The iodometry is a technique that quantifies volumetric analysis an oxidizing agent by titration or indirect iodine titration. It is one of the most common redox titrations in analytical chemistry. Here the species of greatest interest is not properly elemental iodine, I 2 , but its iodide anions, I – , which are good reducing agents.
The I – in the presence of strong oxidizing agents, react rapidly, completely and quantitatively, resulting in an amount of elemental iodine equivalent to that of the oxidizing agent or analyte in question. Thus, by titrating or titrating this iodine with a redox titrant, commonly sodium thiosulfate, Na 2 S 2 O 3 , the concentration of the analyte is determined.
The upper image shows the end point that is expected to be observed in iodometric titrations. However, it is difficult to establish when to stop titration. This is due to the fact that the brown color turns yellowish, and it gradually becomes colorless. That is why the starch indicator is used, to further highlight this end point.
Iodometry allows the analysis of some oxidant species such as the hydrogen peroxides in fats, the hypochlorite in commercial bleaches, or the copper cations in different matrices.
Unlike iodimetry, iodometry is based on species I – , less sensitive to disproportionate or to suffer undesirable reactions. The problem is that, although it is a good reducing agent, there are no indicators that provide end points with iodide. That is why elemental iodine is not left out, but remains a key point in iodometry.
Iodide is added in excess to ensure that it completely reduces the oxidizing agent or analyte, originating elemental iodine, which dissolves in water when it reacts with the iodides in the medium:
I 2 + I – → I 3 –
This gives rise to the triiodide species, I 3 – , which stains the solution brown (see image). This species reacts in the same way as I 2 , so that when titrating the color disappears, indicating the end point of the titration with Na 2 S 2 O 3 (right of the image).
This I 3 – is titled reacting the same as the I 2 , so it is irrelevant which of the two species is written in the chemical equation; as long as the loads are balanced. Generally, this point is often confusing for first-time iodometry learners.
Iodometry begins with the oxidation of iodide anions, represented by the following chemical equation:
A OX + I – → I 3 –
Where A OX is the oxidizing species or the analyte to be quantified. Its concentration is therefore unknown. Next, the I 2 produced is valued or titled:
I 3 – + Holder → Product + I –
The equations are not balanced because they only seek to show the changes that iodine undergoes. The concentration of I 3 – is equivalent to that of A OX , so the latter is being determined indirectly.
The titrant must have a known concentration and quantitatively reduce iodine (I 2 or I 3 – ). The best known is sodium thiosulfate, Na 2 S 2 O 3 , whose titration reaction is:
2 S 2 O 3 2– + I 3 – → S 4 O 6 2– + 3 I –
Note that the iodide reappears and the tetrathionate anion, S 4 O 6 2–, is also formed . However, Na 2 S 2 O 3 is not a primary standard. For this reason, it must be standardized prior to volumetric titrations. Their solutions are evaluated using KIO 3 and KI, which react with each other in an acid medium:
IO 3 – + 8 I – + 6 H + → 3 I 3 – + 3 H 2 O
Thus, the I 3 – ion concentration is known, so it is titrated with Na 2 S 2 O 3 to standardize it.
Each analyte determined by iodometry has its own methodology. However, this section will discuss the procedure in general terms to perform this technique. The quantities and volumes required will depend on the sample, the availability of reagents, the stoichiometric calculations, or essentially the way the method is performed.
Preparation of sodium thiosulfate
Commercially, this salt is found in its pentahydrated form, Na 2 S 2 O 3 · 5H 2 O. The distilled water with which your solutions will be prepared must be boiled first, so that microbes that can oxidize it are eliminated.
Likewise, a preservative such as Na 2 CO 3 is added , so that when in contact with the acidic medium it releases CO 2 , which displaces the air and prevents oxygen from interfering by oxidizing the iodides.
Starch indicator preparation
The more dilute the concentration of the starch, the less intense the resulting dark blue color will be when coordinated with the I 3 – . Because of this, a small amount of it (about 2 grams) dissolves in a volume of one liter of boiling distilled water. The solution is stirred until clear.
Sodium thiosulfate standardization
Once the Na 2 S 2 O 3 is prepared, it is standardized. A determined quantity of KIO 3 is placed in an Erlenmeyer flask with distilled water and an excess of KI is added. A volume of 6 M HCl is added to this flask, and it is immediately titrated with the Na 2 S 2 O 3 solution .
To standardize Na 2 S 2 O 3 , or any other titrant, iodometric titration is performed. In the case of the analyte, instead of adding HCl, H 2 SO 4 is used . Some analytes require time to oxidize I – . In this time interval, the flask is covered with aluminum foil or left to stand in the dark so that the light does not induce undesirable reactions.
When the I 3 – is titrated , the brown solution will gradually turn yellowish, indicative point to add a few milliliters of the starch indicator. Immediately, the dark blue starch-iodine complex will form. If added earlier, the high concentration of I 3 – would degrade the starch and the indicator would not work.
Keep adding Na 2 S 2 O 3 until the dark blue color lightens like the image above. Just when the solution turns a light purple color, the titration is stopped and other drops of Na 2 S 2 O 3 are added to check the exact moment and volume when the color disappears completely.
Iodometric titrations are frequently used to determine the hydrogen peroxides present in fatty products; hypochlorite anions from commercial bleaches; oxygen, ozone, bromine, nitrite, iodates, arsenic compounds, periodates, and the content of sulfur dioxide in wines.
Day, R., & Underwood, A. (1989). Quantitative Analytical Chemistry . (fifth ed.). PEARSON Prentice Hall.