Subject: Chemistry    / General Chemistry


Chemistry 250 – Analytical Chemistry

The Determination of Ascorbic Acid in Vitamin Tablets


The ascorbic acid, or vitamin C (C6H8O6), content of vitamin tablets can be determined by a

variety of techniques. This experiment will utilize a series of redox steps for analysis. Very

interesting oxidation/reduction chemistry can be studied along the way. A back titration will be

necessary for the final step.

This experiment will require three standard solutions: KIO3, KBrO3, and Na2S2O3. The first two

can be obtained as solids in pure forms that can be easily dried (without decomposition) and

weighed and therefore make excellent primary standards. Sodium thiosulfate, however, is

usually obtained in the pentahydrate form, Na2S2O3


H2O. The crystalline lumps are usually at

least partially opaque due to significant loss of the waters of hydration, which increases (but not

quantitatively) with heating. Therefore, since the exact composition is not known, it cannot be

used as a primary standard and must be titrated to determine the concentration.


In the first part of the experiment, a standardized sodium thiosulfate solution is prepared using

KIO3 as the primary standard. A standard amount of KIO3 is added to excess KI to generate the

tri-iodide ion according to the reaction

The I3

– (often referred to as iodine in solution even though elemental I2 is only very slightly soluble

in water) is the major species when I2 is in an aqueous solution of I-


The triiodide ion is a weak (and therefore selective) oxidizing agent

and can be used without an indicator if the concentration is high enough since the triiodide ion is

deep red-brown at high concentrations and the iodide ion is colorless. In dilute solutions,

however, the transition color becomes pale yellow to clear and hard to detect. Starch and the

I2 + I – I3

– K = 270 (2)


– + 2 e- 3 I – E?

= +0.54 V (3)


– + 8 I- + 6 H+ 3 I3

– + 3 H2O (1)2

triiodide ion make a deep blue colored complex, which can be used to enhance the endpoint

detection (now the deep blue to colorless). BEWARE: starch decomposes quickly and MUST be

prepared fresh frequently.

The amount of triiodide ion formed from reaction (1) is determinable if a known mass (and

therefore known number of moles) of the primary standard KIO3 is reacted with excess I- since this

reaction goes quantitatively. The I3

– can then be used to standardize a thiosulfate solution by the

following reaction:

This reaction is run in the presence of starch. Since both the thiosulfate and the tetrathionate



) ions are colorless, the endpoint can be detected as the blue to colorless change due to the

disappearance of the starch-triiodide complex. The I3

starch complex is very strong and therefore

is slow to dissociate as the stoichiometric endpoint nears. Therefore time must be allowed for its

dissociation. The E° value is large enough to ensure quantitative results. NOTE: the blue color

may reappear with time due to the air oxidation of the iodide ion.

Solutions of Na2S2O3 are prepared from the solid and include a small amount of Na2CO3 which

raises the solution pH to improve the stability and to precipitate out any trace amounts of

copper(II) that might be present, since this ion acts as a catalyst in the decomposition of thiosulfate

ion. It is also a good idea to store the solution in the dark to slow its decomposition. Unlike most

standard acids and bases, however, solutions of Na2S2O3 have a limited shelf life.

NOTE: The stoichiometry for the calculations must involve a combination of that from (1) and

(4) with KIO3 as the limiting reagent.


Now that the Na2S2O3 has been standardized, it can be used indirectly to determine the vitamin C

content of vitamin tablets. First the tablets must be crushed and the mass determined. Often

binders are present and remain suspended but do not affect the results. In some tablets, the binder

may be starch so that the characteristic color of the complex with I3

– may be seen early in the

analysis, but this will not affect the results.

An excess of Br- is added to the solution containing vitamin C and a measured amount of standard

KBrO3 solution is added. The product, Br2, reacts with vitamin C to oxidize it.

An excess of BrO3

– is added so that all of the vitamin C reacts and a quantitative excess of Br2

remains. This amount must now be determined.


– + 2 S2O3

2- S4O6

2- + 3 I – E? = +0.46 V (4)


– + 5 Br- + 6 H+ 3 Br2 + 3 H2O (5)

Br2 + (Vit C)red (Vit C)ox + 2 Br- + 2 H+ (6)3

Excess KI is added to this to form the triiodide ion

which is titrated with the thiosulfate standard from Part I above.

The endpoint is visually seen as the disappearance of the blue I3

– starch complex color.

This experimental procedure assumes that the unknown samples will contain about 100 mg of

vitamin C. This is the approximate contents of 1 cup (250 mL) of orange juice, 1 cup of broccoli

(extracted), one multiple vitamin (beware of other species in a multiple vitamin mix that might

interfere with the analysis), or one 100 mg vitamin C tablet. The current RDA (recommended

daily allowance) value for vitamin C is 60 mg.

The calculations are somewhat involved, so an outline for the procedure is given at the end of the

experiment. To make certain that you understand these, try the prelab questions before beginning

the experimental work.


1. If in Part I Step 2, you measured 0.2287 g of KIO3, what would be the molarity of the solution?

2. If in Step 4, 18.21 mL of the sodium thiosulfate solution were used for the titration, what is the

molarity of the standardized thiosulfate solution?

3. If in Part II Step 1, you measured 0.1156 g of KBrO3, what would be the molarity of the solution?

4. If Step 5 took 1.82 mL to back titrate, how many mg of vitamin C were in the sample?

5. Calculate the equilibrium constant for equation 4. (HINT: you know ?°, and from the lecture you

know the relationship between ?°, ?G, and K.)

*** Check your answers with your instructor before proceeding ***

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