EMR info
Waves, light, and energy: Where
chemistry and physics collide
http://imagers.gsfc.nasa.gov/ems/waves3.html
First things first: Waves
a and b represent wavelength (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 (v; I know some of you have
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used f, move on and get with chemisty!):
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the number of cycles per second
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measured in cycles per second (s-1) or Hz (Hertz)
Propagation of electromagnetic waves
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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colour
wavelength(Å) f(*1014 Hz)
Energy (*10-19
J)
Violet
4000---4600
7.5--6.5
5.0--4.3
Indigo
4600---4750
6.5--6.3
4.3--4.2
Blue
4750---4900
6.3--6.1
4.2--4.1
Green
4900---5650
6.1--5.3
4.1--3.5
Yellow
5650---5750
5.3--5.2
3.5--3.45
Orange
5750---6000
5.2--5.0
3.45--3.3
Red6000---8000
5.0--3.7
3.3--2.5
Some equations you need to
know
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λ= c / v
and
E = hv
So….
E = hc / λ
And…
λ = h / mv
When
•  λ= wavelength in m
•  c = speed of light, 3.00E8 m/s
•  v = frequencyin Hz
• (cycles/sec or s-1 or 1/s)
• E= energy in J
• h= Plank’s constant, 6.626E-34 J*s
[Joule(seconds)]
• m= mass of particle in kg
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 = hv
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The energy (E) of a single quantum is
equal to its frequency (v) 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 (or Mr, Read, you
decide), 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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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
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Line spectra formation- go to…..
http://www.mhhe.com/physsci/
chemistry/essentialchemistry/flash/
linesp16.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
a Nitrogen
(z=7)
atom
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.edu/
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
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 that they 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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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
De Broglie wavelength
EMR and the atom: Part Deux
Electron Configurations:
from waves to orbitals
http://imagers.gsfc.nasa.gov/ems/waves3.html
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 because we don’t have enough time or
energy to devote to this)
n  Quantum theory: Treats e-s as quantized
particle-waves of energy that orbit the
nucleus in set paths
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Therefore, each location for an e- has a set energy
EMR and the atom: Part Deux
Electron Configurations:
from waves to orbitals
http://imagers.gsfc.nasa.gov/ems/waves3.html
Warning:
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Much of what we do will not make
sense until it is all put together
Use what we are going now as rules
for what we are going to do with this
information, and know that the
problems are not in this packet
The Bohr Model: What you’ve
seen so far in life
General tutorials for electron
configuration stuff
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some slides in this PowerPoint are from this
site already
http://www.wwnorton.com/chemistry/
overview/ch3.htm
See key equations and concepts (select
from menu on the left), as well as the
looking through the overview where to the
tutorials are listed (links for just those are on
the left, too)
What you’ve seen so far….
Model of
a Nitrogen
(z=7)
atom
Which is really not true- why?
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orbitals- the
“electron cloud” are
3-D, not flat
¡  are not round in most cases
¡  e-s are spread out as much
as possible
¡  (e-s are moving very rapidly)
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Orbitals
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The electrons are spread out in orbitals that
have varying
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Shapes
Energy (distance from nucleus)
The orbitals are described in regards to their
quantum numbers
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Descriptions that are descriptive and hierarchical
There are 4 numbers that describe an orbital
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Written as follows: (#, #, #, ±#)
Principal quantum number (n)
The first number (1, #, #,±#)
n  Describe the
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Values for n are integers
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distance from the nucleus of the orbital
The energy of the orbital
The smallest possible value is 1
As the distance from the nucleus (and
therefore energy) increases, the number
increases
Quantum numbers
The periodic table and n
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The 7 periods on the periodic table
correspond to n values
Each period has a unique n value
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For the 1st period, n=1
For the 2nd period, n=2
And so on….
Angular Momentum (l)
(this is a script l, as in llama)
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Is the shape of the orbital
It is the second number in the description (#,
1,#, ±#)
Range from 0 to n-1 (although we never
deal with anything above l=3)
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s =0
p =1
d =2
f=3
The s orbital
l=0
http://www.sfu.ca/~nbranda/28xweb/images/s_orbital.gif
p orbitals
l=1
d orbitals
l=2
Another look at d orbitals
f orbitals l=3
Blocks and l
The blocks you already know correspond to the orbital of the last
(outermost) e- , or valence e-s occupies
Magnetic number (ml)
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Denote the orbital sublevel that is filled
It is the third number in the description (#,#,
1, ±#)
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s orbitals have one sublevel; a sphere has one
orientation in space
p orbitals have three sublevels; 3 orientations in
space
d orbitals have five sublevels; 5 orientations in
space
f orbitals have seven sublevels; 7 orientations in
space
“Flavors” of ml
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s orbitals have
one sublevel; a
sphere has
one orientation
in space
“Flavors” of ml
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p orbitals have three sublevels; 3
orientations in space
“Flavors” of ml
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d orbitals have
five sublevels;
5 orientations
in space
“Flavors” of ml
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f orbitals
have seven
sublevels; 7
orientations
in space
Magnetic number (ml)
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Denote the orbital sublevel that is filled
It is the third number in the description (#,#,
1, ±#)
Values of –l to l, (integers only)
For
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s
p
d
f
ml = 0 only since l= 0
ml = -1,0,1 since l= 1
ml = -2,-1,0,1,2 since l= 2
ml = -3,-2,-1,0,1,2,3 since l= 3
Spin
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It is the last number in the description
(#,#,#,±½)
Spin is +½ or -½
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Up or down
Summary, excluding spin
How we use this….
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There is a specific order to how the efill the orbitals; it is not random
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Although there are exceptions to the rules
(which is the last thing we’ll do; pretend
for this moment that everything follows
these rules, as most elements do)
The principles of econfiguration
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The Aufbau (next) Principle:
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The Pauli Exclusion Principle:
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That e- fill the lowest energy sublevel before
going to the next sublevel
That e-s are paired according to opposite
spins
Hund’s Rule:
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e-s spread out in equal energy sublevels
before placing electrons in the next level
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The first level to fill is the 1s level
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It is the lowest energy sublevel
It holds two electrons
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They are oppositely paired (up and down- ↑↓)
Each sublevel (each __) holds 2 electrons
Next…
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The second sublevel is the 2s sublevel
It also holds 2 electrons
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because s orbitals have one sublevel,
(not 2 electrons because of the
number 2)*
also oppositely paired ↑↓
*the number is n=, so the number tells how far form the nucleus, not the number
of sublevels or electrons it holds
1s2, 2s2,then comes 2p6
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So, as it states above
¡  1s fills, 2s fills ,then
comes 2p
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It holds up to six
electrons
Because p orbitals
hold 6 electrons in
3 sublevels (_ _ _)
Next…
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From 2p,
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¡
¡
¡
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3s fills with 2e-, then onto
3p, with 6e- then
4s with 2e- followed by
3d with 10e- (because d holds 10e-)
Then 4p with 6e-
Notice, you follow the arrows
Remember, the number of
electrons comes from the letter
(the orbital’s momentum, m)
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The sublevels of the orbitals are first
filled, then you continue onto the
next level (Aufbau)
Also be sure to place one electron in
each sublevel prior to filling the level
(↑ ↑ ↑ and not ↑↓ ↑ _) (Hund)
e-s must be paired with e-s of
opposite spin (↑↓, not ↑↑ or ↓↓)
(Pauli)
Putting it all together…
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Carbon (neutral, so 6 electrons)
What this would look like:
↑↓ ↑↓ ↑ ↑ _
1s 2s 2p
(notice there are 6 arrows for 6 electrons)
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This can also be written as 1s2 2s2 2p2
Notice the superscripts add up to 6
There are some exceptions…
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This is because some energy levels are very close
together
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electrons are able to move between close orbitals in order
to minimize repulsion
Example: the 4s and 3d orbitals are very close in
energy
So exceptions for some period 4 d block elements
occur
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¡
¡
Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Because it takes less energy to split the electrons between
the 5 sublevels than it does to put them together in the 4s
and 3d
Electron configurations
Abbreviated e- configurations
Orbital diagrams