### Spectrum Lines and Ionization Energies

When visible light is passed through a prism, it splits into a partial electromagnetic spectrum: red light is on one side, with blue light on the other. The rest of the invisible spectrum extends on either side of the visible spectrum. However, when a heated or excited element (such as glowing neon gas) is passed through a prism, a set of short, bright lines appears. Each element has its own unique set of spectrum lines; this lets us analyze what a distant object, such as a star, is made of by examining its spectrum.

The hydrogen atom was one of the first elements to have its spectrum analyzed, because of its commonality in the universe. Several series of spectrum lines in hydrogen were observed in the late 19th century, including the Lymann series, the Balmer series, and the Paschen series, named after their discoverers. A scientist named Johannes Rydberg developed an equation that could precisely predict the frequencies of the hydrogen spectrum lines:

ν = 3.288 x 1015 * (1/n1 * 1/n2)

Where n1 and n2 are two whole numbers. However, no one could figure out why these whole numbers should predict the frequencies of the spectrum lines. Neils Bohr, inventor of the then-current model of the atom, finally came up with a solution.

Using the "solar-system" model of the atom, electrons orbit at certain distances, or energy levels, from the nucleus. Exactly why only certain distances are allowed is explained in the next page. When an atom recieves energy, such as from electromagnetic radiation, an electric current, heat energy, or other source, electrons "jump" out farther from the nucleus. Since the electron has to move against the attractive force of the nucleus, the energy is absorbed. When the electron jumps back to its normal energy level, energy of a specific wavelength and frequency is emitted.

Neils Bohr used his atomic theory to explain the spectrum lines: when electrons dropped from the 5th, 4th, 3rd, and 2nd energy levels to the 1st level, electromagnetic radiation of specific energy was released, producing the spectrum lines. When electrons dropped from the 5th, 4th, and 3rd levels to the 2nd, another series of lines resulted. This theory accurately and elegantly explained the mysterious phenomenon of spectrum lines. Other elements, with different electron configurations, had different spectrum patterns.

As we stated earlier, energy of various types (especially electromagnetic radiation) can elevate, or excite, electrons to a higher energy level. If enough energy is added, however, the electron can be removed from the atom altogether. The amount of energy that is required to remove an electron from a mole of atoms is called the ionization energy, and is usually measured in kilojoules per mole. Note that the energy of the EM radiation (using the equation E=hν) must be at least as great as the ionization energy, no matter how much radiation is emitted. If the energy of the EM radiation is higher than the substance's ionization energy, the electrons are directed with greater force (speed). The ionization of a substance when EM energy is directed onto it is called the photoelectric effect, and was studied extensively by Albert Einstein, work for which he received a Nobel Prize.

Once spectrum lines had been accounted for, scientists raised new questions about the structure of electron energy levels: why did electrons only orbit at certain distances from the nucleus? Why is energy of a specific frequency released when electrons jump down to a lower energy level? And why don't electrons eventually lose energy and crash into the nucleus?

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