Experimentally, the spectra were well established, an equation was found to fit the experimental data, but the theoretical foundation was missing. Again, we see the interplay between experiment and theory in physics. Bohr was the first to comprehend the deeper meaning. Balmer first devised the formula for his series alone, and it was later found to describe all the other series by using different values of n f n f. While the formula in the wavelengths equation was just a recipe designed to fit data and was not based on physical principles, it did imply a deeper meaning. Thus, for the Balmer series, n f = 2 n f = 2 and The constant n i n i is a positive integer, but it must be greater than n f n f. There are apparently an unlimited number of series, although they lie progressively farther into the infrared and become difficult to observe as n f n f increases. The Paschen series and all the rest are entirely IR. The Lyman series is entirely in the UV, while part of the Balmer series is visible with the remainder UV. For the Lyman series, n f = 1 n f = 1 for the Balmer series, n f = 2 n f = 2 for the Paschen series, n f = 3 n f = 3 and so on. The constant n f n f is a positive integer associated with a specific series. The observed hydrogen-spectrum wavelengths can be calculated using the following formula: (See Figure 30.15.) These series are named after early researchers who studied them in particular depth. The hydrogen spectrum had been observed in the infrared (IR), visible, and ultraviolet (UV), and several series of spectral lines had been observed. As you might expect, the simplest atom-hydrogen, with its single electron-has a relatively simple spectrum. In some cases, it had been possible to devise formulas that described the emission spectra. (credit for (b): Yttrium91, Wikimedia Commons) The line spectrum for each element is unique, providing a powerful and much used analytical tool, and many line spectra were well known for many years before they could be explained with physics. The discrete lines imply quantized energy states for the atoms that produce them. Part (b) shows the emission line spectrum for iron. (It was a running joke that any theory of atomic and molecular spectra could be destroyed by throwing a book of data at it, so complex were the spectra.) Following Einstein’s proposal of photons with quantized energies directly proportional to their wavelengths, it became even more evident that electrons in atoms can exist only in discrete orbits.įigure 30.14 Part (a) shows, from left to right, a discharge tube, slit, and diffraction grating producing a line spectrum. But, in spite of years of efforts by many great minds, no one had a workable theory. (See Figure 30.14.) Maxwell and others had realized that there must be a connection between the spectrum of an atom and its structure, something like the resonant frequencies of musical instruments. Atomic and molecular emission and absorption spectra have been known for over a century to be discrete (or quantized). (credit: Unknown Author, via Wikimedia Commons) Mysteries of Atomic SpectraĪs noted in Quantization of Energy, the energies of some small systems are quantized. His many contributions to the development of atomic physics and quantum mechanics, his personal influence on many students and colleagues, and his personal integrity, especially in the face of Nazi oppression, earned him a prominent place in history. Bohr’s theory explained the atomic spectrum of hydrogen and established new and broadly applicable principles in quantum mechanics.įigure 30.13 Niels Bohr, Danish physicist, used the planetary model of the atom to explain the atomic spectrum and size of the hydrogen atom. From their sizes to their spectra, much was known about atoms, but little had been explained in terms of the laws of physics. For decades, many questions had been asked about atomic characteristics. In 1913, after returning to Copenhagen, he began publishing his theory of the simplest atom, hydrogen, based on the planetary model of the atom. Bohr became convinced of its validity and spent part of 1912 at Rutherford’s laboratory. The great Danish physicist Niels Bohr (1885–1962) made immediate use of Rutherford’s planetary model of the atom. Describe the triumphs and limits of Bohr’s theory.Illustrate energy state using the energy-level diagram.Explain Bohr’s planetary model of the atom.Explain Bohr’s theory of the hydrogen atom.Describe the mysteries of atomic spectra.By the end of this section, you will be able to:
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