What Led to the Bohr Model of Hydrogen Atom?
In 1897, Sir J.J. Thomson discovered electrons as negatively charged particles present in every element's atom, but without any knowledge of the distribution of electrons, the positive charge, and the mass inside the atom. Subsequently, in 1904, Sir Thomson suggested a model for the atom, also known as the 'plum pudding model,' which stated that the electrons are embedded like plums in a distribution (or pudding) of positive charge within the atom. Thomson's model failed to explain emission spectra and alpha particle scattering. Rutherford came up with another model in which the electrons revolve around the nucleus in different orbits. The revolution is driven by the electrostatic force of attraction between the nucleus and the electrons. But Rutherford's model failed to account for the stability of atoms and the origin of line spectra. To address the shortcomings of these previous models, Prof. Neils Bohr, in 1913, applied Planck's quantum theory and proposed three postulates that came to be known as the Bohr Model of Atom. So, let us discuss the Bohr Model of Hydrogen Atom (class 12) in detail.
Prior to Bohr, there were a number of scientists who were working on the structure of an atom. Rutherford was one of them and his model was the closest one to the Bohr model of the atom. In fact, Niels Bohr had helped to overcome the problem in Rutherford’s model of an atom.
Rutherford’s model of an atom had a major drawback, that it could not explain the stability. It showed that electrons in an atom revolve around a positively recharged centre called the nucleus. However, later on, it was found that any particle in a circular motion would undergo acceleration and thus would lose energy. So the electron would take a spiral path and would finally fall into the nucleus and the atom would collapse. But this is not true because this does not take place in reality.
Bohr said that electrons do revolve around the nucleus but their energy remains fixed. He explained that the energy of the electrons remains fixed because they are restricted to some fixed orbits. Each of these orbits is at a fixed distance from the nucleus and is associated with a fixed amount of energy. These energy levels are represented by the letters K, L, M, N, or the numbers 1, 2, 3, 4, starting from the centre.
So the final model of Neils Bohr was similar to Rutherford’s model of an atom which states that an atom consists of a positively charged centre in which the electrons revolve. The only difference was that in Bohr’s model, electrons revolve around the nucleus in fixed orbits with fixed energy. For his work on the structure of the atom, he got a Nobel Prize in 1922.
Bohr's Theory of Hydrogen Atom and Hydrogen-like Atoms
A hydrogen-like atom consists of a tiny positively-charged nucleus and an electron revolving around the nucleus in a stable circular orbit.
Bohr's Radius:
If 'e,' 'm,' and 'v' be the charge, mass, and velocity of the electron respectively, 'r' be the radius of the orbit, and Z be the atomic number, the equation for the radii of the permitted orbits is given by r = n2 ｘr1, where 'n' is the principal quantum number, and r1 is the least allowed radius for a hydrogen atom, known as Bohr's radius having a value of 0.53 Å.
Limitations or Problems of the Bohr Model
The model does not explain the Heisenberg Uncertainty Principle. The Bohr atomic model theory states that the electrons have both radius and orbit. which means to have a position and momentum simultaneously. This does not match with the Heisenberg Uncertainty Principle.
The theory Bohr devised was a mixture of classical and quantum physics. Quantum physics superseded classical physics, meaning quantum physics has everything classical physics has. This makes the approach of understanding the model of the atom invalid in some aspects.
The model could not explain the various intensities of the spectral lines, which are classified under the Stark effect.
The model could not explain the existence of the hyperfine structure of some spectral lines.
The model makes inaccurate spectral line predictions when it is concerned with larger atoms such as helium, lithium and oxygen or any other element. Bohr’s model only works with hydrogen.
The model does not explain the Zeeman effect, which is the splitting of spectral lines when placed in the magnetic field.
FAQs on Bohr Model of the Hydrogen Atom
1. What is the Bohr Model of the Hydrogen Atom as explained in the CBSE 2025–26 syllabus?
The Bohr Model of the Hydrogen Atom describes the atom as consisting of a tiny, positively charged nucleus with an electron revolving in a fixed, circular orbit around it. According to Bohr, these orbits have discrete (quantized) energy levels. Electrons can jump between these levels by absorbing or emitting specific amounts of energy, resulting in line spectra instead of continuous spectra.
2. What postulates did Niels Bohr introduce for hydrogen’s atomic structure?
Bohr’s major postulates for the structure of the hydrogen atom include:
- Electrons revolve in stable, fixed orbits (energy levels) without emitting energy.
- Each allowed orbit has a definite energy (quantized energy levels).
- Energy is absorbed or emitted only when the electron moves between orbits, corresponding to the energy difference between levels.
- The angular momentum of the electron in orbit is an integral multiple of h/2π, where h is Planck’s constant.
3. What are the key limitations of the Bohr Model of Hydrogen Atom relevant for board answers?
The Bohr Model has several limitations:
- It cannot explain the spectra of multi-electron atoms.
- Fails to address the Zeeman effect (splitting in magnetic field) and Stark effect (splitting in electric fields).
- Cannot explain the intensity variations or fine/hyperfine structures of spectral lines.
- Contradicts the Heisenberg Uncertainty Principle because it assigns definite orbits and momenta simultaneously.
4. How does Bohr’s Model explain the line spectrum of the hydrogen atom?
In the Bohr Model, when an electron jumps from a higher energy level to a lower one, it emits energy as light. This release causes discrete spectral lines (such as those seen in the Balmer or Lyman series) instead of a continuous spectrum. The wavelength of emitted light depends on the energy difference between the two levels and can be calculated using the Rydberg formula.
5. Why was the Bohr Model considered a major improvement over Rutherford’s model of the atom?
The Bohr Model incorporated quantum theory to explain why electrons do not spiral into the nucleus (addressing the stability issue in Rutherford’s model). By proposing fixed energy levels, it explained the existence of stable atoms and discrete atomic spectra—limitations that Rutherford’s classical model could not resolve.
6. What is Bohr’s radius, and how is it significant in the hydrogen atom model?
Bohr’s radius (r1) is the radius of the innermost permitted orbit of the hydrogen atom and is approximately 0.53 Ångstroms. It sets the fundamental scale for electron orbits in hydrogen and can be calculated using the constants for electron charge, mass, the Planck constant, and the fine-structure constant. The formula for any orbit is r = n2 × r1, where n is the principal quantum number.
7. What are some real-world applications or examples of Bohr’s Model’s concepts in modern physics?
Concepts from the Bohr Model are used in:
- Explaining the line spectra of hydrogen in laboratory spectroscopy.
- Understanding atomic structure and emissions in astronomy (hydrogen lines in stars).
- Providing the basis for laser technologies that use controlled electron transitions.
- Teaching foundational ideas of quantum physics before more advanced models are introduced.
8. How did later models (e.g., the Bohr-Sommerfeld model) improve upon or revise Bohr’s theory?
Later models introduced elliptical orbits (Bohr-Sommerfeld) and the concept of quantum numbers for more detailed electron motion. Quantum mechanical models by Schrödinger and Heisenberg further replaced Bohr’s definite orbits with probability distributions (orbitals), providing much more accurate predictions for complex atoms and explaining previously unresolved phenomena such as fine structure and the uncertainty principle.
9. What mistakes do students often make when answering exam questions on the Bohr Model?
Common errors include:
- Confusing Bohr’s quantized orbits with classical elliptical orbits.
- Assuming the Bohr Model applies to all elements, instead of just hydrogen-like atoms.
- Incorrect calculation of radii or energy differences by not using the right formulas.
- Not distinguishing between the limitations of Bohr’s Model and its successes (e.g., failing to note its inability to explain fine structures).
10. If the Bohr Model were strictly true, what predictions would fail in practical experiments?
If only the Bohr Model applied, we would not observe:
- Multiple spectral line intensities (since Bohr does not explain their variation).
- Fine and hyperfine splitting in atomic spectra under magnetic or electric fields.
- The observed uncertainties in position and momentum (contradicting the uncertainty principle).
- Accurate predictions for atoms with more than one electron.
