Download Electron Configurations

Atomic Structure
From Indivisible to Quantum
Mechanical Model of the Atom
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Quantum mechanical
model(Modern Atomic Theory)
SchrÖdinger
Heisenberg
Pauli
Hund
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Heisenberg’s Uncertainty
Principle
Impossible to determine both the
position and the velocity of an e- in an
atom simultaneously with great
certainty.
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SchrÖdinger
e- not in neat orbits, but exist in regions
called orbitals
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Definitions
Orbital  region in space where the
probability of finding an electron is the
highest
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Quantum Numbers
Definition: specify the properties of
atomic orbitals and the properties of
electrons in orbitals
There are four quantum numbers
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Quantum Numbers (1)
Principal Quantum Number, n
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Quantum Numbers
Principal Quantum Number, n



Values of n = 1,2,3,… 
Positive integers only!
Indicates the main energy level occupied
by the electron
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Quantum Numbers
Principal Quantum Number, n


Values of n = 1,2,3,… 
Describes the energy level, orbital size
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Quantum Numbers
Principal Quantum Number, n



Values of n = 1,2,3,… 
Describes the energy level, orbital size
As n increases, orbital size increases.
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Principle Quantum Number
More than one e- can have the same n
value
These e- are said to be in the same eshell
The total number of orbitals that exist
in a given shell = n2
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Orbital Shapes
For a specific main energy level, the
number of sublevels possible is equal to
n.

Ex. n=2, can have two sublevels.
A sublevel is assigned a letter:
s , p , d, f , g, h
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Energy Level and Orbitals
n=1,
n=2,
n=3,
n=4,
only s orbitals
s and p orbitals
s, p, and d orbitals
s,p,d and f orbitals
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Atomic Orbitals
Atomic Orbitals are designated by the
principal quantum number followed by
letter of their subshell


Ex. 1s = s orbital in 1st main energy level
Ex. 4d = d sublevel in 4th main energy level
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The area where an electron can be found,
the orbital,
is defined mathematically,
but we can see it as a specific shape
in 3-dimensional space…
Orbital Shapes
s is spherical.
One possible orientation.
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z
y
x
z
y
The 3 axes represent
3-dimensional space
x
z
y
For this presentation, the
nucleus of the atom is at
the center of the three axes.
x
The 1s orbital is a
sphere, centered
around the nucleus
The 2s orbital is also
a sphere.
The 2s electrons have a
higher energy than the 1s
electrons. Therefore, the 2s
electrons are generally more
distant from the nucleus,
making the 2s orbital larger
than the 1s orbital.
1s orbital
2s orbital
Orbital Shapes
p orbital.
“dumbbell” shape
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There are three
p orbitals
3 possible
orientations
The three 2p orbitals
are oriented
perpendicular
to each other
DEGENERATE
ORBITALS
All three orbitals are
identical of each
other by energy, size
and shape.The only
difference is their
orientation in space.
z
This is
one 2p orbital
(2py)
y
x
z
another 2p orbital
(2px)
y
x
z
the third 2p orbital
(2pz)
y
x
z
The three 2p orbitals,
2px, 2py, 2pz
y
x
3p, 4p, 5p, etc…
have the same
shape and
number, just
larger
Orbital Shapes
d orbital.
“double dumbbell” or four-leaf clover
It has 5 degenerate orbitals
5 possible orientations
The 4d orbitals etc…are the same shape, only larger
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Orbital Shapes
f orbital
It has 7 degenerate orbitals
7 possible orientations
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Energy Level and Orbitals
n=1,
n=2,
n=3,
n=4,
only s sublevel
s and p sublevels
s, p, and d sublevels
s,p,d and f sublevels
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In the same energy level,
energies of orbitals:
s<p<d<f
(because of the amount of repulsion
between electrons)
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Quantum Numbers (4)
Electron Spin Quantum Number,
ms = +1/2, 1/2)
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Electron Spin QN
1. Relates to the spin states of
the electrons.
2. Electrons are –1 charged and
are spinning
3. The two possible spin directions are called +½
and –½
Pauli Exclusion Principle
Wolfgang
Pauli
No 2 e- in an atom can have the
same set of four quantum
numbers (n, l, ml, ms ).
Therefore, no atomic orbital can
contain more than 2 e-.
and they must have opposite
spin.

Like This
Sublevels
There are 4 sublevels(different shaped orbitals)




s (has 1 orbital)
p (has 3 orbitals)
d (has 5 orbitals)
f (has 7 orbitals)
Can hold 2 eCan hold 6 e-
Can hold 10 e-
Can hold 14 e-
Each orbital can hold 2 electrons
Energy
Level
(n)
Sublevels
in Level
1
s
1
1
2
s
1
p
3
4
s
1
p
3
d
5
s
1
p
3
d
5
f
7
3
4
# Orbitals Total # of
in Sublevel Orbitals in
Level
9
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Electron Configurations
Electron Configurations: arrangement of
e- in an atom
There is a distinct electron configuration
for each atom
There are 3 rules for writing electron
configurations:
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Aufbau Principle
Aufbau Principle: an e- occupies the
lowest energy orbital that can receive it.
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Aufbau order:
3d
E 4s
3p
N
E 3s
R
2p
G 2s
Y
1s
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Writing Electron Configurations
Describes
e-
location.
4
3p
Principal
Energy Level
Sublevel
# of e-
Electron Configuration
The total of the superscripts must equal
the atomic number (number of
electrons) of that atom.
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Orbital Diagrams
These diagrams are based on the
electron configuration.
In orbital diagrams:


Each orbital (the space in an atom that will
hold a pair of electrons) is shown.
The opposite spins of the electron pair is
indicated.
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Orbital Diagram Rules
1. Represent each electron by an arrow
2. The direction of the arrow represents the
electron spin
3. Draw an up arrow to show the first electron
in each orbital.
4. Hund’s Rule(the principle of multiplicity):
Distribute the electrons among the orbitals
within sublevels so as to give the most
unshared pairs.
Put one electron in each orbital of a sublevel
before the second electron appears.
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Hund’s Rule
One electron enters each orbital of equal
energy (degenerate orbitals)until all the orbitals
contain one electron with the same spin
direction…
…then they pair up.
p
orbitals



p
orbitals



Like This
Like This


configuration
1s
2s
H
1s1
↑
He
1s2
↑↓
Li
1s22s1
↑↓
↑
Be
1s22s2
↑↓
↑↓
B
1s22s22p1
↑↓
↑↓
↑
C
1s22s22p2
↑↓
↑↓
↑
↑
N
1s22s22p3
↑↓
↑↓
↑
↑
↑
O
1s22s22p4
↑↓
↑↓
↑↓
↑
↑
F
1s22s22p5
↑↓
↑↓
↑↓
↑↓
↑
Ne
1s22s22p6
↑↓
↑↓
↑↓
↑↓
↑↓
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2px
2py
2pz
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Orbital Diagram Examples
H _
1s
Li  _
1s
B
2s
   __ __
1s
2s
2p
N     _
1s
2s
2p
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Orbital filling table

We can use the previous Noble
Gas as an abbreviation to
indicate filled inner orbitals
a. Na = 1s22s22p63s1 or [Ne]3s1
b. Ca = [Ar]4s2
c. Cl = [Ne]3s23p5
d. Rb = [Kr]5s1
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Dot Diagram of Valence
Electrons
When two atom collide, and a reaction takes
place, only the outer electrons interact.
These outer electrons are referred to as the
valence electrons. Valence electrons are
available to be lost, gained, or shared in the
formation of chemical compounds
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Lewis Dot(electron dot) diagrams
A way of keeping track
of valence electrons.
Write the symbol.
Put one dot for each
valence electron
Start at 3 o’clock move
in a counterclockwise
direction Video
X
Distribute one valence electron at a
time

Do not pair (double up) any electrons until
there is one electron in each of the four
directions
Pair up electrons once there is one in
each of the four directions
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The Lewis Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons.
 First we write the symbol.
Then add 1 electron at a
time to each side.
Until they are forced to pair up.

N