1. Electron Configuration

2. Textbook Connection Chapter 4 The Development of a New Atomic Model
Light has characteristics of both particles and waves
Photoelectric Effect
Energy Levels
Bohr Model
The Quantum Model of the Atom
Electrons as Waves
Heisenberg’s Uncertainty Principle
Orbital Probabilities
Quantum Numbers
Electron Configurations
Aufbau Principle
Electron Configuration Forms

3. Textbook Connection Homework – Due November 7th
Page 104 – 1, 2, & 3
Homework – Due Nov. 10th
Page 116 – 1, 2, 3, 4, & 5

4. First, A Little Review:
Electrons are just one of the three major components of Atoms.
The Electron’s mass is so small in comparison to the other subatomic particles that its mass can be ignored in most cases.
Electrons have a charge that is equal but opposite (negative) to that of the Proton (+)

5. Review continued…
The electron was the first subatomic particle to be discovered (J. J. Thomson)

6. Review Continued…
Ernest Rutherford’s “Gold Foil” experiments showed that the atom is mostly empty space with relatively great distances between the electrons and the nucleus.

7. Rutherford Model Electrons existed in set orbits around the nucleus.
Problems with this model is that electrons would fall into the nucleus almost instantly.

8. Modern Atomic Theory Max Planck – “ Energy is not continuous.”
While studying the radiation of heated materials, he attempted to make connections to the radiating atom.
Based on his observations, he determined that energy is always emitted or absorbed in discrete units which he called “quanta.”
Energy = frequency of the radiation * constant
This constant is known as Planck’s constant and is equal to
6.63 x J * s

10. Albert Einstein – Photoelectric Effect
Applied Planck’s work further to the idea such that light energy should be explained as “packets” of energy (quanta) which he called “photons” that could be seen to behave as both wave and particle.
This idea lead to an understanding of why metals lose electrons when they absorb electromagnetic radiation.

11. Niels Bohr– Refined Rutherford atomic model to explain how electrons resist “falling” into the nucleus.
Electrons can only exist in discrete levels or “orbitals” at some distance away from the nucleus based on the amount of energy the electron has.
Electrons can move between levels only if they absorb or radiate a set amount of energy (quanta)

12. Louis de Broglie – Particle/wave duality
Discovered that electrons had a dual nature – similar to both particles and waves which supported Einstein’s concepts of matter and energy.

13. Heisenberg – Uncertainty Principle
Continued to build on the principles of de Broglie and developed formulas connected to the frequencies of spectral lines.
“It is impossible to know both the location and velocity of an electron.”

14. Schrodinger – Mathematical model of quantum mechanics
Viewed electrons as continuous clouds and developed mathematical probabilities of electrons’ existence in the atom.
Schrodinger’s Cat – All explanations of the realities related to a specific situation remain valid until the “box” is opened.

15. James Chadwick – Isolates the Neutron
Using alpha particles he discovers and isolates the neutral atomic particle with a mass close to a proton.

16. Bohr Model
Neils Bohr proposed a model that had electrons existing in orbitals around the nucleus that were differentiated by the amount of energy that they had.

17. Quantum Model
The Quantum Model proposed by Schrodinger’s equations, De Broglie’s wave/particle theory, and Heisenberg’s Uncertainty Principle gives us a much more complex view of the organization of the electron’s of atoms.

18. Quantum Mechanical Model
Unlike the Bohr model, the Quantum model does not restrict an electron to an exact path but instead estimates a probability of finding en electron in a certain position.
Similar to Bohr, electrons finding themselves in predicted areas called “Orbitals” based on having a set amount of energy.

19. Atomic Orbitals
Orbitals are organized by a number of different factors (Quantum Numbers):
Principle Energy Level
Refers to the energy Level, (relative ‘distance’) from the nucleus. Currently, n = 1 – 7
Atomic Orbitals (l angular momentum)
Letters denote the estimated “shape” of the orbital

20. “S” - Orbital
1 orbital “space” = maximum 2 e-

21. Orbitals are organized by a number of different factors (Quantum Numbers):
Magnetic Quantum Number
Tells us the orientation of the space (X, Y, or Z)
Spin Quantum Number
Electrons cannot occupy the same space at the same time so +1/2 or -1/2 to indicate the direction of movement

22. 3 orbital “spaces” = maximum 6 e-
“P” - Orbitals
3 orbital “spaces” = maximum 6 e-

23. 5 orbital “spaces” = maximum 10 e-
“D” - Orbitals
5 orbital “spaces” = maximum 10 e-

24. 7 orbital “spaces” = maximum 14 e-
“F” - Orbitals
                                                              
                                                         
7 orbital “spaces” = maximum 14 e-

25. Sublevels
The number of sublevels is always equal to the principle energy level
n = 1  1 sublevel
n = 2  2 sublevels
n = 3  3 sublevels
n = 4  4 sublevels
n = 5  5 sublevels
The maximum number of electrons that can be found on any energy level can found using the equation “2n2”
Energy level “3”  n = 3  2(3) 2 = 18 maximum electrons

26. Pattern of Electron Configuration
Start 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f g
6s 6p 6d 6f g h
7s 7p 7d 7f g h i

27. Electron Configuration Examples
Start 11 Na
22.989 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2
2s2
2p6
3s1
7 e- - 6 e- = 1 e-
9 e- - 2 e- = 7 e-
11 e- - 2 e- = 9 e-

28. Another Example: Start Co 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f27 Co 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2
2s2
2p6
3s2
3p6
4s2
3d7
27 e- - 2 e- = 25 e-
9 e- - 2 e- = 7 e-
25 e- - 2 e- = 23 e-
23 e- - 6 e- = 17 e-
17 e- - 2 e- = 15 e-
15 e- - 6 e- = 9 e-

29. Now Try These: 12Mg 18Ar 1s2 2s2 2p6 3s2 35Br 1s2 2s2 2p6 3s2 3p6 82Pb
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2