1. The Quantum Model of the Atom

2. Quantum Mechanical Model
The quantum mechanical model makes no attempt to predict the path of an electron around the nucleus
Bohr orbits were replaced with quantum-mechanical orbitals
Orbitals are different from orbits in that they represent probability maps that show a statistical distribution of where the electron is likely to be found

3. Quantum Mechanical Model
In this model, a number and a letter specify an orbital or energy level
As value of “n” increases, orbital becomes larger
Electrons spend more time away from nucleus
Each orbital level can hold a maximum of 2 electrons
“n” = an atom’s principal energy level

4. Principal Energy Sublevels
Sublevels = “s”, “p”, “d”, & “f” distinctions
Correspond to shape of atom’s orbitals
All “s” orbitals = spherical; all “p” orbitals = dumbbell-shaped; “d” and “f” vary

5. Principle Energy Sublevels
The lowest-energy orbital is called the 1s orbital – specified by the number 1 and the letter s.
Two sublevels in P.E.L 2 – 2s and 2p
2s sublevel is same shape, larger size than 1s
2p sublevel has 3 dumbbell-shaped p orbitals of equal energy (2px, 2py, 2pz)

7. Principle Energy Sublevels
P.E.L. 3 contains 3 sublevels (3s, 3p, & 3d)
“s” and “p” sublevels are same
“d” sublevel contains 5 orbitals of equal energy
P.E.L 4 contains 4 sublevels (4s, 4p, 4d, & 4f)
“s”, “p”, & “d” sublevels are same
“f” sublevel contains 7 orbitals of equal energy

8. S-sublevel

9. P-sublevel

10. D-sublevel

11. F-sublevel

12. Review
Bohr predicted that electrons orbited the nucleus like plants orbit the sun
de Broglie hypothesized that electrons behave like waves
Schrodinger predicted existence of atomic orbitals contained electrons
Orbitals existed within the energy levels

13. Electron Configuration
Overview of the arrangement of electrons in an atom
Electrons tend to take position of lowest energy, making low energy systems more stable than high
An atom’s Ground State = most stable configuration
This is the reason that valence electrons are the most active!!

14. Let’s follow the RULES! Aufbau Principle
Each e- occupies lowest energy orbital available
Orbitals are filled completely from lowest to highest energies
s  p  d  f

15. Pauli exclusion principle:
A maximum of 2 e- may occupy a single orbital level
e- spin in one of two directions
Up arrows indicate one direction, down arrows opposite
Atomic orbitals that are FULL are represented as

16. Hund’s rule:
Single e- with same spin must occupy each orbital before e- with opposite spins are added
Good example looks at “p” orbitals

17. Orbital Diagrams Method 1 for showing an atom’s electron configuration
1 box for every orbital
Empty boxes = unoccupied orbitals
Each box labeled with the principal quantum # and sublevel associated with the orbital

18. Let’s work some examples!!H C Ne Mg Ge Au

19. Electron Configuration Notation
Shows P.E.L and energy sublevel associated with each of the atom’s orbitals
Includes superscript representing # of e- in orbital
DOES NOT INCLUDE arrows!!

20. Electron Configuration Notation examples…B Be O Si Nb

21. Noble Gas notation
Represents e- configuration using brackets and symbols of Noble Gases
[He] = 1s2
[Ne] = 1s22s22p6
Using noble gas of most previous one, write e- configuration notation of elements

22. Noble Gas notation examples….Li Na K Si S Eu

23. Valance E- and Configurations
Valence electron = e- in outermost energy level
Determines reactivity
Can use Noble Gas configuration notation to determine valence numbers
S = [Ne]3s23p4
Cs = [Xe]6s1
Fr = [Rn]7s1
Simply add the exponents from the notation