1. Conservation of energy -Energy cannot be created or destroyed-
The difference between each energy level is called a Quantum
When energy is added to an electron, the electron will become “excited” moving to a higher energy level
When the electron comes back down to its ground state energy is released----in the form of light.

2. Electron Configuration
-Each electron is a given set of 4 quantum numbers
Principle Quantum Number or Energy Level
- represent by the letter “n”
n= 1 <2 <3 < 4
- generally, the larger the value of “n” , the further away the electron is from the nucleus.
n=1 can hold a max of 2 e-
n=2 can hold a max of 8 e-
n=3 can hold a max of 18 e-
n= 4 can hold a max of
n = 5 can hold a max of

3. 2nd is L 4 main orbitals—(each orbital can hold 2 electrons)
Provides a code for the shapes of orbital
- represented by the letter “l”
4 main orbitals—(each orbital can hold 2 electrons)
s-orbital –can hold 2 electrons
-lowest in energy level
-spherical in shape
p-orbital-(can hold six electrons)
-3 different p-orbitals (Px, Py, Pz)
- shaped like a dumbbell

4. 3. Magnetic quantum number (ml)
- spatial orientation of the orbital that the electron occupies.
Is it in Px, Py, Pz etc.
4. Spin quantum number (ms)
-indicates the way the electron is spinning
- ±1/2

5. Writing Electron configurations
Copy the following on your periodic table
Or you can use the Aufbau diagram(below)

7. Rules of electron configurations
Pauli Exclusion Principle – No 2 e- can have the same set of four quantum numbers.
The Aufbau Principle
- Electrons fill in lowest energy levels first
1s2 2s2 2p6 3s2 3p6
Hund’s Rule (Bus Back seat Rule)
- electron will fill up empty orbitals before pairing up.

8. At Higher levels there will be an overlap of energy levels-
Writing electron configurations
Ex: F= 1s2 2s2 2p Ar= 1s2 2s2 2p6 3s2 3p6
K= 1s2 2s 22p6 3s2 3p6 4s1
or [ Ar] 4s1
Try these:
Cl:
Na:
Orbital Diagrams ---same concept
Oxygen 8e- 1s2 2s2 2p4