- Fundamentals
- Reactions
- General procedure
- Preparation of sodium thiosulfate
- Starch indicator preparation
- Sodium thiosulfate standardization
- Iodometric titration
- Applications
- References
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.
End point of all iodometric titrations or titrations without addition of starch. Source: LHcheM via Wikipedia.
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 because the brown color is turning yellowish, and this gradually becomes colorless. That is why the starch indicator is used, to further highlight this end point.
Iodometry makes it possible to analyze some oxidizing species such as hydrogen peroxides from fats, hypochlorite from commercial bleaches, or copper cations in different matrices.
Fundamentals
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, causing 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 a brown color (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 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 a source of confusion for first-time iodometry learners.
Reactions
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 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 concentration of I 3 - ions is known, so it is titrated with Na 2 S 2 O 3 to standardize it.
General procedure
Each analyte determined by iodometry has its own methodology. However, this section will address 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 how the method is performed.
Preparation of sodium thiosulfate
Commercially, this salt is 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 the 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 starch concentration, 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.
Iodometric titration
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.
The true end point of an iodometric titration shows a blue color, although lighter, similar to that of this iodine-starch solution. Source: Voicu Dragoș
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.
Applications
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.
References
- Day, R., & Underwood, A. (1989). Quantitative Analytical Chemistry. (fifth ed.). PEARSON Prentice Hall.
- Wikipedia. (2020). Iodometry. Recovered from: en.wikipedia.org
- Professor SD Brown. (2005). Preparation of Standard Sodium Thiosulfate Solution and
- Determination of Hypochlorite in a Commercial Bleach Product. Recovered from: 1.udel.edu
- Daniele Naviglio. (sf). Iodometry and Iodimetry. Federica Web Learning. Recovered from: federica.unina.it
- Barreiro, L. & Navés, T. (2007). Content and Language Integrated Learning (CLIL) Materials in Chemistry and English: Iodometric Titrations. Teacher's material. Recovered from: diposit.ub.edu