Transcript Document

Electrons in Atoms
1
Dalton’s Atomic Theory
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John Dalton (1766-1844) had four theories
1.
All elements are composed of submicroscopic indivisible particles
called atoms
Atoms of the same element are identical. The atoms of anyone
element are different from those of any other element
Atoms of different elements can physically mix together or can
chemically combine w/ one another in simple whole-number
ratios to form compounds
Chemical reactions occur when atoms are separated, joined, or
rearranged. However, atoms of one element are never changed
into atoms of another elements as a result of a chemical reaction
2.
3.
4.
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Thomson’s Atomic Model
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Thomson’s Atomic Model
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Thomson though electrons were like plums
embedded in a positively charged “pudding”, so
his model was called the “plum pudding” model
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Thomson’s Theory
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1.
2.
3.
Thomson stated: The atom had negatively charged
electrons stuck into a lump of positively charged
protons.
Thomson never explained
Number of protons and neutrons
The arrangement of the particles in the atom
The ease with which atoms are stripped of electrons
to form ions
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Rutherford Model
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1.
2.
3.
Rutherford used the Gold Foil Experiment
Rutherford proposed the following:
Thomson model was incorrect
Most of the mass of the atom and all of its positive
charge reside in a very small, extremely dense region,
which he called the nucleus
Most of the total volume of the atom is empty space
in which electrons move around the nucleus
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Rutherford’s Model
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Discovered dense
positive piece at the
center of the atom
Nucleus
Electrons moved
around
Mostly empty space
Bohr Model
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1.
2.
3.
4.
Bohr changed the Rutherford model and
explained how the electrons travel.
Bohr explained the following in his model:
Electrons travel in definite orbits around the
nucleus
Electrons are arranged in concentric circular paths
or orbitals around the nucleus
Electrons don’t fall into the nucleus because
electrons in particular path have fixed energy and
don’t lose energy
His model was patterned after the motion of the
planets around the sun. It is often called the
Planetary model.
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Bohr Model Cont.
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Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Quantum Theory
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1.
2.
3.
Bohr explained how electrons were moving via
Quantum Theory
Key Terms:
Energy Levels- Regions around the nucleus where the
electron is likely moving
Quantum- Amount of energy required to move an
electron from one energy level to the next
Quantum Leap- Abrupt Change
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Bohr’s Model
Increasing energy
Fifth
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Fourth
Third
Second
First
Nucleus
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Further away from
the nucleus means
more energy.
There is no “in
between” energy
Energy Levels
Bohr’s Model cont.
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Energy levels are not equally spaced.
Energy levels more closely spaced further from
the nucleus
Higher energy level occupied by an electron, the
more energetic that electron is.
Amount of energy gained or lost by an electron
is not always the same amount.
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Bohr Model Cont.
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1.
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The Bohr Model did not account for:
Emission spectra of atoms containing more
than one electron.
So comes along the next model:
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The Quantum Mechanical
Model
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move
from one energy level to another.
Since the energy of an atom is never “in between”
there must be a quantum leap in energy.
Schrodinger derived an equation that described the
energy and position of the electrons in an atom
The Quantum Mechanical
Model
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Things that are very small behave
differently from things big enough
to see.
The quantum mechanical model is a
mathematical solution
It is not like anything you can see.
Schrödinger’s Equation
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The wave function is a F(x, y, z)
Actually F(r,θ,φ)
Solutions to the equation are called orbitals.
These are not Bohr orbits.
Each solution is tied to a certain energy
Animation
These are the energy levels
Schrödinger’s Equation
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The Quantum Mechanical
Model
Has energy levels for
electrons.
Orbits are not circular.
It can only tell us the
probability of
finding
an electron a certain
distance from the nucleus.
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The Quantum Mechanical
Model
The atom is found inside
a blurry “electron cloud”
A area where there is a
chance of finding an
electron.
Draw a line at 90 %
Atomic Orbitals
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There are the region of space which there is a high
probability of finding an electron
Within each energy level the complex math of
Schrodinger’s equation describes several shapes.
These are called atomic orbitals
Quantum Numbers- numbers that specify the
properties of atomic orbitals and their electrons
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Quantum Numbers
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1.
2.
3.
4.
There are 4 types of Quantum Numbers
Principal – distance from the nucleus
Angular Momentum- Orbital Shape
Magnetic- Orbital position with respect to the
X, Y, & Z axes.
Spin- Has only two values (+1/2 or –1/2) and
is needed to specify 1 of 2 positional
orientations of an electron
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Principal Quantum Number
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Symbolized by the letter N, indicates the main
energy levels surrounding the nucleus
There are 7 principal quantum numbers
A.K.A. – Shells
Value of N is a whole number ex. 1,2,3 ect..
Main Energy Level – N=1; closest to the nucleus or
ground state
Ground State- state of the lowest energy of the
atom.
As N increases, the distance from the nucleus
increases and the energy increases
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Angular Momentum Quantum
Number
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Indicates the shape of the orbital.
Within each main energy level beyond the first,
orbitals with different shapes occupy different
regions
A.K.A. – Sublevels or Subshells
The number of sublevels = Value of the
Principal Quantum Number
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Magnetic Quantum Number
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I.
II.
III.
IV.
Indicates the orientation of a orbital about
the nucleus
There are 4 types of orbital orientation
S Orbital
P Orbital
D Orbital
F Orbital
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Spin Quantum Number
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Has only two possible values: +1/2 or –1/2.
These values indicate two possible states of an
electron in an orbital
Spin Quantum # is significant because each
single orbital can hold no more than two
electrons, which must have opposite spin.
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S orbitals
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1 s orbital for
energy
Spherical
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals.
every
level
shaped
P orbitals
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Start at the second energy level
3 different directions
3 different shapes
Each can hold 2 electrons
P Orbitals
D orbitals
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Start at the third energy level
5 different shapes
Each can hold 2 electrons
F orbitals
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Start at the fourth energy level
Have seven different shapes
2 electrons per shape
F orbitals
Summary
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
By Energy Level
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First Energy Level
only s orbital
only 2 electrons
1s2
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Second Energy Level
s and p orbitals are
available
2 in s, 6 in p
2s22p6
8 total electrons
By Energy Level
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Third energy level
s, p, and d orbitals
2 in s, 6 in p, and 10 in
d
3s23p63d10
18 total electrons
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Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d,
ahd 14 in f
4s24p64d104f14
32 total electrons
By Energy Level
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Any more than the
fourth and not all the
orbitals will fill up.
You simply run out of
electrons
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The orbitals do not fill
up in a neat order.
The energy levels
overlap
Lowest energy fill first.
Question for You
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How many principal quantum numbers are
there?
What is the maximum number of electrons that
can fill the 3rd energy level?
How many orbitals are in the sublevel F?
What is the total number of orbitals for the 3rd
main energy level?
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Electron Configuration
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1)
2)
The way electrons are arranged in atoms
There are three rules which help dictate how
electrons are arranged in the atoms.
Aufbau Principle- electrons occupy the orbitals of
the lowest energy first
Hund’s Rule- Orbitals of equal energy are each
occupied by one electron before any one orbital is
occupied by a second electron. All electrons in a
single occupied orbital must have the same spin.
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Electron Configuration cont.
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Pauli Exclusion Principle- No two electrons may
occupy any given orbital without having
opposite spin. No two electrons in the same
atom can have the same set of four quantum
numbers.
Let’s determine electron configuration.
Let’s start with Phosphorus.
Need to account for all 15 electrons
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Electron Configurations
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Distribution of all electrons
in an atom
Consist of
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Number denoting the energy
level
Electron Configurations
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Distribution of all electrons
in an atom
Consist of
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Number denoting the energy
level
Letter denoting the type of
orbital
Electron Configurations
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Distribution of all electrons
in an atom.
Consist of
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Number denoting the energy
level.
Letter denoting the type of
orbital.
Superscript denoting the
number of electrons in those
orbitals.
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Orbital Diagrams
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Each box represents one
orbital.
Half-arrows represent the
electrons.
The direction of the arrow
represents the spin of the
electron.
Hund’s Rule
“For degenerate orbitals,
the lowest energy is
attained when the
number of electrons with
the same spin is
maximized.”
Exceptions to Electron
Configuration
Orbitals fill in order
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Lowest energy to higher energy.
Adding electrons can change the energy of the
orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order
Write these electron
configurations
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Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons 1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!
Chromium is actually
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1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
Copper’s electron configuration
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Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled
orbital.
Remember these exceptions
Shortcuts for Electron
Configuration
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There are two short handed methods of writing the
electron configuration.
The 1st method is called the outer-level configuration.
That tells you the outer-most configuration for that
element.
The 2nd method is called the Noble Gas Notation. This
tells you the complete notation using Noble Gases.
Let’s start with outer-level notation!!!
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H
Li
1
3
Na
11
K
19
Rb
37
Cs
55
Fr
87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f14
5d106p67s1
1s2 He 2
Ne
2
2
6
1s 2s 2p
10
1s22s22p63s23p6 Ar18
1s22s22p63s23p64s23d104p6 Kr
36
1s22s22p63s23p64s23d104p65s24d105p6 Xe
54
1s22s22p63s23p64s23d104p65s24d10 Rn
5p66s24f145d106p6 86
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S- block
s1
s2
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Alkali metals all end in s1
Alkaline earth metals all end in s2
really have to include He but it fits better
later.
He has the properties of the noble gases.
The P-block
p1 p2
p3
p4
p5
p6
Transition Metals -d block
d1 d2 d3
s1
d5
s1
d5 d6 d7 d8 d10 d10
F - block
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inner transition elements
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1
2
3
4
5
6
7
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Each row (or period) is the energy level for s
and p orbitals.
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D orbitals fill up after previous energy level so first
d is 3d even though it’s in row 4.
1
2
3
4
5
6
7
3d
1
2
3
4
5
6
7
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4f
f orbitals start filling at 4f
5f
Summary Outer-Level
Configuration
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Writing Electron
configurations the
easy way
Yes there is a shorthand
Electron Configurations repeat
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The shape of the periodic table is a
representation of this repetition.
When we get to the end of the column the
outermost energy level is full.
This is the basis for our shorthand.
The Shorthand
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Write the symbol of the noble gas before the
element.
Then the rest of the electrons.
Aluminum - full configuration.
1s22s22p63s23p1
Ne is 1s22s22p6
so Al is [Ne] 3s23p1
More examples
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Ge = 1s22s22p63s23p64s23d104p2
Ge = [Ar] 4s23d104p2
Hf=1s22s22p63s23p64s23d104p65s2
4d105p66s24f145d2
Hf=[Xe]6s24f145d2
The Shorthand Again
Sn- 50 electrons
The noble gas
before it is Kr
Takes care of 36
Next 5s2
Then 4d10
Finally 5p2
[ Kr ] 5s2 4d10 5p2
Quantum Numbers
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Solving the wave equation gives a set of wave
functions, or orbitals, and their
corresponding energies.
Each orbital describes a spatial distribution of
electron density.
An orbital is described by a set of three
quantum numbers.
The Wave-like Electron
The electron propagates
through space as an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
The Quantum Mechanical Model
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A totally new approach
De Broglie said matter could be like a wave.
De Broglie said they were like standing waves.
The vibrations of a stringed instrument
What’s possible?
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You can only have a standing wave if you have
complete waves.
There are only certain allowed waves.
In the atom there are certain allowed waves
called electrons.
1925 Erwin Schroedinger described the wave
function of the electron
Much math, but what is important are the
solutions
Schrödinger’s Equation
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The wave function is a F(x, y, z)
Actually F(r,θ,φ)
Solutions to the equation are called orbitals.
These are not Bohr orbits.
Each solution is tied to a certain energy
Animation
These are the energy levels
What does the wave Function
mean?
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nothing.
it is not possible to visually map it.
The square of the function is the probability
of finding an electron near a particular spot.
best way to visualize it is by mapping the
places where the electron is likely to be
found.
Probability
Distance from nucleus
Distance from nucleus
Sum of all Probabilities
Defining the size
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The nodal surface.
The size that encloses 90% to the total electron
probability.
NOT at a certain distance, but a most likely
distance.
For the first solution it is a a sphere.
Quantum Mechanics
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Erwin Schrödinger
developed a mathematical
treatment into which both
the wave and particle nature
of matter could be
incorporated.
It is known as quantum
mechanics.
© 2009, Prentice-Hall, Inc.
Quantum Mechanics
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The wave equation is designated
with a lower case Greek psi ().
The square of the wave
equation, 2, gives a probability
density map of where an
electron has a certain statistical
likelihood of being at any given
instant in time.
© 2009, Prentice-Hall, Inc.
Quantum Numbers
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Solving the wave equation gives a set of wave
functions, or orbitals, and their corresponding
energies.
Each orbital describes a spatial distribution of
electron density.
An orbital is described by a set of three quantum
numbers.
© 2009, Prentice-Hall, Inc.
Quantum Numbers
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There are many solutions to Schrödinger’s
equation
Each solution can be described with quantum
numbers that describe some aspect of the
solution.
Principal quantum number (n) size and energy
of an orbital
Has integer values >0
s Orbitals
Observing a graph of
probabilities of finding an
electron versus distance
from the nucleus, we see
that s orbitals possess n−1
nodes, or regions where
there is 0 probability of
finding an electron.
© 2009, Prentice-Hall,
Inc.
p Orbitals
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The value of l for p orbitals is 1.
They have two lobes with a node between them.
© 2009, Prentice-Hall, Inc.
d Orbitals
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The value of l for a d
orbital is 2.
Four of the five d
orbitals have 4 lobes;
the other resembles a
p orbital with a
doughnut around the
center.
© 2009, Prentice-Hall, Inc.
Quantum numbers
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Angular momentum quantum number l
shape of the orbital
integer values from 0 to n-1
l = 0 is called s
l = 1 is called p
l =2 is called d
l =3 is called f
l =4 is called g
Magnetic Quantum Number (ml)
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The magnetic quantum number describes the
three-dimensional orientation of the orbital.
Allowed values of ml are integers ranging
from -l to l:
−l ≤ ml ≤ l.
Therefore, on any given energy level, there
can be up to 1 s orbital, 3 p orbitals, 5 d
orbitals, 7 f orbitals, etc.
© 2009, Prentice-Hall, Inc.
Magnetic Quantum Number (ml)
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Orbitals with the same value of n form a shell.
Different orbital types within a shell are subshells.
© 2009, Prentice-Hall, Inc.
Spin Quantum Number, ms
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This led to a fourth
quantum number, the spin
quantum number, ms.
The spin quantum number
has only 2 allowed values:
+1/2 and −1/2.
© 2009, Prentice-Hall, Inc.
Pauli Exclusion Principle
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No two electrons in the
same atom can have exactly
the same energy.
Therefore, no two electrons
in the same atom can have
identical sets of quantum
numbers.
© 2009, Prentice-Hall, Inc.