1. Stoichiometry and the Mole (Part 1) Formula Mass and Molar Mass

2. For Atomic Structure, Tables Such as this were Used to Specify It ElementSymbolAt. #ProtonsNeutronElectrons Mass Number Oxygen 16 Ba81 26 56 17 35 1734 Nitrogen 87

3. The units were amu’s, which stand for atomic mass units. ElementSymbolAt. #ProtonsNeutronElectrons Mass Number Oxygen 16 Ba81 26 56 17 35 1734 Nitrogen 87

4.  Using atomic masses in amu’s and chemical formulas, one may calculate the formula mass of both ionic and covalent compounds.  The general procedure is: (A) Count (B) Multiply (C) Add For Ionic and Covalent Compounds, Formulas are Used to Specify Chemical Compositions

5.  Use the Periodic Table to find Atomic Mass.  Round to the nearest tenth of a unit.  Count:1 P @ 31.0 amu 3 H @ 1.0 amu  Multiply:(1 x 31.0) = 31.0 amu (3 x 1.0)= 3.0 amu  Add:31.0 + 3.0 = 34.0 amu Calculate the Formula Mass of PH 3

6.  Use the Periodic Table to find Atomic Mass.  Round to the nearest tenth of a unit.  Count:2 Al @ 27.0 amu 3 O @ 16.0 amu  Multiply:(2 x 27.0) = 54.0 amu (3 x 16.0)= 48.0 amu  Add:54.0 + 48.0 = 102.0 amu Calculate the Formula Mass of Al 2 O 3

7.  One amu = 1.66 x 10 -24 grams  One wants to measure gram quantities in the laboratory.  To scale the amu to grams, one needs the concept of the mole.  By definition, one mole contains 6.02 x 10 23 particles.  602 000 000 000 000 000 000 000 amu’s are too small to be useful laboratory units.

8.  By multiplying the atomic masses or formula masses (in amu’s) by Avogadro’s number, one gets quantities that one can measure in grams in the laboratory.  The concept is defined with respect to pure carbon-12 (the isotope of carbon with 6 protons and 6 neutrons).  Thus, one mole of carbon-12 has a mass of 12.0000000000000000000000 grams. 6.02 x 10 23 is called Avogadro’s Number

9.  For elements, the particles are atoms. Examples: Mg, Fe, and B  For covalent compounds (nonmetal with nonmetal), the particles are molecules. Examples: CO 2, SO 3, Br 2  For ionic compounds (usually, but not only, metal with nonmetal), the particles are called formula units. Examples: NaCl, Mg 3 (PO 4 ) 2, (NH 4 ) 2 SO 4 What does one mean by particles?

10.  1 mole contains 6.02 x 10 23 particles.  1 mole of carbon contains 6.02 x 10 23 atoms.  1 mole of lithium contains 6.02 x 10 23 atoms.  1 mole of H 2 contains 6.02 x 10 23 molecules.  1 mole of CO 2 contains 6.02 x 10 23 molecules.  1 mole of NaCl contains 6.02 x 10 23 formula units.  1 mole of (NH 4 ) 2 CO 3 contains 6.02 x 10 23 formula units. Examples of what one mole means

11.  Using atomic masses in grams/mole, chemical formulas, and the mole concept, one may calculate the molar mass of both ionic and covalent compounds.  The general procedure is as before: (A) Count (B) Multiply (C) Add  The only difference between the formula mass and the molar mass is the units. Calculating Molar Masses

12.  Use the Periodic Table to find Atomic Mass.  Round to the nearest tenth of a unit.  Count:1 P @ 31.0 g/mol 3 H @ 1.0 g/mol  Multiply:(1 x 31.0) = 31.0 g/mol (3 x 1.0)= 3.0 g/mol  Add:31.0 + 3.0 = 34.0 g/mol Calculate the Molar Mass of PH 3