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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
No, the cobalt coating does not provide cathodic protection to the iron object against corrosion. This is because the standard electrode potential of iron (\( E^0(Fe^{2+}/Fe) = - 0.44 V \)) is more negative than that of cobalt (\( E^0(Co^{2+}/Co) = - 0.28 V \)). As a result, iron is more easily oxidized (corroded) than cobalt, and the iron will still corrode preferentially over the cobalt coating.

Step by step solution

01

Understand Cathodic Protection

Cathodic protection is a technique used to control the corrosion of a metal surface by making it the cathode of an electrochemical cell. This is usually achieved by connecting the metal to be protected to a more easily corroded "sacrificial metal" that will act as the anode. The sacrificial metal then corrodes preferentially, protecting the cathode (the metal we want to preserve) from corrosion.
02

Electrochemical Series of Metals

To determine if cobalt can provide cathodic protection to iron, we need to look at the electrochemical series of metals. The electrochemical series is a list of metals arranged in order of their standard electrode potentials. In general, when two metals are in contact with each other in an electrolyte, the metal with the lower (more negative) electrode potential will be the anode, meaning that it will corrode preferentially.
03

Check Iron and Cobalt's Position on Electrochemical Series

Now, let's find the standard electrode potentials of both iron and cobalt from the electrochemical series: Iron(Fe): \[ E^0(Fe^{2+}/Fe) = - 0.44 V \] Cobalt(Co): \[ E^0(Co^{2+}/Co) = - 0.28 V \]
04

Determine if Cobalt can provide Cathodic Protection for Iron

From the electrochemical series, we can see that the standard electrode potential of iron (Fe) is -0.44 V, which is more negative than the standard electrode potential of cobalt (Co) which is -0.28 V. This means that in an electrochemical cell, iron is more easily oxidized (corroded) than cobalt. So, when a cobalt coating is applied on an iron object, the iron will still corrode preferentially over cobalt. Thus, the cobalt coating does not provide cathodic protection to the iron object against corrosion.

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Most popular questions from this chapter

For each of the following reactions, write a balanced equation, calculate the standard emf, calculate \(\Delta G^{\circ}\) at \(298 \mathrm{~K},\) and calculate the equilibrium constant \(K\) at \(298 \mathrm{~K}\). (a) Aqueous iodide ion is oxidized to \(\mathrm{I}_{2}(s)\) by \(\mathrm{Hg}_{2}^{2+}(a q) .\) (b) In acidic solution, copper(I) ion is oxidized to copper(II) ion by nitrate ion. (c) In basic solution, \(\mathrm{Cr}(\mathrm{OH})_{3}(s)\) is oxidized to \(\mathrm{CrO}_{4}^{2-}(a q)\) by \(\mathrm{ClO}^{-}(a q)\).

For a spontaneous reaction \(\mathrm{A}(a q)+\mathrm{B}(a q) \longrightarrow \mathrm{A}^{-}(a q)+\) \(\mathrm{B}^{+}(a q),\) answer the following questions: (a) If you made a voltaic cell out of this reaction, what halfreaction would be occurring at the cathode, and what half-reaction would be occurring at the anode? (b) Which half-reaction from (a) is higher in potential energy? (c) What is the sign of \(E_{\text {cell }}^{\circ}\) [ Section 20.3 ]

From each of the following pairs of substances, use data in Appendix \(\mathrm{E}\) to choose the one that is the stronger reducing agent: (a) \(\mathrm{Fe}(s)\) or \(\mathrm{Mg}(s)\) (b) \(\mathrm{Ca}(s)\) or \(\mathrm{Al}(s)\) (c) \(\mathrm{H}_{2}(g,\) acidic solution \()\) or \(\mathrm{H}_{2} \mathrm{~S}(g)\) (d) \(\mathrm{BrO}_{3}^{-}(a q)\) or \(\mathrm{IO}_{3}^{-}(a q)\)

If the equilibrium constant for a two-electron redox reaction at \(298 \mathrm{~K}\) is \(1.5 \times 10^{-4},\) calculate the corresponding \(\Delta G^{\circ}\) and \(E_{\text {red }}^{\circ}\).

The Haber process is the principal industrial route for converting nitrogen into ammonia: $$\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)$$ (a) What is being oxidized, and what is being reduced? (b) Using the thermodynamic data in Appendix \(\mathrm{C}\), calculate the equilibrium constant for the process at room temperature. (c) Calculate the standard emf of the Haber process at room temperature.
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