Mandatory Experiment 4.6
Estimation of iron(II) in an iron tablet using a standard
solution of potassium manganate(VII)
Student Material
Theory
To estimate the iron(II) content of an iron tablet, a small number of tablets are first
dissolved in dilute sulfuric acid. This solution is then titrated against previously
standardised potassium manganate(VII) solution. The reaction is represented by the
equation:
MnO4- + 8H+ + 5Fe+2  Mn+2 + 5Fe+3 + 4H2O
Chemicals and Apparatus
0.005 M potassium manganate(VII) solution
Iron tablets
1.5 M sulfuric acid solution
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Deionised (or distilled) water
Safety glasses
Electronic balance
Clock glass
Dropping pipette
Mortar and pestle
Pipette (25 cm3)
Pipette filler
Burette (50 cm3)
Conical flask (250 cm3)
White card
White tile
Retort stands
Boss-head
Clamp
250cm3 beakers
Wash bottle
Filter funnel
Graduated cylinder (100 cm3)
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Procedure
NB: Wear your safety glasses.
1. Find the mass of five iron tablets. (Note that if Feospan tablets are used, each capsule
should be opened, and the contents weighed.)
2. Crush the weighed tablets in a mortar and pestle. Transfer all the ground material to a
beaker where it is dissolved in about 100 cm3 of dilute sulfuric acid.
3. All of this solution (including washings) is transferred to a 250 cm3 volumetric flask
and the solution made up to the mark with deionised water. The volumetric flask
should be stoppered and inverted several times. This is the solution containing
iron(II) ions.
4. Wash the pipette, burette and conical flask with deionised water. Rinse the burette
with the potassium manganate(VII) solution and the pipette with the iron(II)
solution.
5. Using a pipette filler, fill the pipette with the iron(II) solution and transfer the
contents of the pipette to the conical flask. Acidify this solution by adding about 10
cm3 of dilute sulfuric acid.
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6. Using a funnel, fill the burette with potassium manganate(VII) solution, making sure
that the part below the tap is filled before adjusting to zero. Because of the intense
colour of KMnO4 solution, readings are taken from the top of the meniscus.
7. With the conical flask standing on a white tile, add the solution from the burette to
the flask. Swirl the flask continuously and occasionally wash down the walls
of the flask with deionised water using a wash bottle.
8. The end-point of the titration is detected by 'the first persisting pink colour'. Note the
burette reading.
9. Repeat the procedure two or three times, adding the potassium manganate(VII)
dropwise approaching the endpoint. These accurate titres should agree to within 0.l
cm3.
10. Calculate the concentration of the iron(II) solution, and from this calculate the mass
of iron in an iron tablet.
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Table of results
Copy this table into your practical report book.
Mass of iron tablets
Rough titre
Second titre
Third titre
Average of accurate titres
Volume of iron(II) solution used in each titration
Concentration of potassium manganate(VII) solution
Concentration of iron(II) solution
Volume of iron(II) solution (in total)
Moles of iron in this volume
Mass of iron present in the tablets
Percentage of iron in the tablets
Mass of iron in an iron tablet
=
=
=
=
=
=
=
=
=
=
=
=
=
Questions relating to the experiment
1. In this experiment why is dilute sulfuric acid used rather than deionised water to
dissolve the iron tablets?
2. Why are burette readings taken from the top of the meniscus?
3. How is the end-point of the titration detected?
4. Why is a rough titration carried out?
5. Why is more than one titration carried out subsequently?
6. Prior to the titration, what steps are taken to minimise error?
7. If a brown precipitate appears during the titration, what does this indicate,
and how can it be remedied?
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Teacher Material

Not all iron tablets are suitable for use in this experiment, but those with the
Ferrograd and Feospan brand names, which contain iron(II) sulfate, work well.

Ferrograd tablets contain over twice as much iron(II) sulfate – approximately 325
mg per tablet - as Feospan – approximately 150 mg per tablet. Accordingly, a
more concentrated solution of potassium manganate(VII) solution should be used
- 0.01 M is a suitable concentration.

Some of the constituents of both of these types of iron tablets will not dissolve in
dilute sulfuric acid, but this does not matter – all of the iron(II) sulfate in each
tablet will dissolve.

When Feospan tablets are dissolved, a green solution is obtained. When Ferrograd
tablets are used, a yellow solution is obtained. During the titration, these colours
grow paler, and there is no difficulty detecting the end-point - the first persisting
pink colour.

In step 2 of the procedure, to ensure that all of the ground iron tablets are
transferred to the beaker, rinse the mortar and pestle with some of the 100 cm3 of
dilute sulfuric acid. Transfer all rinsings to the beaker, add the remainder of the
acid, rinse the mortar and pestle with deionised water, and again transfer all
rinsings to the beaker.

Nitric acid and hydrochloric acid are not suitable for use instead of sulfuric acid in
this experiment. The nitric acid is itself an oxidising agent, and the hydrochloric
acid reacts with potassium manganate(VII) solution, being oxidised itself in the
reaction.

A fuller description of titration procedure is to be found in the Student Material
relating to Mandatory Experiment 4.2.
Preparation of reagents
Potassium manganate(VII) is not a primary standard. Even the best grade of potassium
manganate(VII) is contaminated with manganese dioxide. It is therefore not possible to
make up solutions of exact concentration directly from the solid reagent. Instead a
solution must be made up of approximately the required concentration and standardised
prior to use. The potassium manganate(VII) solution for this experiment can be
standardised in mandatory experiment 4.5. Alternatively a standard solution may be
purchased. Note that solutions of potassium manganate(VII) do not keep well unless the
container is scrupulously clean and kept shielded from light. Manganese dioxide,
manganese(II) compounds, light, heat, acids and bases all hasten the decomposition of
solutions of potassium manganate(VII).
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An approximately 0.005 M solution of potassium manganate(VII) solution may be
made up as follows:
Measure out 1.58 g of potassium manganate(VII) into a beaker. Add about 500 cm3 of
deionised water. Stir and warm gently to dissolve the crystals. Pour the solution into a 2
litre volumetric flask. Add more water to the beaker to dissolve any remaining crystals of
potassium manganate(VII). Transfer the solution formed, with washings, to the
volumetric flask. Repeat this process until all crystals have been dissolved and transferred
to the volumetric flask. Carefully fill the flask to the calibration mark with deionised
water. Stopper the flask and shake to ensure a homogeneous solution.
The 1.5 M sulfuric acid solution is prepared as follows:
(Always dilute sulfuric acid by adding the acid to water and not the other way round.) 85
cm3 of the concentrated acid is added slowly to about 700 cm3 of deionised water
containing about 20 ice cubes. The mixture is stirred and made up to 1 litre in a
volumetric flask with deionised water. The flask is stoppered and inverted a number of
times.
Iron tablets containing iron(II) sulfate (ferrous sulfate) can be purchased in a pharmacy.
Quantities per working group
200 cm3 0.005 M potassium manganate(VII) solution
Iron tablets
100 cm3 1.5 M sulfuric acid
Deionised water
Safety considerations

Safety glasses must be worn.
 The use of gloves is recommended.
Chemical hazard notes
Concentrated sulfuric acid
is very corrosive to eyes and skin. Due to its very
considerable heat of reaction with water, it is essential that the acid be added to water
when it is being diluted.
Dilute sulfuric acid
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is harmful to eyes and an irritant to skin.
Solid potassium manganate(VII)
n is a powerful oxidising agent. Contact
with combustible materials may cause fire, and it forms an explosive product with
concentrated sulfuric acid. It is harmful and should not be allowed to be inhaled, ingested
or come into contact with the skin.
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Potassium manganate(VII) solution is an oxidising agent and can be a skin irritant. If it
is washed off, it may leave a brown stain that will slowly disappear.
Ammonium iron(II) sulfate is harmful if ingested in quantity, and is an eye irritant.
Sodium metabisulfite
n produces the toxic gas sulfur dioxide on contact with acids.
Using this material in the disposal of waste material from the experiment should therefore
be carried out in a fume cupboard.
Sodium carbonate (anhydrous)
lungs.
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is an irritant to eyes and skin, and its dust irritates
Disposal of wastes
Dilute with water, neutralise with anhydrous sodium carbonate, and flush to foul water
drain.
To dispose of unused potassium manganate(VII) solution, add some dilute sulfuric acid.
Then, using a fume cupboard, add sodium metabisulfite with stirring until the solution is
colourless. Dilute with excess water, and flush to foul water drain.
Specimen results
Mass of iron tablets
Rough titre
Second titre
Third titre
Average of accurate titres
Volume of iron(II) solution used in each titration
Concentration of potassium manganate(VII) solution
= 1.81 g
= 17.0 cm3
= 16.7 cm3
= 16.7 cm3
= 16.7 cm3
= 25.0 cm3
= 0.005 M
Specimen Calculations
(a) First principles method
Volume of MnO4- solution used
Moles of MnO4- used
= 16.7 cm3
= 16.7 x 0.005 / 1000
= 0.0000835 moles
Balanced equation:
MnO4- + 5Fe2+ + 8H+  Mn2+ + 5Fe3+ + 4H2O
1 mole 5 moles 8 moles 1 mole 5 moles 4 moles
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Moles of Fe2+ reacting
Volume of Fe2+ solution used
Moles/cm3 of Fe2+ used
Moles/250 cm3 of Fe2+
Mass of iron in this volume
Percentage of iron in the tablets
Mass of iron in each tablet
= 0.0000835 x 5
= 0.0004175 moles
= 25.0 cm3
= 0.0004175 / 25.0
= 0.0000167
= 0.004175
= 0.004175 x 56 g
= 0.2338 g
= mass of iron x 100
mass of tablets
= 0.2338 x 100
1.81
= 12.92%
= 0.2338 / 5
= 46.76 mg.
(b) Formula method
VA x MA x nB
= VB x MB x nA
25.0 x MA x 1 = 16.7 x 0.005 x 5
MA = 16.7 x 0.005 x 5 / (25.0 x 1)
= 0.0167 M
Volume of Fe2+ solution in total
Moles of iron in this volume
Mass of iron in this volume
Percentage of iron in the tablets
Mass of iron in each tablet

= 250.0 cm3
= 0.0167 / 4
= 0.004175
= 0.004175 x 56 g
= 0.2338 g
= mass of iron x 100
mass of tablets
= 0.2338 x 100
1.81
= 12.92%
= 0.2338 / 5
= 46.76 mg.
Since tablet packets usually give the amount of iron(II) sulfate per tablet, if a
comparison of results is required then the calculations should be repeated using
the molar mass of FeSO4 (152) instead of that of Fe (56). Alternatively, the
amount of iron in the tablet can be calculated from the data on the packet by
multiplying the supplied figure by 56 / 152.
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Solutions to student questions
1. In this experiment why is dilute sulfuric acid used rather than deionised water to
dissolve the iron tablets?
If deionised water were used, the Fe2+ in the tablets would be almost immediately
oxidised to Fe3+. The sulfuric acid prevents this occurring.
2. Why are burette readings taken from the top of the meniscus?
Because the very dark colour of the manganate(VII) solution makes the meniscus
difficult to see.
3. How is the end-point of the titration detected?
When the first permanent pale pink colour forms in the solution in the conical
flask.
4. Why is a rough titration carried out?
To determine the approximate end-point. This can then be used to get accurate
results in the subsequent titrations.
5. Why is more than one titration carried out subsequently?
To reduce experimental error, by getting the mean of the accurate titres.
6. Prior to the titration, what steps are taken to minimise error?
All glassware is washed with deionised water. The burette and pipette
respectively are rinsed with the solution they are to contain. The tap of the burette
is opened briefly to fill the part of the burette below the tap.
7. If a brown precipitate appears during the titration, what does this indicate,
and how can it be remedied?
Mn(IV) is formed, because of incomplete reduction of the Mn(VII). This should
only happen if there is insufficient sulfuric acid in the conical flask. The remedy
is to add more dilute sulfuric acid to the flask, or, preferably, to repeat the
experiment with sufficient acid present in the flask.
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