An acidified solution of iron(II) ions is titrated with a standardized solution of…

CHEMISTRY · Multi-part · A Level · Chemistry · Oxidation, reduction and redox reactions · Oxidation, reduction and redox reactions

Problem

An acidified solution of iron(II) ions is titrated with a standardized solution of cerium(IV) ions. The titration reaction converts

\text{Fe}^{2+}

\text{Fe}^{3+}

while cerium is reduced from

\text{Ce}^{4+}

\text{Ce}^{3+}

. During the titration, the deep yellow color of

\text{Ce}^{4+}

fades, making it possible to detect the endpoint without an external indicator. (a) Write the balanced ionic equation for the reaction that occurs in the titration medium (acidic solution). Use the lowest possible whole-number coefficients. (b) From your balanced equation, determine the number of moles of electrons transferred per mole of

\text{Ce}^{4+}

that reacts. (c) In a particular analysis,

22.60\,\text{mL}

0.100\,\text{M}

\text{Ce}^{4+}

solution is required to titrate

25.00\,\text{mL}

of the unknown

\text{Fe}^{2+}

solution. Calculate the molarity of

\text{Fe}^{2+}

in the original sample. (d) Explain why an external redox indicator is unnecessary for this titration, referring to observable changes that accompany the redox process.

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