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The Bohr Atomic Theory
Moving toward a new atomic model….
Back to Planck and Einstein…
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Once the ideas of Planck and Einstein (re:
quanta) had been accepted, the study of atomic
and molecular structure began.
With the invention of the spectroscope (Bunsen &
Kirchhoff, Fig 2/175) the quantitative analysis of
the atomic spectra of elements began.
The range of frequencies of light emitted or
absorbed by an atom is called its spectrum (pl.
spectra).
Spectroscope:
If the light emitted by an gaseous element is passed
through a spectroscope, it produces a bright line
atomic spectrum that is characteristic to that element
(fingerprint).
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When an element is excited, atoms absorb
energy, then as e-s fall back to “ground state”,
energy is released in the form of light
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Visible line spectra called emission spectra of
discrete(distinct) coloured lines - specific
wavelengths = specific energies, not a
continuous spectrum
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Atoms absorbing energy - discrete black
lines - absorption spectrum.
Emission Spectra:
Emission Spectrum (Na)
Absorption Spectrum (Na)
http://jersey.uoregon.edu/vlab/elem
ents/Elements.html
Recall:
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Problems with Rutherford’s planetary
model re stability
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Atomic spectra not explained - b/c if
Rutherford’s model was correct,
continual acceleration means
continuous spectrum of light - spectral
evidence shows this is not the case!
The Confusion About Bright Line Spectra
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Niels Bohr – Danish physicist, working for
Thomson (raisin bun guy) suggested that line
spectra could be explained using ideas from the
quantum theory!
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Bohr began to work with Rutherford (1913) modified Rutherford's atomic theory.
Bohr's Model of the Atom
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54 years after line spectra were first
described qualitatively by Bunsen and
Kirchhoff.
28 years after line spectra were
described quantitatively by Jacob
Balmer.
Bohr's Postulates
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The atom has only specific allowable energy
levels (stationary states) & each stat.state
corresponds to the e-s occupying fixed circular
orbits around nucleus.
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While in stat.state, atoms do not emit energy.
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An e- changes stat. state by emitting or
absorbing a specific quantity of energy that is
exactly equal to the difference in energy between
the two stat. states. (Fig 4/176)
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Higher energy level  lower: loss of energy in
the form of light
Lower energy level  higher: energy is absorbed
Only certain quanta of light can be emitted or
absorbed by an atom, so the energy of the
electrons inside the atom is also quantized. In
other words, an electron can only have certain
allowable energies. (staircase analogy).
http://www.upscale.utor
onto.ca/PVB/Harrison/
BohrModel/Flash/Bohr
Model.html
H-spectra example
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If an e- that has been excited to the 3rd NRG
level & then falls from 3rd to 2nd NRG level,
emits a photon of red light with wavelength
656nm & red line is visible on emission spectrum.
Each line on e. spectrum corresponds to a
specific energy transition (I.e. 4th to 2nd = green,
5th to 2nd = blue)
Hydrogen
Evaluating Bohr's Model:
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Was reasonably successful at explaining
Mendeleev's Periodic Table
Periods result in filling an additional energy level
(i.e. Na: 3 levels)
A period ends when the maximum # of electrons
in the valence level is reached (i.e. 2, 8, 8, 18)
BUT - could only explain one-electron systems
(so still more complex than he thought).
Bohr Energy-Level Diagrams
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Same rationale as orbit diagrams, but the focus
is on the energy that the electrons have, rather
than the position of the electrons, so orbits are
not used.
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For a Bohr energy level diagram, start from the
bottom and work your way up!
Example:
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5e8e2e15p+
P
phosphorous
(3rd energy level)
(2nd energy level)
(1st energy level)
(protons)
(symbol)
(name)
(from group 15)
(from 8 elements in period 2)
(from 2 elements in period 1)
(from atomic number)
(from PT)
(lowercase)
Note: The diagram is drawn with the energy levels
appearing to be evenly spaced, BUT we know this is not the
case, b/c line spectra evidence tells us that energy levels are
increasingly closer together at higher energy levels!!
Homework:
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180 #1-5,8,9