Sir J J Thompson proposed that the positive charge which corresponded to most of the mass of an atom was smeared throughout the volume of the atom. The electrons were tiny lumps of negative charge embedded in this positively&endash;charged body of the atom. He described the electrons as "plums in a pudding" so this is called Thompson's "plum pudding model" of the atom.
In 1910 and 1911 graduate students Geiger and Marsden, working for Ernest Lord Rutherford at the Cavendish Laboratory in Cambridge, conducted experiments which brought startling results and required great changes in the ideas of atomic structure. Positively-charged alpha particles, from a radioactive source, were collimated into a beam which then struck a gold foil. Gold was chosen because it is more malleable than any other material; this allowed the gold foil to be made extremely thin so the alpha particles would be scattered by a single gold atom as they passed through the foil. The effects of multiple scattering, involving several gold atoms, were greatly minimized by the extreme thinness of the foil. Deflected alphas were detected by the tiny bit of light that was emitted when an alpha struck a zinc sulfide plate. Geiger and Marsden sat in the dark for hours and watched for these tiny specks of light through a simple microscope.
According to Thompson's plum pudding model of atomic structure, the incident alpha particle, which was known to be positively charged, should see a mostly neutral atom. Only when it happened very close to a negative electron would it be deflected. The electric charges within the atom were so dispersed that the deflections were expected to be very small. They were expecting all observations to show deflections of less than four degrees. This is sketched below,
Rutherford's atomic scattering experiments
of 1910 and 1911 drastically changed our view of atomic structure.
Most of the alphas did, indeed, pass through the foil with no scattering or with very little scattering. However, some were observed with scattering angles far in excess of the four degree maximum that had been expected. Some alphas were backscattered almost fully back in the direction from which they had come! This was entirely counter to their expectations.
To explain these large deflections, Rutherford proposed a nuclear model of the atom. In this new model, all of the positive charge and almost all of the mass was located in a very tiny, very dense nucleus at the center of the atom. To agree with the observed deflection data, this tiny speck of positive charge must have a radius of about 5 x 10-15 m. This is incredibly small compared to the radius of an atom which is about 1 x 10-11 m. The nuclear radius must be only about 1/10 000 of the atomic radius; the rest of the atom is "empty space" occupied by the negatively charged electrons carrying less than 1/2 000 of the mass.
Rutherford's model--required to explain experimental results--was quite a departure from Thompson's earlier model. How could the tiny electrons occupy all the remaining volume of the atom? Electric forces would attract the electrons to the positively-charged nucleus. But what could prevent the electrons from just collapsing into the nucleus? Planets are attracted to our Sun yet do not collapse into the Sun because they are in motion, orbiting the Sun. Perhaps something similar occurs within an atom. Orbiting electrons would find their centripetal force provided by the electric force of attraction between themselves and the nucleus. However, classical electromagnetic theory predicts that any accelerating electric charge will radiate energy in the form of EM waves and lose energy itself. Could that lead to predictions of the hydrogen spectrum? No, that would mean the orbiting electrons would immediately radiate EM waves and spiral in to the nucleus! Rutherford's model--startling though it was--provided only a starting point.
(c) Doug Davis, 1997; all rights reserved