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Today is Thursday,September 24th, 2015
In This Lesson:Unit 2
Electrons, Orbitals, and Atomic Model
History(Lesson 1 of 4)
Pre-Class:In your notebooks, draw a picture of electrons moving around the atom’s
nucleus. Include arrows to show direction.
You’re going to put this on a whiteboard shortly, so grab a SMALL paper towel.
Stuff You Need:Periodic TablePaper Towel
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. 127 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.
“Like Howitzer shells bouncing off of tissue
paper.”
Ernest Rutherford
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).
http://www.pkc.ac.th/kobori/Assets/ChemistryMahidol1/www.il.mahidol.ac.th/course/ap_chemistry/atomic_structure/picture/bild_goldstein.jpg
Eugen Goldstein
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
Schrödinger’s Atom (1926)• In 1923, Louis de Broglie discovered that
particles as small as electrons have some wave-like properties (as opposed to strictly particle-like).– 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 three-dimensional
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.
http://1.bp.blogspot.com/_GVA115I1I8Y/TT6_AHLks3I/AAAAAAAABWo/uvD4LGMKRgY/s1600/Broglie_Big.jpgLo
uis
de B
rogl
ieEr
win
Sch
rödi
nger
http://upload.wikimedia.org/wikipedia/commons/thumb/2/26/Erwin_Schrödinger.jpg/220px-Erwin_Schrödinger.jpg
Chadwick and the Neutron
• Chadwick discovered the neutron in 1932 and won the Nobel Prize three years later for it.
http://www.nobelprize.org/nobel_prizes/physics/laureates/1935/chadwick.jpghttp://www.dnahelix.com/jimmy/jnmov_jn_ext_shrinkray.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.
“One cannot simultaneously determine both the position
and momentum of an electron.”
http://www.wired.com/images_blogs/underwire/2012/09/heisenberg_660.jpg
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=1Energy 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=1Energy 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.
My Sophomore Year of College
e-e-
e-
e-
e-
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.”
s SublevelO
rbita
l
e- e-
s Sublevels
• Shape: Sphere• Appears: n=1 and
above.• # of Orbitals: 1• Capacity: 2 e-
p SublevelO
rbita
l
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-
d SublevelO
rbita
l
e- e- e- e-e- e-
e-e- e
-e
-
d Sublevels
• Shape: Double Dumbbells (4) and Dumbbell Doughnut
• Appears: n=3 through n=6.
• # of Orbitals: 5• Capacity: 10 e-
f SublevelO
rbita
l
e- e- e- e-e- e-
e-e-
e-e-
e-
e-
e-
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?
1s
2s 2p
3s 3p 3d
4s 4p 4d 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 orbital start to appear?• n=4
Summary TableEnergy Level
(n) Sublevels Orbitals Per Sublevel
Electrons Per Sublevel
Electrons Per Energy Level
(2n2)
1 s 1 2 2
2 sp
13
26 8
3spd
135
26
1018
4spdf
1357
26
1014
32
Floor NumberType of
Rooms/Suites on Floor
Rooms per Type of Room/Suite
Capacity of Each Type of Room/Suite
Capacity of Each Floor
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 Tables1
s2
s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
5d
6d
4f
5f
d
f
ps
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.– 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 Nes1
s2
s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
5d
6d
4f
5f
d
f
psNe: 1s2 2s2 2p6
Electron ConfigurationElement Electron Configuration
Hydrogen 1s1
Helium 1s2
Lithium 1s22s1
Beryllium 1s22s2
Boron 1s22s22p1
Carbon 1s22s22p2
Nitrogen 1s22s22p3
Oxygen 1s22s22p4
Fluorine 1s22s22p5
Neon 1s22s22p6
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 Ws1
s2
s2
p1 p2 p3 p4 p5 p6
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
d1 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
5d
6d
4f
5f
d
f
ps
W: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f14 5d4
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 (right-most column) 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 NotationElement 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.
1s2s 2p
3s 3p 3d4s 4p 4d 4f
5s 5p 5d 5f6s 6p 6d
The Full Hotel
7s 7p
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.
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-.
Quantum Number Practice
• Quantum Number Practice Worksheet– 13 is a CHALLENGE.
Summary Table – Quantum Numbers
Principal Quantum
Number (n)
Possible Angular Quantum Numbers (l)
Possible Magnetic Quantum Numbers (m)
1 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)
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.
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.
Chemistry versus Hogs
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.
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.
Chemistry versus Hogs
Hund’s Rule
• Let’s explain Hund’s Rule with an example: Oxygen.• Oxygen is atomic number 8, so it has 8 electrons.
O ____ ____ ____ ____ ____ 1s 2s 2p
Electrons Left:
8
Electrons Left:
7
Electrons Left:
6
Electrons Left:
5
Electrons Left:
4
Electrons Left:
3
Electrons Left:
2
Electrons Left:
1
Electrons Left:
0
• 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:
• 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.
He ____ 1s
1s2
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.
Homework
• In your textbooks, read pages 127-136.• Answer problems 27, 29-39 on page 149.
