When an electron falls from a higher level to a lower one, the energy it loses leaves as a single photon — and the gap sets the colour exactly. Each jump is a spectral line; the full set is an element's fingerprint. Watch hydrogen drop, level to level, and land its Balmer series as real lines of light.
An electron sitting in level n₂ can fall to a lower level n₁. It must shed exactly the energy difference, and it does so as one photon whose colour is fixed by that gap: bigger drop → bluer light. For hydrogen falling to n=2 (the Balmer series, the visible one):
| jump | wavelength | colour | name |
|---|---|---|---|
| 3 → 2 | 656.3 nm | red | H-α |
| 4 → 2 | 486.1 nm | blue-green | H-β |
| 5 → 2 | 434.0 nm | violet | H-γ |
| 6 → 2 | 410.2 nm | deep violet | H-δ |
Long before orbitals were understood, spectroscopists sorted these line-series by how they looked: the sharp series, the principal series, the diffuse series, the fundamental series. Those four initials became the orbital names — s, p, d, f — and after f they just ran the alphabet (g, h, …). So the orbital letters are literally named after the appearance of spectral lines. The lines came first; the structure was read out of them.
The honest line: spectral lines are transitions — the electron jumping — which is a different fact from how many electrons a shell holds. This page is the motion; the rest of the cluster is the count. Together they're the same ladder seen two ways: the rungs (levels) and the falls between them.