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Measuring Matter
How would we describe how much of a substance we have?
 _________________________________
 _________________________________
 _________________________________
Mole (mol): SI unit for measuring the ________________________________________________________.
___mol = ________________ representative particles
Avogadro’s Number: _________________________________
Representative particles: _____________ unit that still has all the ________________ of that substance.
Practice - What is the representative particle of a/an:
1. Element (ex. Cu):
___________________
o Exception: The representative particle of the 7 diatomic elements is a
__________________. (ex. H2)
2. Covalent compound (ex. H2O): ___________________
3. Ionic Compound (ex. NaCl):
___________________
Practice - Convert the following:
4 moles Ca = ________________ atoms Ca?
5 x 1018 atoms Cu = ________________ moles Cu?
9.2 moles F2 = ________________ molecules F2?
9.2 moles F2 = ________________ atoms F?
3.4 moles C2H4 = ________________ atoms total?
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Molar Mass: The __________ of ____________ mole of an element or compound.
o Molar mass of a compound = the sum of the masses of the atoms in the formula (use the
atomic masses in grams on the periodic table).
___ mol = molar mass (in ________)
Practice - Find the molar mass:
1) Sr
2) MgBr2
3) Ba3(PO4)2
Practice - Convert the following:
5.3 moles LiOH = ___________ grams LiOH
68 grams F2 = ___________ moles F2
Standard Temperature and Pressure (STP): ________________ and ________________
Molar Volume of a Gas
Avogadro’s Hypothesis: equal volumes of gases at the _________ temperature and pressure contain
equal numbers of particles.
o At STP, 1 mole of any gas occupies a volume of _____________.
___ mol = ______ L at STP (gases only!!!)
Practice - Convert the following:
5.4 mol He = ___________ L He at STP
5.4 mol CH4 = ___________ L CH4 at STP
560 L SO3 = ___________ mol SO3
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Molar Mass  Density
Say you are given the density of a gas (grams/liters) at STP and you have to find the molar mass
(grams/mole)…
grams
liters

grams
mole
Practice:
1. A gaseous compound composed of sulfur and oxygen has a density of 3.58 g/L at STP.
What is the molar mass of this gas?
2. What is the density of krypton gas at STP?
Law of Definite Proportions: In samples of any chemical compound, the ____________ of the
elements are always in the ________ proportion. (Allows us to write chemical formulas.)
Finding Chemical Formulas (from the % composition)
Example Problem
Percent
# grams in
How many moles
Composition
100 grams
of each element?
18.8% Na
29.0% Cl
52.2% O
Empirical Formula =
Divide by smallest
#moles
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Empirical Formula: lowest whole-number ratio.
 The formula for an ionic compound will always be the empirical formula.
 The formula for a covalent compound will not always be the empirical formula.
Molecular Formula: either the same as the empirical formula (as for ionic compounds) or a
simple whole-number multiple of the empirical formula.
Practice Problems: Determining Chemical Formulas
1. Suppose the mass percents of a compound are 40% carbon, 6.70% hydrogen, and 53.3%
oxygen. Determine the empirical formula for this compound.
Since this compound is covalent, the actual formula may not be the simplest ratio of elements.
If the molar mass of the compound is experimentally shown to be 90.0 g/mol, what is the
molecular formula of this covalent compound?
2. Find the molecular formula of ethylene glycol (CH2O) if its molar mass is 180 g/mol.
3. The percent composition of methyl butanoate is 58.8% C, 9.8% H, and 31.4 % O and its molar
mass is 102 g/mol.
a. What is its empirical formula?
b. What is its molecular formula?
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Stoichiometry: The calculation of ____________ of substances involved in ____________________.
Mole-Mole Conversions
2A + B  3C + 7D
Examples:
Mole-Mole Conversions
N2 (g) + 3H2 (g)  2NH3 (g)
1. If you have 2 moles of N2, how many moles of NH3 will be produced?
2. If you want 5 moles of product, how many moles of hydrogen gas do you need?
3. How many moles of nitrogen are needed to react completely with 8 moles of hydrogen?
Volume-Volume Conversions
1. How many liters of oxygen are required to burn 3.86 L of carbon monoxide?
2CO (g) + O2 (g)  2CO2 (g)
2. How many liters of PH3 are formed when 0.42 L of hydrogen reacts with phosphorus?
P4 (s) + 6 H2 (g)  4 PH3 (g)
Mass-Mass Conversions
N2 (g) + 3H2 (g)  2NH3 (g)
1. How many grams of hydrogen gas are required for 3.75 g of nitrogen gas to react completely?
2. What mass of ammonia is formed when 3.75 g of nitrogen gas react with hydrogen gas?
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Practice Problems
1. The combustion of propane, C3H8, a fuel used in backyard grills and camp stoves, produces
carbon dioxide and water vapor.
C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (g)
What mass of carbon dioxide forms when 95.6 g of propane burns?
2. Solid xenon hexafluoride is prepared by allowing xenon gas and fluorine gas to react.
Xe (g) + 3F2 (g)  XeF6 (s)
a. How many grams of fluorine are required to produce 10.0 g of XeF6?
b. How many grams of xenon are required to produce 10.0 g of XeF6?
3. How moles of CO2 are produced when 52.0 g C2H2 burns?
2C2H2 (g) + 5O2 (g)  4CO2 (g) + 2H2O (g)
4. How many liters of hydrogen gas are formed from 50 grams of potassium?
2K (s) + 2H2O (l)  2KOH (aq) + H2 (g)
5. How many molecules of oxygen are produced by the decomposition of 6.54 g of potassium
chlorate (KClO3)?
2KClO3 (s)  2KCl (s) + 3O2 (g)
6. How many grams of nitrogen dioxide must react with water to produce 5.00 x 1022 molecules
of nitrogen monoxide?
3NO2 (g) + H2O (l)  2HNO3 (aq) + NO (g)
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Limiting Reagent
Limiting Reagent: The reactant that limits the amount of product that can be formed in a reaction. The
reaction will stop when all of this reactant is used up.
Excess Reagent: You have more than you need of this reactant. The reaction will stop before all of
this reactant is used up. You will have some of this reactant leftover.
** The amount of product formed in a reaction can be determined from the given amount of
limiting reagent.
Real Life Example
You have 1 loaf of bread (containing 14 slices of bread),
4 jars of peanut butter, and 2 jars of jelly.
a) How many peanut butter and jelly sandwiches can you make?
b) What is the limiting reagent?
c) What is in excess?
LIMITING REAGENT EXAMPLE PROBLEMS
1. If 2.70 mol C2H4 is reacted with 6.30 mol O2, what is the limiting reagent?
C2H4 (g) + 3O2 (g)  2CO2 (g) + 2H2O (g)
2. Identify the limiting reagent when 6.00 g HCl reacts with 5.00 g Mg.
Mg (s) + 2HCl (aq)  MgCl2 (aq) + H2 (g)
How many grams of MgCl2
are produced in this reaction? ___________________
Which reactant is in excess? ___________________
3. How many grams of water can be produced by the reaction of 2.40 mol C2H2 with 7.4 mol O2?
Limiting Reagent? ______________
Excess Reagent? _______________
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Percent Yield
Theoretical Yield: the maximum amount of product that could be formed from the given amounts of
reactants (ideal conditions).
o Calculated using stoichiometry.
Actual Yield: the amount of product that actually forms in a lab.
o Actual yield is usually less than theoretical yield.
Percent Yield =
Actual Yield
x 100%
Theoretical Yield
PERCENT YIELD EXAMPLE PROBLEM
If 3.75 g of nitrogen completely react, what is the theoretical yield of NH3?
N2 (g) + 3H2 (g)  2NH3 (g)
If the actual yield is 3.86 g, what is the percent yield?
Practice Problem
Find the percent yield if 84.8 g of iron (III) oxide reacts with an excess of carbon monoxide to produce
55.0 g of iron.
Fe2O3 (s) + 3CO (g)  2Fe (s) + 3CO2 (g)