Stoichiometry and limiting reactant
Introduction
In this experiment you will observe the reaction between metallic iron and a solution of
copper(II) sulfate. This reaction produces metallic copper, which is seen precipitating as a finely
divided red powder. The reaction in which one metal replaces another metal from a solution of
one of its salts is known as a displacement reaction. A metal capable of displacing another metal
from a solution of the metal’s salt is said to be “more active” than the displaced metal. In this
experiment, iron is more active than copper.
Iron forms two types of ions, Fe2+ and Fe3+. We can use stoichiometric principles to determine
which of these ions is formed in the reaction between iron and copper (II) sulfate solution.
In this experiment, an excess of copper (II) sulfate solution (to make sure that all the iron is
reacted) will be added to a known amount of iron. The metallic copper produced will be
weighed. Your task is to write the balanced chemical equations representing both possible
reactions and calculate the molar ratio between the Fe consumed and the Cu formed to find out
which one of these two equations is consistent with the results of your experiment.
The student kit contains:
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A balance
A hotplate
100mL beaker
~ 1 gram of iron powder
~35 mL of 1.0 M CuSO4
150 mL erlenmeyer flask,
2×250 mL beaker for waste
Two labeled waste containers for students to return the waste material to the college for
proper disposal
10 mL graduated cylinder
A bottle of distilled water
Acetone, several mL
PROCEDURE
1. Weigh a clean, dry 100mL beaker.
2. Accurately weigh approximately 1.0 gram of iron powder into the beaker.
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3. Measure 30 mL of 1.0 M CuSO4 solution into a graduated cylinder. Pour it into an erlenmeyer
flask, and heat gently to almost boiling.
4. Slowly add the hot CuSO4 solution to the beaker containing the iron powder.
5. Swirl the flask to insure the reaction goes to completion. When the reaction is completed,
allow the copper product to settle. Then carefully decant the liquid from the copper, (pour off the
liquid and leave the solid behind).
6. Add about 10 mL of distilled water to the solid copper and swirl to wash any remaining ions
from the copper.
7. Decant the wash water from the copper and add 10 more mL of distilled water, swirl and
decant again. Put the liquid from these two washes in the labeled waste container provided in
your kit.
8. Now wash the copper with several mL of acetone (Be careful! Acetone is very flammable)
Swirl and allow to stand a few minutes. Decant off the acetone. Repeat with a second portion of
acetone. Discard the acetone in a different waste labeled waste container. The acetone readily
dissolves the water and removes it.
9. Shake the beaker gently to spread the copper in an even layer on the bottom of the beaker.
Allow the acetone to evaporate entirely.
10. When the copper is dry, reweigh to find the mass of copper formed.
11. Calculate the moles of iron used and the moles of copper formed.
Data and Resulats
Beaker size (mL)
Mass of empty beaker
Mass of iron used
Moles of iron used
Mass of beaker plus copper
Mass of copper formed
Moles of copper formed
Moles of Cu divided by moles of Fe
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Calculations
1. From your data, write the balance equation which gives the correct stoichiometry for this
reaction? Explain your answer.
2. What visible evidence indicates that a chemical reaction has occurred?
3. State three reasons why you may not recover the theoretical amount of copper in this
experiment.
Problems:
1. A solution labeled as “Solution 1” contains 2.00 grams of Sodium Chloride.
How many grams of Silver Nitrate must be added to the solution to completely react with
Sodium Chloride according to the reaction below?
NaCl(aq) + AgNO3(aq)  NaNO3(aq) + AgCl(s)
2. A solution labeled as “Solution 2” contains 3.00 grams of AgNO3. When NaCl is added
to the solution, the following reaction occurs:
NaCl(aq) + AgNO3(aq)  NaNO3(aq) + AgCl(s)
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If excess NaCl is added to the solution, how many grams of AgCl(s) will be formed?
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