1. Electron Orbital Diagrams

2. Placing Electrons in Orbitals
Every orbital may hold up to 2 electrons
These 2 electrons have opposite spins: an up spin and a down spin
You can determine how many electrons are possible for an energy level by 2n2
If n=1, only 2 electrons are possible
If n=2, 8 electrons are possible
If n=3, 18 electrons are possible

3. Placing Electrons in Orbitals
There are rules for placing electrons in orbitals.
Hund’s Rule
Aufbau Principle
Pauli’s Exclusion Principle
There is also an order which we follow in placing electrons in an orbital energy diagram.

4. Orbital Diagrams The number is the energy level (n value)
The letter is the orbital type or the sublevel type
The number of lines beside the sublevel is how many orbitals there are on that sublevel
Do you notice any pattern?

5. Orbital Diagrams
For every n value, the s orbital is the lowest energy.
ns < np < nd < nf
It looks like a pattern starts, but then it is broken as 3d is higher in energy than the 4s.
So you memorize this!
How do you place the electrons in these orbitals? Let’s try.

6. Orbital Diagrams
For ions like F- or Ca2+, their orbital diagram will be just like a Noble Gas.
This is an empirical observation called the Octet Rule:
Atoms tend to lose, gain, or share electrons in order to have 8 valence electrons.
What are valence electrons?
Valence electrons are the electrons in the outermost energy level (highest n value) which are furthest from the nucleus.

7. Orbital Diagrams
You can easily count the number of valence electrons from the Periodic Table for Groups 1-2 and (Representative or Main Block Groups)
This means that you can predict the ion charges for these groups as well!

8. Orbital Diagram Exceptions
There are 19 exceptions to the rules when drawing Orbital Diagrams!
You are only responsible for knowing 5: Cr, Mo, Cu, Ag, Au.
These 5 elements all do the same thing: as the (n-1)d orbital is VERY close in energy to the ns orbital, a ns electron is “promoted” to a “higher-energy” (n-1)d orbital.
This will either fill or half-fill the d orbitals.

9. Orbital Diagram Exceptions
Because the d orbitals are filled or half-filled, the atom gains stability overall.
These “exceptions” are thus the ground-state for these atoms.
If you were to draw other exceptions to the rules, these would be classified as “excited” states.

10. Electron Configurations and Noble Gas Configurations

11. Electron Configurations
Orbital energy diagrams tell us what electrons/orbitals an element fills in its ground state (most stable state).
But they are tedious to draw and they must be in a figure.
So how can we get the same info but be able to write it in a sentence?
Electron configurations let us do this.
All we do is translate the vertical orbital diagram into a horizontal line. (well, almost all)

12. Figure: UN
Title:
Filling the orbitals
Caption:
Follow the aufbau principle and Hund's rule in placing electrons in the appropriate orbitals of an atom.

13. Electron Configurations
What about the electron configuration for Sr?
When you write the correct electron configuration for the ground state of Sr, notice that you group the n values together.
This is different from how you wrote the orbital diagram which is strictly based on filling order.
Why do we write the electron configuration this way?
Convention and because that way the valence electrons are at the end of the configuration: and valence electrons are lost first!

14. Noble Gas Configurations
But even electron configurations get long and tedious to write.
Write the electron configuration for gold!
How can we make this shorter?
Well, the Noble Gases have filled s and p subshells, so they are good reference elements.
So we write Noble Gas configurations, which are based on the LAST Noble Gas BEFORE the element we are working with.

15. Figure: UN
Title:
Shorthand notation for electron configuration
Caption:
To simplify the writing of lengthy electron configurations, the symbol of a noble gas may be substituted for the appropriately filled shells and orbitals.

16. Periodic Table and Electron Configurations
Guess what?
The Periodic Table is laid out in order of increasing Atomic Number.
It is also organized by the order in which the electrons fill! (Except for the exceptions.)
Thus, we have the s-block, the p-block, the d-block, and the f-block.

17. Figure: 05-18
Title:
Orbital-filling blocks of the periodic table
Caption:
Figure 5.18  Blocks of the periodic table, corresponding to filling the different kinds of orbitals. Beginning at the top left and going across successive rows of the periodic table provides a method for remembering the order of orbital filling: 1s --> 2s --> 2p --> 3s --> 3p --> 4s --> 3d --> 4p, and so on.

18. Electron Configurations and Periodic Properties

19. Periodic Properties & Valence Electrons
Members of a group have similar chemical properties. Why?
They have the same number of valence electrons, or their outermost electron configurations are the same.
And valence electrons are responsible for chemical properties.

20. Periodic Properties & Valence Electrons
Valence electrons are outer shell electrons and so are furthest from the nucleus (generally).
Thus they are not as tightly held to the nucleus.
So they are removed more easily than core electrons (inner shell electrons).

21. Periodic Properties & Valence Electrons
Also, the core electrons with their negative charges block or insulate or "shield" the valence electrons from feeling the full nuclear charge.
So the shielded valence electrons don't feel the full nuclear charge, Z, they instead feel what we call Zeff, or the effective nuclear charge.
Think of picking up iron filings or paper clips with a magnet, eventually the magnet can’t attract any more filings as it is “shielded”.

22. Figure: 05-16
Title:
Effective nuclear charge (Z)
Caption:
Figure 5.16  The origin of electron shielding and Zeff . Outer electrons are attracted toward the nucleus by the nuclear charge but are pushed away by the repulsion of inner electrons. As a result, the nuclear charge actually felt by outer electrons is diminished, and we say that the outer electrons are shielded from the full charge of the nucleus by the inner electrons.

23. Periodic Properties & Valence Electrons
Members of the same group have the same number of valence electrons and have similar properties as those valence electrons behave similarly.
For example, all of the Alkali Metals have 1 valence electron, with an outer electron configuration of ns1.
This is why Alkali metals are so reactive, their 1 valence electron is easily removed in chemical rxns.
And what is their e- configuration then?

24. Periodic Properties & Valence Electrons
Alkaline Earth metals all have an ns2 outer electron configuration, or they have 2 valence electrons.
What is the outer electron configuration for halogens?

25. Atomic Radii
But chemical reactivity is not the only periodic property or trend which we can see in the Periodic Table.
Another important Periodic Trend or Property is the atomic radii, or the radius of an atom, or 1/2 the diameter of the atom.
We may measure the radius of an atom.
It is also common to measure the distance between 2 nuclei in a diatomic element.
Then the atomic radius would be half of this distance.

26. Figure: UN
Title:
Atomic radii
Caption:
The atomic radius of an atom is taken to be one-half the distance between the nuclei of two of the same atoms bound to one another. Atomic radii for individual atoms can then be used to predict the bond length between two different atoms.

27. Atomic Radii
There are 2 trends in atomic radii that you can explain and see from the Periodic Table:
As you go down a Group, the atomic radius increases.
Now you are adding an energy shell which by definition is further from the nucleus!

28. Atomic Radii
As you go across a Period from left to right, the atomic radius of the elements DECREASES.
This is because you are adding protons so Zeff is increasing. As Zeff increases, the outer electrons are pulled closer to the nucleus, so the radius shrinks.
This is particularly true as you add s and p electrons in the same energy level. They don't shield each other very well, so Zeff increases as you add protons.

29. Atomic Radii
What about the different orbital types, do they have different attractions, or different Zeff?
For example, for n = 4, there are 4s, 4p, 4d, and 4f orbitals. They don’t have the same energy, do they have different Zeff?
They do!
In general, as the subshell energy increases on the same principal energy level, the Zeff decreases.
So for the same n value, the s has the highest Zeff.

30. Figure: UN
Title:
Relationship between Zeff and orbital type
Caption:
An electron's attraction to the nucleus is correlated to the energy level and orbital type occupied by the electron.

31. Figure: 05-20
Title:
Periodicity of atomic radius, atomic number, and Zeff
Caption:
Figure 5.20  Plots of Zeff for the highest-energy electron and atomic radius versus atomic number. As Zeff increases, the valence-shell electrons are attracted more strongly to the nucleus, and the atomic radius therefore decreases.

32. Figure: 05-19
Title:
Atomic radii of the elements
Caption:
Figure 5.19  Atomic radii of the elements in picometers.