Rutherford and Bohr Atomic Model

Rutherford Atomic Model:

Atoms are the basic building blocks that combine to produce everyday matter in the world around us.

In 1914, Rutherford proposed that most of the mass of an atom and positively charged protons were concentrated in an extremely small volume of the atom. He called this region the nucleus, with negatively charged electrons surrounding it. He also proposed that these electrons revolve around the nucleus at a very high speed in circular paths, which he named orbits. Since electrons are negatively charged and the nucleus is densely populated with positively charged protons, there is a strong electrostatic force that holds together the nucleus and electrons.

Rutherford’s model was widely accepted till the second half of the 19th century when Maxwell developed the classical theory of electromagnetism. The classical electromagnetic theory postulates that a charged particle changing either speed, direction or both would emit electromagnetic radiation in the form of light. An orbiting electron in Rutherford’s model continuously emits radiation.

As a result, the electron must lose energy and will be pulled more and more strongly by the nucleus and eventually crash into it. In practice, however, the electron doesn’t collapse into the nucleus. Also, there is no continuous emission of radiation. On the contrary, the energy emission is confined to a discrete wavelength called a line spectrum, which could not be explained by the Rutherford model.

Bohr Atomic Model:

Bohr ignored some of the concepts of classical physics utilized by Rutherford in his theory of atomic structure. Instead, he used experimentally observed facts and Plank’s idea of energy quantization to propose a new model in 1913. He developed his theory of the atom based on the hydrogen (one electron) atom and used the following postulates:

1. The electrons in an atom move in a circular orbit around the positively charged nucleus under the influence of the electrostatic attraction between an electron and the nucleus. This force is balanced by the centrifugal force due to the velocity of an electron in its orbit. The orbits used by the electrons are a specific distance away from the nucleus. The electron in a hydrogen atom normally stays in the first orbit, known as the ground state, which is closest to the nucleus.

2. The energy of the electron remains constant as long as it remains in a permitted orbit.

3. Radiation is emitted or absorbed when the electron makes a transition from an allowed orbit to another of lower or higher energy respectively. The energy difference between the two orbits ΔE is given below:

E₁ – E₂ = |ΔE| = hf

When an electron moves from a higher (energy) orbit E₂ to a lower on E₁ thereby losing energy, the lost energy manifests as a photon, that is light. The energy of the emitted light is related to its colour.

The electron loses the same energy it absorbed when it jumped from a lower orbit to a higher one. Bohr noticed that all of the possible combinations of jumps the electron could make to return to the ground state from the higher orbits account for the existence of all the special lines in the hydrogen atom.