Measuring Iron in Supplement Tablets

Iron supplement tablets are sold to treat and prevent iron-deficiency anaemia. The label states how much iron each tablet contains, often as a mass of iron(II) sulfate and as a mass of elemental iron. In this practical you’ll check that claim using a redox titration with potassium manganate(VII) (potassium permanganate), one of the classic quantitative experiments in advanced chemistry.

The chemistry

The tablets contain iron in the +2 oxidation state (Fe²⁺). In acidic solution, manganate(VII) ions oxidise Fe²⁺ to Fe³⁺ and are themselves reduced to Mn²⁺:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺

Key points:

• The mole ratio is 1 MnO₄⁻ : 5 Fe²⁺. Getting this ratio right is the heart of the calculation.
• No indicator is needed. MnO₄⁻ is deep purple; Mn²⁺ is almost colourless. While Fe²⁺ remains, each drop of permanganate is decolourised. At the end point, the first permanent pale pink colour shows a tiny excess of MnO₄⁻.
• Sulfuric acid provides the H⁺ ions. Hydrochloric acid isn’t used, because permanganate would oxidise chloride ions to chlorine, using up extra titrant. Nitric acid isn’t used either, since it’s itself an oxidising agent and could oxidise Fe²⁺.

See oxidation and reduction and volumetric analysis.

Equipment and reagents

• iron supplement tablets (containing iron(II) sulfate)
• 1 mol/dm³ sulfuric acid
• standard potassium manganate(VII) solution, about 0.00500 mol/dm³ (exact concentration known)
• balance (0.01 g or better), mortar and pestle
• 250 cm³ beaker, glass rod, 250 cm³ volumetric flask, funnel
• 25.00 cm³ pipette and filler, burette, conical flasks, white tile
• distilled water, wash bottle

Safety

• Eye protection throughout.
• Sulfuric acid (1 mol/dm³) is corrosive to eyes and skin, and an irritant.
• Potassium manganate(VII) is an oxidiser and harmful; it stains skin and clothes brown.
• Dispose of solutions containing manganese as instructed, not down the sink unless permitted. See disposing of chemical waste.
• Keep supplement tablets away from young children; iron overdose is dangerous.

Method

Preparing the tablet solution

1. Weigh five tablets together and record the mass. (Using several tablets averages tablet-to-tablet variation and gives a larger titre.)
2. Grind them to a fine powder with a mortar and pestle.
3. Transfer the powder to a beaker and add about 100 cm³ of 1 mol/dm³ sulfuric acid. Stir to dissolve. (Some coating or filler may not dissolve; that’s fine, as the iron salt dissolves.)
4. Transfer the solution and all washings from the beaker, rod and mortar into a 250 cm³ volumetric flask.
5. Make up to the mark with distilled water, stopper and invert several times to mix.

The acid also slows the oxidation of Fe²⁺ by air, which would otherwise make the result too low. Carry out the titrations promptly.

Titration

1. Rinse and fill the burette with the potassium manganate(VII) solution. Because the colour is so intense, read the top of the meniscus (you can’t see the bottom), and read it the same way every time.
2. Pipette 25.00 cm³ of the tablet solution into a conical flask. Add about 10 cm³ of 1 mol/dm³ sulfuric acid to ensure excess acid.
3. Titrate, swirling constantly, until the first permanent pale pink colour lasting about 30 seconds appears.
4. Do a rough titration, then accurate ones, until you have concordant titres (within 0.10 cm³). See averages and anomalies in lab data.

Sample results

• Mass of five tablets: 2.50 g
• Concentration of KMnO₄: 0.00500 mol/dm³ (a more dilute titrant gives a titre near 20–25 cm³ for typical tablets)

Titration	Rough	1	2	3
Final reading (cm³)	23.80	23.30	46.65	23.45
Initial reading (cm³)	0.00	0.00	23.30	0.20
Titre (cm³)	23.80	23.30	23.35	23.25

Mean of concordant titres: (23.30 + 23.35 + 23.25) ÷ 3 = 23.30 cm³

Calculation

Step 1: moles of MnO₄⁻ n(MnO₄⁻) = 0.00500 × 23.30 ÷ 1000 = 1.165 × 10⁻⁴ mol

Step 2: moles of Fe²⁺ in 25.00 cm³ (ratio 1 : 5) n(Fe²⁺) = 5 × 1.165 × 10⁻⁴ = 5.825 × 10⁻⁴ mol

Step 3: moles of Fe²⁺ in 250 cm³ (scale up × 10) n(Fe²⁺) = 5.825 × 10⁻³ mol in five tablets

Step 4: per tablet 5.825 × 10⁻³ ÷ 5 = 1.165 × 10⁻³ mol

Step 5: mass of iron per tablet (Fe = 55.85) 1.165 × 10⁻³ × 55.85 = 0.0651 g = 65.1 mg of iron

Step 6: as iron(II) sulfate (FeSO₄ = 151.91) 1.165 × 10⁻³ × 151.91 = 0.177 g = 177 mg of FeSO₄ (anhydrous basis)

Step 7: percentage of iron by mass in the tablets Mass per tablet = 2.50 ÷ 5 = 0.500 g % Fe = 0.0651 ÷ 0.500 × 100 = 13.0% The rest of the tablet is water of crystallisation, sulfate, fillers, binders and coating. Sense check: the iron salt (177 mg) is well below the tablet mass (500 mg), as it must be.

See titration calculations.

Comparing with the label

Supplement labels may state:

• mass of dried ferrous sulfate (partly dehydrated FeSO₄, roughly FeSO₄·1–1.5H₂O), and/or
• mass of elemental iron.

Always compare iron with iron. For example, a label stating “200 mg dried ferrous sulfate, equivalent to 65 mg iron” should be checked against your mass of iron per tablet. Differences of a few per cent are normal; manufacturers work within tolerance limits.

Sources of error

Error	Effect on calculated iron	Improvement
Fe²⁺ oxidised by air before titrating	too low	titrate promptly; keep acidified
Not enough acid in the flask	brown MnO₂ forms; end point unclear, titre unreliable	add excess sulfuric acid
Overshooting the end point	too high	add dropwise near the end point, swirling
Loss of solid during transfer	too low	rinse all equipment into the volumetric flask
Reading the burette inconsistently	random error	always read the top of the meniscus for permanganate
Permanganate concentration drifts	systematic	standardise KMnO₄ against a primary standard (e.g. sodium ethanedioate)

Uncertainty

With a 25.00 cm³ pipette (±0.06), a 250 cm³ flask (±0.3) and a titre of 23.30 cm³ (±0.10):

• pipette: 0.24%
• flask: 0.12%
• titre: 0.43%

Total ≈ 0.8% (plus the uncertainty of the KMnO₄ concentration). So the result is roughly 65.1 ± 0.5 mg of iron per tablet, which agrees well with a label stating 65 mg. See measurement uncertainty practice problems.

Alternative method: colorimetry

Iron can also be measured colorimetrically. Fe²⁺ forms an intense orange-red complex with 1,10-phenanthroline, or Fe³⁺ forms a blood-red complex with thiocyanate ions. A calibration graph from standard iron solutions gives the concentration in a diluted tablet solution. This method is more sensitive (suitable for small amounts) but usually less precise than a well-performed titration at these concentrations. See analysing food dyes with a colorimeter for the calibration technique.

Questions

1. Why is no separate indicator needed?
2. Why is dilute sulfuric acid used rather than hydrochloric acid?
3. Why were five tablets used rather than one?
4. What would happen to the result if the tablet solution were left open to air for a day before titrating?
5. The label says each tablet contains 65 mg of iron. A student finds 58 mg. Suggest two possible reasons.

Key takeaways

• MnO₄⁻ oxidises Fe²⁺ to Fe³⁺ in acidic solution in a 1 : 5 ratio.
• The titration is self-indicating: the end point is the first permanent pale pink.
• Use sulfuric acid, not hydrochloric or nitric acid.
• Scale up from the pipetted portion to the whole solution, then divide by the number of tablets.
• Compare iron with iron when checking against the label, and evaluate errors such as air oxidation.