Transcript Line Spectrum

Line Spectra
When 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 Spectra
Atomic Structure
Shell model of the atom (Bohr): electrons orbit the nucleus in hierarchy of
stable orbits, each corresponding to a specific amount of electron energy.
E1
E2
E3
Line Spectra
Atomic Structure
The 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.
E1
E2
E3
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Atomic Structure
Energy Minimum Principle: An electron in orbit around a nucleus will orbit in
the lowest available energy orbit.
E1
E2
E3
Line Spectra
Atomic Structure
Exclusion Principle: No two electrons can exist in the same orbit in an atom.
Allowed
E1
E2
E3
Not Allowed
Line Spectra
Atomic Structure
Electrons 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.
E1
E2
E3
Line Spectra
Emission Spectrum
In 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.
E1
E2
E3
E3 – E2 = h f32
Line Spectra
Emission Spectrum
It will continue to cascade down until it reaches the ground state.
E1
E2
E3
E2 – E1 = h f21
Line Spectra
Emission Spectrum
The electron can also bypass intermediate orbits.
E1
E2
E3
E3 – E1 = h f31
Line Spectra
Emission Spectrum
The atom will “glow” at frequencies determined by the difference in energy
between the various orbits in the atom.
E1
E2
E3
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Hydrogen Spectrum
All atoms have an infinite
number of energy levels
(orbits)
The energy corresponding to
the ground state is the
lowest electron energy
E4
E1
E2
The energy corresponding to
the first excited state is the
second lowest energy.
Etc
E3
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Hydrogen Spectrum
The energy difference
between orbits gets smaller
and smaller as you go to
higher and higher orbital
energies.
E4
E1
E2
E3
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Hydrogen Spectrum
For hydrogen
-13.6 eV
En =
n2
E4
E1
E2
Therefore, the energy of
emitted photons is
Ephoton = 13.6 eV (1/n2 – 1/p2)
n is the quantum number of the
final orbit
E3
p is the quantum number of the
starting orbit
Line Spectra
Hydrogen Spectrum
Note: Negative electron energy means that the
electron is bound to the nucleus
Line Spectra
Hydrogen Spectrum
Ephoton = 13.6 eV ( 1/12 - 1/p2 )
p = 2,3,4 …
Lyman Series
Ephoton = 13.6 eV ( 1/22 - 1/p2 )
p = 3,4,5,…
Balmer Series
p = 4,5,6,..
Paschen Series
Ephoton = 13.6 eV ( 1/32 - 1/p2 )
Ephoton = hf = hc/λ
Line Spectra
Hydrogen Spectrum
Lyman series is
Ultraviolet
Ladder (Energy Level) Diagram
Energy
-13.6 /42
-13.6 /32
-13.6 /22
Lyα
-13.6 /12
Lyβ
Lyγ
Etc.
Line Spectra
Hydrogen Spectrum
Balmer series is visible.
Simply called the “H”
lines rather than Balmer
Lines
Ladder (Energy Level) Diagram
Energy
-13.6 /42
-13.6 /32
Hα
-13.6
/22
-13.6 /12
Hβ
Hγ
Etc.
Line Spectra
Hydrogen Spectrum
Line
λ (nm)
Line
λ (nm)
Lyα
122
UV
Hα
656
Visible
Lyβ
103
UV
Hβ
486
Visible
Lyγ
97
UV
Hγ
434
Visible
Line Spectra
Hydrogen Spectrum
Line Spectra
Hydrogen Spectrum
Intensity
Hγ
Hβ
Hα
Wavelength
(nm)
434
486
656
Line Spectra
Emission Spectrum of Neutral and Singly
Ionized Iron
Line spectra become VERY complicated as the number of electrons in
orbit (and therefore the number of protons in the nucleus) grow
http://www.nat.vu.nl/~dennis/elements/iron.html