EMR info
Waves, light, and energy: Where
chemistry and physics collide
http://imagers.gsfc.nasa.gov/ems/waves3.html
Before we get started….
1.  What is light?
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List as many interactions of
light and matter as you can.
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3.
Is it matter?
What forms of light exist?
think how light changes matter,
and how matter changes light
What are some uses of light?
First things first: Waves
a and b represent different wavelengths (λ)- the distance of a
wave from crest to successive crest; measured in meters
http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec.html
Waves: amplitude
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The height of a wave from crest to midline or trough to
midline; measured in meters
Terms you need to know:
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Wavelength (λ)
Amplitude
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Frequency (ⱱ) ; I know some of you have
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used f, move on and get with chemistry! :)
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the number of cycles (oscillations) per second
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measured in cycles per second (s-1) or Hz (Hertz)
Waves on a string
http://www.lbl.gov/MicroWorlds/ALSTool/EMSpec/EMSpec2.html
http://micro.magnet.fsu.edu/primer/lightandcolor/images/electromagneticfigure1.jpg
http://lepus.physics.ualr.edu/~tahall/EXAM2/emspec.jpg
http://www.arpansa.gov.au/images/emsline2.gif
Visible Light
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ⱱ (*1014 Hz)
Energy (*10-19 J)
Violet
400---460
7.5--6.5
5.0--4.3
Indigo
460---475
6.5--6.3
4.3--4.2
Blue
475---490
6.3--6.1
4.2--4.1
Green
490---565
6.1--5.3
4.1--3.5
Yellow
565---575
5.3--5.2
3.5--3.45
Orange
575---600
5.2--5.0
3.45--3.3
Red600---800
5.0--3.7
3.3--2.5
color
wavelength(nm)
Some equations you need to
know
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λ= c / ⱱ
and
E = hⱱ
So….
E = hc / λ
And…
λ = h / mv*
When
•  λ= wavelength in m
•  c = speed of light, 3.00E8 m/s
•  ⱱ (nu)= frequency in Hz
• (cycles/sec or s-1 or 1/s)
• E= energy in J
• h= Planck’s constant, 6.626E-34 J*s
[Joule(seconds)]
• m= mass of particle in kg
• V*= velocity in m/s
What the h? Planck’s Constant
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When metals are heated, they glow
1800s- physicists were trying to determine
the relationship between the color
(wavelength) and intensity of the glow
Max Planck (1900)- energy can be released
or absorbed only in little chunks (packets) of
energy “of some minimal size”
Max Planck and the h
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The chunks of energy were dubbed
“quantum” (“fixed amount”), which is
the smallest amount that can be
emitted or absorbed as EMR.
Proposed: E = hⱱ
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The energy (E) of a single quantum is
equal to its frequency (ν) times a
constant
Planck and the Nobel
(Physics)
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Planck determined that h= 6.626E-34 J-s
Energy is always released in multiples
of hv (1hv, 2hv, 3hv, etc)
h is so small that we cannot see the
effects of this in our daily lives
Analogous to…
Planck won the 1918 Nobel Prize in
physics for his work
Einstein & Bohr: Perfect
Together
Einstein, left
Bohr, above
Einstein:
The Photoelectric
Effect
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Einstein discovered
that one could cause
electrons to be ejected
from the surface of a
metal if the energy of
the light wave was
strong enough
He treated the light
needed to do this as a
piece of matter- a
photon, if you will
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This ejection of eis the
photoelectric
effect
The Photoelectric Effect
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Only light of a certain energy could
knock off an electron from the metal
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Intense light of a weaker wavelength
would not work, but even a low intensity
of the correct wavelength would work
(the energy of the light is transferred to
the kinetic energy of the electron)
Hmmm… light acting as a particle and
as a wave…..
The photoelectric effect…
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Online animations
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PhET
http://www.lewport.wnyric.org/mgagnon/
Photoelectric_Effect/
photoelectriceffect1.htm
http://www.xmission.com/~locutus/
applets/Photoelectric.html
Getting to Bohr….
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Light of a given
wavelength is
monochromatic
(one color)
Most common EMR
sources are
polychromatic, but
we see only one
color
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These can be reduced
to a spectrum when the
different wavelengths
are separated out
Spectral Emissions
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Continuous spectrum: shows all colors
of the rainbow
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Bright line spectrum:
only certain
wavelengths are
visible (the rest do not
appear at all)
Different elements
have different bright
line spectrum when
they are heated
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Na is yellow
Ne is orange-red
Line spectrum
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Ne
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I2
http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm
Hydrogen Spectra
Emission Spectra
Absorption Spectra
http://www.mhhe.com/physsci/astronomy/applets/Bohr/content_files/section1.html
http://www.cartage.org.lb/en/themes/Sciences/Astronomy/Modenastronomy/Interactionoflight/AtomicAbsorption/AtomicAbsorption.htm
Color and what you see:
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Absorption: the wavelengths that are
absorbed by an object are not
available for us to see, as we see the
wavelengths of light that are reflected
off of an object
This is not the same as those
wavelengths that are emitted by an
object that is emitting radiant
energy.
Color and what you see…
Chlorophyll absorption spectra
Perception of color
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Line spectra formation- go to…..
http://www.mhhe.com/physsci/
chemistry/essentialchemistry/flash/
linesp16.swf
http://www.mhhe.com/physsci/
chemistry/animations/chang_7e_esp/
pem1s3_1.swf
Bohr Model and Spectral Emissions
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Bohr proposed that the emission of
light energy from an (electrically or
thermally) excited atom corresponds to
the orbit of the electron around the
nucleus of the atom
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That energy can only be achieved by
being a specific distance from the
nucleus
What you’ve seen so far….
Model of an Iodine atom (atomic number =53)
Bohr Model and moving
electrons
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http://
www.colorado.edu/
physics/2000/
quantumzone/
bohr.html
Energy levels- Bohr Model
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Electrons travel within set
energy levels that have a
particular energy associated
with each level
After all, the e-s are moving
around the nucleus
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think KE here
Each shell has a number
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Closest to the nucleus is n=1
For each successive level add
1 to n
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n=2, n=3, ect….
Energy increases as the distance
from the nucleus increases
Bohr Model and moving
electrons
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http://
www.colorado.e
du/physics/
2000/
quantumzone/
bohr.html
Electron config in energy level
SO…
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We know that the e-’s are free to move
around the nucleus
They also can move from one energy
level to the next (and fall) back when
energy is added
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Move from ground state (“home” level) to
a higher level (the “excited” state)
Returning back to the ground state
releases energy
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This emission is how we see colors:
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the wavelengths of EMR released from
an atom when it has been excited by
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Heat energy
Electrical energy
Chemical energy
Think glowing red hot metal, or fireworks
Determining Energy for n
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To determine the energy for a given
energy level, use the equation:
En=(-RH)(Z/n2)
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RH = 2.18E-18J,
Z= the atomic number of the atom
n=1, 2, 3, 4….
So En=(-2.18E-18J)(Z/n2)
To determine E emitted or absorbed:
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To determine the change in energy for
a given energy transition:
ΔE=Ef-Ei
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*Remember E=hⱱ, so ΔE=hⱱ
so ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i
Remember that + values mean E that
is absorbed, and – values mean
released
E changes continued
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*Remember E=hν, so ΔE=hⱱ to get the
frequency of the light emitted or absorbed
If ΔE is positive
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since Ef >Ei
E is absorbed
The e- was going from ground state to an
excited state
If ΔE is negative
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since Ef < Ei
E is released
The e- was going to ground state from an
excited state
To determine E emitted or absorbed:
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What is the change in energy
associated with an electron dropping
from n=5 to n=1 in a Hydrogen atom?
ΔE=Ef-Ei
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so ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i
ΔE=[(-2.18E-18J)(1/12)]f- [(-2.18E-18J)(1/52)]I
ΔE = -2.09E-18 J
Which means 2.09E-18J are released
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Makes sense; an e- is dropping from 5 to1, E is
released when e- drop
Back to basics EMR calcs…
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That released Energy can be used to
determine the wavelength and
frequency of the EMR emitted.
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Remember that you need to treat the
energy as positive to do this!
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The sign only gives direction of energy flow
There is no negative energy, only energy
leaving
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If you used – energy, you’d get a -λ or -ⱱ
This isn’t possible!
Also…life after Einstein and
Bohr
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We know that electrons have
characteristics of both light (waves)
and matter, so we say that they have a
dual nature
De Broglie
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De Broglie proposed that an electron moving
about the nucleus had a wave-like behavior,
so it has a particular wavelength associated
with it. This wavelength depends upon the
mass and velocity of the electron.
¡  λ = h / mv
¡  mv = the momentum of the particle
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Mass* velocity = p
momentum = p so
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therefore λ = h / p
p = mv
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This matter-wave idea applies to all
matter, not just to electrons
However, the mass is so large, and the
wavelength so small, that we cannot
see it in macroscale objects
This matter-wave theory led to
applications like the electron
microscope
Scanning electron microscope image of a leaf from a Black Walnut tree.
Image shows a cross-section of a cut leaf, itsupper epidermal layer,
mesophyll layer with palisade cells and vascular bundles, and lower
epidermal layer. The protrusion at center is just over 50 microns tall.
(Dartmouth Electron Microscope Facility/Dartmouth College)#
Pollen from a variety of common plants: sunflower, morning glory, hollyhock,
lily, primrose and caster bean. The largest one at center is nearly 100
microns wide. (Dartmouth Electron Microscope Facility/Dartmouth College) #
De Broglie wavelength
Heisenberg:
The Uncertainty Principle
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We can’t determine information about
small scale objects the same way we
can for large scale objects
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Case in point: a ball rolling down a rampwe can get position, direction, and speed
at the same time
We can’t for electrons
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Hence, the uncertainty principle
Heisenberg, cont’d
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It is inherently impossible for us to simultaneously
know both the exact momentum and exact location
of an electron
This is because anything we do to determine the
location or momentum of the electron moves it from
its original path and location; this can’t be reduced
past a certain minimal level
We can know only momentum or location- not both
We can talk probability of the location/ momentum
of an electron
Which brings us to this
question:
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What the heck does all of this have to
do with electron configuration and how
matter behaves?
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On to electron configuration, courtesy of
Schrödinger and company (enter math
that we’ll skip)
n  Quantum theory