1. Modern Atomic Model Sometimes called:
Charge Cloud Model
Wave Mechanical Model
Quantum Theory
Theory created to help describe spectral line signatures of multi-electron elements
Crash Course: History of Atomic Theory (9:45)

2. How is it Different from the Planetary Model?
We can’t tell exactly where an electron is!!
Heisenberg’s Uncertainty
Principle:
We can’t know exact location and momentum of an electron at the same time.

3. Electrons exist in “orbital clouds”
Denser the cloud region the higher the probability of finding an electron there.
Can’t tell exactly where an electron is, but can tell region of higher probability of finding it.

4. Excited Atoms and Charge Cloud
Excited electrons shift to probability regions further from the nucleus

5. Comparison to Bohr’s Electron Model
Charge Cloud = area of highest probability of finding electron, but can’t know exact location
Bohr Model = electrons in fixed orbit like planet around the sun at fixed distance from nucleus

6. STOP HERE
HONORS ONLY MATERIAL LIES BEYOND

7. Extra Videos
Quantum Mechanics and the Bohr Model (6 minutes)
**Developing Modern Atomic Theory (6min)
Rutherford to Bohr to the Modern Model
The Uncertainty Principle (6min) Honors
Probability & Chance of Finding an e- and Orbital Shapes

8. How are Electrons Organized?
Energy Levels
Sublevels
Orbitals
Spin
The Electron Hotel

9. Energy Levels (1-7) Energy Levels 1 2 3 4 5 6 7
Electrons exist at different distances from the nucleus.
Energy Levels
Lowest energy Highest energy
Closest to nucleus Farthest from Nucleus

10. Sublevels (s, p, d, f) Energy levels have certain number of sublevels.
Energy Level Sublevels Possible
1 s
2 s, p
3 s, p, d
4 s, p, d, f
5 s, p, d, f, (g)
6 s, p, d, f, (g, h)
7 s, p, d, f, (g, h, i)
Theoretical sublevels

11. Energy of Sublevels Sublevels have different levels of energy. s p d f
Lowest energy Highest energy

12. Orbitals in Sublevels Sublevels contain different numbers of orbitals.
Maximum 2 electrons can exist in an orbital.
Sublevel # of Orbitals Max e- in Sublevel
s e-
p e-
d e-
f e-

13. Aufbau Diagram Shows: Follow the “diagonal rule”
order of electron filling
order of electron energy
Follow the “diagonal rule”

14. Writing Electron Configurations
Let’s write some electron configurations!
Ex: Magnesium
Follow the diagonal rule

15. Overlap Some sublevels “overlap”
Results in certain sublevels having higher energy than others
Ex: 3d has higher energy than 4s

16. Electron Configurations for Atoms
Extremely Corny Song About Electron Configurations

17. Electron Spin Electrons in an orbital spin in opposite directions
Pauli Exclusion Principle:
In order for two electrons to occupy the same orbital, they must have opposite spins.
Otherwise they create a magnetic field!

18. Orbital Diagrams Show electrons in individual orbitals
s = 1 orbital, p = 3 orbitals, d = 5 orbitals, f = 7 orbitals
“Hund’s Rule for Orbital Filling”
When filling orbitals in a sublevel, place one electron in each orbital before adding the second.

20. Paramagnetic vs. Diamagnetic
Diamagnetic: only paired electrons
in orbitals
Paramagnetic: at least one orbital
with an unpaired electron.
Diamagnetic atoms repel magnetic fields.
The unpaired electrons of paramagnetic atoms realign in response to external magnetic fields and are therefore attracted. Paramagnets do not retain magnetization in the absence of a magnetic field, because thermal motion randomizes the spin orientations.

21. Shapes of Orbitals Orbitals come in different shapes and sizes.
Region of highest probability of finding an electron.
Orbitals Shape & Energy & Spectral Line

22. Probability cloud has a spherical shape
s Orbital
Probability cloud has a spherical shape

23. p Orbitals (px, py, pz) “Dumbell” shape
Three p orbitals can exist, on the x, y, z axis in space

24. Clouds and Probability

25. Five possible d orbitals exist

26. Seven possible f orbitals exist

27. Valence Electrons
Usually found in the s and p sublevels of highest occupied energy level.
How many valence electrons?
Draw a Lewis Dot Diagram of this element.
1s22s22p63s23p64s23d104p2

28. Kernel 1s22s22p63s23p3 All electrons except the valence
How many kernel electrons?
How many valence?
1s22s22p63s23p3

29. Excited vs. Ground State
Electron configuration you would normally write by following the order of filling
Lowest to highest energy.
Excited State:
one or more electrons have jumped up to a higher energy level.
Ex: 1s22s22p63s23p54s23d104p3

30. Atom vs. Ion Configurations
Ions: atoms that have gained or lost electrons.
Figure out how many electrons the ion has then write configuration.
Ex: 20Ca+2 has 18 electrons
1s22s22p63s23p6 = 18 electrons DONE!
Electron Configurations for Ions

31. Isoelectronic Species
Atoms and ions that have the same number of electrons.
Ex: Ar, K+1, Ca+2, P-3, O-2, Cl-1
All have 18 electrons!

32. Impossible Configurations
Break the rules.
Ex: 1s22s22p63s22d103p64s23d104p2
Ex: 1s22s32p63s23p6

33. s, p, d, f, “Blocks”
Indicates what sublevel is being filled last in the atom

34. Some Exceptions to Orbital Filling (HONORS)
When d and f sublevels get filled near the end of a configuration we sometimes see exceptions.
It is more stable for the orbitals of the d and f sublevels to be half filled or filled completely than to be one shy.
Electrons from the sublevel below get “promoted” up to make the atom more stable
Ex: Copper

35. Crash Course Chemistry: The Electron
(13 minutes)

36. Quantum Numbers (Honors)
Set of 4 numbers that help to describe most probable location of each of an atom’s electrons.

37. Principal Quantum Number: (n)
describes principle energy level electron is in
values n = 1, 2, 3, ...7
n=2 can describe the 8 electrons in 2nd energy level

38. Azimuthal Quantum Number: (ℓ)
Describes sublevel electron is in (s, p, d, f)
Values of (ℓ)  range from: 0 to (n-1)
Ex:
If n=4 ℓ can be 0, 1, 2, 3 (representing s, p, d, f sublevels)
If n=2 ℓ can be 0, 1 (representing s, p sublevels)
n = 2 ℓ = 1 means the electrons in 2”p”
n = 3 ℓ = 2 means the electrons in 3 “d”

39. Magnetic Quantum Number: (mℓ)
Describes the orbital the electron is in on sublevel
Relates to it’s general orientation in space
Values from (-ℓ….0….+ℓ) **note the middle orbital is ZERO
Number of possible m values within a sublevel is = to the number of orbitals within a sublevel
Ex: ℓ= 1 (p sublevel), mℓ= -1, 0, +1
(representing px, py, pz orbitals)
Ex: n= 3, ℓ= 2, mℓ= -1
(describes the 2 electrons in the 2nd orbital in 3d)

40. Spin Quantum Number: (ms )
Values of +1/2 or -1/2 (arrow up or arrow down)
1st electron in orbital is “clockwise” = +1/2
2nd electron in orbital is “counterclockwise” = -1/2
Ex: n=2, ℓ = 1, mℓ = +1, ms = -1/2
Represents 2p6 electron