The World of Atom (Part IX)

ReferenceA Logical Approach to Theoretical Physics



Chapter 43: Strange Results from α-Particle Scattering (Hans Geiger 1882 – 1945, Ernest Marsden 1889 – 1970)

The discovery of the electron in 1897 by J. J. Thompson meant that the atoms have a structure. Since the atom is electrically neutral it was assumed that small discrete electrons were embedded in a jelly like viscous sphere of positive electricity. Thompson made a study based on this model. He assumed that the angle of deviation suffered by the charged particle was always caused by a large number of collisions with many atoms. When Geiger and Marsden conducted their scattering experiments, they found that a thickness of 6 x 10-5 cm of gold foil was sufficient to produce a deflection of 90 degrees or more. Rutherford contended that the large deviations observed experimentally must have been caused by single direct collisions, and he began looking for an alternative to explain the data.

Rutherford’s research gave us the nuclear atom, which is consistent with the “whirlpool” model of the atom.

Chapter 44: The Nuclear Atom (Ernest Rutherford 1871 – 1937)

Rutherford was impressed about some of the α-particles scattering back toward their source in the work of Geiger. To him this could only mean one thing: there was an enormous force in the atom, that it would take one hundred electron charges on the gold nucleus to give the observed result. He then proceeded to analyze the theory of single collisions on the basis of a model of the atom that is radically different from the Thomson model. Rutherford was particularly interested in the question, “How close to the nucleus can an alpha particle approach?” He got this answer from his theory and the measurements of the scattered α-particles. The answer, 3 x 10-10 cm., showed him how small and compact a nucleus is, and thus the nuclear atom was born. In this Rutherford model the positive electricity is not distributed over a large volume but instead is concentrated in a very small nucleus at the center of the atom. A model of this sort cannot be in static equilibrium, yet this kind of dynamical equilibrium was in serious contradiction with classical electrodynamics.

This contradiction with classical electrodynamics was resolved later.

Chapter 45: Atomic Structure (Niels Bohr 1885 – 1962)

Rutherford had a planetary model in mind for his nuclear atom, but he could not determine a precise orbit for the electron since the classical electrodynamic theory showed that electron would radiate energy and it must ultimately spiral into the nucleus. Bohr reasoned that if the electron can change its state of motion only in discrete steps it must then stay in a particular orbit until it emits or absorbs enough energy in one single process to go from one orbit to another. Bohr introduced Planck’s constant, imposing the condition that the electron must not radiate continuously but rather in the form of “distinctly separated emissions.” Bohr then introduced the concept of “stationary states” to describe the discrete orbits. With his assumptions Bohr obtained correct expression for the frequencies of the spectral lines in the Balmer series in the spectrum of hydrogen.

Classical electrodynamics assumes electron to be a mass particle with charge, which then makes it spiral into the nucleus. But the atom is stable, and the dynamics of an electron occur naturally. That means that the electron cannot be a mass particle with charge. It is a substance of a very different consistency, and that fact is responsible of its dynamics and charge. Quantum mechanics simply introduces this new dimension of consistency of substance that determines its energy of interaction. The geometry of orbits stratifies this consistency. Transitions among these stratifications of consistency generate the spectral lines.

Chapter 46: The Quantum Theory is Tested (James Franck 1882 – 1964, Gustav Hertz 1887 – 1975)

It was not clear from Bohr’s theory of atomic spectra alone whether the quantum theory could be applied to ordinary mechanical energy of motion, or whether it was limited to the emission and absorption of radiant energy. In 1914 James Franck collaborated with Gustav Hertz and completed basic experiments on the collision of electrons with atoms, which demonstrated that an atom could take on energy from collisions only in discrete amounts, in agreement with Bohr’s theory. These experiments demonstrated that a particle like an electron would transfer its energy in a collision only in multiples of a fundamental quantum. From this point it was clear that quantum theory would have to be taken into account in all processes.

In truth electrons are just as massless as radiant energy, the only difference is that electron’s consistency, and hence its inertia, is much higher than the consistency and inertia of radiant energy.

Chapter 47: The discovery of Isotopes (Frederick Soddy 1877 – 1956)

It was recognized that some atomic weights, such as chlorine, were by no stretch of imagination integral. There was speculation that the atoms of a given element might not be identical and that the atomic weight of an element might be the average of the weight of several unequally massive atoms having the same chemical properties. As experience with radioactivity increased, Frederick Soddy (1910) realized that the radioactive transformations produced atoms of the same chemical species but of different weights. We now recognize such chemically non-separable atoms as “isotopes.”

Chapter 48: The Positive Rays (J. J. Thomson 1856 – 1940)

The canal rays were discovered in 1886. They consisted of positively charged particles, but their composition was found to be very complex. The problem was to develop an experimental method to separate out the various particle components. In a lecture in 1913, J.J. Thomson described a simple device, involving electric and magnetic fields, that is the forerunner of the modern mass spectrometer. In this way a spectrum of the canal ray is obtained on a photographic plate. Each point on the spectrum represents a particular value of the speed and of the ratio of charge to mass. The parabola corresponding to hydrogen ions (protons) is the one that is deflected the most. Thomson was thus able to catalogue the atomic and molecular weights of the various elements and compounds in the tube.

Chapter 49: Transmutation of an Element (Ernest Rutherford 1871 – 1937)

After his basic analysis of the experiments on the scattering of α-particles by heavy nuclei, which finally led to the nuclear model of the atom and to Bohr’s theory of atomic spectra, Rutherford began a series of experiments in 1915 on the scattering of α-particles by light atoms. These experiments resulted in the first artificially induced nuclear transformation of an element in 1919. He discovered that α-particles, in passing through air, collided with nitrogen atoms and knocked protons out of these atoms. The importance of Rutherford’s results for the future of atomic physics was of a twofold nature. First, he demonstrated experimentally that nuclei of atoms contain individual protons. Second, his experiment showed that nuclei could be disrupted and changed into other nuclei; this was the first example of the artificial transmutation of chemical elements.

Chapter 50: The Diversity of Atoms (Francis William Aston 1877 – 1945)

J. J. Thomson’s positive ray apparatus opened the door in 1912 for investigations into elements for their isotopes. F.W. Aston assisted Thomson in finding the isotopes of Neon. In 1919 Aston developed a mass spectrograph as an improvement of Thomson’s parabola method. This spectrograph was used to analyze some fifty elements in the following six years, revealing the almost universal existence of isotopes. Mass spectrometers of increasing precision followed, and Aston found that neither the whole number rule nor Prout’s hypothesis could be substantiated. The addition of masses, equal to that of hydrogen, did not give the masses of the succeeding elements in the periodic table. The elemental masses as measured by the spectrograph showed a “mass defect” as compared to the sum of the masses of their free constituent particles. The greater was the mass defect the more stable was the nucleus.



  1. Large deflection of α-particles from very thin gold foil gives rise to the idea of a nuclear atom.
  2. The size of the nucleus is of the order of 10-10 cm., compared to the size of the atom of 10-8 cm.
  3. The positive charge is concentrated in a very small nucleus at the center of the atom.
  4. Nuclear atom cannot be in static equilibrium, and dynamical equilibrium contradicts classical electrodynamics.
  5. This necessitates the postulate that electron can change its state of motion in discrete steps only.
  6. This then leads to discrete orbits, and transitions from one such orbit to another give rise to a discrete spectrum.
  7. This was confirmed by correct expression for the frequencies of the spectral lines for hydrogen.
  8. It was further confirmed with atoms taking on energy from collisions only in discrete amounts.
  9. The radioactive transformations produce atoms of the same chemical species but of different weights (isotopes).
  10. Particles in a ray can be separated by their masses.
  11. Nuclei of atoms contain individual protons.
  12. Nuclei can be disrupted and changed into other nuclei (artificial transmutation).
  13. The elemental masses show a “mass defect” as compared to the sum of the masses of their free constituent particles. 
  14. The greater is the mass defect the more stable is the nucleus.

The atom consists of a small massive nucleus surrounded by a large volume of rapidly revolving electrons. The  electrons have momentum but no mass. They form discrete stationary orbits around the nucleus that have variations in consistency and energy. Exchanges among these stationary orbits generate the spectral lines. The nucleus is made up of protons but there is a mass defect as protons combine. There are isotopes and artificial transmutation can be induced.


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