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ELECTROMAGNETIC
RADIATION
The Wave Nature of Light
Much of our present understanding of
the electronic structure of atoms has
come from analysis of
the light emitted or absorbed
by substances
Electromagnetic Radiation
Radiant energy which carries energy
through space.
All types of electromagnetic radiation
move through a vacuum at
a speed of 3.00 x 108 m/s
Wave-like Nature of
Electromagnetic Radiation
Electromagnetic radiation is measured in
wavelenghts.
Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultaviolet radiation
Node
Since all electromagntic radiation travels
at the same velocity in vacuum, c,
its frequency, n, is
inversely proportional to
its
wavelength, l.
n l=c
Electromagnetic Radiation
• Waves have a frequency
• Use the Greek letter “nu”, n, for frequency,
and units are “cycles per sec”
• All radiation: l • n = c
• where c = velocity of light = 3.00 x 108 m/sec
• Note that long wavelength = small frequency
• Short wavelength = high frequency
Electromagnetic Spectrum
Indicates the wavelenghts of
electromagnetic radiation
characteristic of various regions of the
electromagnetic spectrum
Electromagnetic Radiation
Note that long wavelength = small frequency
Short wavelength = high frequency
See Screen 7.4
increasing
frequency
increasing
wavelength
Atomic Line Spectra and
Niels Bohr
Bohr’s greatest contribution
to science was in building
a simple model of the
atom. It was based on an
understanding of the
Niels Bohr
(1885-1962)
SHARP LINE
SPECTRA of excited
atoms.
Bohr’s Model of the Hydrogen
Atom
Line Spectra
Produced when gases are placed under
reduced pressure in a tube and a high voltage
is applied
- colored lines, separated by black regions are
produced
Line Spectra
In 1885, Johann Balmer observed that
the four lines of the hydrogen spectrum
fit a formula
Line Spectra
of Excited Atoms
High E
Short l
High n
Visible lines in H atom spectrum are
called the BALMER series.
Low E
Long l
Low n
n = C( 1/22 -1/n2 ) n = 3,4,5,6
C = 3.29 x 1015 s-1
Predicts the frequency of each line of the
hydrogen line spectra
Bohr also assumed the electron could
“jump” from one allowed energy state
to another.
• Energy is absorbed when electron
moves to a higher energy state.
• Energy is emitted when when
electron moves from higher to a
lower energy state
Orbital Energies
En = (-RH)(1/n2) n = 1,2,3,4….
RH = Rydberg constant
(2.18 x 10-18 J)
n = principle quantun number
Line Spectra
of Excited Atoms
• Excited atoms emit light of only
certain wavelengths
• The wavelengths of emitted light
depend on the element.
Atomic Spectra and Bohr
One view of atomic structure in early 20th
century was that an electron (e-) traveled
about the nucleus in an orbit.
+
Electron
orbit
1. Any orbit should be
possible and so is any energy.
2. But a charged particle
moving in an electric field
should emit energy.
Bohr Model stated that
electrons can only exist in certain
discrete orbits — called
stationary states.
Each electron is restricted to
QUANTIZED energy states.
n = quantum no. = 1, 2, 3, 4, ....
Atomic Spectra and Bohr
• Only orbits where n = integral no. are
permitted.
• Results can be used to explain atomic
spectra.
Atomic Spectra and Bohr
If electrons are in quantized energy
states, then DE of states can have only
certain values.
This explain sharp line spectra.
E = -C (1/2 2 )
E = -C (1/1 2 )
n=2
n=1
E
N
E
R
G
Y
E = -C ( 1 / 2 2 )
E = -C ( 1 / 1 2 )
n=2
n=1
Calculate DE for an electron “falling”
from high energy level (n = 2) to low
energy level (n = 1).
DE = Efinal - Einitial = -C[(1/12) - (1/2)2]
DE = -(3/4)C
E
N
E
R
G
Y
E = -C ( 1 / 2 2 )
E = -C ( 1 / 1 2 )
n=2
n=1
DE = -(3/4)C
C has been found from experiment (and is now
called R, the Rydberg constant)
R (= C) = 1312 kJ/mol or 3.29 x 1015
cycles/sec
so, E of emitted light
= (3/4)R = 2.47 x 1015 sec-1
and l = c/n = 121.6 nm
Atomic Line Spectra and
Niels Bohr
Niels Bohr
(1885-1962)
Bohr’s theory was a great
accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
• theory only successful for
H.
• introduced quantum idea
artificially.