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Bohr Model of the Atom Electrons in Atoms nucleus (+) electron (-) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Hydrogen Spectral Lines
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Bohr Model of the AtomElectrons in Atoms
nucleus (+)electron (-)Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Used by permission of Christy Johannesson
Atomic SpectrumHow color tells us about atoms
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PrismWhite light is made up of all the colors of the visible spectrum.Passing it through a prism separates it.
Author: Thomas V. Green Jr.
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If the light is not whiteBy heating a gas or with electricity we can get it to give off colors.Passing this light through a prism does something different.
Author: Thomas V. Green Jr.
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Atomic SpectrumEach element gives off its own characteristic colors.Can be used to identify the atom.How we know what stars are made of.
Author: Thomas V. Green Jr.
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These are called line spectraunique to each element.These are emission spectraThe light is emitted given off.
Author: Thomas V. Green Jr.
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Line-Emission Spectrumground stateexcited stateENERGY IN PHOTON OUTCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
656 nm486 nm410 nm434 nmWavelength (nm)
37.bin
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Bohr Modelelectrons exist only in orbits with specific amounts of energy called energy levelsThereforeelectrons can only gain or lose certain amounts of energyonly certain photons are producedCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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Bohr Model
123456Energy of photon depends on the difference in energy levels
Bohrs calculated energies matched the IR, visible, and UV lines for the H atom
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chemnucleus
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Other ElementsEach element has a unique bright-line emission spectrum.
i.e. Atomic FingerprintHeliumBohrs calculations only worked for hydrogen!
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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Bohrs ExperimentKelter, Carr, Scott, Chemistry A Wolrd of Choices 1999, page 76
Animation by Raymond Chang All rights reserved.
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Copyright 2007 Pearson Benjamin Cummings. All rights reserved.(a) Electronic absorption transition(b) H2 emission spectrum (top), H2 absorption spectrum (bottom)
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Copyright 2007 Pearson Benjamin Cummings. All rights reserved.(a) Electronic absorption transition(b) H2 emission spectrum (top), H2 absorption spectrum (bottom)Lower-energyorbitHigher-energyorbitPhotone-e-
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continuous spectrumabsorption spectrumemission spectrumhot sourcegasabsorption spectrumemission spectrum
*Source: http://abyss.uoregon.edu/~js/21st_century_science/lectures/lec11.html
Hydrogen Spectral Lines
Lyman series(ultraviolet)Balmer series(visible)Paschen series(infrared)Frequency(hertz)101610151014
7 6 5 4 3 2 1 n =
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Copyright 2007 Pearson Benjamin Cummings. All rights reserved.(ultraviolet)(visible)(infrared)HYDROGEN SPECTRAL LINES
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Hydrogen Spectral LinesA B C D E F Lyman series (UV)A B C D E Balmer (Visible)A B C D Paschen (IR)E1E2E3E4E5E6EnergyBohrs model of the atom accounted mathematically for the energy of each of the transitions shown.
IRregion UVregion656 nm
486 nm
434 nm
410 nmDavis, Metcalfe, Williams, Castka, Modern Chemistry, 1999, page 97
ionization
*The energy of the light that is emitted from an atom is equal to the difference in energy between the excited state and the ground state. As Bohr predicted, the colors of light indicated by his theory matched those in the line spectrum of hydrogen.
In 1885, Johann Balmer showed that the frequencies of the lines observed in the visible region of the spectrum of hydrogen fit a simple equation: = constant (1/22 1/n2) where n = 2, 4, 5, 6, and these lines are known as the Balmer series.
Johannes Rydberg restated and expanded Balmers result in the Rydberg equation:
= (1/n21 1/n22),
where n1 and n2 are integers, n2 > n1, and , the Rydberg constant, has a value of 1.09737 x 107m-1.
Electronic Transitions in the Excited Hydrogen Atom
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Used by permission of Christy Johannesson*
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*Source: http://abyss.uoregon.edu/~js/21st_century_science/lectures/lec11.html*
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*The energy of the light that is emitted from an atom is equal to the difference in energy between the excited state and the ground state. As Bohr predicted, the colors of light indicated by his theory matched those in the line spectrum of hydrogen.
In 1885, Johann Balmer showed that the frequencies of the lines observed in the visible region of the spectrum of hydrogen fit a simple equation: = constant (1/22 1/n2) where n = 2, 4, 5, 6, and these lines are known as the Balmer series.
Johannes Rydberg restated and expanded Balmers result in the Rydberg equation:
= (1/n21 1/n22),
where n1 and n2 are integers, n2 > n1, and , the Rydberg constant, has a value of 1.09737 x 107m-1.*
