1. THE MOLE 10/1/14 – CHAPTER 3

2. RELATIVE MASS VS AVERAGE ATOMIC MASS RM = the standard - atomic mass of carbon-12 is used - 1 amu = 1/12 mass of 1 carbon-12 atom - refer to masses of sub-atomic particles proton & neutron

3. Isotopes differ in their atomic mass, but not their chemical behaviors. Average atomic mass = most elements occur as a mixture of isotopes - weighted average of atomic masses of all the naturally occurring isotopes of an element. - depends of the mass and relative abundance of isotope – see sample calc p. 82

4. WHY THE MOLE (MOL) http://chemistry.about.com/cs/generalchemistry/f/blmole.htm Is a standardized grouping. Technically, the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12 It’s just a way to represent a group of something / a standard. Here’s an example – a dozen

5. AVOGADRO’S # That 12 g carbon-12 sample has been measured out to contain 6.0221415 x 10 23 atoms This number represents the number of particles in amount of a substance that we call a mole. Relationship between mass and a number of atoms Use 6.022 x 10 23 in our calculations

6. MOLAR MASS This is the approximate mass of 1 mole of a substance. Is the ratio of grams / mol Need to use the periodic to determine the atomic mass of an element of interest

7. Ex: Molar mass of Li is 6.94 g/mol It is the custom to round the decimal to 2 places. You can determine the mass of 1 mol of a substance Ex: 1 mol of Li is 6.94 g

8. MOLAR MASS = CONVERSION FACTOR TOOL Used to convert between mass of an element, the amount of an element, and number of atoms in an element Refer to Fig 11 p. 84 Sample Problems A – E provide good practice for these conversions and the Answers are in Appendix E of your textbook.