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In this experiment you will determine the formula of iron(III) ammonium sulfate, FeNH_4(SO_4)_2·xH_2O, where x is the number of molecules of water of crystallization. A known mass of this iron(III) compound reacted with excess acidified potassium iodide to produce iodine. You will determine the amount of iodine produced by titrating the mixture with sodium thiosulfate.
FA 1 is 0.900 mol dm^{-3} sodium thiosulfate, Na_2S_2O_3.
FA 2 is a solution of iodine, I_2, produced as outlined in the paragraph above. Use starch indicator.
(a) Method
Diluting FA 1
- Pipette 25.0 cm^3 of FA 1 into the 250 cm^3 volumetric (graduated) flask.
- Make the solution up to the mark using distilled water.
- Shake the flask to mix the solution thoroughly before using it for your titrations.
- Label this diluted solution of sodium thiosulfate FA 3.
- Rinse the pipette with distilled water.
Keep FA 1 for use in Question 3.
Titration
- Fill the burette with FA 3.
- Use the pipette to transfer 25.0 cm^3 of FA 2 into a conical flask.
- Add FA 3 from the burette into the conical flask until the mixture becomes pale yellow.
- Then add 10 drops of starch indicator to give a blue-black colour.
- Continue adding FA 3 until this blue-black colour disappears. This is the end-point of the titration.
- Perform a rough titration and record your burette reading in the space below.
The rough titre is ............... cm^3.
- Carry out as many accurate titrations as you think necessary to obtain consistent results.
- Make sure any recorded results show the precision of your practical work.
- Record in a suitable table below, all of your burette readings and the volume of FA 3 added in each accurate titration.
(b) From your accurate titration results, obtain a suitable value to be used in your calculations. Show clearly how you obtained this value.
25.0 cm^3 of FA 2 required ............... cm^3 of FA 3. [1]
(c) Calculations
Show your working and appropriate significant figures in the final answer to each step of your calculations.
(i) Using information on page 2, calculate the concentration, in mol dm^{-3}, of sodium thiosulfate in FA 3.
concentration of Na_2S_2O_3 in FA 3 = ............................. mol dm^{-3}
(ii) Calculate the number of moles of sodium thiosulfate present in the volume of FA 3 calculated in (b).
moles of Na_2S_2O_3 = ............................. mol
(iii) Use the equation below to calculate the number of moles of iodine that reacted with the sodium thiosulfate in (ii).
\[ I_2 + 2Na_2S_2O_3 \rightarrow 2NaI + Na_2S_4O_6 \]
moles of I_2 = ............................. mol
(iv) Calculate the concentration of I_2, in mol dm^{-3}, in FA 2.
concentration of I_2 = ............................. mol dm^{-3}
(v) The iodine in FA 2 was produced by the reaction of iron(III) ions with excess potassium iodide. Balance the equation for this reaction.
\[ ...... Fe^{3+}(aq) + ...... I^-(aq) \rightarrow ...... Fe^{2+}(aq) + ...... I_2(aq) \]
Use your answer to (iv) and this equation to calculate the number of moles of iron(III) ions that reacted to produce the iodine in 1.00 dm^3 of FA 2.
moles of Fe^{3+} = ............................. mol
(vi) The formula of the iron(III) compound is FeNH_4(SO_4)_2·xH_2O.
38.56 g of this compound was weighed out and added to excess aqueous acidified potassium iodide.
FA 2 was made by making the resulting solution of iodine up to 1.00 dm^3 with distilled water.
Use this information and your answer to (v) to calculate the number of moles of water of crystallization, x, in one mole of the iron(III) compound.
\[ A_r: H, 1.0; N, 14.0; O, 16.0; S, 32.1; Fe, 55.8 \]
\[ x = ............................... \] [6]
534
577
562
In this experiment you will determine the enthalpy change, $\Delta H$, for the reaction of zinc with iron(II) sulfate.
$\text{Zn(s) + FeSO}_4\text{(aq)} \rightarrow \text{Fe(s) + ZnSO}_4\text{(aq)}$
In order to do this, you will determine the enthalpy changes for the reactions of zinc and iron with aqueous copper(II) sulfate. Excess of the two metals will be used during the determinations.
Then you will use Hess’ Law to calculate the enthalpy change for the reaction above.
FA 4 is zinc, Zn.
FA 5 is iron, Fe.
FA 6 is 0.500 mol dm$^{-3}$ copper(II) sulfate, CuSO$_4$.
(a) Determination of the enthalpy change for the reaction of zinc, FA 4, with aqueous copper(II) sulfate, FA 6.
Method
• Support a plastic cup inside the 250 cm$^3$ beaker.
• Use the measuring cylinder to transfer 25 cm$^3$ of FA 6 into the plastic cup.
• Measure and record the initial temperature of the solution in the space below.
• Add all the FA 4 from the container to the FA 6 in the plastic cup.
• Stir constantly until the maximum temperature is reached.
• Measure and record the maximum temperature. Tilt the cup if necessary to ensure the thermometer bulb is fully immersed.
• Calculate and record the temperature rise.
(b) Calculations
Show your working and appropriate significant figures in the final answer to each step of your calculations.
(i) Calculate the energy produced during this reaction.
[Assume that 4.2 J are needed to raise the temperature of 1.0 cm$^3$ of solution by 1.0°C.]
energy produced $=$ ...................... J
(ii) Calculate the number of moles of copper(II) sulfate in 25 cm$^3$ of FA 6.
moles of CuSO$_4$ $=$ ...................... mol
(iii) Calculate the enthalpy change, in kJ mol$^{-1}$, for the reaction below.
$\text{Zn(s) + CuSO}_4\text{(aq)} \rightarrow \text{Cu(s) + ZnSO}_4\text{(aq)}$
enthalpy change $=$ ...... ...................... kJ mol$^{-1}$
(c) Determination of the enthalpy change for the reaction of iron, FA 5, with aqueous copper(II) sulfate, FA 6.
Method
• Support the second plastic cup inside the beaker.
• Use the measuring cylinder to transfer 25 cm$^3$ of FA 6 into the plastic cup.
• Measure and record the initial temperature of the solution in the space below.
• Add all the FA 5 from the container to the FA 6 in the plastic cup.
• Stir constantly until the maximum temperature is reached.
• Measure and record the maximum temperature. Tilt the cup if necessary to ensure the thermometer bulb is fully immersed.
• Calculate and record the temperature rise.
Keep solution FA 6 for use in Question 3.
(d) Calculations
Show your working and appropriate significant figures in the final answer to each step of your calculations.
(i) Calculate the energy produced during this reaction.
[Assume that 4.2 J are needed to raise the temperature of 1.0 cm$^3$ of solution by 1.0°C.]
energy produced $=$ ...................... J
(ii) Calculate the enthalpy change, in kJ mol$^{-1}$, for the reaction below.
$\text{Fe(s) + CuSO}_4\text{(aq)} \rightarrow \text{Cu(s) + FeSO}_4\text{(aq)}$
enthalpy change $=$ ...... ...................... kJ mol$^{-1}$
(e) Use your values for the enthalpy changes calculated in (b)(iii) and (d)(ii) to calculate the enthalpy change for the reaction below.
$\text{Zn(s) + FeSO}_4\text{(aq)} \rightarrow \text{Fe(s) + ZnSO}_4\text{(aq)}$
Show clearly how you obtained your answer by drawing a Hess’ Law energy cycle.
(If you were unable to calculate the enthalpy changes, assume that the value in (b)(iii) is $-210$ kJ mol$^{-1}$ and the value in (d)(ii) is $-144$ kJ mol$^{-1}$. Note: these are not the correct values.)
enthalpy change $=$ ...... ...................... kJ mol$^{-1}$
(f) (i) Calculate the maximum percentage error in the temperature rise in (c).
percentage error $=$ ...................... %
(ii) Apart from using a more accurately calibrated thermometer, suggest one improvement that could be made to this experiment that would increase the accuracy.
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543
566
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(a) FA 1, used in Question 1, is an aqueous solution of sodium thiosulfate.
FA 6, used in Question 2, is an aqueous solution of copper(II) sulfate.
FA 7 is solid sodium thiosulfate.
Carry out the following tests and record your observations in the table below.
[Table_1]
test observations
(i) Using a spatula, place 2 or 3 crystals of FA 7 into a hard-glass test-tube. Heat gently for several seconds, then
heat strongly.
(ii) To a 1 cm depth of dilute sulfuric acid in a boiling tube, add a few crystals of FA 7. Observe until no further change occurs, then
warm the mixture, gently and carefully.
(iii) To a 1 cm depth of aqueous potassium iodide in a test-tube, add a few drops of FA 6, aqueous copper(II) sulfate, then
add FA 1, aqueous sodium thiosulfate, to the mixture until no further change occurs.
(iv) Using your observations in (ii), complete the equation below by giving the formulae of the other two products.
Na₂S₂O₃ + H₂SO₄ → Na₂SO₄ + H₂O + .................... + ....................
(b) FA 8 and FA 9 are aqueous solutions. Each contains one cation and one anion from those listed on pages 14 and 15. Carry out the tests and record all your observations in the table. For each test, use a 1 cm depth of FA 8 or FA 9 in a test-tube.
[Table_2]
test observations
FA 8 FA 9
(i) Add aqueous sodium hydroxide.
(ii) Add a few drops of aqueous silver nitrate followed by aqueous ammonia.
(iii) Add aqueous ammonia.
(iv) Add an equal depth of dilute sulfuric acid.
(v) Add an equal depth of FA 9.
(vi) Using your observations, identify three of the ions present in FA 8 and FA 9.
Write ‘unknown’ next to the ion that you cannot identify directly from your observations.
ions in FA 8: cation ............................................. anion .............................................
ions in FA 9: cation ............................................. anion .............................................
(vii) The ‘unknown’ ion in (vi) can be identified by elimination, using the lists of ions on pages 14 and 15.
Choose one positive test that would confirm the identity of this ion. Name the reagent(s) you would use and state what you would observe if the test was positive. Do not carry out this test.
‘unknown’ ion tested for ............................................
reagent(s) .......................................................................................................
observation(s) ......................................................................................................
554
591
521
