AP Chemistry
Oxidation-Reduction
Corrosion of Iron Experiment
Purpose
Observe the corrosion of iron and investigate conditions related to corrosion.
Pre-Lab Discussion
Corrosion is a oxidation-reduction (redox) reaction in which a free metal is oxidized, or
“corroded,” by some oxidizing agent. In nature, the oxidizing agent is frequently
atmospheric oxygen dissolved in water. Perhaps the most familiar example of corrosion
is the rusting of iron. In this reaction, free iron, Fe, is oxidized to Fe2+ and Fe3+ ions. In
the reduction half-reaction, oxygen, O2, reacts with water to form OH- ions. See
Chemistry The Central Science (Brown, LeMay, Bursten, Murphy) pages 874 – 876 for a
detailed description of the corrosion of iron.
The rusting of iron usually is considered to be a destructive change, and considerable
time and money are expended to prevent it. One method which you will observe in this
experiment, make use of a “sacrificial” metal. If iron is placed in contact with a more
active metal (one that is more easily oxidized) the more active metal will be corroded
instead of the iron. Metals situated below iron on the table of standard electrode
potentials can function as sacrificial metals.
In this experiment, indicators will be used to signal the formation of Fe2+ ions and OHions as products of the corrosion of iron. One such indicator, the hexacyanoferrate(III)
ion, turns blue in the presence of Fe2+ ions. The other, phenolphthalein, turns pink in the
presence of OH- ions. This experiment should help promote a better understanding of
corrosion and of oxidation-reduction reactions.
Pre-Lab Questions:
1. What type of reaction is the rusting of iron?
2. Write out the half-reactions, complete reaction, half-reaction voltages, and
reaction voltage for the production of rust in this experiment? Which is the anode
and which is the cathode?
3. Define the term “sacrificial” metal?
4. Why is iron rust considered destructive and copper rust is not destructive, but
beneficial and attractive? (What is the difference in the properties of each?)
Equipment
Balance
Tirrill Burner or Hot Plate
Graduated cylinder, 100 mL
Beaker, 250 mL
Ring stand
Iron ring
Wire gauze
Glass stirring rod
Dropper pipet
Petri dishes (2)
Emery paper
Metric ruler
Safety goggles
Lab apron
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Revised: 14 March 2013
AP Chemistry
Oxidation-Reduction
Materials
0.1 M K3Fe(CN)6 [potassium hexacyanoferrate(III)]
Agar powder
1% phenolphthalein solution
Distilled water
Copper wire, 5 cm
Zinc strip, 5 cm
10d Iron nails (4)
Safety
Observe all general lab safety precautions. Always wear safety goggles and a lab apron
when working in the lab.
Procedure
1. Set up a ring stand, ring, and wire gauze. Add 200 mL of distilled water to the 250
mL beaker and heat to boiling. While the water is being heated, clean four nails with
steel wool or emery paper.
2. Using the balance, measure out 2.00 g of agar powder. When the water is boiling
vigorously, put out the flame and slowly add the agar to the water, stirring constantly.
When the agar has dissolved, add 10 drops of 0.1 M K3Fe(CN)6. Rinse the dropper
and pipet and add five (5) drops of phenolphthalein solution.
3. Place one bent iron nail and one straight iron nail into one of the Petri dishes. Be sure
the nails do not touch and are positioned as shown in Figure 1.
Figure 1
4. Tightly wrap one of the remaining nails with copper wire and the other with the zinc
strip. Place these nails in the other Petri dish, as shown in Figure 1.
5. When the agar solution is cool enough, pour it into each Petri dish to a depth of 0.5
cm (make sure the nails are completely covered).
6. Make and record observations in the time remaining.
7. At the end of the lab period, place the dishes in your drawer to be observed on the
following day.
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Revised: 14 March 2013
AP Chemistry
Oxidation-Reduction
Observations and Data
(Write a complete description of what you see in each Petri dish and for each nail)
Petri Dish (a):
Petri Dish (b):
Lab Report
A brief, but complete report is expected for this lab. Please include the following:
The Introduction should be a discussion of the corrosion of iron, including the halfreactions that occur at the anode and cathode. Explain how a “sacrificial” metal works to
prevent corrosion of iron. Describe the experimental set-up, and explain the purpose of
the two indicators.
The Results should be a drawing or picture and a complete description of the qualitative
results. (Where are the pink and blue regions?)
The Conclusion should be a complete interpretation of the results
 What is the significance of the pink and blue regions around each nail?
 What conclusions can you draw from each nail? Compare and contrast the two
nails in the two dishes.
 Why did corrosion occur at the blue sites in Dish (a)?
 What is the whitish substance surrounding the zinc-wrapped nail in Dish (b)?
 Explain the difference in corrosion between the two nails in Dish (b).
 What effect does stressing the nail have on the corrosion of the nail?
 What other metals could be used as sacrificial metals?
 Did the experiment produce expected results?
 Are there any suggestions for improvements?
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Revised: 14 March 2013
AP Chemistry
Oxidation-Reduction
Teaching Notes
Equipment
Balance
Tirrill Burner or Hot Plate
Graduated cylinder, 100 mL
Beaker, 250 mL
Ring stand
Iron ring
Wire gauze
Glass stirring rod
Dropper pipet
Petri dishes (2 per student)
Emery paper
Metric ruler
Safety goggles
Lab apron
Materials
0.1 M K3Fe(CN)6 [potassium hexacyanoferrate(III)] FW 329.26 g
[use 3.2926g / 100mL]
Agar powder
1% phenolphthalein solution
Distilled water
Copper wire, 5 cm
Zinc strip, 5 cm [Use electrode strips cut]
10d Iron nails (4 per student)
Pre-Lab Questions:
1. What type of reaction is the rusting of iron?
Oxidation-reduction reaction.
2. Write out the half-reactions, complete reaction, half-reaction voltages, and
reaction voltage for the production of rust in this experiment? Which is the
anode and which is the cathode?
2 Fe(s)  2Fe2+(aq) + 4eEº= 0.44 V (Anode)
O2(g) + 2H2O(l) + 4e-  4OH-(aq) Eº= 0.41 V (Cathode)
2Fe(s) + O2(g) + 2H2O(l)  2Fe(OH)2(s) Eº= 0.85 V (white, green solid)
Further oxidation to red flakes:
More acidic water increases corrosion. If the pH is very low the hydrogen ions will
consume the electrons anyway, making hydrogen gas instead of water:
2H+(aq) + 2e- --> H2(g)
The iron(II) ions also react with hydrogen ions and oxygen to produce iron(III) ions:
4Fe2+(aq) -->4Fe3+(aq) + 4e- E° = -0.77 v (Anode)
4e- + 4H+(aq) + O2(aq) -->2H2O(l) E° = 0.815 v (Cathode)
4Fe2+(aq) + 4H+(aq) + O2(aq) -->4Fe3+(aq) + 2H2O(l)
E° = 0.145 v
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Revised: 14 March 2013
AP Chemistry
Oxidation-Reduction
The loose porous rust or Fe(OH)3 can slowly transform into a crystallized form written as
Fe2O3.H2O the familiar red-brown stuff that is called "rust" forming tubercles.
4Fe(OH)2(s) + O2(g)  2Fe2O3H2O(s) + 2 H2O(l)
3. Define the term “sacrificial” metal?
A metal that is more easily oxidized. Fe is above on the activity series (Zn, Al, Mg, Na,
Ca, K)
4. Why is iron rust considered destructive and copper rust is not destructive,
but beneficial and attractive? (What is the difference in the properties of
each?)
Iron rust does not bond to the base metal iron and is slightly water soluble. Therefore,
iron rust will rub off the metal and expose more metal to oxidation. Copper rust adheres
to the copper metal and protects the copper from other oxidation.
Lab Report
The Introduction should be a discussion of the corrosion of iron, including the halfreactions that occur at the anode and cathode. Explain how a “sacrificial” metal works to
prevent corrosion of iron. Describe the experimental set-up, and explain the purpose of
the two indicators.
The Results should be a drawing or picture and a complete description of the qualitative
results. (Where are the pink and blue regions?)
The Conclusion should be a complete interpretation of the results
 What is the significance of the pink and blue regions around each nail?
 What conclusions can you draw from each nail? Compare and contrast the two
nails in the two dishes.
 Why did corrosion occur at the blue sites in Dish (a)?
 What is the whitish substance surrounding the zinc-wrapped nail in Dish (b)?
 Explain the difference in corrosion between the two nails in Dish (b).
 What effect does stressing the nail have on the corrosion of the nail?
 What other metals could be used as sacrificial metals?
 Did the experiment produce expected results?
 Are there any suggestions for improvements?
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Revised: 14 March 2013
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Corrosion of Iron Experiment