Download Predicting Moles of Reactants/Products

5/07/09
Chemistry
Unit 8: Stoichiometry
Key Learning
1
2
3
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Chemistry Standard/Outcomes
Use of Mole Ratios
Calculating Theoretical Yields
Percent Yield of a Reaction
Determining the Limiting Reactant
Mole Ratios
 Be able to balance reactions.
 Be able to write mole ratios from the coefficients of balanced reactions.
 Be able to use the mole ratio to convert moles of a reactant/product to moles of
another reactant/product.
Calculating Theoretical Yields
 Be able to convert grams to moles (and the reverse) for a substance.
 Be able to calculate the number of grams of product that would result from grams
of reactant.
Percent Yield
 Be able to calculate percent yield when the actual and calculated yields are given.
 Be able to calculate percent yield when the calculated yield needs to be
determined from the amount of reactant given.
Limiting Reactant
 Know the definition of a limiting reactant.
 Be able to identify the limiting reactant by calculating the theoretical yield
produced by both reactants.
 Be able to determine the amount of product that would be produced in a reaction.
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3
Name: _______________
Skill Practice:
Interpreting a Balanced Chemical Equation
In a chemical equation, the coefficients tell us how many molecules are
reacting.
Example:
1H2
1 molecule
+
1Cl2
→
2 HCl
1 molecule
2 molecules
We can represent what is happening with a diagram:
Notice that 1 hydrogen molecule and 1 hydrogen molecule make 2 hydrogen
chlorides.
Complete the blanks:
N2
+
____ molecule(s)
3 H2
→
____ molecule(s)
2 NH3
____ molecule(s)
Draw a diagram like the example above that represents what is happening in this
reaction:
4
A balanced equation is like a recipe. You can double or triple the recipe if you
like. Notice that any change in one coefficient will mean the other coefficients
change in the same ratio!
2 H2
+
→
O2
__4___ molecule(s)
2 H2O
__2__ molecule(s)
__4__ molecule(s)
Complete the blanks:
H2
1.
+
__3__ molecule(s)
N2
2.
3.
_____ molecule(s)
→
3 H2
_20__ molecule(s)
2 NH3
____ molecule(s)
+
____ molecule(s)
→
O2
____ molecule(s)
H2
2 HCl
_____ molecule(s)
+
2 H2
→
Cl2
2 H2O
_1000_ molecule(s)
+
Cl2
4. __6.02 x 1023__ molecule(s)
→
____ molecule(s)
2 HCl
_____ molecule(s)
_____ molecule(s)
Since Avogadro’s number = 6.02×1023 molecules = 1 mole
2 H2
5.
__2__ mole
+
O2
____ mole
→
2 H2O
____ mole
5
The coefficients in a chemical equation give the mole relationships of reactants
and products in a reaction.
Give the mole relationships for each of the following:
H2
6.
____ mole(s)
C3H8
7.
Cl2 →
+
+
____ mole(s)
2 HCl
____ mole(s)
5 O2
→
____ mole(s)
____ mole(s)
3 CO2
+
____ mole(s)
4 H2O
____ mole(s)
Mole Ratios:
2 H2
+
O2
→
2 H2O
A mole ratio is the mole relationship between two specific chemicals:
Example:
 The mole ratio between H2/H2O is: 2 mole H2 = 2 mole H2O
 The mole ratio between H2/O2 is: 2 mole H2 = 1 mole O2
 The mole ratio between O2/H2O is: 1 mole O2 = 2 mole H2O
9. Using the reaction in problem 7 above:
 What is the mole ratio for C3H8/O2? ____________ = ____________
 What is the mole ratio for CO2/H2O? ____________ = ____________
 What is the mole ratio for H2O/O2? ____________ = ____________
6
Skill Practice: Using Mole-Mole Relationships
Example:
C3H8
+
5 O2
→
3 CO2 +
4 H2O
1. How many moles of O2 are needed to completely react with 1 mole of C3H8?
Answer: 5 moles
2. How many moles of CO2 form when 5 moles of O2 react? Answer: 3 moles
3. How many moles of H2O form when 1 mole of C3H8 react? Answer: 4 moles
4. How many moles of C3H8 are required to produce 8 moles of H2O?
Answer: 2 moles.
Practice:
2 Mg
+
O2 
2 MgO
1. How many moles of O2 are needed to completely react with 6 moles of Mg?
Answer: ________
2. How many moles of MgO form when 3 moles of O2 react? Answer: _____
3. How many moles of MgO form when 8 moles of Mg react? Answer: ______
4. How many moles of O2 are required to produce 20 moles of MgO?
Answer: ________
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Mole Ratio Worksheet
Name: _____________
*Balance all equations first!
1) Given this equation: ___N2 + ___H2 ---> ___NH3, write the following
molar ratios:
a) N2 / H2 _____________ = ______________
b) N2 / NH3 _____________ = ______________
c) H2 / NH3 _____________ = ______________
2) Given the following equation: ___H2 + ___S8 ---> ___H2S, write the
following molar ratios:
a) H2 / H2S _____________ = ______________
b) H2 / S8 _____________ = ______________
c) H2S / S8 _____________ = ______________
3) Answer the following questions for this equation:
___ H2 + ___O2 ---> ___ H2O
a) What is the H2 / H2O molar ratio? _____________ = ______________
b) Suppose you had 20 moles of H2 on hand, how many
moles of H2O could you make? ______________
c) What is the O2 / H2O molar ratio? _____________ = ______________
d) Suppose you had 20 moles of O2, how many moles of H2O
could you make? ____________
4) Use this equation: ___N2 + ___H2 ---> ___NH3 for the following
problems.
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
be required? _________
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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Balanced Equations and Mole Ratios
Name: ____________
1. Balance each equation.
2. Write a mole ratio between the two chemicals with a star (*). Follow
the example in the first problem.
Mole ratio
1.
*
*
__1___NH4NO2 --> __2___H2O
2.
_____H2
3.
_____MgCO3 --> _____MgO + _____CO2
4.
_____P4 + _____Cl2 --> _____PCl5
5.
_____CrO3
6.
_____IF5
7.
_____NH3 + _____O2 --> _____NO + _____H2O
8.
*
*
_____HBrO3 --> _____Br2O5
9.
*
_____NO2 + _____H2O
*
+ ___1__N2
*
+ _____N2 --> _____NH3
*
*
*
*
*
*
--> _____Cr2O3 + _____O2
*
+ _____H2O
*
--> _____HF
+ _____HIO3
*
10.
1 mole NH4NO2 = 2 mole H2O
*
+ _____H2O
*
--> _____HNO3
*
*
_____NH4NO3 --> _____N2O + _____H2O
+ _____NO
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Using Balanced Chemical Equations
Predicting Moles of Reactants/Products
Answer the following questions in the space provided. Use the chemical reaction and
show your dimensional analysis!
TiCl4 + 2H2O  4HCl + TiO2
1. How many moles of HCl would you produce if you started with 3.11
moles of TiCl4?
2. How many moles of TiO2 would you produce if you started with 3.11
moles of TiCl4?
3. If you needed 0.003 mol of TiO2 how many moles of TiCl4 would you
need to start with?
4. If you needed 0.003 mole of TiO2, how many moles of H2O would
you need to start with?
5. If you needed 4.91 moles of HCl, how many moles of TiCl 4 should
you start with?
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Converting Grams to Moles
Name: ______________
Use dimensional analysis and show your work!
1. How many grams of NH3 are in 12 moles of NH3?
2. How many grams of K2CO3 are in 2.4 moles?
Use this equation to solve the problems below. (Balance it first!)
____CaH2 + ____H2O  ____Ca(OH)2 + ____H2
3. If you started with 2.50 g of CaH2, how many grams of H2 would be produced?
What mole ratio will you use? ___________=___________
 What molar masses will you use? 1mole = ________ g CaH2
1mole = ________ g H2
Solve the problem: (Fill in the missing numbers)

2.50 g CaH2
1 mole CaH2
____mol H2
_____ g CaH2 ____ mol CaH2
___g H2 =
1 mole H2
4. If you started with 0.112 g of H2O, how many grams of H2 would be produced?


What mole ratio will you use? ___________=___________
What molar mass will you use? 1mole = ________ g H2O
1mole = ________ g H2
Solve the problem: (Watch out!!! The mole ratio is different for this one!)
5. If you wished to produce 4.7 g of Ca(OH)2, how many grams of H2O would be
required?


What mole ratio will you use? ___________=___________
Record your molar masses here:
Solve the problem:
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Use this equation to solve the problems below. (Balance it first!)
____NaN3  ____Na + ____N2
6. If you started with 3.8 g of NaN3, how many grams of Na would be produced?


What mole ratio will you use? ___________=___________
Record your molar masses here:
Solve the problem:
7.
If you started with 0.86 g of NaN3, how many grams of N2 would be produced?


What mole ratio will you use? ___________=___________
Record your molar masses here:
Solve the problem:
8. If you wished to produce 5.7 g of N2, how many grams of NaN3 would be
required?


What mole ratio will you use? ___________=___________
Record your molar masses here:
Solve the problem:
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Introduction to Percent Yield:
You have learned from laboratory experience that there is sometimes a
difference between an experimental result and a mathematical result.
Using stoichiometry, we can mathematically determine the amount of a
product that should be formed during an experiment, yet we sometimes
find that we don't end up with exactly the right amount of product.
Percentage yield tells us how close to “perfect” we were. Specifically,
it allows us to calculate what percent of the expected product we are
able to account for by the end of our experiment. The formula that we
use is:
actual amount of product
percentage yield = ------------------------------------------- x 100
expected amount of product
Example 1: This is just like finding your grade after a test. You scored 28
questions correct out of 30. We compare the “actual” number of questions correct
(28) to the number you “should have answered correctly” (30):
Answer:
Questions correct (28)
percentage yield = ------------------------------------- x 100 =
Questions on the test (30)
93 %
Example 2: A student conducts a reaction that produces 2.755 grams of
copper. Mathematically he determines that 3.150 grams of copper should
have been produced. Calculate the student's percentage yield.
Answer:
actual amount of product: 2.755 g
expected amount of product: 3.150 g
actual amount of product
percentage yield = ------------------------------------------- x 100
expected amount of product
2.755g
percentage yield = --------------- x 100
3.150g
percentage yield = 87.5 %
Percent Yield Practice
Name:_________________
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Be sure to show how you plugged the numbers into the formula as was shown in the answer
to example #2. No credit will be given for answers only.
1.
In the reaction shown below, a student produces 22 grams of water.
Mathematically he determines that 29 grams of water should have been produced.
Calculate the student's percentage yield.
2H2 + O2  2H2O
2.
The actual amount of MgO synthesized in the reaction shown below was 0.77 g.
If the theoretical amount of MgO was calculated to be 1.30 g, what is the percent yield of
the reaction?
2Mg + O2  2MgO
3.
A chemist starts with 11.2 g of KClO3 and calculates that it should be possible to
make 7.6 g of KCl with that much KClO3. After having conducted the reaction, she finds
that she recovered only 6.8 g of KCl. What is her percent yield?
2KClO3  2KCl + 3O2
4.
You have 6.2 g of NH3 in the cabinet. You do a short calculation and figure out
that this much NH3 will yield 0.92 g of H2 in the reaction shown below. You conduct the
reaction and find that you actually get 0.90 g of H2. What is your percent yield?
2NH3  N2 + 3H2
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15
16
17
Limiting Reagents (mole comparisons)
Name: _______________
In each case, identify the limiting reagent. (Show your work!)
1.
C2H4 + 3O2  2CO2 + 2H2O
a) 2.00 moles of C2H4 are reacted with excess O2
b) 3.00 moles of C2H4 are reacted with 5.00 moles of O2
c) 1.45 moles of C2H4 are reacted with 5.00 moles of O2
2.
Zn
+ 2 HCl 
ZnCl2
+
H2
a) 27 moles of Zn are reacted with excess HCl.
b) 12 moles of Zn are reacted with 12 moles of HCl.
c) 1.49 moles of Zn are reacted with 2.81moles of HCl.
3.
2 Mg
+
O2 
2 MgO
a) 0.8 moles of Mg are reacted with excess O2.
b) 0.72 moles of Mg are reacted with 0.22 moles of O2.
c) 2.61 moles of Mg are reacted with 1.26moles of O2.
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Limiting Reagents
Name: _______________
In each case, identify the limiting reagent and the amount of product that would be
produced. (Show your work!)
Strategy:
 Convert g of each reactant to g of product. (Do two separate train track
problems.)
 The one that makes less product was the limiting reagent.
1.
For this equation: C2H4 + 3O2  2CO2 + 2H2O
What is the limiting reagent if 2.00 g of C2H4 are reacted with 5.00 g of O2? How
much CO2 would be made in this reaction?
What is the limiting reagent?
2.
For this equation: Zn
+ 2 HCl 
How much CO2 would be produced?
ZnCl2
+
H2
What is the limiting reagent if 27 g of Zn are reacted with 0.79 moles of HCl?
How much H2 would be produced in this reaction?
What is the limiting reagent?
How much H2 would be produced?
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3. _____ Cu(s) + _____ AgNO3(aq)  _____ Cu(NO3)2(aq) + _____ Ag(s)
If 2.5 g of copper and 5.5 moles of silver nitrate are available to react, what is
the limiting reactant?
4. _____ CaO(s) + _____ H2O(l)  _____ Ca(OH)2(aq)
How many grams of calcium hydroxide will be formed in this reaction when
4.44 g of calcium oxide and 7.77 g of water are available to react? Also
identify the limiting reagent.
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Introduction to Stoichiometry Lab
Name: ____________________
Goal: To make 0.5 g of CuO. (And to make an educated guess as to the amount of
starting material that will give us 0.5 g of CuO.)
Safety: Tie back long hair and loose clothing near the Bunsen burner. Copper carbonate
is toxic. Do not ingest it and wear goggles at all times. Be careful with the flame. Turn
it off if you are leaving your station for any amount of time.
Waste: After they cool down, the products can go in the garbage.
Carbonates are a special class of compounds that do this reaction when heated:
CuCO3  CuO + CO2
What kind of reaction is this? ________________ (choose from synthesis,
decomposition, double and single displacement.)
What you need to do: We want to make exactly 0.5 g of the CuO product. To do this
we need to decide how much CuCO3 to start with.

Weigh between 0.5 and 1.5 g of CuCO3 into your evaporating dish.
Mass of evaporating dish empty ________ g
Mass of evaporating dish filled with CuCO3 ________ g
Mass of CuCO3 that you are starting with ________ g

Set up a Bunsen burner as shown below:

Put the evaporating dish on the wire guaze and start heating. Stir the
powder with your glass stirring rod until all of the green has turned to
black.
Let the dish cool until you can safely touch it. Weigh the evaporating dish
again. ________ g

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


We need to be sure that all of the stuff has changed to product. Heat it
again for another minute, then cool, and weigh the dish again. ________ g
Did the mass change from the last time you weighed it? If so repeat
heating until it doesn’t change by more than 0.1 g. Record the final mass
of the evaporating dish here: _________ g
Mass of empty evaporating dish __________g (You already weighed this
above.)
Mass of black CuO after you subtract out the dish: _________ g
Did you get exactly 0.5 g? (0.45-0.55g is okay.) If you didn’t, predict how much
CuCO3 you should start with to get exactly 0.5 g of CuO in the end. Try it again with
your new predicted amount.
Record the data that you collected for your second attempt here:
Mass of CuCO3 at start
Mass of the CuO at the end.
Remember units!
Conclusion:

What mass of CuCO3 do you need to start with to get 0.5 g of CuO at the end?
_____g (SHOW THIS TO YOUR INSTRUCTOR FOR CONFIRMATION.)

Why does the mass change going from reactants to products?