2.  The percent composition of fuel is 81.7% carbon and 18.3% hydrogen. What is the empirical formula of the fuel?  Start by assuming 100 g of the compound (81.7 g) / (12.01 g/mol) = 6.80 mol C (18.3 g) / (1.008 g/mol) = 18.2 mol H

3.  Recall that many compounds could potentially have the same empirical formula, but the molecular formula is specific to a given compound

4.  Molecular formula subscripts are a whole number multiple of the empirical formula subscripts  Similarly, the molar mass of a compound is a whole number multiple of the molar mass of the empirical formula  So, molar mass of the compound = n x molar mass of the empirical formula, where n = 1, 2, 3…

5.  A chemist determines that a substance has a empirical formula of CH and a molar mass of 78 g/mol. What is the molecular formula of the substance?  The empirical formula of CH is 13.0 g/mol  The ratio of the empirical formula to the molecular formula is = 6  Therefore, we multiply the subscripts of the empirical formula by the ratio (6) to get the molecular formula: C 6 H 6 78 13

6.  The empirical formula for ribose (a sugar) is CH 2 O. A mass spectrometer determined the molar mass of ribose to be 150 g/mol. What is the molecular formula of ribose?

8.  Many ionic compounds crystallize from water solutions and end up with water incorporated into their structures, forming a hydrate (hydrated = contains water, anhydrous = without water)  For example, the chemical name for Epsom salts is magnesium sulfate heptahydrate (use prefixes to tell the number of water molecules in the compound)  MgSO 4 · 7H 2 O  The dot represents a weak bond  The molar mass of a hydrated compound must incorporate the mass of the water molecules in the compound

9.  A 50.0 g sample of a hydrate of barium hydroxide contains 27.2 g of Ba(OH) 2. What is the percent, by mass, of water in the compound?

10.  What is the formula for the hydrated barium oxide compound ?  1. Moles of Ba(OH) 2 ?  2. Moles of H 2 O?  3. Molar ratio?

11.  The mass of an atom is expressed in atomic mass units (u or amu)  Relative measure  One atom of carbon 12 has a mass of 12 u (ie. One u has a mass of 1/12 that of a C-12 atom)  Average atomic mass (what we see on the periodic table) takes into account all isotopes (and their relative abundances)  One mole of an element has a mass (in grams) numerically equivalent to the element’s average atomic mass (in amu)  Take home: grams are numerically equivalent to amu.