1. Correlation with Electronic Structures
   
   1. The elements in a given group have the same outer electron configuration group.

      1	2	3	4	5	6	7	8
      ns1	ns2	ns2p1	ns2p2	ns2p3	ns2p4	ns2p5	ns2p6

      
      Where n = quantum number of outermost level = period number of the element.

      Examples:
      Au = [Xe] 6s² 4f¹⁴ 5^d9
      Cu = [Ar] 4s¹ 3d¹⁰

   2. Elements in groups 1 and 2 fill the s sublevel
      Elements in groups 3 thru 8 fill the p sublevels
      Transition metals (10 in each series) fill the d sublevels
      Lanthanides (14 in series) fill 4f sublevel
      Actinides (14 in series) fill 5f sublevel
      

2. Atomic Radius
   
   Measure of the size of the atom. Atoms are assumed to be spherical (combination of all orbitals).

   Atomic Radius: defined as the distance between bonded atoms (i.e. Cl₂)

   1. In general, atomic radius decreases going across a period from left to right, increases going down a group.
   2. Horizontal trends can be expressed in terms of effective nuclear charge experienced by outer-most electrons.

      Effective nuclear charge = z - s
      where z is the nuclear charge and s (shielding effect) is the number of inner electrons.
      The greater the effective nuclear charge, the stronger the attractive force on the electrons (decreasing atomic radii).

3. Ionization Energy
   
   Penetration effect s > p > d > f
   E =
   
   1. First ionization energy is equal to the energy that must be absorbed to convert an atom to a +1 ion.

      Na_(g) Na⁺_(g) + e^-: 496
      Hydrogen: 1312

      1. Ionization Energy increases as you go across the Periodic Table to the right.
      2. Ionization Energy decreases as you go down the Periodic Table.

   2. Successive ionization energies increase steadily.
      Large jumps in ionization energy occur when electron is removed from inner level (*chart on page 321)

4. Metallic Character
   
   1. Metallic character decreases across the periodic table from left to right.
      Metallic Character increases going down the periodic table.
      
   2. Physical Properties of metals (80% applicable)
      
      1. High electrical conductivity, Metals (compared to nonmetals) have a large atomic radii and small ionization energy.
      2. High thermal conductivity
      3. Ductile (Can be pulled into wires), malleability(can be pounded into sheets)
      4. Luster
   3. Flame Tests

5. Electron Affinity
   E to add an electron, is negative for nonmetals. See page 324.*
   
   Electron affinity increases going up the Periodic Table.
   Electron affinity increases going across the Periodic Table from left to right.

6. Electronegativity increases across the periodic table from left to right. It increases going up the periodic table. See Chapter 8: Chemical Bonding for more about electronegativity
   
   

*All page references are from Chemistry Third Edition by Steven S. Zumdahl.

This page was made by Erik Epp.