Doc Brown's Chemistry Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 9. Group 7/17 The Halogens

9.9 Volumetric analysis involving halogens or halide ions

Volumetric titration analysis questions involving halogens or halide ions e.g. iodine titrated with sodium thiosulfate, chloride ion titrated with silver nitrate solution.

PLEASE NOTE KS4 Science GCSE/IGCSE/O Level GROUP 7 HALOGENS NOTES are on a separate webpage

9.9 Volumetric Analysis – titrations involving halogens or halide ions

(1) Titrating chloride with silver nitrate * (2) Titrating iodine with sodium thiosulphate

The questions below are from the Volumetric (non–redox) Titration Calculations Page with ANSWERS!

• Theory of reaction and using potassium chromate (VI) indicator.

  • As the silver nitrate is run into the chloride solution, white silver chloride is precipitated.

    • Ag⁺(aq) + Cl^–(aq) ==> AgCl(s)

  • When all the chloride has been precipitated, the first drop of excess silver nitrate solution causes a dark brick red precipitate of silver chromate to form, thus indicating the end–point.

    • 2Ag⁺(aq) + CrO₄^2–(aq) ==> Ag₂CrO₄(s)

  • The indicator is chosen on the basis that the solubility product of silver chloride is exceeded before that of silver chromate(VI).

(a) Give the ionic equation for the reaction of silver nitrate and chloride ion.

(b) Calculate the moles of chloride ion in the titrated 25cm³ aliquot.

(c) Calculate the molarity of chloride ion in the diluted seawater.

(d) Calculate the molarity of chloride ion in the original seawater.

(e) Assuming that for every chloride ion there is a sodium ion, what is the theoretical concentration of sodium chloride salt in g dm^–3 in seawater?

19.7 cm³ was required to titrate all the chloride ion. [Atomic masses: Na = 23, Cl = 35.5]

(a) How many moles of chloride ion was titrated?

(b) What mass of sodium chloride was titrated?

(b) What was the % purity of the rock salt in terms of sodium chloride?

 

On titration with 0.1 mol dm^–3 silver nitrate solution, 21.2 cm³ was required to precipitate all the chloride ion. [Atomic masses: Ca = 40, Cl = 35.5]

(a) Calculate the moles of chloride ion titrated.

(b) Calculate the equivalent moles of calcium chloride titrated.

(c) Calculate the equivalent mass of calcium chloride titrated.

(d) Calculate the total mass of calcium chloride in the original 5.0 g of the mixture.

(e) The % of calcium chloride and sodium nitrate in the original mixture.

The questions above are from the Volumetric (non–redox) Titration Calculations Page with ANSWERS!

• The redox theory

• Half–cell reaction data:

  • (i) ¹/₂I₂(aq) + e^– ==> I^–(aq) (E^Ø = +0.54V, reduction of the oxidising agent, iodide gets oxidised)

  • (ii) ¹/₂S₄O₆^2–(aq) + e^– ==> 2S₂O₃^2–(aq) (E^Ø = +0.09V, S₄O₆^2– will act as reducing agent, E^Ø less positive)

• The iodine is reduced by the thiosulphate ion to form iodide, ox. state of I (0) to (–1).

• The thiosulphate ion is oxidised to the tetrathionate ion. In doing so the sulphur atom changes ox. state from an average of four at (+2) in the two S₂O₃^2– ions to an average of four at (+2.5) in the single S₄O₆^2– ion.

  • A bit of an awkward one in analysing sulphur and it is best to reason in terms of an average oxidation state of sulphur.

• 2 x reduction half–cell, (i)	I2(aq) + 2e– ==> 2I–(aq)
  2 x oxidation half–cell, (ii) rev.	2S2O32–(aq) ==> S4O62–(aq) + 2e–
  added gives full equation	2S2O32–(aq) + I2(aq) ==> S4O62–(aq) + 2I–(aq)

• This is used to quantitatively estimate iodine in aqueous solution.

• Indicator theory

• The indicator is a few drops of starch solution which forms a blue–black complex with iodine.

• The end–point is when the solution first becomes colourless with no remaining iodine to form the coloured complex.

The questions below are from the Volumetric Redox Titration Calculations Page with ANSWERS!

Question 2: Given the following two half–reactions

(a) Given (i) S₄O₆^2–(aq) + 2e^– ==> 2S₂O₃^2–(aq)

and (ii) I₂(aq) + 2e^– ==> 2I^–(aq)

construct the full ionic redox equation for the reaction of the thiosulphate ion S₂O₃^2–,and iodine.

(b) what mass of iodine reacts with 23.5 cm³ of 0.012 mol dm^–3 sodium thiosulphate solution.

(c) 25cm³ of a solution of iodine in potassium iodide solution required 26.5 cm³ of 0.095 mol dm^–3 sodium thiosulphate solution to titrate the iodine.

What is the molarity of the iodine solution and the mass of iodine per dm³?

Question 17: 25.0 cm³ of an iodine solution was titrated with 0.1 mol dm^–3 sodium thiosulphate solution and the iodine reacted with 17.6 cm³ of the thiosulphate solution.

(a) give the reaction equation.

(b) what indicator is used? and describe the end–point in the titration.

(c) calculate the concentration of the iodine solution in mol dm^–3 and g dm^–3.

Question 14: Given the half–cell reaction IO₃^–(aq) + 6H⁺(aq) + 5e^– ==> ¹/₂I₂(aq) + 3H₂O(l)(see also Q2)

(a) Deduce the redox equation for iodate(V) ions oxidising iodide ions.

(b) What volume of 0.012 mol dm^–3 iodate(V) solution reacts with 20.0 cm³ of 0.100 mol dm^–3 iodide solution?

(c) 25.0 cm³ of the potassium iodate solution were added to about 15 cm³ of a 15% solution of potassium iodide (ensures excess iodide ion). On acidification, the liberated iodine needed 24.1 cm³ of 0.05 mol dm^–3 sodium thiosulphate solution to titrate it.

(i) Calculate the concentration of potassium iodate(V) in g dm^–3

(ii) What indicator is used for this titration and what is the colour change at the end–point?

Question 18: 1.01g of an impure sample of potassium dichromate(VI), K₂Cr₂O₇, was dissolved in dil. sulphuric acid and made up to 250 cm³ in a calibrated volumetric flask. A 25 cm³ aliquot of this solution pipetted into a conical flask and excess potassium iodide solution and starch indicator were added. The liberated iodine was titrated with 0.1 mol dm^–3 sodium thiosulphate and the starch turned colourless after 20.0 cm³ was added.

(a) Using the half–equations from Q3(a)(ii) and Q2(a)(ii), construct the full balanced equation for the reaction between the dichromate(VI) ion and the iodide ion.

(b) Using the half–equations from Q2(a) construct the balanced redox equation for the reaction between the thiosulphate ion and iodine.

(c) Calculate the moles of sodium thiosulphate used in the titration and hence the number of moles of iodine titrated.

(d) Calculate the moles of dichromate(VI) ion that reacted to give the iodine titrated in the titration.

(e) Calculate the formula mass of potassium dichromate(VI) and the mass of it in the 25 cm³ aliquot titrated.

(f) Calculate the total mass of potassium dichromate(VI) in the original sample and hence its % purity.

The questions above are from the Volumetric Redox Titration Calculations Page with ANSWERS!

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