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Chapter 17: Problem 9

(a) From the galvanic series (Table 17.2), cite three metals or alloys that may be used to galvanically protect 304 stainless steel in the active state. (b) As Concept Check 17.4(b) notes, galvanic corrosion is prevented by making an electrical contact between the two metals in the couple and a third metal that is anodic to the other two. Using the galvanic series, name one metal that could be used to protect a copper-aluminum galvanic couple.

Short Answer

Expert verified
Answer: Magnesium, zinc, and aluminum can be used to galvanically protect 304 stainless steel in the active state. Zinc can be used to protect a copper-aluminum galvanic couple.

Step by step solution

01

(a) Galvanic Protection for 304 Stainless Steel in Active State

To find metals or alloys that can protect 304 stainless steel in the active state, we need to look at the galvanic series, particularly for metals that are more anodic (less noble) than 304 stainless steel in the active state. Based on Table 17.2, some suitable metals include: 1. Magnesium 2. Zinc 3. Aluminum These metals can be used as sacrificial anodes, as they will corrode preferentially to protect the 304 stainless steel.
02

(b) Protecting a Copper-Aluminum Galvanic Couple

To find a metal that can protect a copper-aluminum galvanic couple, we need to find a metal that is more anodic (less noble) than both copper and aluminum in the galvanic series. This will ensure that the protection metal is preferentially corroded before copper and aluminum. Looking at Table 17.2, one suitable metal is: 1. Zinc Zinc is more anodic compared to both copper and aluminum, so it can be used as the sacrificial anode to protect the copper-aluminum galvanic couple from corrosion.

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Most popular questions from this chapter

In the following table, weight gain-time data for the oxidation of some metal at an elevated temperature are given. \begin{tabular}{cc} \hline\(W\left(m g / c m^{2}\right)\) & Time (min) \\ \hline \(1.90\) & 25 \\ \(3.76\) & 75 \\ \(6.40\) & 250 \\ \hline \end{tabular} (a) Determine whether the oxidation kinetics obey a linear, parabolic, or logarithmic rate expression. (b) Now compute \(W\) after a time of 3500 min.

An electrochemical cell is constructed such that on one side a pure nickel electrode is in contact with a solution containing \(\mathrm{Ni}^{2+}\) ions at a concentration of \(3 \times 10^{-3} M\). The other cell half consists of a pure Fe electrode that is immersed in a solution of \(\mathrm{Fe}^{2+}\) ions having a concentration of \(0.1 M\). At what temperature will the potential between the two electrodes be \(+0.140 \mathrm{~V} ?\)

Using the results of Problem 17.13, compute the corrosion penetration rate, in mpy, for the corrosion of iron in citric acid (to form \(\mathrm{Fe}^{2+}\) ions) if the corrosion current density is \(1.15 \times 10^{-5} \mathrm{~A} / \mathrm{cm}^{2}\)

According to Table \(17.3\), the oxide coating that forms on silver should be nonprotective, and yet Ag does not oxidize appreciably at room temperature and in air. How do you explain this apparent discrepancy?

The corrosion rate is to be determined for some divalent metal M in a solution containing hydrogen ions. The following corrosion data are known about the metal and solution: \begin{tabular}{rr} \hline \multicolumn{1}{c}{ For Metal \(M\)} & For Hydrogen \\ \hline\(V_{\left(M M^{2}+\right)}=-0.47 \mathrm{~V}\) & \(V_{\left(\mathrm{H}^{+} / H_{2}\right)}=0 \mathrm{~V}\) \\ \(i_{0}=5 \times 10^{-10} \mathrm{~A} / \mathrm{cm}^{2}\) & \(i_{0}=2 \times 0^{-9} \mathrm{~A} / \mathrm{cm}^{2}\) \\ \(\beta=+0.15\) & \(\beta=-0.12\) \\ \hline \end{tabular} (a) Assuming that activation polarization controls both oxidation and reduction reactions, determine the rate of corrosion of metal \(\mathrm{M}\left(\mathrm{in} \mathrm{mol} / \mathrm{cm}^{2} \cdot \mathrm{s}\right)\) (b) Compute the corrosion potential for this reaction.
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