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5.3 Atomic Emission Spectra and
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Chapter 5
Electrons In Atoms
5.1 Revising the Atomic Model
5.2 Electron Arrangement in Atoms
5.3 Atomic Emission Spectra
and the Quantum
Mechanical Model
1
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CHEMISTRY
& YOU
What gives gas-filled lights their colors?
An electric current
passing through the gas
in each glass tube
makes the gas glow
with its own
characteristic color.
2
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5.3 Atomic Emission Spectra and
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Light and Atomic
Emission Spectra
Light and Atomic Emission Spectra
What causes atomic emission
spectra?
3
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Light and Atomic Emission
Spectra
The Nature of Light
• By the year 1900, there was
enough experimental evidence to
convince scientists that light
consisted of waves.
• The amplitude of a wave is the
wave’s height from zero to the
crest.
• The wavelength, represented by
 (the Greek letter lambda), is the
distance between the crests.
4
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Light and Atomic Emission
Spectra
The Nature of Light
• The frequency, represented by  (the
Greek letter nu), is the number of wave
cycles to pass a given point per unit of
time.
• The SI unit of cycles per second is called
the hertz (Hz).
5
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Light and Atomic Emission
Spectra
The Nature of Light
The product of frequency and wavelength
equals a constant (c), the speed of light.
c = 
 = Wavelength
C = Speed of light
v = frequency
λ= c/ν shows the relationship between
wavelength and frequency of light.
6
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Light and Atomic Emission
Spectra
• The frequency () and wavelength () of
light are inversely proportional to each
other.
• As the wavelength increases, the
frequency decreases.
7
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5.3 Atomic Emission Spectra and
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Light and Atomic Emission
Spectra
The Nature of Light
According to the wave model, light consists of
electromagnetic waves.
• Electromagnetic radiation includes
radio waves, microwaves, infrared
waves, visible light, ultraviolet waves,
X-rays, and gamma rays.
• All electromagnetic waves travel in a
vacuum at a speed of 2.998  108 m/s.
8
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Light and Atomic Emission
Spectra
The Nature of Light
The sun and incandescent light bulbs emit white
light, which consists of light with a continuous
range of wavelengths and frequencies.
• When sunlight passes through a prism, the
different wavelengths separate into a
spectrum of colors.
• In the visible spectrum, red light has the
longest wavelength and the lowest
frequency.
9
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Light and Atomic Emission
Spectra
The electromagnetic spectrum consists of
radiation over a broad range of wavelengths.
Low energy
( = 700 nm)
Frequency  (s-1)
3 x 106
102
Wavelength  (m)
10
High energy
( = 380 nm)
3 x 1012
3 x 1022
10-8
10-14
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5.3 Atomic Emission Spectra and
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Light and Atomic Emission
Spectra
Atomic Emission Spectra
When atoms absorb energy, their
electrons move to higher energy
levels. These electrons lose energy by
emitting light when they return to
lower energy levels.
11
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Light and Atomic
Emission Spectra
Atomic Emission Spectra
A prism separates light into the colors it
contains. White light produces a rainbow
of colors.
Screen
Light
bulb
12
Slit
Prism
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Light and Atomic Emission
Spectra
Atomic Emission Spectra
Light from a helium lamp produces
discrete lines.
Screen
Helium
lamp
13
Slit
Prism
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Light and Atomic Emission
Spectra
Atomic Emission Spectra
• The energy absorbed by an electron for it to move
from its current energy level to a higher energy level
is identical to the energy of the light emitted by the
electron as it drops back to its original energy level.
• The wavelengths of the spectral lines are
characteristic of the element, and they make up the
atomic emission spectrum of the element.
• No two elements have the same emission spectrum.
14
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Sample Problem 5.2
Calculating the Wavelength of Light
Calculate the wavelength of the
yellow light emitted by a
sodium lamp if the frequency of
the radiation is 5.09 × 1014 Hz
(5.09 × 1014/s).
15
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Sample Problem 5.2
Solve for the unknown.
Rearrange the equation to solve for .
c = 
c
= 
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Solve for  by dividing
both sides by :
c

 = 
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5.3 Atomic Emission Spectra and
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Sample Problem 5.2
2 Calculate Solve for the unknown.
Substitute the known values for  and c into
the equation and solve.
c
2.998  108 m/s
–7 m
=
=
=
5.89

10

5.09  1014 /s
17
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5.3 Atomic Emission Spectra and
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Sample Problem 5.2
3 Evaluate Does the answer make sense?
The magnitude of the frequency is much
larger than the numerical value of the
speed of light, so the answer should be
much less than 1. The answer should have
3 significant figures.
18
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5.3 Atomic Emission Spectra and
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What is the frequency of a red laser
that has a wavelength of 676 nm?
19
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5.3 Atomic Emission Spectra and
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What is the frequency of a red laser
that has a wavelength of 676 nm?
c = 
c
=

c 2.998  108 m/s
 =  = 6.76  10–7 m = 4.43  1014 /s
20
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Sample Problem
• A typical yellow light has a
wavelength of 550 nm. Calculate the
corresponding frequency. (speed of light
3.0 x 108 m/s)
v = c/ 
v = 3.0 x 108 m/s / 5.5 x 10-7 m
v = 5.5 x 1014 s-1
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5.3 Atomic Emission Spectra and
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A typical X-ray has a frequency of
3x1019 Hz. Calculate its wavelength
in picometer, pm. (1 pm = 1 x 10-12 m
= c/v
= 3.0 x 108 m/s / 3 x 1019 Hz
= 1 x 10-11 m or 10 pm
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5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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The Quantum Concept
and Photons
The Quantum Concept and Photons
How did Einstein explain the
photoelectric effect?
23
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The Quantum Concept
and Photons
The Quantization of Energy
German physicist Max Planck (1858–1947)
showed mathematically that the amount of
radiant energy (E) of a single quantum absorbed
or emitted by a body is proportional to the
frequency of radiation ().
E
24
 or E = h
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The Quantum Concept
and Photons
The Quantization of Energy
The constant (h), which has a value of 6.626  10–34
J·s (J is the joule, the SI unit of energy), is called
Planck’s constant.
E
25
 or E = h
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The Quantum Concept
and Photons
The Photoelectric Effect
Albert Einstein used Planck’s quantum theory to
explain the photoelectric effect.
In the photoelectric effect, electrons
are ejected when light shines on a
metal.
26
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5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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5.3 Atomic Emission Spectra and
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The Quantum Concept
and Photons
The Photoelectric Effect
Not just any frequency of light will cause the
photoelectric effect.
• Red light will not cause potassium to eject
electrons, no matter how intense the light.
• Yet a very weak yellow light shining on
potassium begins the effect.
28
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The Quantum Concept
and Photons
The Photoelectric Effect
• The photoelectric effect could not be explained
by classical physics.
• Classical physics correctly described light as a
form of energy.
• But, it assumed that under weak light of any
wavelength, an electron in a metal should
eventually collect enough energy to be ejected.
29
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The Quantum Concept
and Photons
The Photoelectric Effect
To explain the photoelectric effect, Einstein
proposed that light could be described as
quanta of energy that behave as if they were
particles.
30
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The Quantum Concept
and Photons
The Photoelectric Effect
These light quanta are called photons.
• Einstein’s theory that light behaves as a
stream of particles explains the photoelectric
effect and many other observations.
31
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The Quantum Concept
and Photons
The Photoelectric Effect
These light quanta are called photons.
• Einstein’s theory that light behaves as a
stream of particles explains the photoelectric
effect and many other observations.
• Light behaves as waves in other situations; we
must consider that light possesses both
wavelike and particle-like properties.
32
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5.3 Atomic Emission Spectra and
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The Photoelectric Effect
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5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
>
The Quantum Concept
and Photons
The Photoelectric Effect
No electrons are ejected
because the frequency
of the light is below the
threshold frequency.
34
If the light is at or above
the threshold frequency,
electrons are ejected.
If the frequency is
increased, the ejected
electrons will travel
faster.
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5.3 Atomic Emission Spectra and
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Sample Problem 5.3
Calculating the Energy of a Photon
What is the energy of
a photon of
microwave radiation
with a frequency of
3.20 × 1011/s?
35
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5.3 Atomic Emission Spectra and
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Sample Problem 5.3
1 Analyze List the knowns and the unknown.
Use the equation E = h ×  to calculate
the energy of the photon.
KNOWNS
frequency () = 3.20 × 1011/s
h = 6.626 × 10–34 J·s
UNKNOWN
energy (E) = ? J
36
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5.3 Atomic Emission Spectra and
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Sample Problem 5.3
2 Calculate Solve for the unknown.
Write the expression that relates the
energy of a photon of radiation and the
frequency of the radiation.
E = h
37
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5.3 Atomic Emission Spectra and
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Sample Problem 5.3
2 Calculate Solve for the unknown.
Substitute the known values for  and h
into the equation and solve.
E = h = (6.626  10–34 J·s)  (3.20  1011/s)
= 2.12  10–22 J
38
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Sample Problem 5.3
3 Evaluate Does the result make sense?
Individual photons have very small
energies, so the answer seems
reasonable.
39
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What is the frequency of a photon
whose energy is 1.166  10–17 J?
40
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5.3 Atomic Emission Spectra and
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>
What is the frequency of a photon
whose energy is 1.166  10–17 J?
E = h
E
=
h
E 1.166  10–17 J
16 Hz
=
=
=
1.760

10
h 6.626  10–34 Js
41
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An Explanation of Atomic
Spectra
An Explanation of Atomic Spectra
How are the frequencies of light
emitted by an atom related to changes
of electron energies?
42
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An Explanation of Atomic
Spectra
When an electron has its lowest possible
energy, the atom is in its ground state.
• In the ground state, the principal quantum
number (n) is 1.
43
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An Explanation of Atomic
Spectra
When an electron has its lowest possible
energy, the atom is in its ground state.
• In the ground state, the principal quantum
number (n) is 1.
• Excitation of the electron by absorbing
energy raises the atom to an excited state
with n = 2, 3, 4, 5, or 6, and so forth.
• A quantum of energy in the form of light is
emitted when the electron drops back to a
lower energy level.
44
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An Explanation of Atomic
Spectra
The light emitted by an electron moving
from a higher to a lower energy level
has a frequency directly proportional to
the energy change of the electron.
45
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An Explanation of Atomic
Spectra
The three groups of lines in the hydrogen spectrum
correspond to the transition of electrons from
higher energy levels to lower energy levels.
46
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5.3 Atomic Emission Spectra and
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CHEMISTRY
& YOU
The glass tubes in lighted signs contain helium, neon,
argon, krypton, or xenon gas, or a mixture of these
gases. Why do the colors of the light depend on the
gases that are used?
47
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5.3 Atomic Emission Spectra and
the Quantum Mechanical Model
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CHEMISTRY
& YOU
The glass tubes in lighted signs contain helium, neon,
argon, krypton, or xenon gas, or a mixture of these
gases. Why do the colors of the light depend on the
gases that are used?
Each different gas has
its own characteristic
emission spectrum,
creating different colors
of light when excited
electrons return to the
ground state.
48
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5.3 Atomic Emission Spectra and
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In the hydrogen spectrum, which of
the following transitions produces a
spectral line of the greatest energy?
A. n = 2 to n = 1
B. n = 3 to n = 2
C. n = 4 to n = 3
49
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In the hydrogen spectrum, which of
the following transitions produces a
spectral line of the greatest energy?
A. n = 2 to n = 1
B. n = 3 to n = 2
C. n = 4 to n = 3
50
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Quantum Mechanics
Quantum Mechanics
How does quantum mechanics differ
from classical mechanics?
51
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Quantum Mechanics
Given that light behaves as waves and
particles, can particles of matter behave as
waves?
52
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Quantum Mechanics
Given that light behaves as waves and
particles, can particles of matter behave as
waves?
• Louis de Broglie referred to the wavelike
behavior of particles as matter waves.
• His reasoning led him to a mathematical
expression for the wavelength of a moving
particle.
53
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Quantum Mechanics
The Wavelike Nature of Matter
Today, the wavelike properties of beams of electrons
are useful in viewing objects that cannot be viewed
with an optical microscope.
54
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Quantum Mechanics
The Wavelike Nature of Matter
Today, the wavelike properties of beams of electrons
are useful in viewing objects that cannot be viewed
with an optical microscope.
• The electrons in an electron
microscope have much smaller
wavelengths than visible light.
• These smaller wavelengths allow a
much clearer enlarged image of a
very small object, such as this pollen
grain, than is possible with an ordinary
microscope.
55
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Quantum Mechanics
Classical mechanics adequately
describes the motions of bodies much
larger than atoms, while quantum
mechanics describes the motions of
subatomic particles and atoms as
waves.
56
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Quantum Mechanics
The Heisenberg Uncertainty Principle
The Heisenberg uncertainty principle
states that it is impossible to know both the
velocity and the position of a particle at the
same time.
57
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Quantum Mechanics
The Heisenberg Uncertainty Principle
The Heisenberg uncertainty principle
states that it is impossible to know both the
velocity and the position of a particle at the
same time.
• This limitation is critical when dealing with
small particles such as electrons.
• But it does not matter for ordinary-sized
objects such as cars or airplanes.
58
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Quantum Mechanics
• To locate an electron, you might strike it with a photon.
• The electron has such a small mass that striking it with a
photon affects its motion in a way that cannot be
predicted accurately.
• The very act of measuring the position of the electron
changes its velocity, making its velocity uncertain.
Before collision:
A photon strikes
an electron
during an attempt
to observe the
electron’s
position.
59
After collision:
The impact
changes the
electron’s
velocity, making it
uncertain.
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The Heisenberg uncertainty principle
states that it is impossible to
simultaneously know which two
attributes of a particle?
60
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The Heisenberg uncertainty principle
states that it is impossible to
simultaneously know which two
attributes of a particle?
velocity and position
61
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Key Concepts and Key
Equations
When atoms absorb energy, their electrons move to
higher energy levels. These electrons lose energy by
emitting light when they return to lower energy
levels.
To explain the photoelectric effect, Einstein proposed
that light could be described as quanta of energy that
behave as if they were particles.
The light emitted by an electron moving from a higher
to a lower energy level has a frequency directly
proportional to the energy change of the electron.
62
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Key Concepts and Key
Equations
Classical mechanics adequately describes the
motions of bodies much larger than atoms, while
quantum mechanics describes the motions of
subatomic particles and atoms as waves.
C = 
E=h
63
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Glossary Terms
1. amplitude: the height of a wave’s crest
2. wavelength: the distance between adjacent
crests of a wave
3. frequency: the number of wave cycles that
pass a given point per unit of time; frequency
and wavelength are inversely proportional to
each other
4. hertz: the unit of frequency, equal to one cycle
per second
64
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Glossary Terms
5. electromagnetic radiation: energy waves that
travel in a vacuum at a speed of 2.998  108
m/s; includes radio waves, microwaves,
infrared waves, visible light, ultraviolet waves,
X-rays, and gamma rays
6. spectrum: wavelengths of visible light that are
separated when a beam of light passes through
a prism; range of wavelengths of
electromagnetic radiation
65
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Glossary Terms
7. atomic emission spectrum: the pattern formed
when light passes through a prism or diffraction
grating to separate it into the different
frequencies of light it contains
8. Planck’s constant: the constant (h) by which
the amount of radiant energy (E) is
proportional to the frequency of the radiation
()
9. photoelectric effect: the phenomenon in which
electrons are ejected when light shines on a
metal
66
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Glossary Terms
10. photon: a quantum of light; a discrete bundle
of electromagnetic energy that interacts with
matter similarly to particles
11. ground state: the lowest possible energy of an
atom described by quantum mechanics
12. Heisenberg uncertainty principle: it is
impossible to know both the velocity and the
position of a particle at the same time
67
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BIG IDEA
Electrons and the Structure of Atoms
• Electrons can absorb energy to move from
one energy level to a higher energy level.
• When an electron moves from a higher
energy level back down to a lower energy
level, light is emitted.
68
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END OF 5.3
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