1. September 1 and September 2 Warm-up
On a blank left side page draw the following diagrams.
Bohr Model Quantum Model

2. Agenda History of Atomic Models
In-Depth look at Bohr and Quantum Models
Comparison of two models
Atomic Orbitals & Sub-Orbitals
Electron Configurations & Rules
How To: Write Electron Configurations
Practice

3. History of the Atomic model and Electrons
1803: John Dalton and the Dalton Model
1897: JJ Thompson and the “plum-pudding” model
1911: Rutherford and the Rutherford Model (First model with a nucleus)
1913: Bohr and the Bohr Model (First model with electron orbitals based on energy)
1926: Schrodinger and the Quantum Model (current)

4. Bohr Model What did it get right?
Electrons have energy levels and can move up as they receive energy or down as they give off energy (as light)
Energy levels are not evenly spaced

5. Quantum Model What changed?
Theoretical calculations and experimental results were inconsistent with Rutherford and Bohr Models.
Schrodinger used new equations to describe the behavior of electrons and “created” Electron Cloud

6. What is the Electron Cloud and the Quantum Model?
The Quantum Model determines allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
A quantum is exact energy needed to move an electron from one energy level to another.
The “Cloud” is more dense where the probability of finding the electron is high.
The “Cloud” is less dense where the probability of finding the electron is low.

7. BREAK TIME

8. Atomic Orbitals
Shrodinger’s equations led to Energy Levels for electrons.
Energy levels use Principle Quantum Numbers (n)
These are assigned values (n=1, 2, 3, 4, …)
These energy levels have sub-levels which describe the shape of the electron cloud (shape of probability)
The MAXIMUM number of electrons that can fill an energy level is described as 2n2
n=1 : 2 electrons n=2 : 8 electrons
n=3 : 18 electrons n=4 : 32 electrons

9. Sub-Orbitals Sub-orbitals are given letter designations (s, p, d, f)
# of Orbitals
Maximum Electrons
First possible energy level
s (spherical) 1 2
p (dumbell) 3 6
d (clover leaf) 5 10
f (complicated) 7 14 4

12. Summary Table Energy Level Number of Sublevels Types of Sublevels n=1
1s (1 orbital)
n=2 2
2s (1 orbital), 2p (3 orbitals)
n=3 3
3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals)
n=4 4
4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals)

13. BREAK TIME

14. Electron Configurations
Electrons fill orbitals based off energy, called the Electron Configuration
Electron configurations follow 3 rules:
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Define each of the rules in your notebook. You will need a textbook/internet. Take 15 minutes to do your research and compare to your neighbors.

15. Aufbau Principle – Electrons occupy lowest energy level first
Pauli Exclusion Principle – Orbital can have two electrons w/opposite spins
Hund’s Rule – One electron occupies each orbital until all are filled (then fill opposites)

17. How To: Writing Electron Configurations
Start with the # of electrons element has
Next, use energy diagram to write out orbitals
1s, then 2s, then 2p, then 3s, etc
Follow pattern until you have used all electrons
Fill orbitals following 3 rules (Aufbau, Pauli, and Hund)

18. BREAK TIME

19. Practice
Hydrogen:
Carbon:
Phosphorus:

20. Is there another way to do this?
There is another method (also called Electron Configuration) that simplifies the Orbitals
Hydrogen: 1s1
Carbon: 1s2 2s2 2p2
Phosphorus: 1s2 2s2 2p6 3s2 3p3

21. How does this method work?
It is the same as the Orbital Filling method, but all similar energy levels are combined.
When all superscripts are added you have the total number of electrons.

22. Practice
Helium:
Oxygen:
Aluminum:

23. Homework Problems 25, 30-32, 34, 50, 57 from Chapter 5 of textbook
Finish Element Research Project