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Chapter 18: Problem 117

Which of the following describes the process of galvanization that protects steel from rusting? (a) Steel is coated with a layer of paint. (b) Iron in steel is oxidized to form a protective oxide coating. (c) Steel is coated with zinc because zinc is more easily oxidized than iron. (d) A strip of magnesium is attached to steel because the magnesium is more easily oxidized than iron.

Short Answer

Expert verified
Option (c) is correct: "Steel is coated with zinc because zinc is more easily oxidized than iron."

Step by step solution

01

Understanding Galvanization

Galvanization is a process used to protect steel from corrosion and rusting. It involves coating the steel with a metal that prevents oxidation.
02

Identifying the Metals Used

The most common metal used in galvanization is zinc. Zinc acts as a protective layer on the surface of steel.
03

Analyzing the Options

We need to determine which option best describes galvanization. Option (c) states that steel is coated with zinc because zinc is more easily oxidized than iron, which matches the description of galvanization.
04

Eliminating Incorrect Options

Let's briefly check why the other options are incorrect. (a) mentions paint, which does not involve a metal coating. (b) involves oxidation of iron, which does not describe galvanization. (d) involves magnesium, which is not the typical method for galvanization.
05

Confirming the Correct Option

Option (c) correctly describes the process of galvanization in which steel is coated with zinc, protecting it from rusting by using zinc's higher tendency to oxidize.

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Key Concepts

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

Corrosion Protection
Corrosion occurs when metals like steel react with environmental elements, leading to rust and deterioration. This process can weaken the metal over time, making it vital to implement protective measures. Galvanization is a critical technique used for corrosion protection by shielding steel with a more reactive metal.
In this process, the metal coating serves as a barrier, preventing water, oxygen, and other chemicals from making contact with the underlying steel.
  • The coating protects the steel physically by sealing it off from the environment.
  • It also serves as a sacrificial layer that corrodes before the steel, extending the life of metal structures.
Overall, effective corrosion protection is essential for prolonging the lifespan of steel products, used widely in construction and manufacturing industries.
Zinc Coating
Zinc is typically used in galvanization due to its unique protective properties. It forms a thin layer on the steel surface, known as a zinc coating, which is fundamental for safeguarding against rust.
This protective action occurs because zinc has a lower reduction potential compared to iron.
  • The zinc coating acts as a physical shield, but there's more to its role.
  • Its ability to undergo oxidation before iron allows it to offer galvanic protection to the steel beneath.
This means that even if scratches or damage expose part of the steel, zinc will corrode preferentially, effectively keeping the steel from rusting itself. Thus, zinc coatings are highly valued in industries that require durable, long-lasting metal surfaces.
Oxidation Process
The oxidation process is central to understanding how galvanization protects metal. Oxidation involves the loss of electrons when a metal is exposed to elements like oxygen and water.
In the context of galvanization, the more active metal - zinc, is deliberately oxidized first. This happens because zinc readily loses electrons compared to iron, which is the primary component of steel.
  • This electron loss results in the formation of zinc oxide, a barrier that immensely slows down further oxidation.
  • By sacrificing itself, zinc provides what is known as cathodic protection to the steel.
Thus, the oxidation process in galvanization is designed to control and redirect corrosion, protecting the structural integrity of steel.

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Most popular questions from this chapter

It has recently been reported that porous pellets of \(\mathrm{TiO}_{2}\) can be reduced to titanium metal at the cathode of an electrochemical cell containing molten \(\mathrm{CaCl}_{2}\) as the electrolyte. When theTiO \(_{2}\) is reduced, the \(\mathrm{O}^{2-}\) ions dissolve in the \(\mathrm{CaCl}_{2}\) and are subsequently oxidized to \(\mathrm{O}_{2}\) gas at the anode. This approach may be the basis for a less expensive process than the one currently used for producing titanium. (a) Label the anode and cathode, and indicate the signs of the electrodes. (b) Indicate the direction of electron and ion flow. (c) Write balanced equations for the anode, cathode, and overall cell reactions.

Predict the anode, cathode, and overall cell reactions when an aqueous solution of each of the following salts is electrolyzed in a cell having inert electrodes: (a) \(\mathrm{Ag}_{2} \mathrm{SO}_{4}\) (b) \(\mathrm{Ca}(\mathrm{OH})_{2}\) (c) KI

A steam-hydrocarbon reforming process is one method for producing hydrogen from fossil fuels for use in a fuel cell. In the first step, steam reacts with hydrocarbons, such as \(\mathrm{CH}_{4}\), at high temperatures in the presence of a catalyst, yielding \(\mathrm{H}_{2}\) and CO. In the second step, the reaction of \(\mathrm{CO}\) and \(\mathrm{H}_{2} \mathrm{O}\), called the water- gas shift reaction, removes toxic carbon monoxide and produces more hydrogen. Step \(1: \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{CH}_{4}(g) \stackrel{1100^{\circ} \mathrm{C} \mathrm{N}_{1} \text { cathmst }}{\longrightarrow} \mathrm{CO}(\mathrm{g})+3 \mathrm{H}_{2}(g)\) Step 2: \(\mathrm{CO}(g)+\mathrm{H}_{2} \mathrm{O}(g) \stackrel{400^{\circ} \mathrm{C}, \text { cualyst }}{\longrightarrow} \mathrm{CO}_{2}(g)+\mathrm{H}_{2}(g)\) (a) In Step 1 , which element is oxidized and which is reduced? (b) In Step 2, which element is oxidized and which is reduced? What is the oxidizing agent and reducing agent? (c) What are drawbacks of steam reforming in the production of hydrogen?

The nickel-iron battery has an iron anode, an \(\mathrm{NiO}(\mathrm{OH})\) cathode, and a KOH electrolyte. This battery uses the following halfreactions and has an \(E^{\circ}\) value of \(1.37 \mathrm{~V}\) at \(25^{\circ} \mathrm{C}:\) $$ \begin{gathered} \mathrm{Fe}(s)+2 \mathrm{OH}(a q) \longrightarrow \mathrm{Fe}(\mathrm{OH})_{2}(s)+2 \mathrm{e}^{-} \\ \mathrm{NiO}(s)+\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{e}^{-} \longrightarrow \mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{OH}(a q) \end{gathered} $$ (a) Write a balanced equation for the cell reaction. (b) Calculate \(\Delta G^{\circ}\) (in kilojoules) and the equilibrium constant \(K\) for the cell reaction at \(25^{\circ} \mathrm{C}\). (c) What is the cell voltage at \(25^{\circ} \mathrm{C}\) when the concentration of \(\mathrm{KOH}\) in the electrolyte is \(5.0 \mathrm{M} ?\) (d) How many grams of \(\mathrm{Fe}(\mathrm{OH})_{2}\) are formed at the anode when the battery produces a constant current of \(0.250 \mathrm{~A}\) for \(40.0 \mathrm{~min}\) ? How many water molecules are consumed in the process?

A constant current of \(100.0 \mathrm{~A}\) is passed through an electrolytic cell having an impure copper anode, a pure copper cathode, and an aqueous \(\mathrm{CuSO}_{4}\) electrolyte. How many kilograms of copper are refined by transfer from the anode to the cathode in a \(24.0 \mathrm{~h}\) period?
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