Chapter 9: Intermolecular Attractions and the Properties of Liquids and Solids

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CHEMISTRY The Molecular Nature of Matter and
Change 3rd Edition
CHAPTER 8 LECTURE NOTES Electron Configuration
and Chemical Periodicity Chem 150 - Ken Marr -
Winter 2003
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8.2 Characteristics of Many-Electron Atoms
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Figure 8.1
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The 4th quantum number, ms

• Electrons can Spin in two directions (like a
  top)
• ms spin quantum number
• can only have two values
• ms -1/2 or 1/2
• Pauli Exclusion Principle
• No 2 electrons can have the same 4 quantum
  numbers
• i.e. Only two electrons per orbital

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Application of the Pauli Exclusion Principle

• How many electrons can occupy a....
• 2s subshell?
• 3d subshell?
• shell with n 3?
• shell with n 4?
• subshell with n 3 and l 1?

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Electronic Structure of Multi-electron Atoms

• Aufbau Principle
• Electrons 1st fill orbitals of lowest energy
• Orbitals of a sublevel half fill before electrons
  start to pair
• Orbital Notation vs Electronic Configuration
• Examples
• Noble gases, Halogens, Alkali Metals, etc.
• Abbreviated electronic Configurations

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Memory Aid for the order of sublevel filling
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Electronic Configurations and the Periodic Table

• Use the periodic table to remember the order of
  orbital filling
• Sublevels found in each
• Group
• Period
• Write Abbreviated electron configurations for
• Mg
• Fe

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Valence Shell Configurations

• Valence Shell Outermost shell (E-level)
• Write the valence shell configurations
  (abbreviated format) for
• Alkali metals
• Halogens
• Group VA
• 26Fe, 39Y

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Unexpected Electron Configurations

• Fully filled and half filled sublevels offer
  stability
• s and d sublevels of the transition metals are
  very close in energy
• Predict which transition elements in the 4th
  period have unexpected electronic configurations.
  
• Write them in abbreviated format

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Periodic Trends in Atomic Radius

• Effective Nuclear Charge, Zeff
• Zeff (Atomic number) (Number of Core
  Electrons)
• Valence electrons do not feel the full effect of
  the nucleus because of screening by core
  electrons
• e.g. Group 1A H ? Fr
• Explain what happens to atomic radius...
• Down a group? Which is more important, n or Eeff?
• Across a period? Which is more important, n or
  Eeff?
• What about the transition metals?

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C Cl Bond
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Size of an Atom vs. its Ion

• How does the size of an atom compare to that of
  its Cation?
• e.g. K vs K
• How does the size of an atom compare to that of
  its Anion?
• e.g. F vs F-

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Ionization Energy

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

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Periodic Trends in Ionization Energy

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

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Magnetic Properties of Atoms

• Paramagnetism
• Due to the spinning of unpaired electrons
• Atoms are attracted to a magnet
• Diamagetism
• Atoms have all electrons paired
• Atoms not attracted to a magnet
• Electric Motors
• Generate a magnetic field by moving electrons in
  curved paths within a coil

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Electron Affinity

• Definition
• Exothermic.......why?
• The most nonmetallic elements have the most
  negative electron affinities
• Equation
• How does electron affinity vary ....
• Across a Period? Why?
• Down a Group? Why?

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Increases
Increases
Ionization Energy