(c) Given the following two reactions: \( \mathrm{MnO}_{4(\cos )}^{-}+8 \mathrm{H}^{+}(\mathrm{ap)})+5 e^{-} \rightarrow \mathrm{Mn}^{2+}((0))+4 \mathrm{H}_{2} \mathrm{O}_{(j)} \) \( \left.F e_{(\text {(n) })}^{3+}+e^{-} \rightarrow F e^{2+}(a)\right) \) (i) Construct a balanced redox equation for the reaction. (2 marks) (ii) Calculate the concentration of \( F e^{2+} \) given that \( 24.30 \mathrm{~cm}^{3} \) of \( 0.020 \mathrm{M} \mathrm{KMnO}_{4} \) reacted with exactly \( 20.00 \mathrm{~cm}^{3} \) of \( \mathrm{Fe}^{2+} \) solution \( (\mathrm{Fe}=55.9) \) (4 marks) (d) (i) A \( 1.51 \mathrm{~g} \) sample of iron wire was dissolved in excess dilute \( \mathrm{H}_{2} \mathrm{SO}_{4} \) acid and the solution made up to \( 250.0 \mathrm{~cm}^{3} \). Determine the percentage of iron in the sample given that a \( 25.00 \mathrm{~cm}^{3} \) aliquot of the prepared solution required \( 25.45 \mathrm{~cm}^{3} \) of \( 0.020 \mathrm{M} \mathrm{KMnO}_{4} \) solution for complete oxidation. (5 marks) (ii) State how the end-point in this titration can be recognized. oxtour change of eopheuponngy ( 2 marks)
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The balanced equation is: \(5\left( \mathrm{Fe}^{3+} + e^{-} \rightarrow \mathrm{Fe}^{2+} \right) \) \(1\left( \mathrm{MnO}_{4}^{-} + 8 \mathrm{H}^{+} + 5 e^{-} \rightarrow \mathrm{Mn}^{2+} + 4 \mathrm{H}_{2}\mathrm{O} \right) \) Adding these two equations, we Show more…
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Hydroxylamine reduces iron (III) according to the equation: $2 \mathrm{NH}_{2} \mathrm{OH}+4 \mathrm{Fe}^{3+} \rightarrow \mathrm{N}_{2} \mathrm{O}(\mathrm{g}) \uparrow+\mathrm{H}_{2} \mathrm{O}+4 \mathrm{Fe}^{2+}+4 \mathrm{H}^{+}$ Iron (II) thus produced is estimated by titration with a standard permanganate solution. The reaction is : $$ \mathrm{MnO}_{4}^{-}+5 \mathrm{Fe}^{2+}+8 \mathrm{H}^{+} \rightarrow \mathrm{Mn}^{2+}+5 \mathrm{Fe}^{3+}+4 \mathrm{H}_{2} \mathrm{O} $$ A $10 \mathrm{~mL}$. sample of hydroxylamine solution was diluted to 1 litre. $50 \mathrm{~mL}$. of this diluted solution was boiled with an excess of iron (III) solution. The resulting solution required $12 \mathrm{~mL}$. of $0.02 \mathrm{M} \mathrm{KMnO}_{4}$ solution for complete oxidation of iron (II). Calculate the weight of hydroxylamine in one litre of the original solution. $(\mathrm{H}=1, \mathrm{~N}=14, \mathrm{O}$ $=16, \mathrm{~K}=39, \mathrm{Mn}=55, \mathrm{Fe}=56)$ [1982 - 4 Marks]
Some Basic Concepts of Chemistry
Topic 2 : Stoichiometry, Equivalent Concept, Neutralization and Redox Titration
A quantity of $25.0 \mathrm{mL}$ of a solution containing both $\mathrm{Fe}^{2+}$ and $\mathrm{Fe}^{3+}$ ions is titrated with $23.0 \mathrm{mL}$ of $0.0200 M \mathrm{KMnO}_{4}$ (in dilute sulfuric acid). As a result, all of the $\mathrm{Fe}^{2+}$ ions are oxidized to $\mathrm{Fe}^{3+}$ ions. Next, the solution is treated with Zn metal to convert all of the $\mathrm{Fe}^{3+}$ ions to $\mathrm{Fe}^{2+}$ ions. Finally, the solution containing only the $\mathrm{Fe}^{2+}$ ions requires $40.0 \mathrm{mL}$ of the same $\mathrm{KMnO}_{4}$ solution for oxidation to $\mathrm{Fe}^{3+}$ Calculate the molar concentrations of $\mathrm{Fe}^{2+}$ and $\mathrm{Fe}^{3+}$ in the original solution. The net ionic equation is $$\mathrm{MnO}_{4}^{-}+5 \mathrm{Fe}^{2+}+8 \mathrm{H}^{+} \longrightarrow_{\mathrm{Mn}^{2+}}+5 \mathrm{Fe}^{3+}+4 \mathrm{H}_{2} \mathrm{O}$$
A 0.213 -g sample of uranyl(VI) nitrate, $\mathrm{UO}_{2}\left(\mathrm{NO}_{3}\right)_{2},$ is dissolved in $20.0 \mathrm{mL}$ of $1.0 \mathrm{M}$ $\mathrm{H}_{2} \mathrm{SO}_{4}$ and shaken with Zn. The zinc reduces the uranyl ion, $\mathrm{UO}_{2}^{2+},$ to a uranium ion, $\mathrm{U}^{n+}$. To determine the value of $n,$ this solution is titrated with $\mathrm{KMnO}_{4} .$ Permanganate is reduced to $\mathrm{Mn}^{2+}$ and $\mathrm{U}^{n+}$ is oxidized back to $\mathrm{UO}_{2}^{2+}$ (a) In the titration, $12.47 \mathrm{mL}$ of $0.0173 \mathrm{M} \mathrm{KMnO}_{4}$ was required to reach the equivalence point. Use this information to determine the charge on the ion $\mathrm{U}^{n+}$. (b) With the identity of $\mathrm{U}^{n+}$ now established, write a balanced net ionic equation for the reduction of $\mathrm{UO}_{2}^{2+}$ by zinc (assume acidic conditions). (c) Write a balanced net ionic equation for the oxidation of $\mathrm{U}^{n+}$ to $\mathrm{UO}_{2}^{2+}$ by $\mathrm{MnO}_{4}^{-}$ in acid.
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