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Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 1

Chesapeake Campus – Chemistry 111 Laboratory

Lab #5 – Limiting Reagent

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Objective

 Use stoichiometry to determine the limiting reactant.

 Calculate the theoretical yield.

 Calculate the percent yield of a reaction.

Introduction

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.

 

Example 1:

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.

 

Example 2:

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

weights.

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.

Percent Yield:

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

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 3

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.

Helpful Hints:

 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

ionic compound.

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.

 

 

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 4

Materials

 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-tongs

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

o Tweezers

 Student balance

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

aluminum.

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 5

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

data table.

Which reactant do you THINK is in excess in each beaker? WHY???? Record this in your data

section.

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

product.

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.

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 6

Name________________________________________________ Date__________________

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.

 

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 7

Name________________________________________________ Date__________________

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.

 

 

Modified from Lumen Learning 2014 CC-BY 4.0 CHM 111 Lab 5 https://lumen.instructure.com/courses/150410/modules Page 8

Name________________________________________________ Date__________________

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

your calculations.

 

 

 

3. Give two errors that could have occurred in your experiment. How would each have affected your percent yield?

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