Nonetheless, his guess that his formula could be extended to predict other lines of hydrogen proved to be correct, leading to the prediction of lines that had not yet been observed.Īnother regularity in atomic spectra was found by Swedish physicist and mathematician Johannes Rydberg. Balmer, however, was unable to provide a reason why his formula worked so successfully. A portion of the hydrogen emission spectrum.Ī Swiss schoolteacher, Johann Jakob Balmer, first expressed the wavelengths of these spectral lines in a single mathematical formula in 1885. The spacing between successive lines becomes smaller and smaller from the first red line to the last ultraviolet one, until the lines become so close that they seem to merge.įigure 4. The most orderly spectrum was that of hydrogen ( Figure 4). Physicists tried to find order in the confusing arrays of spectral lines. When this light was viewed through a spectroscope, a pattern of spectral lines emerged. In the nineteenth century, chemists used optical spectroscopes for chemical analysis. Atomic spectra: Clues to atomic structure Bohr was able to advance to the next step and determine features of individual atoms. Planck had related the frequency of radiated light to energy change in matter without a model of the atom. The frequency of the emitted photon, its color, depends on the magnitude of the jump. So the atom emits a photon whose energy is equal to the difference in energy between the two energy levels, E = hf. This is energy conservation at the atomic level. The amount of energy that boosts an electron to a higher orbit is the same amount of energy carried away when the electron de-excites back to its lower energy state as illustrated in Figure 2. The emitted frequency of radiation is determined by the energy differences in the atom. So an orbiting electron should radiate energy continuously, causing the electron to spiral into the nucleus ( Figure 3).īohr broke with classical physics by stating that the electron doesn’t radiate light while it accelerates around the nucleus radiation of light occurs only when the electron makes a transition from a higher energy level to a lower energy level.

An electron that orbits a nucleus is constantly accelerating. Accelerated electrons, according to James Clerk Maxwell’s theory, radiate energy in the form of electromagnetic waves. The classical view of the atom had a glaring difficulty. Classically, an electron continuously emitting energy should spiral into the nucleus. Bohr took this a step further, hypothesizing that the energy E is the difference in the atom’s energy when an electron moves from one orbit to another.įigure 3. However, he did recognize that the frequency of emitted radiation is determined by E = hf (actually, f = E/h). Interestingly, Bohr didn’t believe in photons when he developed the quantum view of the atom. Five years later Einstein proposed not only that material energy is quantized, but that light itself exists as quantum lumps, or “corpuscles,” later named photons. With the constant of proportionality h (Planck’s constant) we have the familiar E = hf. In 1900 Max Planck postulated that the energy of a radiated quantum of energy is proportional to the frequency of radiation: E ~ f.

In the bohr model of the atom