1. Order of orbitals (filling) in multi-electron atom
Aufbau Principle: an electron occupies the lowest-energy orbital that can receive it.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

2. Outermost subshell being filled with electrons

3. Why are d and f orbitals always in lower energy levels?
d and f orbitals require LARGE amounts of energy
It’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy
This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER!

4. in the orbital or subshell
Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.
number of electrons
in the orbital or subshell
1s1
principal quantum
number n
angular momentum
quantum number l
Orbital diagram
1s1 H

5. Shorthand Notation
Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].
Step 2: Find where to resume by finding the next energy level.
Step 3: Resume the configuration until it’s finished.

6. What is the electron configuration of Mg?
Mg 12 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s2
= 12 electrons
Abbreviated as [Ne]3s2
[Ne] 1s22s22p6
What are the possible quantum numbers for the last (outermost) electron in Cl?
Cl 17 electrons
1s < 2s < 2p < 3s < 3p < 4s
1s22s22p63s23p5
= 17 electrons
Last electron added to 3p orbital
n = 3
l = 1
ml = -1, 0, or +1
ms = ½ or -½

7. Hund’s Rule: the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli exclusion principle.
F 9 electrons
Ne 10 electrons
C 6 electrons
O 8 electrons
N 7 electrons
Ne 1s22s22p6
C 1s22s22p2
N 1s22s22p3
O 1s22s22p4
F 1s22s22p5

8. Exceptions to the Aufbau Principle
Remember d and f orbitals require LARGE amounts of energy
If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)
There are many exceptions, but the most common ones are
d4 and d9
For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.

9. Exceptions to the Aufbau Principle
d4 is one electron short of being HALF full
In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.
For example: Cr would be [Ar] 4s2 3d4, but since this ends exactly with a d4 it is an exception to the rule. Thus, Cr should be [Ar] 4s1 3d5.
Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d.

10. Keep an Eye On Those Ions!
Tin
Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of the highest energy level, not the highest energy orbital!