Iodometric and Iodimetric titrations

Iodometric and Iodimetric titrations are two subtypes of redox titration. As their names suggest, these are titrimetric analyses involving iodine (I₂). Iodine is a halogen from Group VII A of the Periodic Table that appears as a black solid at room temperature and pressure. However, iodine changes color as it is reduced into iodide ions. The different colors of iodine in its reduced and oxidized forms, as well as the complexation of iodine with other color-giving chemical compounds such as starch, lays the foundation of the two types of redox titrations-Iodometry and Iodimetry.

Let’s find out more about these two types of redox titrations in the proceeding sections.

What is Iodometric titration

Iodometric titration, also known as Iodometry, is a type of indirect titration.

The analyte solution is an oxidizing agent in Iodometric titrations. It is reacted with an excess amount of potassium iodide (KI). The oxidizing agent oxidizes iodide (I^–) ions into iodine (I₂). The iodine produced is then titrated with a known concentration of sodium thiosulfate (Na₂S₂O₃). Starch is used as an indicator.

The concentration of I₂ is determined, which is consequently used to determine the unknown concentration of the analyte solution used in the primary redox reaction.

What is Iodimetric titration

Iodimetric titration, unlike Iodometric titration, is a direct titration method. It is also called Iodimetry.

The analyte solution is a reducing agent in Iodimetric titrations. The analyte solution is titrated with a standard iodine solution of known concentration in the presence of starch as an indicator. The reducing agent gets oxidized while reducing I₂ into iodide (I^–) ions. The indicator changes color at the endpoint.

The volume of iodine used then helps determine the unknown analyte concentration as per stoichiometric principles.

How to perform Iodometric titration-Procedure

Step I: A specific volume of the analyte solution is pipetted out in the titration flask.

• A special titration flask is used for titrations involving iodine. This is because iodine is slightly volatile in nature; therefore, a titration flask with a funnel-shaped top and a cap is used. It is known as the iodine flask.

• A specific volume, such as 10 mL of the analyte solution, is measured and transferred to the iodine flask using a pipette.
• A small volume of an acidic solution, such as 5 mL of 6 M HNO₃ is added to it.
• Redox reactions require an acidic medium.
• The sample solution is slightly warmed on a hot plate.

Example: Copper sulfate (CuSO₄) solution is the analyte solution, so we will base this Iodometric titration on finding the exact amount of Cu²⁺ ions.

Step II: Add an excess amount of potassium iodide (KI) into the analyte solution.  

• An excess amount, such as 4 g KI is added to the iodine flask containing the analyte solution.
• Cu²⁺ ions in the analyte solution reduce iodide ions from KI and liberate iodine.
• The solution is kept in dark until further process.

Primary redox reaction of Iodometry  

The reaction of the analyte (oxidizing agent) with KI (reducing agent) :

     Net ionic equation:

In the above redox titration, the copper ions are reduced to copper solid, while the iodide ions are oxidized to iodine. The amount of iodine liberated can be used to determine the concentration of Cu²⁺ ions initially present in the solution. This is known as an indirect or back titration method.

Step III: Preparation of the standard sodium thiosulfate (Na₂S₂O₃) solution (hypo).  

• Na₂S₂O₃ solution is prepared by dissolving an accurately weighed amount of sodium salt in boiled water.
• 2.5 grams Na₂S₂O₃.5H₂O dissolved in 100 mL water gives a 0.10 M sodium thiosulfate solution.
• Boiling the water is important because any microorganisms present in it may otherwise destroy thiosulfate ions.
• A small amount of Na₂CO₃ is added to this solution to maintain a slightly basic pH.
• The solution volume is raised up to the mark with distilled water in a volumetric glass flask.
• This solution is then standardized against a primary standard such as potassium iodate (KIO₃).
• This standard solution of sodium thiosulfate is also called hypo.
• 10 mL of the mixture obtained in step II is immediately titrated with this sodium thiosulfate solution (hypo) from the burette.

Step IV: A few drops of the starch indicator are added into the iodine flask mid-way through the titration.

• The starch indicator should be freshly prepared for greater sensitivity.
• For best results, the starch indicator is added when the solution in the iodine flask obtains a faint yellow color i.e., a few moments after starting titration with sodium thiosulfate solution.
• Starch gives a blue-black color in the presence of iodine in the iodine flask.

• The blue-black color disappears as all the iodine in the flask is consumed.
• This is the equivalence point for the Iodometric titration.
• The color change marks the endpoint of the titration, and the volume of Na₂S₂O₃ used is noted.

Secondary redox reaction of Iodometry 

The redox reaction between sodium thiosulfate (reducing agent) and iodine (oxidizing agent) in the titration flask.

Net ionic equation

In the above reaction, the iodine is reduced to iodide ions while the thiosulfate ion is oxidized. 2 moles of thiosulfate react with 1 mole of iodine in the balanced equation above.

Stoichiometric calculations

How to perform Iodimetric titration-Procedure

Step I: A standard solution of iodine is prepared.  

• Solid iodine is not very water-soluble. Therefore, the standard iodine solution is prepared by dissolving iodine solid in the presence of potassium iodide (KI).
• Iodine (I₂) reacts with Iodide (I^–) ions to produce the triiodide (I₃^–) ions.

• Iodine solubilizes fairly well in the form of triiodide, and a standard solution of known concentration (such as 0.10 M) is prepared using a graduated volumetric flask.
• This solution is filled in the burette.

Step II: Titration of analyte solution with the standard iodine solution.  

• A specific analyte volume, such as 10 mL hydrogen peroxide (H₂O₂), is pipetted out in the iodine flask.
• A small volume of an acid is added.
• A few drops of the starch indicator are also added.
• The analyte mixture is titrated against the iodine solution from the burette.
• A redox reaction occurs between the analyte (reducing agent) and triiodide (oxidizing agent).
• The equivalence point is reached when all the reducing agent present in the iodine flask gets consumed.
• The indicator changes color from yellow brown to blue-black as an excess of iodine is liberated in the titration mixture. This marks the end point of the Iodimetric titration.
• The titre volume is recorded and used to determine the unknown concentration of H₂O₂.

Redox reaction of Iodimetry 

Stoichiometric calculations

What is the difference between Iodometric and Iodimetric titrations

Iodometric titration	Iodimetric titration
Indirect titration	Direct titration
The analyte is an oxidizing agent	Analyte is a reducing agent
Iodine is used as a weak reducing agent	Triiodide is involved as an oxidizing agent
Two redox reactions take place	One redox reaction takes place
Iodide ions are oxidized to release iodine. The iodine released is titrated and used to determine the initial analyte concentration	A standard iodine solution is directly used against the titrand (analyte). Iodine is reduced to iodide, and the analyte concentration is consequently determined
More commonly used	Less commonly used

Practice a quiz here to test your knowledge of Iodometry and Iodimetry.

Check out our other articles on titrations:

• Potentiometric titration
• Complexometric titration
• Acid-base titration

References

1. Madhu. 2011. “Difference Between Iodometry and Iodimetry.” In. Differencebetween.com

2. Meyiwa, Benson. 2020. ‘Iodometric and Iodimetric titration methods ‘, Journal Wetenskap Health, 1: 5-8

3. Seely, Oliver. 2020. “Iodometric determination of Cu in Brass ” In.: LibreTexts Chemistry