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Chapter 20: Problem 9

How does a zinc coating on iron protect the iron from unwanted oxidation? [Section 20.8]

Short Answer

Expert verified
A zinc coating on iron protects the iron from unwanted oxidation by acting as a sacrificial anode, undergoing oxidation preferentially due to its higher reactivity compared to iron. This process, called cathodic protection, leads to the formation of a stable, protective layer of zinc oxide on the surface of the iron, preventing the iron from rusting. The application of the protective zinc coating to iron is known as galvanization.

Step by step solution

01

Understand oxidation process

Oxidation occurs when a material loses electrons to another material, often involving oxygen. In the case of iron, when it comes into contact with oxygen and moisture, it undergoes oxidation, forming iron oxide (rust). This can lead to structural damage and degradation of the iron as the rust flakes off, exposing more iron to the oxidation process.
02

Role of zinc in the oxidation process

Zinc is a more reactive metal compared to iron. Thus, when zinc is used as a coating on iron, it becomes the primary target for oxidation. Zinc undergoes oxidation by losing electrons to the oxygen and forming zinc oxide, which acts as a stable and protective layer on the surface of the iron.
03

Cathodic protection of iron by zinc

When zinc is used as a coating on iron, it acts as a sacrificial anode, a process called cathodic protection. The zinc coating willingly undergoes oxidation, protecting the iron beneath it from being oxidized. This is because zinc, being more reactive, has a greater tendency to lose electrons and form positive ions, as compared to iron. Thus, oxygen and other oxidizing agents will preferentially react with the zinc coating instead of the underlying iron surface.
04

Explain galvanization

Galvanization is the process of applying a protective zinc coating to iron or steel to prevent rusting. This process can be done in various ways, most commonly by hot-dip galvanizing, which involves immersing the iron or steel into molten zinc to form a zinc coating on the surface. This layer serves to protect the iron from the unwanted oxidation, ensuring the structural integrity of the iron remains intact.
05

Summarize the key points

In summary, a zinc coating on iron protects the iron from unwanted oxidation because zinc is more reactive than iron and undergoes oxidation preferentially. This process, known as cathodic protection, preserves the integrity of the underlying iron by creating a stable, protective layer of zinc oxide at the surface. The application of this protective zinc coating to iron is called galvanization.

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Most popular questions from this chapter

(a) What is meant by the term reduction? (b) On which side of a reduction half-reaction do the electrons appear? (c) What is meant by the term reductant? (d) What is meant by the term reducing agent?

The capacity of batteries such as the typical AA alkaline battery is expressed in units of milliamp-hours (mAh). An "AA" alkaline battery yields a nominal capacity of 2850 mAh. (a) What quantity of interest to the consumer is being expressed by the units of \(\mathrm{mAh}\) ? (b) The starting voltage of a fresh alkaline battery is \(1.55 \mathrm{~V}\). The voltage decreases during discharge and is \(0.80 \mathrm{~V}\) when the battery has delivered its rated capacity. If we assume that the voltage declines linearly as current is withdrawn, estimate the total maximum electrical work the battery could perform during discharge.

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Cytochrome, a complicated molecule that we will represent as \(\mathrm{CyFe}^{2+}\), reacts with the air we breathe to supply energy required to synthesize adenosine triphosphate (ATP). The body uses ATP as an energy source to drive other reactions. (Section 19.7) At \(\mathrm{pH} 7.0\) the following reduction potentials pertain to this oxidation of \(\mathrm{CyFe}^{2+}\) : $$ \begin{aligned} \mathrm{O}_{2}(\mathrm{~g})+4 \mathrm{H}^{+}(a q)+4 \mathrm{e}^{-}--\rightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) & E_{\mathrm{red}}^{\mathrm{o}}=+0.82 \mathrm{~V} \\ \mathrm{CyFe}^{3+}(a q)+\mathrm{e}^{-}--\rightarrow \mathrm{CyFe}^{2+}(a q) & E_{\mathrm{red}}^{\circ}=+0.22 \mathrm{~V} \end{aligned} $$ (a) What is \(\Delta G\) for the oxidation of \(\mathrm{CyFe}^{2+}\) by air? (b) If the synthesis of \(1.00\) mol of ATP from adenosine diphosphate (ADP) requires a \(\Delta G\) of \(37.7 \mathrm{~kJ}\), how many moles of ATP are synthesized per mole of \(\mathrm{O}_{2}\) ?

(a) What is a standard reduction potential? (b) What is the standard reduction potential of a standard hydrogen electrode?
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