1. Intro to Chemical Bonding Ch. 5 and Ch. 12

2. Unit Outline  Ch. 5- Nomenclature Naming chemical compounds and molecules  Ch. 12- Chemical Bonding Types of chemical bonds Stable electron configurations Lewis structures VSEPR Theory--molecular geometry

3. Nomenclature: Learning the Language  In the past, there was no system to name chemical compounds Common names used  Quicksilver  Laughing Gas  Need a systematic way to name chemical compounds

4. Naming Binary Compounds  Ionic compounds: a compound that contains ions  Ion: A charged atom Cation: Positively charged atom (lost 1 or more electrons) Anion: Negatively charged atom (gained 1 or more electrons)  Metals usually form cations (lose electrons), nonmetals usually form anions (gain electrons)

5.  Cations take the name of the parent element  Ex) Sodium metal Na(s) vs sodium ion Na +  Single atom anions: change element ending to -ide. The chlorine anion, Cl - is named the chloride ion

6. How do you get the charge?  Look to the P.T.  For metals (not transition metals), Group # (or Group # - 10) is the positive charge  For nonmetals (18 - Group # is negative charge)  Transition metals can have multiple positive charges

7. Naming Type I ionic compounds  Type I: Only one charge on the cation  Rules:  1. The cation is always listed first  2. The cation takes the name of the element  The anion: take of ending of element and add -ide.

8. Examples  NaCl  KI  CaS  CsBr  sodium chloride  potassium iodide  calcium sulfide  cesium bromide

9. Naming Type II Ionic Compounds  Type II: Cation can have multiple charges Transition Metals  Consult ion list for charges  Stock System** vs. Classic Stock system = Name of element(charge as Roman numerals)  Iron(III) = Fe 3+ Classic = Latin or traditional name of element  Ferric ion = Fe 3+

10. Examples  CuCl  HgO  MnO 2  PbCl 4  Copper(I) Chloride  Mercury(II) Oxide  Manganese(IV) Oxide  Lead(IV) Chloride

11. Naming Molecular (Non-metal) Compounds (Type III)  If all elements in the compound are nonmetals, use the following rules: 1.The first element in the formula is named first. (Full element name used) 2.2nd element named as if it were a single atom anion 3.Prefixes denote the number of atoms present. 4.The prefix mono- is never used in naming the first element.

12. Examples  BF 3  NO N2O5N2O5  boron trifluoride  nitrogen monoxide  dinitrogen pentoxide

13. Polyatomic Ions  Polyatomic = more than one atom  Use same naming techniques as in ionic compounds

14. Examples  Na 2 SO 4  Fe(NO 3 ) 3  NH 4 ClO 3  Mn(OH) 2  sodium sulfate  iron(III) nitrate  ammonium chlorate  manganese(II) hydroxide

15. Naming Acids  If the chemical formula starts with an H, it is an acid  Look at what is after the H to name it  If the anion ends in -ide Hydro________ic acid  HCl---chloride anion Hydrochloric acid  If the anion ends in -ate _________ic acid  HNO 3 -------nitrate anion Nitric acid  If the anion ends in -ite _________ous acid  HNO 2 -------nitrite anion Nitrous acid “-ate” something “-ic”y In sp“-ite” of “-ous”

16. Writing Formulas from Names  Ionic Compounds 1.Write the symbol for each ion 2.Balance charges with subscripts  If you get stuck, use the crisscross method  Example, write the formula for Iron(III) Oxide

17. Crisscross Method Fe 3+ O 2- Fe 2 O 3 Two Fe(III) have a charge of 6+Three O 2- have a charge of 6-

18.  Make sure that the formula has subscripts with the least whole number ratio  You never put in a subscript of 1.  Writing formulas for molecular compounds 1.Write the element symbol 2.Prefix becomes subscript

19. Write the following chemical formulas  potassium hydroxide  sodium carbonate  nitric acid  cobalt(III) nitrate  tetraphosphorus hexoxide  sulfur dioxide  hydrosulfuric acid  KOH  Na 2 CO 3  HNO 3  Co(NO 3 ) 3 P4O6P4O6  SO 2 H2SH2S