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The History and Introduction of Nuclear Chemistry and the Atomic Model


Ever since the beginning of human curiosity, mankind has contemplated about the universe, more specifically about the very essence of the environment around them and what the environment consists of. For instance, Democritus wrote, approximately 400 B.C., that indivisible “atoms” must exist and can combine to form substances that today are called molecules. However, Democritus’ theory was based on the principle that all matter cannot be “indefinitely” subdivided, even with the “sharpest imaginable knife;” because eventually, matter would be too small to be able to divide (Harvey 1).

It was not until the 19th century that the basis of the atomic theory was discovered. However, this theory assumed “compounds consisted of molecules,” and each compound has a “fixed” number of atoms. This theory of component ratios was the start of the atomic theory, even though it was relatively vague. It was not until the beginning of the 20th century that “the atomic theory was firmly established” (Harvey 2). However, even though a lot of the atomic theory has been more recently discovered, there are some complexities and uncertainties of nuclear chemistry of atoms that are still very puzzling.


The first part of the atom to be discovered was electrons and their properties. In 1897, J. J. Thomson discovered electrons “in the rays emitted from the cathode of a discharged tube filled with gas at low pressure.” Since these electrons were found to be emitted from all gases in discharged tubes, Thomson postulated that electrons are found in all atoms and each electron has a negative electrical charge (Harvey 2). This charge was found to be approximately 1.6 E19 coulomb (Gillespie 50). However, atoms were already known to have neutral electrical charges, this each atom must have positively charged particle(s) somewhere inside it to neutralize the negatively charged electron(s) in each atom (Harvey 2).

It was not until 1910 that Thomson theorized about the structure of atoms. Thomson knew that almost all the mass of atoms were from the positively charged particles, because the electrons had relatively insignificant mass compared to the rest of the atom (Harvey 2). The proton, the positively charged particle, has a mass approximately equal to 1.67 E-27 kilograms, yet the mass of an electron is about 9.1 E-31 kilograms. Thus, the proton has a mass of almost 2000 times that of an electron (Gillespie 50). Therefore, Thomson believed electrons to “revolve within” positively charged spheres (Harvey 2). Unfortunately, this theory was not substantiated and did not explain several other properties of atoms (Harvey 3).

In 1911, Rutherford theorized a new atomic model where the electrons are located farther away from the positively charged particles, so atoms would not be “completely neutralized.” This new model could explain why alpha particles (helium ions with a positive electrical charge) have a high tendency of being deflected by the “electrostatic repulsion between the positive charges contained in the atom and the alpha particles.” Thus, the new atomic model was illustrated by having “an extremely small nucleus” with a positive charge and almost all the mass of the atom. Alternatively, electrons are in a high orbit surrounding the nucleus with relatively insignificant mass (Harvey 3).

In 1900, the reason why electrons were not completely attracted to the positively charged particles, thus possibly resulting with a neutralization of each other, was finally unraveled by Planck’s Quantum Theory. Planck theorized that an oscillating electron emits energy, as light, by changing from a “definite energy level” to a “second definite level of lower energy.” Thus, electrons have “certain definite quantities of energy,” that are directly “responsible for the emission of light from heated bodies.” As a result, Planck proved that the change in energy (E) is equal to a constant (h), approximately equal to 6.6.252 E-27 erg-seconds and called the Planck’s constant, times the frequency (v) of the light emitted.

E = h * v

Einstein further deduced, since these electrons emitted light in “discrete ‘jumps’ from one ‘allowed’ energy value to another, the light emitted in the jumps must itself consist of separate particles,” that are called photons (Harvey 6). Even though this conclusion was contradictory to the wave theory of light, it could not be explained otherwise. This paradox was further upheld by the photoelectric effect that showed that “the emission of electrons from many substances under the influence of light” (Harvey 7).

During 1913, Bohr merged Planck’s Quantum Theory and Rutherford’s atomic model by proposing specific electron orbits that were stable, where electrons maintained a constant amount of energy “for an indefinite period of time.” Even though this idea was not substantiated, Bohr was able to mathematically prove that electrons could be raised to a higher energy orbit “by the absorption of a [photon] by the atom;” and vice versa, electrons could be decreased to a lower energy orbit by the emission of a photon. Thus, Bohr extenuated Planck’s Quantum Theory equation, with the use of initial and final energies of the atom, E2 and E1 respectively.

v = (E2 - E1) / h

As a result, information about electrons and their orbits were learned about; however, this only caused more questions, such as why electrons do not continuously radiate energy as the revolve in their orbits (Harvey 7).


In 1924, de Broglie suggested that light waves also act as if they have particle properties; so particles, such as elections, should have some wave properties. However, it was not until 1927 that Schrodinger realized the similarities of vibrating system, such as violin and guitar strings, that contain both the properties of wavelength and frequency (Harvey 10).

About 1925, Pauli illustrated how all atoms can be signified by four distinct quantum numbers (Harvey 16). However, Heisenberg’s uncertainty principle stated that it was impossible to determine the exact “dynamical variables,” quantum numbers, of a system at the same time; thus making electrons impossible to locate at a specific time (Harvey 18).

It has long been known that the nucleus of the hydrogen atom consists of a positively charged particle with relatively large mass. This particle, called the proton, has the exact inverse of the electrical charge of an electron. However, not all the mass of all atoms is just composed of protons. During 1932, Chadwick discovered another type of particle with significant mass in polonium when “bombard[ed]” with beryllium. This particle was observed to have a mass almost equal to that of a proton, yet with no electrical charge. Thus, Chadwick rightfully named the particle a neutron. With this new discovery, the neutron was “incorporated” into the composition of the atomic nuclei. As a result of much research, the atomic number of each atom is equivalent to the number of protons it carries. And for each proton in an atom, there exists a complementary neutron, with the exception of the hydrogen atom that has one proton and no neutrons (Harvey 19).

In conclusion, as scientists learn more about the complexities of nuclear chemistry of the atomic model, more questions arise about the uncertainties of the atomic model. However even with the knowledge scientists now have, the world desperately needs “the wisdom and initiative to help the world avoid the fate of nuclear conflict… and to safely exploit the peaceful benefits of nuclear processes” (Arnikar Foreward).

by Phil for Humanity
on 20091001


References

  1. Arnikar, H. J. Essentials of Nuclear Chemistry. Second Edition. John Wiley and Sons: New York, 1982.
  2. Gillespie, Ronald J. Chemistry. Second Edition. Prentice Hall: New Jersey, 1986.
  3. Harvey, Bernard G. Introduction to Nuclear Physics and Chemistry. Prentice Hall: New Jersey, 1962.


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