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Chemistry 11 PLO Examples Sheet 4
Unit V. Mole Concept
1.   Explain why units of atomic mass are not practical in laboratory measurement:

2.   Define the mole:

3.   Determine the molar mass of


     oxygen gas

     ammonium phosphate

4.   Calculate the following:

     the number of moles of methane molecules in 100.0g of methane

     the number of magnesium hydroxide molecules in 4.50g

     the volume of 2.5 mol of pure water
MOLE CONCEPT (Molar Volume of Gases)

1.   State Avogadro’s hypothesis:

2.   A Hoffman apparatus is used to separate water into hydrogen and oxygen. Using this apparatus, 1.50g
     of water is found to yield 405mL of hydrogen gas at room temperature and pressure. Calculate the
     volume of one mole of hydrogen gas for these conditions of temperature and pressure.

3.   State the molar volume of a gas at STP

4.   calculate the following:

     the number of moles of chlorine gas in 10.0L at STP

     the mass of 5.0L of propane gas at STP
MOLE CONCEPT (Percent Composition)

1.   How is an empirical formula different from a molecular formula?

     What do the molecular and empirical formula of a compound have in common?

2.   Determine the percent composition by mass of the following:



3.   Determine the empirical formula for a compound with percent composition as follows: 46.2%C,
     7.69%H, 46.2%O.

4.   Determine the molecular formula of a compound whose empirical formula is CH2F and whose
     molecular mass is 69.0g/mol:

1.   Describe what is meant by the term molarity:

2.   Describe the procedure used to prepare a standard solution of 1.0M sodium nitrate:

3.   Calculate the following:

     The number of grams of CuCl2 needed to prepare 250mL of a 2.5M solution

     The molar concentration of a solution, 1.50L of which contains 4.60g of NaOH

     The volume needed to make up a solution of 1.0M NH4Cl, using 50.0g of solid NH4Cl

4.   Calculate the resulting concentration when 75mL of 0.45M HCl is diluted with water to make up 1.0L