Bohrs Atomic Theory

FIGURE 20 Niels Henrik David Bohr (1885-1962). Photo and permission from Edgar Fahs Smith Collection. 

According to Bohr s opinion, the rare-earth group consisted of elements where the four-quantum level was gradually filled up from 18 to 32 electrons. The number of electrons in the five- and six-quantum levels on the other hand remained imchanged. Bohr s quantum theory thus served as a useful explanation for the pronoimced similarity between the chemical and physical properties of the rare-earth elements. He mentioned that their mutual similarity must be ascribed to the fact that we have here to do with the development of an electron group that lies deeper in the atom. He moreover emphasized that lutetium (Z = 71) had to be considered the last rare-earth element. Element 72 on the other hand did not belong to 

FIGURE 21 Bohr s periodic table (1922). Reproduced with permission (Copyright Nobel Foundation, 1922). 

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Whereas zirconium was discovered in 1789 and titanium in 1790, it was not until 1923 that hafnium was positively identified. The Bohr atomic theory was the basis for postulating that element 72 should be tetravalent rather than a trivalent member of the rare-earth series. Moseley s technique of identification was used by means of the x-ray spectra of several 2ircon concentrates and lines at the positions and with the relative intensities postulated by Bohr were found (1). Hafnium was named after Hafma, the Latin name for Copenhagen where the discovery was made. 

N. Bohr, Atomic Theory and the Description of Nature, Cambridge University Press, 1934. 

The concept of chemical periodicity is central to the study of inorganic chemistry. No other generalization rivals the periodic table of the elements in its ability to systematize and rationalize known chemical facts or to predict new ones and suggest fruitful areas for further study. Chemical periodicity and the periodic table now find their natural interpretation in the detailed electronic structure of the atom indeed, they played a major role at the turn of the century in elucidating the mysterious phenomena of radioactivity and the quantum effects which led ultimately to Bohr s theory of the hydrogen atom. Because of this central position it is perhaps not surprising that innumerable articles and books have been written on the subject since the seminal papers by Mendeleev in 1869, and some 700 forms of the periodic table (classified into 146 different types or subtypes) have been proposed. A brief historical survey of these developments is summarized in the Panel opposite. 

The discovery of hafnium was one of chemistry s more controversial episodes. In 1911 G. Urbain, the French chemist and authority on rare earths , claimed to have isolated the element of atomic number 72 from a sample of rare-earth residues, and named it celtium. With hindsight, and more especially with an understanding of the consequences of H. G. J. Moseley s and N. Bohr s work on atomic structure, it now seems very unlikely that element 72 could have been found in the necessary concentrations along with rare earths. But this knowledge was lacking in the early part of the century and, indeed, in 1922 Urbain and A. Dauvillier claimed to have X-ray evidence to support the discovery. However, by that time Niels Bohr had developed his atomic theory and so was confident that element 72 would be a... 

In chemistry, perhaps because of the significance in visualizing molecular strac-ture, there has been a focus on how students perceive three-dimensional objects from a two-dimensional representation and how students mentally manipulate rotated, reflected and inverted objects (Stieff, 2007 Tuckey Selvaratnam, 1993). Although these visualization skills are very important in chemistry, it is evident that they are not the only ones needed in school chemistry (Mathewson, 1999). For example, conceptual understanding of nature of different types of chemical bonding, atomic theory in terms of the Democritus particle model and the Bohr model, and... 

The units we use in daily life, such as kilogram (or pound) and meter (or inch) are tailored to the human scale. In the world of quantum mechanics, however, these units would lead to inconvenient numbers. For example, the mass of the electron is 9.1095 X J0 31 kg and the radius of the first circular orbit of the hydrogen atom in Bohr s theory, the Bohr radius, is 5.2918 X 10 11 m. Atomic units, usually abbreviated as au, are introduced to eliminate the need to work with these awkward numbers, which result from the arbitrary units of our macroscopic world. The atomic unit of length is equal to the length of the Bohr radius, that is, 5.2918 X 10 n m, and is called the bohr. Thus 1 bohr = 5.2918 X 10"11 m. The atomic unit of mass is the rest mass of the electron, and the atomic unit of charge is the charge of an electron. Atomic units for these and some other quantities and their values in SI units are summarized in the accompanying table. 

See the analysis in Helge Kragh, "Bohr s Atomic Theory and the Chemists, 19131925," Rivista di storia della scienza 2 (1985) 463485. 

Proc.RSL A113 (1927) 621641 Pascual Jordan, "Ueber eine neue Begriindung der Quantenmechanik," ZP 40 (1927) 809838 Werner Heisenberg, "Ueber den anschaulichen Inhalt der quantentheoretischen Kinematik und Mechanik," ZP 43 (1927) 172198. Bohr first discussed the principle of complementarity at a conference in Como in 1927 see Niels Bohr, "The Quantum Postulate and the Recent Development of Atomic Theory," Nature 121 (1928) 580590. 

Having recognized the close quantitative connection between optical properties and collision dynamics at this stage of development of atomic theory is a remarkable example of Bohr s physical intuition. 

The discovery of the rare earth elements provide a long history of almost two hundred years of trial and error in the claims of element discovery starting before the time of Dalton s theory of the atom and determination of atomic weight values, Mendeleev s periodic table, the advent of optical spectroscopy, Bohr s theory of the electronic structure of atoms and Moseley s x-ray detection method for atomic number determination. The fact that the similarity in the chemical properties of the rare earth elements make them especially difficult to chemically isolate led to a situation where many mixtures of elements were being mistaken for elemental species. As a result, atomic weight values were not nearly as useful because the lack of separation meant that additional elements would still be present within an oxide and lead to inaccurate atomic weight values. Very pure rare earth samples did not become a reality until the mid twentieth century. 

Bohr next applied his theory to helium ions—helium atoms in which one of the two electrons is removed—and again the predictions of the theory exactly matched results obtained in experiments. The scientific world was convinced. For example, when Einstein heard of the results, he reversed himself and said, This is a tremendous achievement—Bohr s theory must be right. ... 

With reference to the minima of the radial distribution function D r), SCF analyses [61] using the near-Hartree-Fock wavefunctions of dementi [64] indicate that the numbers of electrons found in the inner shell extending up to the minimum of D r) amount to = 2.054 e (Be), 2.131 (C), 2.186 (O), 2.199 (F) and 2.205 electron (Ne). The results of Smith et al. [65] bearing on the boundaries in position space that enclose the exact number given by the Aufbau principle support the idea of physical shells compatible with that principle. The maxima of D r), on the other hand, also appear to be topological features indicative of shells, their positions correlate well with the shell radii from the Bohr-Schrodinger theory of an atom... 

THE DIELECTRIC CONSTANT OF ATOMIC HYDROGEN FROM THE POINT OF VIEW OF BOHRS QUANTUM THEORY... 

Spectroscopy was to prove indispensable in unlocking the structure of atoms, particulary their electronic stmcture— but those developments would depend on other, later researchers. Max Planck s analysis of blackbody radiation and Bohr s theory of the hydrogen spectrum are just two examples. 

In the modern model of the atom, based on wave mechanics, the conception of electronic orbits in the old model is replaced by the idea of the probability of the occurrence of an electron at a given point. The conclusions, however, which can be drawn from the older model remain the same in the newer conception, and it is important to remember that in this new model the essential points of Bohr s theory have not been discarded, but merely interpreted differently and very greatly refined. 

The periodic system developed from Bohr s atomic theory is of the greatest importance in chemical science because it demonstrates that the properties of the elements depend on their positions in the system. It is immediately apparent that chemical valency depends on the number of loosely-bound electrons in the atom. Thus, the alkali metals have one such electron while the divalent alkaline-earth metals have two, etc. Valency is therefore closely connected with electronic structure and provides the foundation for the modern theory of the chemical bond, the basis of which is to be found in the coupling or transfer of the valency electrons. 

Millikan s experiment did not prove, of course, that (he charge on the cathode ray. beta ray, photoelectric, or Zeeman particle was e. But if we call all such particles electrons, and assume that they have e/m = 1.76 x Hi" coulombs/kg. and e = 1.60 x 10" coulomb (and hence m =9.1 x 10 " kg), we find that they fit very well into Bohr s theory of the hydrogen atom and successive, more comprehensive atomic theories, into Richardson s equations for thermionic emission, into Fermi s theory of beta decay, and so on. In other words, a whole web of modem theory and experiment defines the electron. The best current value of e = (1.60206 0.00003) x 10 g coulomb. 

By 1903. llie wave theory of light based oil Maxwell s equations was well established, but certain phenomena would not fit in. It seemed that emission and absorption of hght occur discontinuously. This led Einstein to (lie view that the energy is concentrated in discrete particles. It was a revolutionary idea, very hard to understand, as the successes of the wave theory were undeniable. It seemed that light had to be understood sometimes as waves, sometimes as particles, and physicists had to get used to it, The idea was incorporated into Bohr s theory of the hydrogen atom and forms an essential part of it. 

N. Bohr. The Theory of Spectra and Atomic Calculation, Cambridge University Press, Cambridge, 1922. 

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. 

N. Bohr 3 discussed the fitness of configurations of the electrons in various atoms for the formation of ions. N. V. Sidgwiek has extended Bohr s theory to the electronic structure of atoms in co-ordination compounds. The subject was also discussed by J. D. M. Smith, and others at the Faraday Society s discussion on The Electronic Theory of Valency. A. Job discussed the catalyzed reaction NH3+HC1—NH4CI on the assumption that an unstable electronic system is formed as an intermediate product. 

Bohr s theory was extended in various ways, especially by Somerfeld, who showed how to deal with elliptical orbits. There was a certain amount of qualitative success in applying the theory to atoms with several electrons. These developments in what is now called the old quantum theory were important as they laid much of the groundwork necessary for a correct theory. Ultimately, they were unsuccessful. Bohr s theory did not really explain what is going on why should only some orbits be allowed Where does the quantization condition (eqn 4.12) come from Following the developments of... 

Students will demonstrate an understanding of the five basic atomic theories—the Dalton atom, the Thomson atom, the Rutherford atom, the Bohr atom, and the Schrodinger electron cloud model—and illustrate this understanding in a two-dimensional work of art. 

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