1. Molecules Molecule – two or more atoms covalently bound together Diatomic molecule – two of the same atom bound together

2. Diatomic Molecules Br I N Cl O F H or the Gang of 7 plus 1 These atoms never exist alone. They need a buddy They always come in pairs For example: –Br  Br 2 –I  I 2 –N  N 2 –Cl  Cl 2 –H  H 2 –O  O 2 –F  F 2

3. Binary Molecular Compounds Binary covalent compounds contain 2 nonmetals This is different from ionic compounds that contain a metal & nonmetal, metal & a polyatomic ion, or 2 polyatomic ions

4. Naming Binary Covalent Compounds Before you can name binary covalent compounds, you MUST know the prefixes!

5. Prefixes Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca 1 2 3 4 5 6 7 8 9 10

6. Rules for Writing Covalent Compounds 1.Find the cation (nonmetal) and write down symbol 2.Look at the prefix and write it as a subscript after the symbol 3.Find the anion (nonmetal) and write down symbol 4.Look at the prefix and write it as a subscript after the symbol *NO CHARGES *NO REDUCING *If there is no prefix, then it only a 1  don’t write 1’s

7. Examples 1.Carbon dioxide 2.Phosphorus trifluoride 3.Nitrogen monobromide 4.Hexaselenium pentaiodide 5.Dicarbon monosulfide CO 2 PF 3 NBr Se 6 I 5 C 2 S

8. Naming Covalent Compounds 1.Name the prefix for # of atoms in the 1 st element -If the prefix is mono, drop it 2. Write the name for the 1 st element 3. Name the prefix for the # of atoms in the 2 nd element 4. Take the root name of the 2 nd element and add “ide” *NO CHARGES

9. Examples 1. SO 2 2. N 2 F 5 3. CO 4. S 3 Cl 4 5. FI 3 Sulfur dioxide Dinitrogen pentafluroide Carbon monoxide Trisulfur tetrachloride Flurorine triiodide

10. Hydrates Some compounds trap water crystals when they form. These are hydrates. Both the name and the formula needs to indicate how many water molecules are trapped. In the name we add the word hydrate with a prefix that tells us how many water molecules.

11. Hydrates In the formula you put a dot and then write the number of molecules. Calcium chloride dihydrate = CaCl 2  2  Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O

12. Acids Acids are in aqueous solution (aq) For the purposes of this class, we will assume that if it begins with H, we will name it according to the rules of naming acids

13. Remember… ide  hydro (root) + ic acid ate  (root) + ic acid ite  (root) + ous acid

14. Rule #1 - naming acids If the anion ends in –ide, the acid will be named… Hydro (root) – ic acid This is usually for H plus one element

15. For example HCl Hydrochloric acid HI Hydroiodic acid H 2 S Hydrosulfuric acid

16. Rule #2 – naming acids If you have an H plus an anion ending in –ate, the acid will be named… (root) – ic acid

17. Examples H 2 SO 4 Sulfuric acid HNO 3 Nitric acid H 3 PO 4 Phosphoric acid

18. Rule # 3 – naming acids If you have an H plus an anion ending in –ite, the acid will be named… (root) – ous acid

19. Examples H 2 SO 3 Sulfurous acid HNO 2 Nitrous acid H 3 PO 3 Phosphorous acid

20. Writing formulas for acids When writing formulas for acids you MUST look at the charges and bring them down!

21. Examples HBr Hydrogen + one element Hydrobromic acid HClO 3 H + chlorate ate  ic Chloric acid

22. More examples H 2 SO 3 H 2 CO 3 HF Nitrous acid Perchloric acid Iodic acid Sulfurous acid Carbonic acid Hydrofluoric acid HNO 2 HClO 4 HIO 3

23. Mixed examples (remember to figure out what type of compound it is 1 st !) KClO 2 CO 2 H 2 SO 4 NH 4 Br CuCO 3 Fe 2 O 3 HClO Potassium chlorite Carbon dioxide Sulfuric acid Ammonium bromide Copper (II) carbonate Iron (III) oxide Hypochlorous acid

24. More Mixed Examples Carbon tetrachloride Phosphorous pentachloride Aluminum oxide Copper (II) nitrate Chlorous acid Hydrophosphoric acid Iron (III) hydroxide CCl 4 PCl 5 Al 2 O 3 Cu(NO 3 ) 2 HClO 2 H 3 P Fe(OH) 3