1. Molecular Compounds and Acids

2. Molecular (covalent bonding) compounds A covalent bond results from the sharing of electrons. The octet rule still applies Covalent bonds generally occur when elements are close to each other on the periodic table. The majority of covalent bonds form between nonmetallic elements. (remember that ionic bonds form between metals and non metals)

3. Example of a covalent bond

4. Octet Rule w/ covalent bonding

5. Properties of molecular compounds Usually have a lower MP (than ionic, metallic) Generally soft Non-conductors (in any state & aqueous) Molecular Compounds can exist in all 3 states: Solids – sugar, ice, aspirin Liquids – water, alcohols Gases – O 2, CO 2, and N 2 O (laughing gas) Very important to organic chemistry, biochemistry, and pharmaceutics.

6. Naming Molecular (covalent bonding) Compounds 1.The 1 st element is named first, using the entire element name 2.The 2 nd element is named using the root of the element & adding the suffix –ide 3.Prefixes are used to indicate the # of atoms of each type that are present. *Exception: The 1 st element in a formula never uses the prefix mono-

7. Prefixes in Molecular (covalent) Compounds # of atoms Prefix# of atoms Prefix 1mono-6hexa- 2di-7hepta- 3tri-8octa- 4tetra-9nona- 5penta-10deca-

8. Examples NH 3 nitrogen trihydride (ammonia) N 2 H 4 dinitrogen tetrahydride H 2 Odihydrogen monoxide (common name water)

9. Practice CO 2 Carbon dioxide CO Carbon monoxide PCl 5 Phosphorus pentachloride N 2 O 5 Dinitrogen pentoxide NOTE – Be sure to drop the last vowel of the prefix if there would be any a-o, o-o, or a-a combinations (pentoxide, not pentaoxide).

10. Practice Dichlorine monoxide Cl 2 O Tetraiodine nonoxide I 4 O 9 Sulfur hexafluoride SF 6

11. Diatomic Elements (elements that exist in pairs) HydrogenH 2 OxygenO 2 NitrogenN 2 FluorineF 2 ChlorineCl 2 BromineBr 2 IodineI 2

12. Ionic vs. Covalent Look at the 1 st element: 1 st element metal: ionic compound 1 st element non-metal: covalent compound Exception: ammonium (NH 4 + ): ionic compound

13. Practice AlCl 3 : Al is a metal, so ionic name Aluminum chloride SiCl 4 : Si is a non-metal, so covalent name Silicon tetrachloride

14. (Arrhenius) Acids Arrhenius acids are compounds which lose H + ion in H 2 O. The general form for an acid is “HA” where H is hydrogen and A is either a monoatomic or polyatomic anion. Here are the rules for naming acids:

15. 1.If the anion part normally ends in –ide (binary acid), then the acid name begins with the prefix hydro and ends with –ic. Ex. HCl is hydrochloric acid 2.If the anion part ends in –ate (polyatomic) then NO hydro is used and the ending is –ic. Ex. HNO 3 is nitric acid (notice – no hydro). 3.If the anion part normally ends in –ite no hydro is used and the ending is –ous. Ex. HNO 2 is nitrous acid

16. For a couple of polyatomic ions, part or even none of the ending is dropped before adding the acid ending. H 2 SO 4 is sulfuric acid, not sulfic acid H 3 PO 4 is phosphoric acid, not phosphic acid

17. Acid formulas Acid formulas always start with an H. To write formulas for acids just use the number of H’s as subscripts equal to the negative charge of the anion (since each H is +1). Ex. Carbonic acid – no hydro is used so the anion must be polyatomic. The acid name ends in –ic so the anion must end in –ate, i.e. carbonate. Since carbonate is CO 3 2- two H’s are necessary and the formula is H 2 CO 3.

18. Practice HBr: Anion ends in –ide, becomes -ic: hydro added Hydrobromic acid Acetic acid Anion ends in –ic (no hydro), anion must be –ate: acetate; Charge on acetate = -1, so HC 2 H 3 O 2