Bohr theory of the atom

The Bohr theory of the atom was further developed with great ingenuity to explain the complexities of atomic line spectra, but the significant advance came with the formulation of wave mechanics. 

In the classical theory of electrodynamics, electromagnetic radiation is emitted when an electron moves in its orbit but, ac cording to the Bohr theory of the atom,... 

Although valency strokes have been customary in chemical formulae for a century, one could not until recently attach to them any real notion about their true nature. On the patient paper one operated with them as with hooks which were undone, rotated etc. at will. Even the Rutherford-Bohr theory of the atom furnished no explanation, not even for the bonding of two hydrogen atoms to form a hydrogen molecule. The successful octet theory and the Lewis and Langmuir theory of the electron-pair bond associated with it was also still purely formal, but later was seen to be essentially correct. 

The Bohr theory of the atom destroyed the last pockets of resistance to the quantum concept. Yet the wave attributes of light were there too. The nature of light took on a dual aspect. This duality in the nature of light is accepted now, though to some in the beginning it was a bitter pill. 

Hg. 3.12 Tlie orbital for a Is electron in the hydrogen atom. The radius of the sphere in which it is 90% likely that the electron will be found, is about 100 pm. The single radius at which the Is electron is most likely to be found is a distance 52.9 pm from the nucleus. This may be compared with the Bohr theory of the atom, where it was assumed that the electron was certain to be found at a radius of 52.9 pm. 

Stark effect The splitting of lines in the spectra of atoms due to the presence of a strong electric field. It is named after the German physicist Johannes Stark (1874-1957), who discovered it in 1913. Like the normal Zeeman effect, the Stark effect can be understood in terms of the classical electron theory of Lorentz. The Stark effect for hydrogen atoms was also described by the Bohr theory of the atom. In terms of quantum mechanics, the Stark effect is described by regarding the electric field as a perturbation on the quantum states and energy levels of an atom in the absence of an electric field. This application of perturbation theory was its first use in quantum mechanics. 

Stark effect The splitting of atomic spectral lines due to the presence of an external electric field. This phenomenon was discovered by the German physicist Johannes Stark (1874-1957) in 1913 and can be explained using classical electron theory and the BOHR THEORY of the atom. 

The quantity, a . is referred to as the Bohr radius (52.917726 pm) because it is identical to the radius of the orbit of the 1 s electron in the Bohr atom". The eariy Bohr theory of the atom invoked eiectron orbits of definite radii, but these are invaiid since they vioiate the Heisenberg uncertainty principle. The Bohr radius is used as the atomic unit of iength. 

Q8 Which one of the following atoms has the greatest number of unpaired electrons in the ground state A Mn B Fe C Ni D Sc Q9 Why was the Bohr theory of the atom developed A To account for changes in gas volumes with temperature,... 

G.N. Lewis was North America s leading physical chemist when he warned physicists that their most recent theoretical proposals — particularly the Bohr theory of the atom — would not suffice as a conceptual or causal foundation for chemistry. In so far as the philosophy of chemistry is born out of that troubling dependence upon physics which is never quite enough to confirm chemistry s reduction to physics,... 

To make an informed guess for your first value of ot, you may wish to reread the section on the Bohr theory of the hydrogen atom and the Schroedinger wave functions for the hydrogen atom in a good physical or general chemistry book (see Bibliography). 

This section started with the discovery of Soddy and Fajans on radioactive decay around 1910 and the relationship of radioactive decay to the periodic table. At this point in the history, we understand the periodic table and we understand the role of isotopes in the periodic table. We have not yet understood the structure of the modern Table, i.e. first row two elements, second row eight elements, etc. That understanding can be based on Bohr theory of the hydrogen atom originally developed in 1911 and is summarized in Bohr s famous article in Zeitschrift fur Physik (Bohr 1922). 

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. 

The Bohr model of the atom took shape in 1913. Niels Bohr (1885-1962), a Danish physicist, started with the classic Rutherford model and applied a new theory of quantum mechanics to develop a new model that is still in use, but with many enhancements. His assumptions are based on several aspects of quantum theory. One assumption is that light is emitted in tiny bunches (packets) of energy call photons (quanta of light energy). 

It was the apparent violation of this requirement that led to the postulate of quantization in the Bohr theory of the hydrogen atom. However, we are concerned here with the classical result in which the charge does radiate. Our objective is to describe the emitted radiation some distance r from the emitter. 

A more detailed account of the Bohr theory of the hydrogen atom is given in Chap. 2 and Apps. II and III. 

Quantum Number (Principal). A quantum number that, in the old Bohr model of the atom, determined the energy of an electron in one of the allowed orbits around the nucleus, In the theory of quantum mechanics, the principal quantum number is used most commonly to describe the atomic shell in which tlie elections are located, In a somewhat general way, it is related to the energy of the electronic states of an atom, The symbol for the principal quantum number is n. In x-ray spectral terminology, a -shell is identical to an n = 1 shell, and an L-shell to an n = 2 shell, etc. 

While h is quite small in the macroscopic world, it is not at all insignificant when the particle under consideration is of subatomic scale. Let us use an actual example to illustrate this point. Suppose the Ax of an electron is 10-14 m, or 0.01 pm. Then, with eq. (1.2.1), we get Apx = 5.27 x 10-21 kg m s-1. This uncertainty in momentum would be quite small in the macroscopic world. However, for subatomic particles such as an electron, with mass of 9.11 x 10-31 kg, such an uncertainty would not be negligible at all. Hence, on the basis of the Uncertainty Principle, we can no longer say that an electron is precisely located at this point with an exactly known velocity. It should be stressed that the uncertainties we are discussing here have nothing to do with the imperfection of the measuring instruments. Rather, they are inherent indeterminacies. If we recall the Bohr theory of the hydrogen atom, we find that both the radius of the orbit and the velocity of the electron can be precisely calculated. Hence the Bohr results violate the Uncertainty Principle. 

A theory of these wave numbers was worked out by Bohr before the discovery of wave mechanics or the wave particle theory. Bohr s theory was based on Einstein s idea that light consists of photons and on Rutherford s nucleus theory of the atom. 

E In the Bohr model of the atom an electron may orbit the atom s nucleus only at certain radii corresponding to certain energies. He called this the quantum theory of the atom. 

In a note presented to the National Academy of Sciences11 have given some computations of the K critical absorption frequencies of the chemical elements based on the Rutherford-Bohr theory of the structure of atoms and the mechanism of radiation. In these computations I have assumed that the electrons were distributed in circular orbits, which did not lie in planes passing through the nucleus of the atom. 

Bohr theory (4.2) The first theory of the atom to propose that electrons in atoms were in definite energy levels, boiling point (14.2) The temperature at which a liquid changes to a gas at the prevailing pressure, boiling-point elevation (15.6) An increase in the boiling point of a solvent due to the presence of a solute, bond See covalent bond and ionic bond. bond order See total bond order. 

---