In the TV show "Breaking Bad" the characters attempt to use HF acid to dissolve guns (among other things). Here we consider instead dissolving guns (which we will assume are pure iron) with sulfuric acid.
Complete the balanced reaction for reacting iron in dilute sulfuric acid to form aqueous FeSO4. Do not worry about formatting subscripts (i.e. O2 to represent diatomic oxygen gas is fine).
Fe + H2SO4 → FeSO4 + __H2(g)___
How many liters of 1 molar sulfuric acid would be required to dissolve 1 kg of iron? Assume the reaction from the previous part goes to completion. The molecular mass of Fe is 55.85 g/mol.
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It is apparent that they must use an outside source of electrical power to drive the dissolution of the iron. They use a small current so that the dissolution proceeds with the minimum voltage required. Assume standard values for the electrochemical potentials.
How much electrical energy supplied this way is thus required to dissolve an additional 1 kg of iron? Give your answer in kJ.
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The characters realize that their hideout has been discovered by the police and they still have a last handgun that weighs 0.25 kg to dissolve. The cops will get there in an hour, so they have to speed up the reaction, by driving it at a higher current. What is the minimum total voltage in volts they'll need to drive the reaction at to get rid of the gun in time? Consider excess potential because of activation losses only and the exchange current I0 to be 1 A for the reaction over the surface of the entire tank (not a current density). α is 0.5 and everything is being done at room temperature. Assume standard electrochemical potentials.
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