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Fireworks!
1. Do you enjoy watching fireworks? Why or
why not?
2. What are the colors and shapes and sounds
of a firework that have most impressed you?
Today’s assignment:
Re-read the article and find 2-3 interesting facts.
Research 2-3 amazing fireworks facts or stories that
you would have included in the article.
Resources are listed on the handout.
Then in your group, design your own firework
display.
<google “fireworks graphic design”>
Here’s one for you:
The electron configuration of carbon is
1s22s22p2.
How many electrons are present in carbon?
What is the atomic number of carbon?
Bellwork
1) The electron configuration of nitrogen is
1s2 2s2 2p3.
What is the atomic number of nitrogen?
How many electrons are in a nitrogen atom?
Write the orbital notation for nitrogen.
Bellwork – 1/27
The electron configuration of boron is
1s2 2s2 2p1.
What is the atomic number for boron?
How many electrons are present in an
atom of boron?
Which period of the periodic table is it in?
Which elements are in the same period?
At the end of class today, 1/26
you should be able to:
Write the orbital diagram and electron
configuration for any 1st or 2nd period element.
Write the noble gas configuration for any
3rd or 4th period elements!
Bellwork 1/13
1) What is the total number of orbitals (not
shapes) in the 3rd main energy level?
2) What’s the difference between a shell, a
subshell and an orbital?
3) How many electrons can a p sublevel hold?
Bellwork 1/15
Which of the principles do each of the following
electron energy diagrams violate?
1)
2)
3)
At the end of class today you should
be able to :
State in your own words the Aufbau principle,
the Pauli exclusion principle, and Hund’s rule.
Target Practice!
Show what cha know!
Bellwork 1/28
Draw and label (with the arrows) the orbital
diagram for fluorine.
Identify these elements:
[Kr] 5s1
[Ar] 4s2 3d5
1/13 Target – turn it in!
In your own words:
1. State the Aufbau principle.
2. State the Pauli exclusion principle.
3. State Hund’s rule…
Section 3:
Electron Configurations
Previously in Chapter 4…
Electrons live in little houses called orbitals.
Bohr thought that a whole bunch of electrons
could live in one house at a time, but the
quantum model made it clear that, like a studio
apartment with an unfriendly landlord, only two
can live in one at the same time.
In this section:
We will list the total number of electrons needed to
fully occupy each main energy level.
State the Aufbau principle, the Pauli exclusion
principle, and Hund’s rule.
Describe the electron configurations for the atoms
of any element using orbital notation, electronconfiguration notation and noble-gas notation.
Finally!
The quantum model of the atom improves on
the Bohr model because it describes the
arrangements of electrons in atoms – other than
hydrogen…
The arrangement of electrons in an atom is
known as the atom’s electron configuration.
Unique!
Because atoms of different elements have
different numbers of electrons, a unique
electron configuration exists for the atoms of
each element.
Like all systems in nature, electrons want to be
arranged with the lowest possible energy… just
like high school students!
Flashback!!
What is an atom’s ground state???
Remember ground state?
The element’s ground state configuration is the
lowest-energy arrangement of electrons for that
element.
A few simple rules combined with quantum
numbers allow us to determine the ground state
configuration.
This is carbon’s electron configuration:
2
1s
2
2s
2
2p
The integers are n (main energy levels).
The letters are l (shape/orbital).
The superscripts are how many electrons are in
each level.
Electrons fill the lowest-energy orbitals
first.
Step 1: Figure out the energy levels of the
orbitals.
Step 2: Electrons are added to the orbitals, one
by one, according to 3 basic rules.
Rule #1: Aufbau principle
Rule #1: According to the Aufbau principle,
electrons fill up lower energy orbitals first, and
then start filling up higher energy orbitals.
In a ground-state H atom, the electron is in the
1s orbital.
Interesting sidenote:
The word “aufbau” isn’t somebody’s name. It
means something in German, but it doesn’t
really matter… 
And back to the lesson…
Chapter 4
Configurations
Relative Energies of Orbitals
Rule #2: People who are remembered
for only one thing…
Wolfgang Pauli stated that all electrons in an
element have unique sets of quantum numbers,
an idea known as the Pauli Exclusion Principle.
Because quantum numbers are the
mathematical description of each electron,
making sure all elements have unique sets of
quantum numbers allows us to tell them apart.
And that’s the one thing we remember Pauli for.
Rule #2: Pauli Exclusion Principle
Rule #2 reflects the importance of the spin
quantum number.
According to the Pauli exclusion principle, no
two electrons in the same atom can have the
same set of four quantum numbers.
An orbital only holds 2 electrons of opposite
spin states.
Rule #3: Hund’s Rule
According to Hund’s rule, electrons prefer to be
unpaired whenever possible.
This means they will always enter an empty
orbital before they pair up.
Hund’s rule also states that all single electrons in
orbitals must have the same spin state.
Hund’s Rule in Action…
Show 2 e- filling the p orbitals
Show 3 e- filling the p orbitals
Show 4 e- filling the p orbitals
Did you mind the rules?!
With a partner, take turns to state Aufbau
principle, Pauli Exclusion principle and Hund’s
Rule.
Use the 5 orbital boxes in your notes to explain
how 3, then 5, then 8 then 10 electrons will fill
these orbitals.
Show how 3, 5, 8 and 10 electrons fill
these orbitals.
Try it yourself and let’s check it!
All you need is right here!
Find the period/row… which equals “n”. Then
use what you know about quantum numbers to
fill in the electrons.
Orbital diagrams are easy!
Orbital diagrams are pictures that show what
electron configurations look like when you draw
them.
Because each term in an electron configuration
is lower in energy than the one before it, we
pretty much just draw a bunch of lines and fill
‘em full of electrons.
Each electron is represented by an arrow,
and the up and down arrows in a single
orbital represent the two different
electrons in this orbital.
• In this method, each line (or box)
represents a single orbital, which is why
there’s one line for s-orbitals (there’s
only one s-orbital for every energy level),
• three lines for p-orbitals (there are 3 porbitals for every energy level – except
n=1, right??),
Let’s start at the start…
H and He
What is the atomic number?
How many electrons?
Draw the orbital diagram then the electron
configuration.
Let’s try some 2nd period notations
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Li
Be
B
C
N
O
F
Ne
Noble gas notation
• If we want to make our lives A LOT simpler, we
don’t have to write out all of these terms that
are in normal electron configurations.
• Instead of writing out a big string of orbitals,
just abbreviate it so you start at the most
recent noble gas to the element you’re
interested in.
For example:
Ne is 1s2 2s2 2p6, right?
In the 3rd period, Na is 1s2 2s2 2p6 3s1
Using noble configuration it becomes
Na = [Ne] 3s1
Mg = [Ne] 3s2
This is sooooooo much simpler!
Use noble gas notation for the 3rd
period elements
Na
[Ne] 3s1
Mg [Ne] 3s2
Al
Si
P
S
Cl
Ar
Let’s look at 4th period elements
K [Ar] 4s1
Ca [Ar]4s2
Sc [Ar] 4s2 3d1
Just keep following the row in the periodic table
BUT REMEMBER, the d block is LOWER than the
s block.
From 4s1 and 4s2… it is 3d1, 3d2, 3d3…3d10
Ti [Ar] 4s2 3d2
V [Ar] 4s2 3d3
Whoa! Look!
Cr [Ar] 4s1 3d5
Why is Cr: [Ar] 4s1 3d5??
Shouldn’t that be
Cr [Ar] 4s2 3d4???
Let’s look at the orbital diagram of chromium
and sort this out!
Compare these:
Cr: [Ar] ___
4s
___ ____ ___ ___ ____
3d
Cr: [Ar] ___
4s
___ ____ ___ ___ ____
3d
4s1 3d5 is more stable than 4s2 3d4, even
though it follows our pattern…
Let’s pick up the pattern again…
Mn [Ar] 4s2 3d5
Fe
Co
Ni
Watch out for copper!!
It happens one more time in the 3rd
period…
Copper!
Cu
[Ar] 4s1 3d10
Cu: [Ar] ___
___ ____ ___ ___ ____
There isn’t a simple explanation…
Everybody else behaves and follows
our neat pattern…
Zn
Ga
Ge
As
Se
Br
Kr
[Ar] 4s2 3d10
[Ar] 4s2 3d10 4p1
[Ar] 4s2 3d10 4p2
Two definitions for you…
The highest occupied energy level of an electron is
the energy level with the highest principal quantum
number.
For example H, n=1 and the highest energy level is
the first main energy level.
For Be, n=2 and the highest energy level is the
second main energy level.
The inner shell electrons are the electrons that
are not in the highest occupied energy level.
That totally makes sense!
This is also a good time to introduce
Valence electrons are the electrons in the
outermost (highest) principal energy level of an
atom.
For example, N 1s2 2s2 2p3 has e- in n=1 and
n=2.
N=2 is the outermost so the valence e- are the
2s and 2p electrons.
For sodium, 1s2, 2s2, 2p6 3s1 or [Ne] 3s1,
the valence e- is the electron in the
3s orbital because n=3 is the outermost
level.
Valence electrons are vewwy, vewwy
important!
Valence electrons are the most important
because they are the ones involved when atoms
form bonds (attach to each other).
And did you happen to notice a trend in the
periodic table??
Indeed!
The atoms of elements in the same group
(vertical column) have the same number of
electrons in a given type of orbital.
The only difference is the principal energy level
of that sublevel.
Thank you, Mr. Mendeleev
Remember that elements were originally organized
into groups on the periodic table based on similar
chemical properties.
Now we can understand the reason behind these
groupings!
Elements with the same valence electron
arrangement show very similar chemical behavior!
Sample B pg. 115
Now it’s your turn!
Chapter 4
Section 3 Electron Configurations
Sample Problem B
• a. Write both the complete electronconfiguration notation and the noble-gas
notation for iron, Fe.
• b. How many electron-containing orbitals are in
an atom of iron? How many of these orbitals are
completely filled? How many unpaired electrons
are there in an atom of iron? In which sublevel
are the unpaired electrons located?
Chapter 4
Section 3 Electron Configurations
Sample Problem B Solution
•
a. The complete electron-configuration notation of iron is
1s22s22p63s23p63d64s2. Iron’s noble-gas notation is
[Ar]3d64s2.
•
•
b. An iron atom has 15 orbitals that contain electrons.
They consist of one 1s orbital, one 2s orbital, three 2p
orbitals, one 3s orbital, three 3p orbitals, five 3d orbitals, and
one 4s orbital.
Eleven of these orbitals are filled, and there are four
unpaired electrons.
They are located in the 3d sublevel.
The notation 3d6 represents 3d
•
•
•
Practice Problems
1a) Write the noble gas notation for I:
1b) How many unpaired e- are there?
2a) Write the noble gas notation for Sn:
2b) How many unpaired e- are there?
Answers Practice B
1a) 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5,
[Kr]4d10 5s2 5p5
1 unpaired eb) 27, 26, 1
2 a) [Kr]4d10 5s2 5p2, 2
b) 10, germanium
3a) 1s2 2s2 2p6 3s2 3p6 3d5 4s2
b) manganese
4 a) 9, 1s2 2s2 2p6 3s2 3p6
b. argon
Let’s practice together:
1a) How many inner shell electrons does an
atom of silicon, Si, contain?
b) How many electron containing orbitals are in
an atom of silicon?
How many of these orbitals are filled?
How many unpaired electrons are there in an
atom of silicon?
c) How many electron containing p orbitals are
there in an atom of silicon?
Answers:
a) 10
b) 8, 6, 2
c) 5
Let’s try this one from the book!
Sample C, pg. 116
Chapter 4
Section 3 Electron Configurations
Sample Problem C
• a. Write both the complete electronconfiguration notation and the noblegas notation for a rubidium atom.
• b. Identify the elements in the second,
third, and fourth periods that have the
same number of highest-energy-level
electrons as rubidium.
Chapter 4
Section 3 Electron Configurations
Sample Problem C Solution
•
a. 1s22s22p63s23p63d104s24p65s1, [Kr]5s1
•
b. Rubidium has one electron in its
highest energy level (the fifth). The
elements with the same outermost
configuration are,
in the second period, lithium, Li;
in the third period, sodium, Na;
and in the fourth period, potassium, K.
•
•
•
Practice problems 1 – 2
Let’s get it a go!
Answers in Appendix E, in the back of the book…
Next, we need to try Section 3, Formative
Assessment… which can only lead to
ONE THING:
For practice:
Textbook pg. 120
#26-31.
Homework answers:
Aufbau principle
26a) An electron occupies the lowest energy
orbital that can receive it.
b) The lowest energy orbital is filled first.
Electrons are then added to the orbital with the
next lowest energy and so on until all of the
electrons in the atom have been placed in
orbitals.
Hund’s rule
27a) Orbitals of equal energy are each occupied
by one electron before any orbital is occupied by
a second electron.
b) By placing as many single electrons as
possible in separate orbitals in the same energy
level, electron-electron repulsion is minimized
and favorable lower energy arrangements result.
28a) No two electrons in the same atom can
have the same four quantum numbers.
b) The 2 different values of the spin quantum
number permit two electrons of opposite spin
states to occupy the same orbital.
29a) the highest occupied energy level in an
atom is the electron containing main energy
level that has the highest principal quantum
number.
b) Inner shell electrons are electrons that are
not in the highest occupied energy level.
30a) first main energy level n=1
b) 2nd main energy level n=2
c) 3rd main energy level n=3
d) 4th main energy level n=4
e) 5th main energy level n=5
Practice time!
1a) 1s2 2s2 2p6 3s2 3p6 3d2 4s2; [Ar]3d2 4s2
b) 12: one 1s orbital, one 2s orbital, three 2p
orbitals, one 3s orbital, three 3p orbitals, one 4s
orbital and two 3d orbitals;
10
2