The iodometry is a volumetric analysis technique that quantifies an oxidizing agent by indirect titration or titration with iodine. It is one of the most common redox titrations in analytical chemistry. Here the species of greatest interest is not properly elemental iodine, Itwo, but their iodide anions, I-, which are good reducing agents.
The I- in the presence of strong oxidizing agents, they react rapidly, completely and quantitatively, resulting in an amount of elemental iodine equivalent to that of the oxidizing agent or analyte in question. Thus, titrating or titrating this iodine with a redox titrant, commonly sodium thiosulfate, NatwoStwoOR3, 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 titling. This is because 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 hydrogen peroxides from fats, hypochlorite from commercial bleaches, or copper cations in different matrices.
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Unlike iodimetry, iodometry is based on species I-, less sensitive to disproportionate or 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:
Itwo + I- → I3-
This gives rise to the triiodide species, I3-, which stains the solution a brown color (see image). This species reacts in the same way as the Itwo, Therefore, when titrating, the color disappears, indicating the end point of the titration with NatwoStwoOR3 (right of image).
This I3- It is titled reacting the same as the Itwo, 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 a source of confusion for first-time iodometry learners..
Iodometry begins with the oxidation of iodide anions, represented by the following chemical equation:
TOOX + I- → I3-
Where toOX it is the oxidizing species or the analyte to be quantified. Its concentration is therefore unknown. Next, the Itwo produced is valued or titled:
I3- + Holder → Product + I-
The equations are not balanced because they only seek to show the changes that iodine undergoes. The concentration of I3- is equivalent to that of AOX, so the latter is being indirectly determined.
The titrant must have a known concentration and quantitatively reduce iodine (Itwo I heard3-). The best known is sodium thiosulfate, NatwoStwoOR3, whose evaluation reaction is:
2 StwoOR3two- + I3- → S4OR6two- + 3 I-
Note that the iodide reappears and the tetrathionate anion, S, is also formed.4OR6two-. However, the NatwoStwoOR3 it is not a primary pattern. For this reason, it must be standardized prior to volumetric titrations. Your solutions are assessed using KIO3 and KI, which react with each other in an acid medium:
IO3- + 8 I- + 6 h+ → 3 I3- + 3 HtwoOR
Thus, the ion concentration I3- is known, so it is titled with NatwoStwoOR3 to standardize it.
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..
Commercially this salt is in its pentahydrated form, NatwoStwoOR35HtwoO. The distilled water with which your solutions will be prepared should be boiled first, so that microbes that can oxidize it are eliminated.
Likewise, a preservative such as Na is addedtwoCO3, so that when in contact with the acidic medium it releases COtwo, which displaces the air and prevents oxygen from interfering by oxidizing the iodides.
The more dilute the starch concentration, the less intense the resulting dark blue color will be when coordinated with the I3-. Because of this, a small amount of it (about 2 grams) dissolves in a volume of one liter of boiling distilled water. Stir the solution until clear.
Prepared the NatwoStwoOR3 it proceeds to standardize it. A specified amount of KIO3 It 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 solution.twoStwoOR3.
To standardize the NatwoStwoOR3, or any other titrant, the iodometric titration is carried out. In the case of the analyte, instead of adding HCl, HtwoSW4. 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 is titled3-, the brown solution will 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 large concentration of I3- it would degrade the starch and the indicator would not work.
Na continues to be addedtwoStwoOR3 until the dark blue color lightens like the picture above. Just when the solution turns a light purple color, the titration is stopped and other drops of Na are added.twoStwoOR3 to check the exact time and volume when the color completely disappears.
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.
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