Electron Config of Iron

Iron has atomic number 26, meaning it has 26 electrons to arrange across its orbitals. Its ground-state electron configuration is:

Full notation: `1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²`

Shorthand notation: `[Ar] 3d⁶ 4s²`

This configuration places Iron in the D-block of the periodic table — Period 4, Group 8. The last subshell filled (the d subshell) determines its block.

SPDF notation tells you exactly: which subshell each electron occupies, how many electrons are in it, and the energy level of each group. This is far more detail than the simpler Bohr model, which only shows shell totals.

The Aufbau (building-up) principle states electrons fill the lowest available energy subshell first. For Iron (Z=26), the filling stops at the 4s² subshell.

Standard Aufbau sequence:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

After filling, Iron's configuration ends at 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s², with 8 valence electrons in its outermost subshell. Note: Iron is a D-block element, so watch for possible Aufbau anomalies driven by extra stability of half-filled or fully-filled d subshells.

The orbital diagram of Iron expands the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² into individual orbital boxes:

- Each s subshell holds max 2 electrons (1 orbital)

- Each p subshell holds max 6 electrons (3 orbitals)

- Each d subshell holds max 10 electrons (5 orbitals)

- Each f subshell holds max 14 electrons (7 orbitals)

Hund's Rule dictates that within any subshell, electrons fill each orbital singly (spin up ↑) before pairing. This avoids electron–electron repulsion. Iron's D-block placement confirms its last orbitals are d type.

The interactive diagram above shows Iron's complete subshell breakdown with orbital boxes for every energy level.

Follow these steps to write Iron's electron configuration from scratch:

Step 1: Identify the atomic number: Z = 26 — this is the total number of electrons to place.

Step 2: Follow the Aufbau sequence, filling the lowest energy subshells first:

> 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → ...

Step 3: Apply Hund's Rule inside each subshell — one electron per orbital before pairing begins.

Step 4: Apply the Pauli Exclusion Principle — each orbital holds at most 2 electrons with opposite spins.

Step 5: After filling all 26 electrons, your result should match:

> 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

Shorthand: Replace the preceding noble gas core with its symbol:

> [Ar] 3d⁶ 4s²

⚠️ Common mistake: Iron is a d-block element. Verify your d-subshell count carefully — anomalies from expected Aufbau order are possible.

Iron has 8 valence electrons — the electrons in its highest occupied principal energy level.

As a D-block element, Iron's valence electrons reside in d orbitals and d/f orbitals. These are the only electrons involved in chemical bonding.

| Block | Type | Max Valence e⁻ |

|---|---|---|

| s-block | Groups 1–2 | 1–2 |

| p-block | Groups 13–18 | 3–8 |

| d-block | Groups 3–12 | up to 10 |

| f-block | Lanthanides/Actinides | up to 14 |

Iron sits in this table as a d-block element with 8 valence electrons.

→ See Iron's valence electrons in the Bohr model for the shell-based view.

→ Electronegativity of Iron — how strongly it attracts these electrons.