
The study of the quantitative relationships between
reactants and products in a reaction
It is used to answer questions like; If I have this much
reactant, how much product can I make?
 If I want this much product, how much reactant do I need?




These questions have real life application,
particularly in manufacturing.
It allows us to convert the mass of a substance to the
number of particles (atoms, ions or molecules) it
contains.
These numbers can be really large, so they are
counted in groups
 Much like when we count a lot of pennies we stack them in 10’s
and count by 10
Atoms are very tiny, so small
that the grouping we use to count
them must be very large
 MOLE; the group (unit of measure) used to count
atoms, molecules, formula units or ions of a
substance
 1 mole of a substance has a particular number of
particles in it!

Much like 1 dozen always means 12; whether it is
12 eggs
12 oranges
or
12 gold bars

The number of particles in a mole = 6.02 x 10 23
or 602,000,000,000,000,000,000,000 !
This is known as Avogadro’s Number
Using this, We can easily count the number of
particles in all kinds of things !
There are 6.02 x 10 23 Carbon atoms in a mole of carbon
There are 6.02 x 10 23 CO2 molecules in a mole of CO2
There are 6.02 x 10 23 sodium ions in a mole of sodium
There are 6.02 x 10 23 marbles in a mole of marbles
That’s a lot of marbles!
The Size of a mole of a substance changes, the bigger
the substance the more space a mole of the
substance takes up, but the number of particles in a
mole is always the same!

Chemicals do not come bundled in moles, like a dozen
eggs comes in a 1 dozen or 1 ½ dozen package so we
use the mole as a grouping unit.
The mass of 1 mole of a pure substance called it’s molar
mass


If I want to produce 500g of methanol using the
following equation, CO2 +3H2  CH3OH + H20 how
many grams of CO2 and H2 do I need?
These are the questions stoichiometry answers!
If I want to produce 500g of methanol using
the following equation;
CO2 +3H2  CH3OH + H20
How many grams of CO2 and H2 do I need?
This equation relates the molecules of reactants and
products, NOT THEIR MASSES!

1 molecule of CO2 and 3 molecules of H2 will make 1
molecule of CH3OH
We need to relate the masses to the number of molecules.
Remember; The average atomic
masses of the elements are found on the
Periodic Table!

We can use the atomic masses on the PT to
relate the mass of the compound to the
mass of a mole!
Molar mass: The mass (in grams)of one
mole of a molecule or a formula unit
Molecular mass: mass in atomic mass units of just
one molecule
Formula Mass: mass in atomic mass units of one
formula unit of an ionic compound
Steps
1.
Find the average Atomic Mass of the element
on the PT. (state it in grams instead of atomic
units)
a)
b)
2.
Example: molar mass of Fe = 55.847 g
Example: molar mass of Pt = 195.08 g
If the element is a molecule, count the number of
atoms in the molecule then multiply the atomic
mass by the number of atoms.
a)
Example: O2, the mass of O =16.0g There are 2
atoms of O in the O2 molecule , 2 atoms X 16.0g =
32.00g is the molar mass of the molecule.
Calculate the molar mass of each of the
following:
1. N2
2. Cl2
3. Br2
4. I2
5. H2
6. F2
Calculate the molar mass of each of the
following:
1. N2 = 14.007g X 2 =28.014 g/mol
2. Cl2 = 35.453g X 2 =70.906 g/mol
3. Br2 = 79.904g X 2 =159.808 g/mol
4. I2 = 126.904g X 2 =253.808 g/mol
5. H2 = 1.008g X 2 =2.016 g/mol
6. F2 = 18.998g X 2 =37.996 g/mol
Steps
1. Count the number and type of atoms
2. Find the Atomic Mass of each atom type, on
the periodic table. Write it in grams.
3. Multiply the mass times the # of Atoms. Then
add the totals
1.
Count the number and type of atoms
Ethanol (C2H5OH)
2.
3.
Atom type
Amount of each atom
C
2
H
6
O
1
Find the Atomic Mass of each atom type, on the periodic table. Write
it in grams.
Atom type
Amount of atom
Ave. Atomic Mass in g
C
2
12.0
H
6
1.00
O
1
16.0
Multiply The mass X the # of Atoms. Then add the totals.
Atom type
Amount of atom
Ave. Atomic Mass in g
Total
C
2
12.0
=24.0
H
6
1.00
=6.0
O
1
16.0
=16.0
Molar Mass Of Ethanol (C2H5OH)
= 46.0g/mole
Example: Calcium Chloride (CaCl2 )
Atom
Types
Amount of
Atoms
Ave. Atomic
Mass in g
Total
Ca
Cl
1
2
40.1
35.5
40.1
71.0
Mass of 1 mol of CaCl2 (molar mass)
111.1 g/mole
What is the molar mass of each of the following?
1.
Fe2 O3
2.
H2O
3.
CO2
4. NaCl
5. NH3
6. BaI2
Fe2 O3 = 55.85g X 2= 111.7 g
16.0g X 3 = 48.0g
= 159.7 g/mol
_______________________________________________
H2O = 1.01g X 2 = 2.02
16.0g X 1 = 16.0
= 18.02 g/mol
_______________________________________________
CO2 = 12.01g X 1 = 12.01
16.0g X 2 = 32.0
= 44.01 g/mol
________________________________________________
NaCl = 22.99 gX1 = 22.99
35.45g X1 = 35.45
= 58.44 g/mol
________________________________________________
NH3 =14.01g X 1 = 14.01
1.01g X 3 = 3.03
= 17.04 g/mol
________________________________________________
BaI2 = 137.33g X 1 = 137.33
126.90g X 2 = 253.80
= 391.13 g/mol
If I want to produce 500g of ethanol using
the following equation;
6CO2 +17H2  3C2H5OH + 9H20
How many grams of CO2 and H2 do I need?
The Molar Mass Of Ethanol (C2H5OH)
= 46.0g/mole

Now we need to find the number of atoms in the
sample.
How many molecules of ethanol are in 500g?
Steps to finding the number of atoms in a given
mass of a sample
1. Use PT to find the molar mass of the substance
2. Convert the mass of the substance to number of
moles in the sample (convert using mass of one
mole as conversion factor)
3. Use the number of atoms in a mole to find the
number of atoms in the sample
4. Solve and check answer by canceling out units
The mass of an iron bar is 16.8g. How many iron(Fe)
atoms are in the sample?
Step 1: Use PT to find the molar mass of the
substance : The molar mass of Fe =55.8g/mole
Step 2: Convert the given mass of the substance to
number of moles in the sample: Fe =55.8g/mole
(16.8g Fe) (1 mol Fe) (6.022 X 1023 Fe atoms)
(55.8g Fe)
(1 mol Fe)
=
1.81 X 10
23 Fe atoms
Step 3: Use the number of atoms in a mole to find
the number of atoms in the sample = 1.18 X 1023
1.
25.0 g silicon, Si
2.
1.29 g chromium, Cr
(
25.0 g Si
(
1.29 g Cr
1 mol Si
) ( 28.1g Si ) (
6.02 X 1023 Si atoms
1
= 5.36 X1023 atoms Si
1 mol Cr
) ( 52.0g Cr ) (
1
= 1.49 X1022 atoms Cr
1 mol Si
)
6.02 X 1023 Cr atoms
1 mol Cr
)
1.
2.
3.
4.
98.3g mercury, Hg
45.6g gold, Au
10.7g lithium, Li
144.6g tungsten, W
1 mol Hg
1. 98.3 g Hg
1
200.6g Hg
= 2.95 X1023 atoms Hg
(
)(
)(
6.02 X 1023 Hg atoms
1 mol Hg
23 Au atoms
45.6
g
Au
1
mol
Au
6.02
X
10
2.
1
197.0g Au
1 mol Au
= 1.39 X1023 atoms Au
(
)(
)(
1 mol Li
3. 10.7 g Li
1
6.94g Li
= 9.28 X1023 atoms Li
(
)(
)(
1 mol W
4. 144.6 g W
1
183.8g W
= 4.738 X1023 atoms W
(
)(
6.02 X 1023 Li atoms
1 mol Li
)(
)
6.02 X 1023 W atoms
1 mol W
)
)
)
Steps
1. Use the PT to calculate the molar mass of one
formula unit
2. Convert the given mass of the compound to
the number of molecules in the sample (use
the molar mass as the conversion factor)
3. Multiply the moles of the compound by the
number of the formula units in a mole
(Avagadro’s number) and solve
4. Check by evaluating the units
Calculate the molar mass (Fe2O3)
2 Fe atoms 2X 55.8 =
111.6
3 O atoms 3 X 16.0 =
+48.0
molar mass
159.6 g/mol
(given mass X 1 mole per molar mass X Form Units per 1 mole)
1.
(
16.8 g Fe2O3
1
)(
1 mol Fe2O3
6.02 X 1023 Fe2O3 Formula units
159.6g Fe2O3
1 mol Fe2O3
)(
= 6.34 X1022 Fe2O3 Formula units
)
1.
89.0g sodium oxide (Na2O)
2.
10.8g boron triflouride ( BF3)
89.0g sodium oxide (Na2O)
Calculate the molar mass (Na2O)
2 Na atoms 2X 23.0 =
46.0
1 O atoms 1 X 16.0 = +16.0
molar mass
62.0 g/mol
1.
(given mass X 1 mole per molar mass X molecules per 1 mole)
(
89.0 g Na2O
1
)(
1 mol Na2O
)(
62.0g Na O
6.02 X 1023 Na2O
2
= 8.64 X1023 Na2O Formula units
Form Units
1 mol Na2O
)
10.8g boron trifloride ( BF3)
Calculate the molar mass (Na2O)
1 B atom 1X 10.8 =
10.8
3 F atoms 3 X 19.0 =
+57.0
molar mass
67.8 g/mol
2.
( given mass X 1 mole per molar mass X molecules per 1
mole)
(
10.8 g BF3
1
)(
1 mol BF3
6.02 X 1023 BF3 Form units
67.8g BF3
1 mol BF3
)(
= 9.59 X1022 BF3 Formula units
)
Steps
1. Determine the molar mass
2. Change given mass to moles by
using molar mass as the conversion
factor. (1/molar mass)
Calculate the number of moles in 6.84g sucrose (C12H22O11)
12 C atoms 12 X 12.0 = 144.0
22 H atoms 22 X 1.0 =
22.0
11 O atoms 11 X 16.0 = +176.0
molar mass
342.0 g/mol
(given mass/1) X (1 mole/ molar mass)
(
6.84 g sucrose
1
1 mol sucrose
) ( 342.0g sucrose)
= 2.0 X10-02 moles of sucrose
1.
2.
3.
16.0g sulfur dioxide, SO2
68.0g ammonia, NH3
17.5g copper(II) oxide, CuO
1.
16.0g sulfur dioxide, SO2
(16.0g/1) (1mole/64.1g ) = 0.250 mol SO2
2.
68.0g ammonia, NH3
( 68.0g/1) (1 mole/ 17.0g) = 4.00 mol NH3
3.
17.5g copper(II) oxide, CuO
( 17.5g/1) (1 mole/ 79.1g) = 0.22 mol CuO
Steps:
1. Find the molar mass of the compound
2. Use the molar mass to convert the given
number of moles to a mass (moles) X (g/mol)
3. Solve
4. Check using dimensional analysis (make sure
units cancel and leaves only grams)
1.
2.
2.
3.
Find the molar mass of the compound (H2O)
H - 2 atoms – 1.0 = 2.0
O - 1 atom - 16.0 = 16.0
18.0 g/mol
Use the molar mass to convert the given number of
moles to a mass (moles) X (g/mol)
(7.5 mol H2O) ( 18.0 g H2O)
( 1 mol H2O)
Solve : 7.5 X 18.0g H2O = 135 g H2O
Check using dimensional analysis (make sure units
cancel and leaves only grams) “mol H2O” cancel each
other out, units are correct!
1.
2.
3.
4.
3.52 mol Si
1.25 mol aspirin, C9H8O4
0.550 mol F2
2.35 mol Barium Iodide, BaI2
(moles) X (g/mol)
1.
What mass of Si = 3.52 mol Si
(3.52 mol Si) (28.1g Si) = 98.9g Si
1
(1 mole Si)
2.
What mass of C9H8O4 = 1.25 mol aspirin, C9H8O4
C -9 atoms – 12.0 – 108.0
H- 8 atoms – 1.0 8.0
O – 4 atoms – 16.0 - 64.0
180.0g/mol
(1.25 mol C9H8O4) (180.0g C9H8O4) = 225.0g C9H8O4
1
(1 mole C9H8O4)
3.
What mass of F2 = 0.550 mol F2
F- 2 atoms – 19.0 = 38.0 g/mol
(0.550 mol F2 ) (38.0 g F2) = 20.9g F2
1
(1 mole F2)
4.
What mass of BaI2 = 2.35 mol Barium Iodide, BaI2
Ba-1 atom – 137.3 - 137.3
I – 2 atoms – 126.9 - 253.8
391.1g/mol
(2.35 mol BaI2) (391.1g BaI2) = 919.1g BaI2
1
(1 mole BaI2)
Know:
1.
What stoichiometry is
2.
What a mole is
3.
How to calculate molar mass of an element
and of a compound
1.
How to determine the number of atoms or
formula units in a given mass of sample
2.
How to determine the number of moles in a given
mass of a sample
3.
How to determine the mass of a given molar
quantity
Review of Calculation Rules
To Find molar (atomic mass of each atom) X (amount of each
mass (g/mol)
atom)
Then add together mass of all atoms
(g/mol)
To Find the #
atoms
in a given
mass
(given mass) X (1mole) /(molar mass (g)) X
(# atoms) /(1 mole)
To Find the #
moles in a
given mass
(given mass) X (1mole)/(molar mass (g)) X
(#atom)/(1mole)
To Find the
mass (g) of a
given molar
quantity
(#moles) X (grams/1 mole) (from molar mass)

Balanced chemical equations relate moles of
reactants to moles of products




Just like when baking, reactants have to be mixed in the
proper proportions to make a certain amount of the desired
product
Specific amounts of reactants produce specific
amounts of product
We can use balanced chemical equations and moles
to PREDICT the masses of reactants or products
When one of the reactants in a reaction are used up,
the reaction stops.

The reactant that is used up is called a limiting reagent
You can not move directly from the mass of one
substance to the mass of the second
Steps
1.
Write a balanced equation
2.
Convert the given mass to moles first!
2.
The coefficients of balanced reactions tell you
the NUMBER OF MOLES of each chemical in
the reactant. (these are used as the conversion
factor)
3.
Once you know the number of moles of any
reactant or product, use the coefficients in the
equation to convert the moles of the other
reactants and products to mass

Ammonia gas is synthesized from nitrogen gas and
hydrogen gas according to the balanced equation :
N2 + 3H2  2NH3
How many grams of hydrogen gas are required for 3.75g
of nitrogen gas to react completely? What mass of
ammonia is formed?
 Reactants and products are related in terms of moles
 The amount of H2 needed depends on the moles of N2
present in 3.75g and the ratio of moles of H2 to moles
of N2 in the equation.
 The amount of ammonia formed depends on the ratio
of moles N2 to moles of ammonia
How many grams of Hydrogen are required for 3.75g of nitrogen to react
completely? What mass of Ammonia is formed?
N2 + 3H2  2NH3
Convert the given mass to moles
Find the # of moles of N2 using molar mass
(3.75g N2) (1 mol N2)
(28.0 g N2)
To find the mass of H2 needed :
The coefficients of the balanced equation shows 3 mol of H2 react with 1
mole of N2 . Multiply moles of N2 by this ratio.
2.(3.75g N2) (1 mol N2) ( 3 mol H2)
(28.0 g N2) (1 mol N2)
Once you know the number of moles of any reactant H2 use the coefficients
in the equation to convert the moles of the other reactants and products
To find the mass of hydrogen, multiply the moles of H2 by the mass of 1
mole of H2.
(3.75g N2) (1 mol N2) ( 3 mol H2) ( 2.0g H2) =3.75 X1 X3 X2.0=0.084 g H2
(28.0 g N2) (1 mol N2) (1 mol H2)
28.0
To find the mass of ammonia produced:
1.
Use the mole ratio of ammonia molecules to
nitrogen molecules to find the moles of ammonia
formed.
(3.75g N2) (1 mol N2) ( 2 mol NH3)
(28.0 g N2) (1 mol N2)
2.
Use the molar mass of ammonia, 17.0g to find the
mass of ammonia formed.
(3.75g N2) (1 mol N2) ( 2 mol NH3) (17.0g NH3) = 3.75 X1 X2 X17.0=4.55gNH3
(28.0 g N2) (1 mol N2) (1 mol NH3)
28.0
When potassium chlorate (KClO3) is heated, it
decomposes to form potassium chloride and
oxygen gas 2KClO3  2KCl + 3O2
a) How many grams of KCl are formed when 28.0g
of KClO3 decompose?
(28.0g KClO3) (1 mol KClO3) ( 1mol KCl)
(122.6 g KClO3) (1 mol KClO3)
b)
(74.6g KCl) = 17.0gKCl
(1 mol KCl)
Use the mass of KCl you determined in part a to
calculate the mass of oxygen gas produced.
(17.0g KCl) (1 mol KCl) ( 3 mol O2)
(74.6 g KCl) (2 mol KCl)
(32g O2) = 10.9g O2
(1 mol O2)
1.
2.
3.
The combustion of propane (C3H8), a fuel used in
backyard grills, produces carbon dioxide and
water vapor. C3H8 + 5O2  3CO2 + 4H2O
What mass of carbon dioxide forms when 95.6g
of propane burns?
Solid xenon hexafluoride is prepared by
allowing xenon gas and fluorine gas to react.
Xe + 3F2  XeF6
How many grams of fluorine are required to
produce 10.0 g of XeF6?
Using the previous reaction, how many grams of
xenon are required to produce 10.0g of XeF6 ?
1.
2.
3.
287 g CO2
4.65 g F2
5.35 g Xe




Avogadro’s principle states that equal volumes of
gasses at the same temperature and pressure
contain equal numbers of moles of gasses.
The molar volume of a gas is the volume that a
gas occupies at 1 atmosphere( 101 kPa, or 760 mm
Hg) of pressure and a temp of 0.0° C or 273°K
(STP).
At STP, the volume of 1 mole of any gas is 22.4L,
(there masses may be different ).
Molar volume is often used in calculations, BUT
BE SURE YOU ARE AT STP!
In the space shuttle, exhaled CO2 is removed from the
air by passing it through canisters of lithium
hydroxide. The following reaction takes place;
CO2 + 2LiOH  Li2CO3 + H2O
How many grams of lithium hydroxide are required
to remove 500.0L of carbon dioxide gas at 101 kPa
pressure and 25.0°C?
Use the molar volume to find the number of moles.
FIRST you must convert the volume of the gas at
25.0°C to a volume at standard temperature!
V = (500.0L CO2) (273 K) = 458 L CO2
(298 K)

Now find the number of moles of CO2
(458 L CO2) ( 1 mol CO2 )
(22.4 L CO2 )
the molar ratio of LiOH to CO2 is 2 to 1 (from the
equation)
 Determine the number of moles of LiOH
(458 L CO2) ( 1 mol CO2 ) ( 2 mol LiOH)
(22.4 L CO2 ) ( 1 mol CO2)

To Convert the number of moles of LiOH to
mass, use the molar mass, 23.9g/mol.
(458 L CO2) ( 1 mol CO2 ) ( 2 mol LiOH) (23.9g LiOH) = 977 g LiOH
(22.4 L CO2 ) ( 1 mol CO2) (1 mol LiOH)
What mass of glucose (C6H12O6) must be broken
down in your body to produce 2.50 L of CO2 at
273°C and 1 atmosphere?
C6H12O6 + 6O2  6H2O + 6 CO2

(2.50L CO2) (1 mol CO2) (1 mol C6H12O6) (180g C6H12O6) = 3.35 g C6H12O6
( 22.4 L CO2) ( 6 mol CO2) (1 mol C6H12O6 )
When steel wool burns in air, this reaction occurs:
4Fe + 3O2  2 Fe2O3
What volume of oxygen, measured at 725 mm Hg (760
mm Hg is 1 atm) and 25.0°C, is required to react
with 100.0g of iron?
(100.g Fe) (1 mol Fe) ( 3 mol O2) (22.4 L ) (298K) (760 mm Hg) = 34.5 L O2
(55.8 g Fe) (4 mol Fe) (1mol O2) (273 K) (725 mm Hg)
What mass of sulfur must burn to produce
3.42 L of SO2 at 273°C and 101 kPa?
The reaction is S + O2  SO2
Answer: 2.45g S
2. What volume of hydrogen gas can be
produced by reacting 4.20 g of sodium in
excess water at 50.0°C and 106 kPa?
The reaction is 2Na + 2H2O  2NaOH + H2
Answer:2.31 L H2
1.
Review of Calculation and Rules
To Predict the (grams) X (moles/gram) X (mole ratio reactants
mass of
to products) X (moles/ mass )
Reactants and
products
Molar Volume The molar volume of a gas is the volume that a
gas occupies at 1 atmosphere( 101 kPa, or 760
mm Hg) of pressure and a temp of 0.0° C or
273°K (STP).
To Determine Convert gas volume to volume at STP
the # of grams (gas volume at STP)X(1mol/22.4L) X(molar ratio)
of a reactant to X (molar mass (g/mol)) = grams
make a
volume of gas