Niels Bohr: Biography & Atomic Theory
Aug 04, · What Were Niels Bohr's Experiments? Niels Bohr investigated the structure of the atom and won the Nobel Prize in Physics in for discovering the Bohr model. He developed the liquid drop model of the atomic nucleus, proposed the idea of complementarity and identified the U isotope essential for nuclear fission. Bohr's Experiment and Reasoning In Bohr analysed the spectral lines of hydrogen. This meant that he observed the coloured lines (wavelengths given by a spectroscope). Since the lines were not continuous like rainbow, the electrons were not giving of energy continuously. Bohr concluded that the laws of basic physics did not abide to objects on an atomic scale.
Niels Bohr investigated the structure of the atom and won the Nobel Prize in Physics in for discovering the Bohr model. He developed the liquid drop model of the atomic nucleus, proposed ddid idea of complementarity and identified the U isotope essential for nuclear fission.
After fleeing the German invasion of Diid, he worked on the development of the atomic bomb at Los Alamos in the United States. The Bohr model proposes electrons orbit the nucleus because of electromagnetic forces rather than gravity. Because of this difference, electrons do not collapse into the nucleus but travel from one orbital shell to another as energy levels change.
As a Danish citizen, Bohr resisted involvement with German nuclear power research and aided many Jewish scientists in their escape what did martin luther king Nazi power. He fled Denmark in and eventually made his way to London, where he was involved with experiments for a viable nuclear fission bomb. Bohr was adamant that atomic research should be shared, causing Winston Churchill to view him as a security risk.
He organized the Atoms for Peace Conference in Geneva in Bohr's son Aage Niels Bohr also contributed to the world of physics. Aage won the Nobel Prize in Physics in What Were Niels Bohr's Experiments? More From Reference. What Is Aristocracy? What Do Stars Symbolize?
The discovery of the electron and radioactivity in the late 19th century led to different models being proposed for the atom’s structure. In , Niels Bohr proposed a theory for the hydrogen atom, based on quantum theory that some physical quantities only take discrete values. Electrons move around a nucleus, but only in prescribed orbits, and If electrons jump to a lower-energy orbit, the difference is . When Bohr was awarded the Nobel Prize for his work in , the Hungarian physical chemist Georg Hevesy, together with the physicist Dirk Coster from Holland, were working at Bohr’s institute to establish experimentally that the as-yet-undiscovered atomic element 72 would behave as predicted by Bohr’s theory. They succeeded in , thus proving both the strength of Bohr’s theory and the truth in . Bohr's experiment was simple. He heated hydrogen gas confined in a tube which caused the gas inside it to excite, releasing energy. When viewed through a diffraction grating, he saw that the single ray of light was actually formed by different colours. Using this information, Bohr came up with his atomic theory.
In Bohr proposed his quantized shell model of the atom see Bohr atomic model to explain how electrons can have stable orbits around the nucleus. The motion of the electrons in the Rutherford model was unstable because, according to classical mechanics and electromagnetic theory, any charged particle moving on a curved path emits electromagnetic radiation; thus, the electrons would lose energy and spiral into the nucleus.
To remedy the stability problem, Bohr modified the Rutherford model by requiring that the electrons move in orbits of fixed size and energy. The energy of an electron depends on the size of the orbit and is lower for smaller orbits.
Radiation can occur only when the electron jumps from one orbit to another. The atom will be completely stable in the state with the smallest orbit, since there is no orbit of lower energy into which the electron can jump. A stable atom has a certain size so that any equation describing it must contain some fundamental constant or combination of constants with a dimension of length. The classical fundamental constants—namely, the charges and the masses of the electron and the nucleus—cannot be combined to make a length.
Bohr noticed, however, that the quantum constant formulated by German physicist Max Planck has dimensions which, when combined with the mass and charge of the electron, produce a measure of length.
Numerically, the measure is close to the known size of atoms. Planck had introduced his constant in in a formula explaining the light radiation emitted from heated bodies. According to classical theory, comparable amounts of light energy should be produced at all frequencies.
This is not only contrary to observation but also implies the absurd result that the total energy radiated by a heated body should be infinite. The energy quantum is related to the frequency of the light by a new fundamental constant, h.
When a body is heated, its radiant energy in a particular frequency range is, according to classical theory, proportional to the temperature of the body. If the radiant energy is less than the quantum of energy, the amount of light in that frequency range will be reduced.
He postulated that the angular momentum of the electron is quantized—i. He assumed that otherwise electrons obey the laws of classical mechanics by traveling around the nucleus in circular orbits.
Because of the quantization, the electron orbits have fixed sizes and energies. The orbits are labeled by an integer, the quantum number n.
With his model, Bohr explained how electrons could jump from one orbit to another only by emitting or absorbing energy in fixed quanta. For example, if an electron jumps one orbit closer to the nucleus, it must emit energy equal to the difference of the energies of the two orbits. Conversely, when the electron jumps to a larger orbit, it must absorb a quantum of light equal in energy to the difference in orbits. The model also explains the Balmer formula for the spectral lines of hydrogen.
The light energy is the difference in energies between the two orbits in the Bohr formula. Bohr himself noted that the formula also applies to the singly ionized helium atom, which, like hydrogen, has a single electron. The nucleus of the helium atom has twice the charge of the hydrogen nucleus, however.
Two of those powers stem from the charge on the nucleus; the other two come from the charge on the electron itself. Bohr modified his formula for the hydrogen atom to fit the helium atom by doubling the charge on the nucleus.
German physicists James Franck and Gustav Hertz confirmed the existence of quantum states in atoms in experiments reported in They made atoms absorb energy by bombarding them with electrons. The atoms would only absorb discrete amounts of energy from the electron beam. When the energy of an electron was below the threshold for producing an excited state, the atom would not absorb any energy. Except for the spectra of X-rays in the K and L series, it could not explain properties of atoms having more than one electron.
The binding energy of the helium atom, which has two electrons, was not understood until the development of quantum mechanics. Several features of the spectrum were inexplicable even in the hydrogen atom.
High-resolution spectroscopy shows that the individual spectral lines of hydrogen are divided into several closely spaced fine lines. In a magnetic field the lines split even farther apart. The quantization of the orientation of the angular momentum vector was confirmed in an experiment in by other German physicists, Otto Stern and Walther Gerlach.
Their experiment took advantage of the magnetism associated with angular momentum; an atom with angular momentum has a magnetic moment like a compass needle that is aligned along the same axis.
The researchers passed a beam of silver atoms through a magnetic field, one that would deflect the atoms to one side or another according to the orientation of their magnetic moments.
In their experiment Stern and Gerlach found only two deflections, not the continuous distribution of deflections that would have been seen if the magnetic moment had been oriented in any direction. Thus, it was determined that the magnetic moment and the angular momentum of an atom can have only two orientations.
The discrete orientations of the orbits explain some of the magnetic field effects—namely, the so-called normal Zeeman effect , which is the splitting of a spectral line into three separate subsidiary lines.
These lines correspond to quantum jumps in which the angular momentum along the magnetic field is increased by one unit, decreased by one unit, or left unchanged.
Spectra in magnetic fields displayed additional splittings that showed that the description of the electrons in atoms was still incomplete. In Samuel Abraham Goudsmit and George Eugene Uhlenbeck , two graduate students in physics at the University of Leiden in the Netherlands, added a quantum number to account for the division of some spectral lines into more subsidiary lines than can be explained with the original quantum numbers. Goudsmit and Uhlenbeck postulated that an electron has an internal spinning motion and that the corresponding angular momentum is one-half of the orbital angular momentum quantum.
Independently, Austrian-born physicist Wolfgang Pauli also suggested adding a two-valued quantum number for electrons, but for different reasons. He needed this additional quantum number to formulate his exclusion principle , which serves as the atomic basis of the periodic table and the chemical behaviour of the elements. According to the Pauli exclusion principle, one electron at most can occupy an orbit, taking into account all the quantum numbers. Pauli was led to this principle by the observation that an alkali metal atom in a magnetic field has a number of orbits in the shell equal to the number of electrons that must be added to make the next noble gas.
These numbers are twice the number of orbits available if the angular momentum and its orientation are considered alone. It could not explain the number of fine spectral lines and many of the frequency shifts associated with the Zeeman effect. Most daunting , however, was its inability to explain the rich spectra of multielectron atoms. In fact, efforts to generalize the model to multielectron atoms had proved futile , and physicists despaired of ever understanding them.
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