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ATOMIC
STRUCTURE
© 2009 Brooks/Cole - Cengage
Atomic Structure
• Much of what we know about the very nature of
matter and the universe around us is due to the
work of pioneering chemists, mathematicians and
physicists in the late 19th and early 20th centuries
» No Computers, calculators, Starbucks, cell phones or even
ELECTRICITY
• This knowledge sprang from studies on light
» What is it?
» How do atoms interact with it?
» How is it made?
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Electromagnetic Radiation
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• In 1864
James Maxwell developed a mathematical
way to describe radiation
(what was going on in America at the time?),
• He said that radiation is a wave with electric and magnetic
fields at right angles to each other move together
• Since it is a wave, it has the following characteristics of all
waves
– Wavelength: λ (lambda) = The distance between successive crests of a wave.
It is measured in units of distance (nanometers, micrometers, meters)
– Frequency: ν (nu) = The number of waves that pass a given point in some
amount of time (usually per second). It is measured in Hertz (Hz) or s-1
• The speed of an electromagnetic wave is defined as:
c = λν
where c is the Universal Constant = 3.00 x108 m/s
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Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultraviolet radiation
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Node
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Electromagnetic Radiation
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Electromagnetic Radiation
Long wavelength , small frequency v
Short wavelength , high frequency v
increasing
frequency
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increasing
wavelength
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Electromagnetic Radiation
Red light has  = 700 nm. Calculate the frequency.
 1 x 10-9 m 
 = 7.00 x 10-7 m
700 nm 
1 nm


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Freq =
3.00 x 10 m/s
-7
7.00 x 10
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m
 4.29 x 1014 sec-1
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Electromagnetic Radiation
Short wavelength 
high frequency
high energy
Long wavelength 
small frequency
low energy
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Electromagnetic Spectrum
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Let’s look at an object being heated
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• The emitted light from a heated object comes from a
collection of oscillators
– Some at high energy, some at intermediate energy, some at
low energy
• In 1879, Josef Stefan determined that the total
intensity of all radiation emitted from a heated
object increases as the fourth power of
temperature
– Intensity=(5.67x10-8Wm-2K-4) · T4
» 1 Watt (W) = 1J/sec
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Wien’s Law
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Wilhelm Wien studied the relationship between
temperature and the wavelength of maximum intensity
in a black body emitter.
He found that as Temperature INCREASES, the
wavelength of maximum emission DECREASES
We can summarize this in Wien’s Law:
Tmax = Constant = 2.9 K•mm
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Quantization of Energy
(Planck and Einstein)
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• For centuries people have observed that as you heat
an object, it goes from red to orange-yellow to white
– The phrase “white hot” comes from this
• What are we actually observing?
• This emitted light is an indicator of the heat given off
by the object
• The problem scientists in the 1800’s had was that it
was theorized that the more heat you put into an
object, the higher the intensity of radiation that would
be emitted at decreasing wavelength
“The Ultraviolet Catastrophe”
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The Ultraviolet Catastrophe
• According to Classical Physics at the time,
having a cookout should turn into a
nightmare.
• The grill should be emitting x-ray and gamma
ray radiation
• But we know this doesn’t happen.
• Max Planck studied this and found…
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Quantization of Energy
•Planck deduced that
energy would be
quantized and this
explained the
“Catastrophe”
•With quantization, only
radiation of certain
energies would be
emitted
See Chem & Chem Reactivity, Figure 6.3
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Quantization of Energy
An object can gain or lose energy by absorbing or
emitting radiant energy in QUANTA.
Energy of radiation is proportional to frequency
E = h·
h = Planck’s constant = 6.6262 x 10-34 J·s
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Quantization of Energy
E = h·
Light with large  (small ) has a small E.
Light with a short  (large ) has a large E.
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Let’s look at an object being heated
• As we heat a metal bar, the atoms in the bar vibrate
faster
• The atoms are called oscillators
• As they drop back down to a lower vibrational state
they emit some radiation
• Each oscillator has a fundamental frequency and the
energy of the emitted radiation is a multiple of this
frequency (this is where n comes into play)
• For a single energy level change, the equation
becomes:
E=hν
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(Planck’s Equation)
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Energy of Radiation
Energy of 1.00 mol of photons of red light.
E = h·
= (6.63 x 10-34 J·s)(4.29 x 1014 s-1)
= 2.85 x 10-19 J per photon
E per mol =
(2.85 x 10-19 J/ph)(6.02 x 1023 ph/mol)
= 172 kJ/mol
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Photoelectric Effect
Experiment demonstrates the particle nature of light.
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Photoelectric Effect
Classical theory said that E of
ejected electron should increase
with increase in light intensity—not
observed!
• No e- observed until light of a
certain minimum E (or frequency, remember
Placnk’s equation?) is used.
– Once this value is reached, electrons are
immediately ejected
• Number of e- ejected depends on
light intensity.
• The kinetic energy of the ejected
electrons increases with the
frequency of the incident radiation
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A. Einstein (18791955)
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Photoelectric Effect
Understand experimental observations
if light consists of particles called
PHOTONS of discrete energy.
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Photoelectric Effect
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• Einstein explained the observations of
photo-electric experiments by combining
Planck’s equation with a new concept
– Light has particle-like properties
– Massless packets of energy are called PHOTONS (hv)
and the energy of the packets is proportional to their
frequency
• No electrons are ejected by the metal if the
incident photons do not have a high
enough energy
• If the frequency is high enough, the energy
is high enough and an electron is knocked
off
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A. Einstein (18791955)
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Photoelectric Effect
Let’s look at this in more detail:
If we have a stream of photons colliding with a metal object, some of
those photons are going to collide with the electrons in the metal
The photons have an energy associated with them (hv) but this value
must be above a certain minimum to eject an electron from the metal.
Different metals do not release electrons with the exact same incident
photons
The metals want to hold onto the electrons and have a characteristic
energy value associated with them called a WORK FUNCTION, 
If the energy of the incident photons is greater than , then the metal
releases electrons
1/2me
Ek of ejected
electron
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v2
= hv - 
Work function of
metal
Energy of incident
photon
Atomic Line Emission Spectra
and Niels Bohr
• It has long been known that applying high voltage
to a tube containing a gas would result in the gas
giving off light
• However, if we split the light into its component
wavelengths with a prism, we’ll see a small
number of lines at specific colors (wavelengths)
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Line Emission Spectra
of Excited Atoms
• Excited atoms emit light of only
certain wavelengths
• The wavelengths of emitted light
depend on the element.
QuickTime™ and a
Graphics decompressor
are needed to see this picture.
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Line Emission Spectra
of Excited Atoms
High E
Short 
High 
Low E
Long 
Low 
Visible lines in H atom spectrum are
called the BALMER series.
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Line Spectra of Other
Elements
Why do elements emit at certain characteristic wavelengths?
Balmer and Rydberg developed an explanation for the line emission
behaviour (Rydberg Formula)
 1
1 
  R 2  2 
n1 n 2 
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R = Rydberg Constant =
1.0974x10-3 m-1
Line Spectra
 1
1 
  R 2  2 
n1 n 2 
R = Rydberg Constant =
1.0974x10-3 m-1
When n1=2 (and n2=2, 3, 4…) You can calculate the Balmer Series of lines
When n1=1 (and n2=2, 3, 4…) You can calculate the Lyman Series of lines

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What do Line Spectra Tell Us?
The characteristic line spectra of each
element tells us that electrons can only
have certain SPECIFIC energies
(that’s what those n values mean, but more on that in a minute)
Each element has a unique configuration
of electrons as evidenced by their unique
line spectra
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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.
1.
Any orbit should be possible
and so is any energy.
2.
But a charged particle moving
in an electric field should emit
energy.
End result should be destruction!
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Atomic Spectra and Bohr
Bohr said classical view is wrong.
Need a new theory — now called QUANTUM
or WAVE MECHANICS.
e- can only exist in certain discrete orbits —
called stationary states.
e- is restricted to QUANTIZED energy states.
Energy of state = - Rhc/n2
where n = quantum no. = 1, 2, 3, 4, ....
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Atomic Spectra and Bohr
Energy of quantized state = - Rhc/n2
• Only orbits where n = some positive integer
are permitted.
• The energy of an electron in an orbit has a
negative value
• An atom with its electrons in the lowest
possible energy level is at GROUND STATE
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Atomic Spectra and Bohr
If e-’s are in quantized energy
states, then ∆E of states can have
only certain values. This explain
sharp line spectra.
PLAY MOVIE
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Energy Adsorption/Emission
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Origin of Line Spectra
Balmer series
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Atomic Line Spectra and
Niels Bohr
Niels Bohr
(1885-1962)
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Bohr’s theory was a great
accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
• theory only successful for H.
• introduced quantum idea artificially.
• So, we go on to QUANTUM or WAVE
MECHANICS
Wave-Particle Duality
L. de Broglie
(1892-1987)
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de Broglie (1924) proposed
that all moving objects have
wave properties.
For light: E = mc2
E = h = hc / 
Therefore, mc = h / 
and for particles
(mass)(velocity) = h / 
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Wave-Particle Duality
QuickTime™ and a
Graphics decompressor
are needed to see this picture.
Experimental proof of wave
properties of electrons
Baseball (115 g) at
100 mph
 = 1.3 x 10-32 cm
e- with velocity =
1.9 x 108 cm/sec
 = 0.388 nm
•The mass times the velocity of the ball is very large, so the wavelength is very
small for the baseball
•The deBroglie equation is only useful for particles of very small mass
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