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Chapter 20: Problem 80

(a) Based on standard reduction potentials, would you expect copper metal to oxidize under standard conditions in the presence of oxygen and hydrogen ions? (b) When the Statue of Liberty was refurbished, Teflon spacers were placed between the iron skeleton and the copper metal on the surface of the statue. What role do these spacers play?

Short Answer

Expert verified
(a) Yes, copper metal would oxidize under standard conditions in the presence of oxygen and hydrogen ions, as the overall reaction is spontaneous with a positive 螖E掳 value of 1.57V. (b) The Teflon spacers placed between the iron skeleton and the copper surface of the Statue of Liberty prevent galvanic corrosion by eliminating physical and electrical contact between the two metals. This intervention ensures the structural stability and longevity of both the iron skeleton and the copper surface of the statue.

Step by step solution

01

Determine the relevant half-reactions

Before analyzing whether copper will oxidize in the presence of oxygen and hydrogen ions, we need to identify two half-reactions. The copper half-reactions are: 1. Copper oxidation: \(Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-\) 2. Oxygen reduction in the presence of hydrogen ions (forming water): \(O_2(g) + 4H^{+}(aq) + 4e^- \rightarrow 2H_2O(l)\)
02

Check the standard reduction potentials

To determine if a reaction is spontaneous under standard conditions, we need to check the standard reduction potentials for both half-reactions. The values are given as: 1. Copper oxidation: \(E掳(Cu^{2+}/Cu) = +0.34 V\) (this value represents the reduction half-reaction, our reaction is the reversed one) 2. Oxygen reduction: \(E掳(O_2/H_2O) = +1.23 V\)
03

Check if the overall reaction is spontaneous

We can determine if the overall reaction is spontaneous by comparing the standard reduction potentials of both half-reactions. In this case, \(E掳(O_2/H_2O) > E掳(Cu^{2+}/Cu)\), implying that the oxygen reduction reaction is more likely to happen compared to copper oxidation. To get the net voltage, we then need to subtract the higher reduction potential from the lower value: \(螖E掳 = E掳(O_2/H_2O) - (-E掳(Cu^{2+}/Cu)) = 1.23 - (-0.34) = 1.57V\) Since 螖E掳 is positive, it indicates that the overall reaction is spontaneous under standard conditions, and copper will oxidize in the presence of oxygen and hydrogen ions. (a) The answer is YES. Copper metal would oxidize under standard conditions in the presence of oxygen and hydrogen ions.
04

Explain the role of Teflon spacers

(b) The Teflon spacers play an essential role in preventing galvanic corrosion between the iron skeleton and the outer copper surface of the Statue of Liberty. Galvanic corrosion occurs when two dissimilar metals, such as iron and copper, are in physical or electrical contact with each other and an electrolyte (such as saltwater or rain). The electrolyte facilitates electron transfer between the two different metals, resulting in corrosion of the least corrosion-resistant metal, which in this case would be iron. By placing Teflon spacers between the iron skeleton and the copper surface, physical and electrical contact between the two metals is prevented. This intervention eliminates the galvanic corrosion process, ensuring the structural stability and longevity of both the iron skeleton and the copper surface of the statue.

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Most popular questions from this chapter

From each of the following pairs of substances, use data in Appendix \(\mathrm{E}\) to choose the one that is the stronger oxidizing agent: (a) \(\mathrm{Cl}_{2}(g)\) or \(\mathrm{Br}_{2}(l)\) (b) \(\mathrm{Zn}^{2+}(a q)\) or \(\mathrm{Cd}^{2+}(a q)\) (c) \(\mathrm{BrO}_{3}^{-}(a q)\) or \(\mathrm{IO}_{3}^{-}(a q)\) (d) \(\mathrm{H}_{2} \mathrm{O}_{2}(a q)\) or \(\mathrm{O}_{3}(g)\)

The \(K_{s}\) value for \(\mathrm{PbS}(s)\) is \(8.0 \times 10^{-28} .\) By using this value together with an electrode potential from Appendix \(\mathrm{E}\), determine the value of the standard reduction potential for the reaction $$ \mathrm{PbS}(s)+2 \mathrm{e}^{-}--\rightarrow \mathrm{Pb}(s)+\mathrm{S}^{2-}(a q) $$

(a) A voltaic cell is constructed with all reactants and products in their standard states. Will this condition hold as the cell operates? Explain. (b) Can the Nernst equation be used at temperatures other than room temperature? Explain. (c) What happens to the emf of a cell if the concentrations of the products are increased?

Gold exists in two common positive oxidation states, \(+1\) and \(+3\). The standard reduction potentials for these oxidation states are $$ \begin{aligned} \mathrm{Au}^{+}(a q)+\mathrm{e}^{-}-\longrightarrow \mathrm{Au}(s) & E_{\mathrm{red}}^{\circ}=+1.69 \mathrm{~V} \\ \mathrm{Au}^{3+}(a q)+3 \mathrm{e}^{-}-\ldots \mathrm{Au}(s) & E_{\mathrm{red}}^{\circ}=+1.50 \mathrm{~V} \end{aligned} $$ (a) Can you use these data to explain why gold does not tarnish in the air? (b) Suggest several substances that should be strong enough oxidizing agents to oxidize gold metal. (c) Miners obtain gold by soaking goldcontaining ores in an aqueous solution of sodium cyanide. A very soluble complex ion of gold forms in the aqueous solution because of the redox reaction $$ \begin{array}{r} 4 \mathrm{Au}(s)+8 \mathrm{NaCN}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)-\longrightarrow \\ 4 \mathrm{Na}\left[\mathrm{Au}(\mathrm{CN})_{2}\right](a q)+4 \mathrm{NaOH}(a q) \end{array} $$ What is being oxidized, and what is being reduced, in this reaction? (d) Gold miners then react the basic aqueous product solution from part (c) with Zn dust to get gold metal. Write a balanced redox reaction for this process. What is being oxidized, and what is being reduced?

A voltaic cell similar to that shown in Figure \(20.5\) is constructed. One electrode compartment consists of a silver strip placed in a solution of \(\mathrm{AgNO}_{3}\), and the other has an iron strip placed in a solution of \(\mathrm{FeCl}_{2}\). The overall cell reaction is $$ \mathrm{Fe}(s)+2 \mathrm{Ag}^{+}(a q) \longrightarrow \mathrm{Fe}^{2+}(a q)+2 \mathrm{Ag}(s) $$ (a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two electrode compartments. (c) Which electrode is the anode, and which is the cathode? (d) Indicate the signs of the electrodes. (e) Do electrons flow from the silver electrode to the iron electrode, or from the iron to the silver? (f) In which directions do the cations and anions migrate through the solution?
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