1. CHE 111 - Module 3 CHAPTER 3 LECTURE NOTES

2. STOICHIOMETRY • Stoichiometry is the study of the quantitative relationships between the amounts of reactants and products in chemical reactions. • We use BALANCED equations to understand stoichiometric relationships of the elements and compounds within a chemical reaction.

3. The Balanced Equation 2Al(s) + 3Br2(l) Al2Br6(s) 2mol of Al : 3mol of Br2 : 1mol of Al2Br6 Therefore the ratio of Al to Br2 to Al2Br6 is 2:3:1 for the chemical reaction to occur.

4. A Closer Look at the Equation 2Al(s) + 3Br2(l) Al2Br6(s) • The chemicals on the left are the reactants and the right are the products. • The coefficient in front of the chemical denotes the stoichiometric relationship.

5. Numerical Subscripts 2Al(s) + 3Br2(l) Al2Br6(s) • The numerical subscriptrepresents the number of atoms present in the molecule • ex. Br2 means that an atom of Br is bonded to another atom of Br • Therefore: Br-Br = Br2

6. Denoting the Phase of Matter 2Al(s) + 3Br2(l) Al2Br6(s) The subscript letters in parenthesis denote the phase of matter that the chemical is in.

7. Formulas and Models of Ethanol • Molecular Formulas C2H6O • Condensed Formulas C2H5OH H H • Structural Formulas H-C-C-O-H H H • Molecular Models (classroom models)

8. Molecular Models • Cache program - models - organic models - ethanol • CD-ROM screen 3.4 • Model of ice

9. Ions and Ionic Compounds • Ions are atoms or groups of atoms that have lost or gained electrons resulting in an overall positive or negative charges. • Ionic compounds are compounds formed by the combination of (+) and (-) ions. (+) ions are called cations (-) ions are called anions

10. Formation of Ions Formation of a cation by a loss of electrons Li atom  Li+ + 1 e- released (3p and 3e-)  (3p and 2e-) Formation of an anion by gaining electrons F atom + 1 e- added  F (9p and 9e-)  (9p and 10e-)

11. Ions and the Periodic Table Metals of group 1A, 2A & 3A form +1, +2, and +3 ions; and non-metals of group 5A, 6A, and 7A form -3, -2, and -1 respectively.

12. Polyatomic Ions • Table 3.1 - page 89 • CD-ROM Screen 3.6 • Hand out • Flash Cards

13. carbonate ion CO3-2 sulfate ion SO4-2 sulfite ion SO3-2 hydroxide OH- phosphate PO4-3 permanganate MnO4- chromate CrO4-2 dichromate Cr2O7-2 ammonium NH4+ oxalate C2O4-2 bicarbonate HCO3- cyanide ion CN- acetate C2H3O2- peroxide O2-2 thiosulfate S2O3-2 bisulfite HSO3- 0 Common Polyatomic Ions

14. 0 Oxoanions A polyatomic anion containing oxygen is called an oxoanion and is named as follows: • Greater # of O atoms has the suffix -ate. • Lesser # of O atoms has the suffix -ite. Ex. NO3- is called nitrate ion NO2- is called nitrite ion

15. 0 Naming Oxoanions More than 2 ions in an oxoanion grouping are named as follows: • Largest # of O atoms has a prefix of per- and a suffix of -ate • Next larger # of O atoms has a suffix -ate • Smaller # of O atoms has a suffix -ite • Smallest # of O atoms has a prefix of hypo- and a suffix of -ite

16. 0 Naming Oxoanions Ex. ClO4- is called perchlorate ClO3- is called chlorate ClO2- is called chlorite ClO- is called hypochlorite

17. 0 Ionic Compounds Ca+2 + 2Cl- CaCl2 Each ion comes together based on charge to form an overall neutral ionic compound. 3Ca+2 + 2PO4-3 Ca3(PO4)2 The cation and the polyatomic ion come together based on charge to form an overall neutral ionic compound.

18. 0 Naming Ionic Compounds • Naming Positive Ions – Cations • Cations are named first in the compound and as follows: • Monatomic cations are mostly metals and are named directly as they are on the periodic table. • Transition metals are named according to their ionic charge • Polyatomic cation, NH4+ is named ammonium directly

19. 0 Naming Ionic Compounds • Naming Negative Ions – Anions • Anions are named lastly and have specific naming rules as follows: • Monatomic ions are named with an –ide after its atomic name • Polyatomic ions are named as memorized dropping the word ion.

20. 0

21. 0

22. 0 Naming Molecular Compounds • 1 mono • 2 di • 3 tri • 4 tetra • 5 penta • 6 hexa • 7 hepta • 8 octa

23. Formula & Molecular Weights • Review of spectra lab - MW calculations • CD-ROM Screen 3.14 • Definition: The total mass of the formula unit or molecule with consideration to the mass of each component element that makes up the overall unit.

24. Calculating Formula & MW • Remember we said that: 1 mole C = 12.011g C = 6.02x1023 atoms C • If we add up the number atoms present of each element in a molecule or formula unit and multiply each by its atomic weight on the periodic table, • Then the resultant sum of each element added together will give you the formula or molecular weight.

25. Example of MW Calculation • Determine the MW of H20 • 1 O @ 15.999g/mole • 2 H @ 1.008g/mole • Therefore 2 x 1.008 = 2.016g/mole • and 1 x 15.999 = 15.999g/mole Total molar mass = 18.015g/mole Determine the MW of ethanol: C2H5OH

26. Converting Mass to Moles • Question: How many moles of H2O are in 42.0g of water? • Answer: First you determine the MW of water as we did on the previous slide, then you convert 42.0g H2O x 1 mole H2O = 18.016g H2O

27. Percent Composition • Calculate the percent composition of NH3 • First determine the atomic weights of each N and H from the periodic table • Then calculate the MW of the ammonia molecule • Take the mass of each element and divide by the MW and multiply 100% • CD-ROM Screens 3.14 and 3.16

28. Hydrated Compounds • Definition: Ionic compound that has water molecules incorporated within its crystal structure • Ex. CuCl2•2H2O • Where we name this compound copper(II) chloride dihydrate • When calculating MW, we calculate the two waters into the overall mass of the compound