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Chapter 20: Problem 86

An iron object is plated with a coating of tin (Sn) to protect against corrosion. Does the tin protect iron by cathodic protection?

Short Answer

Expert verified
Tin does not protect iron by cathodic protection, as its standard oxidation potential (-0.14V) is less negative than iron (-0.44V), making it less likely to oxidize and act as a sacrificial metal. However, tin can still protect iron from corrosion by forming a barrier between the iron and the environment, preventing oxygen and moisture from reaching the iron surface.

Step by step solution

01

Determine the standard oxidation potentials of iron and tin

We must first find the standard oxidation potentials of both iron and tin to determine the likelihood of each metal undergoing corrosion. The standard oxidation potential is a measure of the tendency of a metal to lose electrons and oxidize. Standard oxidation potentials for Iron (Fe) and Tin (Sn) are as follows: \[ E°_{Fe^{2+}/Fe} = -0.44V \] \[ E°_{Fe^{3+}/Fe} = +0.04V \] \[ E°_{Sn^{2+}/Sn} = -0.14V \]
02

Compare the standard oxidation potentials of iron and tin

Now we compare the standard oxidation potentials of iron and tin. The more negative the standard oxidation potential, the more likely the metal is to oxidize. In this case, iron has a more negative value (-0.44V) compared to tin (-0.14V).
03

Determine if tin provides cathodic protection to iron

Since tin has a less negative standard oxidation potential than iron, it is less likely to oxidize and therefore would not act as a sacrificial metal for iron. Hence, tin cannot provide cathodic protection to iron. However, tin can still protect iron from corrosion by creating a barrier between the iron and the environment, preventing oxygen and moisture from reaching the iron surface. In conclusion, tin does not protect iron by cathodic protection, but it can still serve as a barrier to protect the iron from environmental factors that cause corrosion.

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Most popular questions from this chapter

Using the standard reduction potentials listed in Appendix E, calculate the equilibrium constant for each of the following reactions at \(298 \mathrm{~K}\) : (a) $\mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+2 \mathrm{Ag}(s)$ (b) $3 \mathrm{Ce}^{4+}(a q)+\mathrm{Bi}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 3 \mathrm{Ce}^{3+}(a q)+ \mathrm{BiO}^{+}(a q)+2 \mathrm{H}^{+}(a q)$ (c) $\mathrm{N}_{2} \mathrm{H}_{5}^{+}(a q)+4 \mathrm{Fe}(\mathrm{CN})_{6}^{3-}(a q) \longrightarrow \mathrm{N}_{2}(g)+ 5 \mathrm{H}^{+}(a q)+4 \mathrm{Fe}(\mathrm{CN})_{6}^{4-}(a q)$

(a) What is meant by the term oxidation? (b) On which side of an oxidation half-reaction do the electrons appear? (c) What is meant by the term oxidant? (d) What is meant by the term oxidizing agent?

During the discharge of an alkaline battery, \(4.50 \mathrm{~g}\) of \(\mathrm{Zn}\) is consumed at the anode of the battery. (a) What mass of \(\mathrm{MnO}_{2}\) is reduced at the cathode during this discharge? (b) How many coulombs of electrical charge are transferred from \(\mathrm{Zn}\) to \(\mathrm{MnO}_{2} ?\)

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 \(1.00 M\) and the cell generates an emf of \(+0.22 \mathrm{~V},\) what is the concentration of \(\mathrm{Pb}^{2+}\) in the anode half-cell? \((\mathbf{b})\) If the anode half-cell contains \(\left[\mathrm{SO}_{4}^{2-}\right]=1.00 M\) in equilibrium with \(\mathrm{PbSO}_{4}(s),\) what is the \(K_{s p}\) of \(\mathrm{PbSO}_{4} ?\)

(a) Suppose that an alkaline battery was manufactured using cadmium metal rather than zinc. What effect would this have on the cell emf? (b) What environmental advantage is provided by the use of nickel-metal hydride batteries over nickel-cadmium batteries?
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