This document discusses electron configurations and principles including: 1) The Aufbau principle which states that electrons fill orbitals from lowest to highest energy levels. 2) Hund's rule which states that electrons fill orbitals singly before pairing up. 3) The Pauli exclusion principle which states that no two electrons in an atom can have the same set of four quantum numbers and that paired electrons must have opposite spins.

1. Atom video
http://www.youtube.com/watch?v=x
qNSQ3OQMGI&feature=share

2. Basic Principle:
electrons occupy
lowest energy
levels available

4. Aufbau Principle -- “Bottom Up Rule”

5. Electron spin
How could an orbital hold two electrons
without electrostatic repulsion?
 
Stern-Gerlach
Experiment

6. 1
1s
value of energy level
sublevel
no. of
electrons
spdf NOTATION
for H, atomic number = 1
spdf Notation
Orbital Box Notation
Arrows show
electron spin
(+½ or -½)
ORBITAL BOX NOTATION
for He, atomic number = 2
1s
2
1s 
2 ways to write electron configurations

7. Example:
Determine the electron configuration and orbital
notation for the ground state neon atom.
An orbital can contain a maximum of 2 electrons,
and they must have the opposite “spin.”
Pauli exclusion principle

8. Hund’s Rule -
Write the ground state configuration and the
orbital diagram for oxygen in its ground state

9. Outer electron configuration for the elements

10. 11 p+
12 n°
2e– 8e– 1e–
Na
8 p+
8 n°
2e– 6e–
O
4 p+
5 n°
Be
5 p+
6 n°
B
13 p+
14 n°
Al

11. Using the periodic table to know configurations
Period
1
2
3
4
5
6
7
Ne
Ar
Kr
Xe

12. Valence e’s for “main group” elements

13. Rules for Filling Orbitals
Bottom-up
(Aufbau’s principle)
Fill orbitals singly before doubling up
(Hund’s Rule)
Paired electrons have opposite spin
(Pauli exclusion principle)
Basic Principle:
electrons occupy
lowest energy
levels available

14. Identify examples of the following principles:
1) Aufbau 2) Hund’s rule 3) Pauli exclusion

15. Shorthand notation practice
Examples
● Aluminum: 1s22s22p63s23p1
[Ne]3s23p1
● Calcium: 1s22s22p63s23p64s2
[Ar]4s2
● Nickel: 1s22s22p63s23p64s23d8
[Ar]4s23d8 {or [Ar]3d84s2}
● Iodine: [Kr]5s24d105p5 {or [Kr]4d105s25p5}
● Astatine (At): [Xe]6s24f145d106p5
{or [Xe]4f145d106s26p5}
[Noble Gas Core] + higher energy electrons

16. Electron configuration for As

17. Note: Not written according to Aufbau, but grouping according to n

18. Orbital energy ladder
s
p
n = 2
s
d
p n = 3
f
s
d
p
n = 4
s n = 1
Energy

19. Phosphorus
Symbol: P
Atomic Number: 15
Full Configuration: 1s22s22p63s23p3
Valence Configuration: 3s23p3
Shorthand Configuration: [Ne]3s23p3



  
  
        
1s 2s 2p 3s 3p
Box Notation

20. Quantum numbers and orbital energies
Each electron in an atom has a unique set of quantum numbers to
define it
{ n, l, ml, ms }
n = principal quantum number
– electron’s energy depends principally on this
l = azimuthal quantum number
– for orbitals of same n, l distinguishes different shapes
(angular momentum)
ml = magnetic quantum number
– for orbitals of same n & l, ml distinguishes different
orientations in space
ms = spin quantum number
– for orbitals of same n, l & ml, ms identifies the two
possible spin orientations

21. Energy level Sublevel # of orbitals/sublevel
n = 1 1s (l = 0) 1 (ml has one value)
n = 2 2s (l = 0) 1 (ml has one value)
2p (l = 1) 3 (ml has three values)
n = 3 3s (l = 0) 1 (ml has one value)
3p (l = 1) 3 (ml has three values)
3d (l = 2) 5 (ml has five values)
n = principal
quantum
number
(energy)
l = azimuthal
quantum
number
(shape)
ml = magnetic
quantum
number
(orientation)
Quantum numbers and orbital energies
Each atom’s electron has a unique set of quantum numbers to define it {
n, l, ml, ms }

22. 22
Concept: Each electron in an atom has a unique set of quantum numbers
to define it
{ n, l, ml, ms }

23. Quantum numbers: unique set for each e-
s orbitals p orbitals d orbitals f orbitals
l = 0 l = 1 l = 2 l = 3
ml = 0 ml = -1, 0, 1 ml = -2, -1, 0, 1, 2 ml=-3,-2,-1,0,1,2,3
An s subshell A p subshell A d subshell An f subshell
One s orbital Three p orbitals Five d orbitals Seven f orbitals
For n=1 l=0 an s subshell (with 1 orbital)
For n=2 l=0,1 an s subshell and a p subshell (with 3 orbitals)
For n=3 l=0,1,2 an s subshell, a p subshell, a d subshell (with 5 orbitals)
For n=4 l=0,1,2,3 an s subshell, a p subshell, a d subshell, an f subshell (with 7 orbitals)

24. Electronic configuration of Br
1s2 2s22p6 3s23p63d10 4s24p5
[Ar] 3d104s24p5
[Ar] = “noble gas core”
[Ar]3d10 = “pseudo noble gas core”
(electrons that tend not to react)
Atom’s reactivity is determined by valence electrons
valence e’s in Br: 4s24p5
highest n electrons

25. Valence e- shells for
transition metals v. main group elements
d orbitals sometimes
included in valence shell
d orbitals not included
in valence shell
(pseudo noble gas cores)

26. Rule-of-thumb for valence electrons
Examples
● Sulfur: 1s22s22p63s23p4 or [Ne]3s23p4
valence electrons: 3s23p4
● Strontium: [Kr]5s2
valence electrons: 5s2
● Gallium: [Ar]4s23d104p1
valence electrons: 4s24p1
● Vanadium: [Ar]4s23d3
valence electrons: 4s2 or 3d34s2
Identify all electrons at the highest
principal quantum number (n)
Use on exams,
but recognize
limitations
Use Table 8.9
for online HW

27. Selenium’s valence electrons
Pseudo noble gas core includes:
 noble gas electron core
 d electrons (not very reactive)
Written for increasing energy:

28. Core and valence electrons in Germanium
Pseudo noble gas core includes:
 noble gas core
 d electrons
Written for increasing energy:

29. d-block: some exceptions to the Aufbau principle
Fig. 8.9: Use this table for online homework

30. Paramagnetic: atoms with unpaired electrons
that are weakly attracted to a magnet.
Diamagnetic: atoms with paired electrons
that are not attracted to a magnet.
Electron spin & magnetism
For the ground state oxygen atom:
spdf configuration:
orbital box notation:

31. Apparatus for measuring magnetic properties