Atomic Structure

Thomson Model of Atom

J. J. Thomson proposed this model in 1898. As per this model, an atom is spherical in shape in which positive charge is uniformly distributed. Electrons are embedded in this positively charged sphere in a manner that the atom gets the most stable electrostatic arrangement. This model is called plum pudding model, raisin pudding model or watermelon model. This model was able to explain the neutrality of atom but it was not consistent with later experiments.

Rutherford’s Model of Atom

Rutherford's Alpha particle experiment

Rutherford conducted his famous α-particle scattering experiment in which a stream of high energy α-particles was directed at a thin foil of gold. The gold foil had a circular fluorescent zinc sulphide screen around it. Whenever α-particles struck the screen, a tiny flash of light was produced at that point.

The results of this experiment were quite unexpected. As per Thomson model of atom, the mass of each gold atom in foil should have been evenly distributed over the entire atom. So, α-particles should have passed through such a uniform distribution of mass. Particles would have slowed down and change direction by small angles. But contrary to this perception, following observations were made:

On the basis of these observations, Rutherford came with following conclusions:

Rutherford calculated and showed that the volume of nucleus is negligibly small compared to the volume of the atom. The radius of the atom is about 10-10 m, while radius of nucleus is about 10-15 m.

Key points of Rutherford’s Nuclear Model of Atom:

Atomic Number (Z) = No. of protons in atom
= No. of electrons in an electrically neutral atom

Mass Number (A) = No. of protons + No. of neutrons in atom

Isobars: Atoms with same mass number but different atomic numbers are called isobars.

Examples of isobars: 146C and 147N

Isotopes: Atoms with same atomic number but different mass numbers are called istotopes.

Examples of isotopes: Protium (11H), Deuterium (21D) and Tritium (31T)

Chemical properties of atoms are controlled by the number of electrons, which are determined by number of protons in nucleus. Hence, chemical properties of isotopes are same.

Drawbacks of Rutherford’s Model

You have learnt that Rutherford’s nuclear model of atom is similar to the solar system, albeit on a much smaller scale.

When we apply the classical mechanics to the solar system, we see that the planets describe well-defined orbits around the sun. The gravitational force between the planets is given by following expression.


Where, m1 and m2 are masses of given bodies and r is the distance of separation of masses. G is the gravitational constant. This theory can also calculate precisely the planetary orbits and these are agreement with the experimental measurements.

The similarity between the solar system and nuclear model suggests that electrons should move in well defined orbits. Moreover, the coulomb force between electron and nucleus is mathematically similar to the gravitational force. This force is given by following expression.


Where q1 and q2 are charges, r is the distance of separation of charges and k is the constant of proportionality.

But, when a body is moving in an orbit, it undergoes acceleration; because of changing direction on circular path. So, electron (moving in an orbit) is an object under acceleration. According to electromagnetic theory of Maxwell, charged particles should emit electromagnetic radiation when they are accelerated. So, an electron in an orbit will emit radiation (lose energy) and the orbit will continue to shrink. Calculations show that it should take only 10-8 second for an electron to spiral into nucleus. But it does not happen in real life and atom does not lose its existence. In other words, Rutherford model could not explain the stability of an atom.

Rutherford model says nothing about distribution of electrons around the nucleus and the energies of these electrons.

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