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Aim
To familiarize oneself with some more common instances of corrosion.
Reagents
Hydrochloric acid 0.1 M; sulphuric acid 0.1 M; zinc and aluminium granules; copper wire ; solution of copper(II) sulphate; solution of copper(II) chloride; solution of iron (II) sulphate; solution of potassium hexacyanoferrate(III); zinc-plated iron sheet; tin-plated iron sheet; iron paper clips; solid NaCl; urotropin.
Equipment
Test tubes, small beaker (50 cm3), centrifuge tube .
Experiments
1. Making of a galvanic pair
1.1. Place a zinc granule on the bottom of the centrifuge tube and pour solution of hydrochloric acid on top. Write down the equation of reaction which occurs. Which compound is the oxidant, which is the reductant?
Oxidant is H:
Reductant is Zn:
Next, insert the copper wire into the same solution in the centrifuge tube, ensuring that it will not touch the zinc granule. Observe whether any hydrogen evolves on the copper surface? Explain , why copper did not react with diluted hydrochloric acid!
Adding copper wire doesnt make H extract. Copper is less active than H. H starts to extract when connect copper wire to Zinc, because zinc corrodes when connected to more active metal .
Now put the copper wire into contact with the zinc and observe whether hydrogen will evolve on the copper surface. By putting copper and zinc into contact with each other in the solution of hydrochloric acid (an electrolyte), you have created a galvanic pair. Zinc as the metal with more negative potential will be the anode and copper, with more positive potential, the cathode. Which of the two metals will dissolve (corrode)? Write the chemical equations of the reactions occurring on the anode and the cathode!
Reaction on anode:
Reactions on cathodes:
1.2. Place a zinc granule in a test tube and pour ~3 cm3 of CuSO4 solution on top of it. After a couple of minutes pour the solution out of the tube and wash the granule 2-3 times carefully with small amount of distilled water. What has happened to the granule? Write the equation of reaction!
Zinc granule turn black because copper settled on the surface
Reaction equation:
Place a clean zinc granule into another test tube and add ~3 cm3 of HCl solution into both test tubes. Compare the speeds of reaction in the two tube! Explain!
The activity of reaction increased after addinf HCl on zinc granule because Cl ions boosts the reactions.
1.3. Place an aluminium granule into each of two test tubes. Pour ~3 cm3 of CuSO4 solution on top of one of them , and an equal amount of CuCl2 solution onto the other one. Compare the intensity of the reactions in the two test tubes! Which of the anions (Cl– or SO42–) speeds up the reaction?
Reaction in tube where is CuCl2 solution included is more active. Cl boosts the reaction
Check the conclusion you just made by adding some solid NaCl into the test tube containing CuSO4 and observe whether the speed of reaction changes . Write the chemical equation of the reaction between aluminium and copper(II) chloride!
Reaction intensifies when adding to tube of CuSO4 some solid NaCl.
2. Proving presence of Fe2+ ions in solution
In order to prove presence of Fe2+ ions in solution, the solution of potassium hexacyanoferrate(III) is used. If Fe2+ ions are present in the solution, then after adding K3[Fe(CN)6], the blue -colored Fe3[Fe(CN)6]3 will precipitate. Copy down the chemical equation into the protocol :
3FeSO4 (aq) + 2K3[Fe(CN)6] (aq) = Fe3[Fe(CN)6]2 (s,aq) + 3K2SO4 (aq) (5.23)
In order to carry out the proof reaction and observe the colour of the product, pour ~2 cm3 distilled water into a test tube, add three drops of solution of iron(II) sulphate and then two drops of K3[Fe(CN)6] solution.
New compound took colour of dark blue. It means that there is Fe2+ ions in solution.
3. Metallic protective coatings
Pour ~3 cm3 of solution of sulphuric acid into each of two test tubes and add two drops of proving solution for Fe2+ to each (K3[Fe(CN)6]). Place a piece of zinc-plated iron sheet into one of the test tubes, and a piece of tin-plated iron sheet into another. Observe the formation of blue colour around the edges of one of the metal sheets (which one?).
Blue colours occurs around the edges of zinc covered ferrum.
This colour indicates that iron is corroding (Fe2+ ions are produced) in the corresponding system. What is the anode and what is the cathode in each of these cases? Write down equations of reactions on the anode and cathode for both cases!
First case :
Anode (Zn):
Cathode (Fe):
Second case:
Cathode (Fe):
Anode (Sn):
In which case the protection can be called anodic, and in which case, cathodic? In which case would damage to the protective layer be more of a problem?
On first case there is anodic protective layer and defections of this layers are more problematic. In second case it is cathodic protection.
4. Sacrificial anode protection
Pour ~5 cm3 of solution of sulphuric acid into a test tube, and pour about 1 cm layer of the same solution to the bottom of the beaker. Add two drops of Fe2+ proving solution (K3[Fe(CN)6]) to each. Place some iron wire (paper clip) and a zinc granule into the beaker in such a way that they would not touch each other. Physically connect another piece of iron wire to a zinc granule and place these into the test tube. Observe the difference in speed of corrosion (formation of blue colour) in the two systems. Which one is faster? Explain!
More intense reaction occurs in beaker where zinc granule and clipper are divided . Thats because ferrum is anode. If we connect ferrum to element with more negative potential, it turns into anode.
In case of connected metals, which one is the anode, which one is the cathode? Write down the equations of reactions occurring on the anode and cathode!
Anode (Zn):
Cathode (Fe):
5. Action of an inhibitor
Pour ~5 cm3 of solution of sulphuric acid into each of two test tubes. Add two drops of Fe2+ proving solution (K3[Fe(CN)6]) to each. With the aid of a spatula, add some solid corrosion inhibitor (urotropin) into the second test tube and shake. Place a piece of clean iron wire (paper clip) into each test tube, trying to insert them at the same time. Observe, in which order and with which speed the blue colour forms in the solutions ! Did the inhibitor reduce the rate of corrosion of iron?
Blue colour occours first in tube including K3[Fe(CN)6]. It means that inhibiitor accelerates reaction.

Conclution

We can conclude that inhibitor slowers strongly the corrosion of ferrum.