As the names Iodometry and Iodimetry suggest, they relate to a process where Iodine is involved. In fact, both these terms refer to different methods of using Iodine in titrations to determine the concentration of an analyte under investigation. They differ in their approach. Iodometry is an indirect titration method whereas iodimetry is a direct titration method. This is the main difference between Iodometry and Iodimetry.
What is Iodometry
As mentioned above, Iodometry is an indirect method. In this case, the Iodine, which was produced due to a prior redox reaction, is quantified through a separate titration and the concentration of the analyte that produced the Iodine is determined. The technique of Iodometry is commonly used in experiments where the amount of oxidizing agents in a water body needs to be quantified.
What happens here is, an excess amount of Iodide solution (typically Potassium Iodide) is mixed with a sample of the water that needs to be tested. Due to the oxidizing agents present in the water body, the Iodide ions get oxidized to Iodine, while the oxidizing agents get reduced. This is the initial redox reaction. Then the produced Iodine is titrated with a reducing agent such as sodium thiosulfate solution. Here, the Iodine reduces to Iodide ions while the thiosulfate ions get oxidized further. This is the second redox reaction and it is the reaction used for the titration. This is performed in the presence of a starch indicator to make it easier to recognize the end point. Iodine forms a deep-blue colour complex with starch and as the Iodine breaks down to Iodide ions, the colour disappears.
Starch granules stained with iodine – through microscope
What is Iodimetry
As mentioned in the definition, this is a direct titration method. The analyte under investigation needs to be the reducing agent. And this is directly titrated with a standard Iodine solution at the presence of a suitable indicator. Therefore, by determining the end point of the reaction, equations can be derived to obtain information of the stoichiometry and other necessary relationships between the reducing agent and Iodine which acts as the oxidizing agent in this case.
Therefore, in this case, only one redox reaction takes place unlike in the case of the Iodometric titration. However, it is more common to use Iodometric methods for analysis rather than Iodimetric methods.
Definition
In Iodometric titrations, the Iodine which has been produced as a result of a previous redox reaction is been titrated with a reducing agent such as thiosulfate ions.
In Iodimetric titrations, an Iodine solution is directly titrated with a reducing solution.
Route of Titration
An Iodometric titration is an indirect method of analysis.
Iodimetry is a direct analysis method.
No. of Redox Reactions
In Iodometry, two redox reactions occur.
In Iodimetry, only one redox reaction process takes place.
Behaviour of Iodine
In Iodometry, Iodine gets oxidized first and then reduced by a reducing agent.
In Iodimetry, Iodine only gets reduced.
Usage
Iodometry is more commonly seen in experiments.
Iodimetry is less common when compared to Iodometry.
Image Courtesy:
“Wheat starch granules” by Kiselov Yuri – Own work. Licensed under (Public Domain) via Commons
“School level titration demonstration” by UCL – Flickr. (CC BY 2.0) via Commons
Iodometry, known as iodometric titration, is a method of volumetric chemical analysis, a redox titration where the appearance or disappearance of elementary iodine indicates the end point.
Note that iodometry involves indirect titration of iodine liberated by reaction with the analyte, whereas iodimetry involves direct titration using iodine as the titrant.
Redox titration using sodium thiosulphate, Na2S2O3 (usually) as a reducing agent is known as iodometric titration since it is used specifically to titrate iodine. The iodometric titration is a general method to determine the concentration of an oxidising agent in solution. In an iodometric titration, a starch solution is used as an indicator since it can absorb the I2 that is released. This absorption will cause the solution to change its colour from deep blue to light yellow when titrated with standardised thiosulfate solution. This indicates the end point of the titration. Iodometry is commonly used to analyse the concentration of oxidizing agents in water samples, such as oxygen saturation in ecological studies or active chlorine in swimming pool water analysis.
Dilute solutions containing iodine–starch complex. Using starch as an indicator can help create a sharper color change at the endpoint (dark blue to colorless). The color above can be seen just before the endpoint is reached.
To a known volume of sample, an excess but known amount of iodide is added, which the oxidizing agent then oxidizes to iodine. Iodine dissolves in the iodide-containing solution to give triiodide ions, which have a dark brown color. The triiodide ion solution is then titrated against standard thiosulfate solution to give iodide again using starch indicator:
I−3 + 2 e− ⇌ 3 I− (E0 = +0.54 V)Together with reduction potential of thiosulfate:[1]
S4O2−6 + 2 e− ⇌ 2 S2O2−3 (E0 = +0.08 V)The overall reaction is thus:
I−3 + 2 S2O2−3 → S4O2−6 + 3 I− (Ereaction = +0.46 V)For simplicity, the equations will usually be written in terms of aqueous molecular iodine rather than the triiodide ion, as the iodide ion did not participate in the reaction in terms of mole ratio analysis. The disappearance of the deep blue color is, due to the decomposition of the iodine-starch clathrate, marks the end point.
The reducing agent used does not necessarily need to be thiosulfate; stannous chloride, sulfites, sulfides, arsenic(III), and antimony(III) salts are commonly used alternatives[2] at pH above 8.
At low pH, the following reaction might occur with thiosulfate:
S2O2−3 + 2 H+ → SO2 + S + H2OSome reactions involving certain reductants are reversible at certain pH, thus the pH of the sample solution should be carefully adjusted before the performing the analysis. For example, the reaction:
H3AsO3 + I2 + H2O → H3AsO4 + 2 H+ + 2 I−is reversible at pH below 4.
The volatility of iodine is also a source of error for the titration, this can be effectively prevented by ensuring an excess iodide is present and cooling the titration mixture. Strong light, nitrite and copper ions catalyse the conversion of iodide to iodine, so these should be removed prior to the addition of iodide to the sample.
For prolonged titrations, it is advised to add dry ice to the titration mixture to displace air from the Erlenmeyer flask so as to prevent the aerial oxidation of iodide to iodine. Standard iodine solution is prepared from potassium iodate and potassium iodide, which are both primary standards:
IO−3 + 8 I− + 6 H+ → 3 I−3 + 3 H2OIodine in organic solvents, such as diethyl ether and carbon tetrachloride, may be titrated against sodium thiosulfate dissolved in acetone.[clarification needed]
Iodine standard solution, sealed in an ampoule for iodometric analysis.
Iodometry in its many variations is extremely useful in volumetric analysis. Examples include the determination of copper(II), chlorate, hydrogen peroxide, and dissolved oxygen:
Available chlorine refers to chlorine liberated by the action of dilute acids on hypochlorite. Iodometry is commonly employed to determine the active amount of hypochlorite in bleach responsible for the bleaching action. In this method, excess but known amount of iodide is added to known volume of sample, in which only the active (electrophilic) can oxidize iodide to iodine. The iodine content and thus the active chlorine content can be determined with iodometry.[3]
The determination of arsenic(V) compounds is the reverse of the standardization of iodine solution with sodium arsenite, where a known and excess amount of iodide is added to the sample:
As2O5 + 4 H+ + 4 I− ⇌ As2O3 + 2 I2 + 2 H2OFor analysis of antimony(V) compounds, some tartaric acid is added to solubilize the antimony(III) product.[2]
Determination of hydrogensulfites and sulfites
Sulfites and hydrogensulfites reduce iodine readily in acidic medium to iodide. Thus when a diluted but excess amount of standard iodine solution is added to known volume of sample, the sulfurous acid and sulfites present reduces iodine quantitatively:
SO2−3 + I2 + H2O → SO2−4 + 2 H+ + 2 I−HSO−3 + I2 + H2O → SO2−4 + 3 H+ + 2 I−(This application is used for iodimetry titration because here Iodine is directly used)
Determination of sulfides and hydrogensulfides
Although the sulfide content in sample can be determined straight forwardly as described for sulfites, the results are often poor and inaccurate. A better, alternative method with higher accuracy is available, which involves the addition of excess but known volume of standard sodium arsenite solution to the sample, during which arsenic trisulfide is precipitated:
As2O3 + 3 H2S → As2S3 + 3 H2OThe excess arsenic trioxide is then determined by titrating against standard iodine solution using starch indicator. Note that for the best results, the sulfide solution must be dilute with the sulfide concentration not greater than 0.01 M.[2]
Determination of hexacyanoferrate(III)
When iodide is added to a solution of hexacyanoferrate(III), the following equilibrium exists:
2 [Fe(CN)6]3− + 2 I− ⇌ 2 [Fe(CN)6]4− + I2Under strongly acidic solution, the above equilibrium lies far to the right hand side, but is reversed in almost neutral solution. This makes analysis of hexacyanoferrate(III) troublesome as the iodide and thiosulfate decomposes in strongly acidic medium. To drive the reaction to completion, an excess amount of zinc salt can be added to the reaction mixture containing potassium ions, which precipitates the hexacyanoferrate(II) ion quantitatively:
2 [Fe(CN)6]3− + 2 I− + 2 K+ + 2 Zn2+ → 2 KZn[Fe(CN)6]− + I2The precipitation occurs in slightly acidic medium, thus avoids the problem of decomposition of iodide and thiosulfate in strongly acidic medium, and the hexacyanoferrate(III) can be determined by iodometry as usual.[2]
- ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
- ^ a b c d Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. (2000), Vogel's Quantitative Chemical Analysis (6th ed.), New York: Prentice Hall, ISBN 0-582-22628-7
- ^ "Chlorine by Iodometry". National Environmental Methods Index. U.S. Geological Survey.[permanent dead link]
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