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) + 3Br 2(l)  Al 2 Br 6(s) 2 mol of Al : 3 mol of Br 2 : 1 mol of Al 2 Br 6 Therefore the ratio of Al to Br 2 to Al 2 Br 6 is 2:3:1 for the chemical reaction to occur.

4. A Closer Look at the Equation 2Al (s) + 3Br 2(l)  Al 2 Br 6(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) + 3Br 2(l)  Al 2 Br 6(s) The numerical subscript represents the number of atoms present in the molecule –ex. Br 2 means that an atom of Br is bonded to another atom of Br –Therefore: Br-Br = Br 2

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

7. Formulas and Models of Ethanol Molecular Formulas C 2 H 6 O Condensed Formulas C 2 H 5 OH H H Structural FormulasH-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. Common Polyatomic Ions carbonate ion CO 3 -2 sulfate ion SO 4 -2 sulfite ion SO 3 -2 hydroxide OH - phosphate PO 4 -3 permanganate MnO 4 - chromate CrO 4 -2 dichromate Cr 2 O 7 -2 ammonium NH 4 + oxalate C 2 O 4 -2 bicarbonate HCO 3 - cyanide ion CN - acetate C 2 H 3 O 2 - peroxide O 2 -2 thiosulfate S 2 O 3 -2 bisulfite HSO 3 -

14. 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. NO 3 - is called nitrate ion NO 2 - is called nitrite ion

15. 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. Naming Oxoanions Ex. ClO 4 - is called perchlorate ClO 3 - is called chlorate ClO 2 - is called chlorite ClO - is called hypochlorite

17. Ionic Compounds Ca +2 + 2Cl -  CaCl 2 Each ion comes together based on charge to form an overall neutral ionic compound. 3Ca +2 + 2PO 4 -3  Ca 3 (PO 4 ) 2 The cation and the polyatomic ion come together based on charge to form an overall neutral ionic compound.

18. 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, NH 4 + is named ammonium directly

19. 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.

21. Cr +2 Cr +3 Cr +6 Chromium II,III&VI Mn +2 Mn +3 Mn +4 Mn +6 Mn +7 Fe +2 Fe +3 Fe +6 Iron II, III & VI Co +2 Co +3 Cobalt II & III Ni +2 Ni +3 Nickel II & III Cu +1 Cu +2 Copper I & II Pb +2 Pb +4 Lead II & IV Sn +2 Sn +4 Tin II & IV Au +1 Au +3 Gold I & III Hg +1 Hg +2 Mercury I & II Zn +2 Zinc Ag +1 Silver

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

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 = 6.02x10 23 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 H 2 0 –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: C 2 H 5 OH

26. Converting Mass to Moles Question: How many moles of H 2 O 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 H 2 O x 1 mole H 2 O = 18.016g H 2 O

27. Percent Composition Calculate the percent composition of NH 3 –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. CuCl 2 2H 2 O –Where we name this compound copper(II) chloride dihydrate –When calculating MW, we calculate the two waters into the overall mass of the compound