1. Prentice Hall © 2003Chapter 3 Combination reactions have fewer products than reactants: 2Mg(s) + O 2 (g)  2MgO(s) Mg has combined with O 2 to form MgO Categories: 1.Metal + O 2  Metal oxide (Mg + O 2 ) 2.Nonmetal + Nonmetal  covalent compound (ammonia formation) 3.Metal + Nonmetal (other than O 2 )  salt (Mg + Cl 2 ) 4.Metal oxide + H 2 O  metal hydroxide (NaOH) 5.Nonmetal oxide + H 2 O  oxyacid (CO 2 + H 2 O) 6.Metal oxide + Nonmetal oxide  salt (Na 2 O + SO 2 ) 3.2: Some Simple Patterns of Chemical Reactivity no change in charge on central atom

2. Prentice Hall © 2003Chapter 3 Decomposition Reactions have fewer reactants than products: 2NaN 3 (s)  2Na(s) + 3N 2 (g) (the reaction that occurs in an air bag) NaN 3 has decomposed into Na and N 2 gas Categories: 1.Binary Compound  elements (NaCl  ) (don’t forget diatomics!) 2.Metal Carbonate  metal oxide + CO 2 (Na 2 CO 3  ) 3.Metal hydroxide  metal oxide + H 2 O (NaOH  ) 4.Metal chlorate  chloride salt + O 2 (NaClO 3  ) 5.Oxyacid  nonmetal oxide + H 2 O (H 2 CO 3  ) These are the opposite of the combination reactions

3. Prentice Hall © 2003Chapter 3 Combination and Decomposition Reactions Text, P. 80

4. Prentice Hall © 2003Chapter 3 Combustion in Air Combustion is the burning of a substance in oxygen from air: C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(l) Hydrocarbon + O 2  CO 2 + H 2 O (complete combustion)

5. Prentice Hall © 2003Chapter 3 Examples # 13 & 11 (WS) and classification/prediction practice (WS) NaCl

6. Prentice Hall © 2003Chapter 3 Formula and Molecular Weights Formula weights (FW): sum of AW for atoms in a formula FW (H 2 SO 4 ) = 2AW(H) + AW(S) + 4AW(O) = 2(1.01 amu) + (32.06 amu) + 4(16.00 amu) = 98.08 amu Molecular weight (MW) is the weight of the molecular formula MW(C 6 H 12 O 6 ) = 6(12.01 amu) + 12(1.01 amu) + 6(16.00 amu) = 180.18 amu 3.3: Formula Weights

7. Prentice Hall © 2003Chapter 3 Percentage Composition from Formulas Percent composition is the atomic weight for each element divided by the formula weight of the compound multiplied by 100: Don’t forget hydrates!

8. Prentice Hall © 2003Chapter 3 Examples # 15, 17, 19

9. Prentice Hall © 2003Chapter 3 Mole: convenient measure chemical quantities. 1 mole of something = 6.0221367  10 23 of that thing Experimentally, 1 mole of 12 C has a mass of 12 g Molar Mass Molar mass: mass in grams of 1 mole of substance (units of g/mol, g*mol -1 ) Mass of 1 mole of 12 C = 12 g 3.4: The Mole

10. Text, P. 88

11. Prentice Hall © 2003Chapter 3 Examples # 25 & 26

12. Prentice Hall © 2003Chapter 3 Text, P. 88

13. Prentice Hall © 2003Chapter 3 This photograph shows one mole of solid (NaCl), liquid (H 2 O), and gas (N 2 ).

14. Prentice Hall © 2003Chapter 3 Interconverting Masses, Moles, and Number of Particles Molar mass: sum of the molar masses of the atoms: molar mass of N 2 = 2  (molar mass of N) Periodic table Formula weights are numerically equal to the molar mass Text, P. 90

15. Prentice Hall © 2003Chapter 3 Relating particles, Mass and Moles Molar mass (g) 6.02x10 23 particles 1 mole Molecules (molecular compounds) Formula Units (ionic compounds) Atoms (elements) ions are made of atoms are made of

16. Prentice Hall © 2003Chapter 3 Macroscale vs. Microscale Interpretations of the Mole 1 formula unit of NaCl 1 mole of Cl - ions 1 mole of Na + ions1 ion of Na + 1 ion of Cl - Microscale 1 mole of NaCl Macroscale

17. Prentice Hall © 2003Chapter 3 Examples # 27, 29, 33 & 35