Rutherford model of the atom

Can the Rutherford model of the atom explain the emission spectra of elements ... 

The 3rd Solvay Conference in Physics took place in 1921, after a long interruption due to the First World War. Its theme was Atoms and Electrons. 20 It was centered on the Rutherford model of the atom and Niels Bohr s atomic theory. Bohr, however, was not able to attend the conference because of illness. 

Just as the Rutherford model of the atom developed in 1911 was scientifically startling with its revelation of the atom as mostly empty space, so was the Bohr model of the atom introduced in 1913 with its definition of the location of the electron within the atom. As Bohr and others realized that the atomic spectrum of each element is caused by electrons changing energy levels, a different picture of the atom emerged. The new picture of the atom had electrons at various energy levels within the empty space of Rutherford s model (Figure 8.6). This space can still be said to be empty because the mass of the electrons is extraordinarily small in comparison with that of the whole atom. 

The Rutherford model of the atom, in turn, was replaced only two years later by a model developed by Niels Bohr, a Danish physicist. The Bohr model, which is shown in Figure 16, describes electrons in terms of their energy levels. 

Rutherford was able to determine the magnitudes of the positive charges on the atomic nuclei. The pictme of atomic structure that he developed is called the Rutherford model of the atom. 

The Rutherford model of the atom is consistent with the evidence presented so far, but it has some serious limitations. It does not answer such important questions as Why do different elements have such different chemical and physical properties Why does chemical bonding occur at all IVhy does each element form compounds with characteristic formulas How can atoms of different elements give off or absorb light only of characteristic colors (as was known long before 1900) ... 

In 1911, Rutherford s alpha-particle scattering experiments were controversial. In the Rutherford model of the atom, all of the positive charge was crammed into the dense, tiny nucleus. Like charges repel, so the nucleus of the atoms should not be stable, yet it was. The relationships of classical physics that worked so well in explaining large-scale systems did not work on atom-sized systems. Thus, someone had to develop a new approach to understanding the atom. The breakthrough that was needed was the development of the field of study now known as quantum mechanics. 

The Rutherford model of the atom was an improvement over previous models, but it was inoomplete, It did not explain how the atom s negatively oharged electrons are distributed in the space surrounding its positively charged nucleus. After all, it was well known that oppositely charged particles attract each other. So what prevented the negative electrons from being drawn into the positive nucleus ... 

For nearly half a century, Mendeleev s periodic table remained an empirical compilation of the relationship of the elements. Only after the first atomic model was developed by the physicists of the early twentieth century, which took form in Bohr s model, was it possible to reconcile the involved general concepts with the specificity of the chemical elements. Bohr indeed expanded Rutherford s model of the atom, which tried to connect the chemical specificity of the elements grouped in Mendeleev s table with the behavior of electrons spinning around the nucleus. Bohr hit upon the idea that Mendeleev s periodicity could... 

The first detailed model of the atom, proposed by J. J. Thomson in 1898, was based upon the expectation that the atom was a sphere of positive electricity in which electrons were embedded like plums in a pudding. This picture of the atom was not particularly satisfying because it was not useful in predicting or explaining the chemical properties of the atom. Finally, in 1911, a series of experiments performed in the McGill University laboratory of Ernest Rutherford showed that Thomson s picture of the atom had to be abandoned. 

FIGURE 1.6 Rutherford s model of the atom explains why most u particles pass almost straight through the platinum foil, whereas a very few—those scoring a direct hit on the nucleus—undergo verv large deflections. Most of the atom is nearly empty space thinly populated by the atom s electrons. The nuclei are much smaller relative to their atoms than shown here. 

The discoveries of Becquerel, Curie, and Rutherford and Rutherford s later development of the nuclear model of the atom (Section B) showed that radioactivity is produced by nuclear decay, the partial breakup of a nucleus. The change in the composition of a nucleus is called a nuclear reaction. Recall from Section B that nuclei are composed of protons and neutrons that are collectively called nucleons a specific nucleus with a given atomic number and mass number is called a nuclide. Thus, H, 2H, and lhO are three different nuclides the first two being isotopes of the same element. Nuclei that change their structure spontaneously and emit radiation are called radioactive. Often the result is a different nuclide. 

Schematic drawing of an atom, showing a central, positive nucleus surrounded by a cloud of electrons. This model of the atom is consistent with the results of Rutherford s scattering experiments.

Sir Ernest Rutherford (1871-1937 Nobel Prize for chemistry 1908, which as a physicist he puzzled over) was a brilliant experimentalist endowed with an equal genius of being able to interpret the results. He recognized three types of radiation (alpha, beta, and gamma). He used scattering experiments with alpha radiation, which consists of helium nuclei, to prove that the atom is almost empty. The diameter of the atomic nucleus is about 10 000 times smaller than the atom itself. Furthermore, he proved that atoms are not indivisible and that in addition to protons, there must also be neutrons present in their nucleus. With Niels Bohr he developed the core-shell model of the atom. 

The first steps toward the understanding of the nature of the chemical bond could not be taken until the composition and structure of atoms had been elucidated. The model of the atom that emerged from the early work of Thomson, Rutherford, Moseley, and Bohr was of... 

A little earlier, in 1903 (Lenard 1903), Philipp Eduard Anton von Lenard (1862-1947) had carried out some scattering experiments in which he bombarded various metallic foils with high-energy cathode rays. He observed that the majority of electrons passed through the foils undeflected - from this he concluded that the majority of the volume occupied by the metallic atoms must be empty space. This idea was more fully developed by Rutherford (1911), who proposed the nuclear model of the atom which, despite much further elaboration, we still use today for the most basic explanations. 

In the early part of the twentieth century, then, a simple model of atomic structure became accepted, now known as the Rutherford nuclear model of the atom, or, subsequently, the Bohr-Rutherford model. This supposed that most of the mass of the atom is concentrated in the nucleus, which consists of protons (positively charged particles) and neutrons (electrically neutral particles, of approximately the same mass). The number of protons in the nucleus is called the atomic number, which essentially defines the nature of... 

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. 

According to Rutherford s model of the atom and nineteenth-century physics, electrons could emit or absorb any quantity of energy. 

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