1. Ions, Electrons and Periodicity

2. Why do Ions form?
In order to understand why ions form, you need to know how Electrons are arranged!
Where are electrons located???

3. Quick Review of Electron Configuration
Electrons are found in the electron cloud, but they’re not just placed anywhere.
Electrons are found on different energy levels around the nucleus, and each energy level has a maximum number of electrons that it can hold
1st level = 2 2nd level = 8
3rd level = th level = 32
Keep in mind, these are the maximum they can hold and not always the most stable configuration.

4. So how do we know when its stable?
Electrons have a certain amount of energy to them which describes how far away from the nucleus they tend to be
Bohr’s model put these into “shells” or “energy levels”

5. Which energy level this is depends on the atom
The electrons on the outermost energy level are called valence electrons.
Which energy level this is depends on the atom
The group number on the Periodic Table tells us the number of valence electrons for those atoms!
Groups 13-18, just take off the front “1”, so group 13 has 3 valence electrons, etc.

6. Octet Rule
An atom is stable when it has a full valence energy level, called the octet rule
Most atoms want to have 8 valence electrons (hence octet), some exceptions in just a sec
They can do this by either gaining more electrons to fill their current valence shell
Or they can lose electrons to empty the current shell and make the lower energy level the valence shell.

7. 5 6 Take the element Boron! How many Protons? How many Neutrons?
How many Electrons?
Where will they go? 5 6

8. 10 10 Take the element Neon! How many Protons? How many Neutrons?
How many Electrons?
Where will they go? 10 10

9. Unlocking the secrets of bonding, one electron at a time!
So, Atoms want to always be stable, but how can we determine what they will do?
It’s all based on their number of valence electrons.
Lets go back to Boron!

10. Look at Boron in it’s normal state. Is it a stable atom?
Not stable! – NO OCTET!
So how would it become stable? What is easier and quicker for this atom?
Gain 5 more or lose 3?
And the winner is…..
STABLE 1st ENERGY LEVEL

11. How many valence electrons?
Boron only ended up with 2 valence electrons after becoming an ion. Why is that ok?
The first 5 elements actually follow the Duet Rule because they prefer using their first energy level for stability, which is filled with 2 electrons instead of 8.

12. Oxidation Number
By determining how Boron becomes stable, You figured out the oxidation number for Boron!
Oxidation number is the number of electrons an atom gains or loses when bonding.
Get out your best colored pen and get ready to write these simple steps to determine the oxidation number!
Ox-a-what?!

13. All the steps you need
1) Find the element and figure out how many electrons it normally would have. (Remember it’s the same as Protons)
2) Ask yourself, how many valence electrons does this atom have? (Those are the ones on the outer energy level)
3) Would it be quicker to add and get 8 valence electrons, or lose them and go down to a stable lower level? (I don’t think you need a hint here, it’s literally that easy)

14. Nitrogen B) Fluorine C) Phosphorus D) Chlorine
Practice
Nitrogen B) Fluorine
C) Phosphorus D) Chlorine
sph

15. What did you notice?
What did you see about atoms that are in the same groups?
Remember that these elements in the same columns have similar electron configurations!!!!
SO…..Every atom in the same group has the same electron configuration! Ready for a big helping hint??? Get out your Periodic Table and check this out!

17. Electron Configuration

18. Why we Study Electrons
Knowing the position of electrons helps us with many topics:
Why chemical reactions occur
Why some atoms are more stable than others
Why some elements react with only certain atoms
Remember… An atom is made up of
Protons- located in the nucleus
Neutrons- located in the nucleus
Electrons- located in the electron cloud!

19. “Electron Cloud” - Energy Levels ( 1 – 7)
- Sublevels (S,P,D,F – Shapes)
- Orbitals (Specific Pathways – X/Y/Z axis)

20. - S Sublevel – Think Sphere
What the orbitals look like…
- S Sublevel – Think Sphere
- P Sublevel – Figure 8
P Sublevel- X Axis
P Sublevel- Y Axis
Nucleus
P Sublevel- Z Axis
S Sublevel

21. Orbital Diagrams
Orbital Diagrams are Visual illustrations that show the order that electrons will follow as they fill in the electron cloud.
We start from the bottom and work your way up. That symbolizes starting on the first energy level and working outward.

22. There are Rules for Drawing Orbital Diagrams!
When filling in an orbital diagram, it’s important to follow three rules.
You need to know these by name and what each specifically says, so get ready to write them!

23. Rule #1 Aufbau Principle
Electrons must fill the lowest sublevel available. (Letters - S,P,D,F)
You cannot skip any sublevels.
It’s like the Alphabet!

24. Rule #2 Pauli Exclusion Principle
Each orbital (each box on the diagram) can hold only TWO electrons, and the electrons have opposite spins. This is noted by the oppositely pointed arrows!

25. The 3rd and Final Rule WRONG RIGHT Hund’s Rule
Within a sublevel, place one electron in each orbital with the same spin before pairing them with a second electron. - “Empty Seat Rule”
WRONG
RIGHT

26. A Few Things to Remember…
You are drawing arrows, and they represent electrons. (Neon has 10 electrons so you will need to draw 10 arrows.)
Divided up into different sublevels containing a specific number of orbitals. KNOW THIS INFO!
- The S sublevel has only 1 orbital= total of 2 e-
- The P sublevel has 3 orbitals = total of 6 e-
- The D sublevel has 5 orbitals = total of 10 e-
- The F sublevel has 7 orbitals = total of 14 e-

27. Practice!
Fill in the orbital diagram for Magnesium

28. Electron Configuration
An atom’s electron configuration…
Describes the location of electrons within the atom
Identifies the shape of the electron clouds
- regions where the electrons are held.
Uses numbers and letters to describe electrons’ location and which electron cloud it is part of.

29. 1s2 The Format: The “s” tells you the electron’s cloud shape.
In this case it’s a spherical shaped cloud. It is called the s sublevel. Sublevel = shape
1s2
The “2” simply tells you how many electrons are in this cloud.
The “1” tells you how far from the nucleus the electrons can go.
In this case 2 electrons are creating the cloud.
In this case, its in the 1st energy level, which is the closest level to the nucleus.
Note: Every orbital can only hold 2 electrons. They must have opposite spins.

30. Connection b/w Orbital Diagram and Electron Configuration1s 2s 2p
Electron Configuration
1s2 2s2 2p4

31. Noble Gas Configuration Steps:
Shorthand Configuration
This is a shorter way to do e- configurations
Steps…
1) Find the noble gas (group 18) that precedes and comes before the element you are doing.
2) To begin, put the symbol of that noble gas in [brackets]
3) Starting with that noble gas, continue the electron configuration until you get to the element in question.

32. 1s2 2s2 2p3 Nitrogen: [He] 2s2 2p3 Nitrogen: Valence Electrons
Longhand Configuration
Nitrogen:
1s2
2s2
2p3
Valence Electrons
Core Electrons
Shorthand Configuration
Nitrogen: [He] 2s2 2p3

33. Example - Sulfur
[Ne]
3s2
3p4

34. Example - Germanium
[Ar]
4s2
3d10
4p2

35. Practice these in shorthand
Tin (Sn) Uranium (U)

36. Draw this! This will help!! It is a way to remember
the order in which electrons
fill their orbitals.
When you reach the far
Left side, you reach a
“wall” and must go back to
The right hand side.