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Chapter 17: Q44E (page 961)

Aluminium\(\left( {{\bf{E}}_{{\bf{A}}{{\bf{l}}^{{\bf{3 + }}}}{\bf{/Al}}}^{\bf{^\circ }}{\bf{ = - 2}}{\bf{.07\;V}}} \right)\) is more easily oxidized than iron \(\left( {{\bf{E}}_{{\bf{F}}{{\bf{e}}^{\bf{3}}}}^{\bf{^\circ }}{\bf{/F}}{{\bf{e}}^{\bf{ - }}}{\bf{ = - 0}}{\bf{.477\;V}}} \right){\bf{,}}\) and yet when both are exposed to the environment, untreated aluminium has very good corrosion resistance while the corrosion resistance of untreated iron is poor. Explain this observation.

Short Answer

Expert verified
  • Aluminium reacts with atmospheric oxygen to form a stable aluminium oxide layer which is very tightly bound to the surface of the aluminium.
  • Iron on the other hand in reaction with atmospheric oxygen forms rust which can flake off the surface of iron, therefore, exposing the inner metal to further oxidation and corrosion.

Step by step solution

01

Describe Aluminum is more corrosion-resistant than iron:

  • Unlike iron and steel, aluminium does not rust or corrode under damp environments.
  • Its surface is protected by a natural coating of aluminium oxide.
  • This keeps the metal underneath from coming into touch with air and oxygen.
02

Find the observation of aluminium:

  • Aluminium is indeed a more reactive element than iron, but unline iron aluminium gives a stable corrosion product that protects the aluminium metal from further corrosion.
  • Aluminium reacts with atmospheric oxygen to form a stable aluminium oxide layer which is very tightly bound on the surface of the aluminium.
  • The oxide itself is not very reactive so it protects the aluminium metal underneath it, it is said that the oxide layer pacifies aluminium.
  • Iron on the other hand in reaction with atmospheric oxygen forms rust which can flake off the surface of iron, therefore, exposing the inner metal for further oxidation and corrosion.

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Most popular questions from this chapter

For each reaction listed, determine its standard cell potential at \({25\circ }{\rm{C}}\) and whether the reaction is spontaneous at standard conditions.

(a)\({\mathop{\rm Mn}\nolimits} (s) + {\rm{N}}{{\rm{i}}^{2 + }}(aq) \to {{\mathop{\rm Mn}\nolimits} ^{2 + }}(aq) + {\rm{Ni}}(s)\)

(b)\(3{\rm{C}}{{\rm{u}}^{2 + }}(aq) + 2{\rm{Al}}(s) \to 2{\rm{A}}{{\rm{l}}^{3 + }}(aq) + 3{\rm{Cu}}(s)\)

(c)\({\rm{Na}}(s) + {\rm{LiN}}{{\rm{O}}_3}(aq) \to {\rm{NaN}}{{\rm{O}}_3}(aq) + {\rm{Li}}(s)\)

(d) \({\rm{Ca}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}(aq) + {\rm{Ba}}(s) \to {\rm{Ba}}{\left( {{\rm{N}}{{\rm{O}}_3}} \right)_2}(aq) + {\rm{Ca}}(s)\)

Determine \({\bf{\Delta G}}\) and \({\bf{\Delta G}}^\circ \) for each of the reactions in the previous problem.

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 \({\bf{2}}{{\bf{5}}^{\bf{^\circ }}}{\bf{C}}\)

What mass of zinc is required to galvanize the top of a 3.00 m × 5.50 m sheet of iron to a thickness of0.100 mm of zinc? If the zinc comes from a solution of \(Zn{\left( {N{O_3}} \right)_2}\) and the current is 25.5 A, how long will it take to galvanize the top of the iron? The density of zinc is 7.140 g/cm3

Use cell notation to describe the galvanic cell where copper(II) ions are reduced to copper metal and zinc metal is oxidized to zinc ions.
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