Chemistry ELECTRONIC STRUCTURE Electron Configurations

1. ChemistryELECTRONIC STRUCTUREElectron Configurations DO NOW: what are the four quantum numbers? Calculators not required Periodic tables not required

2. Quantum Numbers and electron placement • From last time… • No two electrons may be doing the same two things: Pauli exclusion principle • Each orbital can hold two electrons • s subshell: ____ electrons • p subshell: ____ electrons • d subshell: ____ electrons • f subshell: ____ electrons • So, how do we determine which electrons go in which orbitals? • Works best if we think of orbitals as rooms in a hotel…

3. AUFBAU PRINCIPLE • The Aufbauprinciple states that electrons will enter orbitals in order from lowest energy to highest energy. • If you are booking a room in a hotel, you would prefer to spend less money for your room. • The “atomic currency” is energy. Less energy to spend is preferred.

4. Orbital Splitting • The Aufbau principle is necessary because the orbitals split in energy when there are many electrons repelling against other in the atom. • This makes some orbitals in an energy level higher in energy than others. • Or, some rooms on the same floor cost more to book.

5. Atomic “price schedule” • To the right is the energy “price schedule” of the atomic orbitals. • To continue our analogy… • Energy level = floor # • Sublevel = type of room (suite, penthouse, deluxe, fantastic) • Orbital = single room • Spin = bed in room • 2 beds per room

6. Predicting the price schedule • The Aufbau principle has a set pattern. • Following and re-creating the diagram to the right will help you memorize the “price schedule” and which orbitals are lower in energy. • 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

7. Orbital BOX Diagrams • Each box in the diagram represents one orbital. • Half-arrows represent the electrons. • The direction of the arrow represents the relative spin of the electron.

8. Hund’s Rule • The lowest energy is attained when the number of electrons with the same spin is maximized. • Make all orbitals in a subshell (s, p, d, etc.) half-filled before adding the opposite spin electrons. • Half-filled subshells are especially stable. • In other words… • Electrons will always pick the lowest price room, but will prefer to have single rooms until they are forced to double up.

9. Drawing Orbital Box Diagrams • Follow the atomic “price schedule” for oxygen:

10. Magnetism • Magnetism results from unpaired electrons in the ground state (as opposed to the excited state, like we do in the hydrogen gas light). • When there are multiple unpaired electrons (like in iron), this results in a ferromagnet, or permanent magnet, or a paramagnet, or temporary magnet.

11. Practice • Draw the full orbital box diagram for iron. • See if you can tell whether or not iron is magnetic.

12. Anatomy of an Electron Configuration • An electron configuration gives the orbital placement of every electron in an atom. • This is one piece of an electron configuration.

13. Electron Configurations • Each component consists of • A number denoting the energy level;

14. Electron Configurations • This shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital,

15. Electron Configurations • This shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital, • A superscript denoting the number of electrons in those orbitals.

16. Sample Electron Configuration • Remember, we can only fit so many electrons in each type of orbital, and that each energy level only has certain subshells: • n =1 only has s 2 electrons in s • n = 2 can have s and p 6 in p • n = 3 has s, p, and d 10 in d • n = 4 has s, p, d, f 14 in f • You keep putting electrons in orbitals from low to high energy until you run out. Example: Neon (10 electrons) • 1s22s22p6 (total of 10 electrons)

17. Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table (shaded in different colors here) correspond to different types of orbitals.

18. Practice • Give the ground-state electron configuration for sodium. • Note: ground-state is when every electron is at its lowest energy orbital. If an electron is at a higher energy orbital, we say it is in an excited state, because energy had to be put in to make it go up.

19. Shortened Electron Configurations • As a shortcut, we can use a noble gas (He, Ne, Kr, etc..) as a starting point for the electron configuration. • Sodium’s full ground-state electron configuration: • 1s22s22p63s1 • Neon: 1s22s22p6 • Shortened ground-state electron configuration: • [Ne]3s1

20. Practice • Give the shortened (noble gas) ground-state electron configurations for uranium.

21. Exception: half-filled and filled d orbitals • When a subshell is half-filled or completely filled, it lowers the overall energy of the atom. • In other words, this is favored. • This usually doesn’t effect electron configurations, except in the d and f subshells. • Example: Copper (29 electrons) • Normal electron configuration: 1s22s22p63s23p64s23d9 • Added stability if 3d is full (electron gets promoted to 3d from 4s) • 1s22s22p63s23p64s13d10

22. Exception: half-filled and filled d orbitals • The same happens with half-filled d orbitals: • Example : Chromium • 1s22s22p63s23p64s23d4 not as stable as a half-full d-subshell • Becomes: 1s22s22p63s23p64s13d5 • In summary: • Configurations that end in s2d9 become s1d10. • Configurations that end in s2d4 become s1d5. • This should also be reflected in orbital box diagrams.