Classical Atomic Theories

By 1910 it was clear from experiment that atoms contained negatively charged electrons, and that the number of electrons in an atom was roughly A/2 where A was its chemical atomic weight. Since atoms are usually neutral, there must have been an equal amount of positive charge within the atom. Since electrons were known to have a very small mass compared with the mass of the atom, the mass of the atom must have been associated with the positive charge.

The model of J. J. Thomson described the atom as having the electrons embedded in a sphere of positive charge. This model was qualitatively successful in that electromagnetic emission from the atom could be explained in terms of classical harmonic oscillation of the electrons of the atom. Evidence of agreement between predicted and observed emission spectra was not as convincing, however. Emission spectra were found to occur at discrete frequencies characteristic of the atom. Further, scattering of a particles (doubly ionized helium) indicated that the positive charge was concentrated in a very small volume 'nucleus' at the center of the atom.

Figure 1.2. Thomson's plum-pudding model of the atom showing scattering of an a particle.

Rutherford's model [1] of the hydrogen atom in 1911 had one electron orbiting a small heavy proton nucleus, attracted to the nucleus by a Coulomb interaction. This model appeared to agree with the physical structure of the atom, but the model implied that the hydrogen atom would be inherently unstable and there were still problems predicting the electromagnetic emission spectra of the atom. The electron could not have been stationary as it would simply be Coulomb attracted into the nucleus of the atom, resulting in an atomic diameter four orders of magnitude smaller than observed. An electron in a curved orbit is accelerated so must radiate electromagnetic energy. As energy is radiated, the radius of the electron orbit should decrease until the electron collapses into the nucleus. The Rutherford model predicted a typical lifetime of a hydrogen electron to be on the order of 10-10 seconds [2] and the radiated energy should have been continuous in energy.

Figure 1.3. Visible portion of the hydrogen spectrum. After Finkelnburg.

Atoms placed in an excited state, such as by passing electric current through a gas, were found to emit electromagnetic radiation concentrated at discrete frequencies. By passing the light emitted from an excited gas and then diffracting the light through a prism, this characteristic electromagnetic radiation appeared as lines on photographic plates. Similarly, atoms were found to absorb incident electromagnetic radiation at the same characteristic frequencies. The spectra of emitted radiation were found to be quite complicated for some atoms, but quite regular for the hydrogen atom. Attempts were made to identify series of progressively spaced lines in all elements, and empirical equations were derived to describe these series.

The discreteness of atomic electron energy levels was directly demonstrated in 1914 by the Frank-Hertz experiment. Passing electrons through a gas of a given atom to a grid at a given acceleration potential and then collecting the electrons at an anode at a lesser potential, drops in current were found when the acceleration voltage was increased beyond critical values. These drops in current corresponded to points where the electrons were accelerated to sufficient energy to excite electrons of the atom, losing energy in inelastic collisions. The electron energies remaining after the collision were then insufficient to overcome the reverse anode potential.

Clearly a model of the atom was needed that would at least predict the characteristic electromagnetic emission and absorption spectra observed in nature.

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