Why does iron lose its 4s electrons before its 3d electrons?
Iron’s ground state is [Ar] 3d⁶ 4s². The 4s subshell filled before 3d, following the Madelung ordering, and the filling diagram below shows it.
So when iron forms Fe²⁺, the expectation is that the last electrons in should be the first out, giving [Ar] 3d⁴ 4s². The actual configuration of Fe²⁺ is [Ar] 3d⁶ — the 4s electrons leave, and the 3d electrons stay.
Students find this contradictory, and they are right to. The resolution is that the Madelung rule describes the order in which subshells are filled across the periodic table, not the relative energies of the orbitals in any particular atom.
The 4s orbital is lower in energy than 3d at the point where 4s begins filling — for potassium and calcium, which have no 3d electrons at all. Once 3d starts filling, the picture changes. Each added 3d electron screens the nucleus poorly from the other 3d electrons but pulls the whole 3d set downward in energy relative to 4s. By the time you reach iron, 3d sits below 4s.
The 4s orbital also extends further from the nucleus. Its electrons are more weakly held and more exposed, so they are the ones that go when the atom ionises.
The general rule for all transition metals follows: the outermost s electrons are lost first. Fe³⁺ is [Ar] 3d⁵ — a half-filled d subshell, which is part of why iron(III) is so stable and so common in nature.
Why does stellar fusion stop at iron?
The standard statement is that iron-56 has the highest binding energy per nucleon, so fusing it releases no energy and stellar nucleosynthesis halts. The statement is very nearly right, and the small correction is interesting.
The nucleus with the highest binding energy per nucleon is nickel-62, at about 8.794 MeV. Iron-58 is next, and iron-56 is third at about 8.790 MeV. The differences are tiny, and the group is usually called the iron peak.
Iron-56 nevertheless dominates the end products, for a reason that is about production rather than stability. Silicon burning in a massive star proceeds largely through nickel-56, which is doubly magic and easy to assemble from alpha particles. Nickel-56 is radioactive: it decays to cobalt-56 and then to iron-56, with half-lives of about 6 and 77 days. The gamma rays from that decay chain are what powers the visible light curve of a type Ia supernova for weeks after the explosion.
So the iron in your blood and in the Earth’s core was mostly nickel when it was made.
Past the iron peak, fusion costs energy rather than releasing it. A star’s core can no longer support itself against gravity, and for a sufficiently massive star the collapse and rebound is the supernova. Elements heavier than iron are built by neutron capture, not fusion — mostly during that explosion and in neutron star mergers.
Iron is the most abundant element on Earth by mass, largely because the core is mostly iron and nickel, and the fourth most abundant in the crust.
How do you read iron’s electron configuration?
Iron’s ground state is [Ar] 3d⁶ 4s²: an argon core of 18 electrons, six in 3d, two in 4s. The electrons per shell are 2, 8, 14, 2.
Six electrons across five 3d orbitals means one orbital holds a pair and four hold single electrons, by Hund’s rule. Those four unpaired electrons are the origin of iron’s magnetism — each atom carries a magnetic moment, and in the metal the moments of neighbouring atoms align spontaneously through exchange coupling.
Iron is one of only three elements that are ferromagnetic at room temperature, alongside cobalt and nickel, and it is by far the most important industrially. Heat it above 770 °C and thermal agitation overcomes the coupling; the alignment collapses and iron becomes ordinary paramagnetic material. That temperature is the Curie point.
The partly filled d subshell also explains iron’s variable oxidation states. Iron commonly forms +2 and +3, and less commonly +4, +5 and +6. Compounds with different d-electron counts absorb different wavelengths, which is why iron compounds span such a range of colours — from the pale green of iron(II) sulfate to the deep red-brown of iron(III) oxide.
The ground state term symbol is ⁵D₄, and the spin multiplicity of 5 confirms the four unpaired electrons.
Why is iron the metal civilisation was built on?
Copper and bronze came first, by a wide margin, and the reason is temperature. Copper smelts around 1085 °C; iron melts at 1538 °C. Early furnaces could not reach that.
Iron working began instead with the bloomery, which never melts the ore at all. It reduces iron oxide to a spongy mass of metal and slag at around 1200 °C, which is then hammered repeatedly to drive out the slag. The process is laborious, and the product is inferior to bronze unless carburised.
What made iron win was abundance. Bronze requires copper and tin, and tin deposits are rare and geographically concentrated — Bronze Age economies depended on long, fragile trade routes. Iron ore is almost everywhere. Once the technique spread, iron could be made locally by anyone, and the bottleneck moved from access to raw material to skill.
Much of that ore is the banded iron formation laid down when oxygen first accumulated in Earth’s atmosphere: dissolved iron in the anoxic oceans reacted with the new oxygen and precipitated out. The iron in a modern steel beam was put there by cyanobacteria roughly two billion years ago.
What does iron do in the body?
An adult human contains around four grams of iron, and most of it is in haemoglobin, the protein that carries oxygen in red blood cells.
Each haemoglobin molecule holds four haem groups, and each haem group holds a single iron atom at the centre of a flat organic ring. Oxygen binds directly to that iron. What makes the arrangement work is that iron in haem binds oxygen reversibly — firmly enough to pick it up in the lungs, loosely enough to release it in tissue. Free iron would simply be oxidised and stay that way; the protein scaffold tunes its behaviour.
Binding also changes haemoglobin’s shape, making the remaining sites bind more readily. This cooperativity is why haemoglobin loads and unloads oxygen over a narrow pressure range rather than gradually, and it is one of the most studied examples of protein allostery.
Iron deficiency is the most common nutritional deficiency worldwide, affecting an estimated two billion people. Too little iron means too little haemoglobin, and the resulting anaemia presents as fatigue and breathlessness.
Excess iron is equally dangerous and harder to fix, because humans have no mechanism for excreting it. Haemochromatosis, a common genetic disorder, causes iron to accumulate in the liver, heart and pancreas over decades. Treatment is still, essentially, bloodletting.