Iron reacts with a wide range of elements, including oxygen, the halogens (fluorine, chlorine, bromine), sulfur, carbon, nitrogen, and the hydrogen in acids and steam. It sits in the middle of the reactivity series, above copper but below more reactive metals like zinc and magnesium, which means it readily gives up electrons to many nonmetals but won’t displace stronger metals from their compounds.
Oxygen and Moisture
The most familiar reaction iron undergoes is rusting. When iron is exposed to both oxygen and water, it forms iron hydroxide, the flaky reddish-brown substance you see on old nails and car panels. The process requires moisture to get started, which is why iron left in dry conditions can stay rust-free for years. In simple terms, four atoms of iron combine with three molecules of oxygen and six molecules of water to produce iron(III) hydroxide.
At higher temperatures, the reaction changes character. When iron is heated and exposed to steam rather than liquid water, it produces a mixed iron oxide (a black, magnetic compound called magnetite) and releases hydrogen gas. This reaction kicks in around 570 °C and was historically used in industrial hydrogen production. The takeaway: iron reacts with oxygen under almost any conditions, but temperature and moisture determine exactly what you get.
Halogens: Fluorine, Chlorine, and Bromine
Iron reacts vigorously with the halogen family. Cold iron wool will spontaneously burn in fluorine gas without any external heat, producing iron(III) fluoride, a white or pale green powder. Chlorine requires a bit more energy. When chlorine gas contacts hot iron, it forms iron(III) chloride, which appears as black crystals in its pure, dry form. Even a trace of moisture will turn those crystals reddish-brown.
Bromine follows a similar pattern, reacting with iron to produce iron(III) bromide. The trend across the halogens is consistent: fluorine reacts most aggressively, chlorine next, and bromine least. In each case, iron loses three electrons per atom, reaching the +3 oxidation state. Iodine, the least reactive halogen, can also react with iron but typically only forms iron(II) iodide under normal conditions because iodine is too weak an oxidizer to push iron to the +3 state.
Sulfur
Iron and sulfur react when heated together to form iron sulfide, a compound with the formula FeS. You may have seen this demonstrated in a chemistry class: iron filings and sulfur powder are mixed and heated in a test tube until the bottom glows red-hot. Once the reaction ignites, it releases enough energy to sustain itself without further heating. The product, iron sulfide, is a dark, brittle solid that looks nothing like either of the starting materials. In nature, this compound shows up as the mineral pyrrhotite. A related iron-sulfur mineral, pyrite (FeS₂), forms under different geological conditions and is sometimes called “fool’s gold.”
Acids
Iron dissolves in most common acids, producing a salt and hydrogen gas. Drop a piece of iron into dilute hydrochloric acid or dilute sulfuric acid and you’ll see bubbles of hydrogen forming on the metal’s surface. In dilute sulfuric acid, iron is oxidized to its +2 state, producing iron(II) sulfate and hydrogen gas. The acid’s negative ions are essentially bystanders in this reaction; swapping sulfuric acid for hydrochloric acid gives the same type of result, just with iron(II) chloride instead.
Concentrated sulfuric acid tells a different story. Because concentrated sulfate ions are strong enough to act as oxidizers on their own, the reaction goes further. Some of the iron is pushed from the +2 state all the way to the +3 state, and sulfur dioxide gas is released alongside water. So concentration matters: dilute acid produces hydrogen gas and a simple iron(II) salt, while concentrated acid generates a mix of iron(II) and iron(III) salts plus toxic sulfur dioxide fumes.
One important exception is concentrated nitric acid. Rather than dissolving iron, it forms a thin, protective oxide layer on the surface that blocks further reaction. This phenomenon, called passivation, is actually useful in industrial settings where iron needs to resist corrosion.
Carbon
Iron doesn’t react with carbon the way it does with oxygen or chlorine. Instead, carbon atoms are small enough to slip into the gaps between iron atoms in the crystal lattice, forming what chemists call an interstitial solid solution. This is the basis of steel. The amount of carbon dissolved in the iron determines the steel’s properties: low-carbon steel (under about 0.3% carbon) is soft and easy to shape, medium-carbon steel is stronger, and high-carbon steel (up to about 2%) is hard but brittle. Above roughly 2% carbon by weight, the material starts behaving like cast iron, which melts at a noticeably lower temperature than pure iron because the iron-carbon mixture reaches a eutectic point where melting begins more easily.
Nitrogen
Iron reacts with nitrogen, but only under specific conditions. At room temperature, nothing happens. In industrial nitriding, iron or steel parts are exposed to ammonia gas at temperatures around 590 °C. The ammonia breaks down and releases nitrogen atoms, which diffuse into the iron surface and form iron nitride compounds. These nitride layers are extremely hard, typically 10 to 25 micrometers thick, and dramatically improve a part’s resistance to wear and fatigue. This is a surface treatment rather than a bulk reaction, so it changes the skin of the metal while leaving the core unchanged.
Phosphorus and Silicon
Iron also interacts with phosphorus and silicon, though these reactions are most relevant in metallurgy rather than everyday chemistry. During smelting, phosphorus atoms wedge themselves into interstitial positions at the boundaries between iron grains. This makes the iron brittle, which is why steelmakers work hard to remove phosphorus during refining. Silicon behaves differently: it substitutes directly for iron atoms in the crystal structure rather than squeezing between them. Small amounts of silicon (typically 1 to 3%) are deliberately added to some steels to improve strength and magnetic properties, as in the electrical steel used in transformer cores.
Water and Steam
Pure liquid water reacts with iron very slowly at room temperature, which is why iron pipes can last decades if kept free of dissolved oxygen. The real trouble starts when oxygen is also present, which brings us back to rusting. Steam, however, is a different matter. When steam passes over iron heated above 570 °C, the reaction proceeds in stages. Iron first converts to wüstite (an intermediate oxide), then to magnetite, releasing hydrogen gas at each step. Below 570 °C, the intermediate stage is skipped and magnetite forms directly. This steam-iron reaction was one of the earliest industrial methods for producing hydrogen and is still studied for clean energy applications today.
Where Iron Sits in the Reactivity Series
Iron’s position in the reactivity series explains all of these reactions. It sits above hydrogen, which is why it can displace hydrogen from acids and steam. It sits above copper, which is why an iron nail dipped in a copper sulfate solution will develop a coating of metallic copper. But it sits below zinc, magnesium, and aluminum, which means those metals can protect iron from corrosion by reacting preferentially. This is the principle behind galvanizing: coating iron with zinc so the zinc corrodes first, sacrificing itself to keep the iron intact.