Electron configuration of Iron (Fe)

Configuration at a glance

Iron has atomic number 26, so a neutral atom has 26 electrons. The configuration 1s2 2s2 2p6 3s2 3p6 3d6 4s2 is a compact record of where those electrons sit by subshell. The number before each letter is the principal energy level, the letter names the subshell, and the superscript tells you how many electrons are assigned there.

The shorthand [Ar] 3d6 4s2 replaces the filled inner core with the previous noble gas. That notation keeps attention on the electrons that usually explain bonding, ions, and periodic trends.

Iron follows the standard filling order

Iron is a period 4 transition metal that follows the standard Aufbau sequence: 1s2 2s2 2p6 3s2 3p6 4s2 3d6. Note that 4s fills before 3d, but when iron forms ions, the 4s electrons are lost first. This ordering is critical for correctly predicting Fe2+ and Fe3+ configurations.

Step-by-step filling walkthrough: 1. Start with atomic number 26. 2. Follow the Aufbau sequence: fill 1s, 2s, 2p, 3s, 3p, then continue through 4s, 3d, 4p in order (or 5s, 4d, 5p for period 5). 3. Keep a running total of electrons after each subshell. 4. Write the noble-gas shorthand by replacing the filled core with the previous noble gas in brackets. 5. Verify: add every superscript and confirm the total equals 26.

Real-world applications

Iron's electron configuration directly enables its role in biology and industry. The 3d6 subshell can lose different numbers of electrons, giving multiple stable oxidation states:

• Hemoglobin contains iron that binds oxygen via its 3d/4s orbitals. Fe2+ in heme reversibly binds O2, enabling blood to transport oxygen.
• Steel production relies on iron's ability to form alloys with carbon.
• Iron catalysts are used in the Haber-Bosch process for ammonia production.
• Permanent magnets (Fe3O4, magnetite) depend on unpaired 3d electrons.

Common ions: Fe2+ ([Ar] 3d6) and Fe3+ ([Ar] 3d5). The Fe3+ configuration has a half-filled 3d subshell (3d5), providing extra stability.

Common exam questions

- Q: Write the electron configurations of Fe2+ and Fe3+. A: Fe2+ = [Ar] 3d6 (lose the 4s2 electrons first). Fe3+ = [Ar] 3d5 (lose 4s2 then one 3d electron). Fe3+ with 3d5 has a half-filled d-subshell, which is especially stable.

- Q: How many unpaired electrons does neutral iron have? A: Four. In 3d6, five d orbitals fill with one electron each before pairing (Hund's rule). That gives 4 unpaired in 3d plus 2 paired in 4s.

- Q: Why is Fe3+ stable despite removing three electrons? A: The 3d5 half-filled subshell provides extra exchange energy stability, similar to the d10 full-subshell stability seen in copper.

How to read the subshells

Read the configuration in filling-order language first, not as a simple shell list. For transition metals, the s subshell of the outermost shell fills before the d subshell of the previous shell, but ions always lose the outer s electrons first. Keeping these two ideas separate prevents many wrong answers on ion configuration questions.

For Iron, the valence clue is 4s2. The element sits in period 3 and belongs to the d-block.

Orbital diagram

An orbital diagram turns each subshell into boxes and arrows. Each orbital receives one electron before pairing when several equal-energy orbitals are available (Hund's rule).

For Iron, draw the diagram by filling each subshell in Aufbau order and applying Hund's rule within p and d subshells. Count the unpaired electrons — this determines magnetic properties and influences common oxidation states.

Common mistakes

The first mistake is counting the superscripts incorrectly. Add every superscript and confirm the total equals the atomic number for a neutral atom. The second mistake is writing orbitals in shell order (e.g., 3d before 4s) and ignoring the filling sequence. The third mistake is forgetting that when writing ion configurations, the outer s electrons are removed before the inner d electrons, even though the s subshell filled first.

Element-specific practice problems

Problem 1: Write the full electron configuration for neutral Iron and verify the electron count against 26.

Problem 2: Convert the full configuration to noble-gas shorthand notation.

Problem 3: Predict the electron configurations of Fe2+ and Fe3+. Explain why Fe3+ (3d5) is more stable than Fe2+ (3d6) despite having fewer electrons.

Problem 4: Iron in hemoglobin is Fe2+. What would happen to the configuration if it oxidised to Fe3+ (methemoglobin)?

Problem 5: Compare the configuration of Iron with Manganese (Z=25) and Cobalt (Z=27). How does the d-electron count change across these three neighbours?