Bohr model
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Picture the solar system: planets circling the Sun on fixed tracks. Niels Bohr proposed in 1913 that an atom looks much like this — the electron orbits the nucleus on certain allowed circular paths, like planets confined to particular lanes. The radical twist was that only specific orbits are permitted, and an electron jumping between them flashes out a packet of light.
The Bohr model is an early picture of the atom in which a single electron travels in fixed circular orbits around the nucleus, each orbit having a definite radius and energy. The electron does not radiate while in an orbit; it emits or absorbs light only when it jumps between orbits, releasing or taking in a photon whose energy equals the gap between levels. This neatly explained hydrogen's spectrum for the first time.
The model was a milestone — it correctly predicted hydrogen's spectral lines and introduced quantised energy levels. But its caveat is large: it treats the electron as a particle on a definite track, which the later wave-mechanical picture overturned. It fails for atoms with more than one electron and is best seen as a historic stepping-stone, not the modern truth.
Bohr's model assigns hydrogen's lowest orbit a radius of about 0.053 nanometres — the so-called Bohr radius — and an energy of −13.6 electronvolts. Plugging the gaps between his allowed orbits into a simple formula reproduces the exact wavelengths of hydrogen's visible spectral lines.
Electrons on fixed orbits, light emitted on jumps — right for hydrogen, wrong as a literal picture.
The Bohr model gets hydrogen's energy levels right but for the wrong reasons; the modern quantum-mechanical model replaces fixed orbits with probability clouds (orbitals) and works for all atoms.