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Light Spectra

Skills to develop

  • Explain the terms white light, line, emission, & absorption spectra
  • Interpret the regularity of hydrogen spectra
  • Light Spectra

    The intensity of light as a function of wavelength is a spectrum. To see a spectrum, a beam of electromagnetic radiation is spread according to wavelength or colour so that the intensity as a function of wavelength is represented. A visible spectrum has many colors, and a rainbow is a typical spectrum.

    Visible light is only in a small range of the total electromagnetic radiation spectrum and a partial spectrum may be in regions that are not visible. Classified according to region of radiation, we have infrared-, ultraviolet-, X-ray-, and gamma-ray- spectra.

    Kinds of Spectra

    Light beams (radiations) from (hot) solids may be spread (by a prism) into a continuous display of color, and such beams give continuous spectra or white spectra.

    Light beams from a gaseous sample usually show colored lines called line spectra or discontinued spectra. These continuous and line spectra are called emission spectra. The spectra of H, Hg, and Ne are line spectra.

    When a white light beam passes through a medium, light of some colors will be absorbed. As a result, the spectrum shows dark (absent) lines. Such a spectrum is called an absorption spectrum.

    The dark lines in the absorption spectrum of a gas correspond exactly to some of the bright lines in the emision spectrum of that gas.

    Spectra of the H Atom

    Since the invention of the spectrometer to analyze light, the light source from hydrogen gas has been intensely studied. Lines were observed in the infrared, visible, and ultraviolet regions. The wavelengths or frequencies of these lines had been known for a long time.

    The frequency (Hz) wave number (number per meter) and wavelength (nm) of some of the lines in the visible region are given below:

    in 1e14 Hz
    Wave number
    in e-19 J

    Many people tried to find the rules that govern their variation, and in the process many mathematical techniques have been employed. Among these attempts, there has been suggestions that the wavenumbers vary according to

    --- - K,

    where RH is the Rydberg constant, n an integer, and K some constant.

    At the time when Bohr worked on this problem, Rutherford had shown that atoms consisted of small dense nuclei surrounded by very light electrons. Thus, Bohr thought the atom might be a minature solar system with electrons revolving around (like the planet) the nucleus (the Sun). He also applied Planck's quantum concept and implied that when the angular momentum of the revolving electron is of a certain value, the electron orbital is stable. He derived a formula that showed the energy of the photons of these lines vary according to

                  1       1
      E = - RH ( ---  -  --- )
                 n2       4
    and RH = 2.179E-18 J, is the Rydberg constant.

    Depending on the units used for RH, the energy E of this formula can be wavenumbers or frequencies.

    Furthermore, the formula also applies to lines of the hydrogen spectrum in the infrared and ultraviolet regions if the 4 is replaced by the square of some other integer, n'. Thus, the formula has a general form of

                1       1
      E = - R ( ---  -  --- )
                n2       n'2
    Depending on the value of n', we have the following series: n' = 1, Lyman series (ultraviolet)
    n' = 2, Balmer series (visible) Wavelength vary from 400 to 700 nm
    n' = 3, Ritz-Paschen series (short wave infrared)
    n' = 4, Pfund series (long wave infrared)

    Thus, this formula agrees with all observed lines in the hydrogen spectrum. The above formula has been employed to calculate the spectra of the hydrogen atom.

    Confidence Building Questions