Slide 1Slide 2Slide 3Slide 4Slide 5Slide 6Slide 7Slide 8Slide 9Slide 10Slide 11Slide 12Slide 13Slide 14Slide 15Slide 16Slide 17Slide 18Slide 19Slide 20Slide 21Line SpectraWhen the particles in the solid, liquid, or gas accelerate, they will produce EM waves.Electron orbit to orbit transitions in atoms (gasses)Applicable to the study of stars (gaseous objects)Line SpectraAtomic StructureShell model of the atom (Bohr): electrons orbit the nucleus in hierarchy of stable orbits, each corresponding to a specific amount of electron energy.E1E2E3Line SpectraAtomic StructureThe normal state of the atom has the electron in the lowest energy orbit, called the ground state. The energy associated with the ground state is called the ground state energy.E1E2E3Line SpectraAtomic StructureEnergy Minimum Principle: An electron in orbit around a nucleus will orbit in the lowest available energy orbit.E1E2E3Line SpectraAtomic StructureExclusion Principle: No two electrons can exist in the same orbit in an atom.E1E2E3AllowedNot AllowedLine SpectraAtomic StructureElectrons can move, as a result of energy inputs to the atom, to a higher energy orbit. In this case, the electron is said to be in an excited state.E1E2E3Line SpectraEmission SpectrumIn accordance with the Energy Minimum Principle, the electron will then “jump” to a lower energy state. In doing so, it will give up a specific amount of energy through the emission of a photon.E1E2E3E3 – E2 = h f32Line SpectraEmission SpectrumIt will continue to cascade down until it reaches the ground state.E1E2E3E2 – E1 = h f21Line SpectraEmission SpectrumThe electron can also bypass intermediate orbits.E1E2E3E3 – E1 = h f31Line SpectraEmission SpectrumThe atom will “glow” at frequencies determined by the difference in energy between the various orbits in the atom.E1E2E3Line SpectraHydrogen SpectrumE1E2E3E4All atoms have an infinite number of energy levels (orbits)The energy corresponding to the ground state is the lowest electron energyThe energy corresponding to the first excited state is the second lowest energy.EtcLine SpectraHydrogen SpectrumE1E2E3E4The energy difference between orbits gets smaller and smaller as you go to higher and higher orbital energies.Line SpectraHydrogen SpectrumE1E2E3E4For hydrogenEn = -13.6 eVn2Therefore, the energy of emitted photons isEphoton = 13.6 eV (1/n2 – 1/p2)n is the quantum number of the final orbitp is the quantum number of the starting orbitLine SpectraHydrogen SpectrumNote: Negative electron energy means that the electron is bound to the nucleusLine SpectraHydrogen SpectrumEphoton = 13.6 eV ( 1/12 - 1/p2 ) p = 2,3,4 …p = 2,3,4 …Lyman SeriesLyman Seriesp = 3,4,5,…p = 3,4,5,…Balmer SeriesBalmer Seriesp = 4,5,6,..p = 4,5,6,..Paschen SeriesPaschen SeriesEEphotonphoton = hf = hc/λ = hf = hc/λEphoton = 13.6 eV ( 1/32 - 1/p2 ) Ephoton = 13.6 eV ( 1/22 - 1/p2 )Line SpectraHydrogen SpectrumLadder (Energy Level) DiagramEnergy-13.6 /12 -13.6 /22 -13.6 /32 -13.6 /42 Lyα LyγLyβEtc.Lyman series is UltravioletLine SpectraHydrogen SpectrumLadder (Energy Level) DiagramEnergy-13.6 /12 -13.6 /22 -13.6 /32 -13.6 /42 Hα HγHβ Etc.Balmer series is visible. Simply called the “H” lines rather than Balmer LinesLine SpectraHydrogen SpectrumLineLineλ (nm)λ (nm)LineLineλ (nm)λ (nm)LyαLyα122122UVUVHαHα656656VisibleVisibleLyβLyβ103103UVUVHβHβ486486VisibleVisibleLyγLyγ9797UVUVHγHγ434434VisibleVisibleLine SpectraHydrogen SpectrumLine SpectraHydrogen SpectrumHαHβHγWavelength(nm)Intensity656656486486434434http://www.nat.vu.nl/~dennis/elements/iron.htmlLine spectra become VERY complicated as the number of electrons in orbit (and therefore the number of protons in the nucleus) growLine SpectraEmission Spectrum of Neutral and Singly Ionized
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