Part 2

De Morgan evidently did not keep up with the latest developments in science, however, because two years before his birth, John Dalton, an English schoolteacher, had changed the atomic theory of matter from a philosophical speculation into a firmly established principle. The evidence that convinced Dalton and many other contemporary scientists of the reality of atoms came from quantitative chemical analysis.

Dalton knew that many chemical substances could be separated into two or more simpler substances. Chemicals that could be separated further were called compounds; those that could not were called elements. Careful experiments by Dalton and others showed that whenever two or more elements combined chemically to make a compound the relative amounts of the elements had to be carefully adjusted to fit a definite proportion in order to have no elements left over after the reaction was finished. For example, if hydrogen and oxygen were combined to form water, the weight of oxygen had to be eight times the weight of hydrogen; otherwise, either some hydrogen or some oxygen would be left over.

This fundamental truth is now called the Law of Definite Proportions. Another important principle, called the Law of Multiple Proportions, is illustrated by hydrogen peroxide, which is made up of the same two elements that are found in water. The weight of oxygen in hydrogen peroxide, however, is 16 times the weight of hydrogen or exactly twice the relative weight found in water.

These principles of chemical combination convinced Dalton that each chemical element consists of small, indivisible units, all just alike, called atoms, and that each chemical compound also has basic units, called molecules, which cannot be divided without reducing the compound into its elements—that is, destroying it as a compound. He visualized a molecule of a compound as formed by the uniting of individual atoms of two or more elements. It was obvious to him that in any molecule of a compound, the weight of each atom of a component element bore a proportionate relationship to the weight of the entire molecule which was equal to the proportion, by weight, of all that element in the compound. And although Dalton had no idea how heavy any individual atom really was, he could tell how many times heavier or lighter it was than an atom of another element.

Incidentally, Dalton mistakenly thought that one atom of oxygen was eight times as heavy as one atom of hydrogen instead of 16 times as heavy. He assumed a water molecule to be HO instead of H₂O.

Cathode Rays Show Atoms Contain Smaller Parts

Curiosity about the fundamental nature of matter was matched by equally avid curiosity about the fundamental nature of electricity. Before 1850 much had been learned about the behavior of electric charge and electric currents flowing through solids and liquids. Real progress in understanding electric charge, however, had to wait for the development of highly efficient vacuum pumps.

About 1854 Heinrich Geissler, a German glassblower, developed an improved suction pump, and also succeeded in sealing into a glass tube two wires attached to metal electrodes inside the tube. Experimenters were then able to study the flow of electricity through a near-vacuum. A Geissler tube is diagramed in Figure 1.

By the 1890s it had become clear that the flow of electricity through a highly evacuated tube consisted of a negative electric charge moving at a very high speed along straight lines between sealed-in electrodes. Since it originated at the negative electrode, or cathode, the invisible stream of charge was named “cathode rays.”

CURRENT SOURCE CATHODE (-) STREAM OF ELECTRONS VACUUM PUMP ANODE (+)

Although many investigators contributed to knowledge about cathode rays, the experiments of Joseph J. Thomson, a British physicist, are generally considered to have been the most enlightening. Thomson arranged a cathode-ray tube so that the rays could be deflected by magnets and by electrically charged metal plates. By applying certain well-known principles of physics, he was able to confirm an impression already held by physical chemists, namely, that electric charge, like matter, was “atomized”—the stream of charge consisted of a swarm of very small particles, all alike. He succeeded also in determining that the speed of the particles was about one-tenth the speed of light.

Probably Thomson’s most significant result was determining the ratio of the charge of each little particle to its weight. He was able to do this by measuring the magnetic force required to divert a stream of charged particles. (You can do this experiment yourself with relatively simple equipment.) This charge-to-weight ratio proved to be nearly 2000 times greater than the already known charge-to-weight ratio for a positively charged hydrogen atom, or ion, which until then was thought to be the lightest constituent of matter. It remained to be determined whether charge or weight caused the difference. Further experimentation showed that the charges were approximately the same amount in the two cases. It was therefore proven that the weight of the hydrogen atom, lightest of all the atoms, was nearly 2000 times as great as the weight of one of the little negative particles.

The name “electron” was given to the small negative particles identified by Thomson. Since the electrons had come from the cathode, it was apparent that the atoms in the cathode must contain electrons. Thomson reasoned that electric current in a wire is a stream of electrons passing successively from atom to atom and that the difference between an electrically charged atom and a neutral atom is that the charged one has gained or lost one or more electrons.

Radioactive Atoms Discovered

In 1896 the French physicist Henri Becquerel was investigating the relation between fluorescence and X rays, a puzzling kind of penetrating radiation discovered a few months earlier by the German, Wilhelm Roentgen. Various chemical compounds fluoresce, or glow, when exposed to ultraviolet rays and other types of radiation. While experimenting with a large number of chemicals, Becquerel discovered, quite by accident, that a compound containing the element uranium can, without being exposed to any kind of radiation, darken a photographic plate completely wrapped in heavy black paper.

Although no one realized it at the time, Becquerel had discovered that atoms of some elements will at random times transform themselves into atoms of a different element by emitting certain extremely high-speed charged particles. Atoms that can do this are said to be radioactive, and it was the radiation from transforming uranium atoms that darkened Becquerel’s photographic plate.

Rutherford Finds the Atomic Nucleus

We are greatly indebted to the imagination and experimental skill of the British physicist Ernest Rutherford for the interpretation of radioactivity in terms of the structure of atoms.

Rutherford, born and educated in New Zealand, moved to England to work under Thomson at Cambridge University in 1895. Shortly afterward, Wilhelm Roentgen in Germany discovered X rays, Becquerel in France discovered radioactivity, and Thomson proved the existence of the electron.

During the next few years, curiosity about the fundamental nature of radioactivity led a number of people to do a great deal of work. The element thorium was found to be radioactive, and Marie and Pierre Curie discovered two new elements, polonium and radium, that were also radioactive. The radiation from radioactive materials was found to be of three kinds called alpha rays, beta rays, and gamma rays. Alpha rays were first detected by Rutherford, who later identified them as positively charged helium atoms. Becquerel demonstrated that beta rays, like cathode rays, consist of negatively charged electrons. The highly penetrating gamma rays were proved by Rutherford and E. N. da C. Andrade to be electromagnetic radiation similar to X rays.

Rutherford, in collaboration with the English chemist Frederick Soddy, brought order out of a chaos of puzzling discoveries by establishing the general behavior of radioactive atoms. He determined that certain naturally occurring atoms of high atomic weight can spontaneously emit an alpha or a beta particle and thereby convert themselves into new atoms. These new atoms, being also radioactive, sooner or later convert themselves into still different atoms, and so on. Each time an alpha particle is emitted in this sequence, the new atom is lighter by the weight of the alpha particle, or helium atom. The disintegration process proceeds from stage to stage until at last a stable atom is produced. The end product in this “decay” process in naturally occurring radioactive elements is lead.

One experiment by Rutherford and his co-workers had a most profound effect on the understanding of atomic structure. What they did was to direct a stream of alpha particles at a thin piece of gold foil. The results were astonishing. Almost all the particles passed straight through the foil without changing direction. Of the few particles that did ricochet in new directions, however, some were deflected at very sharp angles. (See Figure 2.)

As a result of this experiment, Rutherford proposed a concept of the atom entirely different from the one which prevailed at this time. The prevailing notion was one advanced by Thomson which conceived of an atom as a blob of positive electric charge in which were imbedded, in much the same way as plums are in a pudding, enough electrons to neutralize the positive charge. Rutherford’s concept, which quickly set aside Thomson’s “plum pudding” model, was that an atom has all of its positive charge and virtually all of its mass concentrated in a tiny space at its center. (Collisions with this center, which came to be known thereafter as the nucleus, had been responsible for the sharp changes in direction of some of the alpha particles.) The space surrounding this nucleus is entirely empty except for the presence of a number of electrons (79 in the case of the gold atom), each about the same size as the nucleus.

To illustrate Rutherford’s concept, let us imagine a gold atom magnified so that it is as large as a bale of cotton. The nucleus at the center of this large atom would be the size of a speck of black pepper. If this imaginary bale weighed 500 pounds, the little speck at its center would weigh 499¾ pounds; the surrounding cotton (corresponding to empty space in Rutherford’s concept) containing the 79 electrons would weigh but ¼ pound. To express this idea another way, any object such as a gold ring, as dense and solid as it may seem to us, consists almost entirely of nothing!

The Proton Is Recognized

Rutherford’s discovery aroused intense curiosity about the nature and possible structure of this extremely small, but all-important, part of an atom. It was assumed that the positive charge carried by the nucleus must be a whole-number multiple of a small unit equal in size but opposite in sign to the charge of an electron. This conclusion was based on the information that all atoms contain electrons and that an undisturbed atom is electrically neutral. Since it was known that a neutral atom of hydrogen contains just one electron, it appeared that the charge on a hydrogen nucleus must represent the fundamental unit of positive charge, some multiple of which would represent the charge on any other nucleus. Several lines of investigation combined to establish quite firmly that nuclei of atoms occupying adjacent positions on the periodic chart of the elements differed in charge by this fundamental unit. Since the hydrogen nucleus seemed to play such an important role in making up the charges of all other nuclei, it was given the name proton from the Greek “protos,” which means “first.”

Isotopes Are Discovered

At a historic meeting of the British Association for the Advancement of Science held in Birmingham, England, in 1913, two apparently unrelated lines of investigation were reported, each of which showed that some atomic nuclei have identical electric charges but different weights.