CoreChem:4.3 The Nuclear Atom
From ChemEd Collaborative
4.3 THE NUCLEAR ATOM
Radioactivity
Just prior to the turn of the twentieth century, additional observations were made which contradicted parts of Dalton’s atomic theory. The French physicist Henri Becquerel (1852 to 1928) discovered by accident that compounds of uranium and thorium emitted rays which, like rays of sunlight, could darken photographic films. Becquerel’s rays differed from light in that they could even pass through the black paper wrappings in which his films were stored.
Although themselves invisible to the human eye, the rays could be detected easily because they produced visible light when they struck phosphors such as impure zinc sulfide. Such luminescence is similar to the glow of a psychedelic poster when invisible ultraviolet (black light) rays strike it.
Further experimentation showed that if the rays were allowed to pass between the poles of a magnet, they could be separated into the three groups shown in Fig. 4.2. Because little or nothing was known about these rays, they were labeled with the first three letters of the Greek alphabet. Upon passing through the magnetic field, the alpha rays (α rays) were deflected slightly in one direction, beta rays (β rays) were deflected to a much greater extent in the opposite direction, and gamma rays (γ rays) were not deflected at all.
Deflection by a magnet is a characteristic of electrically charged particles (as opposed to rays of light). From the direction and extent of deflection it was concluded that the β particles had a negative charge and were much less massive than the positively charged α particles. The γ rays did not behave as electrically charged particles would, and so the name rays was retained for them. Taken together the α particles, β particles, and γ rays were referred to as radioactivity, and the compounds which emitted them as radioactive.
Study of radioactive compounds by the French chemist Marie Curie (1867 to 1934) revealed the presence of several previously undiscovered elements (radium, polonium, actinium, and radon). These elements, and any compounds they formed, were intensely radioactive. When thorium and uranium compounds were purified to remove the newly discovered ele- ments, the level of radioactivity decreased markedly. It increased again over a period of months or years, however. Even if the uranium or thorium compounds were carefully protected from contamination, it was possible to find small quantities of radium, polonium, actinium, or radon in them after such a time.
To chemists, who had been trained to accept Dalton’s indestructible atoms, these results were intellectually distasteful. The inescapable conclusion was that some of the uranium or thorium atoms were spontaneously changing their structures and becoming atoms of the newly discovered elements. A change in atomic structure which produces a different element is called transmutation. Transmutation of uranium into the more radioactive elements could explain the increased emission of radiation by a carefully sealed sample of a uranium compound.
During these experiments with radioactive compounds it was observed that minerals containing uranium or thorium always contained lead as well. This lead apparently resulted from further transmutation of the highly radioactive elements radium, polonium, actinium, and radon. The lead found in uranium ores always had a significantly lower atomic weight than lead from most other sources (as low as 206.4 compared with 207.2, the accepted value). Lead associated with thorium always had an unusually high atomic weight. Nevertheless, all three forms of lead had the same chemical properties. Once mixed together, they could not be separated. Such results, as well as the reversed order of elements such as Ar and K in the periodic table, implied that atomic weight is not the fundamental determinant of chemical behavior.
The Electron
Near the middle of the nineteenth century the English chemist and physicist Michael Faraday (1791 to 1867) established a connection between electricity and chemical reactions. He already knew that an electric current flowing into certain molten compounds through metal plates called electrodes could cause reactions to occur. Samples of different elements would deposit on the electrodes. Faraday found that the same quantity of electric charge was required to produce 1 mol of any element whose valence was 1. Twice that quantity of charge would deposit 1 mol of an element whose valence was 2, and so on. Electric charge is measured in units called coulombs, abbreviated C. One coulomb is the quantity of charge which corresponds to a current of one ampere flowing for one second. It was found that 96 500 C of charge was required to deposit on an electrode l mol of an element whose valence is l.
Faraday’s experiments strongly suggested that electricity, like matter, consists of very small indivisible particles. The name electron was given to these particles, and an electric current came to be thought of as a flow of electrons from one place to another. When such a current flows into a chemical compound, one electron is required for each atom of a univalent element deposited on an electrode, two electrons for each atom of an element whose valence is 2, and so on. Thus an electric charge of 96 500 C corresponds to 1 mol of indivisible electric particles (electrons).
The relationship between electricity and atomic structure was further clarified by experiments involving cathode-ray tubes in the 1890s. A cathode-ray tube can be made by pumping most of the air or other gas out of a glass tube and applying a high voltage to two metal electrodes inside. If ZnS or some other phosphor is placed on the glass at the end of the tube opposite the negatively charged electrode (cathode), the ZnS emits light. This indicates that some kind of rays are streaming away from the cathode. When passed between the poles of a magnet, these cathode rays behave the same way as the β particles described earlier. The fact that they were very small electrically charged particles led the English physicist J. J. Thomson (1856 to 1940) to identify them with the electrons of Faraday’s experiments. Thus cathode rays are a beam of electrons which come out of the solid metal of the cathode. They behave exactly the same way no matter what the electrode is made of or what gas is in the tube. These observations allow one to conclude that electrons must be constituents of all matter.
In addition to being deflected by a magnet, the electron beam in a cathode-ray tube can be attracted toward a positively charged metal plate or repelled from a negative plate. By adjusting such electrodes to exactly cancel the deflection produced by a magnet of known strength, Thomson was able to determine that the ratio of charge to mass for an electron is 1.76 × 108 C/g. This is a rather large ratio. Either each electron has a very large charge, or each has a very small mass. We can see which by using Faraday’s result that there are 96 500 C mol–1 of electrons
Thus the molar mass of an electron is 5.48 × 10–4 g mol–1, and if we think of the electron as an “atom“(or indivisible particle) of electricity, its atomic weight would be 0.000548—only [[Image:]] that of hydrogen, the lightest element known. In 1909 the American physicist Robert A. Millikan (1863 to 1953) was able to determine the charge on an electron independently of its mass. His value of 1.6 × 10–19 C can be combined with Thomson’s charge-to-mass ratio to give an independent check on the molar mass for the electron
thus confirming that the electron has much less mass than the lightest atom. (The quantity 1.6 × 10–19 C is often represented by the symbol e. Thus the charge on a single electron is
–e = –1.6 × 10–19 C. The minus sign indicates that the electron is a negatively charged particle.)
The Nucleus
The results of Thomson’s and other experiments implied that electrons were constituents of all matter and hence of all atoms. Since macroscopic samples of the elements are found to be electrically neutral, this meant that each atom probably contained a positively charged portion to balance the negative charge of its electrons. In an attempt to learn more about how positive and negative charges were distributed in atoms, Ernest Rutherford (1871 to 1937) and his coworkers performed numerous experiments in which α particles emitted from a radioactive element such as polonium were allowed to strike thin sheets of metals such as gold or platinum.
It was already known that the α particles carried a positive charge and traveled rapidly through gases in straight lines. Rutherford reasoned that in a solid, where the atoms were packed tightly together, there would be numerous collisions of α particles with electrons or with the unknown positive portions of the atoms. Since the mass of an individual electron was quite small, a great many collisions would be necessary to deflect an α particle from its original path, and Rutherford’s preliminary calculations indicated that most would go right through the metal targets or be deflected very little by the electrons. In 1909, confirmation of this expected result was entrusted to Hans Geiger and a young student, Ernest Marsden, who was working on his first research project.
The results of Geiger and Marsden’s work (using apparatus whose design is shown schematically in Fig. 4.3) were quite striking. Most of the α particles went straight through the sample or were deflected very little. These were observed by means of continuous luminescence of the ZnS screen at position 1 in the diagram. Observations made at greater angles from the initial path of the a particles (positions 2 and 3) revealed fewer and fewer flashes of light, but even at an angle nearly 180° from the initial path (position 4) a few α particles were detected coming backward from the target. This result amazed Rutherford. In his own words, “It was quite the most incredible event that has ever happened to me in my life. 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. On consideration, I realized that this scattering backwards must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of an atom was concentrated in a minute nucleus.”1 Rutherford’s interpretation of Geiger and Marsden’s experiment is shown schematically in Fig. 4.4.
Quantitative calculations using these experimental results showed that the diameter of the nucleus was about one ten-thousandth that of the atom. The positive charge on the nucleus was found to be + Ze, where Z is the number which indicates the position of an element in the periodic table. (For example, H is the first element and has Z = 1. He is the second element and Z = 2. The twentieth element in Table 4.2 or Fig. 4.1 is Ca, and the nucleus of each Ca atom therefore has a charge of + 20e = 20 × 1.60 × 10–19 C = 32.0 × 10–19 C.) In order for an atom to remain electrically neutral, it must have a total of Z electrons outside the nucleus. These provide a charge of –Ze to balance the positive nuclear charge. The number Z, which indicates the positive charge on the nucleus and the number of electrons in an atom, is called the atomic number.
The significance of the atomic number was firmly established in 1914 when H. G. Moseley (1888 to 1915) published the results of experiments in which he bombarded a large number of different metallic elements with electrons in a cathode-ray tube. Wilhelm Roentgen (1845 to 1923) had discovered earlier that in such an experiment, rays were given off which could penetrate black paper or other materials opaque to visible light. He called this unusual radiation x-rays, the x indicating unknown. Moseley found that the frequency of the x-rays was unique for each different metal. It depended on the atomic number (but not on the atomic weight) of the metal. (If you are not familiar with electromagnetic radiation or the term frequency, read Sec. 21.1 where they are discussed more fully.) Using his x-ray frequencies, Moseley was able to establish the correct ordering in the periodic table for elements such as Co and Ni whose atomic weights disagreed with the positions to which Mendeleev had assigned them. His work confirmed the validity of Mendeleev’s assumption that chemical properties were more important than atomic weights.
1 Ernest Rutherford, the Development of the Theory of Atomic Structure, in J. Needham and W. Pagel (eds.) “Background to Modern Science,” The Macmillan Company, New York, 1938.
