1. Today is Wednesday, February 22nd, 2017
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.

2. 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.

3. 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.

4. Guiding Video
TED: George Zaidan and Charles Morton – The Uncertain Location of Electrons

5. 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.

6. 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.

7. 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.

8. In shorter terms… Electrons are important because: They create ions.
They lead to bonding.
They determine how atoms behave.

9. Thomson’s Atom (1897)
Called the Plum Pudding Model, as Thomson thought that electrons were like plums sitting in a positive pudding.
JJ Thomson

10. 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

11. 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

12. Rutherford’s Atom (1913)
After the Rutherford experiment, the atom model looked like this:
Looked like the Infinity Ward logo, but it’s wrong.

13. 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

14. 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…

15. 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

16. 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:

17. 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.
Louis de Broglie
Erwin Schrödinger

18. Chadwick and the Neutron
Chadwick discovered the neutron in 1932 and won the Nobel Prize three years later for it.
James Chadwick

19. 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.

20. 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.

21. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
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.

22. 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.

23. 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!

24. 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.

25. 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.

26. Electron Energy Levels (Shells)
Energy Level 1  n=1
Energy Level 2  n=2
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.

27. n n is the Principal Quantum Number.
Energy Level 1  n=1
Energy Level 2  n=2
n is the Principal Quantum Number.
To determine how many electrons fit into a given energy level, use this formula:
Electrons = 2n2

28. 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.

29. 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.

30. 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.

31. My Freshman Year of College
I had the basic two bed/one roommate set up.
Also, my roommate was awful but that’s besides the point.

32. My Sophomore Year of College
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.”

33. s Sublevel
Orbital e-

34. s Sublevels Shape: Sphere Appears: n=1 and above. # of Orbitals: 1
Capacity: 2 e-

35. p Sublevel
Orbital e- e- e-

36. p Sublevels Shape: Dumbbell (3) Appears: n=2 and above.
# of Orbitals: 3 (x, y, z)
Capacity: 6 e-

37. d Sublevel
Orbital e- e- e- e- e-

38. d Sublevels Shape: Double Dumbbells (4) and Dumbbell Doughnut
Appears: n=3 through n=6.
# of Orbitals: 5
Capacity: 10 e-

39. f Sublevel
Orbital e- e- e- e- e- e- e-

40. f Sublevels Shape: Flowers…and stuff. Appears: n=4 through n=5.
# of Orbitals: 7
Capacity: 14 e-

41. And the “hotel” as a whole?4p 4d 4f 3s 3p 3d 2s 2p 1s

42. 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.

43. 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.

44. 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

45. Electrons Per Sublevel Electrons Per Energy Level (2n2)
Summary Table
Energy Level (n)
Sublevels
Orbitals Per Sublevel
Electrons Per Sublevel
Electrons Per Energy Level (2n2) 1 s 2 p 3 6 8 d 5 10 18 4 f 7 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

46. 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.

47. 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.

48. Electron Configuration Tables1 s2 1s s2 p1 p2 p3 p4 p5 p6 2s 2p p s 3s d1 d2 d3 d4 d5 d6 d7 d8 d9
d10 3p 4s 3d d 4p 5s 4d 5p 6s 5d 6p 7s 6d d1 f1 f2 f3 f4 f5 f6 f7 f8 f9
f10
f11
f12
f13
f14 f 5d 4f 6d 5f

49. 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:

50. 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.

51. 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 s p

52. 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.

53. 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).

54. Electron Configuration for Nes1 s2 1s s2
Ne: 1s2 2s2 2p6 p1 p2 p3 p4 p5 p6 2s 2p p s 3s d1 d2 d3 d4 d5 d6 d7 d8 d9
d10 3p 4s 3d d 4p 5s 4d 5p 6s 5d 6p 7s 6d d1 f1 f2 f3 f4 f5 f6 f7 f8 f9
f10
f11
f12
f13
f14 f 5d 4f 6d 5f

55. 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

56. Use this for the next slide’s questionsp1 p2 p3 p4 p5 p6 2s 2p p s 3s d1 d2 d3 d4 d5 d6 d7 d8 d9
d10 3p 4s 3d d 4p 5s 4d 5p 6s 5d 6p 7s 6d d1 f1 f2 f3 f4 f5 f6 f7 f8 f9
f10
f11
f12
f13
f14 f 5d 4f 6d 5f

57. 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.

58. Electron Configuration for Ws1 s2
W: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f14 5d4 1s s2 p1 p2 p3 p4 p5 p6 2s 2p p s 3s d1 d2 d3 d4 d5 d6 d7 d8 d9
d10 3p 4s 3d d 4p 5s 4d 5p 6s 5d 6p 7s 6d d1 f1 f2 f3 f4 f5 f6 f7 f8 f9
f10
f11
f12
f13
f14 f 5d 4f 6d 5f

59. On your worksheets…
Try the first page of the worksheet labeled, “Electron Configurations Orbital Notation.”
We’ll do the first one (Mg) together.

60. 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

61. 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

62. 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?

63. 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.

64. Electron 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

65. 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.

66. The Full Hotel 7s 7p 5s 5p 5d 5f 6s 6p 6d 1s 2s 2p 3s 3p 3d 4s 4p 4d

67. 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!
AND FOLLOWING

68. 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).

69. 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.

70. Breaking Down The Code e- e- 2p e- e- e- e- e- e- e- 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. 2p z y x z y x
e- e- e- e- e- e- e- e- e-
So then I add that ml = 1. That identifies just one orbital, which still contains two electrons.
Even after providing two quantum numbers, there are still up to six electrons in three orbitals remaining.
That identifies just one electron within one orbital within one sublevel on one energy level. Thus, four quantum numbers always indicate just one electron.
Finally, I add that ms = +½.
Suppose I were to give you just one quantum number: n = 2.
So then I tell you l = 1.
If n = 2, then l = 0 or l = 1.

71. 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)

72. 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.

73. 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-.

74. Cracking the Code Another way to look at it…
Quantum Numbers
n = 3
l = 2
ml = -2
ms = +½
Quantum Numbers
n = 3
l = 2
ml = -2
ms = +½
____ ____ ____ ____ ____ 3d
Orbital Notation
____ ____ ____ ____ ____ 3d
Orbital Notation

75. Quantum Numbers Summary Image

76. Quantum Number Practice
Quantum Number Practice Worksheet
13 is a CHALLENGE.

77. Summary Table – Quantum Numbers
Principal Quantum Number (n)
Possible Angular Quantum Numbers (l)
Possible Magnetic Quantum Numbers (m) 1
0 (s)
(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)

78. 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

79. 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

80. 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. ↿ ⇂

81. 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.

82. 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.

83. 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

84. 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.

85. Hund’s Rule
Each student will take a seat by himself until there are no free seats left. Only then will they pair.

86. 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.

87. Hund’s Rule Let’s explain Hund’s Rule with an example: Oxygen.
Oxygen is atomic number 8, so it has 8 electrons.
Electrons Left: 1
Electrons Left: 3
Electrons Left: 6
Electrons Left: 7
Electrons Left: 2
Electrons Left: 8
Electrons Left:
Electrons Left: 5
Electrons Left: 4
O ____ ____ ____ ____ ____
1s s p
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.

88. 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

89. Electron Configuration
Element
Electron Configuration
Orbital Notation
Shorthand
Notation Li
1s22s1
____ ____ ____ ____ ____
1s s p
[He]2s1 Be
1s22s2
[He]2s2 B
1s22s22p1
[He]2s22p1 C
1s22s22p2
[He]2s22p2 N
1s22s22p3
1s s p
[He]2s22p3 O
1s22s22p4
[He]2s22p4 F
1s22s22p5
[He]2s22p5 Ne
1s22s22p6
[He]2s22p6

90. 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.