Measuring Vitamin C in Fruit Juice

Which contains more vitamin C: fresh orange juice, carton orange juice, or a vitamin-C-enriched squash? And what happens to the vitamin C if the juice is left open or heated? This practical answers those questions with a redox titration. Vitamin C (ascorbic acid) is a good reducing agent, and it reacts rapidly and completely with iodine. By measuring how much iodine solution a juice sample can decolourise, you can calculate how much vitamin C it contains.

Purpose

To determine the concentration of vitamin C in fruit juices by titration with iodine solution, using starch as the indicator.

Principle

The redox reaction

Ascorbic acid, C₆H₈O₆, is oxidised by iodine to dehydroascorbic acid, C₆H₆O₆, while iodine is reduced to iodide ions:

C₆H₈O₆ + I₂ → C₆H₆O₆ + 2H⁺ + 2I⁻

• Ascorbic acid loses two hydrogen atoms (two electrons and two protons) — it’s oxidised.
• Iodine gains two electrons — it’s reduced (oxidation number 0 → −1).

The reaction ratio is 1 : 1: one mole of ascorbic acid reacts with one mole of I₂. (See oxidation and reduction and vitamin C chemistry.)

Iodine is only slightly soluble in water, so the titrant is iodine dissolved in potassium iodide solution, where it exists mainly as the triiodide ion, I₃⁻. This behaves the same way in the calculation.

The indicator

As long as ascorbic acid remains in the flask, every drop of iodine added is instantly reduced to colourless iodide. Once all the ascorbic acid has reacted, the next drop of iodine stays unreacted. Starch detects it: iodine (as I₃⁻ and longer polyiodides) slips inside the amylose helix and forms an intense blue-black complex (see the starch–iodine test). The endpoint is the first permanent blue-black colour that lasts for about 20 seconds.

Equipment and reagents

• 50 cm³ burette, clamp and stand, white tile
• 25 cm³ (or 10 cm³) volumetric pipette and filler
• 250 cm³ conical flasks
• Iodine solution, about 0.005 mol dm⁻³ in potassium iodide (supplied)
• Standard ascorbic acid solution (e.g. 1.00 g dm⁻³, freshly made) for standardisation
• 1 % starch solution (freshly made)
• Fruit juices: fresh-squeezed, carton, “juice drink”, etc.
• Filter funnel and paper or muslin (for pulpy juice)
• Eye protection

Safety

• Wear eye protection.
• Iodine solution stains skin and clothing and is an irritant; clean spills promptly.
• Don’t drink any of the juice samples in the laboratory.

Method

Part A: standardising the iodine solution

Iodine solution slowly loses strength (iodine is volatile), so its exact concentration should be found on the day, using a known ascorbic acid solution.

1. Rinse and fill the burette with iodine solution. Record the initial reading to 0.05 cm³.
2. Pipette 25.0 cm³ of standard ascorbic acid solution into a conical flask. Add about 1 cm³ of starch solution.
3. Titrate with iodine, swirling constantly, until the first permanent blue-black colour. Record the final reading.
4. Repeat until you have concordant titres (within 0.10 cm³ of each other). See titration errors for good technique.

Part B: titrating the juices

5. Filter pulpy juice through muslin or filter paper.
6. Pipette 25.0 cm³ of juice into a clean conical flask. Add 1 cm³ of starch solution.
7. Titrate with the standardised iodine solution to the first permanent blue-black. With orange juice, the endpoint is a change from orange to a dark brown-blue; hold the flask over the white tile and compare it with an untitrated sample.
8. Repeat to get concordant titres. Repeat for each juice.

Carry out titrations promptly after opening or preparing each juice, because vitamin C oxidises in air.

Calculation

Step 1: concentration of iodine

Standard ascorbic acid: 1.00 g dm⁻³. M(C₆H₈O₆) = 176.12 g mol⁻¹ (check with the molar mass calculator).

Concentration = 1.00 ÷ 176.12 = 5.678 × 10⁻³ mol dm⁻³.

Moles in 25.0 cm³ = 5.678 × 10⁻³ × 0.0250 = 1.420 × 10⁻⁴ mol.

Suppose the mean titre was 27.40 cm³. Since the ratio is 1 : 1, moles of I₂ = 1.420 × 10⁻⁴ mol.

[I₂] = 1.420 × 10⁻⁴ ÷ 0.02740 = 5.18 × 10⁻³ mol dm⁻³.

Step 2: vitamin C in the juice

Suppose 25.0 cm³ of fresh orange juice needed a mean titre of 11.20 cm³.

Moles of I₂ = 5.18 × 10⁻³ × 0.01120 = 5.80 × 10⁻⁵ mol = moles of ascorbic acid.

Mass of ascorbic acid in 25.0 cm³ = 5.80 × 10⁻⁵ × 176.12 = 0.01022 g = 10.2 mg.

Per 100 cm³: 10.2 × 4 = 40.9 mg per 100 cm³.

That’s in the typical range for fresh orange juice (often quoted at around 30–60 mg per 100 cm³). A 200 cm³ glass would contain about 80 mg — close to or above the daily reference intake in many countries. For more practice with calculations like these, see titration calculations.

Results table (example)

Juice	Mean titre / cm³	Vitamin C / mg per 100 cm³
Freshly squeezed orange	11.20	40.9
Carton orange (just opened)	8.95	32.7
Carton orange (opened 3 days)	6.10	22.3
Apple juice (not fortified)	0.60	2.2

(Illustrative values — your results will differ.)

Sources of error

• Other reducing agents: juices contain other substances that can react with iodine (some polyphenols, sulfur dioxide preservatives), which makes the vitamin C appear higher than it is.
• Coloured juices mask the endpoint. Diluting the juice with a known volume of water, or using a white tile and a comparison flask, helps.
• Air oxidation: vitamin C oxidises when exposed to air, light and heat, especially with traces of copper or iron ions. Delays make results lower.
• Iodine loss: iodine evaporates from the solution, so its concentration drifts — hence standardisation on the day.
• Pulp blocking the pipette or trapping vitamin C; filter first.
• Starch solution that is old or added too early in large amounts can give a sluggish endpoint.

How precise is the result?

Every reading carries an uncertainty. A typical 25.0 cm³ pipette is accurate to about ±0.06 cm³, and each burette reading to ±0.05 cm³. Because a titre is the difference between two readings, its uncertainty is ±0.10 cm³.

For the orange juice titre of 11.20 cm³:

• burette: 0.10 ÷ 11.20 × 100 = 0.89 %
• pipette: 0.06 ÷ 25.0 × 100 = 0.24 %

Adding these gives roughly ±1.1 % from the glassware alone, or about ±0.5 mg per 100 cm³. That’s small compared with the errors from other reducing agents in the juice, air oxidation and judging the endpoint. So when two juices differ by only 1–2 mg per 100 cm³, the difference may not be real. Bigger titres give smaller percentage errors, which is why a very dilute iodine solution is used for low-vitamin juices.

Improvements and extensions

• Titrate a juice before and after heating to 80 °C for 10 minutes, or after standing open for 24 hours, to see how vitamin C breaks down.
• Compare fresh fruit (juiced and filtered) with cartons and cordials.
• Use DCPIP (a blue dye decolourised by vitamin C) as an alternative method and compare the results.
• Calculate the percentage uncertainty in the burette and pipette readings and combine them.

Key takeaways

• Vitamin C is a reducing agent; it reacts with iodine in a 1 : 1 redox reaction.
• Starch turns blue-black at the endpoint, when excess iodine appears.
• Standardise the iodine solution against a known ascorbic acid solution on the day.
• Moles of I₂ = moles of ascorbic acid; multiply by 176.12 g mol⁻¹ for mass.
• Vitamin C is easily lost to air, heat and time, which you can measure with this method.