Electron Configuration & Orbital Tools

An electron configuration is the address list of every electron in an atom — which shells, which subshells, which orbitals they live in. Electrons in atoms are not visitors: chemistry is almost entirely the story of the outermost electrons finding lower-energy arrangements, so learning where electrons sit is learning why the periodic table has its shape, why sodium bonds once and carbon four times, and why iron is magnetic. The counting works inward from the outside: neutral atoms have exactly Z electrons, and Z alone determines the whole list.

Shells are the energy levels labeled n = 1, 2, 3...; each shell holds subshells labeled s, p, d, f; each subshell is made of orbitals — regions where an electron is most likely to be found. An orbital is not a track: it is a probability cloud, spherical for s, dumbbell-shaped for p, clover-shaped for most d orbitals (the Shapes panel above renders them). Capacity stacks up quickly:

The four quantum numbers give each electron a unique address: n (shell), l (subshell), m_l (which orbital, from −l to +l), m_s (spin, ±½). Pauli's exclusion principle is simply the statement that no two electrons in an atom share all four.

The Aufbau order 1s → 2s → 2p → 3s → 3p → 4s → 3d... is not a count-down of shells — the first 3d orbital sits above 4s in energy, so potassium (Z=19) and calcium (20) fill 4s before the 3d block begins. Iron (Z=26) lands on [Ar]4s²3d&sup6; with four unpaired electrons, which is exactly why it is strongly paramagnetic — the page computes this in one click. About twenty elements then cheat on the predictions: chromium takes [Ar]4s¹3d⁵ and copper [Ar]4s¹3d¹⁰, trading one 4s electron for a half-filled or completely-filled d subshell, both unusually stable. Ion formation reverses the trick's direction: 4s electrons leave first, so Fe²⁺ is [Ar]3d⁶. The periodic table's shape is this order drawn large: each block is one subshell filling, and chemistry repeats because valence configurations repeat.

Core electrons (the inner shells, buried in the noble-gas shorthand like [Ar]) are spectators; valence electrons — those in the outermost s and p (plus partially-filled d) — do all the bonding. The group number of the main-group elements is the valence count: 1 for Na, 2 for Mg, ... 7 for Cl. This is why the configuration panel's shorthand notation matters: [Ar]4s²3d⁶ tells you at a glance that iron's chemistry lives in that 3d shell. Unpaired valence electrons also decide the magnetic verdict: any unpaired electron (Fe has 4) makes the atom paramagnetic, attracted into a magnetic field; all-paired atoms like Zn ([Ar]3d¹⁰4s²) are diamagnetic and weakly repelled.

• Electrons orbit like planets. An orbital is a probability cloud — the 90% boundary of where the electron is likely to be — and electrons "move" only in the sense that the cloud is not static in time.
• 4s fills first, so it ionizes last. Once 3d electrons arrive they push 4s up in energy, and cations lose 4s first (Fe²⁺, Cu⁺).
• Full subshells cause paramagnetism. The opposite: full means every electron paired, hence diamagnetic — Zn ([Ar]3d¹⁰4s²) is diamagnetic, while atomic Cu keeps its 4s¹ electron unpaired and is paramagnetic, exactly as the page's Cu sample reports.
• Half-filled shells are magic for everyone. The stability bonus applies to degenerate orbitals (like the five 3d); s, p, d, f always fill by the same three rules — there is no separate "quantum rule" per element.

Related tools: Periodic Table (see the blocks the filling order creates), Bonding & VSEPR (what those valence electrons build), and Nuclear Chemistry (the nucleus that Z counts).