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Chapter 17: Q45E (page 961)

If a sample of iron and a sample of zinc come into contact, the zinc corrodes but the iron does not. If a sample of iron comes into contact with a sample of copper, the iron corrodes but the copper does not. Explain this phenomenon.

Short Answer

Expert verified

The metal which has the lowest reduction potential is the one that will be the anode and therefore the one which will corrode.

Step by step solution

01

Metal with the lowest reduction potential:

  • It has a significant proclivity to remove valence shell electrons and a low proclivity to take electrons and prefers oxidation.
  • The metal with the lowest electron-acceptance propensity has the lowest standard reduction potential.
  • Lithium, as a result, has the lowest standard reduction potential.
02

Find the metal which has the lowest reduction potential:

  • Corrosion is an electrochemical reaction where metal is the anode and oxygen is reduced at the cathode.
  • The product of the overall reaction is a metal oxide which we call "rust".
  • The metal which has the lowest reduction potential is the one that will be the anode and therefore the one which will corrode.
  • If we have two metals in contact, the first one to corrode will be the one that has the lower reduction potential.
  • We can look up the standard reduction potential in the table in the Appendixof the book. We can see that, if we compare the three metals in question (zinc, iron, and copper), zinc has the lowest reduction potential, copper the highest and iron is in the middle of the two.
  • This means that out of iron and zinc, zinc will corrode and iron will not; and out of iron and copper, iron will corrode but copper will not, as it says in the question.

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Most popular questions from this chapter

Why is a salt bridge necessary in galvanic cells like the one in Figure 17.4?

From the information provided, use cell notation to describe the following systems:

(a) In one half-cell, a solution of \({\rm{Pt}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}\)forms Pt metal, while in the other half-cell, Cu metal goes into a \({\rm{Cu}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}\)solution with all solute concentrations 1M.

(b) The cathode consists of a gold electrode in a \(0.55{\rm{MAu}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_3}\) solution and the anode is a magnesium electrode in \(0.75{\rm{MMg}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}\) solution.

(c) One half-cell consists of a silver electrode in a \(1{\rm{MAgN}}{{\rm{O}}_3}\) solution, and in the other half-cell, a copper electrode in \(1M{\rm{Cu}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}\) is oxidized.

Use the data in Appendix \({\rm{L}}\) to determine the equilibrium constant for the following reactions. Assume 298.15\({\rm{K}}\) if no temperature is given.

(a) \({\bf{AgCl(s)}}\rightleftharpoons {\bf{A}}{{\bf{g}}^{\bf{ + }}}{\bf{(aq) + C}}{{\bf{l}}^{\bf{ - }}}{\bf{(aq)}}\)

(b) \({\bf{CdS(s)}}\rightleftharpoons {\bf{C}}{{\bf{d}}^{{\bf{2 + }}}}{\bf{(aq) + }}{{\bf{S}}^{{\bf{2 - }}}}{\bf{(aq)}}\) at \({\bf{377\;K}}\)

(c) \({\bf{H}}{{\bf{g}}^{{\bf{2 + }}}}{\bf{(aq) + 4B}}{{\bf{r}}^{\bf{ - }}}{\bf{(aq)}}\rightleftharpoons {\left[ {{\bf{HgB}}{{\bf{r}}_{\bf{4}}}} \right]^{{\bf{2 - }}}}{\bf{(aq)}}\)

(d) \({{\bf{H}}_{\bf{2}}}{\bf{O(l)}}\rightleftharpoons {{\bf{H}}^{\bf{ + }}}{\bf{(aq) + O}}{{\bf{H}}^{\bf{ - }}}{\bf{(aq)}}\) at \({\bf{2}}{{\bf{5}}^{\bf{^\circ }}}{\bf{C}}\)

Aluminium metal can be made from aluminium ions by electrolysis. What is the half-reaction at the cathode? What mass of aluminium metal would be recovered if a current of 2.50 × 103 A passed through the solution for 15.0 minutes? Assume the yield is 100%

For each of the following balanced half-reactions, determine whether an oxidation or reduction is occurring.

(a) \({\bf{C}}{{\bf{l}}^{\bf{ - }}}{\bf{ + 3}}{{\bf{e}}^{\bf{ - }}} \to {\bf{C}}{{\bf{l}}_{\bf{2}}}\)

(b) \({\bf{M}}{{\bf{n}}^{{\bf{2 + }}}} \to {\bf{Mn}}{{\bf{O}}_{\bf{2}}}\)

(c) \({{\bf{H}}_{\bf{2}}} \to {{\bf{H}}^{\bf{ + }}}\)

(d) \({\bf{NO}}_{\bf{3}}^{\bf{ - }} \to {\bf{NO}}\)
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