Why did Bohr’s theory fail to explain the spectra of a multi-electron atom?

Bohr’s model was a great feat as it provided the best explanation possible for the structure an atom. It failed to account for all multi-electron types. These are his limitations.

Limitation of Bohrs Theory

I. Bohr claims that radiation is caused by an electron jumping from one energy orbit into another. However, Bohr.II does not explain how this radiation happens. Bohrs theory has been able to explain the observed spectra of hydrogen atoms and hydrogen-like ions (e.g. He+, Li2+, Be3+ etc. However, it was unable to explain the line spectra for multi-electron atoms.
III. Bohr’s model had two dimensions, while an atom has three.
V. He couldn’t explain the Zeeman effect (magnet effect) in line spectra.
VI. VI. He cannot explain Stark effect (electric Effect) on line spectra.
VII. Bohr’s equation clearly defined the position and momentum of an electron, as well as its revolving circle around the nucleus. However, Heisenberg’s Uncertainty principle states that it is impossible to accurately measure the momentum and position of electrons. The position can only be measured with the greatest precision. There will not be any uncertainty in the momentum’s value and vice versa.
IX. Acc. Bohr states that the angular momentum (mvr), of an electron in an nth orbit is equal or greater than nh/2p. However, he cannot explain this concept. (Later provided by de Broglie).

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