Light from Atoms
When atoms absorb energy — from heat, electrical discharge, or photon absorption — their electrons jump to higher energy levels. These excited states are unstable; within nanoseconds, electrons fall back to lower levels, releasing the energy difference as photons of specific wavelengths. The resulting emission spectrum is a set of bright lines against a dark background, each line corresponding to a particular quantum transition.
The Hydrogen Blueprint
Hydrogen, with a single electron, produces the simplest and most historically important emission spectrum. The Balmer series (transitions to n=2) spans the visible range: Hα at 656 nm (red), Hβ at 486 nm (teal), Hγ at 434 nm (violet), converging toward 365 nm. Johann Balmer's empirical formula (1885) and Niels Bohr's quantum model (1913) transformed these lines into proof that energy is quantized.
Multi-Electron Complexity
Beyond hydrogen, atoms with multiple electrons produce far richer spectra. Electron-electron repulsion splits energy levels, spin-orbit coupling creates fine structure, and nuclear effects add hyperfine splitting. Iron alone has thousands of catalogued spectral lines. This simulation uses simplified Rydberg-like models scaled by atomic number and excitation energy to generate representative spectra for elements up to zinc.
From Flame Tests to Stellar Chemistry
Emission spectroscopy identifies elements with extraordinary sensitivity — parts per billion in inductively coupled plasma (ICP) instruments. Kirchhoff and Bunsen discovered cesium and rubidium through their spectral lines in 1860. Today, the same principle reveals the chemical composition of stars billions of light-years away, connecting quantum mechanics on the smallest scales to cosmology on the largest.