1. Quantum Mechanical Model of the Atom Orbitals and Electron Configuration
Mrs. Hayes
Chemistry

2. Quantum Mechanical Model

3. Orbitals and the Quantum Mechanical Model of the Atom
Since we know that Bohr’s model of hydrogen will not work for all atoms, we need another way to predict where electrons are located.
We use “orbitals” not “orbits” to make this prediction.

4. s, p, d orbital shapes

5. Quantum Mechanical Model

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

7. e- Movement is Based on 3 Factors
Their energies.
Their orientation in space.
How wildly they swing about in their path (some swing back and forth like waving your arms instead of making a full circle).

8. Quantum Numbers Describe the locations of the e-’s around the nucleus.
Quantum #’s are sort of like a home address for the electron.

9. Quantum Numbers
Principal Q. #: Describes the distance that the electron is from the nucleus. The bigger the number, the farther away the electron is. Example: (1=closest, 2, 3, 4...farther away)
Magnetic Q. #: tells how many orientations in 3-D there are about the nucleus for each orbital shape.
s = 1 orientation
p= 3 orientations... (x, y, and z)
d= 5 orientations
f= 7 orientations

10. Quantum Numbers
Angular Momentum Q. #: Describes the shape of the electron’s path around the nucleus with a letter: (s, p, d, & f) These are sometimes called “sublevels”.
Spin Q. #: describes how the electron in an orientation is spinning around the nucleus. Some like to imagine it spinning “clockwise” and “counterclockwise”.
The spin is represented as an arrow in the direction of the spin.

11. Electron Configurations
“Configuration” means the arrangement of the parts of something.
Electron configuration is the arrangement of electrons in space around the nucleus.

12. Rules for e- configurations
Rule #1 (Aufbau Principle): Electrons fill lowest energy orbitals first.
Rule #2: Only 2 electrons can fit into each orientation.
Rule #3 (Pauli Exclusion Principle): Electrons in the same orientation have opposite spin.
Rule #4 (Hund’s Rule): “Monopoly rule”---> Every “□” in an orbital shape gets an electron before any orientation gets a second e-.

13. Electron Configurations
The orientations are represented with a line or a box.
Examples: ___ This means a spherical orbital at a distance of 1s, “1” (close) to the nucleus. This orbital is centered about the x, y, and z axis.
□ □ □ This represents an ellipsoid orbital with its 4p with 3 possible orientations at a distance of “4 ”from the nucleus

15. Electron Configuration Pattern

16. What do the Numbers and Letters Represent?
The large numbers represent the main energy levels and are determined by the electrons’ average distance from the nucleus.
The letters are the sublevels in which the electrons are found (s, p, d, f)
Each main energy level has a certain number of sublevels.

18. How to Use the Configuration
We move across the periodic table from left to right just like we read a book.
Let’s take the element Lithium (Li, 3)
Lithium has 3 protons, so a neutral atom of Li will have 3 electrons
We need space for 3 electrons:
1s2, 2s1
There are 2 electrons in the first sublevel, and one electron in the second sublevel.
2 + 1 = 3…that’s it!

19. Orbital Diagrams We can DRAW the Configuration, too.
2s orbital can hold 2 electrons, but we have only one to place there.
1s2
2s1
You must fill the first sublevel before going to the next one.
Electrons have opposite “spins” in an orbital, so we have the arrows going in the opposite direction.
Lithium

20. Let’s Try Fluorine…
Fluorine has an atomic number of 9, so it has 9 electrons in a neutral atom.
1s2,2s2, 2p5
The 2p orbital can hold 6 electrons, but we only have five to place there.
1s2
2s2
2p5

21. What About Arsenic? Atomic number = 33, so 33 electrons
Fill up the s orbitals first
Fill up the sublevels until we get to 4p…then we have only three left there.

22. Hund’s Rule. Each orbital of equal energy must be occupied before a second electron can enter. All single electrons have the same spin.
4p3
3d10
4s2
3p6
3s2
2p6
Pauli says no two electrons in orbital can be alike, so opposite spins.
2s2
1s2
Aufbau principle

23. We Also Have Shorthand- The “Noble Gas Configuration”
Use the nearest noble gas BEFORE the element to replace the sublevel orbitals in your configuration.
Example for Magnesium:
Neon, 3s2
Neon replaces 1s2, 2s2, and 2p6