1. AP Chemistry Unit 2 Topics:
atomic models: evolutionary evidence, limitations of shell model, evidence of QM idea*
Orbital diagrams; how it relates to the periodic table (and properties)
e- dot notation
isotopes (mass spec): counting p+, no, & e−
periodic table: sections, trends using shell model and Coulomb’s Law, PES data
*Some of these topics will have a corresponding video from the internet to watch instead of “formal notes”.
-- Also, much of Unit 2 contains PowerPoint slides that were taken directly from the 1st year chemistry notes that I use. So if I cover this old material too rapidly, you can listen to the 1st year lecture on my website to refresh your memory of these “review topics”. I will go slower over the “new material” compared to “old material”.

3. Quantum Mechanical Model
Quantum Mechanical Model

7. Quantum Mechanical Model

8. Bohr’s Energy Levels
The energy levels in an atom are sort of like _________ of a ladder.
The more energy an electron has, the __________ away from the nucleus it usually will be.
The energy levels are not evenly spaced. They get ___________ together as you travel farther away.
To move from one “rung” to another requires a “____________” of energy.
steps
farther
closer
quantum

9. Bohr Atomic Model

10. The difference between continuous and quantized energy levels.
continuous energy levels
quantized energy levels

11. Quantum Numbers
Describe the ______________ of the e-’s around the nucleus.
Quantum #’s are sort of like a home _____________ for the electron.
This information about the location of the e-’s in an atom can be used to:
(1) determine chemical & physical _____________ for the elements.
(2) show how the _______________ __________ is organized.
(3) show _____ and _____ elements combine to form compounds.
location
address
properties
Periodic Table
how
why

12. The Four Quantum Numbers
Principal Q. #: Describes the _____________ that the electron is from the nucleus. The bigger the number, the ___________ away the electron is.
Example: (1=closest, 2, 3, 4...farther away)
These distances are sometimes called _______________
______________ ____________.
distance
farther
principal
energy levels 1 2 3
nucleus

13. shape
Orbital Q. #: Describes the __________ of the electron’s path around the nucleus with a letter: (s, p, d, & f) These are sometimes called “_____________”.
s=_____________ cloud; p=_____________ or a 3-D figure 8;
sublevels
spherical
ellipsoid

14. d & f orbital shapes are complex ________- _______________ ellipsoids, and some d’s and f’s are an ellipsoid with a doughnut or two around the middle.
All of these orbital shapes are based on the probability of finding the electron in the cloud.
d - orbitals
criss
crossing
f - orbitals

15. How principal energy levels can be divided into sublevels:f s p d s p s

16. Magnetic Q. #: tells how many _________________ in 3-D there are about the nucleus for each orbital shape.
s=___ orientation p= ___ orientations... (x, y, and z)
d= ___ orientations f= ___ orientations
The orientations are represented with a line or a box.
Examples: ___ This means a __________ orbital at a distance of 1s “__” (close) to the nucleus. This orbital is centered about the x, y, and z axis.
□ □ □ This represents an ___________ orbital with its p ____ possible orientations at a distance of “___” from the nucleus.
orientations 1 3 5 7
spherical 1
ellipsoid 3 4

17. Quantum Mechanical Model

18. Principal energy level 2 shown divided into the 2s and 2p sublevels

19. Spin Q. #: describes how the electron in an orientation is spinning around the nucleus. This spin can be thought of as “____” or “________”. (Some like to imagine it spinning “clockwise” and “counterclockwise”.) The spin is represented as an ___________ in the direction of the spin.
Example: ↑ This represents one electron in a _________ s orbital with spin “____” at a distance of “___” from the nucleus.
Remember, the four quantum numbers tell us the location, or “address” of each electron in an atom.
This information is vital in understanding the layout of the Periodic Table and the reasoning behind why and how atoms form bonds. up
down
arrow
spherical up 2

20. Electron Configurations
Electron configurations are notations that represent the four Quantum #’s for all of the electrons in a particular atom. Here are the rules for these notations:
Rule #1 (Aufbau Principle): Electrons fill ________ energy orbitals first.
Examples: 1s would be filled before ____
3s would fill before ____
lowest 2s 3p

21. Electron Configurations↑ ↑
Silicon ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓
(Energy Level Diagram) ↑ ↓
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p…

22. Rule #2: Only ___ electrons can fit into each orientation.
Example: ___ ___ not ____
1s s s
Rule #3 (Pauli Exclusion Principle): Electrons in the same orientation have ______________ spin.
Example: ___ not ___
1s s
Rule #4 (Hund’s Rule): “_______________ rule” or “Bus Seating” rule---> Every “□” in an orbital shape gets an electron before any orientation gets a second e−.
Example: □□□ not □□□
2p p 2 ↑ ↓ ↑ ↑ ↓ ↑
opposite ↑ ↓ ↑ ↑
Monopoly ↑ ↑ ↑ ↑ ↓ ↑