In 1913, Danish physicist Niels Bohr developed the first model of the atom that used the concepts of quantum theory to explain atomic spectra. The Bohr atom depicted electrons "in orbit" around a heavy nucleus, similar to the motion of planets around the sun. The model was based on Ernest Rutherford's 1911 demonstration that the atom consists of a central nucleus surrounded by a cloud of electrons. The electrons, Rutherford postulated, would have to be in constant motion, since the attraction of opposite charges would otherwise pull them into the nucleus. Because electrons are charged particles, they should radiate electromagnetic energy during orbits, eventually losing momentum and spiraling into the nucleus. The radiated energy would be expected to change continuously across the electromagnetic spectrum as the momentum of the moving particle changed.
This was not the observed behavior of atoms. Under normal conditions, they do not radiate energy at all. When an atom is exposed to an environment in which energy can be transferred to it (a very strong electric field or very high temperature), it does emit light. This light is observed only at particular wavelengths instead of a continuous spectrum. These wavelengths are referred to as the spectral lines of the substance. Each element, when it is excited, gives off certain colors or wavelengths, a color "fingerprint" that identifies the particular atoms emitting the light. These fingerprints are sometimes referred to as spectral lines. The light coming from these atoms does not take the shape of lines but, when separated using an array of lenses and slits, the light produces an image in the shape of a series of lines, each corresponding to a particular wavelength. Every type of atom has a characteristic set of spectral lines at discrete wavelengths.
To explain this phenomenon, Bohr applied the idea of quantized energy previously developed by Planck. His first assumption was that, in contrast to classical mechanics, where an infinite number of orbits are possible, an electron could be in only one of a discrete set of orbits, which he termed stationary states. His proposed atom restricted electrons to orbits of certain "allowed" radii. An electron orbiting in one of these "allowed" orbits has a defined energy state, does not radiate energy, and does not spiral into the nucleus. Even though the electron is in constant motion, it does not emit electromagnetic radiation from a stationary state. Bohr concluded that energy changes in an atom occur only when electrons move from one allowed orbit to another. Energy must be absorbed for an electron to move to a higher state (further from the nucleus) and emitted when the electron moves to an orbit of lower energy (closer to the nucleus).
The overall change in energy associated with this "orbit jumping" is the difference in energy levels between the initial and final orbits. The difference in energy levels corresponds to a particular wavelength of light multiplied by Planck's constant. The spectrum for each atom is defined by these wavelengths. The line spectrum of an atom--even a simple hydrogen atom with only one electron--can be very complex. Although the electron can only move from one energy level to another with no stopping in between, the energy differences between levels are not all the same. This means that a move from the third level to the second will emit a quantum of light energy (now known as a photon) with a different wavelength than that of a move between the second level and the first. To further complicate the spectrum, the electron can move several levels at once, emitting yet another wavelength. Each stable orbit has room for a limited number of electrons. Once it is full, no more electrons can move to that particular energy. A final rule of electron movement states that the electron cannot move beyond the highest energy level into the nucleus.
The Bohr atom model successfully accounted for most known spectral lines for hydrogen and predicted new spectral wavelengths that were later confirmed. The theory also explained the observation that more spectral lines were observed for hydrogen in interstellar space. These lines are the result of transitions from very high energy levels. Electrons can only remain at these levels in the high vacuum of interstellar space, where collisions between atoms are much less frequent. Although the theory provided a good explanation for the spectrum of light from an atom, it did have some major deficiencies: it did not account for intensities of the various lines in the spectrum; it did not explain chemical bonding; and it did not work for atoms with more than one electron. It was only useful for predicting the behavior of atoms with a single electron (H, He+ , and Li2+ions).
Even so, Bohr's theory was a significant advance in understanding atomic structure. It provided several very important features that have survived in the current quantum mechanics models: notably, the existence of stationary, nonradiating states and the relationship of radiation frequency to the energy difference between the initial and final states in a quantum transition. After its introduction by Bohr in 1913, the model was expanded and modified by Bohr and other scientists, most notably Arnold Sommerfeld. In 1922, Niels Bohr was awarded the Nobel Prize in physics. Bohr's model of the atom was important because it introduced quantized energy states for the electrons. Although the model itself was eventually replaced, the concept of quantized energy states remained as a key element of all later models.
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