Atomic Spectra

In 1666 Newton allowed a thin beam of sunlight to fall upon a prism. The light spread into a rainbow of colors providing a "pleasing divertissement to view the vivid and intense colors produced thereby". He then recombined the colored rays using another prism and again got back the original white light. Clearly, white light was really a mixture of all colors combined. By combining a telescope with a prism, Fraunhofer, in 1814,  was able to spread sunlight out enough to observe dark lines mixed within the rainbow of colors, and began the study of spectroscopy. Further work by Bunsen and Kirchoff in 1859 explained the dark lines. They noted that when a flame containing sodium was observed with a spectrometer, two bright yellow lines (among others) was always present. These lines, previously noticed by Fraunhofer in a candle flame, were the identifying mark of sodium. From this came the fundamental laws of spectroscopy: Each element has a unique set of spectral lines, and these can exist as bright colored lines from a source emitting light (e.g. a gas excited by an electrical discharge, as in a neon sign), or as dark lines against a continuum of colors (e.g. when white light is shown through a cool vapor of neon).

It was time when Newton's laws had solidified mechanics and Maxwell's electromagnetic theory had united the fields of electricity, magnetism and light. There was, in the minds of some physicists, left little remaining to do in their field. The matter of understanding the fundamental principles behind the spectral patterns arising from the elements was the first step that led to a new and totally unexpected physics.

The Electron

As the 19th century was coming to a close, three new phenomena had caught the attention of physicists. Rontgen's X-rays, created when high voltages were imposed on a vacuum tube, Becquerel's radioactivity from which new rays were emitted from uranium and the newly identified (by Marie Curie) element radium, and the electrical discharge in rarified atmospheres that eventually led to the discovery of the electron by J. J. Thompson. The radioactive rays were named alpha rays, beta rays and gamma rays. Gamma rays were quickly associated with the X-rays of Rontgen, but the other two had not been observed elsewhere.

Thompson's experiments were built on the popular but unexplained experiments using "Crook's tubes": A glass tube has the air pumped from it, and a high voltage connected between electrodes at either ends of the tube. As the voltage is increased the tube glows brightly. If a mask is inserted in the tube its shadow appears on one end, a clear indicator that a beam of particles is being projected from one end of the tube (the negative, or cathode end) to the other (positive, or anode) end. Because these "cathode rays" carried an electric charge, Thompson was able to deflect  with electric and magnetic fields. From this he conjectured that these cathode rays were in fact particles of matter that were "atoms, or molecules, or matter in a still finer state of subdivision?", and was able to determine their ratio of charge to mass. In a later elegant experiment Robert Millikan balanced the downward gravitational force on tiny charged droplets of oil by an upward electrical force and thereby measured the electrical charge. Once this was done the mass was immediately known and the beginnings of a quantitative assessment of the structure of the atom was under way.

Thompson's Atom - and Rutherford's Correction

If each atom had its own signature in the form of spectral lines, what was the source of these lines? Thompson thought he had the answer. Maxwell had shown that an accelerating charge radiated energy. Of cathode rays Thompson wrote "I can see no escape from the conclusion that they are charges of negative electricity carried by particles of matter." These rays (now particles in Thompson's view, later named electrons) when made to oscillate could account for the atomic spectra. 

Building on a model proposed by Kelvin in 1902, Thompson's model of the atom consisted of a region of positive charge within which were embedded the negatively charged electrons. The positive distribution of charge just balanced the charge of the electrons as was required to assure overall neutrality of the atom. When the atom was excited (by a flame, perhaps) the resulting oscillations gave rise the spectra. If the excitation were great enough (as in the case of Crook's tube experiments) cathode rays would result. Many problems arose when this model was put to the test, but it was a significant beginning. Thompson's atom was constructed using the laws of Newton and Maxwell, and there where questions unanswered: His model used the mutually repelling force of the electrons within the positive cloud to establish stable atoms, but such a model was hard pressed to explain the necessary radiation patterns needed to match his theory with the observed spectral lines. A theory must be mathematically as well as conceptually plausible. Thompson's work was done while he was director of the Cavendish Laboratory at Cambridge University where, in 1890 he was joined by Ernest Rutherford, a young physics student from New Zealand.

Thompson's Atom

Rutherford began to investigate the recently discovered properties of radioactivity. One characteristic of the radioactive materials discovered by Henri Becquerel was that when the emissions from a radioactive source were formed into a beam, and the beam allowed to pass through electric and/or magnetic fields, the beam would be separated into three distinct portions labeled alpha, beta and gamma. The deflections resulting from the fields determined that the alpha beams consisted of massive, positively charged particles (later determined to be Helium nuclei), the beta beams were light, negatively charged particles (Thompson's electrons) and the gamma rays were uncharged (and so difficult to characterize, but eventually were determined to be high energy electromagnetic radiation). A diagram of the apparatus is shown to the right.

The alpha particles carried a charge twice that of the electron and were about 8,000 times heaver. Because of their mass and charge they were used as projectiles to bombard thin targets and so react with the atoms of the target material. Rutherford directed a beam of alpha particles at a thin film of gold, in an attempt to explore the character of Thompson's atomic model. one would expect that the atoms of the foil would deflect the alpha particles over a range of angles as they passed through the positive boy of the atoms. Instead, most of the particles went through with only small deflections, but an occasional alpha particle would be deflected almost directly back towards the source. This was completely foreign to what was predicted by the Thompson model, which caused Rutherford to write "It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." To explain the effect he proposed a different model for the atom in which the positive charge was concentrated in a tiny region at the center, and the electrons orbited about, much as planets orbit the Sun. This model, familiar to all school children, is shown here.

While Rutherford's model was successful in explaining his scattering experiment, it had a fateful flaw: Unlike planets orbiting our Sun, the electrons orbiting Rutherford's nucleus carried an electrical charge. As these charges go around, they accelerate and so, according the electromagnetic theory, must radiate energy. This process is paramount in, for example, the transmission of radio and television signals. This flaw was to be removed only at the cost of the sanctity of the laws of Newton and Maxwell.

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