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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: \[ E掳_{Fe^{2+}/Fe} = -0.44V \] \[ E掳_{Fe^{3+}/Fe} = +0.04V \] \[ E掳_{Sn^{2+}/Sn} = -0.14V \]
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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Key Concepts

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

Standard Oxidation Potential
The standard oxidation potential is crucial in evaluating how likely a metal will lose electrons, thereby corroding. Essentially, it measures the readiness of a metal to oxidize.
In electrochemistry, the more negative the standard oxidation potential of a metal, the higher its tendency to undergo oxidation. For example, in the given data:
  • The standard oxidation potential of iron (Fe) when it transitions from Fe虏鈦� to its solid form is \( E^\circ_{Fe^{2+}/Fe} = -0.44V \).
  • Meanwhile, tin (Sn) goes from Sn虏鈦� to Sn solid with a standard oxidation potential of \( E^\circ_{Sn^{2+}/Sn} = -0.14V \).
With iron's oxidation potential being more negative than tin's, iron oxidizes, or corrodes, more readily than tin.
This is a key consideration in corrosion studies and protective measures.
Corrosion Prevention
Corrosion prevention is a vital aspect in prolonging the life of metal objects. It involves protecting metals from adverse environmental conditions that lead to wear and degradation through rusting or oxidation.
When it comes to iron, one way to prevent corrosion is by coating it with another metal, like tin. Although tin doesn't protect iron through cathodic protection, it serves as a physical shielding layer.
Here's how tin acts:
  • It doesn't corrode as easily as iron because of its less negative standard oxidation potential.
  • The tin coating forms a barrier blocking moisture and oxygen, which are two primary causes of corrosion.
  • Even if there's a scratch on the tin layer, it tends to self-heal to continuously protect the underlying iron.
Nonetheless, regular maintenance and inspections are necessary to ensure the protective layer remains intact.
Electrochemistry
Electrochemistry plays a pivotal role in understanding metal corrosion and protection strategies. This scientific field deals with the movement of electrons and the electrical properties of chemicals.
A core component of electrochemistry is the electrochemical series, which ranks metals by their standard oxidation potentials.
  • The series helps predict and understand how different metals interact with each other when in contact.
  • It is instrumental in designing protection systems involving cathodic protection or protective coatings.
By studying electrochemical interactions, scientists can devise methods to prevent unnecessary metal wear, ensuring structures and components last longer.
In practical applications, tools like galvanic cells highlight how metal interactions can produce electrical energy, while underlying principles aid in innovative corrosion protection solutions.

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Most popular questions from this chapter

(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?

At \(298 \mathrm{~K}\) a cell reaction has a standard cell potential of \(+0.63 \mathrm{~V}\). The equilibrium constant for the reaction is \(3.8 \times 10^{10}\). What is the value of \(n\) for the reaction?

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 \(\mathrm{CyFe}^{2+}\) $$ \begin{aligned} \mathrm{O}_{2}(g)+4 \mathrm{H}^{+}(a q)+4 \mathrm{e}^{-} & \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) & & E_{\mathrm{red}}^{\circ}=+0.82 \mathrm{~V} \\\ \mathrm{CyFe}^{3+}(a q)+\mathrm{e}^{-} & \longrightarrow \mathrm{CyFe}^{2+}(a q) & E_{\mathrm{red}}^{\circ} &=+0.22 \mathrm{~V} \end{aligned} $$ (a) What is \(\Delta G\) for the oxidation of \(\mathrm{CyFe}^{2+}\) by air? \((\mathbf{b})\) If the synthesis of \(1.00 \mathrm{~mol}\) 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) Write the half-reaction that occurs at an oxygen electrode in acidic aqueous solution when it serves as the cathode of a voltaic cell. (b) Write the half-reaction that occurs at an oxygen electrode in acidic aqueous solution when it serves as the anode of a voltaic cell. (c) What is standard about the standard oxygen electrode?

Hydrazine \(\left(\mathrm{N}_{2} \mathrm{H}_{4}\right)\) and dinitrogen tetroxide \(\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)\) form a self-igniting mixture that has been used as a rocket propellant. The reaction products are \(\mathrm{N}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\). (a) Write a balanced chemical equation for this reaction. (b) What is being oxidized, and what is being reduced? (c) Which substance serves as the reducing agent and which as the oxidizing agent?
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