The Bohr model, proposed by Niels Bohr in 1913, pictures an atom's electrons orbiting its nucleus in fixed, well-defined circular paths, not unlike planets orbiting the sun. Even at the time it was proposed, physicists recognised it couldn't be a fully accurate physical description — classical electromagnetism predicted that an electron moving in a circular orbit should continuously radiate energy and spiral into the nucleus almost instantly, something atoms obviously don't do. Bohr's model simply assumed, without a physical explanation, that certain orbits were exempt from that expected radiation. It's still taught today, not because the physical picture holds up, but because it produces genuinely useful, correct numerical predictions for a specific important case.
An assumption that worked, without yet explaining why it worked
Bohr's key move was to propose, essentially as a working assumption rather than something derived from more fundamental principles, that an electron could only occupy certain specific, quantised orbits, and that it wouldn't radiate energy while remaining in one of those allowed orbits — only jumping between them, absorbing or emitting energy in the process. This assumption had no complete justification within physics as it was understood at the time; it was adopted because it worked, in the sense that it correctly predicted the specific pattern of light wavelengths emitted by hydrogen atoms, matching experimental spectroscopic data with genuinely striking accuracy for what was, physically, still an incomplete and somewhat ad hoc model.
Quantum mechanics later explained why the model worked, and where it stops
The fuller, physically consistent explanation for why quantised orbits exist came only later, through the development of quantum mechanics, which replaced Bohr's picture of electrons following well-defined orbital paths with a fundamentally different description involving probability distributions and wave-like behaviour, not classical, planet-like orbits at all. Quantum mechanics also revealed the specific limits of Bohr's simpler model: it makes accurate predictions for hydrogen, which has just a single electron, but fails to correctly predict the spectra of atoms with more than one electron, where electron-electron interactions the Bohr model has no way to account for become significant. Bohr's model survives in introductory teaching specifically because it offers an intuitive, visually simple picture and gets hydrogen's spectrum right, even though physicists have known since not long after quantum mechanics matured that it isn't the accurate physical description of what's actually happening inside an atom.
What we're still unsure about
That the Bohr model correctly predicts hydrogen's spectral lines while resting on physically inconsistent assumptions, later superseded by quantum mechanics, is well established and not disputed history and physics. What's more a matter of pedagogical judgement than settled fact is exactly how long, and in how much depth, physics education should continue relying on the Bohr model as an introductory stepping stone before moving to the more physically accurate but considerably more mathematically demanding quantum mechanical picture — different courses and textbooks make different choices about that balance, reflecting a genuine ongoing tradeoff between physical accuracy and accessible intuition rather than one universally agreed teaching sequence.
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