Chemistry Honors – Corrosion Lab
Introduction:
Corrosion is the redox process by which metals (most commonly iron, or the iron alloy steel) are oxidized by elemental oxygen in the presence of moisture. If anything of yours has ever rusted after you left it outside, then you have first-hand experience with the damaging effects of corrosion!
Corrosion is responsible for the loss of billions of dollars annually in metal products. It is also a safety issue. For example, corrosion of support elements of a bridge could cause collapse. (The I-35W bridge collapse in Minneapolis in 2007 was determined to have been caused by a design flaw, but that event focused attention on the condition of many bridges where corrosion has the potential to cause a similar disaster.)
Studies have shown that oxidation of metals occurs most readily at points of strain. Thus a steel nail, which is mostly iron, first corrodes at the tip and head. A bent piece of metal is most susceptible to corrosion at the bend.
A point of strain acts as an anode, where the iron is oxidized to the ferrous ion and pits are formed. The electrons then flow to areas of the metal exposed to dissolved O2. These act as cathodes. The oxygen is reduced with water to form hydroxide ions. At the cathode, the OH- ions turn phenolphthalein pink.
When dissolved potassium hexacyanoferrate(III), K3Fe(CN)6, is present at the anode, this reaction occurs:
2+ 3- 3 Fe (aq) + 2 Fe(CN)6 (aq) → Fe3[Fe(CN)6]2 (aq), an intensely blue-colored compound.
Metals above Fe on the activity series, when attached to Fe, may prevent corrosion of Fe by supplying 2+ 2+ electrons to replace those lost by oxidation of the Fe. For example: Mg (s) + Fe (aq) → Mg (aq) + Fe (s). So the magnesium corrodes instead of the iron. A block of magnesium bolted to the hull of a ship protects against corrosion of the hull. The magnesium block is called a "sacrificial anode".
One may also notice a white precipitate in an area of corrosion. This may happen if the hydroxide ions combine with the metal ions.
The purpose of this lab is to determine the ability of various metals to protect iron from corrosion. You can choose two of these metals to test: copper, zinc, magnesium, tin, lead.
Prelab questions:
1. Corrosion is the process by which iron is oxidized by elemental oxygen. Therefore oxygen is the (oxidizing / reducing) agent and (gives / accepts) electrons.
2. Write the reduction half-reaction for the reduction of elemental oxygen in the presence of water, producing hydroxide ions.
Procedures:
1. Measure out approximately 100 ml of 0.1 M sodium nitrate solution into a graduated cylinder, pour into a 250 ml beaker and heat to approximately 80 °C. Make sure that the solution does not boil!
Be ready (with beaker tongs) to remove the beaker if it looks like it is going to boil over.
2. While the solution is warming up, tare an empty weigh boat, and measure out approximately 1.50 g of powered agar.
3. When it is ready, remove the solution from the hot plate, and add the agar to the warm solution, and stir the mixture until the agar forms a suspension. Re-heat as necessary to dissolve the agar.
4. Add 10 drops of 0.1 M potassium ferricyanide, and then 10 drops of phenolphthalein indicator, to the agar suspension. Stir to mix thoroughly. Your “reaction medium” is now ready!
5. Clean four 6 cm iron strips with steel wool.
6. Bend one of these strips into an L shape and place it on the bottom of the first petri dish. Place a second strip, unbent, next to the bent piece. Do not let the strips touch each other in the petri dish.
7. Decide which two metals you and your lab partner want to choose as experimental “sacrificial anodes.” Take one piece of each metal, and polish them with steel wool.
8. Lightly wrap one metal around the third 6 cm iron strip.
9. Lightly wrap second metal around the fourth 6 cm iron strip.
10. Place the wrapped iron strips on the bottom of the second petri dish. Do not allow the two wrapped strips to touch each other.
11. Before adding the agar, record “day 1 experimental conditions” in your data table.
12. Pour enough warm agar in both petri dishes to cover all of the iron strips to a depth of about 1 mm.
Pour carefully! You do not want the metal strips to be knocked into each other!
13. Cover both dishes, label with your group name, period and experimental conditions, and let stand for about 24 hours.
14. On day 2, observe the petri dishes against both white and black backgrounds and record your results.
15. Clean up on day 2:
Use tongs and forceps to remove metal pieces. Iron strips can be cleaned with water in the sink. Dry them thoroughly. Other metal pieces can be rinsed and put in separate waste containers as designated. Agar pieces can be thrown in the trash. Petri dishes can be cleaned and laid out to dry on the lab desk.
Data table:
You should prepare a data table that looks like this with plenty of room to write observations and sketch, using colored pencils.
Experimental conditions on day 1 Day 2 observations
Petri dish 1
Petri dish 2