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PowerPoint Answers to Review - Naming Chemical Compounds - PowerPoint - PowerPoint


  • pg 1
+   –    0   0

H    Cl     H   H
     The basic units: ionic vs. covalent
• Ionic compounds form repeating units.
• Covalent compounds form distinct molecules.
• Consider adding to NaCl(s) vs. H2O(s):

       Na Cl Na Cl                       O
                                     H        H
      Cl Na Cl Na                              O
                                             H   H

• NaCl: atoms of Cl and Na can add individually
  forming a compound with million of atoms.
• H2O: O and H cannot add individually, instead
  molecules of H2O form the basic unit.
   I’m not stealing, I’m sharing unequally
• We described ionic bonds as stealing electrons
• In fact, all bonds share – equally or unequally.
• Note how bonding electrons spend their time:

H2      H H         HCl      H Cl         LiCl [Li]+[    Cl ]–

   0     0            +     –           +        –

  (non-polar)        polar covalent             ionic
• Point: the bonding electrons are shared in each
  compound, but are not always shared equally.
• The greek symbol  indicates “partial charge”.
• Recall that electronegativity is “a number that
  describes the relative ability of an atom, when
  bonded, to attract electrons”.

• The periodic table has electronegativity values.

• We can determine the nature of a bond based on
  EN (electronegativity difference).

• EN = higher EN – lower EN
  NBr3: EN = 3.0 – 2.8 = 0.2 (for all 3 bonds).
                     Bond Type

Basically: a EN below 0.5 = non-polar covalent,
  0.5 - 1.7 = polar covalent and above 1.7 = ionic

Determine the EN and bond type for these: HCl, CrO,
 Br2, H2O, CH4, KCl
       Electronegativity Answers

HCl: 3.0 – 2.1 = 0.9     polar covalent
CrO: 3.5 – 1.6 = 1.9     ionic
Br2:   2.8 – 2.8 =   0   covalent
H2O: 3.5 – 2.1 = 1.4     polar covalent
CH4: 2.5 – 2.1 = 0.4     covalent
KCl: 3.0 – 0.8 = 2.2     ionic
                  Holding it together
Q: Consider a glass of water. Why do molecules
   of water stay together?
A: there must be attractive forces.

                                               forces are much

Intramolecular forces occur       Intermolecular forces occur
      between atoms                   between molecules

• We do not consider intermolecular forces in ionic bonding
  because there are no molecules.
• The type of intramolecular bond determines the type of
  intermolecular force.
Inter-molecular Forces (IMF’s)
• London Dispersion Forces

• Dipole – Dipole Forces

• Hydrogen Bonding
     London Dispersion Forces
• London dispersion forces result from
  instantaneous non-permanent dipoles created by
  random electron motion. London dispersion forces
  are present in all molecules and are directly
  proportional to molecular size.

• LDF’s are the weakest of the IMF’s
         Dipole – Dipole Forces
• Dipole-dipole interaction is the
  attraction between a partially
  negative portion of one molecule
  and a partially positive portion of a
  nearby molecule.
• Dipole-dipole interaction occurs in
  any polar molecule as determined by
  molecular geometry.
             Hydrogen Bonding
• Hydrogen bonding is the unusually
  strong dipole-dipole interaction that
  occurs when a highly
  electronegative atom (N, O, or F) is
  bonded to a hydrogen atom.

• Hydrogen bonding is the strongest
  of the 3 IMF’s
 Effects of IMF’s on Physical Properties
• The strength of intermolecular forces present in a substance
  is related to the boiling point and melting point of the
  substance. Stronger intermolecular forces cause higher
  melting and boiling points.

• CH4 - Methane: has only very weak London dispersion
  forces (lowest b.p. & m.p.)
• CHCl3 - Chloroform: has dipole-dipole interaction (moderate
  b.p. & m.p.)
• NH3 - Ammonia: has hydrogen bonding and dipole-dipole
  interaction (high b.p. & m.p.)
Name / Formulate the following Compounds

1.    PbS                   1.    Lead (II) Sulfide
2.    HClO4                 2.    Hydrogen Perchlorate
3.    Zn(OH)2               3.    Zinc Hydroxide
4.    HBr(aq)               4.    Hydro Bromic Acid
5.    (NH4)2CO3             5.    Ammonium Carbonate
6.    sulfur hexafluoride   6.    SF6
7.    nitrous acid          7.    HNO2(aq)
8.    hydrogen chloride     8.    HCl
9.    lead(II) chloride     9.    PbCl2
10.   zinc sulfate          10.   ZnSO4

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