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Chapter 20: Problem 86

An iron object is plated with a coating of cobalt to protect against corrosion. Does the cobalt protect iron by cathodic protection? Explain.

Short Answer

Expert verified
Cobalt does not provide cathodic protection to iron objects because it has a lower standard reduction potential (-0.277V) than iron (-0.036V), meaning it is less likely to be reduced and cannot act as an effective cathode. Thus, the cobalt coating will not protect iron objects against corrosion in this manner.

Step by step solution

01

Understand the Corrosion process and Cathodic Protection

Corrosion is an electrochemical process where a metal, such as iron, is oxidized to form ions, while other substances, such as oxygen or water, are reduced. This process occurs in the presence of an electrolytic solution. Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. This works by providing a more easily reduced substance that won't corrode, which will keep the original metal from corroding.
02

Compare Standard Reduction Potentials

To determine if cobalt protects iron through cathodic protection, we need to compare the standard reduction potentials of cobalt and iron. Standard reduction potentials are measures of the tendency of a chemical species to acquire electrons and be reduced. A higher standard reduction potential implies that a species is more likely to be reduced. The standard reduction potential of iron (Fe) is: Fe鲁鈦� + 3e鈦� 鈫� Fe E掳 = -0.036V The standard reduction potential of cobalt (Co) is: Co虏鈦� + 2e鈦� 鈫� Co E掳 = -0.277V
03

Determine if Cobalt can provide Cathodic Protection

Since Co has a lower reduction potential than Fe, it is less likely to be reduced. However, for cathodic protection to occur, we need cobalt to act as an effective cathode, which means it must be more prone to reduction than iron. In this case, cobalt is not more prone to reduction than iron, which means that cobalt will not provide cathodic protection to iron objects.
04

Explain the Conclusion

Cobalt does not provide cathodic protection to iron objects because it has a lower standard reduction potential than iron. Consequently, cobalt is less likely to be reduced than iron, meaning that it cannot act as an effective cathode to protect the iron. Instead, iron will still be prone to corrosion even with a cobalt coating, so it is not an ideal method for protecting iron against corrosion.

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Most popular questions from this chapter

Elemental calcium is produced by the electrolysis of molten \(\mathrm{CaCl}_{2}\), (a) What mass of calcium can be preduced by this process if a current of \(7.5 \times 10^{3} \mathrm{~A}\) is applied for \(48 \mathrm{~h}\) ? Assume that the electrolytic cell is \(68 \%\) efficient. (b) What is the minimum voltage needed to cause the electrolysis?

(a) Write the anode and cathode reactions that cause the corrosion of iron metal to aqueous iron(II). (b) Write the balanced half-reactions involved in the air oxidation of \(\mathrm{Fe}^{2+}(a q)\) to \(\mathrm{Fe}_{2} \mathrm{O}_{3} \cdot 3 \mathrm{H}_{2} \mathrm{O}\),

A voltaic cell is constructed that is based on the following reaction: $$ \mathrm{Sn}^{2+}(a q)+\mathrm{Pb}(s) \longrightarrow \mathrm{Sn}(s)+\mathrm{Pb}^{2+}(a q) $$ (a) If the concentration of \(\mathrm{Sn}^{2+}\) in the cathode half-cell is \(1.00 \mathrm{M}\) and the cell generates an emf of \(+0.22 \mathrm{~V}\), what is the concentration of \(\mathrm{Pb}^{2+}\) in the anode half-cell? (b) If the anode half-cell contains \(\left[\mathrm{SO}_{4}{ }^{2-}\right]=1.00 \mathrm{M}\) in equilibrium with \(\mathrm{PbSO}_{4}(s)\), what is the \(K_{4 p}\) of \(\mathrm{PbSO}_{4}\) ? Batteries and Fuel Cells (Section 20.7)

A voltaic cell that uses the reaction $$ \mathrm{T1}^{3+}(a q)+2 \mathrm{Cr}^{2+}(a q) \longrightarrow \mathrm{Tr}^{+}(a q)+2 \mathrm{Cr}^{3+}(a q) $$ has a measured standard cell potential of \(+1.19 \mathrm{~V}\). (a) Write the two half-cell reactions. (b) By using data from Appendix \(\mathrm{E}\), determine \(E_{\mathrm{ed}}^{0}\) for the reduction of \(\mathrm{Ti}^{3+}(a q)\) to \(\mathrm{Ti}^{+}(a q)\). (c) Sketch the voltaic cell, label the anode and cathode, and indicate the direction of electron flow.

A voltaic cell is constructed that uses the following half-cell reactions: $$ \begin{aligned} \mathrm{Cu}^{*}(a q)+\mathrm{e}^{-} & \longrightarrow \mathrm{Cu}(s) \\ \mathrm{l}_{2}(s)+2 \mathrm{c}^{-} & \longrightarrow 2 \mathrm{I}^{-}(a q) \end{aligned} $$ The cell is operated at \(298 \mathrm{~K}\) with \(\left[\mathrm{Cu}^{+}\right]=0.25 \mathrm{M}\) and \(\left[1^{-}\right]=3.5 \mathrm{M}\). (a) Determine \(E\) for the cell at these concentrations. (b) Which electrode is the anode of the cell? (c) Is the answer to part (b) the same as it would be if the cell were operated under standard conditions? (d) If \(\left[\mathrm{Cu}^{+}\right]\)were equal to \(0.15 \mathrm{M}\), at what concentration of I \({ }^{-}\)would the cell have zero potential?
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