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Chapter 20: Problem 82

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
No, cobalt does not provide cathodic protection to iron objects. Cobalt has a standard reduction potential of -0.28 V, which is less negative than iron's SRP values (-0.44 V and -0.04 V). Based on this comparison, cobalt is less likely to corrode than iron, so iron would corrode preferentially instead of experiencing cathodic protection. However, the cobalt coating can still act as a barrier against corrosive species.

Step by step solution

01

Understanding Cathodic Protection

Cathodic protection is a technique used to protect metal objects against corrosion by converting the metal to be protected into the cathode of a redox reaction. By reducing the metal's tendency to corrode, another material (the anode) will corrode preferentially instead of the protected metal. This sacrificial material needs to be more active (higher tendency to corrode) than the protected metal.
02

Compare the Standard Reduction Potentials (SRP) of Iron and Cobalt

We will refer to the SRP values for Iron and Cobalt. To compare their tendencies to corrode, we need to look at their standard reduction potentials. Iron (Fe) has two standard reduction potentials: \(Fe^{2+}(aq) + 2e^{-} \rightarrow Fe(s); E^{\circ} = -0.44 V\) \(Fe^{3+}(aq) + 3e^{-} \rightarrow Fe(s); E^{\circ} = -0.04 V\) Cobalt (Co) has the standard reduction potential: \(Co^{2+}(aq) + 2e^{-} \rightarrow Co(s); E^{\circ} = -0.28 V\)
03

Determine Cathodic Protection

For cobalt to provide cathodic protection to the iron object, it must have a more negative standard reduction potential (higher tendency to corrode) than iron. However, cobalt's SRP value is -0.28 V, which is less negative than the SRP values of iron. Therefore, based on the comparison of SRP values, cobalt cannot provide cathodic protection to an iron object. Instead, iron would corrode preferentially. In conclusion, cobalt does not provide cathodic protection to iron objects. However, it is worth noting that cobalt can still protect the iron surface by acting as a barrier against corrosive species (such as oxygen and water) that could attack the iron surface, but this is not due to cathodic protection.

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Most popular questions from this chapter

A plumber's handbook states that you should not connect a copper pipe directly to a steel pipe because electrochemical reactions between the two metals will cause corrosion. The handbook recommends you use, instead, an insulating fitting to connect them. What spontaneous redox reaction(s) might cause the corrosion? Justify your answer with standard emf calculations.

In each of the following balanced oxidation-reduction equations, identify those elements that undergo changes in oxidation number and indicate the magnitude of the change in each case. (a) \(\mathrm{I}_{2} \mathrm{O}_{5}(s)+5 \mathrm{CO}(g)-\mathrm{I}_{2}(s)+5 \mathrm{CO}_{2}(g)\) (b) \(2 \mathrm{Hg}^{2+}(a q)+\mathrm{N}_{2} \mathrm{H}_{4}(a q) \longrightarrow\) \(2 \mathrm{Hg}(l)+\mathrm{N}_{2}(g)+4 \mathrm{H}^{+}(a q)\) (c) \(\begin{aligned} 3 \mathrm{H}_{2} \mathrm{~S}(a q)+2 \mathrm{H}^{+}(a q)+2 \mathrm{NO}_{3}^{-}(a q) & \\ & 3 \mathrm{~S}(s)+2 \mathrm{NO}(g)+4 \mathrm{H}_{2} \mathrm{O}(l) \end{aligned}\) (d) \(\mathrm{Ba}^{2+}(a q)+2 \mathrm{OH}^{-}(a q)+\) \(\mathrm{H}_{2} \mathrm{O}_{2}(a q)+2 \mathrm{ClO}_{2}(a q)--\rightarrow\) \(\mathrm{Ba}\left(\mathrm{ClO}_{2}\right)_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)\)

During a period of discharge of a lead-acid battery, \(402 \mathrm{~g}\) of \(\mathrm{Pb}\) from the anode is converted into \(\mathrm{PbSO}_{4}(s) .\) What mass of \(\mathrm{PbO}_{2}(s)\) is reduced at the cathode during this same period?

A voltaic cell utilizes the following reaction: \(4 \mathrm{Fe}^{2+}(a q)+\mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)-\cdots 4 \mathrm{Fe}^{3+}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)\) (a) What is the emf of this cell under standard conditions? (b) What is the emf of this cell when \(\left[\mathrm{Fe}^{2+}\right]=1.3 \mathrm{M}\) \(\left[\mathrm{Fe}^{3+}\right]=0.010 \mathrm{M}, \mathrm{P}_{\mathrm{O}_{2}}=0.50 \mathrm{~atm}\), and the \(\mathrm{pH}\) of the so- lution in the cathode compartment is \(3.50 ?\)

A common shorthand way to represent a voltaic cell is to list its components as follows: anode|anode solution || cathode solution|cathode A double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change in phase, such as from solid to solution. (a) Write the half-reactions and overall cell reaction represented by \(\mathrm{Fe}\left|\mathrm{Fe}^{2+} \| \mathrm{Ag}^{+}\right| \mathrm{Ag}\); sketch the cell. (b) Write the half-reactions and overall cell reaction represented by \(\mathrm{Zn}\left|\mathrm{Zn}^{2+} \| \mathrm{H}^{+}\right| \mathrm{H}_{2} ;\) sketch the cell. (c) Using the notation just described, represent a cell based on the following reaction: $$ \begin{aligned} \mathrm{ClO}_{3}^{-}(a q)+3 \mathrm{Cu}(s)+& 6 \mathrm{H}^{+}(a q)--\rightarrow \\ & \mathrm{Cl}^{-}(a q)+3 \mathrm{Cu}^{2+}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l) \end{aligned} $$ \(\mathrm{Pt}\) is used as an inert electrode in contact with the \(\mathrm{ClO}_{3}^{-}\) and \(\mathrm{Cl}^{-}\). Sketch the cell.
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