Download Light, Energy, and More

Light, Energy, and More
Chemistry
Recap…
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Electromagnetic
Spectrum
High Energy
Low Energy
Wave Nature of
Light
What’s Going On Here?
When we heat metal what
happens?
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Does the wave model of light explain
these changes?
• Does not explain different wavelengths and
frequencies at different temperatures
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What is light?
• Radiation….what is radiation?
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Particles or rays of energy
What is temperature anyways?
• The measure of the average kinetic energy of
the particles in an object
• Kinetic Energy vs. Potential Energy
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Too many questions….
Max Planck (1900)
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German Physicist
Began to look for
answers
Matter can only
gain or lose energy
in small quantized
amounts
What’s quantized?
Vocab Word!!!
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QUANTUM
• Minimum amount of energy that can be
gained or lost by an atom
• The emitted light from a glowing metal
is a ENERGY…this energy is quantized
If energy is now quantized…how
can we determine the amount of
energy of a quantum?
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What is energy measured in?
What are we observing?
What happens to the color when we
increase the temperature (energy)?
• Proportional or inversely proportional?
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Now we need a constant…
• Planck’s constant, h=6.626 x 10^-34
J*s
Time to put these words into action!
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What is the frequency and
wavelength electromagnetic radiation
that emits 1.68 x 10^-17 J of
energy? What type of
electromagnetic radiation is this?
Wavelength= 1.18 x 10^-8 m
Ultraviolet radiation
Some questions to answer…
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What is the color we see?
What happens to the energy of the
radiation when we increase the
frequency, v, of the radiation
emitted?
Iron at room temp…color and E?
Iron with a little heat…color and E?
Iron with lots of heat…color and E?
According to Planck’s Theory…
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If we have a given v,
matter can emit or absorb
E only in whole number
multiples of hv (1hv, 2hv,
3 hv…)
Matter can ONLY have
specific amounts of energy
Wall of kids building blocks
• We can only add or take
away in increments of
whole blocks…we cannot
remove half a block
The Big Mystery of the 1900’s…
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The Photoelectric
Effect…
• What caused these
color changes in
metals???
Photoelectric Effect
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Electrons (photoelectrons) are
emitted from a metal’s surface when
a light of a certain frequency shines
on the surface
Certain specific amounts of energy
(what’s this called???) needed to
knock out electrons from metal
atoms.
Albert Einstein (1905)
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Added onto Planck’s
Theory…
Called the electron’s
emitted, PHOTONS (the
little energy packets
Planck called quantums)
Now… Ephoton = hv
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Planck paved the
way for the
explanation behind
the mystery
But some one else
came into the
picture…
Now light is not just a wave…
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Einstein’s Dual Nature of Light
• Particle and wave characteristics
• Light is a beam of tiny particles, called
photons, acting like a wave
NEW WORD!!!
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Photon
A particle of electromagnetic
radiation with no mass that carries a
quantum of energy
What Einstein added…
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Energy of a photon has a minimum
or threshold value to eject
photoelectrons
What must happen for the
photoelectric effect to occur?
• Energy of a photon (particle of EM
radiation) must have the minimum
energy requirement to free the electron
from the atom of metal
Mystery Solved!
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No matter how long a
light of a certain
frequency is shone on
metal (intensity),
electrons will not be
ejected unless the
minimum amount of
energy is shone.
Silver metal
• Photoelectrons ejected
when a light with a
frequency of at least
1.14 x 10^15 Hz or
greater is used
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Sodium metal
• Red light
• Violet light
Revised Planck’s Work…
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Einstein piggy-backed off of Planck’s
Theory and we now have…..
Photon
Time to do a little work….
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Tiny water drops in the air disperse
the white light of the sun into a
rainbow. What is the Energy oa a
photon from the violet portion of the
rainbow if it has a frequency of
7.23x10^14 Hz?
E=4.79 x 10^-19 J
• Energy in a photon of violet light
A couple more… 
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A photon has an energy of 2.93 x
10^-25 J. What is its frequency?
What type of electromagnetic
radiation is the photon?
V=4.42 x 10^8 Hz
TV or FM waves
Practice makes perfect… 
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What is the energy of each photon in
the following types of radiation?
• 6.32 x 10^20 Hz
• 9.50 x 10^13 Hz
• 1.05 x 10^16 Hz
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What types of radiation are each?
4.19 x 10^-13 J gamma or x-ray
6.29 x 10^20 J infrared
6.96 x 10^-18 J ultraviolet
How does this work?
(Neon Signs)
What do we know about neon
signs?
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Electricity is passed through tube full of
neon gas
Neon atoms in tube absorb this energy
• What happens when something absorbs
energy?
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Neon atoms in tube become excited
• Stable of Unstable?
• What happens when something is unstable?
• What do we see released energy as?
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Electromagnetic radiation…visible light!!!
EM spectrum
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What happens when we pass sunlight
through a prism?
• Continuous spectrum of colors
• ROYGBIV
What happens when we pass light
from neon gas or hydrogen gas
through prism?
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Separation of colors
Discontinuous spectrum
This is called…
ATOMIC EMISSION SPECTRUM
(AES)
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AES of an element is the set of
frequencies of the electromagnetic
radiation emitted by the atoms of that
element
Individual lines of color
Only certain lines of color appear for
certain elements…
• What does this mean…????
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Every element has a unique AES
Why is this important?
Hydrogen Atom
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Why did scientists want to use
hydrogen?
• How many protons?
• How many electrons?
• Do you think it is easy to use?
• Check out the AES of hydrogen gas…
Neils Bohr (1913)
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Danish Physicist
Worked with
Rutherford
Quantum Model of
Hydrogen atom
• Predicted lines of
Hydrogen AES
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Hydrogen has only one electron but
why do we get different colored lines
on AES???
• We get hydrogen atoms excited…
• Electrons move to excited levels
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H has certain allowable energy
states….
• The lowest energy state is called the
GROUND STATE
Bohr’s Hydrogen Orbits…
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He related H’s energy states to the motion
of an electron in an atom
Single electron in moves around nucleus
in circular orbits
Smaller orbit, smaller radius, closer to
nucleus means…?
• Lower energy level
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Larger orbit, larger radius, farther from
the nucleus means…?
• Higher energy level
Bohr’s Quantum Model
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Assigned quantum numbers, n, to
each orbit
Calculated orbits radius
• Chart on page 127
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1st orbitn=1 (first energy level)
2nd orbitn=2 (second energy level)
3rd orbitn=3 (third energy level)
When we add energy, what
happens to electron?
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Electron excited
Moves to next energy level
Excited=?
• unstable
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What happens when something is
unstable?
• Wants to get back to being stable
• Releases energy
• Goes back down to lower energy level
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Photon is emitted corresponding to the 2
different energy levels associated with the
2 orbits
NEW EQUATION
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/_\ E= E higher e- orbit - E lower e- orbit =E photon=hv
Only certain energies are possible so
only certain frequencies, v, of EM
radiation are emitted
Lets look at the AES of Hydrogen…
•How many lines are there?
•So how many different types of radiations are we seeing?
•There are 4 electron transitions account for lines in the
hydrogen spectrum
•Going from 3rd orbital to 2nd orbital…
•Going from 4th orbital to 2nd orbital…
•Going from 5th orbital to 2nd orbital…
•Going from 6th orbital to 2nd orbital…
Names for these lines…
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Balmer Series
• The 4 visible color lines
• Electrons that drop into n=2
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Other electrons transitions not visible
• Lyman series
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Ultraviolet light
Electrons drop into n=1
• Paschen series
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Infrared
Electrons drop into n=3
Problems with Bohr’s Model
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Predicted AES lines of H but not any
other elements
Did not account for all chemical
behavior
Big problem…
• Electrons don’t move in circular orbits
• Time for a new model…
Louis De Broglie (1924)
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French physics
graduate student
Proposed idea that
accounted for the
fixed energy levels in
Bohr’s model
If waves can have particle like
characteristics, then can
particles, such as electrons, have
wave like characteristics???
What he knew…
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Electrons have wavelike motion
(because it’s a particle)
An electron had restricted orbits
Each orbit had a fixed radius from
the nucleus
Are a wide variety of wavelengths,
frequencies, and energies possible?
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No…there could only be allowed
certain possible frequencies,
wavelengths, and energies in an
atom
De Broglie came up with an equation
for the wavelength of a particle of
mass (m) moving at velocity (v).
De Broglie’s Equation
What does this equation do?
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What are we using?
• Wavelength
• Planck’s constant
• Mass of the particle
• Velocity
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Tells us that all moving particles
have wave-like characteristics
Food for thought…
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Cars?
Baseball?
Do these have
wavelike
characteristics?
Why or why not?
Yes…let’s look at the equation…
λ= h
mv
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The car and the baseball do have a velocity and a
mass…
Using De Broglie’s equation we do get a
wavelength for the movement of a baseball and a
car…
Let’s try the calculation…
Problem time…
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Mass of car= 910 kg
Velocity of car= 25m/s
What is the wavelength of the moving car?
• 2.9 x 10^-38 m
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How big is this?
Can we see or measure this wavelength?
• No, much to small to be detected, even with
the most sophisticated equipment
Another one…
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Electron speed= 25 m/s
Electron mass= 9.11 x 10^-28 g
What is the wavelength of the
moving electron?
• 2.9 x 10^-5 m
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Do you think we can measure this
wavelength and see it?
• Yes, with the right equipment
Practice makes perfect 
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What is the wavelength of an
electron of mass 9.11 x 10-28 kg
traveling at a velocity of 2.00 x 108
m/s? (Planck's constant = 6.63 x 1034 J/Hz.
3.64 x 10-15m.
Werner Heisenberg (1901-1976)
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German
theoretical
physicist
Drew conclusion
from Rutherford,
Bohr, and De
Broglie’s models
Problem with finding the position of
an electron
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Helium balloon in a dark room
How would you determine the
location of this balloon?
Is the balloon going to stay in the
same position?
Energy transfer
What if I gave you a flashlight?
• What happens when we shine a beam of
light on the balloon?
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Photons from light that reflect off of
the balloon reach our eyes and tell
us where the balloon is
Is there a transfer of energy?
• How big is the balloon compared to the
photons?
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Can we do the same thing with
finding the location of an electron in
an atom?
Heisenberg focused on the
interactions between photons and
electrons…
Heisenberg Uncertainty Principle
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It is fundamentally impossible to
know precisely both the velocity and
position of a particle at the same
time
Erwin Schrodinger (1926)
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Austrian physicist
Furthered De
Broglie’s waveparticle theory
Derived equation
that treated
hydrogen’s electron
as a wave
Unlike Bohr’s, his fit
well with atoms of
different elements
Quantum Mechanical Model of the
atom
The Quantum Mechanical Model
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Similar to Bohr’s…
• Limits an electron’s energy to certain
values
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Unlike Bohr’s…
• What did Bohr say about the orbit of an
electron around the nucleus?
• The Quantum Mechanic Model makes no
attempt to describe the electron’s path
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Schrodinger’s wave equation
• Solutions to equation called wave
function
1
Z 3/2 σ
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 
Ψ

e
1s
a
π equation
0
Don’t worry about the
its self…just
know the basics….
• Wave function probability of finding
the electron within a particular volume
of space around the nucleus
• High probability more likely to occur
• Low probability less likely to occur
What the wave function tells us
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The atomic orbital of the
electron
• Atomic orbital 3-D
region around nucleus
Fuzzy Cloud
Density of the cloud at a
given point is proportional to
the probability of finding the
electron at that point
New Word
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Orbital region of space where
there is a 90% probability of finding
an electron of a given energy
“electron cloud”
Orbital
What did Bohr assign to electron
orbitals?
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Quantum numbers
Quantum Mechanical Model does the
same…
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Four Quantum Numbers:
• Specify the “address” (zip code) of each
electron in an atom
First number…Principal Quantum
Number ( n)
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Energy level (associated with the electron)
Size if orbital
• Lowest energy level is assigned principle quantum
number of 1 (n=1)
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Ground state
• What do you think happens as we increase n?
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Orbital becomes larger
Electron spends more time farther away from the
nucleus atom’s energy increases
Principle energy levels
contain…
Energy
Sublevels
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Principle energy level 1 single sublevel
Principle energy level 2 two sublevels
Principle energy level 3 three sublevels
What pattern do you see in the number of
sublevels as we move further away from
the nucleus?
• They increase as n increases (the further we
get from the nucleus)
UPPER LEVEL
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Electron’s are labeled
according to n value
In atom’s with more
than one electron, two
or more electron’s
may have the same n
value
• They are in the same
“electron shell”
Second quantum
number
Angular Momentum Quantum
Number (l)
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Each value of l corresponds to a
different type of orbital with a different
shape
Value of n controls l (subshells
possible)
Angular momentum numbers can equal
0, 1, 2, 3…
l=n-1
• When n=1, l=0 only one possible subshell
• When n=2, l=0,1 two possible subshells
What the number of l means…
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Corresponds to the name of the
subshell
• L=0
• L=1
• L=2
• L=3
subshell
subshell
subshell
subshell
s
p
d
f
S P D F: THE SUBLEVELS
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Each of these 4 sublevels has a unique
shape
Each orbital may contain at most, 2
electrons
LETTERS ORIGINATED FROM
DESCRIPTIONS OF THEIR SPECTRAL
LINES
•
•
•
•
S sharp…spherical
P principal…dumbbell shaped
D diffuse…not all the same shape
F fundamental…not all the same shape
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When principle energy level n=1, then l=0, which means there
is only a single sublevel (one orbital) which is the small,
spherical 1s
When principle energy level n=2, then l can equal 0 or 1, which
means that there are two sublevels (orbitals) 2s and 2p
• 2s sublevel bigger than 1s, still sphere
• 2p sublevel three dumbbell shaped p orbitals of equal energy
called 2px, 2py, and 2pz
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The letters are just there to tell you what axis the electrons go on: x,y, or
z axis
When the principle energy level n=3, then l can equal 0,1, or 2,
which means that there are 3 possible sublevels:
• 3s, sphere, bigger than 1s and 2s
• 3p, dumbbells
• 3d
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Each d sublevel consists 5 orbitals of equal energy
Four d orbitals have same shape but different orientations
Fifth d orbital, 3dz2 is shaped and oriented different from the other four
When the principle energy level n=4, then 1 can equal 0,1,2, or
3 which means l=n-1=4 possible sublevels:
•
•
•
•
•
Seven f orbitals of equal energy ( 2 electrons in each orbital)
4s, sphere
4p, dumbbells
4d,
4f
n
=
# of sublevels per level
n2 =
# of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
Orbitals combine to form a spherical
shape.
2s
2px
2py
2pz
Remember…
1. Principal #
energy level
2. Ang. Mom. #
sublevel (s,p,d,f)
There are two more quantum
numbers (3 and 4) we will
discuss next class
Third Quantum Number
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Ml specifies the orientation of the
orbital in space containing the
electron
Tells us whether the orbital is on the
x, y, or z axis
Fourth Quantum Number
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Ms related to the direction of
the electron spin
Tells us if electron has a
clockwise spin or counter
clockwise spin
Specifies orientation of
electrons spin axis
Recap…
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Bohr?
• Orbits explained hydrogen’s quantized
energy states
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De Broglie?
• Dual particle and wave nature of
electrons
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Schrodinger?
• Wave equation predicted existence of
atomic orbitals containing electrons
Electron Configuration
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Definition: arrangement of electrons in an
atom
Basic rules for filling up orbital's with
electrons
Which is more stable, low energy or high
energy?
• So which orbitals are going to be filled up first?
• We are going to want an arrangement that
gives us the lowest possible energy
Ground state electron configuration
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The most stable, lowest energy
electron arrangement of an atom
Each element has a ground-state
electron configuration
Three Rules for Electron
Arrangement
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Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Aufbau Principle
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Each electron occupies the lowest energy
orbital available
In order to do this, you must learn the
sequence of atomic orbitals from lowest to
highest energy
Aufbau Diagram
• Each box represents an orbital
• Each arrow represents an electron
• Only two arrows per box…
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Only two electrons per orbital
Some important things to
remember about Aufbau…
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All orbitals related to an energy
sublevel are of equal energy
• All three 2p orbitals have the same
energy
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In a multi-electron atom, the energy
sublevels within a principle energy
level have different energies
• All three 2p orbitals are of higher
energy than the one 2s orbital
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In order of increasing energies, the
sequence of energy sublevels within
a principle energy level is s, p, d, f
Orbitals related to energy sublevels
within one principle energy level can
overlap orbitals related to energy
sublevels within another principle
level
• Ex. An orbital related to the atoms 4s
sublevel has a lower energy than the
five orbitals related to 3d sublevel.
Pauli Exclusion Principle
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States that a
maximum on 2
electrons can
occupy a single
atomic orbital but
only if the
electrons have
opposite spins
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Wolfgang Pauli
Austrian Physicist
Observed atoms in
excited states
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Each electron has a spin
Kinda like a spinning top
It can only spin in one of 2 directions
In order for electrons to be together
in an orbital, they must have
opposite spins
Hund’s Rule
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What kind of charge do electrons have?
Do they attract or repel each other?
So……..
Hund’s Rule states that single electrons
with the same spin must occupy all each
energy equal orbital before additional
electrons with opposite spins can occupy
the same orbital
2p orbitals
Read
section 5-3!