Chemistry

In 1913, Niels Bohr introduced the Bohrfs model to account for the experimental spectrum of hydrogen. Bohr postulated that the orbital momentum of an electron around the nucleus is quantized, allowing only ejumpf from one allowed orbit to another, and not spiral into the nucleus.

Bohr also added that each stable orbit has limited occupancy for electrons. The energy associated with these quantum jumps explained both the appearance of bright lines and dark lines at the same wavelengths in the emission-absorption spectrum, seen in different elements, notably the Balmer series of hydrogen. Since the spacing between the lines (energy levels) depends on the size of the Planckfs constant. Bohr rewrote the Balmerfs formula to include the Planckfs constant, h. Because circular orbits are involved, the basic quanta of energy are measured in terms of h/2.

Bohrfs model gave hydrogen atom energy levels:
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RH, Rydberg constant: 2.18 x 10-18 J
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