Rutherford’s atomic model

Rutherford performed several calculations that led him to an inescapable conclusion the atom is made up mainly of empty space, with a small, massive region of concentrated charge at the centre. Soon afterward, the charge on this central region was determined to be positive, and was named the atomic nucleus. Because Rutherford s atomic model, shown in Figure 3.5 on the next page, pictures electrons in motion around an atomic nucleus, chemists often call this the nuclear model of the atom. You may also see it referred to as a planetary model because the electrons resemble the planets in motion around a central body. 

Rutherford s atomic model solved problems inherent in Thomson s atomic model, but it also raised others. For example, an atomic nucleus composed entirely of positive charges should fly apart due to electrostatic forces of repulsion. Furthermore, Rutherford s nuclear atom could not adequately explain the total mass of an atom. The discovery of the neutron, in 1932, eventually helped to settle these questions. 

There was a more significant problem, however. Rutherford s atomic model seemed to contradict the laws of nineteenth-century physics. According to these assumptions, an electron in motion around a central body must continuously give off radiation. Consequently, one should be able to observe a continuous spectrum (a rainbow ) of light energy as the electron gives off its radiation. 

Q WSBM Why did Rutherford s atomic model require modification Be as specific as possible in your answer. 

Summarize Rutherford s atomic model. Which aspects of this model are deficient or false as far as today s understanding is concerned ... 

Rutherford s atomic model Just like the solar system, an atom has a nucleus, consisting of positively charged particles, around which electrons orbit. 

Ellen Gleditsch, Rutherfords atommodel og de radioaktive grundstoffe [Rutherford s atomic model and the radioactive elements]. Archiv for mathematik og naturviden-skab B.XXXIV, no. 6 (1915) 3-28. 

Scientists of the nineteenth century lacked the concepts necessary to explain line spectra. Even in the first decade of the twentieth century, a suitable explanation proved elusive. This changed in 1913 when Niels Bohr, a Danish physicist and student of Rutherford, proposed a new model for the hydrogen atom. This model retained some of the features of Rutherford s model. More importantly, it was able to explain the line spectrum for hydrogen because it incorporated several new ideas about energy. As you can see in Figure 3.8, Bohr s atomic model pictures electrons in orbit around a central nucleus. Unlike Rutherford s model, however, in which electrons may move anywhere within the volume of space around the nucleus, Bohr s model imposes certain restrictions. 

Dalton s atomic model does not include negatively charged (-) electrons and positively charged (+) protons. Thomson discovered the electron in 1897, while the discovery of the proton was made by Rutherford in 1919. We can summarize Thomson s ideas as follow ... 

Rutherford made important contributions to the explanation of atomic structure. He discovered the nucleus in 1911 and the proton in 1919. Prior to Rutherford, Thomson s atomic model was valid. His model stated that the atom was a sphere in which electrons and protons were moved arbitrarily. But there was an important question about how these protons and electrons were distributed. Was there any regularity or were they moving arbitrarily The answer to this question could not yet be seen. In order to get answers to these problems and to verify Thomson s atomic model, Rutherford proposed a model resulting from his a - particle experiment. 

Rutherford s atom consisted of a positively charged center some 10,000 times smaller than the atom itself. This center also carried most of the mass of the atom. For the gold atom, he found the charge at the center to be approximately 100 times the charge of the electron. Surrounding this center of positive charge were the electrons (see Figure 4.2). In March 1911, this new model of the atom was conveyed to the community of science. Later, in October 1912, Rutherford used the term nucleus for the first time. 

Rutherford s nuclear model of the atom explains the results of the gold foil experiment. Most alpha particles pass straight through, being only slightly deflected by electrons, if at all. The strong force of repulsion between the positive nucleus and the positive alpha particles causes the large deflections. 

List the strengths and weaknesses of Rutherford s nuclear model of the atom. (4.2)... 

Many scientists in the early twentieth century found Rutherford s nuclear atomic model to be fundamentally incomplete. To physicists, the model did not explain how the atom s electrons are arranged in the space around the nucleus. Nor did it address the question of why the negatively charged electrons are not pulled into the atom s positively charged nucleus. Chemists found Rutherford s nuclear model lacking because it did not begin to account for the differences in chemical behavior among the various elements. 

We urge you to keep Rutherford s planetary model of the atom (Section 1.4) in mind while reading this chapter. That model, with its consideration of the electrical forces within atoms and molecules, provides the foundation of the entire subject of chemical bonding and molecular structure. 

Rutherford s planetary model of the atom assumes that an atom of atomic number Z comprises a dense, central nucleus of positive charge +Ze surrounded by a total of Z electrons moving around the nucleus. The attractive forces between each electron and the nucleus, and the repulsive forces between the electrons, are described by Coulomb s law. We first discuss Coulomb s law in general terms, and then apply it to the planetary atom. 

Bohr supplemented Rutherford s planetary model of the atom with the assumption that an electron of mass moves in a circular orbit of radius r about a... 

The incompatibility of Rutherford s planetary model, based soundly on experimental data, with the principles of classical physics was the most fundamental of the conceptual challenges facing physicists in the early 1900s. The Bohr model was a temporary fix, sufficient for the interpretation of hydrogen (H) atomic spectra as arising from transitions between stationary states of the atom. The stability of atoms and molecules finally could be explained only after quantum mechanics had been developed. 

The observation of line spectra did not correlate with classical theory for one major reason. As was mentioned in the chapter introduction, if an electron spiraled closer to the nucleus, it should emit radiation. Moreover, the frequency of the radiation should be related to the time of revolution. On the spiral path inward, that time should decrease smoothly, so the frequency of the radiation should change smoothly and create a continuous spectrum. Rutherford s nuclear model seemed totally at odds with atomic line spectra. 

Realizing that Rutherford s planetary model of the atom is incompatible with the classical Maxwell theory of radiation—which stipulates that a charged electron in circular motion will continually emit radiation and thereby lose energy, its orbit will shrink, and it will quickly spiral into the nucleus—Niels Bohr in 1913 (see Fig. 3.25) asserted that an electron in an atomic orbit simply does not radiate in other words Maxwell s theory does not apply at this level. Bohr s main contribution was to make two nonclassical assumptions. ... 

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