28.2 Atomic Spectra 905l(nm) 400 500 600 700(a)HHgNeHl(nm) 400 500 600 700(b)Figure 28.3 Visible spectra. (a) Line spectra produced by emission in the visible range for the elementshydrogen, mercury, and neon. (b) The absorption spectrum for hydrogen. The dark absorptionlines occur at the same wavelengths as the emission lines for hydrogen shown in (a).K. W. Whitten, R. E. Davis, M. L. Peck, and G. G. Stanley, General Chemistry,7th ed., Belmont, CA, Brooks/Cole, 2004.ions such as He and Li 2 . Further, a thorough understanding of the physics underlyingthe hydrogen atom can then be used to describe more complex atomsand the periodic table of the elements.Suppose an evacuated glass tube is filled with hydrogen (or some other gas) atlow pressure. If a voltage applied between metal electrodes in the tube is greatenough to produce an electric current in the gas, the tube emits light having acolor that depends on the gas inside. (This is how a neon sign works.) When theemitted light is analyzed with a spectrometer, discrete bright lines are observed,each having a different wavelength, or color. Such a series of spectral lines is commonlycalled an emission spectrum. The wavelengths contained in such a spectrumare characteristic of the element emitting the light (Fig. 28.3). Because notwo elements emit the same line spectrum, this phenomenon represents a marvelousand reliable technique for identifying elements in a gaseous substance.The emission spectrum of hydrogen shown in Figure 28.4 includes four prominentlines that occur at wavelengths of 656.3 nm, 486.1 nm, 434.1 nm, and410.2 nm, respectively. In 1885 Johann Balmer (1825–1898) found that the wavelengthsof these and less prominent lines can be described by the simple empiricalequation1 R H 1 [28.1]2 2 1 2 nwhere n may have integral values of 3, 4, 5, . . . , and R H is a constant, called theRydberg constant. If the wavelength is in meters, R H has the valueR H 1.097 373 2 10 7 m 1 [28.2]The first line in the Balmer series, at 656.3 nm, corresponds to n 3 in Equation28.1, the line at 486.1 nm corresponds to n 4, and so on. In addition to theBalmer series of spectral lines, a Lyman series was subsequently discovered in thefar ultraviolet, with the radiated wavelengths described by a similar equation.In addition to emitting light at specific wavelengths, an element can absorblight at specific wavelengths. The spectral lines corresponding to this process formwhat is known as an absorption spectrum. An absorption spectrum can be obtainedby passing a continuous radiation spectrum (one containing all wavelengths)through a vapor of the element being analyzed. The absorption spectrumconsists of a series of dark lines superimposed on the otherwise bright continuousspectrum. Each line in the absorption spectrum of a given element coincides witha line in the emission spectrum of the element. This means that if hydrogen is theλ(nm)364.6410.2 434.1486.1 656.3Figure 28.4 The Balmer series ofspectral lines for atomic hydrogen,with several lines marked with thewavelength in nanometers. The linelabeled 346.6 is the shortest-wavelengthline and is in the ultravioletregion of the electromagneticspectrum. The other labeled lines arein the visible region. Balmer series Rydberg constant
906 Chapter 28 Atomic <strong>Physics</strong>APPLICATIONDiscovery of Heliumabsorbing vapor, dark lines will appear at the visible wavelengths 656.3 nm,486.1 nm, 434.1 nm, and 410.2 nm, as shown in Figures 28.3b and 28.4.The absorption spectrum of an element has many practical applications. For example,the continuous spectrum of radiation emitted by the Sun must passthrough the cooler gases of the solar atmosphere before reaching the Earth. Thevarious absorption lines observed in the solar spectrum have been used to identifyelements in the solar atmosphere, including one that was previously unknown.When the solar spectrum was first being studied, some lines were found that didn’tcorrespond to any known element. A new element had been discovered! Becausethe Greek word for Sun is helios, the new element was named helium. It was lateridentified in underground gases on Earth. Scientists are able to examine the lightfrom stars other than our Sun in this way, but elements other than those presenton Earth have never been detected.Applying <strong>Physics</strong> 28.1Thermal or Spectral?On observing a yellow candle flame, your laboratorypartner claims that the light from the flame originatesfrom excited sodium atoms in the flame. You disagree,stating that because the candle flame is hot, the radiationmust be thermal in origin. Before the disagreementleads to fisticuffs, how could you determine whois correct?Explanation A simple determination could be madeby observing the light from the candle flame througha spectrometer, which is a slit and diffraction gratingcombination discussed in Chapter 25. If the spectrumof the light is continuous, then it’s probably thermalin origin. If the spectrum shows discrete lines, it’satomic in origin. The results of the experiment showthat the light is indeed thermal in origin and originatesfrom random molecular motion in the candleflame.Applying <strong>Physics</strong> 28.2AurorasAt extreme northern latitudes, the aurora borealisprovides a beautiful and colorful display in thenight sky. A similar display occurs near thesouthern polar region and is called the auroraaustralis. What’s the origin of the various colorsseen in the auroras?Explanation The aurora is due to high speed particlesinteracting with the Earth’s magnetic field andentering the atmosphere. When these particles collidewith molecules in the atmosphere, they excite themolecules in a way similar to the voltage in thespectrum tubes discussed earlier in this section. Inresponse, the molecules emit colors of light accordingto the characteristic spectrum of their atomicconstituents. For our atmosphere, the primaryconstituents are nitrogen and oxygen, which providethe red, blue, and green colors of the aurora.+ er– em eFigure 28.5 Diagram representingBohr’s model of the hydrogen atom.The orbiting electron is allowed onlyin specific orbits of discrete radii.Fv28.3 THE BOHR THEORY OF HYDROGENAt the beginning of the 20th century, scientists were perplexed by the failure ofclassical physics to explain the characteristics of spectra. Why did atoms of a givenelement emit only certain lines? Further, why did the atoms absorb only those wavelengthsthat they emitted? In 1913 Bohr provided an explanation of atomic spectrathat includes some features of the currently accepted theory. Using the simplestatom, hydrogen, Bohr developed a model of what he thought must be the atom’sstructure in an attempt to explain why the atom was stable. His model of the hydrogenatom contains some classical features, as well as some revolutionary postulatesthat could not be justified within the framework of classical physics. The basic assumptionsof the Bohr theory as it applies to the hydrogen atom are as follows:1. The electron moves in circular orbits about the proton under the influence ofthe Coulomb force of attraction, as in Figure 28.5. The Coulomb force producesthe electron’s centripetal acceleration.
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An Abbreviated Table of Isotopes A.
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Some Useful Tables A.15TABLE C.3The
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IndexPage numbers followed by “f
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Current, 568-573, 586direction of,
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Index I.5Fissionnuclear, 973-976, 9
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South poleEarth’s geographic, 626
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CreditsPhotographsThis page constit
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PHYSICAL CONSTANTSQuantity Symbol V