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Astronomers and physicists have worked hard to learn the lines that go with each element by studying the way atoms absorb and emit light in laboratories here on Earth. Then they can use this knowledge to identify the elements in celestial bodies. In this way, we now know the chemical makeup of not just any star, but even galaxies of stars so distant that their light started on its way to us long before Earth had even formed.

Energy levels and excitation

Bohr’s model of the hydrogen atom was a great step forward in our understanding of the atom. However, we know today that atoms cannot be represented by quite so simple a picture. For example, the concept of sharply defined electron orbits is not really correct; however, at the level of this introductory course, the notion that only certain discrete energies are allowable for an atom is very useful. The energy levels we have been discussing can be thought of as representing certain average distances of the electron’s possible orbits from the atomic nucleus.

Ordinarily, an atom is in the state of lowest possible energy, its ground state    . In the Bohr model of the hydrogen atom, the ground state corresponds to the electron being in the innermost orbit. An atom can absorb energy, which raises it to a higher energy level (corresponding, in the simple Bohr picture, to an electron’s movement to a larger orbit)—this is referred to as excitation    . The atom is then said to be in an excited state . Generally, an atom remains excited for only a very brief time. After a short interval, typically a hundred-millionth of a second or so, it drops back spontaneously to its ground state, with the simultaneous emission of light. The atom may return to its lowest state in one jump, or it may make the transition in steps of two or more jumps, stopping at intermediate levels on the way down. With each jump, it emits a photon of the wavelength that corresponds to the energy difference between the levels at the beginning and end of that jump.

An energy-level diagram for a hydrogen atom and several possible atomic transitions are shown in [link] . When we measure the energies involved as the atom jumps between levels, we find that the transitions to or from the ground state, called the Lyman series of lines, result in the emission or absorption of ultraviolet photons. But the transitions to or from the first excited state (labeled n = 2 in part (a) of [link] ), called the Balmer series, produce emission or absorption in visible light. In fact, it was to explain this Balmer series that Bohr first suggested his model of the atom.

Energy-level diagrams for hydrogen.

Energy-Level Diagram for Hydrogen and the Bohr Model for Hydrogen. The right hand side (a) of the figure shows the Bohr model with the Lyman, Balmer, and Paschen series illustrated. A small circle representing the nucleus is enclosed by a larger circle for orbit n=1, then another larger circle for n=2 and so on up to n=5. At the top of this diagram are 4 arrows starting at n=2, with one arrow going up to n=3, one to n=4 and one to n=5. As these arrows are moving away from the nucleus, they represent absorption of energy by the atom to move an electron up to each level. Next is the Lyman series, with arrows from each upper orbital pointing down to n=1. As these arrows are pointing toward the nucleus, energy is released from the atom as electrons “fall” from upper levels down to n=1. Finally, there is the Paschen series, with arrows from the upper levels all pointing down to n=3. Again, as these arrows point toward the nucleus, light is emitted as the electron moves closer to the nucleus. The left hand side (b) of the figure shows the movement of electrons from higher to lower energy levels, represented with arrows pointing downward. From left to right, Lyman series has the longest arrows, then Balmer series with arrows about half as long, then Paschen series with arrows about a fourth as long, then Bracket series with arrows about an eighth as long.
(a) Here we follow the emission or absorption of photons by a hydrogen atom according to the Bohr model. Several different series of spectral lines are shown, corresponding to transitions of electrons from or to certain allowed orbits. Each series of lines that terminates on a specific inner orbit is named for the physicist who studied it. At the top, for example, you see the Balmer series, and arrows show electrons jumping from the second orbit ( n = 2) to the third, fourth, fifth, and sixth orbits. Each time a “poor” electron from a lower level wants to rise to a higher position in life, it must absorb energy to do so. It can absorb the energy it needs from passing waves (or photons) of light. The next set of arrows (Lyman series) show electrons falling down to the first orbit from different (higher) levels. Each time a “rich” electron goes downward toward the nucleus, it can afford to give off (emit) some energy it no longer needs. (b) At higher and higher energy levels, the levels become more and more crowded together, approaching a limit. The region above the top line represents energies at which the atom is ionized (the electron is no longer attached to the atom). Each series of arrows represents electrons falling from higher levels to lower ones, releasing photons or waves of energy in the process.
Practice Key Terms 4

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Source:  OpenStax, Astronomy. OpenStax CNX. Apr 12, 2017 Download for free at http://cnx.org/content/col11992/1.13
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