The Mole Notesheets
Name_______________________________
Period______Date_____________________
7.1, 7.3
7.1 Chemical Measurement:
1. Counting units:
pair = _______
dozen = _______
score = _______
gross = _______
ream = _______
mole = _________________________
2. A mole is the ________________ of a substance. It is based on the number of atoms of an
element equal to the number of atoms in exactly 12.0g of _____________________.
3. The number of particles in a mole is called ________________________ ____________________.
4. The value of Avogadro’s constant is ___________________________________________________.
5. Atomic mass: mass of one atom of an element measured in amu (atomic mass units).
Ex #1) H = 1 amu
O = _______amu
C = _______amu
6. Formula mass: mass of all the atoms in a single molecule or formula unit of a compound.
Ex #2) H2O =
Ex #3) H2CO3 =
7. Molar Mass: mass of one mole of an element The units used are _________. Round all elements’
masses to the nearest whole number except _______ + _______.
Ex #4) Elements:
Cu = _______g/mol
1 mole Cu = ___________________atoms
Cl = _______g/mol
1 mole Cl = ____________________atoms
Ex #5) Covalent Compounds:
H2O = _______g/mol
1 mole H2O = ______________________molecules
H2CO3 = _______g/mol
Ex #6) Ionic Compounds:
NaCl = _______g/mol
MgO = _______g/mol
8. Examples: N2 = _________g/mol
2N2 = _________g/mol
=
=
1 mole H2CO3 = ____________________molecules
1 mole NaCl = __________________formula units
1 mole MgO = __________________formula units
_________________________molecules
_________________________molecules
9. Practice: Calculate the molar mass of sucrose, C12H22O11.
Avogadro’s Number
1
10.
1 mole
= _________________________________ particles
=
molar mass
1 mole Ne
= ________________________________________
=
__________g
1 mole CO2
= ________________________________________
=
__________g
1 mole CaF2
= ________________________________________
=
__________g
11. The mole is the _________________ between calculations of # of particles, mass and volume of a gas.
12. Conversion Factors:
1 mole = 6.02 x 1023 particles (atoms, molecules or formula units)
1 mole = molar mass (grams) from the periodic table
May be used in a problem in one of two ways depending on the given units.
 6.02x1023   1 mole 

 or 
23 
 1 mole   6.02x10 
 X grams   1 mole 


 or 
 1 mole   X grams 
Ex #1) Given 11.2 g of NaCl, how many moles does this represent?
 1 mole NaCl 
11.2 g NaCl 
  .191 mol NaCl *3 significant digits because 11.2 is 3
 58.5 g NaCl 


Ex #2) Given 2.50 moles of NaCl, how many grams does this represent?
Ex #3) How many particles are in 2.00 moles of H2O? How many atoms is this?
Ex #4) How many moles of CaCO3 are 8.74 x 1023 formula units?
Multistep Problems:
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Ex #1) How many particles of copper are in 56.3 g of copper?
 1mole Cu
56.3 g Cu 
 63.5 g Cu

  6.02 x 1023 atoms Cu 
23

 = 5.34 x 10 atoms Cu

1mole Cu


We use atoms because copper is an element. There are 3 significant digits because 56.3 has 3.
Ex #2) Given 7.2 x 1023 atoms of calcium, how many grams of calcium is this?
Ex #3) How many sugar particles of sugar (C12H22O11) are in 250g of sugar?
Moles & Gases:
Molar Volume: 1 mole of any gas at Standard Temperature and Pressure (STP = 0oC and 1 atm) has a
volume of 22.4dm3 or __________L
Ex #1) How many particles of CO2 gas are in a 1.0L flask at STP?
 1 mole CO2
1.0 L CO2 
 22.4 L CO2
  6.02 x 1023 molecules CO2 
22

 = 2.7 x 10 molecules CO2
1 mole CO2


Ex #2) How many atoms of radon gas are contained in a 6500. dm3 basement room at STP?
Ex #3) If a room has a volume of 4000. L, how many moles of air is this at STP?
7.3 % Composition, Empirical & Molecular Formulas:
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% Composition: The mass of each element in a compound compared to the entire __________ of the
compound x by 100%. This can be calculated experimentally from the grams of each
element in a compound, or the expected % can be calculated by using the molar masses
of the elements in the compound compared to the molar mass of the compound.
Ex #1) What is the % hydrogen in water? Use the molar masses when not provided the lab data.
%H=
2.0 g H
x 100% = 11%
18 g H20
%O=
The numbers should add up to 100% unless significant digits prevent it.
Ex #2) Find the % composition of a compound that contains 2.30 g of sodium, 1.60 g of oxygen, and
0.100 g of hydrogen.
Ex #3) A sample of an unknown compound with a mass of 0.562 g has the following % composition:
13.0 % carbon, 2.20 % hydrogen, and 84.5 % fluorine. When this compound is decomposed into
its elements, what mass of each element should be recovered?
Empirical Formula: A formula that gives the simplest _____________-number ratio of the atoms of the
elements in the compound.
Molecular Formula: Gives the _____________ number of atoms of each element in a molecular compound.
The molecular formula may be the same as the empirical formula. H20
Ex #1) Hydrogen peroxide
= _______________ molecular formula
= _______________ empirical formula (a ____: ____ ratio)
Ex #2) glucose
=
C6H12O6
molecular formula
= _______________ empirical formula (a___: ___:___ ratio)
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Ex #3) Find the empirical formula of the compound if you have 80.g of carbon and 20.g of hydrogen.
Ex #4) An unknown compound is analyzed. It’s composition was determined to be 51.85% carbon,
8.64% hydrogen and 39.51% oxygen. Find the empirical formula of the compound.
Ex #5) Find the molecular formula of ribose (molecular mass = 150 g/mol). It has a chemical
composition that is 40.0% carbon, 6.67% hydrogen and 53.3% oxygen. Assume a 100 g sample.
Hint: First find the empirical formula:
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