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Chemical Quantities & Stoichiometry
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
What is a mole?
How do find the mass of a compound?
What are the indicators of a chemical
reaction?
What does the subscript, coefficient, and
symbols mean in a formula?
◦ Ex:
5Ba3(PO4)2

Molar Mass: The mass of one mole of an element
or compound.
◦ 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.
Find the molar mass:
1.
Sr
= 87.62 grams/mol
2.
MgBr2
3.
Ba3(PO4)2
24.3 + (2x 79.9) =
Ba =
P
=
O =+
=
184.1 grams/mol
3 x 137.38 g
2 x 30.97 g
8x
16 g
602.08 grams/mol

Calculate the molar mass
1. KCl
2. HC2H3O2
3. CO2
4. Ba(OH)2
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
Law of Definite Proportions:
In samples of any chemical compound, the
masses of the elements are always in the
same proportion.
=> Allows us to write chemical formulas.
Percent Composition by Mass
◦ Worksheet
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
Percent Composition - % by mass of each
element in a compound
Percent =
Part x 100
Whole
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
Al:
S:
O:
Percent Comp = Mass of 1 element
x 100
Mass of compound
Example: Find the mass percent composition
of Al2(SO4)3
2 x 27 = 54
3 x 32 = 96
12 x 16 = 192
54 x 100= 15.8%
% Al:
342
% S:
96 x 100 = 28.1%
342
%O:
192 x 100 = 56.1%
342
Total Comp. = 342

Calculate the percent composition
1. NaOH
Na________
2. Ca3(PO4)2
Ca_______
P________
O________
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

By mass
By volume
By counting the # of atoms/molecules
◦ We can use a word like a “dozen” to specify a certain
quantity.
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Mole (mol): SI unit for measuring the amount of
a substance.
1 mol = 6.02 x 1023 representative particles
Avogadro’s Number: 6.02 x 1023
Representative Particle: smallest unit that has
all the characteristics of that substance.

Atom
Element (ex. Cu): ___________________
◦ Exception: The representative particle of the
7 diatomic elements is a molecule. (ex. H2)

Molecule
Covalent compound (ex. H2O): ____________

Ionic Compound (ex. NaCl):
Formula Unit
_____________
(Molecule)

4 moles Ca =
4 moles Ca
atoms Ca.
6.02 x 1023 atoms Ca
1 mole Ca
= 2.41 x 1024 atoms Ca

5 x 1018 atoms Cu =
Cu.
5 x 1018 atoms Cu
moles
1 mole Cu
6.02 x 1023 atoms Cu
= 8.3 x 10-6 moles Cu

9.2 moles F2 =
9.2 moles
molecules F2?
6.02 x 1023 molecules F2
1 mole
= 5.5 x 1024 molecules F2

9.2 moles F2 =
9.2 moles F2
6.02 x 1023 molecules F2
1 mole
= 1.1 x 1025 atoms F
atoms F?
2 atoms F
1 molecule F2

3.4 moles C2H4 =
atoms?
3.4 moles C2H4
6.02 x 1023 molecules C2H4
1 mole C2H4
= 1.22 x 1025 atoms
total
6 atoms
1 molecules C2H4
1.
0.45 mole He = ________atoms
2.
7.6 x 1023 molecules CH4 = ________moles
3.
15 moles H2O = __________ atoms
1 mol = molar mass (in grams)

68 grams F2 =
68 grams F2
moles F2?
1 mole F2
38 grams
68 / 38 = 1.8 moles F2
1 mol = molar mass (in grams)

0.24 mol CH4 =
0.24 mol CH4
grams CH4?
16.042 grams
1mol
0.24*16.042 = 3.85 grams CH4
1.
25 g NaCl = _________ moles
2.
3.2 moles of MgCO3 = _______ grams
3.
3.4 moles Au = _________atoms

Standard Temperature and Pressure (STP):
0oC, 1 atm
◦ See Reference Tables

Avogadro’s Hypothesis: equal volumes of gases
at the same temperature and pressure contain
equal numbers of particles.
◦ At STP, 1 mole of any gas occupies a volume of 22.4 L.
1 mol = 22.4 L at STP (gases only!!!)

5.4 moles He =
5.4 moles He
L He at STP?
22.4 L He
1 mole He
5.4 x 22.4 = 120.96 L He

5.4 moles CH4 =
5.4 moles CH4
L CH4 gas at STP?
22.4 L CH4
1 mole CH4
5.4 x 22.4 = 120.96 L CH4

560 L SO3 =
560 L SO3
mol SO3
1 mole SO3
22.4 L SO3
560 / 22.4 = 25 mole SO3
1.) 0.75 mole N2 = ________L
2.) 11.2 L NH3 = _______ mole
3.) 100 grams Ba2S = ________ mole
4.) 1.5 x 1020 atoms Cu = _________ moles
22.4 L
at STP
(gases only)
LITERS
1 MOLE
Molar Mass
6.02 x 1023 particles
GRAMS
ATOMS/MOLECULES
5.6 grams of CaCl2 = _________ L CaCl2
6x1024 atoms W = _________grams
1.) 1.5 x 1025 molecules CO2 = ___________
grams
2.) 5g O2 = __________liters
3.) 54 L N2 = __________ molecules
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
Hydrates are ionic compounds (salts) that have a definite
amount of water as part of their structure.
This “water of hydration” is released as vapor when the
hydrate is heated. The remaining salt is known as the
anhydrous salt. The general reaction for heating a hydrate is:
CoCl2·6H2O + heat
->
CoCl2 +
H 2O
Hydrate + heat -> Anhydrous salt + water
Percent
Composition
18.8% Na
# grams in
100 grams
How many moles
of each element?
Divide by smallest
#moles
29.0% Cl
52.2% O
Empirical Formula: ___________________

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.

Calculate the empirical formula of a
compound containing 0.90g Ca and 1.60g Cl.
◦ Step 1: Convert GRAMS to MOLES.
 Ca: 0.90g
1 mole
40.1 g
= 0.0224 mole Ca
 Cl: 1.60g
1 mole
35.5 g
= 0.0451 mole Cl

Step 2: DIVIDE the # of moles of each
substance by the smallest number to get the
simplest mole ratio.
Ca: 0.0224 = 1
0.0224
Cl: 0.0451 = 2.01 ~ 2
0.0224
CaCl2
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
Step 3: If the numbers are whole numbers,
use these as the subscripts for the formula.
If the numbers are not whole numbers,
multiply each by a factor that will make them
whole numbers.
Look for these fractions:
◦ 0.5  x 2
◦ 0.33  x 3
◦ 0.25  x 4
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?
Find the molecular formula of ethylene glycol
(CH3O) if its molar mass is 62 g/mol.
Step 1: CH3O = (12) + (3 x 1) + (16) = 31
Step 2: 62 / 31 = 2
Step 3: 2 (CH3O)  C2H6O2
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.
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◦
What is its empirical formula?
◦
What is its molecular formula?
58.8% C
1 mole C
12 g C
= 4.9 / 1.9 = 2.5 x 2 = 5
9.8% H
1 mole H = 9.8 / 1.9 = 5 x 2 = 10
1gH
31.4%O
1 mole O = 1.9 /1.9 = 1 x 2 = 2
16 g O
C5H10O2  (5x12) + (10x1) + (2x16) = 102 g/mol
Empirical mass = molecular mass, so molecular formula is
the same  C5H10O2
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If a compound is 40% C, 7% H, and 53% O,
what is its empirical formula?
What is the molecular formula for this
element if the molecular mass is 180 g/mol?
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How would you do this???
 grams (Molar Mass)
(Density) grams
liters
mole
Example:
 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?
3.58 g
L
22.4 L
1 mole
3.58 x 22.4 = 80.3 g/mole

What is the density of krypton gas at STP?
83.8 grams Kr
mole
1 mole
22.4 Liters
83.8 / 22.4 = 3.74 g/L Kr
Molarity: the number of moles of solute
dissolved in one liter of solution.
◦ It is a way to express concentration
Molarity= moles of solute
liters of solution
or
M = mol
L
How many grams of NaOH would you need to
add to .35 liters of water to make a 1.50 M
solution?
*NOTICE: No mention of STP and NaOH is not a
gas
Battery acid is generally 3M H2SO4. How many
grams of H2SO4 are in 400. mL of this
solution?
1. What is the density of chlorine gas at STP?
2. How many moles of H2SO4 are in 20 mL of
a10M solution?
3. How many liters of HCl are there if you have
a 2M solution and .15 moles of solute?
Given the following equation:
2 C4H10 + 13 O2 → 8 CO2 + 10 H2O
show all examples of molar ratios (Hint: 5)


Stoichiometry: The calculation of quantities of
substances involved in chemical reactions.
N2 (g) + 3H2 (g)  2NH3 (g)


The above equation could be read:
1 mol of N2 reacts with 3 moles of H2 to yield
2 moles of ammonia.

If you follow these steps, you will be an expert at
stoichiometry
1.) Convert your units to moles
2.) Use a mole ratio
3.) Convert to the units asked for
in the problem
2A + B  3C + 7D

Given the number of moles of reactant A
(ex. 6 moles A), I can find:
◦ 1) The number of moles of reactant B needed to
react completely with 6 moles of A (all 6 moles are
used up).
◦ 2) The number of moles of product C formed.
◦ 3) The number of moles of product D formed.
N2 (g) + 3H2 (g)  2NH3 (g)
There is a 1:3:2 mole ratio
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?
N2 + 3H2  2NH3
If you have 2 moles of N2, how many moles of
NH3 will be produced?
2 mol N2
2 mol NH3
1 mol N2
2 x 2 / 1 = 4 moles NH3
N2 + 3H2  2NH3
If you want 5 moles of product, how many
moles of hydrogen gas do you need?
5 mol NH3
3 mol H2
2 mol NH3
5 x 3 / 2 = 7.5 moles H2
N2 + 3H2  2NH3
How many moles of nitrogen are needed to
react completely with 8 moles of hydrogen?
8 mol H2
1 mole N2
3 mole H2
8 x 1 / 3 = 2.67 mole N2
2H2O2 -> 2H2 + O2
1.
2.
If you have 10 moles of H2, how many
moles of H2O2 did you originally have?
If you started out with 3.5 moles of H2O2,
how many moles of O2 did you create?
****The only way to convert
from one compound to
something totally different in
the reaction is to use the
MOLE TO MOLE RATIO from
the coefficients!!!****
Note – If you don’t have moles already, your
first step is to convert to moles!
Mole Review:
22.4 L
at STP
1
mole
Molar
Mass
6.02 x 1023
particles
LITERS
OF GAS
AT STP
Molar Volume
(22.4 L/mol)
MASS
IN
GRAMS
Molar Mass
MOLES
(g/mol)
6.02 
1023
particles/mol
Molarity (mol/L)
LITERS
OF
SOLUTION
NUMBER
OF
PARTICLES
1.
2.
How many liters of oxygen are required to
burn 3.86 L of carbon monoxide?
2CO (g) + O2 (g)  2CO2 (g)
How many liters of PH3 are formed when
0.42 L of hydrogen reacts with phosphorus?
P4 (s) + 6 H2 (g)  4 PH3 (g)
1.
Find the liters of sugar required to produce
1.82L of carbon dioxide gas at STP from the
reaction described by the following
equation:
C6H12O6 -> 2C2H6O + 2CO2
N2 (g) + 3H2 (g)  2NH3 (g)
1.
2.
How many grams of hydrogen gas are
required for 3.75 g of nitrogen gas to react
completely?
What mass of ammonia is formed when
3.75 g of nitrogen gas react with hydrogen
gas?
N2 + 3H2  2NH3
How many grams of H2 are required to produce
5.0 grams of NH3?
Grams NH3  moles NH3  moles H2  grams H2
5.0 g NH3
1 mole NH3
17 g NH3
3 mole H2
2 g H2
2 mole NH3 1 mole H2
5.0 x 3 x 2 / 17 / 2 = 0.88 g H2
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)
◦ How many grams of fluorine are required to produce 10.0 g of XeF6?
◦ How many grams of xenon are required to produce 10.0 g of XeF6?

How moles of CO2 are produced when 52.0 g C2H2
burns?
2C2H2 (g) + 5O2 (g)  4CO2 (g) + 2H2O (g)
How many liters of hydrogen gas are formed from
50 grams of potassium?
2K (s) + 2H2O (l)  2KOH (aq) + H2 (g)

How many molecules of oxygen are produced by
the decomposition of 6.54 g of potassium chlorate
(KClO3)?
2KClO3 (s)  2KCl (s) + 3O2 (g)


How many molecules of oxygen are produced by
the decomposition of 6.54 g of potassium
chlorate (KClO3)?
2KClO3 (s)  2KCl (s) + 3O2 (g)
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)
C5H12 + 8O2  5CO2 + 6H2O
1)
2)
3)
If you have 55.0g of C5H12, how many grams
of water could you make?
If you start with 6.0 liters of oxygen, what
mass of carbon dioxide could you make?
If you end up with 17 liters of CO2, what
volume of oxygen would you have started
with?

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.
◦ Determines the amount of product that is produced.

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.

You have :
◦ 1 loaf of bread (containing 14 slices of bread)
◦ 4 jars of peanut butter
◦ 2 jars of jelly
A) How many peanut butter and jelly
sandwiches can you make?
B) What is the limiting reagent?
C) What are the reactants in excess?
**The amount of product
that can be formed in a
reaction is always
determined by the limiting
reactant!!**

A s’mores MUST have:
◦ 2 graham crackers
◦ 2 pieces of chocolate
◦ 1 marshmallow

If you had:
◦ 8 graham crackers
◦ 4 pieces of chocolate
◦ 6 marshmallows
How many s’mores could you make?
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? _________________
- How much of your excess reagent do you have
leftover? _____________________
3) How many grams of water can be produced by
the reaction of 2.40 mol C2H4 with 7.4 mol O2?
C2H4 (g) + 3O2 (g)  2CO2 (g) + 2H2O (g)
Limiting Reagent? ______________
Excess Reagent? _______________
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
Turn lab into the bin-Name on the sheet!
Take out your stapled packet and complete
the left side of the second page
(stoichiometry and limiting reagents)
Quiz tomorrow will be based on the left side
1)What is the percent composition of
Ca3(PO4)2?
2)
If you have 15.0 g of oxygen gas (at
STP), how many liters do you have?
How many grams of titanium (IV) oxide
can be produced when 80 g of TiCl4
reacts with 20 g O2 as shown in the
unbalanced equation?
TiCl4 (s) + O2(g)
TiO2(s) + Cl2(g)
3)

Theoretical Yield: the maximum amount of
product that could be formed from the given
amounts of reactants (ideal conditions).
◦ Calculated using stoichiometry.

Actual Yield: the amount of product that
actually forms in a lab.
◦ Actual yield is usually less than theoretical yield.

Percent Yield =
Actual Yield
Theoretical Yield
x 100%
1.
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?

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)
1) What is the formula (molar) mass of Mg(NO3)2?
2) What is the percent oxygen in #1?
3) If the empirical formula is NO2 and the molar
mass is 138 g/mol, what is the molecular
formula?
4) What volume is 65g of N2 gas at STP?
5) If you start with 23 L of O2, what mass of water
could you obtain, based on this reaction:
O2 + 2H2  2H2O
6) If 8 grams of fluorine are mixed with 8 grams of
potassium chloride, and 5.1 grams of potassium
fluoride are produced in the lab, what is the
percent yield?


How many molecules of NaCl are there if you
have 25 grams of NaCl?
If 12.02 grams of Ni(CO)4 yielded 2.53 g Ni
when you performed the experiment,
determine the percent yield.
◦ Ni(CO)4 (g)
Ni(s) + 4 CO (g)


Make sure your answers are on the test itself and
all your work is on the blue sheet
Staple together your work and test in this order:
◦ Test (w/ answers)
◦ Blue paper (work)


Turn into the bin
Pick up the gas laws worksheet
◦ P = pressure (atm, kPa, mmHg, torr-see reference table)
◦ V= volume (L)
◦ T=temperature (must be in Kelvin!)
 K = 273 + °C



Work and answers in bin – MAKE SURE YOUR
NAME IS AT THE TOP!
Test can go on the outside with your
laminated sheets
Finish your textbook questions pg. 379-380
#61-83