Electron configuration

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Atomic Structure andAtomic Structure and
Electron Configuration
HS Ch i t R id L i S i
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HS Chemistry Rapid Learning Series
Wayne Huang, PhD
Kelly Deters, PhD
Russell Dahl, PhD
Elizabeth James, PhD

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Learning Objectives
„Basic structure of atoms.
„How to determine the
By studying this tutorial you will learn…
„How to determine the
number of electrons.
„How to place electrons in
energy levels, subshells
and orbitals.
„How to show electron
configurations using three
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configurations using three
methods.
„How to write and
understand Quantum
Numbers.
Concept Map
Chemistry
Studies
Previous content
New content
Matter
Studies
Atoms
Made of
Electrons
Quantum Numbers
Chemical properties
determined by
Location described by
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Boxes and Arrows
Spectroscopic
Notation
Noble Gas
Notation
3 ways to show configurations

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Atomic Structure
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Definition: Atom
Atom- Smallest piece p
(basic unit) of matter that
has the chemical
properties of the element.
Often called the
Graphical Rendering of an Atom
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Often called the
“Building Block of Matter”.
pg
Protons
Neutrons
Electrons

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What’s in an Atom?
An atom is made of three sub-atomic particles.
Particle Location Mass
1
Charge
Nucleus
Nucleus
Outside the
nucleus
1 amu =
1.67×10
-27
kg
1 amu =
1.67×10
-27
kg
0.00055 amu
9.10×10
-31
kg
+1
0
-1
Proton
Neutron
Electron
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1 amu (“atomic mass unit”) = 1.66 ×10
-27
kg
The Atom
Nucleus
Ch
Electron
Cloud
M
Very small
relative mass
Charge =
-(# of
electrons)
Charge =
# of protons
Mass =
# of protons
+ # of neutrons
Overall Charge =
# of protons
-
#fl t
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# of electrons
Overall Mass =
# of protons
+
# of neutrons

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Protons Versus Electrons
Protons Electrons
+ Charge -Charge
Found in nucleus.
# determines the “identity”
of the atom (atomic
number).
Found outside nucleus.
# and configuration determine how the atom will react.
Contributes to mass of atom. Not contribute significantly to mass of atom.
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Cannotbe lost or gained
without changing which
element it is (nuclear
reaction).
The ratio of protons to electrons determines the charge on the atom (since neutrons are “neutral”).
Canbe lost or gained—
results in an atom with a
charge (ion).
Electron
Locations
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Definition: Electron Cloud
Electron cloud–It is
the area outside of thethe area outside of the
nucleus where the
electrons reside (i.e.
the probability of
finding electrons).
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Electron Clouds
Electron
cloud
Principal
energy levels
Subshells
The electron cloud is made of energy levels (n).
Energy levels are
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Subshells
composed of subshells (l).
Subshells have orbitals (m
l).

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Definition: Subshell and Orbital
Subshell– A set of orbitals with equal
energy.gy
Orbital– Area of probability of an electron
being located.
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Each orbital can hold 2 electrons (spin up and down).
Types of Subshells
Begins in
Number of Total number
There are 4 types of subshells that electrons reside in
under ordinary circumstances.
Subshell
Begins in
energy level
equal energy
orbitals
of electrons
possible
s
p 2
1
3
1
6
2
g
y Increases
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d
f
3
4
5
7
10
14
Ener
g
Subshell Mnemonic: spdf= Smart People Don’t Fail.

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Pictures of Orbitals
1sorbital1sorbital
3 porbitals
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5 dorbitals
Electron
Configuration
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Definition: Electron Configurations
Electron Configurations–
Shows the grouping and gpg
position of electrons in an atom.
Since the number of electrons and their
configuration determines the chemical properties of
the atom, it is important to understand them.
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Box (and Arrow) Notation: Electron
configurations use boxes for orbitals and
arrows for electrons.
Aufbau Principle
Aufbau (building-up) Principle:Electrons must fill
subshells (and orbitals) so that the total energy of 1
The first of 3 rules that govern electron configurations:
() gy
atom is at a minimum.
1
What does this mean?
Electrons must fill the lowest
available subshells and orbitals
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before moving on to the next higher energy subshell/orbital.

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Energy and Subshells
The energy diagram below shows the relative energy
levels.
6p
5d
4f
3s
4s
5s
3p
4p
5p
3d
4d
6s
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1s
2s
2p
Energy
Subshells are filled from the lowest
energy level (1s) to increasing energy
levels (follow the arrows).
Not that this does not always go in
numerical order.
Hund’s Rule
Hund’s Rule:Place electrons in unoccupied
The second of 3 rules that govern electron configurations.
Hund s Rule:Place electrons in unoccupied
orbitals of the same energy level (spin up)
before doubling up.
2
How does this work?
If you need to add 3 electrons to apsubshell
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If you need to add 3 electrons to a psubshell,
add 1 to each (in parallel spins) before
beginning to double up.

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Pauli Exclusion Principle
Pauli Exclusion Principle:Two electrons that
th bit l t h diff t i
3
The last of 3 rules that govern electron configurations.
occupy the same orbital must have different spins.
“Spin” describes the angular
momentum of the electron.
“Spin” is designated with an up
or down arrow.
How does this work?
Spin
Up Spin
Down
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How does this work?
If you need to add 4 electrons to a p
subshell, you’ll need to double up. When
you double up, make them opposite spins.
Down
Determining the Number of Electrons
In order to properly construct an electron configuration,
you must be able to determine how many electrons to
use.
Br
1-
Charge = -1
Charge = # of protons – # of electrons
Atomic number = # of protons
Example: How many electrons does the following have?
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-1 = 35 - Electrons
Atomic number for Br = 35 = # of protons
Electrons = 36 35Br
1-

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Another Example
In order to properly construct an electron configuration,
you must be able to determine how many electrons to
use.
No charge written ÆCharge is 0Cl
Charge = # of protons – # of electrons
Atomic number = # of protons
Example: How many electrons does the following have?
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Electrons = 17
0 = 17 - Electrons
Atomic number for Cl = 17 = # of protons
17Cl
Applying the Rules
Aufbau Principle:Electrons must fill subshells (and orbitals)
so that the total energy of atom is at a minimum.
1
Use the 3 rules of electron configurations.
Hund’sRule:Place electrons in unoccupied orbitals of the
Example: Give the electron configuration for a Cl atom.
No charge written ÆCharge is 0
17Cl
Atomic number for Cl = 17 = # of protons
P
auli Exclusion Principle:Two electrons that occupy the same
orbital must have different spins.
3
Hund s Rule:Place electrons in unoccupied orbitals of the
same energy level before doubling up.
2
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0 = 17 - Electrons Electrons = 17
Place 17 electrons
1s 2s 2p 3s 3p
42 3 1 5 6 7910 11 12 13 14 15 16 17 8
Electron Configuration Rules Mnemonic: Aufbau (stays low); Hund (does not
double up); Pauli (spins up and down) = “Alligator stays low; Hippo does not pair
up and Penguin jumps up and down.”

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Spectroscopic
Notation
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Definition: Spectroscopic Notation
Spectroscopic Notation– Shorthand way of
showing electron configurations.gg
The number of electrons in a subshell are shown
as a superscript after the subshell designation.
Box
26/56
1s 2s 2p 3s 3p
1s
2
2s
2
2p
6
3s
2
3p
5
Spectroscopic
Notation
Box
Notation
1S
2
Principal Energy Level (n)
Subshell (l)
Number of Electrons

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Writing Spectroscopic Notation
Determine the number of electrons to place.1
Follow Aufbau’s Principle for filling order.2
Fill in subshells until they reach their max (s= 2, p= 6, d= 10,
f= 14).
3
The total of all the superscripts is equal to the number of
electrons.
4
Example: Give the spectroscopic notation for S.
No charge writtenÆCharge is 0
S
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0 = 16 - Electrons
No charge written ÆCharge is 0
16S Atomic number for S = 16 = # of protons
Electrons = 16
Place 16 electrons
1s2s2p3s3p
22 624
22624++++ =16
Electron Configurations
and the Periodic Table
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1
2
2
2
2
5
Configurations Within a Group
Look at the electron configurations for the Halogens
(Group 7).
F 1s
2
2s
2
2p
5
F
Cl
Br
1s
2
2s
2
2p
6
3s
2
3p
5
1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
10
4p
5
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I 1s
2
2s
2
2p
6
3s
2
3p
6
4s
2
3d
10
4p
6
5s
2
4d
10
5p
5
All of the elements in Group 7 end with 5electrons in
a psubshell.
Configurations and Periodic Table
In fact, every Group ends with the same number of
electrons in the highest energy subshell.
Each area of the periodic table is referred to by the
hih t bhlltht ti l t
s
1
s
2
p
1
p
2
p
3
p
4
p
5
p
6
highest energy subshell that contains electrons.
d-block
p-blocks-block
IA
IIA
IIIB IVBVBVIBVIIBVIIIBVIIIBVIIIB IB IIB
IIIA IVA VA VIA VIIA
VIIIA
Group Æ
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d
1
d
2
d
3
d
4
d
5
d
6
d
7
d
8
d
9
d
10
p
1
p
2
p
3
p
4
p
5
p
6
f
1
f
2
f
3
f
4
f
5
f
6
f
7
f
8
f
9
f
10
f
11
f
12
f
13
f
14
f-block

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Wondering how to remember the order of filling of
the subshells?
Just use the periodic table as a mnemonic device.
Periodic Table as a Road-Map - 1
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In order to do this, the “f” block needs to be placed in atomic
order.
(It’s usually written below to fit it on the paper).
To see the filling order of subshells, read from left to right, top to bottom!
Periodic Table as a Road-Map - 2
1s 1s
This tool shows that the 3denergy level is filled after the 4senergy level!
2p
3p
4p
5p
6p
3d
4d
5d
6d
4f
5f
1s
2s
3s
4s
5s
6s
7s
1s
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psubshells begin in level 2, so begin the p-block with “2p”.
ssubshells begin in level 1, so begin the s-block with “1s”.
dsubshells begin in level 3, so begin the d-block with “3d”.
fsubshells begin in level 4, so begin the f-block with “4f”.

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Another Tool for Filling Order
1s
There is another tool (mnemonic device) commonly
used to remember orbital filling order.
2s2p
3s3p 3d
4s4p 4d 4f
5s5p 5d 5f
Building-Up Principle:
To read the chart, start
with 1sand follow the
arrows. Move down one
diagonal as far as
possible, then jump to
the top of the next
di l d k
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5s5p 5d 5f
6s6p 6d
7s7p
8s
diagonal and keep
going.
ElectronElectron
Configurations
of Ions
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Definition: Ion
Ion– an atom (or group of
atoms) that has gained or )g
lost electrons resulting in a net charge.
Atoms gain and lose electrons to be in a more stable state.
35/56
Usually, the “more stable state” is a full valence shell.
Outermost shell of electrons
Look at the electron configurations for the following
(#p= # of protons and #e= # of electrons):
Full Valence Shell Ions
1s2s2p
22 6
Br
-
O
2-
1s2s2p3s3p
22 62 6
4s
2
3d
10
4p
6
#p = 35 -1 = 35 - #e #e = 36
#p = 8 -2 = 8 - #e #e = 10
Charge
= p-e
36/56
Na
+
Ca
2+
1s2s2p3s3p
22 62 6
#p = 11 +1 = 11 - #e #e = 10
#p = 20 +2 = 20 - e #e = 18
1s2s2p
22 6

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What do you notice about each of these
configurations?
Full Valence Shell Ions
They all end with full psubshells.
Notice that O
2-
and Na
+
have the
same number and configuration of
1s2s2p22 6
Br
-
O
2-
1s2s2p3s3p
22 62 6
4s
2
3d
10
4p
6
37/56
electrons.
Na
+
Ca
2+1s2s2p3s3p
22 6 2 6
1s2s2p
22 6
This makes them isoelectric.
Noble Gas
Configuration
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Definition: Noble Gas Notation
Noble Gas– Group 8 of the Periodic
Table. They contain full valence shells.
Noble Gas Notation– Noble gas is used
to represent the core (inner) electrons
and only the valence shell is shown.
35Br
39/56
1s2s2p3s3p
22 62 6
4s
2
3d
10
4p
5
4s
2
3d
10
4p
5
[Ar]
35Br
Spectroscopic Notation:
Noble Gas Notation:
The “[Ar]” represents the core electrons and only the valence electrons are shown.
How do you know which noble gas to use to
symbolize the core electrons?
Which Noble Gas Do You Choose?
Think: Price is Right.
Hd i thPiiRiht?How do you win on the Price is Right?
By getting as close as possible without going over.
Choose the noble gas that’s closest without going over!
Noble Gas # of electrons
He 2
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Ne
Ar
Kr
Xe
10
18
36
54

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How do you know where to start off after using a
noble gas?
Use the periodic table!
Where Does the Noble Gas Leave Off?
2p
3p
4p
5p
6p
3d
4d
5d
6d
4f
5f
1s
2s
3s
4s
5s
6s
7s
He
Ne
Ar
Kr
Xe
Rn
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6d5f7s
The noble gas fills the subshell that it’s at the end of.
Begin filling with the “s” subshell in the next row to show
valence electrons.
Noble Gas Notation Example
Determine the number of electrons to place.1
Determine which noble gas to use.2
Start where the noble gas left off and write
spectroscopic notation for the valence electrons.
3
Example: Give the noble gas notation for As.
No charge written ÆCharge = 0
A
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+
0 = 33 - electrons
33
As Atomic number for As = 33 = # of protons.
Electrons = 33 Place 33electrons.
[Ar] 4s3d4p
210 3
18 2 10 3++ =33
Closest noble gas: Ar (18)
Ar (1s
2
2s
2
2p
6
3s
2
3p
6
) is full up through 3p.

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Comparing the
Different Notations
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Pros and Cons of Each Notation
Each notation has it’s advantages and disadvantages.
Pro Con
Shows if electrons
are paired or
unpaired.
Quicker than “Boxes
and arrows”.
Longest method.
Does not show
pairing of electrons.
“Boxes and
arrows”
Spectroscopic
Notation
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Allows focus on the
valence electrons
(that control
bonding).
Quickest method.
Does not show core
electrons.
Does not show
pairing of electrons.
Noble Gas
Notation

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Exceptions to the
Aufbau Rule
45/56
Stability of d Subshells with 5 or 10
dsubshells have 5 orbitals…
They can hold 10 electrons.
According to the Aufbau principle, Cr should have the
following valence electron configuration:
4s
2
3d
4
But a half-fullor completely full dsubshell is more stable
46/56
than the above configuration, so it is:
4s
1
3d
5

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Elements with Exceptions
The following elements are excepts to the Aufbau
Principle:
Element Should be Actually is
4s
2
3d
4
5s
2
4d
4
6s
2
5d
4
4s
2
3d
9
5s
2
4d
9
4s
1
3d
5
5s
1
4d
5
6s
1
5d
5
4s
1
3d
10
5s
1
4d
10
Cr
Mo
W
Cu
Ag
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g
6s
2
5d
9
6s
1
5d
10
Au
They are the two groups on the periodic table that begin with Cr and Cu.
Quantum
Numbers
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Definition: Quantum Numbers
Quantum Numbers– A set of 4
numbers (n, l, m
l, & m
s) that
describes the electron’s
placement in the atom.
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4 Quantum Numbers
2, 1, -1, +½
n m
l
2p
10+1
n = 2
or
m
s= ½(up)
Quantum
Number
Symbol
n
Describes
Shell Number
(Size)
Sbhll
Possible
Numbers
Whole # ≥1Principal
Ai thl
l m
s
-1 0 +1
l= 1
m
l= -1
50/56
l
m
l
m
s
Subshell
Type (Shape)
Whole # < n
(0 Æn-1)
-lÆ+l
+½ or –½
Azimuthal
(Angular)
Magnetic
Spin
Orbital
(Orientation)
Spin (up or down spin)

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Determining Quantum Numbers
n: principal energy level
l: subshell s=0
Give the number of the shell.
4p
3
l: subshell s= 0
p= 1
d= 2
f= 3
m
l: orbital
0s -101p
-2-1012d
-3-2-10123f
Coding system: 0,1… n-1.
Number-line system of identifying orbitals.
0 is always in the middle.
Number line fromlto +l
51/56
Number line from –lto + l.
m
s: spin
Coding system
↑= + ½ (spin up)
↓= - ½ (spin down)
Quantum Number Examples
Give the quantum numbers for the red arrow.Example:
1s 2s 2p 3s 3p
It’s in level“3”
0
It s in level 3.
___, ___, ___, ___3
It’s in subshell “s” - the “code” for “s” is “0”.
0
It’s in orbital “0”.
0It’s a downarrow. -½
Give the quantum numbers for the red arrow.Example:
52/56
1s 2s 2p 3s 3p
It’s in level “2”.
___, ___, ___, ___2
It’s in subshell “p”—the “code” for “p” is “1”.
1
It’s in orbital “-1”.
-1
It’s an uparrow.
+½
-1 0 +1

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Identifying Incorrect Quantum Numbers
Example: What’s wrong with the following sets of quantum numbers?
1, 1, 0, +½
n= 1…OK as n(energy level) can be any whole # > 0
l = 1…subshell is “p”, but if n= 1 so lmust be 0 (i.e. s subshell).
There is nopsubshell in energy level 1
;
2, 1, -2, -½
There is no psubshell in energy level 1.
n = 2…OK as ncan be any whole # >0
l= 1…subshell is “p”.
OK as level 2 has “p”, i.e. “2p”.
m
l= -2…on the “-2” orbital
“p” subshell has 3 orbitals: ___ ___ ___
-1 0 +1
No “-2” orbital in a “p” subshell.
m
lmust be between–landl(i.e.-1, 0,+1),not-2.
;
53/56
1, 0, 0, -1
m
lmust be between land l(i.e. 1, 0, 1), not 2.
n= 1…OK as ncan be any whole # >0
l= 0…subshell is “s”.
OK as level 1 has an “s”.
m
l= 0…on the “0” orbital
OK as “s” has 1 orbital and it’s “0”.
m
s= -1
m
smust be either +½ or -½, not -1.
;
Electron
configurationscan
be shown with
Electron
configurationscan
be shown with
Atoms are made of
protons, neutrons
and electrons. The
configuration of the
Atoms are made of
protons, neutrons
and electrons. The
configuration of the
Quantum numbers
describe the
location of an
Quantum numbers
describe the
location of an
Learning Summary
ff
boxes and arrows,
in spectroscopic
notation, or noble
gas notation.
boxes and arrows,
in spectroscopic
notation, or noble
gas notation.
g
electrons
determines the
chemical properties
of the atom.
g
electrons
determines the
chemical properties
of the atom.
ocat o o a
electron in an atom
and are a series of
4 numbers.
ocat o o a
electron in an atom
and are a series of
4 numbers.
54/56
Electron configurations
are written following the
Aufbau principle, Hund’s
Ruleand the Pauli
Exclusion Principle.
Electron configurations
are written following the
Aufbau principle, Hund’s
Ruleand the Pauli
Exclusion Principle.
Electronsare organized in
levels, subshells and
orbitals.
Electronsare organized in
levels, subshells and
orbitals.

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You have successfully completed
the core tutorial
Atomic Structure and
Electron ConfigurationElectron Configuration
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What’s Next …
Step 1: Concepts – Core Tutorial (Just Completed)
ÆStep 2: Practice – Interactive Problem Drill
Step 3: Recap – Super Review Cheat Sheet
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