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Where are the electrons in an atom?

The A-level picture: subshells, orbitals, the filling order and its famous exceptions.

  1. A copper wire1 mm
  2. One copper crystal10 µm
  3. A row of atoms1 nm
  4. One atom0.3 nm
Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1
What the simulation shows, in words

Choose an element. The orbital picture shows each subshell as a row of boxes, lowest energy at the bottom: one box for s, three for p, five for d, seven for f. Each electron is an arrow. Arrows go one per box with the same spin first, then pair up with opposite spins. Chromium and copper show their exceptions, with one electron in 4s. You can switch to the shell picture at any time.

Shells split into subshells

Each shell is made of subshells, labelled s, p, d and f. Each subshell is made of orbitals, and each orbital holds at most two electrons.

Subshell Orbitals Electrons when full First appears in shell
s 1 2 1
p 3 6 2
d 5 10 3
f 7 14 4

So shell 3 (3s, 3p, 3d) can hold 2 + 6 + 10 = 18. But the subshells don’t all have the same energy. In potassium and calcium, 4s is lower in energy than 3d, so it fills first. That is why potassium is [Ar] 4s¹, and why the transition metals only start at scandium.

Three rules for filling

  1. Lowest energy first (the Aufbau principle). The usual order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p.
  2. At most two electrons per orbital, with opposite spins (the Pauli exclusion principle). That is the up and down arrow in each box.
  3. Spread out before pairing up (Hund’s rule). In a subshell with several orbitals, electrons go one into each orbital, all with the same spin, before any pair up. Look at nitrogen: three p electrons, one in each p box.

The exceptions are real

Chromium is [Ar] 3d⁵ 4s¹, not 3d⁴ 4s². Copper is [Ar] 3d¹⁰ 4s¹, not 3d⁹ 4s². In both, moving one electron from 4s to 3d gives the atom a lower total energy, because a half-full or full d subshell is especially stable. These are not textbook conventions: they are measured, and the simulation uses the NIST data, so they appear by themselves. Further down the table there are more surprises, such as palladium with no 5s electrons at all, and lawrencium, whose outermost electron is in 7p rather than 6d.

Losing electrons: 4s goes first

Once 3d has electrons in it, the 3d orbitals drop below 4s in energy. So when a transition metal atom forms a positive ion, it loses its 4s electrons first. Iron is [Ar] 3d⁶ 4s², so Fe²⁺ is [Ar] 3d⁶ and Fe³⁺ is [Ar] 3d⁵. Writing Fe²⁺ as [Ar] 3d⁴ 4s² is the most common mistake in this topic.

What an orbital looks like

An orbital is not a path. It describes where an electron is likely to be found. An s orbital is a sphere; a p orbital is two lobes either side of the nucleus. The orbital viewer draws real hydrogen orbitals, point by point, from the equations.

The blocks of the periodic table follow the subshell being filled: groups 1 and 2 (and helium) are the s block, groups 13 to 18 (apart from helium) the p block, the transition metals the d block, and the two rows usually printed underneath the f block.

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