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Chapter 20: Problem 87

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

Short Answer

Expert verified
The spontaneous redox reaction that might cause corrosion between a brass pipe and a galvanized steel pipe is \(Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)\). This reaction occurs because zinc is more likely to be oxidized and copper is more likely to be reduced, based on their standard reduction potentials. The calculated standard emf for the reaction is 1.10 V, which indicates that the reaction is spontaneous and will cause corrosion. To prevent this corrosion, an insulating fitting should be used to connect the two pipes, as suggested in the plumber's handbook.

Step by step solution

01

Identify the half-reactions for copper and zinc

To identify the half-reactions for copper and zinc, we have to look into the standard reduction potential table: Copper (Cu): \(Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)\), standard reduction potential, \(E^\circ_{Cu}\) = +0.34 V Zinc (Zn): \(Zn^{2+}(aq) + 2e^- \rightarrow Zn(s)\), standard reduction potential, \(E^\circ_{Zn}\) = -0.76 V Now that we have the possible half-reactions, let's analyze the spontaneous redox reactions.
02

Determine the overall redox reaction

A redox reaction occurs when one element is oxidized while the other is reduced. From the standard reduction potentials, it can be concluded that zinc is more likely to be oxidized due to its lower standard reduction potential, while copper is more likely to be reduced because of its higher standard reduction potential. When one element is oxidized, it loses electrons and when an element is reduced, it gains electrons. So, the overall redox reaction is: Zinc is oxidized: \(Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-\) Copper is reduced: \(Cu^{2+} + 2e^- \rightarrow Cu(s)\) Combining both, we get: \(Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)\)
03

Calculate the standard emf for the reaction

To calculate the standard emf (\(E^\circ_{cell}\)) for the reaction, we'll use the formula: \(E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}\) In our case, the cathode is the Cu electrode, which accepts electrons, and the anode is the Zn electrode, which donates electrons. Therefore, the standard emf for the reaction is: \(E^\circ_{cell} = E^\circ_{Cu} - E^\circ_{Zn} = 0.34 V - (-0.76 V) = 1.10 V\) Since the standard emf of the cell is positive (\(E^\circ_{cell} > 0\)), the redox reaction between copper and zinc is spontaneous and will cause corrosion. #Conclusion# The spontaneous redox reaction between copper in the brass pipe and galvanized steel pipe will cause corrosion. The reaction is \(Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)\) and has a standard emf of 1.10 V. Therefore, we should use an insulating fitting to connect the two pipes, as recommended in the handbook.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Electrochemical Cells
Understanding electrochemical cells is crucial when exploring how different metals interact and can lead to corrosion. An electrochemical cell comprises two metal electrodes immersed in electrolyte solutions that conduct electricity through ion flow. These cells are integral in batteries, galvanic corrosion, and other processes that involve electron transfer between materials.

In the context of the exercise about galvanized steel and brass pipes, the direct physical connection facilitates the setup of a makeshift electrochemical cell, where zinc from the galvanized layer and copper from the brass act as anode and cathode, respectively. The moist environment surrounding the pipes serves as an electrolyte, which allows for ionic movement and electron flow from the anode to the cathode, setting the stage for corrosion through oxidation and reduction reactions.
Standard Reduction Potential
The standard reduction potential is a measure that indicates the tendency of a chemical species to gain electrons and hence be reduced. It is crucial for predicting the direction of redox reactions and is measured in volts (V). Each half-reaction has a standard reduction potential associated with it, and when comparing two potential reactions, the one with the higher potential acts as the cathode (reduction) and the other as the anode (oxidation).

For the brass and galvanized steel example, zinc has a lower standard reduction potential (\(E^\text{o}_{\text{Zn}} = -0.76\text{ V}\)) compared to copper's higher potential (\(E^\text{o}_{\text{Cu}} = 0.34\text{ V}\)), indicating that in the presence of an electrolyte, zinc will preferentially oxidize while copper will reduce, leading to the corrosion of zinc.
Corrosion Prevention
Corrosion prevention is essential to maintain the integrity of metal structures and is achieved through various methods. In pipes, where dissimilar metals come into contact, an insulating fitting can prevent the establishment of an electrochemical cell, thus eliminating the electrochemical reaction that would cause corrosion.

Other methods include coating the metals with protective layers to prevent electrode reactions, using sacrificial anodes that preferentially corrode, or applying corrosion inhibitors which can adsorb on metal surfaces and block the reactive sites. Regular maintenance and material selection, based on understanding the galvanic series, also play a significant role in preventing corrosion.
Galvanic Series
The galvanic series is a list of metals and alloys organized according to their standard reduction potentials in a given environment. Metals at the top have a greater tendency to lose electrons and corrode, acting as anodes, while those at the bottom are more likely to gain electrons and be protected, acting as cathodes.

By consulting the galvanic series, plumbers and engineers can predict which metal combinations are likely to produce galvanic corrosion. The series explains why zinc (closer to the top) would corrode when in contact with copper (closer to the bottom). It underpins the advice in the plumber's handbook, advising against direct contact between brass and galvanized steel to prevent corrosion.

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Most popular questions from this chapter

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

Predict whether the following reactions will be spontaneous in acidic solution under standard conditions: (a) oxidation of \(\mathrm{Sn}\) to \(\mathrm{Sn}^{2+}\) by \(\mathrm{I}_{2}\) (to form \(\mathrm{I}^{-}\) ), (b) reduction of \(\mathrm{Ni}^{2+}\) to \(\mathrm{Ni}\) by \(\mathrm{I}^{-}\) (to form \(\mathrm{I}_{2}\) ), (c) reduction of \(\mathrm{Ce}^{4+}\) to \(\mathrm{Ce}^{3+}\) by \(\mathrm{H}_{2} \mathrm{O}_{2},\) (d) reduction of \(\mathrm{Cu}^{2+}\) to \(\mathrm{Cu}\) by \(\mathrm{Sn}^{2+}\) (to form \(\mathrm{Sn}^{4+}\) ).

A voltaic cell similar to that shown in Figure 20.5 is constructed. One electrode half-cell 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 half-cells. (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?

A voltaic cell similar to that shown in Figure 20.5 is constructed. One half- cell consists of an aluminum strip placed in a solution of \(\mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3},\) and the other has a nickel strip placed in a solution of \(\mathrm{NiSO}_{4}\). The overall cell reaction is $$2 \mathrm{Al}(s)+3 \mathrm{Ni}^{2+}(a q) \longrightarrow 2 \mathrm{Al}^{3+}(a q)+3 \mathrm{Ni}(s)$$ (a) What is being oxidized, and what is being reduced? (b) Write the half-reactions that occur in the two half-cells. (c) Which electrode is the anode, and which is the cathode? (d) Indicate the signs of the electrodes. (e) Do electrons flow from the aluminum electrode to the nickel electrode or from the nickel to the aluminum? (f) In which directions do the cations and anions migrate through the solution? Assume the \(\mathrm{Al}\) is not coated with its oxide.

(a) Which electrode of a voltaic cell, the cathode or the anode, corresponds to the higher potential energy for the electrons? (b) What are the units for electrical potential? How does this unit relate to energy expressed in joules? (c) What is special about a standard cell potential?
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