Unit Five:
Moles and Stoichiometry
Teacher Copy
0
Unit Five: Moles and Stoichiometry
Table of Contents:
Pg. 1
Table of Contents
Pg. 2
Mole Map
Pg. 3
The Mole
Pg. 3-4
Mole Calculations (Assignment on Pg. 18 & 19)
Pg. 5
Molar Mass (Assignment on Pg. 20)
Pg. 5-6
Mole-Mass Conversions
Pg. 6
# of Items- Mass Conversions (Assignment on Pg. 21)
Pg. 7
Percent Composition (Assignment on Pg. 22)
Pg. 7
Empirical Formula (Assignment on Pg. 23)
Pg. 8
Calculating moles of a gas (Assignment on Pg 24)
Pg. 8-9
Concentration (Assignment on Pg. 25)
Pg. 9-10
Stoichiometry (Assignment on Pg 26)
Pg. 10
Volume- Volume Stoich (Assignment on Pg. 27)
Pg. 11
Mass Stoich (Assignment on Pg. 28)
Pg. 12
Mixed Stoich (Assignment on Pg. 29)
Pg. 12-14
Limiting Reagents (Assignment on Pg. 30 & Pg. 31)
Pg. 15-16
Limiting Reagent Lab
1
Mole Map
2
Avogadro’s Number:

Since atoms and molecules are so tiny we cannot work with them individually; instead we
work with moles of atoms and molecules.

A mole = Avogadro’s number = 6.022x1023 (similarly to how 1 dozen=12)

Named after Amedeo Avogadro

The name avogadro’s number was changed to avogadro’s constant when it became an SI unit.

Avogadro’s number is often represented by the symbol n in formulas, but the units are moles
(mol).

Because atoms are so small we work in moles instead of atoms. If we had 1 mole of peas on
the earth it would cover the earth 8 km deep.
Calculate the number of moles:
If you have 24 pencils, how many dozen pencils do you have?

Remember we are actually multiplying by 1, so we are not changing the number.

Unit Conversions: # & units given x units we want
.
units we were given
Map: When converting # of items to moles or vice versa, we must remember that 1 mole = 6.022 x 1023
Ex. If you have 3.60 x 1020 pencils how many moles of pencils is that? (Remember Sig. Figs)
Examples:

If you have 1.20 mols of carbon dioxide how many molecules of carbon dioxide do you have?

If you have 4.678x1020 atoms of carbon, how many moles is that?

Note that Moles can be used to calculate quantity of anything. This means it could be atoms,
molecules, formula units, ions, peanuts, books… anything!
3
More Calculations with Moles:

How many dozen pieces of bread are there in 2 dozen sandwiches? (assume 2 pieces of bread
per sandwich)

How many atoms of carbon are in 1 molecule of CO2? 1

How many atoms of oxygen are in 1 molecule of CO2? 2

How many moles of carbon are in 1 mole of CO2? 1 mol

How many moles of oxygen are there in 2.0 moles of carbon dioxide?

How many moles of oxygen are there in 2.0 moles of acetic acid?
You may need 2 conversion factors if you are going from # of one thing to moles of another thing or
vice versa. For example, if you are given atoms of carbon in carbon dioxide and asked to find the
number of moles of carbon dioxide, you will need 2 conversion factors.
Map:
A solution of calcium chloride contains 4.3 x 1024 chloride ions. How many moles of calcium chloride
are in the solution?
A sample of hydrogen gas contains 5.0 moles of hydrogen molecules. How many hydrogen atoms are in
the sample?
If there are 3.98x1031 atoms of hydrogen in a sample of acetic acid, how many moles of acetic acid are
there?
4
Molar Mass:

Avogadro’s number is not a randomly contrived number. It is based off of the
carbon -12 isotope. There are 6.022x1023 atoms of carbon in 12.00grams of
carbon-12 isotope. This means if there is a mole of something, there is the
same number of that thing as there are atoms in 12 grams of carbon.

Molar mass of an element is the mass that one mole of that element has. For
carbon-12, this would be 12g/mole. But for carbon it is 12.01g/mole, because
it is the average mass of all isotopes (similarly to how the average atomic mass
was the relative average atomic mass of each isotope).

The molar mass of an element is found on the periodic table and has units
grams/mole.
Molar Mass of Atoms:

The molar mass of an atom is found by using the periodic table. The molar mass is the number at
the bottom of the box, but you need to add the units g/mol
Examples:
The molar mass of hydrogen: 1.008g/mol
Aluminum: 26.98g/mol
Oxygen: 15.999g/mol
Scandium: 44.96g/mol
Molar Mass of Compounds:

Molar mass of a compound is the mass that one mole of that compound has. This can be found by
adding up the molar mass of each atom in the compound.
Examples
Hydrogen molecule: H2: H + H= 1.008g/mol+1.008 g/mol=2.016 g/mol
NaOH: Na + O + H= 22.99g/mol+15.99 g/mol+1.008 g/mol= 39.99 g/mol
C6H12O6: 6xC + 12xH + 6xO= 6x12.01 g/mol + 12x1.008 g/mol + 6x15.99 g/mol= 180.16 g/mol
Aluminum Borate: AlBO3: Al+B+3xO= 26.98g/mol + 10.81g/mol + 3x15.99g/mol= 85.76g/mol
Mole-Mass Conversions
You can use the formula: mass = molar mass x moles but I would recommend using the unit conversion
method.
# & units given x units we want
.
units we were given
Map Update
5
Ex. How many grams of Sodium are in 7.124 moles of sodium?
Write down what we are given and the division line for the conversion
factor.
Then we write the units we have on the bottom and the units we want
on the top of the conversion factor.
Then we fill in the appropriate numbers in our conversion factor (molar
mass paired with grams and 1 paired with moles).
Ex. How many moles of carbon are in a 3.98 g sample of carbon?
**Remember: whenever atoms, molecules, formula units or ions is mentioned, you will need to use
Avogadro`s number. When mass or grams or weight is mentioned, you will need to use molar mass in
your conversion factor**
Ex. How many moles of water are in a 250.0 g sample of water?
# of items-Mass conversions

We can use molar mass to convert mass to moles and moles to mass and we can use avogadro’s
number to convert moles to units and units to moles. This means we can now go from mass to
units and units to mass using molar mass and Avogadro’s number.

For these questions we will have to use 2 conversion factors. Whenever you are given something
that is not moles, you must convert to moles first (using a conversion factor).
o You can use the “map” to help you determine the path you need to take, but you will not
be given the “map” on an exam.
Ex. How many atoms of sodium are in 5.34g of sodium?
Ex. How many grams of carbon dioxide are there in 7.398x1021 atoms of carbon dioxide?
Remember: if the problem mentions mass, weight or grams use molar mass.
Percent Composition:

Percent composition tells us the percent by mass that one element takes up in a compound.

This is done by taking the molar mass of the atom and multiplying it by how many of that
atom is in the compound and dividing by the molar mass or the compound.
6

% composition = (# of x atoms  molar mass of element x)
molar mass of compound
Ex. What is the percent composition of each element in sulfuric acid?
Ex. How many grams of hydrogen could be extracted from 100 grams of sulfuric acid?
% composition x mass of compound
Ex. How many grams of carbon could be extracted from 15 grams of methane?
Determining Empirical Formula:

An empirical formula of a compound is the lowest whole number ratio the components or
elements of the can make.

In order to determine the empirical formula, we need the percent composition of each atom
in the compound.
o We then take the % composition of each atom (ignore all units) and divide it by that
atoms molar mass. We do this for each atom in the compound.
o We then take the lowest number we get and divide each other number by this number
and round our answers to the 1’s place. This will give us the lowest ratio that these
elements combine in.
Ex. 24.7% Potassium, 34.8% Manganese, and 40.5% Oxygen.
Ex. 2.8% Hydrogen, 97.2% Chlorine
Calculating moles of a gas:

In 1811, the Amedeo Avogadro proposed the principle that equal volumes of gases at the same
temperature and pressure contain equal numbers of molecules. He determined that at standard
temperature and pressure, one mole of gas occupies 22.41410 L (usually rounded to 22.4 L).

This means that 1 mole of carbon dioxide gas and one mole of hydrogen gas will take up the
same volume if they have the same temperature and pressure. At standard temperature and
pressure, this volume is 22.4L. Standard temperature and pressure is (273K (0oC) and 1 atm).

This means that 1 mole of any gas at STP takes up 22.4L of space. We can use this relation to
7
convert volume of a gas to moles of a gas.
(Update Map)
Ex. How many moles of neon gas (@ STP) are in 6.9L?
Ex. How many litres will 2.83 moles of carbon dioxide gas take up at standard temperature and
pressure?
Ex. How many molecules of oxygen gas are there in 32.3L at STP?
Calculating Concentration:

Solubility is the physical property referring to the ability of a substance (solute) to dissolve in
a solvent.

Molarity = Number of moles of solute dissolved in one liter of solution
o has units “mole/L”, represented by “M” (molar)
o To calculate molarity you must first find the number of moles of solute and divide by
the volume of solution (in liters).
o Molarity=
moles of solute
.
volume of solution (L)
o Concentration questions can be difficult to complete using the conversion factor
method, so it may just be easier to use the formula.
(Update Map)
Examples:
Determine the molarity of a solution created by dissolving 2.0 g of sodium hydroxide in 30mL of
water.
Determine what volume of water you would need to add to 3.5 moles of sodium hydroxide to make a
1.0 M solution.
Determine the mass of sodium hydroxide that would need to be dissolved in 0.500L of water to
create a 0.50 M solution.
8
Stoichiometry

Stoichiometry uses the quantitative relationship between reactants and products in a
chemical reaction to determine the quantities of reactants and products in a certain chemical
reaction.
(Look at map)
Mole-mole problems:

In a balanced equation we can determine how many moles of one reactant or product is
needed to react with or produce another reactant or product. We can use a conversion factor
and use the coefficients for the conversion factor.
Ex. 4P +5O2  2P2O5

If we have 5 molecules of oxygen how many molecules of di-phosphorus pentoxide will be
produced?
How many atoms of phosphorus are required?

If we have 5 dozen molecules of oxygen how many dozen molecules of di-phosphorus
pentoxide will be produced?
How many dozen atoms of phosphorus are required?

If we have 5 moles of oxygen molecules how many moles of di-phosphorus pentoxide
molecules will be produced?
How many moles of phosphorus atoms are required?
Ex. NaCl +F2  NaF + Cl2
How many moles of fluorine molecules are needed to completely react with 3 moles of sodium
chloride?
MAKE SURE EQUATION IS BALANCED!
Ex. Na + H2O  NaOH + H2
How many moles of hydrogen (molecules) are produced from a reaction with 5 moles of water and
excess sodium?
BALANCE!
9
Volume-Volume Stoichiometry Problems:

Remember the volumes of different gases at the same temperature and pressure have the
same volume. For example 1 Litre of Oxygen gas would have the same number of molecules
as 1 Litre of Nitrogen gas at the same temperature and pressure. Also, If we had 1 mole of
oxygen and 1 mole of Nitrogen gas at the same temperature and pressure, they would have
the same volume.
(Update Map)
Ex. H2 + NO  H2O +N2
If we use 2 litres of nitrogen oxide, how many litres of hydrogen would be needed?
How many litres of Water vapour?
How many litres of nitrogen?

The coefficients can be used to relate volume as well as moles, because gases of equal
volumes contain the same number of moles (as long as we have constant temperature and
pressure).
Ex In the above reaction, if we have 7.5 Litres of Hydrogen gas, how many Litres of Nitrogen oxide
are needed to completely react?
Ex. C3H8 + O2  CO2 + H2O
For the reaction above, how many liters of water vapor are produced when 18.6L of propane gas are
reacted with an excess of oxygen?
How many Litres of oxygen gas are required to produce 4.67 Litres of CO2 (assuming we have
unlimited propane)?
Mass Stoichiometry

We know that the mole to mole ratio of any 2 substances involved in the reaction is equal to the
ratio of their coefficients in the balanced equation.

We are also able to convert mass to moles and moles to mass.

This means we are now able to determine the mass of one compound in a reaction, given the
mass of another.
(Look @ Map)
10
Ex.
Iron metal mixes with Copper (II) Sulfate in solution to form a reaction. If we use 2.24g of iron
metal, how much copper (II) sulfate is needed to fully react with the iron?

Balanced equation:

We are going to go from mass of Iron to mass of copper (II) sulfate. To do this we need to go
from mass of iron to moles of iron, then from moles of iron to moles of copper (II) sulfate. Then
we need to go from moles of copper (II) sulfate to mass of copper (II) sulfate. This means we
need THREE conversion factors. (See Map)
Ex. For the reaction between sodium hydroxide and sulfuric acid, if we have a mass of 5.6g of
sodium hydroxide and excess sulfuric acid, what is the mass of sodium sulfate produced?
Ex. KClO3  KCl + O2
In the reaction above, if we produce 5.9g of potassium chloride, what is the mass of potassium
chlorate that we started with?
What mass of oxygen would be produced?
Look at mass of reactants vs. mass of products:
11
Mixed Stoichiometric Problems
Ex. For the reaction between zinc and hydrochloric acid:
If we start with 0.500g of zinc, what volume of 1.0M solution hydrochloric acid do we need?
Ex. For a reaction between vinegar & baking soda (acetic acid and sodium bicarbonate):
Note: A 5.0% solution of vinegar means there are 5.0 g of acetic acid for every 100mL of solution. So
if we use 50mL of vinegar that means we are using 2.5g of acetic acid.
What volume of carbon dioxide will be produced if we mix 50mL of 5% vinegar with excess baking
soda?
First figure out the equation:
Acetic acid + sodium bicarbonate 
And carbonic acid decomposes to produce water and carbon dioxide:
Final reaction:
Calculations:
What mass of sodium bicarbonate is needed to complete this reaction?
12
Limiting Reagents:

In a chemical reaction, you may have one reactant that limits the other one, because there is not
enough of it to react with all of the other reactant.
Sandwiches Analogy:

When making a sandwich you need two pieces of bread and on piece of bologna. If we had 20
pieces if bread and 3 pieces of bologna we could only make 3 sandwiches.
Demonstration
Observe 4 reactions of vinegar and sodium bicarbonate all will have 50mL of vinegar and the mass of
sodium bicarbonate will change.
Predict what you think will happen when 50mL of vinegar are mixed with 1.0g of baking soda, 2.0g of
baking soda, 3.5g of baking soda, and 4.5g of baking soda if we put a balloon over the Erlenmeyer flask
to capture any gas.
Observations: (draw what you see)
Yesterday’s calculations we found out that when we use 50mL of 5% acetic acid solutions, we require
3.5g of sodium bicarbonate to completely react.
Trial 1: use 1.0g of NaHCO3 and 50mL CH3COOH: extra vinegar left over: NaHCO3 all used
Trial 2: use 2.0g of NaHCO3 and 50mL CH3COOH: extra vinegar left over: NaHCO3 all used
Trial 3: use 3.5g of NaHCO3 and 50mL CH3COOH: perfect amount of each: both all used up
Trial 4: use 4.5g of NaHCO3 and 50mL CH3COOH: extra NaHCO3 left over: vinegar all used up
In most chemical reactions one reactant is in excess and one is limiting.
Deciding which reactants are the limiting reagents and the reactants in excess when given
masses of both reactants:
1. Write the balanced chemical equation for the chemical reaction
2. Determine how many moles of each reactant you have.
3. Use the stoichiometric ratio to determine which reactant it limiting. You can complete this step
by choosing one reactant (reactant a) and figuring out how many moles of the other reactant
(reactant b) are required. If you need more moles of reactant be than you have, reactant b is your
limiting reagent. If you have more moles of reactant b than what you calculated as required,
reactant a is limiting.
Remember the limiting reagent is the reactant that you do not have enough of.
4. Use the limiting reagent to determine the amount product produced.
5. If you are asked to determine the amount of excess reagent you have, you must use the mole
quantity of the limiting reagent to determine the mass of the other reactant required. Then you
would subtract the mass required from the mass you have.
13
Examples
Ex 1. _____Zn + _____HCl -----> _____ZnCl2 + _____H2
If we combined 2.0g of zinc with 3.0g of HCl, which of the reactants is limiting? How much of the excess
reagent will be left over? What volume of H2 will be produced?
Zn= 2.0g, HCl= 3.0g
Moles of Zn:
Moles of HCl:
Stoichiometric ratio:
Limiting reagent:
Excess left over:
Volume of H2 produced:
Ex 2. _____CaCO3 + _____HCl -----> _____CaCl2 + _____CO2 + _____H2O
If we combine 2.0g CaCO3 with 1.0g of HCl, which reactant is limiting?

Determine # of moles of each reactant: (Balance equation first)

Use stoichiometric ratio to determine limiting reagent:
Ex 3. For the reaction between 0.287g of magnesium and 10mL of 0.50M hydrochloric acid, determine
the limiting reagent, the mass of excess reagent left over, and the volume of hydrogen gas produced.
14
Date:
Class:
Limiting Reagents Lab
Name:
Partners:
There is no background information for this lab. Instead you should use your notes as background
information.
Purpose: to investigate the stoichiometric relationship of a chemical reaction involving a precipitate.
Materials:
25mL graduated cylinder
2 100mL beakers
Spatula/scoopula
Funnel
Filter paper
Erlenmeyer flask
Distilled water
Sodium carbonate
Calcium chloride dihydrate
Procedure:
Day 1:
1. Gather together the following: a 25 mL graduated cylinder, two 100 mL beakers, and spatula.
2. Wash all equipment. Dry equipment thoroughly with paper towel.
3. Set up your filtration system (funnel in an Erlenmeyer flask).
4. Using tape & a pen, label the 100 mL beakers “sodium carbonate” and “calcium chloride
dihydrate”.
5. Weigh 0.50 g sodium carbonate & 1.00 g calcium chloride dihydrate into separate 100 mL beakers
(place the beaker on the scale, press tare, when the scale reads 0 carefully add the desired mass a
little at a time). Record the masses you’ve used.
Mass of sodium carbonate:__________________________g
Mass of calcium chloride dihydrate:__________________________g
6. Add 10 – 15 mL of distilled water to each 100 mL beaker. Gently swirl the contents of the beakers
to dissolve the solids. All solid should be dissolved before you proceed (this will take 1 – 2
minutes).
7. Pour the calcium chloride dihydrate solution into the sodium carbonate solution (not the other
way around). You will see calcium carbonate form. Swirl this mixture gently for 10 – 20 seconds.
8. Get two pieces of filter paper. Weigh both together and record the mass: ______________________ g
9. With the pieces of filter paper together, fold them to make a cone (you may want to wet the filter
paper). Place the empty beaker under the funnel. Swirl the mixture & pour it into the funnel. Add
3-5 mL of distilled H2O to the beaker that held the mixture. Using a spatula, thoroughly scrape the
sides and bottom of the beaker to dislodge all remaining residue. Pour this liquid into the funnel.
Repeat this rinsing procedure (3-5 mL distilled H2O and scraping) twice more to ensure that you
have all of the calcium carbonate residue. You are not measuring the amount of filtrate (amount of
water in the bottom) so don’t worry about adding more water than recommended.
15
10. If the filtrate appears cloudy, pour the filtrate back into the funnel to filter the liquid for a second
time. If the filtrate is clear, carefully remove the filter paper (with solid) and lay it on a piece of
paper towel. Put your name on the paper towel and leave it on the counter at the side of the room.
11. Wash equipment well with tap water. Return equipment & wipe off your bench.
Day 2
12. Weigh the filter paper containing the dry CaCO3. Be careful not to spill any solid.
13. Calculate the mass of CaCO3 produced: filter paper & CaCO3 ________________ g - filter paper (see
above) _____________________ g = ____________________ g. Keep your filter paper (with product) and
filtrate until you have answered all of the questions below.
14. After your questions are answered, throw the filter paper in the garbage and clean up your area.
Analysis Questions: (show calculations for all questions where calculations are required)
1. You observed the reaction of calcium chloride dihydrate with sodium carbonate to form calcium
carbonate and sodium chloride. Write the balanced chemical equation for this reaction (the
“dihydrate” is not involved in the chemical reaction. On the products side of the equation, you can
simply represent it as two solitary water molecules).
Look at your balanced equation. Only one compound forms a solid. All other compounds are
aqueous. Reflect this in your equation by adding “(s)” or “(aq)” after each compound.
2. Which chemical is the limiting reagent in this lab?
3. Which reactant is there an excess of? What mass of the excess reagent would have been left
unreacted? Where is the excess reagent after the reaction and filtration?
4. How many grams of NaCl should be produced from today’s lab? Where is the NaCl after the reaction
and filtration?
5. How many grams of CaCO3 should be produced from today’s lab?
6. How many grams of CaCO3 would be produced if the other reagent was the limiting reagent?
7. Was the mass of CaCO3 correctly predicted? (In other words, was the actual mass close to the
calculated amount?)
Conclusion:
.
Discuss results compared to hypothesis. Calculate percent yield:
.
.
Actual yield
x 100 = % yield
Theoretical yield
16
Name:_______________________________
# of Items – Mole Conversion Assignment
1. How many molecules are in the quantities below?
a. 2.0 moles
b. 1.5 moles
c. 0.75 moles
d. 15 moles
e. 0.35 mole
2. How many moles are in the number of molecules below?
a. 6.02 x 1023
b. 1.204 x 1024
c. 1.5 x 1020
d. 3.4 x 1026
e. 7.5 a 1019
3. How many Magnesium Bicarbonate molecules are in 7.82 moles?
4. How many molecules of Barium Iodide would be in 3.782 moles?
5. If you have 2.785 x 106 molecules of Magnesium Bromate, how many moles would that be?
6. If you have 6 molecules of Sodium Carbonate, how many moles do you have?
7. If you have 6.65 moles of Aluminum Phosphate, how many molecules of Aluminum Phosphate do you
have?
8. How many moles of Calcium Carbonate are there if there are 2.839 x 1018 molecules?
17
Name:_________________________
Mole-Mole Calculations Assignment
1. Number of moles of sodium in 1 mole of sodium hydroxide.
2. Number of moles of oxygen atoms in 2 moles of sodium hydroxide.
3. Number of moles of oxygen atoms in 1 mole of Chromium (II) Oxalate.
4. If you know there are 4 moles of oxygen in a sample of phosphoric acid, how many moles of
phosphoric acid molecules are there?
5. If there are 6 moles of chlorine in a sample of carbon tetrachloride, how many moles of carbon
tetrachloride molecules are there?
6. How many moles of Carbon are there in 1 mole of CH3COOH
7. Number of Hydrogen atoms in 2 moles of Phosphorus acid.
8. Number of ammonium ions in 3 moles of ammonium borate.
9. How many iodine atoms are in 2 moles of phosphorus triiodide.
10. How many atoms of Hydrogen are in 3.837 moles of Hydrogen gas?
11. How many atoms of Sodium are in 7.82 moles of Sodium Hydroxide?
12. How many atoms of Iodine are in 3.84 moles of Aluminum Iodine?
18
Name:_______________________
Molar Mass Assignment
1. Determine the molar mass of the following elements:
a. Nitrogen:
b. Tin:
c. Gold:
2. Which element has a molar mass of:
a. 9.0122g/mol?
b. 50.9415g/mol?
c. 114.82g/mol?
3. Determine the molar mass of the following compounds:
a. CI4
b. LiNH4
c. K2C2O4
d. KMnO4
e. KCl
f. Na2SO4
g. Ca(NO3)2
h. Al2(SO4)3
i.
(NH4)3PO4
j.
CuSO4  5H2O (copper (II) sulphate pentahydrate: means copper (II) sulphate surrounded
by 5 water molecules)
k. Lead (II) Iodate
l.
Benzoic acid
m. Zinc Sulfite
n. Barium Iodide
o. Iodine gas
p. Magnesium Bromide
q. Copper (II) Iodide
r. Iron (III) Bisulphate
19
Name:__________________________
# of Item – Mass Conversion Assignment
1. Determine the number of moles in each of the quantities below:
a. 25g of NaCl
b. 125g of H2SO4
c. 100. g of KMnO4
d. 35g of CuSO4  5H2O
2. Determine the number of grams in each of the quantities below:
a. 2.5 moles of NaCl
b. 0.50 moles of H2SO4
c. 1.70 moles of KMnO4
d. 3.2 moles of CuSO4  5H2O
3. How many atoms of Aluminum are in 32.32g of aluminum?
4. What is the mass of 3.432 x 1034 molecules of carbon tetrafluoride?
5. How many atoms of magnesium carbonate are there in 6.73g of magnesium carbonate?
6. What is the mass of 7.29x1027 molecules of hydrochloric acid?
7. How many formula units of Mercury (II) Perchlorate are in 0.798Kg?
8. How many molecules of hydrochloric acid are in a 7.32g sample?
9. How many hydrogen atoms are in a 2.50g sample of hydrogen gas?
20
Name:______________________________
Mole-Volume Conversion Assignment
Assume all information given at Standard Temperature and Pressure (STP).
1. What volume will the following quantities of gases occupy at STP?
a. 1.00 mole of H2
b. 3.20 moles of O2
c. 0.750 moles of N2
d. 1.75 moles of CO2
e. 0.50 moles of NH3
f. 5.0g of H2
g. 100. g of O2
h. 28.0g of N2
i.
60. g of CO2
j.
10. g of NH3
2. What is the mass of a sample of carbon dioxide that contains 2.38 moles of carbon dioxide @STP?
3. How many moles of Neon gas are there in 74.2 Liters @STP?
4. How many moles of Carbon Monoxide are in 7.38 Liters @STP?
5. How many milliliters of Iodine gas would 3.782 moles take up at 0oC and 1 atm?
6. What is the mass of 32.8L of CO2 gas?
7. If you have 2.8g of chlorine gas, how many liters would it take up?
21
Name:________________________
Molarity Calculations
1. What is the molarity of a solution in which 58g of NaCl are dissolved in 1.0L of solution?
2. What is the molarity of a solution in which 10.0g of AgNO3 is dissolved in 500. mL of solution?
3. How many grams of KNO3 should be used to prepare 2.00 L of a 0.500 M solution?
4. To what volume should 5.0 g of KCl be diluted in order to prepare a 0.25 M solution?
5. How many grams of CuSO4  5H2O are needed to prepare 100. mL of a .10 M solution?
Name:___________________________________
Mixed Mole Problems
1. How many grams are there in 1.5 x 1024 molecules of CO2?
2. What volume would the CO2 in problem 1 occupy?
3. A sample of NH3 gas occupies 75.0 liters at STP. How many molecules is this?
4. What is the mass of the sample of the NH3 in problem 3?
5. How many atoms are there in 1.3 x 1022 molecules in NO2?
6. A 5.0g sample of O2 is in a container at STP. What volume is the container?
7. How many molecules of O2 are in the container in problem 6? How many atoms of oxygen?
22
Name:_______________________________
Percent Composition
Find the percent composition for the elements in each of the elements in the compounds below.
1. KMnO4
2. HCl
3. Mg(NO3)2
4. (NH4)3PO4
5. Al2(SO4)3
6. How many grams of oxygen can be produced from the decomposition of 100. g of KClO3?
7. How much iron can be recovered from 25.0 g of Fe2O3?
8. How much silver can be produced from 125g of Ag2S?
23
Name:________________________________
Determining Empirical Formula
1. 75% carbon, 25% hydrogen
2. 52.7% potassium, 47.3% chlorine
3. 22.1% aluminum, 25.4% phosphorus, 52.5% oxygen
4. 13% magnesium, 87% bromine
5. 32.4% sodium, 22.5% sulfur, 45.1% oxygen
6. 25.3% copper, 12.9% sulfur, 25.7% oxygen, 36.1% water
24
Name:_________________________________
Stoichiometry: Mole- Mole Problems
1. ____N2 + ____H2  ____NH3
How many moles of hydrogen are needed to completely react with two moles of nitrogen?
2. ____ KClO3  ____ KCl + ____ O2
How many moles of oxygen are produced by the decomposition of six moles of potassium chlorate?
3. ____ Zn + ____ HCl  ____ ZnCl2 + ____ H2
How many moles of hydrogen are produced from the reaction of three moles of zinc with an excess
of hydrochloric acid?
4. ____ C3H8 + ____ O2  ____ CO2 + ____ H2O
How many moles of oxygen are necessary to react completely with four moles of propane?
5. ____ K3PO4 + ____ Al(NO3)3  ____ KNO3 + ____ AlPO4
How many moles of potassium nitrate are produced when two moles of potassium phosphate react
with two moles of aluminum nitrate?
25
Name:_________________________
Volume-Volume Stoichiometry
1. ____N2 + ____ H2  ____ NH3
What volume of hydrogen is necessary to react with five liters of nitrogen to produce ammonia?
(Assume constant temperature and pressure).
2. What volume of ammonia is produced in the reaction in Problem 1?
3. ____ C3H8 + ____ O2  ____ CO2 + ____ H2O
If 20 liters of oxygen are consumed in the above reaction, how many liters of carbon dioxide are
produced?
4. ____ H2O  ____H2 + ____O2
If 30mL of hydrogen are produced in the above reaction, how many milliliters of oxygen are
produced?
5. ____ CO + ____O2  ____ CO2
How many liters of carbon dioxide are produced if 75 liters of carbon monoxide are burned in
oxygen? How many liters of oxygen are necessary?
26
Name:_________________________
Mass-Mass Stoichiometry
1. ____ KClO3  ____ KCl + ____O2
How many grams of potassium chloride are produced if 25 g of potassium chlorate decompose?
2. ____ N2 + ____H2  ____ NH3
How many grams of hydrogen are necessary to react completely with 50.0g of nitrogen in the above
reaction?
3. How many grams of ammonia are produced in the reaction in Problem 2?
4. ____ AgNO3 + ____ BaCl2  ____ AgCl + ____ Ba(NO3)2
How many grams of silver chloride are produced from 5.0 g of silver nitrate reacting with an excess
of barium chloride?
5. How much barium chloride is necessary to react with the silver nitrate in Problem 4?
27
Name:______________________________
Mixed Stoichiometry Problems
1. ____ N2 + ____H2  ____ NH3
What volume of NH3 at STP is produced if 25.0 g of N2 is reacted with excess of H2?
2. ____ KClO3  ____ KCl + ____O2
If 5.0g of KClO3 is decomposed, what volume of O2 is produced at STP?
3. How many grams of KCl are produced in Problem 2?
4. ____ Zn + ____ HCl  ____ ZnCl2 + ____ H2
What volume of hydrogen at STP is produced when 2.5 g of zinc react with an excess of
hydrochloric acid?
5. ____ H2SO4 + ____ NaOH  ____ H2O + ____ Na2SO4
How many molecules of water are produced if 2.0 g of sodium sulfate are produced in the above
reaction?
6. ____ AlCl3  ____ Al + ____ Cl2
If 10.0 g of aluminum chloride are decomposed, how many molecules of Cl2 are produced?
28
Name: __________________________
Assignment: Limiting reagent given the masses of reactants
1. For the reaction:
Zn +
HCl 
H2 +
Zn2Cl
Determine the limiting reagent in the following situations:
a. 1.0g Zn combines with 0.75g of HCl
b. 5.0g Zn combines with 4.5g HCl
2. For the reaction:
FeCl3 + AgNO3  AgCl +
Fe(NO3)3
Determine the limiting reagent in the following situations:
a. 1.0 g of FeCl3 and 1.0g AgNO3
b. 5.0 g of FeCl3 and 4.0g AgNO3
3. For the reaction:
KOH +
H3PO4 
K3PO4 +
H2O
Determine the limiting reagent in the following situations:
a. 1.5 g of H3PO4 and 5.0g KOH
b. 0.87 g of H3PO4 and 1.3g KOH
29
Name:________________________
Limiting Reagent Assignment
1. ____ N2 + ____H2  ____ NH3
How many grams of NH3 can be produced from the reaction of 28g of N2 and 25g of H2?
2. How much of the excess reagent in Problem 1 is left over?
3. _____ Mg + _____HCl  _____ MgCl2 + _____H2
What volume of hydrogen at STP is produced from the reaction of 50.0 g of Mg and the
equivalent of 75 g HCl?
4. How much of the excess reagent in Problem 3 is left over?
5. _____ AgNO3 + _____ Na3PO4  _____ Ag3PO4 + _____ NaNO3
Silver nitrate and sodium phosphate are reacted in equal amounts of 200. G each. How many
grams of silver phosphate are produced?
6. How much of the excess reagent in Problem 5 is left?
30