Chemistry 9701 · AS & A Level · Analytical techniques

Analytical techniques — practice question

The most frequently encountered form of iron(II) sulfate is the heptahydrate, $FeSO_4\cdot7H_2O$. When heated to $90^\circ\text{C}$, it loses part of its water of crystallisation and changes into a different hydrated form of iron(II) sulfate, $FeSO_4\cdot xH_2O$. A $3.40\,\text{g}$ sample of $FeSO_4\cdot xH_2O$ was dissolved in water to make $250\,\text{cm}^3$ of solution. A $25.0\,\text{cm}^3$ portion of this solution was acidified and titrated with $0.0200\,\text{mol dm}^{-3}$ potassium manganate(VII). During the titration, $20.0\,\text{cm}^3$ of the potassium manganate(VII) solution was needed to react completely with the $Fe^{2+}$ ions in the sample.
(a)

In the acidified solution, the $MnO_4^-$ ions in the potassium manganate(VII) oxidise the $Fe^{2+}$ ions.

(a(i))[1]

Explain, in terms of electron transfer, what the term oxidise means in the sentence above.

(a(ii))[3]

Complete and balance the ionic equation for the reaction of the manganate(VII) ions with the iron(II) ions: $MnO_4^-(aq) + 5Fe^{2+}(aq) + \ldots H^+(aq) \rightarrow \ldots (aq) + 5Fe^{3+}(aq) + \ldots H_2O(l)$

(b(i))[1]

Calculate the amount of manganate(VII), in moles, used in the titration.

(b(ii))[1]

Use the equation in part (a)(ii) together with your answer to part (b)(i) to calculate the amount, in moles, of $Fe^{2+}$ in the $25.0\,\text{cm}^3$ sample of solution used.

(b(iii))[1]

Calculate the amount of $FeSO_4\cdot xH_2O$, in moles, in $3.40\,\text{g}$ of the compound.

(b(iv))[1]

Calculate the relative formula mass of the compound $FeSO_4\cdot xH_2O$.

(b(v))[1]

The relative formula mass of anhydrous iron(II) sulfate, $FeSO_4$, is $151.8$. Calculate the value of $x$ in $FeSO_4\cdot xH_2O$.

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