Atomic Models: Thomson, Rutherford, and Bohr’s Contributions

Posted on Mar 19, 2025 in Chemistry

Atomic Models: Thomson, Rutherford, and Bohr

JJ Thomson proposed the first model of the atom: electrons are like plums embedded in a positive charge ‘pudding’. The positive charge of the ‘pudding’ exactly compensated for the negativity of the electrons, making the atom electrically neutral.

Rutherford: When alpha particles are fired at a thin gold foil, some are deflected. This indicates that there is a zone (nucleus) where the positive charge is concentrated, and a mass greater than or comparable to that of the alpha particles. The area where the mass is concentrated and positively charged must be very small compared to the size of the atom. The electrons orbit the nucleus in circles.

This contradicted Maxwell’s electromagnetic theory. According to this theory, any accelerating electrical charge should emit electromagnetic waves. An electron turning in circles around the nucleus should therefore emit electromagnetic waves. This lost energy would cause the electron to spiral into the nucleus. Rutherford’s atomic model was therefore unworkable from the point of view of classical physics.

It did not give a satisfactory explanation of atomic spectra. If a tube is filled with H or I and subjected to high voltages, the gas emits light. If that light is passed through a prism, the constituent colors are separated, giving a light spectrum. It was concluded that light emission could be due to electrons jumping from superior orbits to + inferior orbits, emitting the excess energy in luminous form. This interpretation, however, would suggest that the spectra should be continuous, as orbits at any radius (and energy) are possible. Experience, on the contrary, showed that the atomic spectra are discontinuous, consisting of different colored stripes on a black background.


Bohr’s Postulates

  1. An electron in a specific orbit does not emit energy. These orbits are considered stationary energy states. Each orbit corresponds to a specific energy level; the further away from the nucleus, the higher the energy.
  2. Not all orbits are possible. Only orbits with certain energy values are allowed, defined by the principal quantum number, n. Only orbits where the main quantum number (n) takes integer values are possible: n = 1, 2, 3, … Orbits that do not correspond to integer values of the main quantum number are forbidden.
  3. The energy released when an electron drops from a higher orbit, of energy E2, to a lower orbit, of energy E1, is emitted as light. The frequency (f) of the light is given by the expression:

Bohr’s postulates gave excellent results when interpreting the spectrum of the H atom, but it must be taken into account that they contradicted some of the more established laws of physics:

  1. The first postulate was against Maxwell’s electromagnetic theory, since according to this theory, any accelerated electric charge should emit energy as electromagnetic radiation.
  2. The second postulate was even more surprising. In classical physics, it was unacceptable to assume that the electron could only orbit at certain distances from the nucleus, or could not have certain energy values. The statement was tantamount to assuming that an object describing circles attached to a rope, can not describe those whose radius is not a multiple of two (for example).
  3. The third premise stated that light is emitted in small packets or quanta, which although it had been proposed by Planck in 1900, was still surprising in an era in which the idea that light was a wave was firmly entrenched.