Answer to Question #319977 in General Chemistry for joshuajr

Question

. 2.83 g of a sample of haematite iron ore [iron (III) oxide, Fe₂O₃] were dissolved in concentrated hydrochloric acid and the solution diluted to 250 cm³. 25.0 cm³ of this solution was reduced with tin(II) chloride (which is oxidised to Sn⁴⁺ in the process) to form a solution of iron(II) ions. This solution of iron(II) ions required 26.4 cm³ of a 0.0200 mol/dm³ potassium dichromate(VI) solution for complete oxidation back to iron(III) ions.

(a) given the half–cell reactions 

(i) Sn⁴⁺_(aq) + 2e^– ==> Sn²⁺_(aq)

and (ii) Cr₂O₇^2–_(aq) + 14H⁺_(aq) + 6e^– ==> 2Cr³⁺_(aq) + 7H₂O_(l)

deduce the fully balanced redox equations for the reactions

(i) the reduction of iron(III) ions by tin(II) ions

(ii) the oxidation of iron(II) ions by the dichromate(VI) ion

(b) Calculate the percentage of iron(III) oxide in the ore. 

Answer 1

Ratio 2.03:6.40

PCH4= (2.03/8.43) × 648= 156mmHg

PCO2= 648 - 156= 492mmHg