2.83g of a sample of haematite iron ore [iron(III) oxide Fe2o3] was dissolved in concentrated hydrochloric acid and the solution diluted to 250cm3. 25.00cm3 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 required 26.4 cm3 of 0.02 mol dm-3 potassium dichromate (VI) for oxidation.Given the half-cell reactions(i) Sn4+ and 2e- lt;---gt; Sn2+ AND (ii) Cr2O7 2- + 14H+ + 6e- lt;---gt; 2Cr3+ + 7H20deduce the fully balanced redox equations for the reactions(i) the reduction of iron(II) ions by tin (II) ions AND(ii) the oxidation of iron(II) ions by the dichromatre (VI) ion(b) calculate the percentage of iron (III) oxide in the ore i know thats allot but i raelly could use the help
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Chemical analysis 2.83g iron ore [iron(III) oxide Fe2O3] in 250cm3 25.00cm3 reduced with tin(II) to form iron(II) 26.4 cm3 of 0.02 mol dm-3 K2Cr2O6 for Fe+2 to Fe+3 ai) 2Fe+3 + Sn+2 ----- 2Fe+2 + Sn+4 (Fe+3 --1e and Sn+2--2e) aii) 6Fe+2 + Cr2O7 + 14H+---- 6Fe+3 + 2Cr+3 + 7H2O (Fe+2--1e and Cr--3e) b) g Fe+3 in 250cm3 10 aliquots * (6 mole Fe+3/mole Cr2O7) * (0.02 mole Cr2O7/1000cm3 * 26.4 cm3/aliquot) * atwt Fe g/mole ? % 100 * g Fe+3/2.83g sample Plug and SOLVE Basic mathematics is a prerequisite to chemistry – I just try to help you with the methodology of solving the problem.
