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Chapter 19: Problem 42

How does the tendency of iron to rust depend on the \(\mathrm{pH}\) of solution?

Short Answer

Expert verified
The tendency of iron to rust increases in acidic conditions (lower pH) and decreases in basic or neutral conditions (higher pH). This is because acidic conditions facilitate the removal of electrons from iron (oxidation), promoting rust formation, whereas basic conditions slow down the rust formation by creating a protective layer on iron's surface and the low concentration of hydrogen ions slows down the oxidation process.

Step by step solution

01

Understanding pH

pH is a measure of how acidic or basic water is. It ranges from 0 to 14. A pH of 7 is neutral. A pH less than 7 is acidic. A pH greater than 7 is basic. Acidic solutions have high concentration of hydrogen ions (H+), while basic solutions have high concentration of hydroxide ions (OH−).
02

Effect of pH on rusting process

When iron reacts with oxygen and water, it forms hydrated iron(III) oxide, commonly known as rust. This reaction (corrosion process) be expressed as: \(4Fe(s) + 3O2(g) + 6H_2O(l) \rightarrow 4Fe(OH)_3(s)\), where \(Fe(s)\) is iron, \(O2(g)\) is oxygen, \(H2O(l)\) is water, and \(Fe(OH)_3(s)\) is hydrated iron(III) oxide or rust. This process is faster in acidic conditions (lower pH) because in acidic conditions, the concentration of hydrogen ions (which are positively charged) is increased and these ions promote the removal of electrons from iron (oxidation process), hence helping to form rust.
03

pH Levels and Corrosion Rate

Research studies have found out that the corrosion rate of iron is slower in neutral or basic conditions (higher pH). There are two reasons for this. Firstly, at higher pH, the concentration of hydrogen ions is lower, so the removal of electrons from iron is slower. Secondly, basic conditions often produce a protective layer on the iron's surface that inhibit further corrosion.

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Most popular questions from this chapter

A piece of magnesium ribbon and a copper wire are partially immersed in a \(0.1 \mathrm{M} \mathrm{HCl}\) solution in a beaker. The metals are joined externally by another piece of metal wire. Bubbles are seen to evolve at both the \(\mathrm{Mg}\) and \(\mathrm{Cu}\) surfaces. (a) Write equations representing the reactions occurring at the metals. (b) What visual evidence would you seek to show that \(\mathrm{Cu}\) is not oxidized to \(\mathrm{Cu}^{2+} ?\) (c) At some stage, \(\mathrm{NaOH}\) solution is added to the beaker to neutralize the HCl acid. Upon further addition of \(\mathrm{NaOH},\) a white precipitate forms. What is it?

Industrially, copper is purified by electrolysis. The impure copper acts as the anode, and the cathode is made of pure copper. The electrodes are immersed in a \(\mathrm{CuSO}_{4}\) solution. During electrolysis, copper at the anode enters the solution as \(\mathrm{Cu}^{2+}\) while \(\mathrm{Cu}^{2+}\) ions are reduced at the cathode. (a) Write half-cell reactions and the overall reaction for the electrolytic process. (b) Suppose the anode was contaminated with \(\mathrm{Zn}\) and \(\mathrm{Ag} .\) Explain what happens to these impurities during electrolysis. (c) How many hours will it take to obtain \(1.00 \mathrm{~kg}\) of \(\mathrm{Cu}\) at a current of \(18.9 \mathrm{~A} ?\)

Discuss the advantages and disadvantages of fuel cells over conventional power plants in producing electricity.

How many faradays of electricity are required to produce (a) \(0.84 \mathrm{~L}\) of \(\mathrm{O}_{2}\) at exactly \(1 \mathrm{~atm}\) and \(25^{\circ} \mathrm{C}\) from aqueous \(\mathrm{H}_{2} \mathrm{SO}_{4}\) solution; (b) \(1.50 \mathrm{~L}\) of \(\mathrm{Cl}_{2}\) at \(750 \mathrm{mmHg}\) and \(20^{\circ} \mathrm{C}\) from molten \(\mathrm{NaCl}\); (c) \(6.0 \mathrm{~g}\) of Sn from molten \(\mathrm{SnCl}_{2} ?\)

The oxidation of \(25.0 \mathrm{~mL}\) of a solution containing \(\mathrm{Fe}^{2+}\) requires \(26.0 \mathrm{~mL}\) of \(0.0250 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) in acidic solution. Balance the following equation and calculate the molar concentration of \(\mathrm{Fe}^{2+}\): $$ \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+\mathrm{Fe}^{2+}+\mathrm{H}^{+} \longrightarrow \mathrm{Cr}^{3+}+\mathrm{Fe}^{3+} $$
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