2. Electronegativity ++ –– 00 00 HClHH

3. The basic units: ionic vs. covalent Ionic compounds form repeating units. Covalent compounds form distinct molecules. Consider adding to NaCl(s) vs. H 2 O(s): H O H Cl Na Cl Na H O H H O H NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms. H 2 O: O and H cannot add individually, instead molecules of H 2 O form the basic unit.

4. 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: Point: the bonding electrons are shared in each compound, but are not always shared equally. The greek symbol  indicates “partial charge”. H2H2 HClLiCl ++ –– 00 00 +– covalent ( non-polar ) polar covalentionic HH H Cl [Li] + [ Cl ] –

5. Electronegativity 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 NBr 3 :  EN = 3.0 – 2.8 = 0.2 (for all 3 bonds).

6. 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, Br 2, H 2 O, CH 4, KCl

7. Electronegativity Answers HCl:3.0 – 2.1 = 0.9 polar covalent CrO:3.5 – 1.6 =1.9ionic Br 2 :2.8 – 2.8 =0covalent H 2 O:3.5 – 2.1 =1.4polar covalent CH 4 :2.5 – 2.1 =0.4covalent KCl:3.0 – 0.8 =2.2ionic

8. Holding it together Q:Consider a glass of water. Why do molecules of water stay together? A:there must be attractive forces. Intramolecular forces occur between atoms Intermolecular forces occur 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. Intramolecular forces are much stronger

9. Inter-molecular Forces (IMF’s) London Dispersion Forces Dipole – Dipole Forces Hydrogen Bonding

10. 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

11. 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.

12. 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

13. 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. EXAMPLES: 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.)

14. 1.PbS 2.HClO 4 3.Zn(OH) 2 4.HBr(aq) 5.(NH 4 ) 2 CO 3 6.sulfur hexafluoride 7.nitrous acid 8.hydrogen chloride 9.lead(II) chloride 10.zinc sulfate Name / Formulate the following Compounds 1.Lead (II) Sulfide 2.Hydrogen Perchlorate 3.Zinc Hydroxide 4.Hydro Bromic Acid 5.Ammonium Carbonate 6.SF6 7.HNO2(aq) 8.HCl 9.PbCl2 10.ZnSO4