Chemical Calc[1].Problems[Volumetric Analysis]

1

r

Hydrochloric acid neutralises Sodium hydroxide solution.

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)

25.0 cm3 of hydrochloric acid solution, of unknown concentration, is placed in a conical flask with a suitable indicator. The content was titrated with 0.1 mol dm-3 ( 0.1 M) NaOH solution placed in the burette.

The initial and final burette readings of the titre are given below :

(a) Name a suitable indicator for this titration. State the colour change of the

indicator. Indicator : ………………………………………………………………………….………

Colour changes from…………………..……….…….. to ……………..…………………

(b) Write the net ionic equation for this reaction

……………………………………………………………………………………………

(c) Name a piece of apparatus that could be used to transfer the 25.0 cm3 of HCl solution. ……………………………………………………………………………………………

(d) Use the table to calculate the mean titre of NaOH(give your answer to the correct number of signifcant figures).

Titre = ………………………………. cm3

(e) Calculate the amount of 0.1 mol dm-3 sodium hydroxide used in the titration and hence the concentration of the acid solution?

Volumetric analysis

4 321 25.500.50 24.000.00 Initial burette reading / cm3 48.0023.2046.5023.55Final burette reading / cm3

titre / cm3

2

A solution of HCl acid was titrated against a solution of sodium carbonate of 26.50 g.dm3 ,using Methyl orange indicator. It was found that 25.0 cm3 of the acid solution neutralized 20.0 cm3 of the alkali.

Na2CO3 + 2HCl → 2NaCl(aq) + CO2(g) + H2O(l)

Calculate

(a) the amount (mol) of the alkali used

(b) the amount (mol) of the acid used .

(c) the concentration of acid in mol. dm-3

The titrations of 25.0 cm3 of sodium hydroxide solution neutralise completely 10.0 cm3

of sulphuric acid solution of concentration 0.10 mol/dm3. What is the concentration of the sodium hydroxide solution?

2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)

3

A solution of sodium hydroxide, NaOH, ( Mr = 40.0 ) contains 10.0 g dm-3. (a) What is the molar concentration of this solution?

(b) What volume of this solution would be needed to neutralise completely 25.0 cm3 of 0.20 mol dm-3 hydrochloric acid solution?

What volume of 0.500 M nitric acid solution is required to completely neutralise 3.70 g of calcium hydroxide, Ca(OH)2, according the following equation?

2HNO3 + Ca(OH)2 → Ca(NO3)2 + H2O

4

The titration of H2SO4 with KOH is represented by the equation :

H2SO4 + 2KOH → K2SO4 + 2 H2O What volume of 0.1 mol dm-3 KOH is required to titrate 50.0 ml of 0.200 mol dm-3 H2SO4 ?

In an experiment to determine the molar concentration of a hydrochloric acid solution without using titration technique, a student weighed 2.00g of solid magnesium and then dropped this quantity in 50.0 ml of HCl solution. When the reaction stopped, he weighed the remaining piece of magnesium after drying it. The mass of magnesium left was found to be 0.50g

Mg(s) + 2HCl(aq) → MgCl2 + H2(g)

(a) Calculate the amount of magnesium that reacted with the 50.0 ml of HCl solution.

(b) How would,away from any observation, make sure that all the acid has reacted?

……………………………………………………………………………………...

……………………………………………………………………………………...

(c) Calculate the molar concentation of the hydrochloric acid solution.

[HCl ] = ………………………….. mol dm-3

5

A sample of hydrated sodium carbonate crystals , Na2CO3.×H2O, weighing 4.00g was dissolved in enough water and made to 250 cm3. Portions of 25.0 cm3 of this solution required 22.40 cm3 of HCl acid solution of conc. 0.20 mol.dm-3 for complete neutralization.

Na2CO3 + 2HCl → 2NaCl(aq) + CO2(g) + H2O(l)

(a) Calculate the amount(mol)of HCl used in the titration.

(b) Calculate the amount of the hydrate that reacted . (c) Find × .

A solution of 50 cm3 of hydrogen peroxide, H2O2,was diluted to 1 dm3 of this solution were acidified with dilute sulfuric acid, and titrated with 0.20 mol.dm-3 KMnO4 solution in which 22.0 cm3 was required. Find the concentration of the original hydrogen peroxide.

2MnO4- + 5H2O2 + 6H+ → 2Mn+2 + 5O2 + 8H2O

6

A student crushed 0.500g of dry solid of Vitamin C tablet in a mortar to obtain fine powder.He then dissolved the amount in enough of distilled water and made up the volume to 100 ml in a volumetric flask . The student then titrated 25.0cm3 of this acid solution with with 0.02 mol/dm3 NaOH solution. The average titre of sodium hydroxide was found to be 14.50 cm3

Data

• Formula of ascorbic acid C6H8O6 • Relative molar mass of ascorbic acid Mr (C6H8O6 ) = 176 • The reaction between sodium hydroxide and ascorbic acid:

C6H8O6 (aq) + NaOH(aq) → C6H7O6 Na + H2O(l)

(a) Calculate the amount ( number of moles) of NaOH used in the titration (b) Hence, calculate the number of moles of ascorbic acid present in 25.0 cm3 solution used in the titration. (c) Hence, calculate the % by mass of ascorbic acid present in Vitamin C tablet.

7

Volumetric Analysis 1

The determination of the concentration of a solution of hydrochloric acid A student was provided with a solution (R) of Hydrochloric acid of unknown concentration and was required to determine its concentration. ♦ Procedure

A. Dilution of the hydrochloric acid solution (R)

1. Transfer 10.0 cm3 of solution (R) to a volumetric flask . 2. The solution in the volumetric flask is made up to 100 cm3 with distilled water 3. The dilute solution of HCl prepared in (2) is then labelled as solution (Q).

B. Titration of the diluted hydrochloric acid solution (Q) with NaOH solution

Using a pippete the student placed 25 cm3 of solution ( Q ) of the diltuted hydrochloric acid into a conical flask then added 3 drops of an indicator . Then he titrated the solution with 0.100 mol/dm3 sodium hydroxide solution .

♦ Results

The following results were obtained by the student,

I II III

Initial burette Reading /cm3

0.00

29.05

2.55

Final burette Reading /cm3

15.40

41.55

15.10

Titre / cm3

♦ Equation of the Reaction:

HCl(aq) + NaOH(aq) → NaCl (aq) + H2O(l)

(a) Write the ionic equation of the above reaction

………………………………………………………………………………………

(b) Suggest a suitable indicator for the above reaction and state the colour change

Indicator……………………………………… Colour changes from ……………………….. to …………………………………

8

(c) Name a suitable piece of apparatus to transfer 10.0 cm3 of solution (R) to a volumetric flask ………………………………………………………………………………………

(d) Which two values would you choose to calculate the average (mean) titre

………………………………………………………………………………………. (e) Calculate the amount ( number of moles) of sodium hydroxide used for neutralisation. (f) Hence calculate the concentration of the hydrochloric acid solution.(Q) (g) Hence calculate the concentration of the original hydrochloric acid solution(R).

9

Volumetric Analysis 2 Acid base Titration

The determination of Water of crystallisation in Na2CO3 × H2O

In an exercise to determine the number of moles of water of crystallization, × , in one mole of Sodium carbonate Na2CO3×H2O .A weighed amount of sodium carbonate was dissolved in water and made up exactly to 250 cm3 of solution 25.0 cm3 of the sodium carbonate was then titrated against hydrochloric acid , HCl, of concentration 0.25 mol/dm3. The table of results is shown below The average titre was 10.5 cm3.

The equation of the reaction is : Na2CO3 × H2O + 2HCl → 2 NaCl + CO2 + (× +1) H2O

Mass of empty bottle

m1

Mass of the empty bottle + amount of Na2CO3×H2O

m2

Mass of the emptied bottle ( after you transfer the weighed amount of the solid of Na2CO3×H2O )

m3

= …………………g = …………………g = …………………g

10.50

15.75

12.0

10

(a) Calculate the mass of solid hydrated sodium carbonate dissolved.

Mass of sodium carbonat =………………………...g (b) How many moles of HCl solution were used in the titration ?

Amount of HCl = ……………………………….mol (c) Hence, calculate the amount( number of moles) of sodium carbonate present in

25cm3 of solution. (d) Hence, calculate the amount( number of moles) of sodium carbonate present in

250 cm3 of solution. (e) Hence, calculate the mass of 1 mol of Na2CO3 × H2O and then deduce the value of × .

M = …………………………… × = ……………………………

11

Volumetric Analysis 3 FA1 is a hydrated metal sulphate, XSO4.7H2O. You are required to determine the mass of water of crsytallisation in a weighted sample of FA1 and to calculate the relative atomic mass, Ar of the element X.

(a) Accuraltely weigh the hard glass test-tube provided. Record the mass in Table 1.1 below. Add to the test-tube between 2.00g and 2.50g of FA1 and accurately weigh the test-tube and contents. Record this mass in Table 1.1 below.

Table 1.1 Mass of FA1

Mass of test-tube + FA1 /g

12.80

Mass of empty test -tube /g

10.50

Mass of FA1 /g

…………………..

(b) Heat the test-tube, gently at first then strongly, to drive off the water of crystallisation. The crystals will ' crackle' at first as water is lost and 'steam' ( condensed water vapour) will be seen coming out of the mouth of the tube.

If the crystals are overheated the sulphate can decompose and give off sulphur trioxide, which will be seen as white fumes. If you see white fumes, do not confuse this with steam, stop heating.

Place the test-tube on a heat proof mat and leave to cool. Do not move about the laboratory with a hot test-tube.

When cool , reweigh the test tube and its contents. Record the mass in Table 1.2 below.

Table 1.2 Mass of FA1 after heating

Mass of test-tube + FA1 after heating /g

11.76

Mass of empty test-tube ( from Table 1.1) /g

10.50

Mass of FA 1 after heating /g

[4]

Accuracy [6]

12

(c) By repeating the heating, cooling and rewieghing, show clearly by your results in Table 1.2 that the water of crystallisation has been driven from the crystals, FA1.

(d) Calculate

(i) the mass of anhydrous XSO4 present in the crystals.

(ii) the mass of water driven from the crystals of FA1.

[1] (e) Calculate how many moles of water are present in the sample of FA1 used

[ Ar ; H, 1.0 ; O, 16]

[1]

(f) Use your answer to (e) and the formula XSO4.7H2O to calculate how many moles of XSO4 are present in the sample of FA1 used.

[1]

(g) Use your answers to (d) and (f) to calculate the relative molecular mass , Mr, of XSO4.

[1] (h) Calculate the relative atomic mass, Ar , of the element X.

[ Ar ; O, 16 ; S, 32.0 ]

[1]

13


FA 1 is anhydrous sodium carbonate, Na2CO3, provided in a stoppered tube. FA 2 is an aqueous solution of hydrochloric acid, HCl. Acids and carbonates in solution react as shown in the equation.

2H+(aq) + CO32−(aq) → H2O(l) + CO2(g)

You are to determine the concentration, in mol dm-3, of the hydrochloric acid FA 2.

(a) Weigh the stoppered tube labelled FA 1 and record the mass in Table 1.1

Table 1.1 Weighing of sodium carbonate


13.12

Mass of empty test -tube + residual FA1 /g

11.00

Mass of FA1 used /g

…………………..

Transfer the contents of the weighed tube into a 250 cm3 beaker and dissolve the solid in about 100 cm3 of distilled water. Reweigh the tube and stopper and any residual sodium carbonate and record the mass in Table 1.1. Calculate the mass of sodium carbonate dissolved in the water. (b) Transfer the sodium carbonate solution to the graduated flask labelled FA 3. Rinse the beaker with distilled water several times, adding each rinsing to the graduated flask. This ensures that all of the sodium carbonate has been transferred to the flask. Make up the solution to 250 cm3 with distilled water and mix thorougly. Pipette 25.0 cm3 of FA 3, the sodium carbonate, into a conical flask and place the flask on a white tile. Add a few drops of indicator provided and titrate with FA 2, the hydrochloric acid. Repeat the titration as many times as you think necessary to obtain accurate results. Make certain that the recorded results show the precision of your practical work.

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Table 1.2 Titration of FA 3 with FA2

Final burette reading /cm3 11.50 22.25 33.30 10.20 Initial burette reading / cm3 0.00 12.20 23.20 0.00 Volume of FA 2 used / cm3

Summary 25.0 cm3 of FA 3 reacted with …………………….cm3 of FA 2. Show which results you used to obtain this volume of FA 2 by placing a tick ( ) under the readings in Table 1.2.

You are advised to show full working in all parts of the calculations. (c) Calculate the concentration in mol.dm-3 of the sodium carbonate, Na2CO3, in FA 3.

[ Ar : Na, 23.0 ; C, 12.0; O, 16.0] (d) Calculate how many moles of sodium carbonate, Na2CO3 ,were pipetted into

the conical flask. (e) Calculate how many moles of hydrochloric acid, HCl, have been run from the

burette.

2H+(aq) + CO32-(aq) → H2O(l) + CO2(g)

(f) Calculate the concentration , in mol dm-3 , of HCl in FA 2.

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FA 1 is a mixture of anhydrous sodium carbonate, Na2CO3, and sodium chloride NaCl provided in a stoppered tube. FA 2 is an aqueous solution of hydrochloric acid, 0.10 mol dm-3 HCl. Acids and carbonates in solution react as shown in the equation.

2H+(aq) + CO32-(aq) → H2O(l) + CO2(g)

You are to determine the % , by mass , of NaCl in acid FA 1. Weigh the stoppered tube labelled FA 1 and record the mass in Table 1.1

Table 1.1 Weighing of sodium carbonate and Sodium chloride sample


13.50

Mass of empty test -tube + residual FA1 /g

10.50

Mass of FA1 used /g

…………………..

Transfer the contents of the weighed tube into a 250 cm3 beaker and dissolve the solid in about 100 cm3 of distilled water. Reweigh the tube and stopper and any residual sodium carbonate and record the mass in Table 1.1. Calculate the mass of sodium carbonate dissolved in the water. Transfer the sodium carbonate solution to the graduated flask labelled FA 3. Rinse the beaker with distilled water several times, adding each rinsing to the graduated flask. This ensures that all of the sodium carbonate has been transferred to the flask. Make up the solution to 250 cm3 with distilled water and mix thorougly. Pipette 25.0 cm3 of FA 3, the sodium carbonate, into a conical flask and place the flask on a white tile. Add a few drops of indicator provided and titrate with FA 2, the hydrochloric acid.

Repeat the titration as many times as you think necessary to obtain accurate results. Make certain that the recorded results show the precision of your practical work.

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Table 1.2 Titration of FA 3 with FA2

Final burette reading /cm3 23.50 22.20 2.50 21.50 Initial burette reading / cm3 0.00 43.60 24.10 0.00 Volume of FA 2 used / cm3

Summary 25.0 cm3 of FA 3 reacted with …………………….cm3 of FA 2. Show which results you used to obtain this volume of FA 2 by placing a tick ( ) under the readings in Table 1.2.

You are advised to show full working in all parts of the calculations. (a) Calculate the amount (number of moles) of HCl in the mean titre. (b) Calculate how many moles of sodium carbonate, Na2CO3 ,were pipetted into

the conical flask. (c) Hence, calculate the amount of sodium carbonate, Na2CO3 in 250.0 cm3 of

solution FA 3. (d) Hence, calculate the % , by mass of sodium carbonate in the sample.

[ Ar : Na, 23.0 ; C, 12.0;

17

Volumetric Analysis 6 Dilution

Acid base Titration HCl/ NaOH

You are provided with Solution C A solution of hydrochloric acid (HCl) of Unknown molar conc. Solution B 0.1 mol.dm –3 NaOH {Sodium hydroxide} solution

You are required to

• Find the concentration of the diluted solution (D) and then use the same solution to • Find the concentration of solution (C) of HCl

Procedure

� Part A Dilution of solution (C) of HCl

Transfer 10.0 cm3 of the HCl solution (C) of HCl in a 100 cm3 volumetric flask. Make up the solution up to Mark with distilled water. Stopper the volumetric flask. Invert & shake the flask several times Label it as Diluted HCl or Solution (D).

� Part B Titration of Solution (D) With NaOH solution.

Using a pipette, transfer 25.0 cm3 of the diluted solution (D) of hydrochloric acid into a clean conical flask. Add 2-3 drops of phenolphthalein indicator.

Clean through the burette with distilled water and then with NaOH solution.

Titrate slowly with the NaOH solution with constant swirling until ONE single drop of NaOH solution causes a permanent pink colour (30 seconds) Record the volume} two accurate titrations} in the table below : Repeat this step till you get at least 2 concordant titres.

Rough accurate1 accurate2 Final burette reading /cm3

13.5

26.5

40.8

Initial burette reading /cm3

0.0

14.0

28.5

Titre / cm3

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(a) Name the piece of apparatus that you could use to transfer 10 cm3 of the HCl solution

(C) of HCl into a 100 cm3 volumetric flask. …………………………………………………………………………………………………..

(b) Calculate the average (mean) titre of sodium hydroxide solution from your table

(c) Calculate the amount( number of moles ) of NaOH solution required to neutralise completely the 25 cm3 of the diluted HCl solution (D)

(d) Hence calculate the concentration of solution (D)

[HCl] diluted = …………………….. mol/dm3

(e) Hence calculate the concentration of the solution (C) { original solution of

hydrochloric acid}.

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Acid base Titration

The determination of relative atomic mass of a metal by back-titration .

Total # of moles (HCl ) = moles reacted with NaOH(excess) + moles reacted with metal

n HCl total = n HCl(excess) + n HCl reacted with metal A student was asked to find the relative atomic mass (Ar) of a metal M .

� The student added 1.12 g of the metal M to 100 cm3 of 0.6 mol/dm3 hydrochloric acid solution which is in excess.

� The acid solution remained after the reaction has ceased ( stops) , was placed in burette and titrated with a 0.8 mol/dm3 sodium hydroxide

solution using a suitable indicator. The table of results is shown below

Titration 1

Titration 2

Titration 3

Final burette reading /cm3

26.50

27.50

25.30


0.50

2.0

0.00

Titre / cm3

Equations M(s) + 2HCl(aq) → MCl2(aq) + H2(g)

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

(a) Name a suitable indicator for this titration.

moles of NaOH C×V

moles of excess HCl from eqn.( 1:1)

moles of HCl reacted with metal n reacted = n total – n excess

moles of metal from eqn.

R.A.M of metal by

Ar = nm

20

Indicator :……………………………………………………………………………

Colour in acid : …………………………………………………….………………..

Colour in alkali : …………………………………………………………………….

(b) From the table of results calculate the average titre of NaOH

The average titre …………………… cm3.

Thus ………………..cm3 of 0.8 mol/dm3 NaOH required to react with excess amount of HCl acid solution.

(c) Calculate the number of moles of NaOH in 0.8 mol/dm3 solution.( reacted with excess

amount of HCl acid ). Amount of NaOH reacted with excess HCl = ………………………………….moles (a) Hence calculate the number of moles of excess hydrochloric acid reacted with NaOH solution.

Amount of excess HCl = ………………………………….moles

(b) Calculate the total number of moles (amount) of hydrochloric acid HCl present in the

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original 100 cm3 of 0.6 mol/dm3 HCl solution.

(c) Hence calculate the number of moles (amount) of HCl which reacted with 1.12g of the metal M.

Amount of HCl reacted with the 1.12g metal M ………………………………..moles

(d) Calculate the relative atomic mass, Ar , of the metal M

The relative atomic mass of M = ………………………………………

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FA 1 is a solution containing 5.00 g dm –3 of hydrated ethanedioic acid, H2C2O4.xH2O. FA 2 is a solution containing 2.37 g dm –3 of potassium manganate(VII), KMnO4. You are also provided with 1.00mol dm –3 sulphuric acid, H2SO4. In the presence of acid, potassium manganate(VII) oxidises ethanedioic acid; In the presence of acid, potassium manganate(VII) oxidises ethanedioic acid; 2MnO4–(aq) + 5H2C2O4(aq) + 6H+(aq) → 2Mn2+(aq) + 10CO2(g) + 8H2O(l) You are to determine the value of x in H2C2O4.xH2O. (a) Fill the burette with FA 2. Pipette 25.0 cm3 of FA 1 into a conical flask. Use the measuring cylinder provided to add to the flask 25 cm3 of 1.00 mol dm –3 sulphuric acid and 40 cm3 of distilled water. Heat the solution in the flask until the temperature is just over 65 °C. The exact temperature is not important. Be careful when handling hot solutions. Remove the thermometer and carefully place the hot flask under the burette. If the neck of the flask is too hot to hold safely, use a folded paper towel to hold the flask. Run in about 1 cm3 of FA 2. Swirl the flask until the colour of the manganate(VII) ions has disappeared then continue the titration as normal until a permanent pale pink colour is obtained. This is the end point. Record the burette readings in Table 1.1. If a brown colour appears during the titration, reheat the flask to 65 °C. The brown colour should disappear and the titration can then be completed. If the brown colour does not disappear on reheating, discard the solution and restart the titration. Repeat the titration as many times as you think necessary to obtain accurate results. Make certain that the recorded results show the precision of your practical work.

Table 1.1 Titration of FA 1 with FA 2 Final burette reading/cm3

0.00

0.40

5.00

Initial burette reading/cm3

27.80

27.10

31.70

Volume of FA2 used /cm3

Summary 25.0 cm3 of FA 1 reacted with ………………. cm3of FA 2. Show which results you used to obtain this volume of FA 2 by placing a tick ( ) under the readings in Table 1.1.

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You are advised to show full working in all parts of the calculations. (b) Calculate how many moles of potassium manganate(VII), KMnO4, were run from the

burette during the titration. [Ar: K, 39.1; Mn, 54.9; O, 16.0.]

[2] (c) Calculate how many moles of ethanedioic acid, H2C2O4, reacted with the potassium

manganate(VII) run from the burette. [1]

(d) Calculate the mass of H2C2O4 in each dm3 of FA 1 [Ar: H, 1.0; C, 12.0; O, 16.0.]

[3]

(e) Calculate the mass of water in the 5.00 g of H2C2O4.xH2O. [1]

(f) Calculate the value of x, in H2C2O4.xH2O. [1]

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[Total: 15] FA 1is a solution containing 7.76 g dm –3 of hydrated , sodium thiosulphate Na2S2O3 xH2O. FA 2 is a solution containing 0.010 mol dm–3 iodine This is a good place to introduce this titration. 25.0 cm3 samples of 0.010 mol dm–3 iodine solution are titrated with 7.76 g dm–3 sodium thiosulphate solution, which is delivered from a burette. The end point can be made clearer by adding starch solution when the iodine colour is a pale straw colour. Students could be asked to see if their results are consistent with the equation for the reaction

2Na2S2O3(aq) + I2(aq) → 2NaI(aq) + Na2S4O6(aq) It is also a useful exercise to ask students to work out the oxidation numbers of sulphur in the sodium compounds in the equation.?

Table 1.1 Titration of FA 1 with FA 2 Final burette reading/cm3

0.00

0.30

15.00

Initial burette reading/cm3

12.80

12.80

27.60

Volume of FA1 used /cm3

Summary 25.0 cm3 of FA 2 reacted with ………………. cm3of FA 1. Show which results you used to obtain this volume of FA 1by placing a tick ( ) under the readings in Table 1.1. You are advised to show full working in all parts of the calculations. (b) Calculate how many moles of iodine were run transferred into the conical flask. [2] (c) Calculate how many moles of sodium thiosulphate run from the burette.

[1]

(d) Calculate the molar mass Mr of the hydrated Na2S2O3 xH2O.

25

[3]

(e) Calculate the value of x, [1]

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Volumetric Analysis 10 Redox Titration

KMnO4 / FeSO4

Objectives :

To find the stoichiometric factors a and b in the redox equation given below:

a MnO4 − (aq) + b Fe 2+

(aq) → Products

� Introduction : The intense purple colour of potassium manganate (VII) serves as a visual indicator(hence, no indicator is needed)

FeSO4 → Fe++ + SO4 2−

[FeSO4 ] = [ Fe+2] ( 1 : 1)

Preparation of Standard Iron(II) sulphate solution

1. Weigh about 1.10 to 1.20g (to the nearest 0.1 mg) of dry solid of FeSO4.7 H2O In Table 1

- Record the mass of empty bottle as m1 - Record the mass of the empty bottle with amount of FeSO4.7H2O as m2 - Record the mass of the empty bottle after you transfer the weighed amount of

the solid as m3 2. Transfer the weighed amount of the solid FeSO4.7H2O into a 100 ml clean beaker.

3. Dissolve the amount in about 30-50 ml of distilled water using a glass rod.

4. Transfer all the solution through a filter funnel into a 100 cm3 volumetric ( graduated)

flask. Then wash the funnel and the flask and transfer all washings to the volumetric flask.

5. Make up the volume up to the mark (i.e to 100ml) by adding distilled water.

6. Insert a stopper and invert the flask several times so the solution is well mixed.

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Part B

Titration of FeSO4 solution with 0.01 M potassium Manganate(VII) KMnO4 Solution .

• Using a pipette, transfer 25 cm3 of the FeSO4 into a clean conical flask.

• Add about 15-20 cm3 of sulphuric acid to the conical flask to acidify the FeSO4 solution.

• Clean through the burette with distilled water and then with KMnO4 solution.

• Fill the burette with KMnO4solution. Don’t forget to fill the tip.

• Record the initial reading in the table below.

• Titrate slowly with KMnO4 solution with constant swirling until ONE single drop of

KMnO4

• solution causes a permanent faint pink colour (which lasts for about 30 seconds)

• Record the volume {two accurate titrations} in the table below

• Repeat the titration two more times.

Data & Calculations Table 1

Mass of empty bottle

Mass of the empty bottle with amount of FeSO4.7H2O

Mass of the empty bottle after you transfer the weighed amount of the solid

m1

m2

m3

= 10.5 g = 12.80 g = 11.7 g

Table 2

Rough Accurate1 Accurate2

Final burette reading /cm3

21.50

42.50

22.80


0.00

22.10

2.50

Titre / cm3

(a) Therefore 25 ml of Fe+2 solution required …………..…….……….. ml of KMnO4 of 0.01 M

for complete redox reaction.

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Given

• [ KMnO4 ] = 0.01 mol dm-3

• M = Molar mass of FeSO4.7H2O = 278 g.mol-1

• V = Volume of prepared solution = 100 ml = 0.1 dm3

• m = Mass of dissolved FeSO4.7H2O = ………………….g

• C = Molar concentration of FeSO4.7H2O = [FeSO4.7H2O ]

= [ Fe+2 ]

= Vn

=VM

m×

= ………………………… mol/ dm3

(c) Calculate the average (mean) titre from your table

(a) Using the obtained data , calculate the stoichiometric factors a & b

a MnO4− + b Fe 2+ → Products

a = ………………………. b = ……………………….. (e) Explain briefly , why purple KMnO4 solution was placed in the burette during this

titration rather than in the conical flask ……………………………………………………………………………………… ………………………………………………………………………………………

(f) Complete the following reduction half-equation :

MnO4- → Mn++

Chemical Calc[1].Problems[Volumetric Analysis]

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