1. CHEMISTRY The Molecular Nature of Matter and Change 3 rd Edition CHAPTER 8 LECTURE NOTES Electron Configuration and Chemical Periodicity Chem 150 - Ken Marr - Winter 2003

3. 8.2 Characteristics of Many-Electron Atoms

4. Figure 8.1

5. The 4th quantum number, m s 1. Electrons can Spin in two directions (like a top) 2. m s = spin quantum number »can only have two values m s = -1/2 or +1/2 3. Pauli Exclusion Principle »No 2 electrons can have the same 4 quantum numbers »i.e. Only two electrons per orbital

6. Application of the Pauli Exclusion Principle 1. How many electrons can occupy a.... a.2s subshell? b.3d subshell? c.shell with n = 3? d.shell with n = 4? e.subshell with n = 3 and l = 1?

14. Electronic Structure of Multi-electron Atoms 1. Aufbau Principle Aufbau Principle »Electrons 1st fill orbitals of lowest energy »Orbitals of a sublevel half fill before electrons start to pair 2. Orbital Notation vs Electronic Configuration Orbital Notation vs Electronic Configuration »Examples: – Noble gases, Halogens, Alkali Metals, etc. 3. Abbreviated electronic Configurations Abbreviated electronic Configurations

15. Memory Aid for the order of sublevel filling

16. Electronic Configurations and the Periodic Table 1. Use the periodic table to remember the order of orbital filling Use the periodic table to remember the order of orbital filling 2. Sublevels found in each… a.Group b.Period 3. Write Abbreviated electron configurations for a.Mg b.Fe

20. Valence Shell Configurations 1. Valence Shell = Outermost shell (E-level) 2. Write the valence shell configurations (abbreviated format) for a.Alkali metals b.Halogens c.Group VA d. 26 Fe, 39 Y

21. Unexpected Electron Configurations 1. Fully filled and half filled sublevels offer stability 2. s and d sublevels of the transition metals are very close in energy 3. Predict which transition elements in the 4th period have unexpected electronic configurations. »Write them in abbreviated format

22. Periodic Trends in Atomic Radius 1. Effective Nuclear Charge, Z eff Z eff = (Atomic number) – (Number of Core Electrons) 2. Valence electrons do not feel the full effect of the nucleus because of screening by core electrons e.g. Group 1A: H  Fr 3. Explain what happens to atomic radius... » Down a group? Which is more important, n or E eff ? » Across a period? Which is more important, n or E eff ? » What about the transition metals?

24. C – Cl Bond

28. Size of an Atom vs. its Ion 1. How does the size of an atom compare to that of its Cation? e.g. K vs K + 2. How does the size of an atom compare to that of its Anion? e.g. F vs F -

32. Ionization Energy 1. Ionization Energy »Definition »Endothermic......Why? 2. Equations for »1st Ionization Energy »2nd Ionization Energy »3rd Ionizationenergy

36. Periodic Trends in Ionization Energy 1. Across a Period....Explain why! 2. Down a Group......Explain why! 3. Compare the relative 1st, 2nd, 3rd, etc. Ionization Energies for.... »elements in period 2 »elements in period 3

45. Magnetic Properties of Atoms 1. Paramagnetism »Due to the spinning of unpaired electrons »Atoms are attracted to a magnet 2. Diamagetism »Atoms have all electrons paired »Atoms not attracted to a magnet 3. Electric Motors »Generate a magnetic field by moving electrons in curved paths within a coil

49. Electron Affinity 1. Definition »Exothermic.......why? »The most nonmetallic elements have the most negative electron affinities 2. Equation 3. How does electron affinity vary.... »Across a Period? Why? »Down a Group? Why?

53. Ionization Energy Increases