Introduced by Danish physicist Niels Bohr in 1913, the Bohr model revolutionized the scientific understanding of the atom. Before Bohr, scientists struggled to explain why electrons did not spiral into the positively charged nucleus, which classical physics predicted should happen. Bohr solved this by proposing that electrons could only occupy specific, fixed circular orbits around the nucleus, similar to planets orbiting the sun. This is why it is often called the planetary model.
The most significant contribution of the Bohr model was the concept of quantization. Bohr postulated that each specific orbit corresponds to a specific, discrete energy level. Electrons can only exist on these defined paths and never in the empty space between them.
For an electron to move to a higher orbit, it must absorb a very specific, quantized packet of energy, known as a photon. To drop to a lower orbit, it must emit a photon of a specific energy.
This theory perfectly explained the emission spectrum of hydrogen. When hydrogen gas is excited by electricity, it emits light at only a few specific wavelengths, creating a distinct line spectrum rather than a continuous rainbow. The Bohr model proved that these distinct lines of light correspond exactly to the energy differences between the electron’s fixed orbits.
When an electron drops from a high energy orbit to a lower one, the energy difference is released as visible light.
Despite its massive success with hydrogen, the Bohr model is ultimately fundamentally flawed. It could not accurately predict the emission spectra of atoms with more than one electron, and it incorrectly treated electrons as distinct particles moving on predictable paths. Modern quantum mechanics has since replaced the Bohr model with the electron cloud model, which treats electrons as standing waves in three-dimensional orbitals. However, the Bohr model remains a vital teaching tool for introducing atomic energy levels.