1. Chapter 11 “Electron Configuration & Quantum Numbers”

2. Electron Arrangement in Atoms
OBJECTIVES:
Describe how to write the electron configuration for an atom.

3. Electron Arrangement in Atoms
OBJECTIVES:
Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

4. To find location & energy of e-, you must know the
QUANTUM NUMBERS
(click on the images below to watch an introductory video on Quantum Numbers)

5. QUANTUM NUMBERS
After watching the videos answer the question on paper below for extra credit:
a. What are the 4 quantum numbers?
b. How many sublevels are there?
c. What are the sublevels?
d. How many electrons are there per orbital?
e. What does Ms mean?
f. If n = 3, what does l = ?
g. How many orbitals do the following sublevels have? s, p, d, f.
h. What symbol is used for the electrons spin?

6. QUANTUM NUMBERS
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
1) Principal quantum number (n)
2) Angular momentum quantum number (l)
3) Magnetic quantum number (m)
4) Spin quantum number (s)

7. QUANTUM NUMBERS
“n” ; Principal Quantum # n = 1,2,3,4,... energy levels or periods on the periodic table max. # of e- per level = 2n2

8. 1 2 3 4 5 6 7
Period Number
Each row (or period) is the energy level. There are 7 energy levels on the PTE.

9. QUANTUM NUMBERS (cont)
“l”; Sublevels
s, p, d, f, g,...
# of sublevels per energy level = n
ex. n=1, l =s or n=3, l =s,p,d
**the levels/sublevels fill from low to high energy
Value of l 1 2 3
Type of orbital s p d f

10. Elements in the s - blocks
Alkali metals all end in s1
Alkaline earth metals all end in s2
really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons.

11. The P-block p1 p2 p3 p4 p6 p5

12. Transition Metals - d block
d = n - 1
Note the change in configuration. s1 d5
s1 d10 d1 d2 d3 d5 d6 d7 d8
d10

13. The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4.
d = n - 1 1 2 3 4 5 6 7 3d 4d 5d

14. f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13 F - block
Called the “inner transition elements”
f = n - 2 f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

15. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f, and are 2 less than the period number.
f = n - 2

16. Practice:
Identify the energy level and sublevel of where the following elements are located: Al K Fe U Hg

17. Practice:
Identify the energy level, sublevel and column of where the following elements are located:
Al: 3p1
K: 4s1
Fe: 3d6
U: 5f4
Hg: 5d1

18. Aufbau Principle:
- e- enter levels/sublevels of lowest energy 1st (1s, 2s, 2p, 3s, 3p, 4s, )

19. Practice:
Now identify the energy level and sublevel configuration of the following elements: Al K Fe U Hg

20. Practice:
Now identify the energy level and sublevel configuration of the following elements:
Al: 1s,2s,2p,3s,3p
K: 1s,2s,2p,3s,3p,4s
Fe: 1s,2s,2p,3s,3p,4s,3d
Hg: 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d

21. QUANTUM NUMBERS (cont)
“m”; Orbitals
determines the # of e- held in the sublevels & their position
n2 = total orbitals per energy level
ex. n=1, 1 orbital, n=3, 9 orbitals
Orbitals are filled in order of increasing energy, with no more than two electrons per orbital

22. QUANTUM NUMBERS (cont)
“m”; Orbitals (cont)
s = 1 orbital & 1 pr. of e-
p = 3 orbitals & 3 pr. of e-
d = 5 orbitals & 5 pr. of e-
f = 7 orbitals & 7 pr. of e-
*only 1 pr. of e- per orbital

23. 2 6 10 14 Summary Starts at energy level s 1 1 p 3 2 d 5 3 7 4 f
Maximum electrons
Starts at energy level
# of (orbitals) 2 s 1 1 p 3 6 2 10 d 5 3 7 14 4 f

24. Quantum Number, ml
Describes the three-dimensional orientation of the orbital.
Values are integers ranging from -l to l:
−l ≤ ml ≤ l.
Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.

25. Quantum Number, ml Orbitals with the same value of n form a shell.
Different orbital types within a shell are subshells.

26. QUANTUM NUMBERS (cont)
“s”; Spin
rotation of e-
each pair of e- must have opposite spins:
+1/2 or -1/2 or _ _

27. To find location & energy of e-, you must know the
QUANTUM NUMBERS
(click on the image below to watch a review video on Quantum Numbers – watch to the 4:20 mark)

28. Electron Configurations
The way in which electrons are distributed among the various orbitals is called the electron configuration.
Watch the video for how to write the electron configurations of elements

29. Writing Electron Configuration
Mg=12
1. e- enter energy levels of lowest energy 1st
(1s2, 2s2, 2p6, 3s2)
Try it with these elements:
Al: 1s22s22p63s23p1
K: 1s22s22p63s23p64s1
Fe; 1s22s22p63s23p64s23d6
Hg: 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p6,6s2,4f14,5d10

30. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f14 5d106p67s1
Do you notice any similarity in these configurations of the alkali metals?

31. Do you notice any similarity in the configurations of the noble gases?
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 He
Do you notice any similarity in the configurations of the noble gases? 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

32. Look at the electron configurations for these elements:
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 (expected)
But this is not what happens!!

33. Chromium is actually:
1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals (the others are all still full)
Half full is slightly lower in energy.
The same principal applies to copper.

34. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: d4, d9

35. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

36. Writing Orbital Notation
Watch the video on writing Orbital Notation before going on.

37. Writing Orbital Notation
Try it with these elements: Al K Fe Hg

38. Writing Orbital Notation
Use your PTE
Determine the number of e-’s
Determine the energy levels
Write the electron configuration
Determine the orbitals needed
Write the orbital notation using arrows

39. Hund’s Rule:
When occupying orbitals of equal energy, 1 e- enters each orbital (with same spin) until all orbitals have 1 e-
Only after all orbitals in a sublevel are half- filled do electrons “pair up.”

40. Writing Orbital Notation
Al: 1s22s22p63s23p1
__, __, __ __ __, __, __ __ __
K: 1s22s22p63s23p64s1
__, __, __ __ __, __, __ __ __, __
Fe; 1s22s22p63s23p64s23d6
__, __, __ __ __, __, __ __ __, __, __ __ __ __ __
Hg: 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p6,6s2,4f14,5d10
__, __ __ __ __ __ __ __, __ __ __ __ __

41. Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
To show the different direction of spin, a pair in the same orbital is written as:
Wolfgang Pauli

42. Writing Electron-Dot Diagrams
Mg = 12
Electron Configuration
Orbital Notation
Use only valence electrons and orient to the orbital notation
Watch how-to video:

43. Writing Electron-Dot Diagrams
Try it with these elements: Al K Fe U Hg