Rapid Learning Printable Tutorial - 13
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Atomic Structure and
Electron Configuration
Rapid Learning Core Tutorial Series
Wayne Huang, PhD
Kelly Deters, MA
Russell Dahl, PhD
Rapid Learning Center
www.RapidLearningCenter.com/
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Objectives
By studying this tutorial you will learn…
„ Basic structure of atoms
„ How to determine the
number of electrons
„ How to place electrons in
energy levels, subshells
and orbitals
„ How to show electron
configurations using three
methods
„ How to write and
understand Quantum
Numbers
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Electron Configuration Concept Map
Previous
Previouscontent
content
Chemistry
Chemistry
New
Newcontent
content
Studies
Quantum
QuantumNumbers
Numbers
Matter
Matter
Location described by
Made of
Electrons
Electrons
Chemical properties
determined by
Atoms
Atoms
3 ways to show configurations
Boxes
Boxesand
andArrows
Arrows
Spectroscopic
Spectroscopic
Notation
Notation
Noble
NobleGas
Gas
Notation
Notation
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Atomic Structure
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Definition: Atom
Atom - n. smallest piece of matter
that has the chemical properties of
the element.
Often called the
“Building Block of Matter”
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What’s in an Atom?
An atom is made of three sub-atomic particles.
Particle
Location
Mass
Charge
Proton
Nucleus
1 amu =
1.67×10-27 kg
+1
Neutron
Nucleus
1 amu =
1.67×10-27 kg
0
Electron
Outside the
nucleus
0.00055 amu
9.10×10-31 kg
-1
1 amu (“atomic mass unit”) = 1.66 × 10-27 kg
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The Atom
Electron
cloud
Nucleus
Mass =
# of protons
+ # of neutrons
Charge =
# of protons
Charge =
- (# of electrons)
Very small
relative mass
Overall Charge =
# of protons
- (# of electrons)
Overall Mass =
# of protons
+ # of neutrons
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Protons Versus Electrons
Protons
Electrons
+ Charge
- Charge
Contributes to mass of
atom
Does not contribute
significantly to mass of atom
Found in nucleus
Found outside nucleus
# determines the “identity”
of the atom
# and configuration
determine how the atom will
react
Cannot be lost or gained
without changing which
element it is (nuclear
reaction)
Can be lost or gained—
results in an atom with a
charge (ion)
The ratio of protons to electrons determines the charge on
the atom.
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Electron Locations
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Definition: Electron Cloud
Electron cloud – It is
the area outside of
the nucleus where
the electrons reside.
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Electron Clouds
Electron
cloud
Principle
energy levels
The electron cloud is
made of energy levels.
Subshells
Energy levels are
composed of subshells.
Orbitals
Subshells have orbitals.
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Definition: Subshell and Orbital
Subshell – A set of orbitals with
equal energy.
Orbital – Area of probability of the
electron being located.
Each orbital can hold 2 electrons.
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Types of Subshells
Energy increases
There are 4 types of subshells that electrons reside in
under ordinary circumstances.
Subshell
Begins in
energy level
Number of
equal energy
orbitals
Total number
of electrons
possible
s
1
1
2
p
2
3
6
d
3
5
10
f
4
7
14
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Pictures of Orbitals
s orbital
3 p orbitals
5 d orbitals
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Electron Configuration
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Definition: Electron Configurations
Electron Configurations – Shows the
grouping and 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.
Electron configurations use boxes for orbitals
and arrow for electrons.
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Aufbau Principle
The first of 3 rules that govern electron configurations
1
Aufbau Principle: Electrons must fill subshells
(and orbitals) so that the total energy of atom is at
a minimum.
What does this mean?
Electrons must fill the lowest available subshells and orbitals
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
6s
5p
5d
4f
4d
5s
4p
3d
4s
3p
3s
2p
Energy
2s
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Subshells are filled from the
lowest energy level to increasing
energy levels.
Not that this does not always go
in numerical order.
1s
Hund’s Rule
The second of 3 rules that govern electron configurations.
2
Hund’s Rule: Place electrons in unoccupied
orbitals of the same energy level before doubling
up.
How does this work?
If you need to add 3 electrons to a p subshell, add 1 to each
before beginning to double up.
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Pauli Exclusion Principle
The last of 3 rules that govern electron configurations.
Pauli Exclusion Principle: Two electrons that
occupy the same orbital must have different spins.
3
“Spin” describes the angular
momentum of the electron
“Spin” is designated with an up
or down arrow.
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.
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Determining the Number of Electrons
In order to properly construct an electron configuration,
you must be able to determine how many electrons to
use.
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
Br1-
How many electrons does the following have?
Charge = -1
Atomic number for Br = 35 = # of protons
-1 = 35 - electrons
Electrons = 36
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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.
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
How many electrons does the following have?
No charge written Æ Charge is 0
Cl
Atomic number for Cl = 17 = # of protons
0 = 17 - electrons
Electrons = 17
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Applying the Rules
Use the 3 rules of electron configurations.
1
Aufbau Principle: Electrons must fill subshells (and orbitals)
so that the total energy of atom is at a minimum.
2
Hund’s Rule: Place electrons in unoccupied orbitals of the
same energy level before doubling up.
3
Pauli Exclusion Principle: Two electrons that occupy the same
orbital must have different spins.
Example:
Give the electron configuration for a Cl atom
No charge written Æ Charge is 0
Cl
Atomic number for Cl = 17 = # of protons
0 = 17 - electrons
Place 17 electrons
1s
2s
Electrons = 17
4
9
70
6
5
3
2
1
8
17
16
15
14
13
12
11
2p
3s
3p
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Spectroscopic Notation
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Definition: Spectroscopic Notation
Spectroscopic Notation – Shorthand way
of showing electron configurations.
The number of electrons in a subshell are
shown as a superscript after the subshell
designation.
1s
2s
2p
3s
3p
1s2 2s2 2p6 3s2 3p5
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Writing Spectroscopic Notation
1
Determine the number of electrons to place.
2
Follow Aufbau’s Principle for filling order.
3
Fill in subshells until they reach their max (s = 2, p = 6, d = 10,
f = 14).
4
The total of all the superscripts is equal to the number of
electrons.
Example:
S
Give the spectroscopic notation for S.
No charge written Æ Charge is 0
Atomic number for S = 16 = # of protons
Electrons = 16
0 = 16 - electrons
Place 16 electrons
2 + 2 + 6 + 2 + 4 = 16
1s 2 2s 2 2p 6 3s 2 3p 4
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Electron Configurations
and the Periodic Table
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Configurations Within a Group
Look at the electron configurations for the Halogens
(Group 7).
1s2 2s2 2p5
F
Cl
1s2 2s2 2p6 3s2 3p5
Br
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
I
All of the elements in Group 7 end with 5 electrons in a p
subshell.
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Configurations and the 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
highest energy subshell that contains electrons.
p-block
s-block
d-block
s1 s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f-block
f1
f2
f3
f4
f5
f6
f7
f8
f9 f10 f11 f12 f13 f14
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Periodic Table as a Road-Map
Wondering how to remember the order of filling of
the subshells?
Just use the periodic table.
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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)
Periodic Table as a Road-Map
To see the filling order of subshells, read from left to right,
top to bottom!
This tool shows that the 3d energy level is filled after the 4s energy level!
1s
1s
2p
2s
3p
3s
4s
3d
4p
5s
4d
5p
6p
6s
4f
5d
7s
5f
6d
s subshells begin in level 1, so begin the s-block with “1s”
p subshells begin in level 2, so begin the p-block with “2p”
d subshells begin in level 3, so begin the d-block with “3d”
f subshells begin in level 4, so begin the f-block with “4f”
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Another Tool for Filling Order
There is another tool commonly used to remember
orbital filling order.
1s
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2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
To read the charge,
move down one
diagonal as far as
possible, then jump to
the top of the next
diagonal and keep
going.
8s
Electron Configurations of
Ions
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Definition: Ion
Ion – an atom that has gained or lost
electrons resulting in a net charge.
Atoms gain and lose electrons to be in a more stable state.
Usually, the “more stable state” is a full valence shell.
Outermost shell of electrons
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Full Valence Shell Ions
Look at the electron configurations for the
following:
Br-1
p = 35
-1 = 35 - e
e = 36
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6
O2-
p=8
-2 = 8 - e
e = 10
+1 = 11 - e
e = 10
+2 = 20 - e
e = 18
1s 2 2s 2 2p 6
Na+
p = 11
1s 2 2s 2 2p 6
Ca2+
p = 20
1s 2 2s 2 2p 6 3s 2 3p 6
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Full Valence Shell Ions
What do you notice about each of these
configurations?
They all end with full p subshells.
Br-1
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6
O21s 2 2s 2 2p 6
Na+
1s 2 2s 2 2p 6
Notice that O2- and Na+ have the
same number and configuration of
electrons.
This makes them isoelectric.
Ca2+
1s 2 2s 2 2p 6 3s 2 3p 6
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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.
Br
Spectroscopic
1s 2
2s 2
2p 6
3s 2
3p 6 4s 2 3d 10 4p 5
[Ar] 4s 2 3d 10 4p 5
Noble gas
The “[Ar]” represents the core electrons and only the valence electrons are shown.
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Which Noble Gas Do You Choose?
How do you know which noble gas to use to
symbolize the core electrons?
Think: Price is Right.
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!
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Noble Gas
# of electrons
He
2
Ne
10
Ar
18
Kr
36
Xe
54
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Where Does the Noble Gas Leave Off?
How do you know where to start off after using a
noble gas?
Use the periodic table!
1s
He
2s
3s
2p
Ne
3p
Ar
4s
3d
4p
Kr
5s
4d
5p
Xe
6p
Rn
6s
4f
5d
7s
5f
6d
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.
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Noble Gas Notation Example
1
Determine the number of electrons to place.
2
Determine which noble gas to use.
3
Start where the noble gas left off and write spectroscopic
notation for the valence electrons.
Example:
As
Give the noble gas notation for As.
No charge written Æ Charge is 0
Atomic number for As = 33 = # of protons
0 = 33 - electrons
Electrons = 33
Place 33 electrons
Closest noble gas: Ar (18)
[Ar] 4s 2 3d 10 4p 3
Ar is full up through 3p
18 + 2 + 10 + 3 = 33
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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
“Boxes and
arrows”
Shows if electrons
are paired or
unpaired
Longest method
Spectroscopic
notation
Quicker than “Boxes
and arrows”
Does not show
pairing of electrons
Does not show core
electrons
Noble Gas
notation
Allows focus on the
valence electrons
(that control
bonding)
Quickest method
Does not show
pairing of electrons
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Exceptions to the
Aufbau Rule
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Stability of d Subshells with 5 or 10
d subshells have 5 orbitals…
They can hold 10 electrons.
According to the Aufbau principle, Cr should have the
following valence electron configuration:
4s2 3d4
But a half-full or completely full d subshell is more stable
than the above configuration, so it is:
4s1 3d5
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Elements with Exceptions
The following elements are excepts to the Aufbau
Principle:
Element
Should be
Actually is
Cr
4s2 3d4
4s1 3d5
Mo
5s2 4d4
5s1 4d5
W
6s2 5d4
6s1 5d5
Cu
4s2 3d9
4s1 3d10
Ag
5s2 4d9
5s1 4d10
Au
6s2 5d9
6s1 5d10
They are the two groups on the periodic table that begin with Cr and Cu.
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Quantum
Numbers
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Definition: Quantum Numbers
Quantum Numbers – A set of
4 numbers that describes
the electron’s placement in
the atom.
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4 Quantum Numbers
ml
n
2, 1, -1, + ½
l
Quantum
Number
ms
Symbol
Describes
Possible
Numbers
Principal
n
Shell number
Whole #s ≥ 1
Azimuthal
l
Subshell
type
Whole # < n
Magnetic
ml
Orbital
-lÆ+l
Spin
ms
Spin
+ ½ or – ½
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Determining Quantum Numbers
4p 3
n: principal energy level
Give the number of the shell
l: subshell
s=0
p=1
d=2
f=3
coding system
ml: orbital
s
0
p
-1
0
Number-line system of identifying orbitals.
0 is always in the middle.
Number line from – l to + l
1
f
d
-3
-2
-2
-1
0
1
2
-1
0
1
2
3
↑= + ½
↓=-½
ms: spin
Coding system
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Quantum Number Examples
Example: Give the quantum numbers for the red arrow.
1s
2s
2p
3s
3p
0
It’s in level “3”
It’s in subshell “s”—the “code” for “s” is “0”
It’s in orbital “0”
___,
3 ___,
0 ___,
0 ___
-½
It’s a down arrow
Example: Give the quantum numbers for the red arrow.
1s
2s
It’s in level “2”
2p
3s
-1
3p
0 +1
It’s in subshell “p”—the “code” for “p” is “1”
It’s in orbital “-1”
It’s an up arrow
___,
2 ___,
1 ___,
-1 ___
+½
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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”
There is no p subshell in energy level 1
2, 1, ;
-2, - ½
n = 2…OK as n can be any whole # >0
l = 1…subshell is “p”
OK as level 2 has “p”
ml = -2…on the “-2” orbital
“p” subshell has 3 orbitals: ___ ___ ___
-1 0 +1
No “-2” orbital in a “p” subshell.
ml must be between –l and l
1, 0, 0, ;
-1
n = 1…OK as n can be any whole # >0
l = 0…subshell is “s”
OK as level 1 has an “s”
ml = 0…on the “0” orbital
OK as “s” has 1 orbital and it’s “0”
ms = -1
ms must be either + ½ or – ½
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Summary
Electron
Electron
configurations
configurations can
can
be
be shown
shown with
with
boxes
boxes and
and arrows,
arrows,
in
in spectroscopic
spectroscopic
notation,
notation, or
or noble
noble
gas
gas notation.
notation.
Atoms
Atomsare
are made
made of
of
protons,
protons, neutrons
neutrons
and
andelectrons.
electrons. The
The
configuration
configuration of
of the
the
electrons
electronsdetermines
determines
the
the chemical
chemical
properties
properties of
of the
the
atom.
atom.
Electrons
Electrons are
are
organized
organized in
in
levels,
levels, subshells
subshells
and
and orbitals.
orbitals.
Quantum
Quantum numbers
numbers
describe
describe the
the
location
location of
of an
an
electron
electron in
in an
an
atom
atom and
and are
are aa
series
series of
of 44
numbers.
numbers.
Electron
Electron configurations
configurations
are
are written
written following
following
the
the Aufbau
Aufbau principle,
principle,
Hund’s
Hund’s Rule
Rule and
and the
the
Pauli
Pauli Exclusion
Exclusion
Principle.
Principle.
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Congratulations
You have successfully completed
the core tutorial
Atomic Structure and
Electron Configuration
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