Chem133 Chapter8

  1. What was periodic law?
    • In 1870 Mendeleev saw that arranging elements by atomic mass they exhibit a periodic recurrence of similar properties.
    • Today's periodic table arranges them by atomic number.
  2. What is the difference between the spin quantum number and the other three?
    • The spin quantum number describes the electron, rather than the orbital
    • The spin number can be either ±Image Upload 1
  3. What is Hund's rule?
    When orbitals of equal energy are available, the electron configuration of lowest energy has the maximum number of unpaired electrons with parallel spins.
  4. What is the order you fill orbitals?
    2s 2p
    3s 3p 3d
    4s 4p 4d 4f
    5s 5p 5d 5f
    6s 6p 6d
    7s 7p
    • 1s->2s->2p->3s->3p->4s->3d->4p->5s->4d->5p->6s->4f->5d->6p->7s->5f->6d->7p
    • (Go diagonally down left)
  5. Define inner (core) electrons.
    Those electrons in common with the previous noble gas.
  6. Define outer electrons.
    Those in the highest energy level (highest n value.) They spend most of their time farthest from the nucleus.
  7. Define Valence Electron.
    • Electrons involved in bonding:
    • For main group electrons they are the outer electrons.
    • For transition elements, in addition to the outer ns electrons the (n-1)d electrons are also valence electrons.
  8. What happens to the filling order when you start to fill 3d orbitals?
    • When you add an electron to a 4s2 3p3 you obtain 4s1 3d6
    • Electrons will jump between sub-levels to fill each orbital with parallel spinning electrons.
    • Ths also holds true for f orbitals.
  9. Define metallic radius.
    • one-half the distance between nuclei of adjacent, individual atoms in a crystal of the element.
    • Used mainly for metals.
  10. Define covalent radius.
    • One-half the distance between elements which occur in molecules. 
    • Used for mainly for nonmetals.
    • You can use a known covalent radius to find an unknown.
  11. Describe trends in atomic radii for main group elements.
    • Radii increase going down a group, and decrease going across a period.
    • (transition metals experience similar trends, but are extremely less profound.)
  12. Define ionization energy.
    • The energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions.
    • When removing more than 1 electron: IE1< IE2 < IE3 <...
  13. Explain the trends of ionization energies in the periodic table.
    • As you move down a group, IE decreases because there are more inner electrons shielding the electron being removed.
    • As you move across a period, IE increases because the effective charge the electrons feel from the nucleus increases. Shielding is more dramatic from inner electrons than electrons of the same energy level.
  14. Define Electron Affinity.
    • Energy change accompanying the addition of 1 mol of electrons to 1 mol of gaseous atoms.
    • Most EA1's are negative because the electron is attracted to the nucleus (producing an exothermic reaction.)
  15. Why are noble gases stable?
    Their highest energy level is filled (ns2np6nd10...)
  16. Define isoelectronic.
    They share the same electron configuration as the nearest noble gas.
  17. Why do things like Na2+, Ca3+, F2-, or S3- not exist?
    If you attain an isoelectronic configuration, you can't fuck with it because it's too stable, so removing or adding electrons from a noble-gas configuration takes a ton (too much) of energy.
  18. What is a pseudo-noble gas configuration?
    • When an atom loses all of its highest energy level electrons.
    • SnImage Upload 2Sn4+: [Kr]5s24d105p2Image Upload 3[Kr]4d10
  19. Define inert pair configuration.
    • It leaves the ns2 alone
    • Image Upload 4
Card Set
Chem133 Chapter8