1/2024):
How to Ignite Iron?
Iron Fe is a metal known since ancient times. In its pure form, it is a shiny, silvery, fairly hard, and relatively high-melting metal. Its melting point is 1534.85°C (2794.73°F). Iron is brittle, which is why today it is most commonly used in the form of alloys with carbon C — cast iron and steel — as well as with other metals. These alloys are mechanically stronger than pure iron and exhibit other desirable properties [1] [2]. Iron is also a trace element essential for life, as it is a key component of hemoglobin [3].
Chemically, iron belongs to the group of relatively reactive metals. It readily reacts with acids, forming corresponding salts, and can oxidize with atmospheric oxygen — even at room temperature.
We know that combustion is a rapid form of oxidation. So, is it possible to actually set iron on fire?
Let’s Try It
Take any iron or steel object and heat it intensely in a burner flame. After a while, a thin layer of oxide will form on the metal’s surface — but you certainly won’t achieve ignition. However, if you toss a pinch of fine steel filings into the flame, you’ll see beautiful, sparkling flashes. These are rapidly burning iron particles. Similarly, you can ignite steel or iron wool, commonly used for polishing floors or brass objects. It consists of extremely thin filaments that can easily catch fire, producing a spectacular display of sparks while releasing a significant amount of heat (Photo 1).
Finely powdering the iron increases its surface area, which improves access to oxygen from the air and allows the particles to heat more efficiently — triggering rapid oxidation. To burn iron, then, you simply need to grind it finely enough. It turns out that iron with a sufficiently developed surface becomes pyrophoric, meaning it can spontaneously ignite at room temperature.
Materials Needed
This time, all you need is one chemical compound: iron(II) oxalate FeC2O4. This compound typically appears as a bright yellow powder (Photo 2).
You can, of course, purchase iron(II) oxalate from a chemical supplier, but preparing it yourself is simple enough to make it a rewarding experiment in your own lab. For this synthesis, you’ll need just two substances:
- oxalic acid C2H2O4•2H2O,
- iron(II) sulfate heptahydrate FeSO4•7H2O.
Oxalic acid is the simplest dicarboxylic acid, and its dihydrate form appears as a white crystalline powder. Iron(II) sulfate heptahydrate is typically found as greenish crystals. Other well-soluble iron(II) salts can also be used. Neither substance is highly toxic, although oxalic acid is irritating and may have harmful effects if mishandled.
You’ll need to prepare concentrated solutions of both oxalic acid and iron(II) sulfate, and then mix them together. The reaction proceeds faster when the mixture is gently heated; the liquid becomes cloudy, and a yellow precipitate of poorly soluble iron(II) oxalate begins to form and settle at the bottom. The reaction can be written as:
The precipitate should be washed several times by decantation, then filtered and thoroughly dried. It's best to dry it at room temperature, as excessive heat may decompose the compound.
Demonstration
Add a small amount of thoroughly dried iron(II) oxalate to a test tube. Don’t use too much — just enough to form a 1–2 cm (about 0.4–0.8 in) layer of powder (Photo 3).
Now heat the tube using a spirit or gas burner (Photo 4).
After a short while, you’ll notice that water vapor begins to rise from the solid. This vapor condenses on the cooler upper part of the test tube. It’s best to gently heat this area too, to prevent condensation buildup.
Continue heating until the entire contents of the tube turn into a grayish-black powder (Photo 5).
Now, plug the test tube with a small piece of glass wool and, while it is still warm, place it over a fireproof surface (Photo 6A). In a darkened room, open the test tube and pour the black powder out from a height. Upon contact with air, the substance ignites, creating a dazzling cascade of orange-white sparks (Photo 6B).
Explanation
Iron(II) oxalate is stable at room temperature but decomposes easily upon heating. This decomposition produces water vapor, carbon dioxide, fine iron powder, and iron(II) oxide. The water escapes as steam, while a pyrophoric mixture of elemental iron and iron(II) oxide remains in the test tube. It doesn’t ignite immediately because the generated carbon dioxide displaces the air in the tube. Once the powder is poured out and exposed to atmospheric oxygen, the following reaction occurs:
The iron and iron(II) oxide oxidize into iron(III) oxide, releasing energy in the form of light and heat — and creating the eye-catching display of sparks.
References:
- [1] Bielański A., Chemia ogólna i nieorganiczna, PWN, Warszawa, 1981, pp. 534 back
- [2] Dobrzański L.A., Metaloznawstwo opisowe stopów żelaza, Wydawnictwo Politechniki Śląskiej, Gliwice, 2007, pp. 13-15 back
- [3] Kabata J., Ochrem B., Hellmann A., Badania laboratoryjne i morfologiczne, in: Gajewski P. (red.), Interna Szczeklika, Medycyna Praktyczna, Polski Instytut Evidence Based Medicine, 2020 back
All photographs and illustrations were created by the author.
Addendum
The result of the experiment can be seen in the video:
Marek Ples