Download 4 Unit Packet - SRHSchem

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Unit 4: Holey Moley
Topics/ Daily Outline:
# A Day B Day
1
12.6
12.7
2
12.8
12.9
3 12.12 12.13
4 12.14 12.15
5 12.16 12.19
6 12.20 12.21
7 12.22
1.3
8
1.4
1.5
9
1.6
1.9
10 1.10
1.11
11 1.12
1.13
1.17
1.18
12
1.19
1.20
HW:
1.
2.
3.
4.
5.
6.
Content:
Mole Calculations
Mole Calculations
Empirical and Molecular Formula
Composition of Hydrates
Mole-mole Stoichiometry
Stoichiometry, Percent yield
Limiting Reactant
Limiting Reactant Calculations
Molarity, Precipitate Formation
Can You Make 1 Gram of a Precipitate?
Dimensional Analysis Exam
Quarterly Assessments
TEXT:
10.1, 10.2
10.1, 10.2
10.3
10.3
12.1, 12.2
12.2, 12.3
12.3
12.3
16.2
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CW #:
1
2
3
4
5
6
7
8
9
10
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HW #:
1
-2
-3
4
--
--
--
5
6
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Mole Calculations Practice
The Strange Case of Mole Airlines
Stoichiometry Practice I (Mole to Mole)
Stoichiometry Practice II (Other Units)
Stoichiometry Practice III (Limiting Reactants)
Review for Quarterly Assessment
http://www.srhschem.wikispaces.com
If you are absent, please use this sheet to determine what you missed and collect the materials
from the make-up work bins up front. Get help from a friend, the link above, or the instructor.
Drills
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CW 1: The Mole
The Elephant and the Methane Molecule
One Elephant has one trunk and four legs.
One methane molecule, CH4, contains one carbon atom and four hydrogen atoms.
1 dozen = 12 items
1 mole = 6.02X1023 items = Avogadro ’s Number
1. How many trunks are found in one dozen elephants?
2. How many legs are found in one dozen elephants?
3. How many carbon atoms are found in one dozen methane (CH4) molecules?
4. How many hydrogen atoms are found in one dozen methane molecules?
5. How many trunks are found in one mole of elephants?
6. How many carbon atoms are found in one mole of methane molecules?
7. How many hydrogen atoms are found in one mole of methane molecules?
8. How many legs are found in one mole of elephants?
9. How is “a mole” similar to “a dozen”?
10.A mole is equal to 6.02x1023 items, which is a very large number. Why would chemists
want to use moles as the unit to count atoms in?
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The Mole Road Map
11.Find the molar mass of the following compounds.
a. SO2
d. Phosphoric Acid
b. Iron (III) Chloride
e. Ammonium Sulfate
c. CCl2F2
f. Pb(NO3)2
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30
12.Use the mole road map to explain how to convert from 1.56x10 particles of sodium
chloride to grams of sodium chloride.
13.Set up the conversion factors and solve the problem in question 12.
14.Perform the following conversions.
a. How many oxygen molecules are in 3.36 L of oxygen gas at STP? (Answer:
9.03x1022 molecules)
b. Find the mass in grams of 2.00x1023 molecules of F2. (Answer:12.6 g)
c. Determine the volume in liters occupied by 14 g of nitrogen gas at STP. (Answer:
11.2 L)
d. Find the mass, in grams, of 1.00x1023 molecules of N2. (Answer: 4.65 g)
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CW 2: The Aluminum Mole Lab
Materials: Aluminum foil, electronic scale, metric ruler (0.01 cm precision).
Procedure:
1. Keeping the aluminum foil as flat as possible, measure the short length (from the
torn edge to torn edge) to the correct precision and record.
2. Ball up the foil and measure its mass to a precision of 0.01 grams and record.
3. Perform the calculations below.
4. Record the length you found in the class data table.
Data Table:
Measured Length of Foil (cm)
Mass of Foil (g)
Length of 1 Mole of Foil (cm)
Calculations:
Based on the mass of your sample of aluminum foil and the molar mass of aluminum,
determine how many atoms of aluminum are present in your sample.
Set up a proportion to determine the length of foil that would need to have been torn to have
1 mole of foil.
𝑌𝑜𝑢𝑟 𝑆𝑎𝑚𝑝𝑙𝑒 𝐿𝑒𝑛𝑔𝑡ℎ
𝑥
=
𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝐴𝑡𝑜𝑚𝑠 𝑖𝑛 𝑌𝑜𝑢𝑟 𝑆𝑎𝑚𝑝𝑙𝑒 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝐴𝑡𝑜𝑚𝑠 𝑖𝑛 𝑎 𝑀𝑜𝑙𝑒
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Class Data:
Group
Length of 1 Mol Al foil
Group
1
6
2
7
3
8
4
9
5
10
Length of 1 Mol Al foil
1. Calculate the average of the class data.
2. Determine your % error by comparing your data (experimental value) with the class
mean (accepted value).
(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)
% 𝐸𝑟𝑟𝑜𝑟 = |
| × 100%
𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
3. Consider your percent error. What are some possible sources of error? (NOTE:
“plugging in the numbers wrong” or miscalculations ARE NOT sources of error.)
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CW 3: Finding the Empirical Formula
Model 1: Percent Composition
The percent composition (by mass) of an element in a molecule is the mass of the element in
the molecule divided by the mass of the entire molecule times 100. Or, because the number of
atoms (molecules) is proportional to the number of moles of atoms (molecules),
% 𝐶𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛 𝑜𝑓 𝐸𝑙𝑒𝑚𝑒𝑛𝑡 𝑖 =
Table 1
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝐸𝑙𝑒𝑚𝑒𝑛𝑡 𝑖 𝑖𝑛 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑
× 100%
𝑀𝑜𝑙𝑎𝑟 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑡ℎ𝑒 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑
Percent composition (by mass) of some common organic molecules
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1. Verify that the % composition given for ethyne in Table 1 is correct.
Molar mass of C2H2:
Molar mass of carbon in C2H2:
% Composition Carbon:
Molar Mass of hydrogen in C2H2:
% Composition Hydrogen:
2. Fill in the missing molecular formulas and % compositions in Table 1.
3. Is it possible, given the original data in Table 1, to determine the % composition by
mass of H for 2-butene without using the equation given in the model? If so, how?
4. Based on the data in Table 1, is it possible to determine the molecular formula of a
compound solely from its percent composition? Why or why not?
5. What feature related to composition do all compounds with the same % composition
have?
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Model 2: Empirical Formula
The empirical formula of a compound describes the relative number of each type of atom in
the compound. It is given in terms of the smallest-possible-whole-number ratios (as
subscripts). For example, the empirical formula of ethane, C2H6, is CH3. (Note that the
subscript "1" is omitted.)
6. What feature related to the composition of a compound can be determined solely by
percent composition?
7. Determine the empirical formula of each of the molecules in Table 1.
8. A molecule containing only nitrogen and oxygen contains (by mass) 36.8% N.
a. How many grams of N would be found in a 100 g sample of the compound?
How many grams of O would be found in the same sample?
b. How many moles of N would be found in a 100 g sample of the compound?
How many moles of O would be found in the same sample?
c. What is the ratio of the number of moles of O to the number of moles of N?
d. What is the empirical formula of the compound?
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Practice Problems
9. The molecule 2-hexene has the molecular formula C6H12. Refer to Table 1 and
determine the percent composition of H in this molecule.
10.Determine the percent composition of carbon in acetic acid, HC2H3O2.
11.A compound used as a dry-cleaning fluid was analyzed and found to contain 18.00%
C, 2.27% H, and 79.73% Cl. Determine the empirical formula of the fluid.
12.A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g
H. What is the empirical formula of the compound?
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Model 3: Molecular Formula
The empirical formulas we have calculated in the preceding section express the simplest
atomic ratio between the elements in the compound. An empirical formula does not
necessarily represent the actual numbers of atoms present in a molecule of a compound; it
represents only the ratio between those numbers. The molecular formula of a compound may
be the empirical formula, or it may be a multiple of the empirical formula.
13.The empirical formula of a compound is NO2. Its molecular mass is 92 g/mol. What is
its molecular formula?
a. Determine the molar mass of NO2.
b. Divide the molecular mass by the molar mass. You should get a whole number.
c. Multiply the subscripts in NO2 by the whole number you got in the last step.
14.The empirical formula of a compound is CH2. Its molecular mass is 70 g/mol. What is
its molecular formula?
15.A compound is found to be 40.0% carbon, 6.7% hydrogen and 53.5% oxygen. Its
molecular mass is 60 g/mol.
a. What is the empirical formula?
b. What is the molecular formula?
16.The empirical formula of a compound is C3H7. Its molecular weight is 86.2 g/mol.
What is the molecular formula?
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CW 4: Composition of Hydrates
A hydrate is an ionic compound (salt) with water molecules loosely bonded to its crystal
structure. The water is in a specific ratio to each formula unit of the salt. For example, the
formula Na2S·9H2O indicates that there are nine water molecules for every one formula unit of
Na2S.
1. What percentage of water is found in Na2S·9H2O?
2. A 5.0 g sample of a hydrate of Cu(NO3)2 was heated to constant mass, leaving 3.9 g of
the anhydrous Cu(NO3)2 behind.
a. What percentage of water was in the hydrate?
b. Determine how many moles of Cu(NO3)2 are present.
c. Determine how many moles of H2O were present before heating.
d. Determine the formula of the hydrate (similar to finding the empirical formula,
divide each of the answers in B and C by the smallest number of moles to get
whole numbers).
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3. A group of students wanted to determine the formula of a hydrate of MgCO3. They
measured the mass of an empty crucible, then added some of the hydrate. They
heated the sample until the mass of the sample remained constant.
Mass empty crucible
50.43 g
Mass of crucible with sample
66.10 g
Mass of crucible with sample after heating 58.01 g
a. How many grams of the hydrate did they add to the crucible?
b. How many grams of water did they remove from the hydrate?
c. What is the formula of the hydrate?
Lab Procedure
4. You will be given an unknown hydrate of copper (II) sulfate. Explain the experimental
steps (lab procedure – heating to constant mass) you would need to complete to
determine the formula of the hydrate. Then explain how you would use the data you
collected to solve for the formula of the hydrate (calculations).
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Lab Write Up
Complete the following, showing all work. Neatly write up the answers, and submit them to
your teacher.
1. Perform the following calculations NEATLY SHOWING ALL WORK:
a. Calculate the exact mass of the hydrate that was added to the crucible (before
heating).
b. Calculate the mass of the sample after heating.
c. Calculate the mass of the water that was driven off.
d. Find the percentage of water that was driven off.
e. Determine the formula of the hydrate based on your data.
2. The actual formula of the hydrate is CuSO4 · 5H2O. Based on this formula, calculate the
actual percentage of water.
3. Calculate the percent error, using the experimental value you found in question 1d and
the actual value your found in question 2.
% 𝐸𝑟𝑟𝑜𝑟 = |
(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)
| × 100%
𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
4. Discuss possible sources of error and how they affected your data and calculations.
NOTE: “plugging in numbers wrong” or miscalculations are NOT sources of error, as you
can easily go back and fix these. Think about the data you collected and possible sources
of error.
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CW 5: Mole-Mole Stoichiometry
Chemical Equations
A balanced chemical reaction can be interpreted in two ways.
 First, it can be thought of as describing how many molecules of reactants are consumed
in order to produce a certain number of molecules of products.
 Analogously, it can be thought of as describing how many moles of reactants are
consumed in order to produce the indicated number of moles of products.
___H2(g) + ___O2(g)  ___H2O(g)
(1)
___ Mg(s) + ___H2O(l)  ___Mg(OH)2(aq) + ___H2(g)
(2)
The Mole Ratio
A mole ratio is the ratio between the amounts in moles of any two compounds involved in a
chemical reaction. Mole ratios are used as conversion factors between products and reactants
in many chemistry problems.
1. Write all possible mole ratios for reaction (1). There should be six total, including
inverses.
2 𝑚𝑜𝑙 𝐻2
1 𝑚𝑜𝑙 𝑂2
1 𝑚𝑜𝑙 𝑂2
2 𝑚𝑜𝑙 𝐻2
2. Using you answers to the previous question, fill in the blanks with the correct mole ratio
or the solution to the problem. Use slashes to show units that cancel.
3 𝑚𝑜𝑙 𝐻2
a.
×
= 1.5 𝑚𝑜𝑙 𝑂2
1
b.
c.
d.
2.3 𝑚𝑜𝑙 𝑂2
1
×
1.7 𝑚𝑜𝑙 𝐻2 𝑂
1
8.7 𝑚𝑜𝑙 𝐻2
1
×
×
= 4.6 𝑚𝑜𝑙 𝐻2 𝑂
= ________ 𝑚𝑜𝑙 𝐻2
= _________ 𝑚𝑜𝑙 𝐻2 𝑂
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3. How is the mole ratio used like a conversion factor? What units does it allow you to
convert between?
4. Answer the following using this equation: ___H2 + ___O2  ___H2O
a. What is the mole ratio of H2 to H2O?
b. Suppose you had 20 moles of H2, how many moles of H2O could you make?
c. What is the O2 / H2O molar ratio?
d. Suppose you had 20 moles of O2, and enough H2, how many moles of H2O could
you make?
5. Use this equation to answer following: ___N2 + ___H2  ___NH3
a. If you used 1 mole of N2, how many moles of NH3 could be produced?
b. If 10 moles of NH3 were produced, how many moles of N2 would you need?
c. If 3.00 moles of H2 were used, how many moles of NH3 would be made?
d. If 0.600 moles of NH3 were produced, how many moles of H2 are required?
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CW 6: MORE Stoichiometry
Stoichiometry Road Map
Particles
Given
Particles
Wanted
Moles
Given
Mass Given
Moles
Wanted
Volume
(liters) @
STP Given
____ Li + ____ AlCl3  ____ Al + ____ LiCl
1. How many grams of Li are needed to produce 10 grams of LiCl?
a. Write out a flow map for this problem.
b. Solve.
2. How many grams of Al are consumed if 30.5 grams of LiCl are produced?
a. Write out a flow map for this problem.
b. Solve.
Mass
Wanted
Volume
(liters) @
STP
Wanted
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Tin (II) oxide reacts with nitrogen trifluoride to produce tin (II) fluoride and dinitrogen trioxide.
3. How many grams of NF3 are required to fully react 36.52 grams of SnO?
4. How many particles of N2O3 are produced from 51.3 grams of SnO?
5. How many moles of SnF2 are produced from 17.89 grams of N2O3?
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Theoretical Yield
The theoretical yield is the calculated amount of product that can be produced from a given
amount of reactants, assuming that everything in the lab goes perfectly. In reality, lab work is
never perfect, resulting in an experimental yield that is less than the calculated theoretical
yield.
6. Find the theoretical yield of FeCl3(s) in grams for each of the following.
___Fe(s) + ___Cl2(g)  ___FeCl3(s)
a. What is the theoretical yield in grams of FeCl3 if 3 L of Cl2 gas are used?
b. How many grams FeCl3 can be made from 15.0 grams of iron?
Because errors always occur during chemistry labs, the experimental yield of a product is
always less than the theoretical (calculated) yield. We express this using percent yield. The
highest possible percent yield is 100%, but in practice this will never occur.
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑌𝑖𝑒𝑙𝑑 =
𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑌𝑖𝑒𝑙𝑑
× 100%
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 (𝐶𝑎𝑙𝑐𝑢𝑙𝑎𝑡𝑒𝑑) 𝑌𝑖𝑒𝑙𝑑
7. Using the theoretical yields from question 6 and the experimental yields given below,
determine the percent yield for each of the problems in question 6.
a. Theoretical Yield: (See answer 6a above) g FeCl3 | Experimental Yield: 13.20 g
FeCl3
b. Theoretical Yield: (See answer 6b above) g FeCl3 / Experimental Yield: 40.56 g
FeCl3
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CW 7: Limiting Reactant: S’mores
A delicious treat, known as a S’more is constructed with the following ingredients and
amounts:
1 teddy graham cracker
1 mini chocolate bar
2 mini marshmallows
At the store, these items can only be obtained in full boxes, each of which contains one gross
of items. A gross is a specific number of items, analogous (but not equal) to one dozen. The
boxes of items have the following net weights (the weight of the material inside the box):
Box of teddy graham crackers 9.0 pounds
Box of mini chocolate bars
36.0 pounds
Box of mini marshmallows
3.0 pounds
*Each box contains one gross of items*
1. If you have 100 graham crackers, how many chocolate bars and how many
marshmallows do you need to make S'mores with all of the graham crackers?
2. If you have 1000 graham crackers, 800 chocolate bars, and 1000 marshmallows:
a. How many S'mores can you make?
b. What (if anything) will be left over, and how many of that item will there be?
Chemists refer to the reactant which limits the amount of product that can be made from a
given collection of original reagents as the limiting reagent or limiting reactant.
3. Identify the limiting reactant for question 2. Explain.
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4. Based on the information given, which of the three
ingredients (a teddy graham cracker, a mini chocolate
bar, or a mini marshmallow):
a. Weighs the most? Explain your reasoning.
Box of teddy graham crackers: 9.0 pounds
Box of mini chocolate bars: 36.0 pounds
Box of mini marshmallows: 3.0 pounds
1 box = 1 gross of items
b. Weighs the least? Explain your reasoning.
5. If you have 36.0 pounds of graham crackers, 36.0 pounds of chocolate bars, and 36.0
pounds of marshmallows:
a. Which item do you have the most of?
b. Which item do you have the least of?
c. Determine how many gross you have of each ingredient.
6. If you attempt to make S'mores from the materials in question 5, what will be the
limiting reagent?
a. How many gross of S'mores can you make?
b. How many gross of each of the two leftover items will you have?
c. How many pounds of each of the leftover items will you have?
d. How many pounds of S'mores will you have?
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7. Explain why is it not correct to state that if we start with 36 pounds each of graham
crackers, mini chocolate bars, and mini marshmallows, then we should end up with 3 ×
36 = 108 pounds of S'mores.
8. Write a balanced chemical equation for the formation of a S’more. Use G as the symbol
for teddy graham cracker, Ch for the mini chocolate bar, and M for marshmallow.
9. Using the bag of materials provided to you, complete the following.
a. Assuming that you had plenty of the other ingredients, how many s’mores could
you make from each ingredient?
__________ teddy grahams could make __________ s’mores
__________ mini chocolate bars could make __________ s’mores
__________ mini marshmallows could make __________ s’mores
b. Given the materials in your bag, what is the maximum amount of S’mores you can
actually make?
c. What is the limiting reactant?
d. How much of the excess reactants will be left over?
Page 28 of 45
CW 8: Limiting Reactant: Chemical Quantities
1. Consider the chemical reaction above, in which 2 moles of hydrogen gas (H2) react with
1 mole of oxygen gas (O2) to produce 2 moles of water vapor (H2O).
a. How many hydrogen (H2) molecules are in the chemical reagent container?
b. How many oxygen (O2) molecules are in the chemical reagent container?
c. How many water (H2O) molecules are produced from the molecules found in the
chemical product container?
d. Is there anything left over in the product container? Why?
e. What is the mole ratio of oxygen to water?
f. What is the mole ratio of hydrogen to water?
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2. You react 100.0 g of O2 and 50.0 g of H2 in a container and produce water.
2H2 + O2  2H2O
a. How many moles of oxygen are in 100.0 g? Explain how you found your answer.
b. How many moles of water could be made from 100.0 g of oxygen and excess
hydrogen?
c. How many moles of hydrogen are in 50.0 g? Explain how you found your answer.
d. How many moles of water could be made from 50.0 g of hydrogen and excess
oxygen?
e. Based on your answers to b and d, how many moles of water will be produced
from reacting 100.0 g of O2 and 50.0 g of H2? Explain.
f. Which reactant is the limiting reactant in this scenario? Why?
g. What is the theoretical yield of water in this scenario? Report your answer in
grams of water.
h. What might have happen during this chemical reaction that prevented it from
producing the theoretical yield of water in lab? In other words, why might your
water production be lower than expected?
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3. Given the balanced chemical reaction: 2 NO(g) + O2(g) 2 NO2(g)
a. Calculate the mass of nitrogen dioxide that can be made from 30.0 grams of NO
and 30.0 grams of O2. (HINT: Do two calculations.)
b. Explain why the amount of NO2 that can be made depends on the limiting
reactant.
4. Acetylene gas, HCCH, is commonly used in high temperature torches.
a. Write a chemical equation for the reaction of acetylene (HCCH) with hydrogen gas
to form ethane (C2H6).
b. How many grams of ethane can be produced from a mixture of 30.3 grams of
HCCH and 4.14 grams of H2?
c. How does the above calculation help determine the limiting reactant?
Page 31 of 45
CW 9: Molarity and Precipitates
What is Molarity?
The concentration of a solute in an aqueous solution can be expressed in many ways: grams of
solute per liter of solution; grams of solute per 1000 grams of water; moles of solute per 1000
grams of water; and so on. One of the most frequently used concentration units is molarity.
𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑒 𝑖 =
The unit for molarity is
𝑚𝑜𝑙𝑒𝑠
𝑙𝑖𝑡𝑒𝑟
𝑀𝑜𝑙𝑒𝑠 𝑜𝑓 𝑖
𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑆𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠
and is represented by the symbol M (pronounced “molar”).
1. When one mole of Na2SO4 is dissolved in water:
a. How many moles of sodium ions are found in the solution?
b. How many moles of sulfate ions are found in the solution?
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2. Verify that when 10.0 g of sodium sulfate dissolves in water:
a. There are 7.04 × 10–2 moles of sodium sulfate in the water.
b. There are 7.04 × 10–2 moles of sulfate in the water.
c. There are 14.1 × 10–2 moles of sodium in the water.
3. In the solution pictured above, what is the molarity of the
a. Sodium sulfate?
b. Sulfate ions?
c. Sodium ions?
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4. Which is more concentrated with respect to sodium ions, 50.0 g of NaCl in 0.5 L of
solution or 59.0 g of Na2SO4 in 0.5 L of solution?
5. Which is more concentrated with respect to sodium ions, 0.50 M NaCl or 0.30 M
Na2SO4?
Making a 0.5 Molar Solution
6. Watch the video of Ms. L making a solution using a volumetric flask. Write down the
steps you will need to follow to make your solutions.
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7. Your task, as a group of 4, is to make 50 mL of a 0.5 M solution of CaCl2 or Na2CO3.
a. What is the equation for molarity? What units should the volume be in?
b. Using the molarity equation, solve for the moles of each substance required for
50 mL of a 0.5 M solution.
i. CaCl2:
ii. Na2CO3:
c. Use molar mass to calculate the grams of each substance needed to make your
solutions.
d. If you measured out 10 mL of the 0.5 M CaCl2 solution, how many moles of CaCl2
would that be?
e. If you measured out 15 mL of the 0.5 M Na2CO3 solution, how many moles of
Na2CO3 would that be?
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8. Present your calculations and procedure to Ms. L to obtain your materials. Then make
your solutions.
50 mL of 0.5 M CaCl2
50 mL of 0.5 M Na2CO3
Calculated mass of CaCl2 needed:
Calculated mass of Na2CO3 needed:
Actual mass of CaCl2 measured:
Actual mass of Na2CO3 measured:
Actual molarity of CaCl2 based on actual mass:
Actual molarity of Na2CO3 based on actual mass:
9. What are some possible sources of error that may cause the concentration of your
solutions to be “off”?
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Predicting Precipitates
10.You will be provided with the following 1M solutions: CaCl2, Na2CO3, CuSO4.
a. Select two solutions that will form a solid product. Write out the equation for
your chosen reaction below.
b. How many grams of your precipitate can be made with 50 mL of each of your
reacting aqueous solutions?
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CW 10: Determination of PO4-3 Concentration
Background:
Nutrient pollution in the Chesapeake Bay is one of the biggest environmental challenges
facing Maryland and surrounding states. One of the major sources of nutrient pollution comes
from phosphates, which are applied as fertilizers in the form of sodium phosphate (Na3PO4).
When it rains, excess fertilizer runs off from farmlands and into the Bay. This can lead to algae
blooms and pH changes, which are harmful to the Bay ecosystem.
Environmental chemists use stoichiometry to test water samples to determine
phosphate levels. They can precipitate out the phosphate from a water sample, then use the
mass of precipitate collected to calculate the concentration of phosphate in the original
solution. Such tests determine if the concentration of phosphate is within safe levels and can
help scientists target specific areas for remediation or support legal actions such as litigation
or lawmaking.
Objectives:
 Predict the formation of a precipitate
 Perform stoichiometric calculations to determine the concentration of a compound in a
local water sample
Materials:
 Local water sample
 50 mL graduated cylinder
 Two 250 mL beakers
 Filter paper
 Funnel
 Ring stand with ring
 1 M CaCl2 solution
 1 M KNO3 solution
 1 M NaCl solution
Pre-Lab Calculations:
1. Given the possible solutions above, determine which one you will need to react to form
a precipitate from the sodium phosphate (Na3PO4). The solubility rules/ chart will be
helpful in doing this. Write the balanced chemical equation below.
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2. Suppose that you perform the reaction from question one using 100 mL of the water
sample. You obtain 0.86 grams of the precipitate.
a. Use the grams of the precipitate to determine how many moles of Na3PO4 were in
the sample.
b. Given the volume of 100 mL, what was the molarity of Na3PO4 in the sample?
3. Suppose that you failed to add enough of your chosen solution, leaving some of the
Na3PO4 unreacted in the water.
a. Would you obtain too much or too little of the precipitate? Explain.
b. Would your calculated molarity be too high or too low? Explain.
c. Which reactant do you want to be the excess reactant? Explain.
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4. Initial field tests have determined that the concentration of sodium phosphate is around
0.146 M Na3PO4.
a. How many moles of Na3PO4 are in a 25 mL of 0.146 M Na3PO4?
b. How many moles of your chosen compound are needed to fully react with this
number of moles?
c. Given that the concentration of your chosen solution is 1.0 M, how many mL of
the chosen solution are needed?
d. To ensure that the chosen solution is in excess (and all of the Na3PO4 reacts),
multiply your answer from the last question by 1.3 to get the volume you should
use in the experiment.
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Procedure:
Before you may begin the lab, paraphrase the procedure and complete the required
calculations. Show your work to Ms. L to get approval.
1. Use the 50 mL graduated cylinder to measure 25.0 mL of the water sample to be tested.
Record this as the volume of the sample. Pour this into one of the 250 mL beakers.
2. Rinse down the sides of the empty graduated cylinder using a distilled water wash
bottle, pouring the rinse into the reaction beaker. This ensures that all the reactant in
the graduated cylinder ends up in the reaction mixture.
3. Rinse the graduated cylinder in the sink using tap water, followed by a distilled water
rinse from the wash bottle.
4. Use the rinsed 50 mL graduated cylinder to measure out the correct volume (pre lab
question 4d) of your chosen solution. Pour this into the same 250 mL beaker.
5. Gently swirl the reaction beaker for about 10 seconds to
ensure that the reaction has gone to completion. Record
observations.
6. Obtain a piece of filter paper. Write your names on the paper
using a pencil. Mass the unused filter paper.
7. Using the ring stand, ring, filter paper, funnel, and your
empty 250 mL beaker, create the set up shown here.
8. Filter the contents of the reaction beaker, using a distilled
water wash bottle to rinse down any precipitate that sticks to
the sides of the reaction beaker.
9. Once all liquid has passed through the filter paper, carefully remove the filter paper.
Place the filter paper on the drying surface. Once the filter paper is dry, you will re-mass
it to determine the amount of solid you collected (next class).
10.Pour any waste liquids into the waste beaker by the sink.
11.Rinse ALL glassware using tap water, following by a distilled water rinse from the wash
bottle.
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Data Table:
Volume of water sample
Volume of reacting solution
Mass of filter paper before filtration (with
names written in pencil)
Mass of dried filter paper with precipitate
Mass of precipitate collected (experimental
yield)
Observations
Post Lab Questions:
1. The actual concentration of the Na3PO4 solution was __________. Given that you used
50.0 mL of this solution, how many grams of precipitate should you collect (theoretical
yield)?
2. Determine how each of the following errors would affect the grams of Ca3(PO4)2
precipitate collected and thus the calculated Na3PO4 concentration of the water sample.
a. Not rinsing all of the reactants into the reaction mixture (250 mL beaker).
b. A hole in the filter paper.
c. Not enough of the chosen solution was added to the reaction mixture.
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5: Stoich Point
Determination of Phosphate Concentration
Name:
Partner(s):
Class:
Date:
To complete this lab, you will write a laboratory summary, using this sheet to guide you. Your
final lab report should be typed and submitted on Edmodo. The following sections should be
completed, in order, as they appear below. Please note that only the data and procedure may
be shared between group members. All other parts are in your own words.
COVER SHEET: A cover sheet with:
 Title of the lab
 Your name
 Your partners’ names
 Due date of the lab report
 Class period
PURPOSE: What are we trying to determine/ do in this experiment?
PROCEDURE: A step by step procedure for setting up the experiment and collecting data over
the course of the experiment. Directions should be numbered; and read something like a
recipe. Underline any materials you will need once you have written the procedure.
DATA: Organize ALL data into a neat data table. This means you will need in depth
observations. Things to observe:
 Measurements needed for calculations
 Possible sources of error
 What happened during the reaction
ANALYSIS AND CONCLUSION: Answer the following questions using complete sentences. You
should use at least 5 sentences for each question. For calculations, show all work, including
units.
1. Write and balance the chemical equation for your chosen solutions to make the
precipitate. Include state symbols (s, l, g, aq).
2. Using the volume of your water sample and the mass of Ca3(PO4)2 collected, calculate
the molarity of the Na3PO4 in the water sample.
3. The actual molarity of the sodium phosphate solution is 0.146 M. Calculate the percent
error.
(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)
% 𝐸𝑟𝑟𝑜𝑟 = |
| × 100%
𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒
4. Discuss three specific sources of error and how each error could have affected the
experimental yield.
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Grading Rubric:
Cover Page (IN YOUR OWN WORDS)
 Title of the lab: 0.5 point
 Your name: 0.5 point
 Your partners’ names: 0.5 point
 Due date of the lab report:0.5 point
 Class period: 0.5 point
/2.5
Purpose (IN YOUR OWN WORDS)
 Clear statement about what we set out to do with this lab
/2.5
Procedure (SAME AS GROUP MEMBERS)
 Step by step, repeatable, clear
/5
Data (SAME AS GROUP MEMBERS)
 Table is neat, organized, readable: 3 points
 Quality of observations: 7 points
/10
Conclusion (IN YOUR OWN WORDS)
 Question 1: 5 points
 Question 2: 10 points
 Question 3: 5 points
 Question 4: 10 points
/30
TOTAL:
/50
Reference Materials
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