1. Review Periodic Table
Mendeleev- arranged in order of increasing atomic mass.
Moseley – in order of increasing atomic number
Atomic mass = #protons + # neutrons, atomic number = #protons
Period- row of elements, properties change gradually
Group – column of elements, similar properties
Electron configuration – all the elements in a group have the same configuration of electrons in their outer shell

2. In the beginning
When God started building the atoms of elements he began with a design. We may have discovered that design.
He began adding electrons, protons and neutrons according to the design.
The first element had 1 electron and 1 proton and either 0,1 or 2 neutrons.
The second element had 2 electrons and 2 protons and either 1,2 or 3 neutrons.
Simulation

3. Uncertainty Principle
The problem of defining the nature of electrons in atoms was solved by W. Heisenberg.
We cannot simultaneously define the position and momentum (= m•v) of an electron.
We can define e- energy exactly but we must accept the limitation that we do not know the exact position of the e-.
W. Heisenberg

4. Electron Configuration
Electrons within the electron cloud have different energy levels.
Each electron is assigned 4 different “quantum numbers”. No two electrons within an atom can have the same quantum numbers.
Energy Level quantum number, there are 7 levels (1-7)
Shape quantum number -suborbits for electrons s,p,d,f
Orientation quantum number
Spin quantum number- right handed or left handed

5. QUANTUM NUMBERS n  the principal quantum number or shell
Each orbital is a function of 3 quantum numbers: n, l, and ml
n  the principal quantum number or shell
(or size and energy level)
l  the angular momentum quantum number or
subshell (or shape)
ml  the magnetic quantum number
designates an orbital within a subshell (or the
orientation in space)
In addition each electron has 1 of 2 Spins

6. l (angular) 0, 1, 2, .. n-1 Orbital shape or type (subshell)
QUANTUM NUMBERS
Symbol Values Description
n (major) 1, 2, 3, .. Orbital size and energy (shell)
l (angular) 0, 1, 2, .. n-1 Orbital shape or type (subshell)
ml (magnetic) - l l Orbital orientation
# of orbitals in subshell = 2 l + 1

7. Shells and Subshells When n = 1, then l = 0 and ml = 0
Therefore, in n = 1, there is 1 type of subshell
and that subshell has a single orbital
(ml has a single value  1 orbital)
This subshell is labeled s
Each n value (shell) has 1 orbital labeled s, and it is SPHERICAL in shape.

8. s Orbitals
All s orbitals are spherical in shape.

9. p Orbitals When n = 2, then l = 0 and 1
Therefore, in the n = 2 shell there are 2 types of orbitals — 2 subshells
For l = 0 ml = 0
this is a s subshell
For l = ml = -1, 0, +1
this is a p subshell with 3 orbitals

10. p Orbitals When l = 1, there is a PLANAR NODE through the nucleus.
3 Orientations

11. p Orbitals
The three p orbitals lie 90o apart in space
A p orbital

12. d Orbitals
When n = 3, l = 0, 1, 2 and so there are 3 subshells in the shell.
For l = 0, ml = 0
 s subshell with single orbital
For l = 1, ml = -1, 0, +1
 p subshell with 3 orbitals
For l = 2, ml = -2, -1, 0, +1, +2
 d subshell with 5 orbitals (orientations)

13. d Orbitals
s orbitals have no planar node (l = 0) and so are spherical.
p orbitals have l = 1, have 1 planar node, and are “dumbbell” shaped.
This means d orbitals (with l = 2) have 2 planar nodes

14. d Orbitals

15. f Orbitals For l = 3, ml = -3, -2, -1, 0, +1, +2, +3
When n = 4, l = 0, 1, 2, 3 so there are 4 types of subshells in the shell.
For l = 0, ml = 0
 s subshell with single orbital
For l = 1, ml = -1, 0, +1
 p subshell with 3 orbitals
For l = 2, ml = -2, -1, 0, +1, +2
 d subshell with 5 orbitals
For l = 3, ml = -3, -2, -1, 0, +1, +2, +3
 f subshell with 7 orbitals

16. f Orbitals

17. Arrangement of Electrons in Atoms
Electrons are arranged in
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)
Various
Way of
Describing
Energy
Levels {

18. Arrangement of Electrons in Atoms
Each orbital can be assigned no more than 2 electrons!
This is tied to the existence of a 4th quantum number, the
electron spin, ms.

19. Electron Spin, ms There are two spin directions that are
Any spinning electric charge creates a magnetic field, comparable to N and S.
There are two spin directions that are
designated by ms = +1/2 and -1/2.

20. Pauli Exclusion Principle
No two electrons in the same atom can have the same set of 4 quantum numbers.
That is, each electron has a unique address.

21. Aufbau Principle Page 111
An electron occupies the lowest energy orbital that can receive it.
Fill lower energy orbitals first
Energy levels overlap
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

22. Electron Filling Order
This is confusing- follow a listing of configurations to see just how it works.

23. Hund’s Rule
Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state.

24. Writing Atomic Electron Configurations
There are two ways of writing configurations. One is called spectroscopic notation.

25. Writing Atomic Electron Configurations
The other method of writing e- configurations is called the orbital box notation.
One electron has n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2

26. See Electron Configuration in ChemLand.

27. Lithium
Group 1A
Atomic number = 3
1s22s1  3 total electrons

28. Beryllium
Group 2A
Atomic number = 4
1s22s2  4 total electrons

29. Boron
Group 3A
Atomic number = 5
1s2 2s2 2p1 
5 total electrons

30. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2  6 total electrons
Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place a single electron in each orbital as long as possible.

31. Nitrogen
Group 5A
Atomic number = 7
1s2 2s2 2p3 
7 total electrons

32. Oxygen
Group 6A
Atomic number = 8
1s2 2s2 2p4 
8 total electrons

33. Fluorine
Group 7A
Atomic number = 9
1s2 2s2 2p5 
9 total electrons

34. Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 
10 total electrons
Note that we have reached the end of the 2nd period, and the 2nd shell is full!

35. Electrons in Atoms When n = 1, then l = 0
this shell has a single orbital (1s) to which 2e- can be assigned.
When n = 2, then l = 0, 1
l = 0 is the 2s orbital which has 2e-
l = 1 is the three 2p orbitals which have 6e-
TOTAL = 8e- in 4 orbitals

36. Electrons in Atoms When n = 3, then l = 0, 1, 2
l = 0 is the 3s orbital e-
l = 1 is the three 3p orbitals 6e-
l = 2 is the five 3d orbitals 10e-
TOTAL = e- in 9 orbitals

37. Electrons in Atoms When n = 4, then l = 0, 1, 2, 3 4s orbital 2e-
three 4p orbitals 6e-
five 4d orbitals 10e-
seven 4f orbitals 14e-
TOTAL = 32e-
ChemLand: Atomic Structure
Quantum Numbers