## Origin

In the early 20th century, experiments by Ernest Rutherford established that atoms consisted of a diffuse cloud of negatively charged electrons surrounding a small, dense, positively charged nucleus. Given this experimental data, Rutherford naturally considered a planetary-model atom, the Rutherford model of 1911 – electrons orbiting a solar nucleus – however, said planetary-model atom has a technical difficulty. The laws of classical mechanics (i.e. the Larmor formula), predict that the electron will release electromagnetic radiation while orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus. This atom model is disastrous, because it predicts that all atoms are unstable.Also, as the electron spirals inward, the emission would gradually increase in frequency as the orbit got smaller and faster. This would produce a continuous smear, in frequency, of electromagnetic radiation. However, late 19th century experiments with electric discharges through various low-pressure gases in evacuated glass tubes had shown that atoms will only emit light (that is, electromagnetic radiation) at certain discrete frequencies.

To overcome this difficulty, Niels Bohr proposed, in 1913, what is now called the

*Bohr model of the atom*. He suggested that electrons could only have certain

*classical*motions:

- The electrons can only travel in certain orbits: at a certain discrete set of distances from the nucleus with specific energies.
- The electrons of an atom revolve around the nucleus in orbits. These orbits are associated with definite energies and are also called energy shells or energy levels. Thus, the electrons do not continuously lose energy as they travel in a particular orbit. They can only gain and lose energy by jumping from one allowed orbit to another, absorbing or emitting electromagnetic radiation with a frequency
*ν*determined by the energy difference of the levels according to the*Planck relation*:

*h*is Planck's constant. - The frequency of the radiation emitted at an orbit of period
*T*is as it would be in classical mechanics; it is the reciprocal of the classical orbit period:

*only when restricted by a quantum rule*. Although rule 3 is not completely well defined for small orbits, because the emission process involves two orbits with two different periods, Bohr could determine the energy spacing between levels using rule 3 and come to an exactly correct quantum rule: the angular momentum

*L*is restricted to be an integer multiple of a fixed unit:

*n*= 1, 2, 3, ... is called the principal quantum number, and

*ħ*=

*h*/2π. The lowest value of

*n*is 1; this gives a smallest possible orbital radius of 0.0529 nm known as the Bohr radius. Once an electron is in this lowest orbit, it can get no closer to the proton. Starting from the angular momentum quantum rule Bohr

^{[2]}was able to calculate the energies of the allowed orbits of the hydrogen atom and other hydrogen-like atoms and ions.

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