Mr Cole Chemistry

Redox Titrations Questions

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  1. A student is asked to determine the percentage purity of a sample of impure iron. They dissolve 2.50 g of the impure iron sample in excess dilute sulfuric acid. The resulting solution is then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII). In the titration, 24.50 cm3 of the potassium manganate(VII) solution is required to reach the end point. Determine the percentage purity of the iron sample.
  2. A student is asked to determine the percentage purity of a sample of impure iron (Fe). They dissolve 1.80 g of the impure iron sample in excess dilute sulfuric acid. The resulting solution is made up to 250 cm3 in a volumetric flask with distilled water. 25.0 cm3 portions of this solution are then titrated against a 0.0150 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below. Calculate the percentage purity of the iron sample.
  3. A student is asked to determine the value of X in the hydrated iron(II) sulfate compound, FeSO4⋅XH2O. They dissolve 6.64 g of the hydrated compound in distilled water and make the solution up to 250 cm3 in a volumetric flask. 25.0 cm3 portions of this solution are then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below. Determine the value of X in FeSO4⋅XH2O. X is an integer.
  4. A student is asked to determine the value of X in the hydrated sodium ethanedioate (sodium oxalate) compound, Na2C2O4⋅XH2O. They dissolve 3.50 g of the hydrated compound in distilled water and make the solution up to 250 cm3 in a volumetric flask. 25.0 cm3 portions of this solution are then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below:
  5. A scientist conducts an experiment to determine the concentration of a solution of NaClO. A 25cm3 sample of the solution is reacted with excess KI. The resulting products are reacted against 0.1moldm-3 sodium thiosulfate, where it required 52.0cm3 to react.
     
    1. Write the equation for the reaction between chlorate ions and iodide ions, given their electrode potentials in the table above.
    2. Write the equation for the reaction between iodine and thiosulfate, given their electrode potentials in the table above.
    3. Calculate the concentration of the sodium chlorate(I)
  6. A student is given a sample of impure iron(II) ethanedioate, FeC₂O₄, and is asked to determine the mass of pure FeC₂O₄ in the sample. The student dissolves 2.80 g of the sample in distilled water and makes up the solution to 250.0 cm³ in a volumetric flask. The student then titrates 25.0 cm³ of this solution against 0.0150 mol dm⁻³ potassium manganate(VII) (KMnO₄) in acidic conditions. Calculate the percentage purity of the sample of iron ethanedioate.
  7. A student is given a sample of hydrated copper(II) sulfate, CuSO₄·xH₂O, and is asked to determine the value of x (the number of moles of water of crystallisation). To do this, the student dissolves 2.50 g of the sample in distilled water and makes up the solution to 250.0 cm³ in a volumetric flask. A 25.0 cm³ aliquot of this solution is taken and titrated with 0.0100 mol dm⁻³ EDTA (ethylenediaminetetraacetic acid) solution, finding that 27.40cm3 was required to fully react with the Cu2+ ions.
  8. A 25.0 cm³ sample of hydrogen peroxide solution was titrated against a 0.0200 mol dm⁻³ solution of potassium permanganate in acidic conditions. The titration required 18.5 cm³ of potassium permanganate to reach the endpoint. Determine the concentration of the hydrogen peroxide solution in mol dm⁻³.
  9. A student conducts an experiment to determine the percentage of copper in an alloy. The student:

    – Reacts 1.25 g of the alloy with concentrated nitric acid to form a solution (all of the copper in the alloy reacts to form aqueous copper(II) ions)
    – Pours the solution into a volumetric flask and makes the volume up to 500 cm³ with distilled water
    – Shakes the flask thoroughly
    – Transfers 50.0 cm³ of the solution into a conical flask and adds an excess of potassium iodide
    – It requires exactly 12.50 cm³ of 0.100 mol dm⁻³ sodium thiosulfate (Na₂S₂O₃) solution to react with all the iodine produced.

    The equations for the reactions are:
    – 2 Cu²⁺ + 4 I⁻ → 2 CuI + I₂
    – 2 S₂O₃²⁻ + I₂ → 2 I⁻ + S₄O₆²⁻

    Calculate the percentage of copper by mass in the alloy.
  10. A student conducts an experiment to determine the percentage of zinc in a brass sample. The student:

    – Reacts 1.20 g of the brass sample with concentrated sulfuric acid to form a solution (all of the zinc in the brass reacts to form aqueous zinc ions, Zn²⁺)
    – Pours the solution into a volumetric flask and makes the volume up to 250 cm³ with distilled water
    – Shakes the flask thoroughly
    – Transfers 50.0 cm³ of the solution into a conical flask and adds an excess of sodium hydroxide to precipitate zinc hydroxide (Zn(OH)₂)
    – Dissolves the precipitate in hydrochloric acid and titrates the resulting solution with 0.100 mol dm⁻³ EDTA solution, using exactly 18.50 cm³ to react with all the zinc ions.

    Calculate the percentage of zinc by mass in the brass sample.
  11. A solid mixture contains ethanedioic acid (H₂C₂O₄), sodium ethanedioate (Na₂C₂O₄) and an inert solid. To determine the amount of ethanedioic acid and sodium ethanedioate in the mixture, the following procedure was carried out. Determine the mass (in grams) of ethanedioic acid and sodium ethanedioate in the original 2.50 g solid mixture.
    • A 2.50 g sample of the solid mixture was dissolved in 250 cm³ of water to make a solution.
    • A 25.0 cm³ portion of this solution was titrated with 0.100 mol dm⁻³ sodium hydroxide (NaOH) solution. The end-point was reached after 12.5 cm³ of NaOH was added.
    • A separate 25.0 cm³ portion of the solution was titrated with 0.0200 mol dm⁻³ potassium manganate(VII) (KMnO₄) solution in acidic conditions. The end-point was reached after 15.0 cm³ of KMnO₄ was added.
  12. A 1.664g sample of vanadium(V) oxide (V₂O₅) is dissolved in 250cm3 of dilute sulfuric acid to form a solution containing vanadium(V) ions, VO2+. This solution is then reduced by adding a metal powder, resulting in a solution containing vanadium ions in a lower oxidation state. The reaction is allowed to proceed until completion, and the excess metal is removed by filtration.

    A 25cm3 sample of the vanadium ions present in the solution are titrated against acidified 0.05 mol dm-3 potassium manganate where it required 22.0cm3 KMnO4 to fully react. What was the oxidation state of the vanadium after it reacted with the metal.
  13. A container of H2O2 was tested by a scientist to determine whether it had started to decompose. A 1.00 cm3 sample of the hydrogen peroxide was taken from the container and 24.0 cm3 of water were added to it. This solution was reduced by reacting with potassium iodide. The iodine formed in the reduced back to iodide by reacting it against 2.00 moldm-3 sodium thiosulfate, where it required 42.6 cm3 of sodium thiosulfate to react. In the reaction, the thiosulfate (S2O32-) is oxidised to S4O62-.
    1. Write an equation for the reduction of hydrogen peroxide.
    2. Write an equation for the oxidation of iodide.
    3. Write an equation for the reaction between hydrogen peroxide and iodide.
    4. Write an equation for the oxidation of S2O32- to S4O62-.Write an equation for the reaction between iodide and sodium thiosulfate.
  1. A student is asked to determine the percentage purity of a sample of impure iron. They dissolve 2.50 g of the impure iron sample in excess dilute sulfuric acid. The resulting solution is then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII). In the titration, 24.50 cm3 of the potassium manganate(VII) solution is required to reach the end point. Determine the percentage purity of the iron sample.
    Moles of KMnO4​ used: 4.90×10−4 mol
    Moles of Fe2+ ions in 25.0 cm3 portion: 2.45×10−3 mol
    Total moles of Fe2+ ions in 250 cm3 solution: 2.45×10−2 mol
    Mass of pure iron in the sample: 1.37 g
    Percentage purity of the iron sample: 54.8%
  2. A student is asked to determine the percentage purity of a sample of impure iron (Fe). They dissolve 1.80 g of the impure iron sample in excess dilute sulfuric acid. The resulting solution is made up to 250 cm3 in a volumetric flask with distilled water. 25.0 cm3 portions of this solution are then titrated against a 0.0150 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below. Calculate the percentage purity of the iron sample.

    Concordant results and average titre: 25.825 cm3
    Moles of KMnO4 used: 3.87×10−4 mol
    Moles of Fe2+ ions in 25.0 cm33 portion: 1.935×10−3 mol
    Total moles of Fe2+ ions in 250 cm33 solution: 1.935×10−2 mol
    Mass of pure iron in the sample: 1.08 g
    Percentage purity of the iron sample: 60.0%
  3. A student is asked to determine the value of X in the hydrated iron(II) sulfate compound, FeSO4⋅XH2O. They dissolve 6.64 g of the hydrated compound in distilled water and make the solution up to 250 cm3 in a volumetric flask. 25.0 cm3 portions of this solution are then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below. Determine the value of X in FeSO4⋅XH2O. X is an integer.

    Concordant results and average titre: 24.225 cm3
    Moles of KMnO44​ used: 4.845×10−4 mol
    Moles of Fe2+ ions in 25.0 cm3 portion: 2.4225×10−3 mol
    Total moles of Fe2+ ions in 250 cm3 solution: 2.4225×10−2 mol
    Mass of anhydrous iron(II) sulfate (FeSO4​) in the sample: 3.68 g
    Mass of water in the hydrated compound: 2.96 g
    Moles of water: 0.164 mol
    Value of X in FeSO4⋅XH2O: 6.77 = 7
  4. A student is asked to determine the value of X in the hydrated sodium ethanedioate (sodium oxalate) compound, Na2C2O4⋅XH2O. They dissolve 3.50 g of the hydrated compound in distilled water and make the solution up to 250 cm3 in a volumetric flask. 25.0 cm3 portions of this solution are then titrated against a 0.0200 mol dm−3 solution of potassium manganate(VII) (KMnO4​). The student carries out the titration and records the following results in the table below:

    Mean titre: 22.40 cm³
    Moles of MnO₄⁻: 4.48 × 10⁻⁴ moles
    Moles of Na₂C₂O₄ in 25.0 cm³: 1.12 × 10⁻³ moles
    Moles of Na₂C₂O₄ in 250.0 cm³: 1.12 × 10⁻² moles
    Molar mass of Na₂C₂O₄·xH₂O: 312.5 g mol⁻¹
    Value of 𝑥: 10
    Final hydrated formula: Na₂C₂O₄·10H₂O
  5. A scientist conducts an experiment to determine the concentration of a solution of NaClO. A 25cm3 sample of the solution is reacted with excess KI. The resulting products are reacted against 0.1moldm-3 sodium thiosulfate, where it required 52.0cm3 to react.
     
    1. Write the equation for the reaction between chlorate ions and iodide ions, given their electrode potentials in the table above.
      ClO⁻ + 2I⁻ + 2H⁺ → Cl⁻ + I₂ + H₂O
    2. Write the equation for the reaction between iodine and thiosulfate, given their electrode potentials in the table above.
      I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻
    3. Calculate the concentration of the sodium chlorate(I)
      Moles of thiosulfate: (52/1000) x 0.1 = 0.0052
      Moles of iodine: 2:1 ratio, therefore 0.0052/2 = 0.0026
      Moles of chlorate(I): 1:1 ratio, therefore 0.0026
      Concentration of chlorate(I): 0.0026 / (25/1000) = 0.104 mol dm-3
  6. A student is given a sample of impure iron(II) ethanedioate, FeC₂O₄, and is asked to determine the mass of pure FeC₂O₄ in the sample. The student dissolves 2.80 g of the sample in distilled water and makes up the solution to 250.0 cm³ in a volumetric flask. The student then titrates 25.0 cm³ of this solution against 0.0150 mol dm⁻³ potassium manganate(VII) (KMnO₄) in acidic conditions. Calculate the percentage purity of the sample of iron ethanedioate.

    Mean titre: 21.70 cm³
    Moles of MnO₄⁻: 3.255 × 10⁻⁴ moles
    Moles of FeC₂O₄ in 25.0 cm³: 5.425 × 10⁻⁴ moles
    Moles of FeC₂O₄ in 250.0 cm³: 5.425 × 10⁻³ moles
    Mass of pure FeC₂O₄ in the sample: 0.781 g
    Percentage purity of FeC₂O₄ in the sample: 27.9%
  7. A student is given a sample of hydrated copper(II) sulfate, CuSO₄·xH₂O, and is asked to determine the value of x (the number of moles of water of crystallisation). To do this, the student dissolves 2.50 g of the sample in distilled water and makes up the solution to 250.0 cm³ in a volumetric flask. A 25.0 cm³ aliquot of this solution is taken and titrated with 0.0100 mol dm⁻³ EDTA (ethylenediaminetetraacetic acid) solution, finding that 27.40cm3 was required to fully react with the Cu2+ ions.
    Mean titre: 24.70 cm³
    Moles of EDTA reacted: 2.47 × 10⁻⁴ moles
    Moles of Cu²⁺ in 250.0 cm³ solution: 2.47 × 10⁻³ moles
    Mass of anhydrous CuSO₄ in the sample: 0.394 g
    Mass of water in the sample: 2.106 g
    Moles of water in the sample: 0.117 moles
    Value of x in CuSO₄·xH₂O: 5
  8. A 25.0 cm³ sample of hydrogen peroxide solution was titrated against a 0.0200 mol dm⁻³ solution of potassium permanganate in acidic conditions. The titration required 18.5 cm³ of potassium permanganate to reach the endpoint. Determine the concentration of the hydrogen peroxide solution in mol dm⁻³.
    Step 1: Moles of KMnO₄ = 0.00037 mol
    Step 2: Moles of H₂O₂ = 0.000925 mol
    Step 3: Concentration of H₂O₂ = 0.0370 mol dm⁻³
  9. A student conducts an experiment to determine the percentage of copper in an alloy. The student:

    – Reacts 1.25 g of the alloy with concentrated nitric acid to form a solution (all of the copper in the alloy reacts to form aqueous copper(II) ions)
    – Pours the solution into a volumetric flask and makes the volume up to 500 cm³ with distilled water
    – Shakes the flask thoroughly
    – Transfers 50.0 cm³ of the solution into a conical flask and adds an excess of potassium iodide
    – It requires exactly 12.50 cm³ of 0.100 mol dm⁻³ sodium thiosulfate (Na₂S₂O₃) solution to react with all the iodine produced.

    The equations for the reactions are:
    – 2 Cu²⁺ + 4 I⁻ → 2 CuI + I₂
    – 2 S₂O₃²⁻ + I₂ → 2 I⁻ + S₄O₆²⁻

    Calculate the percentage of copper by mass in the alloy.
    Step 1: Moles of sodium thiosulfate used: 0.00125 mol
    Step 2: Moles of iodine produced: 0.000625 mol
    Step 3: Moles of copper(II) ions: 0.00125 mol
    Step 4: Mass of copper in the sample: 0.07944 g
    Step 5: Total mass of copper in the original alloy: 0.7944 g
    Step 6: Percentage of copper in the alloy: 63.55%
  10. A student conducts an experiment to determine the percentage of zinc in a brass sample. The student:

    – Reacts 1.20 g of the brass sample with concentrated sulfuric acid to form a solution (all of the zinc in the brass reacts to form aqueous zinc ions, Zn²⁺)
    – Pours the solution into a volumetric flask and makes the volume up to 250 cm³ with distilled water
    – Shakes the flask thoroughly
    – Transfers 50.0 cm³ of the solution into a conical flask and adds an excess of sodium hydroxide to precipitate zinc hydroxide (Zn(OH)₂)
    – Dissolves the precipitate in hydrochloric acid and titrates the resulting solution with 0.100 mol dm⁻³ EDTA solution, using exactly 18.50 cm³ to react with all the zinc ions.

    Calculate the percentage of zinc by mass in the brass sample.
    Moles of EDTA used: 0.00185 mol
    Moles of zinc ions: 0.00185 mol
    Mass of zinc in the sample: 0.1209 g
    Total mass of zinc in the original brass sample: 0.6045 g
    Percentage of zinc in the brass sample: 50.38%
  11. A solid mixture contains ethanedioic acid (H₂C₂O₄), sodium ethanedioate (Na₂C₂O₄) and an inert solid. To determine the amount of ethanedioic acid and sodium ethanedioate in the mixture, the following procedure was carried out. Determine the mass (in grams) of ethanedioic acid and sodium ethanedioate in the original 2.50 g solid mixture.

    a) Calculation of ethanedioic acid using NaOH titration:

    From the equation, 1 mole of H₂C₂O₄ reacts with 2 moles of NaOH.
    Moles of NaOH used = 0.100 mol dm−3 × 0.0125 dm3 =1.25×10−3 mol
    Moles of H₂C₂O₄ in 25.0 cm³ = 1.25×10−3 / 2=6.25×10−4 mol

    b) Calculation of total ethanedioate using KMnO₄ titration:

    From the equation, 5 moles of C₂O₄²⁻ (from either H₂C₂O₄ or Na₂C₂O₄) react with 2 moles of KMnO₄.
    Moles of KMnO₄ used = 0.0200 mol dm−3 × 0.0150 dm3 = 3.00×10−4 mol
    Using the stoichiometric ratio:
    Total moles of C₂O₄²⁻ in 25.0 cm³ = (5/2) × 3.00×10−4 = 7.50×10−4 mol

    c) Calculation of sodium ethanedioate:

    Total moles of C₂O₄²⁻ = Moles from H₂C₂O₄ + Moles from Na₂C₂O₄.
    Moles of Na₂C₂O₄ = Total moles of C₂O₄²⁻ – Moles of H₂C₂O₄
    = 7.50×10−4−6.25×10−4=1.25×10−4 mol

    d) Mass of ethanedioic acid and sodium ethanedioate in the original mixture:

    Moles of H₂C₂O₄ in 250 cm³ solution:
    6.25×10−4 mol× (250/25) = 6.25×10−3 mol
    Mass of H₂C₂O₄ = 6.25×10−3 mol × 90.0 g mol−1=0.563 g

    Moles of Na₂C₂O₄ in 250 cm³ solution:
    1.25×10−4 mol × (250/25) =1.25×10−3 mol
    Mass of Na₂C₂O₄ = 1.25×10−3 mol × 134.0 g mol−1 = 0.168 g
  12. A 1.664g sample of vanadium(V) oxide (V₂O₅) is dissolved in 250cm3 of dilute sulfuric acid to form a solution containing vanadium(V) ions, VO2+. This solution is then reduced by adding a metal powder, resulting in a solution containing vanadium ions in a lower oxidation state. The reaction is allowed to proceed until completion, and the excess metal is removed by filtration.

    A 25cm3 sample of the vanadium ions present in the solution are titrated against acidified 0.05 mol dm-3 potassium manganate where it required 22.0cm3 KMnO4 to fully react. What was the oxidation state of the vanadium after it reacted with the metal.
    Moles of KMnO₄ = 0.5mol dm⁻³ × (22/1000) ​dm³ = 0.0011 mol
    MnO₄⁻ is reduced to Mn²⁺ in acidic solution: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O
    Since 1 mole of MnO₄⁻ accepts 5 moles of electrons:
    Moles of electrons transferred = 5 × 0.0011 = 0.00550mol
    Initially, 1.664 g of V₂O₅ was dissolved.
    Moles of V₂O₅: 0.00915mol
    Each mole of V₂O₅ produces 2 moles of VO₂⁺: 0.00915 × 2 = 0.0183 moles of VO₂⁺
  13. A container of H2O2 was tested by a scientist to determine whether it had started to decompose. A 1.00 cm3 sample of the hydrogen peroxide was taken from the container and 24.0 cm3 of water were added to it. This solution was reduced by reacting with potassium iodide. The iodine formed in the reduced back to iodide by reacting it against 2.00 moldm-3 sodium thiosulfate, where it required 42.6 cm3 of sodium thiosulfate to react. In the reaction, the thiosulfate (S2O32-) is oxidised to S4O62-.
    1. Write an equation for the reduction of hydrogen peroxide.
      2H+ + H2O2 + 2e– –> 2H2O
    2. Write an equation for the oxidation of iodide.
      I2 –> 2I– + 2e–
    3. Write an equation for the reaction between hydrogen peroxide and iodide.
      2H+ + H2O2 + 2I– –> 2H2O + I2
    4. Write an equation for the oxidation of S2O32- to S4O62-.Write an equation for the reaction between iodide and sodium thiosulfate.
      2S2O32-–> S4O62-+ 2e–
    5. Write an equation for the reaction between iodide and sodium thiosulfate.
      2S2O32- + I2 –> S4O62-+ 2I–
    6. Given the density of hydrogen peroxide is 1.45, determine whether the hydrogen peroxide had decomposed.
      Moles of thiosulfate: (42.6/1000) x 2 = 0.0852
      Moles of Iodine: 0.0852 / 2 = 0.0426Moles of H2O2: 1:1 ratio, therefore 0.0426
      Mass of H2O2 in the sample (before any decomposition): 1 x 1.45 = 1.45
      Moles of H2O2 in the sample (before any decomposition): 1.45 / 34 = 0.0426
      Therefore there had been no decomposition.
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