Download Unit 7 Homework and Lab Packet

Gen Chem
Unit 7 Stoichiometry
Name ___________________________________ Per: ______________________
Unit 7 Homework and Lab Packet
Stoichiometry (Ch. 12)
Mole Ratio
Practice I
Mass-Volume
Practice II
Percent Yield
Practice
Mole Ratio
Practice II
Volume-Volume
Practice
Limiting Reactant
Practice I
Gram-Gram
Practice I
Stoichiometry
Practice 1
Limiting Reactant
Practice II
Gram-Gram
Practice II
Stoichiometry
Practice II
Stoichiometry
Labs
Mass-Volume
Practice I
Stoichiometry
Practice III
Unit 7 Homework grade: _____________________
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Gen Chem
Unit 7 Stoichiometry
Unit 7 Overall goals: To understand basic stoichiometry calculations and how they’re applied in the
laboratory and your everyday life
After completing this unit, you should understand and be able to explain the following:
 Use the unit analysis method to answer stoichiometry problems
 To state the value for the molar volume of any gas at STP: 22.4 L/mol
 To relate the volume of a gas at STP to its mass and number of particles
 To classify the three basic types of stoichiometry problems: mass-mass, mass-volume, and
volume-volume
 To perform mass-volume stoichiometry calculations
 To perform volume-volume stoichiometry calculations
 To identify the limiting reactant in a chemical reaction
 To calculate the percent yield for a reaction given the actual yield and theoretical yield
Questions to ponder:
o Does the amount of each reactant matter in a chemical reaction?
o What happens if there are reactants leftover?
o What would happen if too much air bag chemical were put into an airbag?
o What would happen if not enough air bag chemical were put into an airbag?
Iowa Core Essential Concepts and Skills
-Design and conduct scientific investigations
-Use technology and mathematics to improve investigations and communications.
-Formulate and revise scientific explanations and models using logic and evidence.
-Understand and apply knowledge of chemical reactions
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Gen Chem
Unit 7 Stoichiometry
Mole Ratio Practice 1
Directions: Using the balanced chemical equations, use unit-analysis to complete the problems. Show
your work.
2 H2S + 3 O2  2 H2O + 2 SO2
1. What mole ratios can be made from this reaction?
Hint: Here’s one: 2 mol H2S : 3 mol O2
2. If 5.0 moles of H2S is consumed during the reaction, how many moles of O2 would be needed?
3. How many moles of SO2 would be produced by the complete reaction of 0.25 moles of H2S?
4. During the formation of 1.5 moles of H2O, how many moles of SO2 would be produced?
2 Al + 6 HCl  2 AlCl3 + H2
5. According to the coefficients in the balanced equation _____ moles of H2 can be produced from 2
moles of Al.
6. The complete reaction of 2.5 moles of Al would require _____ moles of HCl.
7. How many moles of AlCl3 would be made during the production of 10.0 moles of H2?
8. How many moles of HCl must be consumed in order to produce 0.10 moles of H2?
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Gen Chem
Unit 7 Stoichiometry
Mole Ratio Practice II
1. Given the balance equation, calculate the moles of oxygen gas that react with 0.500 mol of
hydrogen gas.
2 H2 + O2  2 H2O
2. How many moles of water are produced from 0.500 mol of hydrogen gas? Use the balanced
equation in #1.
3. How many moles of propane gas, C3H8, react with 1.75 mol of oxygen gas? Balance the equation
before proceeding.
C3H8 + O2  CO2 + H2O
4. How many moles of water are needed to react completely with 12 moles of sodium according to the
balanced equation?
2 Na + 2 H2O  2 NaOH + H2
5. How many grams of lead(II)iodide are produced from 6.0 mol NaI according to the balanced
equation given?
Pb(NO3)2 + 2 NaI  2 NaNO3 + PbI2
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Gen Chem
Unit 7 Stoichiometry
Gram – Gram Practice 1
Using the balanced equation, answer the following questions. Show your work.
2 Al + 3 FeCO3  Al2(CO3)3 + 3 Fe
1. How many moles of Fe will be formed if 200 moles of FeCO3 react with aluminum?
200 mol FeCO3
1
=
mol Fe
2. How many moles of Fe are formed when 50 grams of Al react with iron (II) carbonate?
50 g Al
1
=
mol Fe
A) Convert 50g Al to moles by dividing by the molar mass of Al
B) Take the answer from A and use the mole ratio to get moles of Fe
3. How many moles of Al react to completion with 600 grams of FeCO3?
600 g FeCO3
1
=
mol Al
A) Convert 600 g FeCO3 to moles by dividing by molar mass of FeCO3
B) Take the answer from A and use the mole ratio to get moles of Al
Test Yourself. Use the same basic steps from #1 and 2 to answer:
4. How many moles of Al2(CO3)3 are formed when 300 g of Al react with iron (II) carbonate?
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Gen Chem
Unit 7 Stoichiometry
5. How many grams of Fe are formed when 200 g of FeCO3 react?
200 g FeCO3
1
=
g Fe
A) Convert 200g FeCO3 to moles by _____________ by the molar mass of FeCO3
B) Convert the answer in A to moles of Fe by using the mole ratio
C) Using your answer in B, multiply by the molar mass of Fe to get grams of Fe
6. How many grams of Al2(CO3)3 are formed when 300 g of Al react with iron (II) carbonate?
A) Convert 300g of Al to moles by __________ the molar mass of Al
B) Take the answer to A and convert it to moles of __________ by using the _____________
C) Using your answer to B, convert to grams of Al2(CO3)3 by _____________ molar mass of
Al2(CO3)3
Test yourself!
7. How many grams of FeCO3 reacted to form 10 g Al2(CO3)3?
8. How many grams of Fe are formed if 10 g Al2(CO3)3 are formed?
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Gen Chem
Unit 7 Stoichiometry
Gram – Gram Practice II
1. Use the equation given to determine how many grams of aluminum chloride will be produced from
925 g of Cl2.
2 AlBr3 + 3 Cl2  3 Br2 + 2 AlCl3
2. Use the equation to determine what mass of FeS must react to form 326 g of FeCl2.
FeS + 2 HCl  H2S + FeCl2
3. How many grams of magnesium oxide are needed to produce 264 g of magnesium hydroxide
according to the given equation?
MgO + H2O  Mg(OH)2
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Gen Chem
Unit 7 Stoichiometry
4. Given the balanced equation, calculate the mass of product that can be prepared from 2.36 g of zinc
metal.
2 Zn + O2  2 ZnO
5. How many grams of oxygen gas must react to give 1.28 g ZnO? Use the equation in #4.
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Gen Chem
Unit 7 Stoichiometry
Mass-Volume Practice I (remember that 1000ml = 1 L) Extra Credit if you answer in mL or begin in mL
1. Given the balanced equation, how many liters of carbon dioxide gas at STP are produced from the
decomposition of 1.59 g of ferric carbonate?
Fe2(CO3)3

Fe2O3
+
3CO2
2. Given the balanced equation, how many liters of oxygen gas at STP are released from the
decomposition of 2.57 g of calcium chlorate (Ca(ClO3)2)?
Ca(ClO3)2

CaCl2
+ 3O2
3. Given the balanced equation, how many liters of carbon dioxide gas at STP are released from the
decomposition of 1.59 g of lithium hydrogen carbonate?
2 LiHCO3

Li2CO3
+
H2O
+
CO2
4. Given the balanced equation, how many liters of oxygen gas at STP are released from the
decomposition of 2.50 g mercuric oxide?

2 HgO
2 Hg
+
O2
5. Given the balanced equation, what mass of magnesium metal reacts with sulfuric acid to yield
0.225 L of hydrogen gas at STP?
Mg
+
H2SO4 
MgSO4
H2
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Gen Chem
Unit 7 Stoichiometry
6. Given the balanced equation, how many grams of hydrogen peroxide must decompose to give
5.0 L of oxygen gas at STP?
2 H2O2

2 H2O
+
O2
7. What mass of sodium metal reacts with water to give 0.750 L of hydrogen gas at STP?
Balance the equation before proceeding.
Na
H2O 
+
NaOH
+
H2
8. How many grams of manganese(II) chloride must react with sulfuric acid to release 5.0 L of
hydrogen chloride gas at STP? Balance the equation before proceeding.
MnCl2
+
H2SO4

MnSO4
+
2HCl
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Gen Chem
Unit 7 Stoichiometry
Mass-Volume Practice II
(Assume liters unless otherwise noted)
1. Hydrogen peroxide is always decomposing. How many grams of hydrogen peroxide must
decompose to give 5.0 L of oxygen gas at STP?
2 H2O2  2 H2O + O2
2. Candles are made of paraffin wax (C25H52). When paraffin burns in oxygen, carbon dioxide and
water are produced. If 5.5 g of paraffin wax burn, what volume of carbon dioxide is produced at STP?
C25H52 + 38 O2  25 CO2 + 26 H2O
3. The body metabolizes sucrose, C12H22O11, by burning it with oxygen to produce carbon dioxide,
water, and energy (ATP). If 12 grams of sucrose are burned according to the following equation, what
volume (L)of carbon dioxide is produced at STP? Extra credit if answered in mL.
C12H22O11 + 12 O2  12 CO2 + 11 H2O
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Gen Chem
Unit 7 Stoichiometry
4. Solid lithium hydroxide (LiOH) is used in spacecraft to remove exhaled carbon dioxide from the air.
The lithium hydroxide reacts with carbon dioxide to form lithium carbonate (Li2CO3) and water. How
many liters of carbon dioxide can be absorbed by 74 g of LiOH?
2 LiOH + CO2  Li2CO3 + H2O
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Gen Chem
Unit 7 Stoichiometry
Volume-Volume Practice
1. Assuming all gas volumes are measured at STP, how many liters of oxygen gas react with 2.00 L of
carbon monoxide? Balance the equation before proceeding.
CO
+ O2

CO2
2. How many liters of carbon dioxide are produced from 2.00 L of CO? (use the equation in #1)
3. Assuming all gas volumes are measured at STP, how many liters of iodine vapor react with
12.5 L of hydrogen gas? Balance the equation.
H2
+
I2 
HI
4. How many liters of hydrogen iodide are produced from 1.25 L of H2? (use the equation in #3)
5. Assuming all gas volumes are measured at STP, how many liters of nitrogen gas react to give 4.50
L of ammonia gas? Balance the equation.
H2
+
N2 
NH3
6. How many liters of hydrogen gas must react to give 5.50 L of NH3? (use the equation in #5)
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Gen Chem
Unit 7 Stoichiometry
7. Assuming all gas volumes are measured at STP, how many liters of chlorine gas react to yield 1.75 L
of dichlorine trioxide?
Cl2
+
O2

Cl2O3
8. How many liters of oxygen gas must react to give 1.75 L of Cl2O3? (use the equation in #7)
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Gen Chem
Unit 7 Stoichiometry
Stoichiometry Practice I
1. If 45 g of zinc reacts with hydrochloric acid, how many grams of zinc chloride are produced?
Zn + 2 HCl  ZnCl2 + H2
2. How many moles of chlorine gas must react with an excess of sodium iodide if you wish to produce
10.0 g of sodium chloride?
Cl2 + 2 NaI  2 NaCl + I2
3. How many grams of free silver would be produced by reacting 3.5 moles of free copper with an
excess of silver nitrate to produce copper (II) nitrate?
Cu + 2 AgNO3  2 Ag + Cu(NO3)2
4. How much silver phosphate is produced if 20.0 grams of silver acetate is reacted with an excess of
sodium phosphate?
3 AgC2H3O2 + Na3PO4  Ag3PO4 + 3 NaC2H3O2
5. How many liters of sulfuric acid are required to react with 2.00 g sodium hydroxide?
H2SO4 + 2 NaOH  Na2SO4 + 2 H2O
6. How many moles of oxygen would be needed to react completely with 0.50 moles of hydrogen to
produce water?
O2 + 2 H2  2 H2O
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Gen Chem
Unit 7 Stoichiometry
Stoichiometry Practice II
1. 3.5 g KClO3 are heated over a Bunsen burner. If complete decomposition takes place, how many
moles of oxygen will be produced?
2 KClO3  3 O2 + 2 KCl
2. If 23 grams of NaCl reacts with an excess of silver nitrate, how many grams of silver chloride are
produced?
NaCl + AgNO3  AgCl +
NaNO3
3. How many moles of Cu2S can be produced from 9.9 grams of copper (I) chloride reacting with an
excess of H2S?
2 CuCl + H2S  Cu2S + 2 HCl
4. 3.00 moles of hydrogen was produced when zinc was added to sulfuric acid. How many grams of
zinc were used in order to produce this amount of gas?
Zn + H2SO4  H2 + ZnSO4
5. How many grams of NaAlO2 can be obtained from 1.46 grams of AlCl3?
AlCl3 + 4 NaOH  3 NaCl + 2 H2O + NaAlO2
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Gen Chem
Unit 7 Stoichiometry
Stoichiometry Practice III
1. 4.00 grams of copper (II) sulfate reacted with an iron nail. How much copper can be obtained from
this reaction?
CuSO4 + Fe  FeSO4 +
Cu
2. How much copper can be obtained from a reaction of 5.48 grams of iron metal with copper (II)
sulfate?
CuSO4 + Fe  FeSO4 + Cu
3. What mass of sodium nitrate will be obtained from the reaction of 4.97 grams of lead (II) nitrate with
sodium iodide?
Pb(NO3)2 + 2 NaI  PbI2 + 2 NaNO3
4. What mass of lead (II) iodide will be obtained from the reaction of 4.50 g sodium iodide with lead (II)
nitrate?
Pb(NO3)2 + 2 NaI  PbI2 + 2 NaNO3
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Gen Chem
Unit 7 Stoichiometry
Standardized Test Prep
Test-Taking Tip
Anticipate the answer. Use what you know to predict what you think the answer
should be. Then look to see if your answer, or one much like it, is given as an option.
Select the choice that best answers each question
1. Nitric acid is formed by the reaction of nitrogen dioxide and water.
3 NO2 + H2O  NO + 2 HNO3
How many moles of water are needed to react with 8.4 mol NO2?
a. 2.8 mol
b. 3.0 mol
c. 8.4 mol
d. 25 mol
2. Phosphorus trifluoride is formed from its elements.
P4 + 6 F2  4 PF3
How many grams of fluorine are needed to react with 6.20 g of phosphorus?
a. 2.85 g
b. 5.70 g
c. 11.4 g
d. 37.2 g
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Gen Chem
Unit 7 Stoichiometry
Percent Yield Practice
Percent yield = (actual yield/theoretical yield) x 100
1. Zack the chemistry student prepares acetone by decomposing 31.6 g of calcium acetate. If Zack
collected 10.4 g of acetone and the theoretical yield is 11.6 g, what is the percent yield?
2. Samuel the chemistry student prepares lead(II) iodide from 1.55 g of lead (II) nitrate and potassium
iodide. If Samuel collected 2.01 g of PbI2, and the theoretical yield is 11.6 g, what is the percent yield?
3. Kelly took A 1.50 g sample of sodium nitrate and decomposed it by heating. If the resulting sodium
nitrite has a mass of 1.29 g and the theoretical yield is 1.22g, what is the percent yield?
4. Jessie decomposed 1.000 g of potassium bicarbonate by heating. If the resulting potassium
carbonate has a mass of 0.725 g and the theoretical yield is 0.716 g, what is the percent yield?
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Gen Chem
Unit 7 Stoichiometry
Limiting Reactant Practice 1
1. If 1.00 mol of nitrogen gas and 1.50 mol of oxygen gas react, what is the limiting reactant and how
many moles of NO are produced from the reaction?
N2
+
O2

2 NO
2. If 1.00 mol of nitrogen gas and 0.500 mol of oxygen gas react, what is the limiting reactant and how
many moles of NO are produced? (use the equation in #1)
3. If 1.00 mol of nitrogen monoxide gas and 1.00 mol of oxygen gas react, what is the limiting reactant
and how many moles of NO2 are produced from the reaction?
2 NO
+
O2

2 NO2
4. If 1.00 mol of nitrogen monoxide gas and 0.250 mol of oxygen gas react, what is the limiting reactant
and how many moles of NO2 are produced from the reaction? (use the equation in #3)
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Gen Chem
Unit 7 Stoichiometry
Limiting Reactant Practice 2
1. The principle ingredient of the self-defense agent mace is chloroacetophenone, C6H5COCH2Cl. This
compound is prepared by chlorination of acetophenone, C6H5COCH3. What mass of
chloroacetophenone would be produced by the reaction of 4.00 g of acetophenone with 10.4 g of
chlorine?
C6H5COCH3 + Cl2  C6H5COCH2Cl + HCl
2. Which reactant in #1 is in excess and by how many grams is it in excess?
3. Some household drain cleaners consist of a mixture of granules of sodium hydroxide and powdered
aluminum. When dissolved in water, the sodium hydroxide reacts with the aluminum to produce
hydrogen gas:
2 Al + 2 NaOH + 6 H2O  2 NaAl(OH)4 + 3 H2
The sodium hydroxide helps dissolve grease, and the hydrogen gas provides mixing and scrubbing
action. What mass of hydrogen gas would be formed from a mixture of 2.48 g Al and 4.76 g NaOH?
Which reactant is in excess and by how many grams is it in excess?
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Gen Chem
Unit 7 Stoichiometry
4. Sugar cane produces sucrose (C12H22O11) from carbon dioxide and water through photosynthesis:
12 CO2 + 11 H2O  C12H22O11 + 12 O2
If the reaction consumes 30.0 g of carbon dioxide and produces 6.84 g of sucrose, what is the percent
yield of sucrose? Percent yield = (actual/theoretical) x 100
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Gen Chem
Unit 7 Stoichiometry
How Many Molecules of Water are in Lake Erie?
Problem
How do you determine the number of molecules of water in Lake Erie?
Materials
Water balance
eyedropper
beaker
graduated cylinder
Procedure
Using the equipment listed above and the following helpful information, determine the number of
molecules of water in Lake Erie.
-
One mole of anything contains 6.02 x 1023 units
One mole of water has a mass of 18.0 grams
Lake Michigan has an average depth of 85 meters
Lake Michigan covers 58,000 km2 (5.8x104 km2)
1 mL = 1 cm3
1. How many drops of water are in 2 liters?
2. What value can be calculated about Lake Erie with the data from Step 1?
3. Calculate the volume of water in Lake Erie.
4. How many drops of water are in Lake Erie?
5. Calculate how many molecules of water are in Lake Erie.
6. Calculate how many moles of water are in Lake Erie.
7. The recommended daily allowance of water is about two liters. How many molecules of
water should you drink each day?
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Gen Chem
Unit 7 Stoichiometry
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Gen Chem
Unit 7 Stoichiometry
LAB
Moles of Iron and Copper
Introduction
The mole is a convenient unit for analyzing chemical reactions. The mole is equal to 6.02x1023
particles, or Avogadro’s number of particles. More importantly, however, the mass of a mole of any
compound or element is the mass in grams that corresponds to the molecular formula, or atomic mass.
Simply stated, the atomic mass of copper is 63.5 amu, which means that the mass of one mole of
copper atoms is equal to 63.5 grams. Likewise, the molecular mass of water is equal to 18.09 amu,
and the mass of one mole of water molecules is equal to 18.0 grams.
The mole is the common language in chemical reactions. In this experiment, you will observe
the reaction of iron nails with a solution of copper (II) sulfate and determine the number of moles
involved in the reaction.
Prelab
In your lab notebook, before your summary of the procedure, answer the following questions. Be sure
to include the questions.
1. How many moles are present in a sample of 34.0 g of iron metal?
2. How many atoms of iron are present in 2.00 mol of iron?
3. What is the formula and molar mass of copper(II)sulfate?
4. Define decant: (read the procedure and deduce the definition)
5. Why is the washing of the copper necessary in this experiment?
6. Write a balanced equation for this lab. (assume iron has a charge of +2)
25
Gen Chem
Unit 7 Stoichiometry
Procedure
1. Find the mass of a clean, empty, dry 100 ml beaker. Record the mass to the nearest 0.01 g.
2. Add approximately 4.00 g copper(II)sulfate crystals to the beaker. Find the mass and record it in the
data table.
3. Add 50.0 ml of distilled water to the beaker. Swirl the beaker around to dissolve all of the
copper(II)sulfate crystals.
4. Obtain two clean, dry iron nails from your teacher. If the nails are not clean, use a piece of sand
paper or steel wool to make the surface of the nail shiny. Find the mass of the nails and record.
5. Place the nails in to the copper(II)sulfate solution. Leave them undisturbed for 20 minutes. During
that time, you should see the formation of copper in the beaker. At the same time, some of the iron will
be used up.
6. Use the tongs to carefully pick up the nails, one at a time. Use distilled water in a wash bottle to
rinse off any remaining copper from the nails before removing them completely from the beaker. If
necessary, use a stirring rod to scrape any excess copper from the nails. Set the nails aside on a
paper towel to dry.
7. After the nails are completely dry find the mass of the nails and record.
8. Decant means to pour off only the liquid from a container that is holding both a solid and liquid.
Carefully decant the liquid from the solid. Observe your teacher for proper technique. Pour the liquid
into another beaker so that in case you over pour, you can still recover the solid.
9. After decanting, rinse the solid again with about 25 ml of distilled water. Decant again. Repeat this
step three or four more times.
10. Next, wash the solid with about 25 ml of 1M hydrochloric acid. Decant again. Then, once more,
clean the solid with 25 ml of distilled water.
11. After the final washing with water, place the copper in a drying oven to dry.
12. Allow the copper to become completely dry, then find the mass of the beaker plus the copper and
record.
13. Clean up and put away all materials. Wash hands thoroughly.
Data Table
Before the reaction
Mass of empty, dry beaker (g)
Mass of beaker + copper(II)sulfate (g)
Mass of two iron nails (g)
After Reaction
Mass of two iron nails (g)
Mass of beaker + copper (dry) (g)
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Gen Chem
Unit 7 Stoichiometry
Postlab Discussion
In this experiment, you have reacted some of the iron from the nail at the same time as you
have produced some copper. In order to find the moles of each of these substances, you will need to
divide the mass of the iron used or the copper produced by the molar mass. You can also determine
the moles of copper(II)sulfate that you started with by dividing the mass of the copper(II)sulfate by the
molar mass.
Analysis and Conclusion
1a) How many grams of copper(II)sulfate (CuSO4) did you dissolve in water? Just give the mass.
b) How many moles of copper(II)sulfate is this?
2. How many grams of iron reacted with the copper(II)sulfate? (Hint: You started with full nails and you
had some left over.)
3. From the mass of copper(II)sulfate used, predict the mass of copper produced in this lab.
(This is called your theoretical yield.)
3B. Repeat #3, but start with your total mass of nails instead of copper (II) sulfate.
3C. How do #3 and #3B compare? How should they compare? Why?
4. Compare your theoretical yield to your actual yield of copper by doing a percent yield.
Actual yield
x 100 = % yield
Theoretical Yield
5. If you did not get 100%, explain why. What error or mistakes could have affected your results?
27
Gen Chem
Unit 7 Stoichiometry
6. Was there any evidence that some of the copper(II)sulfate was left in the beaker? Explain.
7. What is the limiting reactant? Fe or CuSO4? Support your answer.
8. Which reactant was in excess? Support your answer.
28
Gen Chem
Unit 7 Stoichiometry
LAB
The School Bus Reaction
Introduction
In this lab, you will be observing the reaction between sodium iodide and lead(II)nitrate. The
question to be answered: How does the amount of product obtained in a chemical reaction compare
with the predicted amount of product?
Procedure
1. In this lab you will use 0.0075 moles lead(II)nitrate, Pb(NO3)2, and 0.013 moles of potassium iodide,
NaI. Calculate the mass of each. Show your calculations below.
2. Label two clean, dry 100 ml beakers- one “A” and one “B”. Mass each.
3. Add 30 ml of lead(II) nitrate solution to beaker A and 30 ml of potassium iodide to beaker B.
4. Slowly add the sodium iodide solution to the lead(II)nitrate solution while stirring (B  A). Wash the
potassium iodide beaker several times with a small amount of distilled water from a wash bottle,
pouring the wash into the mixture.
5. Using a hot plate warm the mixture for five minutes. Do not boil! Let the precipitate settle. If not
settled after five minutes, continue to warm until the precipitate settles.
6. Record the mass of a piece of filter paper. Filter the mixture, collecting the filtrate in beaker B.
Decant the liquid from the mixture prior to pouring in the precipitate. Record the color of the precipitate
(solid left in the filter paper), and the color of the filtered liquid, the filtrate.
7. When filtering is complete, rest the filter paper with precipitate on beaker A. Dry the precipitate and
evaporate the filtrate in a drying oven overnight.
8. When both beakers are dry, measure their masses carefully. Clean the beakers.
Data Table
Mass of 0.0075 mol Pb(NO3)2
Mass of 0.013 mol KI
Color NaI
Color Pb(NO3)2
Mass beaker A (g)
Mass beaker B (g)
Mass of filter paper
AFTER DRYING
Color of precipitate
Color of filtrate
Mass of Beaker A, precipitate, filter
paper
Mass of beaker B, crystals
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Gen Chem
Unit 7 Stoichiometry
Analysis and Conclusion
1. Write a balanced chemical equation for the double replacement reaction occurring in this
experiment.
2. Calculate the following. Show your work. Be sure to include units on all of your answers.
a) Mass of the precipitate recovered (solid left in the filter paper).
b) Mass of the solute in the filtrate after drying (solid left in the beaker).
c) Theoretical mass of the precipitate (PbI2)
d) Theoretical mass of the solute in the filtrate (NaNO3)
d) Percent yield for the precipitate.
Actual yield/ theoretical yield= ____x 100%
2a/2c= ___x100%
e) Percent yield for the solute in the filtrate.
Actual yield/theoretical yield= ___x100%
2c/2d = ___ x100%
3. Add the masses of your two reactants: ____________
Add the masses of your two products: ____________
How do these two sums compare?
How should they compare and why?
30
Gen Chem
Unit 7 Stoichiometry
Stoichiometry of a Volcano
In class, you’ve learned to compute how much of a chemical product you can make when you mix
measured amounts of chemical reactants. In this lab, you will be actually using this information to
predict how much product will be made; you will then calculate the percent yield gained from the
amount that you actually recover.
The reaction you will be working with should be familiar to you from elementary school science fair
volcanoes: You will be mixing baking soda (NaHCO3) with vinegar (CH3COOH) to generate carbonic
acid (H2CO3), which breaks up into water and carbon dioxide gas and sodium acetate (NaCH3COO),
which is a food preservation additive.
Prelab:
1)
If we mix 25 grams of sodium bromide with a large amount of potassium chloride, what will our
theoretical yield of sodium chloride be?
NaBr + KCl  NaCl + KBr
2)
If our actual yield from this reaction was 18 grams of sodium chloride, what would our percent
yield for this reaction be?
Actual yield
x 100%
Theoretical yield
3)
Is the answer in question 2 reasonable? If so, explain why you think this was a reasonable
answer. If not, explain what is wrong with it and discuss possible reasons you might get this
answer in the laboratory.
4)
What are some factors that might cause our percent yield to be greater than 100%? What are
some factors that might cause it to be less than 100%? Make sure you discuss specific cases of
how both might happen.
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Gen Chem
Unit 7 Stoichiometry
Section 1: Computing the amounts of reactants that we need
In this lab, you will need to do a reaction where baking soda will react with an excess of vinegar. By
doing this, you will (hopefully!) ensure that you will get 100% actual yield for the reaction.
For our reaction, we will need to use 0.05 moles of baking soda, which we will call by its chemical
name, sodium hydrogen carbonate, for the rest of this lab. If we use much more than 0.05 moles of
baking soda, the reaction will be too large and we will risk having some of the reaction products pour
over the side of the flask when we mix it with the vinegar (which we will call acetic acid).
In the space below, calculate how much sodium hydrogen carbonate we will need for this lab:
For this lab, we will use _________ grams of sodium hydrogen carbonate.
Section 2: The reaction
In this section of the lab, we will be doing the reaction discussed in the prelab.
1)
Measure out the weight of sodium hydrogen carbonate that you calculated you would need in the
first section of this lab. Make sure that the amount you use is as close as you can make it to this
amount. Write the exact amount of sodium hydrogen carbonate that you used here:
Amount of sodium hydrogen carbonate used: ______________
2)
Dissolve the sodium hydrogen carbonate in about thirty milliliters of water. Stir the solution
until most or all of it is dissolved (if a little won’t dissolve, that’s OK).
3)
Weigh a 500 mL flask. You will need the weight of the empty flask at the end of the lab.
Weight of the empty 500 mL flask: ______________
4)
Add the sodium hydrogen carbonate solution to the preweighed 500 mL flask.
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Gen Chem
Unit 7 Stoichiometry
5)
Obtain 150 mL of acetic acid (a.k.a. vinegar) and slowly add it to the sodium hydrogen carbonate
solution. You will observe the formation of bubbles when the acetic acid is added to the sodium
hydrogen carbonate solution. Wait until the bubbling subsides before adding more acetic acid.
When all of the acetic acid has been added, stir for two minutes before moving on to step 6.
6)
When the solution is again calm (there may be a few bubbles rising from the bottom of the flask
- this is normal), move the flask to a hot plate and heat it to boiling. Be careful that the flask
does not boil over because this will cause errors in your calculations - a good way to prevent this
is to add a boiling stick. Once the flask has started boiling, gently set a watch class on its mouth
to keep any of the liquid inside from splattering.
7)
When all of the liquid in the solution has boiled away, remove the flask from the hot plate. The
powder that you observe inside is the product of the reaction, sodium acetate. Once the flask has
had a few minutes to cool down to room temperature, measure and record its weight.
Weight of the flask, after the reaction: _______________
8)
When this is done, you can rinse out the flask and any other glassware you used. All waste can
go down the sink.
Questions:
1)
Write down the equation of the reaction that we did in this lab. (Hint: The reaction is discussed
in the introduction to this lab.) Is it balanced?
NaHCO3
+ HC2H3O2
 H2O
+ CO2
+ NaCH3COO
2)
Using the exact weight of sodium hydrogen carbonate that you measured in step 1 and the
equation that you wrote in the problem above, what is the theoretical yield (mass) of sodium
acetate in this reaction?
NaHCO3 + HC2H3O2  H2O + CO2 + NaCH3COO
3)
Calculate the actual yield of sodium acetate that you recovered in this lab, using the weight of the
empty flask and the weight of the flask after the reaction.
4)
Using the actual yield of sodium acetate that you measured in step 3 and the theoretical yield of
sodium acetate that you calculated in step 2, calculate the percent yield of sodium acetate
recovered in this lab:
Actual yield
x 100%
Theoretical yield
33
Gen Chem
Unit 7 Stoichiometry
5)
Was your percent yield of sodium acetate 100%? What factors do you think caused any error
that you found? Explain, using specific examples:
6)
Do you think it is common for chemists to get 100% yields for chemical reactions? Why or why
not?
7)
If you had to do this lab again, what would you do differently to improve your answers?
Explain, using specific examples
34