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Chesapeake Campus – Chemistry 111 Laboratory
Lab #5 – Limiting Reagent
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Use stoichiometry to determine the limiting reactant.
Calculate the theoretical yield.
Calculate the percent yield of a reaction.
In lecture you have learned to read chemical equations and evaluate the mol to mol ratios of
reactants and products involved in a chemical reaction. In laboratory experiments it is difficult to
measure out chemicals in the exact ratio necessary for the chemical reaction. For time and speed
reasons, the reaction mixtures in lab will usually have a limiting and an excess reactant.
Limiting reagent (also called limiting reactant) problems use stoichiometry to determine the
theoretical yield for a chemical reaction. The limiting reactant will be completely consumed in the
reaction and limits the amount of product you can make. The limiting reactant also determines the
amount of product you can make (the theoretical yield). The reactant that is left over after the
reaction is complete is called the excess reactant.
Consider the process it takes to make a ham sandwich. You need two slices of bread and one
piece of ham to make each sandwich. How many complete sandwiches could you make if you had
eighteen slices of bread and six slices of ham? Let’s set this up like a chemical equation where the
ham and bread are our reactants and the sandwich is our product:
Two slices of bread + 1 piece of ham = 1 ham sandwich
Now we can use stoichiometry to determine the amount of sandwiches we could make if we
used all of our reactants.
18 𝑠𝑙𝑖𝑐𝑒𝑠 𝑜𝑓 𝑏𝑟𝑒𝑎𝑑 1 𝑠𝑎𝑛𝑑𝑤𝑖𝑐ℎ
2 𝑝𝑖𝑒𝑐𝑒𝑠 𝑜𝑓 𝑏𝑟𝑒𝑎𝑑 = 9 𝑠𝑎𝑛𝑑𝑤𝑖𝑐ℎ𝑒𝑠
6 𝑠𝑙𝑖𝑐𝑒𝑠 𝑜𝑓 ℎ𝑎𝑚 1 𝑠𝑎𝑛𝑑𝑤𝑖𝑐ℎ
1 𝑠𝑙𝑖𝑐𝑒 𝑜𝑓 ℎ𝑎𝑚 = 6 𝑠𝑎𝑛𝑑𝑤𝑖𝑐ℎ𝑒𝑠
Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 2
We have enough bread to make 9 sandwiches. There is not enough ham available to make as
many sandwiches. If we use all the ham, we can only make 6 sandwiches. Since the ham limits the
number of sandwiches we can make, the ham is our limiting reagent and the bread is going to be in
excess when the ham is consumed by the reaction. Additionally, since the theoretical yield depends
on the limiting reactant, we can say that our theoretical yield for the above reaction is 6 sandwiches.
It does not matter that there is enough bread to make 9 sandwiches. Once the ham runs out, it is not
possible to make any more sandwiches. The reaction is complete at this point. Limiting reactant is
completely consumed while the excess reactant (bread) is left over. We could even take this a step
further and determine the amount of excess reagent left over at the end of the reaction. In order to
perform that calculation, use the theoretical yield to calculate the amount of excess reactant used in
the reaction. Then you can subtract the amount of excess reactant used in the chemical equation
from the amount you began with.
For example given the balanced reaction N2 + 3 H2 → 2 NH3 If you began with 28 g of N2
and 2.8 g of H2. Since it is not possible to determine which reactant is the limiting reactant simply
from the masses of the reactants, you must first convert the grams to moles using the molecular
Therefore 28 g N2 = 1 𝑚𝑜𝑙 𝑁2
28.0 𝑔 𝑁2 = 1 mole N2
And 2.8 g H2 = 1 𝑚𝑜𝑙 𝐻2
2.02 𝑔 𝐻2 = 1.4 mole H2
While there is indeed more H2 than N2 based on moles of reactants, this is not the final
answer! You must convert to the mol of product using the mol to mol ratio.
If N2 is completely used 1 mole of N2 2 𝑚𝑜𝑙 𝑁𝐻3
1 𝑚𝑜𝑙 𝑁2 = 2 mol NH3 produced
If H2 is completely used 1.4 mole of H2 2 𝑚𝑜𝑙 𝑁𝐻3
3 𝑚𝑜𝑙 𝐻2 = 0.93 mol NH3 produced
Thus, since H2 will produce less of the product, it is the limiting reagent and N2 is the excess
reagent. Here the theoretical yield is 0.93 mol NH3. Note that for the problems in today’s lab we will
then convert the mol of product to grams using its molar mass.
It is often important to calculate the percent yield of a reaction. If everything goes according to
plan, you will get exactly 100 percent of the theoretical yield produced in your reaction. However,
laboratory errors will often affect this number. Spills, calculation errors, not drying a product and
many other errors affect the mass of product obtained. Here the amount of product actually produced
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in the laboratory experiment is compared to the amount of product that should have been made
theoretically. Percent yield is given by the equation:
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑌𝑖𝑒𝑙𝑑 = 𝐴𝑐𝑡𝑢𝑎𝑙 (𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙) 𝑌𝑖𝑒𝑙𝑑
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑌𝑖𝑒𝑙𝑑 𝑥 100
Guidelines for Limiting Reagent Problems (Calculating Theoretical Yield):
Convert from grams of reactant added to mol using molar mass.
Convert from mol of reactant to mol of product using the coefficients in the balanced equation
(mol to mol ratio).
Convert from mol of product to mol of reactant using the molar mass.
These problems must be worked out stoichiometrically.
You cannot compare masses of reactants. You MUST convert to mol.
You cannot compare mol of one reactant to another UNLESS you consider the mol to mol
ratio. In the reaction between hydrogen and nitrogen, there is technically more mol of
hydrogen added to the reaction vessel. However once the mol to mol ratio is considered, it
becomes apparent that hydrogen is the limiting reactant even though it had more mol.
The molar mass of hydrates MUST include the mass of the water molecules attached to the
In this experiment, you will predict and observe a limiting reactant during the reaction which
involves the reduction of copper (II) chloride dihydrate. You will use the single displacement reaction
of solid aluminum with aqueous copper (II) chloride.
2 Al(s) + 3 CuCl2 • 2 H2O (aq) → 3 Cu (s) + 2 AlCl3 (aq) + 6 H2O (l)
Copper (II) chloride, CuCl2, turns a light blue in aqueous solution. This is due to the Cu2+ ion.
Aluminum chloride is colorless in aqueous solution. You will be able to monitor the reaction’s
progress by evaluating the color change occurring in your beaker. The production of solid copper is
relevant in many industrial processes. Copper is mankind’s oldest metal, dating back more than
10,000 years. The copper (II) chloride reduction reaction has been used in petroleum industries for
sweetening (a refining process used to remove sulfurous gases from natural gas). This process has
also been used for etch bath regeneration. In an etch bath, CuCl2 is used to remove unwanted copper
from printed copper coated wiring boards, leaving just copper “wiring”.
*Note that the hydrate portion of CuCl2 •2 H2O should be included in the molar mass calculation.
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Student tray containing the following:
o 2-250ml beaker
o 1-10ml graduated cylinder
o 1-25ml graduated cylinder
o 1-100ml graduated cylinder
o 1 stir rod
o 1 spatula
o 1 container of 6M HCl w/pipet
o 1 container of Methanol w/pipet
o 1 container of CuCl2 ∙ 2 H2O
o 1 container of Al
o Beaker labeled ACID WASTE
o 1 DI water bottle
o Weigh boats
o Beaker of disposable supplies
Safety and Notes
Review MSDS information on all chemicals before coming to class.
Reactions should be done under hood.
HCl is corrosive; please use protective gear and caution. If you get it on your skin, flush with copious amounts of water and inform your instructor.
Methanol is flammable. Keep away from heat sources.
All waste should be ultimately disposed of in the container labeled CHM 111 WASTE in the back hood.
Remember to clean up any spills, wash glassware and return all items to proper trays.
Do not touch metals with hands.
Use proper labeling techniques.
Experimental Procedure and Data
1. Label a 250 mL beaker as “A.” Weigh beaker and record your measurement in the data section.
2. Using an analytical balance and disposable weigh boats, weigh approximately 0.50 g CuCl2 •2 H2O
and 0.25 g Aluminum foil. Record the exact weight in the table below.
3. Place the Al and CuCl2 •2 H2O into beaker “A”. Make sure that the aluminum foil is unfolded so that
it will completely react.
4. Label a second 250 mL beaker “B”. Weigh and record.
5. Using a balance and a new disposable weigh boat, weigh out 0.70 g of CuCl2 •2 H2O and 0.05 g of
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6. Place the Al and CuCl2 •2 H2O into beaker “B”. Again, make sure that the aluminum foil is unfolded
so that it will completely react.
7. Look at the contents of each beaker. Record the color of substances and any other observations
(odor (waft), bubbling, heat formation, etc.) that are visible at the beginning of the reaction in the
Which reactant do you THINK is in excess in each beaker? WHY???? Record this in your data
8. Using a graduated cylinder, measure 50.0 ml of distilled water and add to each beaker. When water
is added to the beakers, the CuCl2 •2 H2O will dissolve and the reaction will proceed.
9. Stir the substances in the beakers occasionally with the stirring rod. The reaction should take about
30 minutes to complete.
10. Record any color changes or any other observations as the reaction proceeds in the data table.
11. As the reaction proceeds, record your observations (color changes, bubbling, etc) in the data table.
When the reaction has finished, evaluate the beaker: which reactant do you THINK (based on your
observations) is in excess in each beaker? WHY????
12. When the reaction is complete and you no longer notice bubbles forming, if there is excess
aluminum foil still observed in the beakers, add 6 M HCl in 1 mL portions under the hood until the
foil is completely reacted and no longer visible (but do not add more than 5 mL). Stir to dissolve.
13. Allow the solid Cu to settle in both beakers. Decant (pour off the liquid) the solution from the
beakers into a waste container. Be careful not to lose any of the copper.
14. Wash the copper solid with 15 mL of deionized water. Let solid settle. Decant (be careful to pour as
much water off as possible without losing any of the copper solid). Repeat once more
15. Wash the copper solid with 10 mL of methanol. Let solid settle. Decant.
16. Under the hood, heat the beakers on a hot plate at a low setting until dry. Avoid heating at high
temperatures for longer periods of time which may cause the unwanted oxidation of the copper
17. When the product appears dry, carefully place the beaker on wire gauze or paper towels (do NOT
place directly on the counter as the glassware could shatter).
18. After cooling, weigh the beaker and its contents. Record this in your data section.
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Lab Partner Name ______________________________________ Bin #__________________ Note – This pre-lab must be completed before you come to lab.
Pre-Lab Assignment Questions
1. Given the unbalanced equation __ Na + __ Cl2 → __ NaCl
If 10.0 g of Na and 14.0 g of Cl2 are reacted together in a lab experiment:
a) Balance the equation.
b) Calculate the number of moles of each reactant used.
c) What is the limiting reagent? Show your calculations for each reactant.
d) What is the maximum number of moles of NaCl that can be formed?
e) What is the maximum number of grams of NaCl that can be formed?
f) How many moles of the excess reactant will be remaining?
2. What would be the benefit of having a limiting reagent when performing a lab experiment? Why not simply make both reactants go to completion?
3. Can we tell from just the masses which of the two reactants will potentially be the limiting reagent? Explain why or why not? Keep in mind what is happening at the molecular level in a chemical reaction.
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Lab Partner Name ______________________________________ Bin #__________________
Experimental Data and Results
Beaker #A Beaker #B
Mass of Empty Beaker
grams CuCl2 • 2 H2O
MM CuCl2 • 2 H2O
Mol CuCl2 • 2 H2O
Theoretical Yield of Cu (if CuCl2 • 2 H2O is limiting)
Grams of Al
Moles of Al
Theoretical Yield of Cu (if Al is limiting)
Observations Before the Reaction Begins
Observations After Reaction is Complete
Mass of Beaker and Solid Cu
Mass of solid Cu formed during reaction
Moles of solid Cu formed during reaction
*Show an example calculation for every box for Beaker A for full credit.
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Lab Partner Name ______________________________________ Bin #__________________
*Calculations: For full credit, clearly show all calculations with all units labeled.
Beaker A Beaker B
Theoretical mass of Cu produced if CuCl2 • 2 H20 was the limiting reagent
Theoretical mass of Cu produced if Al was the limiting reagent
What is the Theoretical Yield?
What was the limiting reactant?
What was the % yield of the reaction?
Results, Discussions and Post-lab Questions
1. In these experiments, when and why did the reaction stop? Explain your answer at the particle level in regards to reactants available.
2. If we began the experiment with 0.70 g of CuCl2 • 2 H2O, according to the stoichiometry of the reaction, how much Al should be used to complete the reaction without either reactant being in excess? Show
3. Give two errors that could have occurred in your experiment. How would each have affected your percent yield?