Electron Configurations

1. Electron Configurations Chapter 5 Section 3

2. Vocabulary • electron configuration • aufbau principle • Pauli exclusion principle • Hund’s rule • Valence electron • Electron-dot structure

3. Review • Each principal energy level can have the same number of sublevels as the level number • Each sublevel orbital has a different shape • Each orbital can have only 2 electrons • What does this mean?

4. Electron Arraignment follows Rules • Low energy level systems are more stable than high-energy systems • Atoms will assume the electron arrangement that gives the atom the lowest energy • Most stable is the “ground state” (lowest energy) • Three rules/principles for arranging electrons • Aufbau • Pauli Exclusion principle • Hund’s Rule

5. Aufbau Principle • Each electron occupies the lowest energy level available • Learn the sequence of atomic orbitals from lowest to highest:

6. Aufbau Diagram #4 – also do electron DOT diagram, #7 & 8 – SHOW YOUR WORK

7. Aufbau Expanded

8. Using Aufbau • All orbitals related to the same energy level are of equal energy • All 2p orbitals have the same energy • In a multi-electron atom, the energy sublevels within a principal energy level have different energies: • 2p orbitals are higher energy than 2s orbitals • The sequence of sublevels within a principle level in increasing energy is: s, p, d, and f • Orbitals related to energy sublevels within one principle energy level can overlap orbitals related to another principal level • Notice: 4s is lower than 3d

9. Pauli Exclusion Principle • Each electron has an associated spin, like a top • Can spin on bottom or top • Each orbital can hold AT MOST 2 electrons, but only if they have opposite spin • Designated as:

10. Hund’s Rule • Because negatively charged electrons repel each other, they try to get as far away from each other as possible, therefore: • Single electrons with the same spin will occupy each equal energy orbital before additional electrons with opposite spins occupy the same orbitals. • WHAT?

11. 2 1 3 4 Hund’s Rule

12. Representing Atom’s Electrons • Orbital Diagrams • Electron Configuration Notations

13. Orbital Diagrams • Draw one box for each orbital • Empty box means no electrons • Box with single up arrow means orbital with one electron • Box with two arrows (up/down) means orbital with two electrons • Each box is labeled with the principle quantum number and the sublevel • Example:

14. Example Orbital Diagram - N

15. Example Orbital Diagram – C & O

16. Electron Configuration Notation • Specify principal energy level and energy sublevel • Use a superscript to represent the number of electrons • Use the Aufbau diagram to help you remember the level/sublevel filling order • Examples:

17. Electron Configuration Notation

18. Electron Configuration Notation

19. A Variation – Noble Gas Notation • Similar to Electron Configuration Notation, but: • Start with the previous noble gas, and put that symbol in brackets • Add electrons to the noble gas • Example: • Sodium (Na) = [Ne] 3s1 • Ne = 1s2 2s2 2p6 • Na = 1s2 2s2 2p6 3s1

20. Exceptions to Aufbau • Cr (Chromium) = [Ar] 4s1 3d5 • Cu (Copper) = [Ar] 4s1 3d10

21. Valence Electrons • Only Valence Electrons contribute to chemical bonding • Valence Electrons = electrons in outer most principal energy level • Generally highest energy • Examples: • Carbon: [He] 2s2 2p2 has 4 valence electrons • Sulfur: [Ne] 3s2 3p4 has 6 valence electrons

22. Electron Dot Diagrams • Also called Lewis Electron Dot Diagrams or Lewis Electron Dot Structure or Electron Dot Structure • Shows atomic symbol and ONLY valence electrons

23. Electron Dot Diagrams

24. Periodic Table