Bohr's Model: Learn its Postulates, Limitations & Uses

Neils Bohr proposed Bohr’s model of an atom in 1915. He was a Danish physicist who received his PhD from the University of Copenhagen in 1911. He then spent a year in England with J.J Thompson and Ernest Rutherford and was awarded the Nobel Prize in Physics in 1922. Historically, results observed from the studies of interactions of radiations with the matter have provided immense information regarding the structure of atoms and molecules. Neils Bohr used these findings to build on Rutherford’s model, and while his theory is not current quantum mechanics, it can still be used to justify many points in the atomic structure and spectra.

Read on to learn more about the Chemistry of Bohr’s atomic model, its postulates, and its limitations.

Bohr’s Model of Hydrogen Atom

According to Bohr’s model of an atom, a positively charged nucleus is present at the centre of the atom, and the electrons (negatively charged particles) move around the nucleus in orbits with fixed size and energy. He also explained that by gaining or losing energy (in the form of photons), the electron could jump in-between orbits. Bohr explained that the electrons present in shells that are far from the nucleus would have more energy, and the electrons present in shells that are near the nucleus will have less energy. Also, electrons present in lower orbits will be the most stable.

In Rutherford’s model of an atom, the electrons were in random motion, and the model didn’t explain the stability of an atom. Bohr’s atomic model is similar to the structure of the solar system, where planets orbit the sun. The force of attraction between the planets and the sun is gravity, while the force of attraction between electrons and the nucleus is the electrostatic force of attraction.

Bohr’s atomic model diagram is given below:

Postulates of Bohr’s Atomic Model

The following are the key postulates of Bohr’s atomic model:

• An atom consists of a small, heavy and positively charged nucleus in the centre, and the electrons revolve around it in circular orbits.
• Electrons rotate exclusively in orbits with a set energy value, out of a wide variety of circular orbits theoretically feasible around the nucleus. As a result, these orbits are known as energy levels or stationary states. The word stationary does not mean that the electrons are stationary, but it means that the energy of the electron revolving in a particular orbit is fixed and does not change with time. The different energy levels are numbered 1, 2, 3, 4…, etc., or designated as K, L, M, N, O, P…, etc., starting from the shell closest to the nucleus.
• Since the electrons revolve only in those orbitals which have fixed values of energy, they can have only certain definite or discrete values of energy and not any value of their own. It is expressed by saying that the energy of an electron is quantized.
• The angular momentum of an electron in an atom, like energy, can have definite or discrete values and not have any value of its own, and it is likewise quantized.
• When electrons in an atom are at their lowest or normal energy level, they continue to revolve in their orbits without losing energy. Because energy cannot be constantly lost or acquired. It is referred to as an atom’s normal or ground state.
• Only when electrons leap from one orbit to the next can they produce or absorb energy. For example, when energy is provided to an atom by electric discharge or high temperature, an electron in the atom may absorb a specified quantity of energy and jump from its regular energy level, i.e., ground state, to some higher energy level. It is known as the excited state of an atom. Because the electron’s lifespan in the excited state is limited, it returns to the lower energy level instantly by releasing energy in the form of light of an appropriate frequency or wavelength.
• Further, since each energy level is associated with a certain definite amount of energy, therefore energy is always emitted or absorbed in certain discrete quantities called quanta or photons and not any value. It means that the energy of the electron cannot change gradually and continuously but changes abruptly as the electron jumps from one energy level to the other.

Distribution of Electrons in Orbits

In an atom, electrons are distributed in circular orbits revolving around the nucleus. In the ground state of the atoms, the orbitals are filled with electrons in order of their increasing energies. In other words, electrons occupy the lowest-energy orbital accessible to them initially and only enter higher-energy orbitals when the lower-energy orbitals are full.

The maximum number of electrons a shell can hold is given by the formula , where n is the number of energy shells. 1st energy shell is K, the second energy shell is L, the third energy shell is M, and so on. Therefore, the 1st shell (K) can have a maximum of 2 electrons. Second shell (L) can have a maximum of 8 electrons, third shell (M) can have a maximum of 18 electrons, and fourth shell (N) can have a maximum of 32 electrons.

Uses of Bohr’s Model of Atom

Some of the uses of this model of atom are:

• It explains the stability of the atom.
• It explains the line spectrum of hydrogen.
• Bohr’s model helps to calculate the energy of an electron in H-atom and H-like particles.
• This model also helps to deduce Rydberg’s formula put forward empirically.

Limitations of Bohr’s Model

Bohr’s atomic model has many limitations, which are discussed below:

• Bohr’s theory successfully explains the line spectra of hydrogen atoms and hydrogen-like particles containing a single electron only. It did not, however, explain the line spectra of multi-electron atoms. When spectroscopes with better resolving powers were used, it was found that even in the case of the hydrogen spectrum, each line was split up into several closely spaced lines (called fine structure), which could not be explained by Bohr’s model of an atom.
• In the production of line spectrum, if the source emitting the radiation is placed in a magnetic field or an electric field, it is observed that each spectral line splits up into several lines. The splitting of spectral lines in the magnetic field is called the Zeeman effect, while the splitting of spectral lines in the electric field is called the Stark effect. Bohr’s model of the atom was unable to explain the splitting of spectral lines.
• According to Bohr’s atomic model, the electrons move along certain circular paths in one plane. Thus, it gives a flat model of an atom. But it is well established that the atom is three-dimensional and not flat, as Bohr suggested.
• Now, it is well known that in covalent molecules, the bonds have directional characteristics (i.e., atoms are linked to each other in particular directions), and hence they possess definite shapes. Bohr’s model is unable to explain it.
• Another limitation of Bohr’s model is its inability to explain de Broglie’s concept of the dual nature of matter and Heisenberg’s uncertainty principle.
• It is also unable to explain the concept of elliptical orbits.
• Bohr’s model does not consider electron spin energy.
• The assumption that angular momentum is an integral multiple of \( \frac{h}{2\pi} \) was made without any explanation.

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