In an electric power plant, personnel monitor the \(\mathrm{O}_{2}\) content of boiler feed water to prevent corrosion of the boiler tubes. Why does Fe corrode faster in steam and hot water than in cold water?

Short Answer

Expert verified
Fe corrodes faster in hot water and steam due to higher reaction rates and increased oxygen solubility at elevated temperatures.

Step by step solution

01

- Understand Corrosion

Corrosion occurs when iron (Fe) reacts with oxygen (O₂) and water (H₂O) to form iron oxides, commonly known as rust. This reaction is an electrochemical process that requires the presence of both oxygen and water.
02

- Temperature and Reaction Rates

Reaction rates generally increase with temperature. In hot water or steam, the kinetic energy of the molecules is higher, resulting in more frequent and energetic collisions between Fe, O₂, and H₂O molecules. This accelerates the corrosion process.
03

- Solubility of Oxygen

Hot water and steam can hold more dissolved oxygen than cold water. The increased availability of O₂ in hot environments boosts the corrosion rate of Fe. Higher temperatures also often reduce the formation of protective oxide layers on Fe, making it more susceptible to corrosion.
04

- Formulating Conclusion

In summary, Fe corrodes faster in steam and hot water because the elevated temperatures increase reaction rates and allow for greater solubility of oxygen, both of which facilitate the corrosion process.

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Key Concepts

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

Corrosion Process
Corrosion is a natural process where metals deteriorate due to chemical reactions in their environment. When iron (Fe) corrodes, it reacts with oxygen (O₂) and water (H₂O) to form rust, a type of iron oxide. This reaction requires the presence of both oxygen and water.
Corrosion is essentially an electrochemical process. Here's how it works:
  • Iron atoms lose electrons to form iron ions (Fe²⁺).
  • These electrons then react with oxygen and water to form hydroxide ions (OH⁻).
  • Finally, the iron ions combine with hydroxide ions to form iron hydroxides, which further react with oxygen to produce iron oxides (rust).
The entire process is an intricate dance between iron, water, and oxygen. Preventing corrosion involves altering this balance by, for example, reducing oxygen exposure or using corrosion inhibitors.
Reaction Rates
The rate at which reactions occur is crucial for understanding corrosion. Reaction rates typically increase with temperature. This is because higher temperatures give molecules more kinetic energy, making them move faster and collide more often.
In water and steam, the molecules are highly energetic, leading to more frequent and forceful collisions among iron, oxygen, and water molecules. This makes the electrochemical reactions that lead to corrosion happen faster.
This concept can be summarized as follows:
  • Higher temperature leads to higher kinetic energy.
  • Higher kinetic energy results in more collisions.
  • More collisions increase the reaction rate, thereby accelerating corrosion.
Essentially, the hotter the environment, the quicker the corrosion process; hence, iron corrodes faster in hot water or steam compared to cold water.
Solubility of Oxygen
Solubility refers to the amount of a substance (in this case, oxygen) that can be dissolved in a liquid (water). The solubility of oxygen in water increases with temperature up to a certain point. This means that hot water and steam can hold more dissolved oxygen than cold water.
More dissolved oxygen at higher temperatures means more oxygen is available to react with iron, speeding up the corrosion process. Besides, higher temperatures often inhibit the formation of protective oxide layers on the iron's surface, making it more vulnerable to corrosion.
Key points on solubility of oxygen:
  • Hot water holds more dissolved oxygen than cold water.
  • More dissolved oxygen accelerates the corrosion process.
  • Protective oxide layers are less likely to form at higher temperatures, increasing corrosion susceptibility.
Understanding the solubility of oxygen helps to explain why iron corrodes faster in environments with hot water or steam.

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Most popular questions from this chapter

Aqua regia, a mixture of concentrated \(\mathrm{HNO}_{3}\) and \(\mathrm{HCl}\) was developed by alchemists as a means to "dissolve" gold. The process is a redox reaction with this simplified skeleton reaction: $$\mathrm{Au}(s)+\mathrm{NO}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{AuCl}_{4}^{-}(a q)+\mathrm{NO}_{2}(g)$$ (a) Balance the reaction by the half-reaction method. (b) What are the oxidizing and reducing agents? (c) What is the function of \(\mathrm{HCl}\) in aqua regia?

Electrolysis of molten \(\mathrm{MgCl}_{2}\) is the final production step in the isolation of magnesium from seawater by the Dow process (Section 22.4). Assuming that \(45.6 \mathrm{~g}\) of Mg metal forms, (a) How many moles of electrons are required? (b) How many coulombs are required? (c) How many amps will produce this amount in \(3.50 \mathrm{~h} ?\)

A chemist designs an ion-specific probe for measuring \(\left[\mathrm{Ag}^{+}\right]\) in an \(\mathrm{NaCl}\) solution saturated with \(\mathrm{AgCl}\). One half-cell has an Ag wire electrode immersed in the unknown AgCl-saturated \(\mathrm{NaCl}\) solution. It is connected through a salt bridge to the other half-cell, which has a calomel reference electrode [a platinum wire immersed in a paste of mercury and calomel (Hg \(_{2} \mathrm{Cl}_{2}\) ) ] in a saturated KCl solution. The measured \(E_{\text {cell }}\) is \(0.060 \mathrm{~V}\). (a) Given the following standard half-reactions, calculate \(\left[\mathrm{Ag}^{+}\right]\). Calomel: \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}(s)+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Hg}(I)+2 \mathrm{Cl}^{-}(a q) \quad E^{\circ}=0.24 \mathrm{~V}\) Silver: \(\mathrm{Ag}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{Ag}(s)\) $$E^{\circ}=0.80 \mathrm{~V}$$ (Hint: Assume that [Cl \(^{-}\) ] is so high that it is essentially constant.) (b) A mining engineer wants an ore sample analyzed with the \(\mathrm{Ag}^{+}-\) selective probe. After pretreating the ore sample, the chemist measures the cell voltage as \(0.53 \mathrm{~V}\). What is \(\left[\mathrm{Ag}^{+}\right] ?\)

(a) How do the relative magnitudes of \(Q\) and \(K\) relate to the signs of \(\Delta G\) and \(E_{\text {cell }} ?\) Explain. (b) Can a cell do work when \(Q / K>1\) or \(Q / K<1 ?\) Explain. 21.53 A voltaic cell consists of \(\mathrm{A} / \mathrm{A}^{+}\) and \(\mathrm{B} / \mathrm{B}^{+}\) half-cells, where A and \(B\) are metals and the A electrode is negative. The initial \(\left[\mathrm{A}^{+}\right] /\left[\mathrm{B}^{+}\right]\) is such that \(E_{\text {cell }}>E_{\text {cell }}^{\circ}\) (a) How do \(\left[\mathrm{A}^{+}\right]\) and \(\left[\mathrm{B}^{+}\right]\) change as the cell operates? (b) How does \(E_{\text {cell }}\) change as the cell operates? (c) What is \(\left[\mathrm{A}^{+}\right] /\left[\mathrm{B}^{+}\right]\) when \(E_{\text {cell }}=E_{\text {cell }}^{\circ} ?\) Explain. (d) Is it possible for \(E_{\text {cell }}\) to be less than \(E_{\text {cell }}^{\circ} ?\) Explain.

In basic solution, \(\mathrm{Se}^{2-}\) and \(\mathrm{SO}_{3}^{2-}\) ions react spontaneously: \(2 \mathrm{Se}^{2-}(a q)+2 \mathrm{SO}_{3}^{2-}(a q)+3 \mathrm{H}_{2} \mathrm{O}(I) \longrightarrow\) $$2 \mathrm{Se}(s)+6 \mathrm{OH}^{-}(a q)+\mathrm{S}_{2} \mathrm{O}_{3}^{2-}(a q) \quad E_{\mathrm{cell}}^{\circ}=0.35 \mathrm{~V}$$ (a) Write balanced half-reactions for the process. (b) If \(E_{\text {sulfite }}^{\circ}\) is \(-0.57 \mathrm{~V},\) calculate \(E_{\text {selenium }}^{\circ}\)

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