Limitations of Bohr’s Atomic Theory

According to Danish physicist Niel Bohr’s atomic model, an atom includes a small, charged nucleus surrounded by electrons that journeys in round orbits across the nucleus with force supplied through electrostatic forces. This version of Bohr’s theory was entirely based on Planck’s quantum principle of radiation. Bohr’s atomic principle became massively popular in explaining the stability of the atom and the broad spectrum of a hydrogen atom. The Bohr atomic principle made accurate predictions for lesser-sized atoms like hydrogen. 

Additionally, this principle did not explain the Stark impact while the spectral line breaks up into quality strains within the occurrence of an electric-powered discipline. Bohr also gave some postulates to help us understand his atomic model. However, there are numerous limitations to Bohr’s atomic theory. 

Bohr’s Atomic Theory

According to Bohr’s Atomic Model, an atom includes a small, charged nucleus surrounded by electrons that revolve in round orbits due to electrostatic forces across the nucleus. 

Postulates of Bohr’s Atomic Model

The most important postulates of Bohr’s Atomic Model are as follows: 

1. The electrons flow across the nucleus in positive circular paths called orbits.
2. Each orbit is related to a specific quantity of electricity, and therefore, those are called electricity ranges or electricity states. Electricity ranges are numbered as 1,2,3,4, etc. or distinctively as K, L, M, N, etc. The electricity stage closest to the nucleus is numbered 1, distinguished as the K shell. 
3. While shifting a specific electricity stage, an electron neither lost nor gained electricity. The electricity of an electron in a particular state of electricity usually stays constant or stationary. This is known as a regular or floor stage.
4. An electron emits or absorbs electricity. It jumps from one orbit or electricity stage to another. When it jumps from a higher to a lower electricity stage, it emits electricity even as it absorbs electricity while it jumps from a lower electricity stage to a higher electricity stage.
5. The electricity absorbed or released is identical to the distinction between energies of both electricity ranges (E1, E2) and is decided by Plank’s equation.

ΔE = E2-E1 = HV 

Where  E2 and E1 are the energies of electrons inside the higher and lower electricity ranges, respectively, ΔE is the distinction in energies of ranges.  Here, V is labeled as the frequency of radiation emitted or absorbed.

The transition of electrons ranges from lower to higher and vice versa.

1. Like electricity, the angular momentum of an electron in an atom will have positive specific or discrete values and now no longer have any values of its own. The best possible values of angular momentum are given through the expression,mvr= nh/2π, i.e., h/2π could be an essential more than one of angular momentum of the electron. Here, m= mass of the electron, ν= tangential speed of a revolving electron, r= radius of the orbit, h= Planck’s regular, and n= integer 1,2,3,

All the above postulates hold significant importance in Bohr’s atomic theory model, and without these, he would not be able to explain the theory effectively.

Limitations

The limitations of Bohr’s atomic theory are as follows:

• Bohr’s version of an atom could not explain the broad spectra of atoms containing multiple electrons, known as multielectron atoms. According to Bohr’s principle, one and the best spectral line can originate from an electron among any given electricity. 
• Positive single strains break into some very carefully associated strains when practical spectroscopy is used. The life of one of these lines could not be defined on the premise of Bohr’s principle.
• Bohr’s principle did not account for the impact of magnetic discipline on the spectra of atoms or ions. It changed after the atom emitting radiations were positioned in a robust magnetic field. Every spectral line is in addition broken up into some strains. This phenomenon is called Zeeman’s effect.
• As a result, it is no longer difficult to explain why the spectral line is broken up into numerous additives in the presence of a magnetic discipline. While the hydrogen fuel line changed into an excited in a magnetic field, the produced emission spectrum changed into a break-up. Bohr’s version couldn’t account for this.
• Another limitation of Bohr’s atomic theory was that it failed to explain the impact of an electric-powered discipline called the Stark impact at the spectra of atoms. When a fabric with a line emission spectrum is placed in an outside electric discipline, its strains spill into some carefully spaced strains. The relative intensities of spectral strains cannot be predicted using this principle.
• Bohr’s principle does explain the shapes of molecules that form as a result of directional bonding between atoms. In terms of size, this version may be very restrained. Poor spectral predictions are obtained when large atoms are involved.
• The Bohr Model now no longer accounts for the truth that accelerating electrons no longer emit electromagnetic radiation. According to Bohr, radiation is emitted while an electron jumps from one electricity orbit to every other electricity orbit. However, how this radiation takes place isn’t always defined through Bohr. It violates Heisenberg’s Uncertainty principle. 
• The Bohr Model regards electrons to have a respective radius and orbit, which is not in line with Heisenberg’s principle. Bohr assumes that an electron in an atom is placed at a particular detachment from the nucleus and is revolving spherically with a specific speed. This is more in line with Heisenberg’s Uncertainty principle.

Conclusion 

Bohr’s atomic theory talks about an atom that includes a small, charged nucleus surrounded by citrons. Electrons revolve around orbits across the nucleus with force applied through electrostatic forces. The Bohr model is important because it was the first to hypothesize the quantization of electron routeways in titles. The limitations of Bohr’s atomic theory were that it could not explain the broad spectra of atoms containing multiple electrons, known as multielectron atoms. Moreover, it did not account for the impact of magnetic discipline on the spectra of atoms or ions. This principle could not explain the impact of an electric-powered field called the Stark impact at the spectra of atoms.It no longer offers any clue to explain the shapes of the molecules bobbing up out of the directional bonding among atoms.