In This Lesson:
Unit 2
Electrons,
Orbitals, and
Atomic Model
History
(Lesson 1 of 4)
Stuff You Need:
Periodic Table
Today is Wednesday,
February 22nd, 2017
Pre-Class:
In your notebooks, draw a picture of
electrons moving around the atom’s nucleus.
Include arrows to show direction.
Today’s Agenda
• A little history review…
• Electron Configuration
– Also known as “Where the electrons at?”
•
•
•
•
•
•
Electron Orbitals and Quantum Numbers
Heisenberg Uncertainty Principle
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
And coloring!
• Where is this in my book?
– P. 128 and following…
– Oh, by the way, quantum numbers aren’t in there. You heard me.
By the end of this lesson…
• You should be able to describe the Quantum
Mechanical Model of the atom.
• You should be able to indicate the
arrangement and locations of electrons in
multiple formats.
Guiding Video
• TED: George Zaidan and Charles Morton – The
Uncertain Location of Electrons
In the beginning…
• There was Democritus, a Greek
professor (460 BC – 370 BC).
– He came up with the term “atom” to
describe the tiny particles he suggested.
• Then there was John Dalton (1803).
– He studied combinations of elements in
chemical reactions.
– His atomic model was just a solid ball.
Discovery of the Electron
• In 1897, JJ Thomson used
a cathode ray tube to
deduce the presence of a
negatively charged
particle.
– Cathode ray tubes pass
electricity through a gas
that is contained at a very
low pressure.
Conclusions from Studying Electrons
• Cathode rays have identical properties regardless
of the element used to produce them. All
elements must contain identically charged
electrons.
• Atoms are neutral, so there must be positive
particles in the atom to balance the negative
charge of the electrons.
• Electrons have so little mass that atoms must
contain other particles that account for most of
the mass.
In shorter terms…
• Electrons are important because:
– They create ions.
– They lead to bonding.
– They determine how atoms behave.
Thomson’s Atom (1897)
• Called the Plum Pudding Model, as Thomson
thought that electrons were like plums sitting
in a positive pudding.
JJ Thomson
Rutherford and the Nucleus
• Ernest Rutherford fired α
particles (helium nuclei)
at an extremely thin
sheet of gold foil.
• He recorded where the
particles “landed” after
striking (or passing
through) the gold.
Ernest Rutherford
“Like Howitzer shells
bouncing off of tissue
paper.”
Rutherford’s Findings
• Because most particles passed through and only a
very few were significantly deflected, Rutherford
concluded that the nucleus:
– Is small
– Is dense
– Is positively charged
Rutherford’s Atom (1913)
• After the Rutherford
experiment, the atom model
looked like this:
• Looked like the Infinity Ward
logo, but it’s wrong.
http://www.epa.gov/radiation/images/ruthbohr.jpg
http://www.thatvideogameblog.com/wp-content/uploads/2010/08/infinity-ward-logo.jpg
Eugen Goldstein and the Proton
• Eugen Goldstein is
sometimes credited with
the discovery of the
proton.
– Other times it goes to
Wilhelm Wien who
performed other critical
measurements of the
proton using an anode ray
(somewhat like Thomson’s
cathode ray).
Eugen Goldstein
http://www.pkc.ac.th/kobori/Assets/ChemistryMahidol1/www.il.mahidol.ac.th/course/ap_chemistry/ato
mic_structure/picture/bild_goldstein.jpg
Jimmy Neutron and the Rutherford Atom?
• Even Jimmy Neutron has an image of the
Rutherford Model on his shirt!
– Not so “boy genius” after all…
Bohr’s Atom (1913)
• Bohr thought of electrons
moving around the nucleus
like planets around the Sun.
• His was a flat model of the
atom.
• In reality, the electrons
actually move around the
nucleus like bees around a
hive.
Niels Bohr
The Bohr Model
• Niels Bohr, among other
things, proposed the Bohr
Model.
• Unlike Rutherford’s atom,
which had electrons all at
approximately the same
distance from the nucleus,
Bohr’s model showed them
orbiting in a flat space but at
different, fixed distances:
http://www.thephysicsmill.com/blog/wp-content/uploads/bohr_model_no_emission.png
– More on this in our next lesson.
• In 1926, Erwin Schrödinger develops
equations that lead to the electron cloud
model of the atom.
– Electrons around found in a threedimensional space around the nucleus and
are more likely to be found closer-in.
• Combined, these two discoveries do away
with the Bohr model but require a more
complex model of the atom.
Erwin Schrödinger
• In 1923, Louis de Broglie discovered that
particles as small as electrons have some
wave-like properties (as opposed to strictly
particle-like).
Louis de Broglie
Schrödinger’s Atom (1926)
http://upload.wikimedia.org/wikipedia/commons/thumb/2/26/Erwin_Schrödinger.jpg/220px-Erwin_Schrödinger.jpg
http://1.bp.blogspot.com/_GVA115I1I8Y/TT6_AHLks3I/AAAAAAAABWo/uvD4LGMKRgY/s1600/Broglie_Big.jpg
Chadwick and the Neutron
• Chadwick
discovered the
neutron in 1932
and won the Nobel
Prize three years
later for it.
http://www.dnahelix.com/jimmy/jnmov_jn_ext_shrinkray.jpg
http://www.nobelprize.org/nobel_prizes/physics/laureates/1935/chadwick.jpg
James Chadwick
Modern Atomic Theory
• All matter is composed of atoms.
• Atoms cannot be subdivided, created, or
destroyed in ordinary chemical reactions.
However, these changes can occur in nuclear
reactions!
• Atoms of an element have a characteristic
average mass which is unique to that element.
• Atoms of any one element differ in properties
from atoms of another element.
The Quantum Mechanical Model
• The currently-accepted model is the Quantum
Mechanical Model of the atom.
• In it, mathematical models determine the
most likely positions of electrons around the
nucleus.
– Sound complicated? It is.
• Instead of exploring the laws, we’re going to
look at some of the “results” of them.
– But first, an actual look at atoms on camera.
• NOVA video.
Heisenberg Uncertainty Principle
• Werner Heisenberg
discovered that you can
find out where an electron
is, but not where it’s going.
• Alternatively, you can find
out where it’s going but not
where it is.
– Not both.
http://www.wired.com/images_blogs/underwire/2012/09/heisenberg_660.jpg
“One cannot simultaneously
determine both the position
and momentum of an
electron.”
Heisenberg Uncertainty Principle
• To be able to see things, light must strike an object and
then bounce off of it, returning to your eye.
• For objects like, say, bowling balls, light strikes it and
the bowling ball just sits there.
• For electrons, however, they have so little mass that
when light strikes them, they move in a different
direction.
http://cdn2-b.examiner.com/sites/default/files/styles/image_full_width/hash/0d/2e/0d2e398879c6b94255370961648165a2.jpg
Guiding Example
• Now, before we dive face-first into electron
orbitals, we’re going to use a “guiding example”
from something not-so-scientific to understand
the concepts behind them.
• The Hog Hotel!
• Remember, as we explore this analogy, the goal
of this entire lesson is to learn how electrons
configure themselves around the nucleus.
– It’s a big game of hide and seek with electrons!
The Hog Hotel Analogy
• Imagine you’re the manager of a towering
hotel (for pigs) and you have a list of pigs that
want to stay there.
• Here are the rules you need to follow:
– Rooms must be filled from the ground up.
– Only singles first. No pig gets a roommate until all
rooms on one floor are filled.
– If two pigs are staying in the same room, they will
face opposite directions. Weird.
The Hog Hotel Analogy
• On your Hog Hotel worksheets, try the first
page and #2 on the back of the first page.
• Then we’ll go over it.
• Then we’ll do the rest of the back page.
Electron Energy Levels (Shells)
• Rising up from the lobby
of the hotel are the
various floors hogs
might occupy.
• Moving away from the
nucleus are the various
energy levels electrons
might occupy.
• These energy levels are
symbolized by n.
Energy Level 1  n=1
Energy Level 2  n=2
n
• n is the Principal
Quantum Number.
• To determine how
many electrons fit into
a given energy level,
use this formula:
Electrons = 2n2
Energy Level 1  n=1
Energy Level 2  n=2
Aufbau Principle
• In German, aufbau means “building up.”
• The Aufbau Principle states that electrons,
when not excited, will fill energy levels starting
at the lowest energy.
– In the Hog Hotel, this was the rule that the hogs
are lazy and prefer rooms on the lowest floors
possible.
Orbital Shapes
• Imagine that each room in the hotel, even on
the same floor, has a different shape.
• In the atom, on the energy level are sublevels
consisting of orbitals where there is a 90%
probability of finding an electron.
– An orbital is like a specific room (indicated
sometimes by a direction).
• Orbitals can hold up to 2 electrons.
– A sublevel is like a group of rooms or a suite
(indicated by a letter – also called subshells).
• Sublevels can hold 1, 3, 5, or 7 orbitals.
Orbital Hotel Rooms?
• For the next few slides, I’m going to show you
pictures of orbitals.
• Think of these as rooms in a weird atomic
hotel.
– Some are basic rooms, holding only two electrons.
– Some are like suites, with individual rooms
comprising a larger room.
• They don’t all appear on every floor, however.
• I’ll explain what I mean with a look back at
two of my dorm rooms from college.
My Freshman Year of College
e-
e-
I had the basic two bed/one roommate set up.
Also, my roommate was awful but that’s besides the point.
e-
e-
e-
My Sophomore Year of College
ee-
We had what our school called a suite, which was an arrangement of
mini-rooms. Let’s compare this to the atom and its “rooms.”
Orbital
s Sublevel
e-
e-
s Sublevels
• Shape: Sphere
• Appears: n=1 and
above.
• # of Orbitals: 1
• Capacity: 2 e-
Orbital
p Sublevel
e-
e-
e-
e-
e-
e-
p Sublevels
•
•
•
•
Shape: Dumbbell (3)
Appears: n=2 and above.
# of Orbitals: 3 (x, y, z)
Capacity: 6 e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
Orbital
d Sublevel
e-
d Sublevels
• Shape: Double
Dumbbells (4) and
Dumbbell Doughnut
• Appears: n=3 through
n=6.
• # of Orbitals: 5
• Capacity: 10 e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
e-
Orbital
f Sublevel
e-
f Sublevels
• Shape:
Flowers…and
stuff.
• Appears: n=4
through n=5.
• # of Orbitals: 7
• Capacity: 14 e-
And the “hotel” as a whole?
4s
4p
3s
4d
3p
2s
3d
2p
1s
4f
After f?
• Right now there are no elements in existence
that have electrons at energy levels higher
than 7.
• There are also no sublevels beyond f.
• However, if somehow we were to create an
atom that had so many electrons we filled the
f sublevel on the n=5 energy level, what would
be next?
• g, then h and so on in alphabetical order.
You Should Know…
• You may be feeling a little overwhelmed.
• If you understand this, you’re in good shape:
– Around the atom are energy levels, like floors in a hotel
room. The farther out, the higher energy.
– Each energy level has sublevels, like “types of rooms” in a
hotel.
– Each sublevel has one or more orbitals, which are like
individual rooms. For example, s sublevels have one
orbital, whereas p sublevels have three orbitals.
– These orbitals each can hold two electrons and show the
90% likely location of those two electrons at any time.
Quick Review
• How many electrons can fit into that s sublevel?
• 2
• Which energy level is farther from the nucleus,
n=2 or n=5?
• 5
• How many electrons can fit at the 2nd energy
level? (n=2)
• 8 (remember 2n2?)
• In which energy level does the f sublevel start to
appear?
• n=4
Summary Table
Sublevels
Orbitals Per
Sublevel
Electrons Per
Sublevel
Electrons Per
Energy Level
(2n2)
1
s
1
2
2
2
s
p
1
3
2
6
8
3
s
p
d
1
3
5
2
6
10
18
4
s
p
d
f
1
3
5
7
2
6
10
14
32
Floor Number
Type of
Rooms/Suites
on Floor
Rooms per
Type of
Room/Suite
Capacity of
Each Type of
Room/Suite
Capacity of
Each Floor
Energy Level
(n)
Putting It All Together
• Let’s try the third and fourth pages of the hog
hotel worksheet.
• It’s the same thing we’ve been doing, only
using “up arrows” and “down arrows” instead
of forward and backward letters.
Orbital Notation
• What you have just learned (the arrow way of
writing electrons) is called orbital notation.
• As it turns out, there’s a pattern to finding the
orbitals in which the electrons are placed.
– Mendeleev was on to something!
• Let’s do some color-coding so we can predict
what orbitals to write.
Electron Configuration Table
s1
1s
s2
p1 p2 p3 p4 p5 p6
s2
2s
2p
3s
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 3p
4s
5s
s
3d
4d
6s
5d
7s
6d
4p
d
5p
p
6p
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
5d 4f
6d 5f
f
Inner Transition Metals
• Below the table are the inner transition
metals (f block).
• They look disconnected, but really they are
“within” the transition elements (d block).
• Expanded, the table would look like this:
d and f Sublevels
• Uh, wait a second…
• It looks like according to the table we just shaded,
d and f sublevels are going out of order.
– In the n=6 row, it’s 5d and 4f.
– What’s the deal?
• d and f sublevels exist at lower energy levels than
p sublevels (starting at n=4), so they’ll be filled
first according to the Aufbau Principle.
– Stick with me here – I’ll teach you an easy way to
remember that.
Writing Configurations
• Chemists need to be able to effectively record the
electron configurations of various atoms. Consider
Neon, the first element on the last page of the Hog
Hotel.
• Neon is in the second row (n=2), so there are
electrons in n=1 and n=2.
–1
2
• There are electrons in sublevels 1s, 2s, and 2p.
– 1s 2s 2p
• Finally, there are two electrons in sublevel 1s, two in
subshell 2s, and 6 in subshell 2p.
– 1s2 2s2 2p6 (electron configuration)
– ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ (orbital notation)
1s
2s
2p
Two Ways to Figure This Out…
• It can be hard to remember the
order of the various quantum
numbers and subshells.
• You can figure out the electron
configuration of an element two
ways.
– The easy way and the hard way.
– Just kidding. They’re just
different.
• One way is the diagonal rule.
• This:
• The other way is hard to explain
in writing, but I like it better.
Directions for Using the Cheat Sheet
• Target your element.
• Starting with hydrogen, move left to right across
the rows, moving down one each time you reach
the end.
• Every time you either A) reach the end of a row
or B) change blocks, write down the “address” of
the last element in that section.
• Stop when you get to your element.
• Check your work! You should be able to count
the same number of electrons (more on that in a
bit).
Electron Configuration for Ne
s1
1s
s2
s2
Ne:
1s2
2s2
p1 p2 p3 p4 p5 p6
2p6
2s
2p
3s
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 3p
4s
5s
s
3d
4d
6s
5d
7s
6d
4p
d
5p
p
6p
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
5d 4f
6d 5f
f
Electron Configuration
Element
Electron Configuration
Hydrogen
1s1
Helium
1s2
Lithium
1s22s1
Beryllium
1s22s2
Boron
1s22s22p1
Carbon
1s22s22p2
Nitrogen
1s22s22p3
Oxygen
1s22s22p4
Fluorine
1s22s22p5
Neon
1s22s22p6
Use this for the next slide’s questions
s1
1s
s2
p1 p2 p3 p4 p5 p6
s2
2s
2p
3s
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 3p
4s
5s
s
3d
4d
6s
5d
7s
6d
4p
d
5p
p
6p
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
5d 4f
6d 5f
f
Let’s try a few practice elements…
• Cobalt (Co):
– 1s2 2s2 2p6 3s2 3p6 4s2 3d7
• Europium (Eu):
– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1
4f6
• Tungsten (W):
– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14
5d4
– Notice how we had to do a little rearranging at the
end of the electron configuration for Tungsten.
Electron Configuration for W
s1
1s
W: 1s2 2s2 2p6 3s2 3p6
4s2 3d10 4p6 5s2 4d10
5p6 6s2 5d1 4f14 5d4
s2
2s
3s
4s
5s
s
s2
p1 p2 p3 p4 p5 p6
2p
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 3p
3d
4d
6s
5d
7s
6d
4p
d
5p
p
6p
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
5d 4f
6d 5f
f
On your worksheets…
• Try the first page of the worksheet labeled,
“Electron Configurations Orbital Notation.”
• We’ll do the first one (Mg) together.
Things to Check
• Suppose you’ve just written Magnesium’s
electron configuration:
– 1s2 2s2 2p6 3s2
• To make sure you’re right, check how many
electrons magnesium has.
– 12
• Do the “exponents” in the configuration add
up to the same amount?
– 1s2 2s2 2p6 3s2 = 12
Ions
• Writing electron configurations of ions is easy:
• Step 1: Figure out how many electrons the ion
has. [remember, protons – electrons = charge]
• Step 2: Make that number the new atomic
number.
• Step 3: Target an element with that new atomic
number.
– Example: Oxygen with a charge of -2 (O2-) has two
extra electrons.
– It’s basically like diagramming a Neon atom.
– O2- = 10 e- = 1s2 2s2 2p6
Noble Gas Notation
• Try this: Write the electron configuration for
Neon in your notebooks.
– 1s22s22p6
• Now try this: Write the electron configuration
for Sodium underneath.
– 1s22s22p63s1
• Notice anything?
Shorthand Notation
• Notice that the configurations build on one another.
• To save time, scientists use Shorthand Notation (or
Noble Gas Notation) to condense the writing.
• Start from the last noble gas (Key: right-most column
of elements) prior to your element and put it in
brackets.
• Then, simply write the new configuration after it.
– Example: Sodium is [Ne] 3s1
• NOTE: Noble gases themselves can still be written in
shorthand. Just use the previous noble gas and go
from there. Helium does NOT have a shorthand
configuration.
Shorthand Notation
Element
Electron Configuration
Shorthand Notation
Hydrogen
1s1
--
Helium
1s2
--
Lithium
1s22s1
[He]2s1
Beryllium
1s22s2
[He]2s2
Boron
1s22s22p1
[He]2s22p1
Carbon
1s22s22p2
[He]2s22p2
Nitrogen
1s22s22p3
[He]2s22p3
Oxygen
1s22s22p4
[He]2s22p4
Fluorine
1s22s22p5
[He]2s22p5
Neon
1s22s22p6
[He]2s22p6
Exceptions
• Unfortunately, there are some exceptions to the
electron configuration rule. Copper and Chromium are
two good examples of this. Try diagramming them.
– Cr: 1s2 2s22p6 3s23p64s23d4
– Cu: 1s2 2s22p6 3s23p64s23d9
• Contrary to what you may have come up with, in reality
their configurations are:
– Cr: 1s2 2s22p6 3s23p64s13d5
– Cu: 1s2 2s22p6 3s23p64s13d10
• The reason for this is that filled sublevels are the most
stable. Half-filled sublevels are not as stable as filled,
but more stable than others.
The Full Hotel
7s 7p
6s 6p 6d
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
Quantum Numbers
• Because the current model atom is three
dimensional and based on mathematics, we use a
series of descriptions (numbers) to denote
electrons.
• This system allows us to combine Electron
Configuration and Orbital Notation into one.
• The descriptions are called quantum numbers, and
they include the principal quantum number (n).
– KEY: Think of these as mathematical code language for
stuff like “3d10.” You already know this!
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html AND FOLLOWING
Quantum Numbers
• n = Principal Quantum Number
– Indicates energy.
• l = Angular Quantum Number
–
•
•
•
•
Indicates sublevel:
0=s
1=p
2=d
3=f
• ml = Magnetic Quantum Number
– Indicates orbital.
• ms = Spin Quantum Number
– Indicates particular electron by its spin (more to come).
Quantum Number Rules
• n is from 1-7 (you knew that already).
• l is from 0 to n-1.
– This should make sense to you because:
• On n=1, only s (0) sublevels appear.
• On n=4, s (0), p (1), d (2), and f (3) sublevels appear.
• ml is from –l to +l.
– Each ml value represents a different orbital.
• When l = 1, we’re talking about a p sublevel.
• In that case, ml can be either -1, 0, or 1, each representing one of
the three “dumbbells” in space.
• ms (spin) is either -½ or ½.
– In short, one direction or another.
– This indicates a single electron.
n=2
l=1
ml = 1
ms = ½
Breaking Down The Code
So you could be
talking about 2s
and its single
orbital…
…or you could be
talking about 2p
and its three
orbitals.
y
y
2p
z
x
e-
ee-
e-
x
z
IfFinally,
nm
=give
2,
then
lthat
=within
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=1.+½.
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So
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Even
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= quantum
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identifies
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sublevelto
onsix
one
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level.
quantum numbers
n = Thus,
2.orbitals
which
still contains
twofour
electrons.
always indicate just one electron.
Breaking Down The Code
• If I described something as having these quantum
numbers, what am I really saying?
•
•
•
•
n=3
l=2
ml = 2
ms = ½
• Translated:
•
•
•
•
n = 3 (so third energy level)
l = 2 (so it’s a d sublevel – we’re talking about 3d)
ml = 2 (so one particular 3d orbital)
ms = ½ (so one electron in one orbital in 3d)
Putting It Into Code
• Alternatively, what if I wanted to refer to two
electrons in the 2px orbital? How would it be
written in quantum numbers?
– n = 2 (that’s an easy one)
– l = 1 (because when l = 1, that’s code for p)
– ml = -1 (because we just want one p orbital/dumbbell)
• For our purposes, we could have also picked 0 or 1.
– ms is not needed because we’re talking about two
electrons.
Quantum Number Practice
• What combinations of l and m can there be when n =
3?
– l can be 0, 1, or 2 (reflecting s, p, or d orbitals)
– m can be -2, -1, 0, 1, or 2 (reflecting the orientation of either
one s orbital, three p orbitals, or all five d orbitals.
• Describe the 3p sublevel using quantum numbers.
– n=3, l=1, m=-1, 0, 1
• How many electrons am I describing if I indicate
quantum numbers of n=4, l=2, m=2?
– n indicates a set of 2n2 electrons (32).
– l indicates a d sublevel, so that cuts us down to 10 electrons.
– m indicates the orientation of one of the d orbitals, so 2 e-.
Cracking the Code
Another way to look at it…
Quantum Numbers
n=3
l=2
ml = -2
ms = +½
____
____
____
____
3d
Orbital Notation
____
Quantum Numbers Summary Image
Quantum Number Practice
• Quantum Number Practice Worksheet
– 13 is a CHALLENGE.
Summary Table – Quantum Numbers
Principal
Quantum
Number (n)
Possible Angular
Possible Magnetic Quantum
Numbers (m)
Quantum Numbers (l)
0 (s)
0
(up to 1 orientation for s)
2
0, 1 (s, p)
-1, 0, 1
(up to 3 orientations for p)
3
0, 1, 2 (s, p, d)
-2, -1, 0, 1, 2
(up to 5 orientations for d)
4
0, 1, 2, 3 (s, p, d, f)
-3, -2, -1, 0, 1, 2, 3
(up to 7 orientations for f)
1
The “Rules”
• We’ve already learned one “rule:”
– Aufbau Principle – non-excited electrons fill energy
levels from the lowest level up.
• Now let’s learn the other two:
– Pauli Exclusion Principle
– Hund’s Rule
Pauli Exclusion Principle
• No more than two
electrons can occupy the
same orbital (not sublevel,
though).
• Each must have opposite
spins within a magnetic
field. This is the fourth
quantum number – ms.
• +½
• -½
Wolfgang Pauli
PEP and Orbital Notation
• In electron configuration,
there is no indication of spin.
• In the Hog Hotel, electrons in
the same orbital were
illustrated by opposite-facing
hogs.
• In orbital notation, scientists
use up and down arrows to
describe electrons’ opposite
spins.
↿⇂
Chemistry versus Hogs
Hog Hotel
Chemistry
Fill floors from the ground
up. Hogs hate to go up
stairs if they can avoid it.
Aufbau Principle – Fill
energy levels from lowest to
highest.
Only two hogs per room.
They face opposite ways.
One hog per room until
forced to put two in. Hogs
hate to go up stairs.
Chemistry versus Hogs
Hog Hotel
Chemistry
Fill floors from the ground
up. Hogs hate to go up
stairs if they can avoid it.
Aufbau Principle – Fill
energy levels from lowest to
highest.
Only two hogs per room.
They face opposite ways.
Pauli Exclusion Principle –
Only two electrons per
orbital. Electrons spin
opposite ways.
One hog per room until
forced to put two in. Hogs
hate to go up stairs.
Hund’s Rule
• Two electrons can occupy
a given orbital only after
all other orbitals have
been filled with one.
• In the Hog Hotel, Hund’s
rule was illustrated by the
“singles only” concept.
Friedrich Hund
Hund’s Rule
• You can also think of it
with a plain English
example:
– Imagine a school bus
being filled with students
who all dislike each
other.
http://www.instructables.com/image/FUPUHD6FGH3UFIV/Removing-School-Bus-Seats.jpg
Hund’s Rule
• Each student will take a seat by himself until
there are no free seats left. Only then will
they pair.
Chemistry versus Hogs
Hog Hotel
Chemistry
Fill floors from the ground
up. Hogs hate to go up
stairs if they can avoid it.
Aufbau Principle – Fill
energy levels from lowest to
highest.
Only two hogs per room.
They face opposite ways.
Pauli Exclusion Principle –
Only two electrons per
orbital. Electrons spin
opposite ways.
One hog per room until
forced to put two in. Hogs
hate to go up stairs.
Hund’s Rule – One electron
per orbital until forced to
put two in.
Hund’s Rule
• Let’s explain Hund’s Rule with an example:
Oxygen.
• Oxygen is atomic number 8, so it has 8 electrons.
O
•
•
•
•
Electrons
Left:
80
7
6
3
2
1
5
4
____ ____ ____ ____ ____
1s
2s
2p
First, fill the 1s shell with electrons.
Then, fill the 2s shell with electrons.
Then, begin filling the 2p shell, but only put one electron in
each orbital (keep ‘em all spinning the same way).
Finally, place a second electron in each shell.
Putting It All Together
• Using the three rules (Aufbau Principle, Pauli
Exclusion Principle, Hund’s Rule), let’s draw some
electron diagrams!
• Let’s start with Helium:
He
____
1s
1s2
• Notice that Helium has a full 1s shell (like a full
first floor), with no other electrons occupying any
other energy level.
• This comes into play on the next slide.
Element
Electron
Configuration
Orbital Notation
Shorthand
Notation
Li
1s22s1
____
____ ____
1s
2s
____ ____
2p
[He]2s1
Be
1s22s2
____
____ ____
1s
2s
____ ____
2p
[He]2s2
B
1s22s22p1
____
____ ____
1s
2s
____ ____
2p
[He]2s22p1
C
1s22s22p2
____
____ ____
1s
2s
____ ____
2p
[He]2s22p2
N
1s22s22p3
____
____ ____
1s
2s
____ ____
2p
[He]2s22p3
O
1s22s22p4
____
____ ____
1s
2s
____ ____
2p
[He]2s22p4
F
1s22s22p5
____
____ ____
1s
2s
____ ____
2p
[He]2s22p5
Ne
1s22s22p6
____
____ ____
1s
2s
____ ____
2p
[He]2s22p6
Putting It All Together
• Finally, using all that we’ve learned, let’s do
the following:
• Complete the Electron Configurations and
Orbital Notation sheet.
• Complete the Electron Configuration
Evaluation Worksheet
– If you can do all this, you’re ready.