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Niels Bohr's model of the hydrogen atom was one of the first great successes of quantum theory. Its predictions of the wavelengths of the hydrogen spectrum are within a few tenths of one percent or better of the actual values. It failed however in the explanation of the spectra of helium and the higher elements. Nevertheless it is still a valuable theory for providing insights into the mechanism of electron transition phenomena on the subatomic level.
The great success of the Bohr model had been in explaining the spectra of hydrogenlike atoms;i.e., ions with a single electron around a positive nucleus. The Schroedinger equation replicated this explanation in a more sophisticated manner and the Bohr analysis was considered obsolete. But the Schroedinger equation approach can be solved for only a very limited number of models. Beyond this limited set the Schroedinger equation approach gives no insights, whereas the Bohr model does provide insights into diverse cases. In particular the Schroedinger equation approach cannot be applied to the case which takes into account the relativistic effects. On the other hand, the Bohr analysis can. See Relativistic Bohr Model.
The details of the application of the Bohr analysis to a hydrogen like atom in the
nonrelativistic regime are given here. The significant initial assumption is
that the angular momentum is quantized in units of Planck's constant divided by 2π,
h. The angular momentum of the electron is mvr,
where r is the radius of the elecron's orbit, v is its orbital velocity and m is its mass. Thus
where n is a positive integer which is called the quantum number of the electron.
The potential energy of an electron, V(r), is given by −κ/r, where κ is a constant equal to the force constant for electrostatic attraction times the square of the charge of an electron. The attractive force is given by −κ/r².
In a circular orbit the balance of the attractive force and the centrifugal force requires that:
But from the quantization of angular momentum
Equating the two expressions for v² gives
This is the quantization condition for the orbit radius. The quantization of the other characteristics of the state of the electron follow from that for r.
Orbital velocity is given by
Kinetic energy K is given by
The potential energy V is then
From the expressions for K and V it is seen that
If total energy decreases by ΔT half of the decrease goes into increased kinetic energy and the other half goes into an emitted photon. Thus the energy γ is given by
where n_{I} is the initial quantum number for the electron and n_{F} is its final quantum number. The photon energy U is converted into wavelength λ via the relationship
where c is the speed of light.
Comparison of Spectral Wavelengths Computed from Bohr Model of the Hydrogen Atom with the Measured Wavelengths 


Measured  Bohr Model  Proportional Deviation  
Upper  Lower  wavelength (nanometers)  wavelength (nanometers)  Of 1%  
2  1  121.566  121.551  0.01234  
3  2  656.28  655.987  0.04465  
3  1  102.583  102.549  0.03314  
4  3  1875.01  1878.509  0.18661  
4  2  486.133  486.202  0.01419  
4  1  97.254  97.24  0.0144  
5  4  4050  3999.407  1.24921  
5  3  1281.81  1278.161  0.28468  
5  2  434.04  433.502  0.12556  
5  1  94.976  94.932  0.04633  
6  5  7400  7293.036  1.44546  
6  4  2630  2637.907  0.30065  
6  3  1093.8  1097.182  0.3092  
6  2  410.174  410.535  0.08801  
6  1  93.782  93.712  0.07464  
7  3  1004.98  1007.981  0.29861  
7  2  397.002  397.377  0.09446  
8  3  954.62  953.705  0.09585  
8  2  388.9049  388.657  0.06374  
9  2  383.5384  383.844  0.07968 
As can be seen in the table the errors are typically a small fraction of one percent.
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