4.6 - Determination of the Amount of Iron in an Iron Tablet

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Experiment Summary

This experiment estimates the iron(II) content in an iron tablet using a titration with potassium manganate(VII) ( $KMnO₄$ ).

Iron(II) is oxidised to iron(III) by potassium manganate(VII) in an acidic medium.
The amount of potassium manganate(VII) required to react with the iron(II) ions is used to calculate the iron content of the tablets.

The reaction is:

\text{MnO}_4^- + 8\text{H}^+ + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}

Materials and Apparatus Required

Chemicals

0.005 M potassium manganate(VII) solution
Iron tablets
1.5 M sulfuric acid
Deionised water

Apparatus

Electronic balance
Clock glass
Mortar and pestle
Burette (50 cm³)
Pipette (25 cm³) and pipette filler
Conical flask (250 cm³)
Beakers (250 cm³)
Wash bottle
Filter funnel
White tile
Graduated cylinder (100 cm³)
Retort stand, boss-head, and clamp
Safety glasses

Safety Precautions

Wear safety glasses at all times.
Potassium manganate(VII) is a strong oxidising agent and can irritate the skin. Avoid contact and inhalation.
Sulfuric acid is corrosive. Add acid to water when diluting, not the other way around.
Wash off any chemical spills immediately with water.

Method

Weigh 5 iron tablets using the electronic balance.
Crush the tablets using a mortar and pestle, then transfer the crushed powder to a beaker.
Dissolve the crushed tablets in about 100 cm³ of 1.5 M sulfuric acid.
Stir until fully dissolved.
Transfer the solution to a 250 cm³ volumetric flask.
Rinse the beaker with deionised water and add the washings to the flask.
Make up to the 250 cm³ mark with deionised water.
Stopper and invert the flask several times.
Rinse the pipette with the iron(II) solution and transfer 25 cm³ of it into a conical flask.
Add about 10 cm³ of dilute sulfuric acid to the flask.
Fill the burette with potassium manganate(VII) solution, ensuring the portion below the tap is filled.
Record the initial burette reading.
Titrate the iron(II) solution by adding potassium manganate(VII) dropwise while swirling the conical flask.
Continue until a pale pink colour persists.
Record the final burette reading.
Repeat the titration until you obtain two concordant results (within 0.1 cm³).
Use the titration results to calculate the iron content in the tablets.

Results

Measurement	Value
Mass of iron tablets	1.81 g
Rough titre	17.0 cm³
Second titre	16.7 cm³
Third titre	16.7 cm³
Average of accurate titres	16.7 cm³
Volume of iron(II) solution used	25.0 cm³
Concentration of potassium manganate(VII)	0.005 M

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Sample Calculation

Using the formula:

\text{Moles of MnO}_4^- = \frac{16.7 \times 0.005}{1000} = \highlight[0.0000835 \, \text{moles}]

From the balanced equation:

\text{MnO}_4^- + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+}

Moles of $Fe²⁺$ reacting:

0.0000835 \times 5 = \highlight[0.0004175 \, \text{moles}]

Mass of iron in the tablets:

\text{Mass of Fe} = 0.0004175 \times 56 = 0.02338 \, \text{g} = \highlight[23.38 \, \text{mg}]

Percentage of iron in the tablets:

\frac{0.2338}{1.81} \times 100 = \highlight[12.92\%]

Mass of iron in each tablet:

\frac{0.2338}{5} = \highlight[46.76 \, \text{mg}]

Example Questions with Answers

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Q1: Why is sulfuric acid used instead of water to dissolve the iron tablets?

Sulfuric acid prevents the oxidation of $Fe²⁺$ to $Fe³⁺$ during dissolution.

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Q2: How is the endpoint of the titration detected?

The endpoint is reached when a pale pink colour persists in the conical flask.

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Q3: Why are multiple titrations performed?

To ensure accuracy by averaging the concordant results and reducing experimental error.

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Q4: What happens if a brown precipitate forms during the titration?

A brown precipitate indicates insufficient sulfuric acid.

This can be remedied by adding more dilute sulfuric acid.

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Q5: Why are burette readings taken from the top of the meniscus in this titration?

The deep colour of the potassium manganate(VII) solution makes the bottom of the meniscus difficult to see.

- Determination of the Amount of Iron in an Iron Tablet Simplified Revision Notes for Leaving Cert Chemistry

4.6 - Determination of the Amount of Iron in an Iron Tablet

Experiment Summary

Materials and Apparatus Required

Chemicals

Apparatus

Safety Precautions

Method

Results

Sample Calculation

Example Questions with Answers

Q1: Why is sulfuric acid used instead of water to dissolve the iron tablets?

Q2: How is the endpoint of the titration detected?

Q3: Why are multiple titrations performed?

Q4: What happens if a brown precipitate forms during the titration?

Q5: Why are burette readings taken from the top of the meniscus in this titration?

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