Electron Configuration

1. Electron Configuration Writing e- configurations Drawing orbital notations

2. Quantum Mechanics Model- • describes the electron cloud around the nucleus as a 3-D wave in which the electrons move. There are different energy levels in this cloud with limits the number of electrons in each level.

3. Heisenberg Uncertainty Principle- • states that it is impossible to know both the velocity and position of a moving body (such as the electron or photon) at any one time.

4. Energy levels- • Referred to as principal energy levels (denoted by the letter “n”). • Total of 7 energy levels (look at the periodic table…there are 7 rows called PERIODS or SERIES)

5. Energy levels- • Lowest energy level, n=1, is closest to the nucleus and has a low amount of electron energy. • The electron energy is called quanta. • As the value of “n” increases farther from the nucleus, the energy also increases.

6. Sublevels- • Each energy level is broken into sublevels (denoted by the letter “l”). • There are four available sublevels- s, p, d, and f.

7. Sublevels- • “s” can hold up to 2 electrons • “p” can hold up to 6 electrons • “d” can hold up to 10 electrons • “f” can hold up to 14 electrons

9. Electron Configuration- • Is the arrangement of electrons in an atom. [It is like writing an address for the electrons in their energy levels.] Note: Remember the number of electrons that are being counted come from the atomic number…if the atom is neutral, the Protons = electrons

10. Energy Diagram

12. Rules to using the Energy Diagram : • Always start at the lowest energy level (1s). This is known as the Aufbau Principle. • Follow the diagonal arrows after leaving 1s. • You must fill up a sublevel before moving on to the next available sublevel.

13. Guided Practice: Write out the electron configuration for magnesium, phosphorus, zinc, and bromine.

14. Drawing orbital notation- • Follow the energy diagram to write out the electron configuration. • Under the configuration, draw the boxes to represent the orbitals. • Use boxes to represent an orbital (where the electrons are found within the sublevels.)

15. Drawing orbital notation- • Only 2 electrons can fit in an orbital and they must having opposite spins (called the Pauli Exclusion Principle.) • Electrons will be represented with arrows, ↑ and ↓ .

16. “s” can hold up to 2 electrons, it will have 1 box • “p” can hold up to 6 electrons, it will have 3 boxes • “d” can hold up to 10 electrons, it will have 5 boxes • “f” can hold up to 14 electrons, it will have 7 boxes

17. Hund’s Rule states that equal energy electrons must be spread out evenly in the orbital (boxes). In other words, place all up arrows in the boxes before filling in with the down arrows. See pg.136.

18. Guided Practice: Using the electron configuration for Mg, P, Zn, and Br, draw the orbital notation for each.

19. Valence Electrons • The outermost electrons; • involved in chemical bonding (we will use these in later chapters to form ions and form bonds); • represented by the electrons found in the “s” and “p” sublevels.

20. Valence Electrons • After writing the electron configuration out, pick the highest energy level you have written… • The electrons found in these levels will be either “s” or “s and p” electrons… • These are the valence electrons…never exceeding a maximum of 8…(octet rule).

21. Valence Electrons • A quick way to tell the valence electrons of the main group elements (the tall columns) is to look at the roman numeral of the group…

22. Valence Electrons Ex. I A sodium….1 valence electron II A magnesium…2 valence electrons III A aluminum…3 valence electrons IV A carbon…4 valence electrons V A nitrogen…5 valence electrons VI A oxygen…6 valence electrons VII A fluorine…7 valence electrons VIII A neon…8 valence electrons