1. Electron Configurations
HW: read CH 11

2. Atoms and Energy
Radiant Energy: knowledge led to refinements of atomic model
A. Wave particle: light behaves as both a wave and a particle-strange!
B. Electromagnetic Radiation (EMR): form of energy that exhibits behavior (light) as it travels through space (X-rays, visible light, radio waves)
All EMR has a wavelength, frequency, and amplitude that determines its energy

3. Atoms and Energy Electromagnetic Spectrum
Different types of radiation have different wavelengths
Different colors have different wavelengths

4. Emission of Energy by Atoms
When atoms receive energy from some source and become excited, they release energy by emitting light
Light is different colors for different elements – must be because of various energy levels

5. Emission of Energy by Atoms
Emission line spectra is like a chemical fingerprint!
Each element has a different atomic structure so their emission line spectra is unique!
Why do larger gases produce more color bands than smaller elements??? More energy levels!
Argon (swapped)

6. Bohr Model Limitations
Why is the Bohr model of the atom no longer accepted?
1. Does not explain behavior of more than 1 electron
2. NOT 3D representation
3. Impossible to know location of every electron
8ROHpZ0A70I (4 min)

7. Modern Theory: Atomic Orbitals
Every energy level has sub levels:
1st = s (lowest energy)
2nd = s and p
3rd = s and p and d
4th = s and p and d and f (highest energy)
s = 2 electrons
p = 6 electrons
d = 10 electrons
f = 14 electrons

8. Whiteboards!

9. Tricks and Hints
All the superscript values must add up to the total number of electrons.
Be very careful with inner transition metals. Some periodic tables are “incorrect” (textbook + whiteboard = bad)
Sometimes you can write in shorthand notation. Called abbreviated configuration or noble gas core method.
Write the symbol of the noble gas that comes before the desired element in a bracket. Write the rest of the configuration that comes after the noble gas
Example: Selenium: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
Abbreviated: [Ar] 4s2 3d10 4p4

11. YOU TRY! Write the full electron configuration for: A) Mn B) Ce
Write the abbreviated electron configuration for:
C) Sn
D) Cf

12. Orbital shapes
Shapes indicate the region around the nucleus of an atom where an electron is likely to be found (90% chance)
s = sphere/circle
p = dumbbell
d = dumbbells on 3 planes
f = 2 dumbbells on 5 planes

13. s orbital has 1 sub-orbital
p orbital has 3 sub-orbitals
d orbital has 5 sub-orbitals
f orbital has 7 sub-orbitals

14. Rules for electrons and orbitals
1. Aufbau Principle: electrons are added 1 at a time to the lowest energy level until all electrons have been used. So 1s, 2s, 2p, ...
2. Pauli Exclusion Principle: an atomic orbital may hold at most 2 electrons. Each electron must spin in opposite directions (arrows)
3. Hund’s Rule: orbitals of equal energy are occupied by 1 electron before any orbital is occupied by a 2nd electron (all must have 1 electron before having 2)
Crash course (12 min)

16. Electron Config Examples
(HINT: the superscript total should equal atomic number)
Nitrogen – atomic number 7
Electron config: 1s2 2s2 2p3
Orbital Notation

17. Electron Config Examples
Silver – atomic number 47
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9

18. Electron Config Examples
YOU TRY: Magnesium – atomic number 12
YOU TRY: Iridium – atomic number 77
(you can do abbreviated version of Ir)

19. Of course there will always be exceptions (fancy AP info)
Chromium SHOULD be [Ar] 4s2 3d4 BUT is really [Ar] 4s1 3d5
Copper SHOULD be [Ar] 4s2 3d9 BUT is really [Ar] 4s1 3d10
This minimizes electron repulsions

20. More exceptions
Isoelectric elements have the same electron configuration
Ca 1s2 2s2 2p6 3s2 3p6 4s2
Ti2+ 1s2 2s2 2p6 3s2 3p6 4s2

21. How would you write this??
Write the abbreviated version and show orbitals for the following:
Ag 2+
Br 1-
Co 3-

22. Periodic trends – Atomic Radius
Atoms get larger going down a group
Atoms get smaller left to right across a period
more protons as you move to the right; the more protons give more atomic pull. The strong attractive force shrinks orbitals so the atom is smaller

23. Periodic trends – Ionic Radius
Recall: Ions are elements that gain or lose electrons 
More electrons the size becomes larger
more repulsion's, spread out electrons, increase size
Fewer electrons the size becomes smaller
reduces repulsion's, electrons pulled closer to nucleus

24. Periodic trends – Ionization Energy
Energy needed to remove one of its electrons
Atoms with high I.E. hold onto their electrons tightly
Atoms with low I.E. easily lose electrons
I.E decreases as you move down a group
I.E. increases as you move from left to right 
think of octet rule for optimum number of electrons

25. Ionization Energy exceptions (more fancy AP info)
It requires less energy to remove a p4 than a p3 due to increased repulsion with p4 (Note: half full shells = happy is NOT a correct answer)
Ionization energy decreases a little between Be/B and Mg/Al and Zn/Ga because it goes from an s orbital (close to nucleus) to a p orbital (farther away from nucleus)

26. Periodic trends - Electronegativity
Ability to attract electrons, related to Ionization Energy
Fluorine most electronegative (wants electron the most)
Left side of table least electronegative (not want electrons)