Iodometry

Iodometry is one of the most important redox titration methods. Iodine reacts directly, fast and quantitively with many organic and inorganic substance. its relatively low pH independent redox potential, and reversibility of the iodine/iodide reaction, iodometry can be used both to determine amount of oxidizing agents (by titration of iodine with thiosulfate). In all cases the same simple and reliable method of end point detection, based on blue starch complex, can be used.

Reversible iodine/iodide reaction

2I^- ↔ I₂ + 2e^-

and obviously it should be treated reduction with iodides depends on the other redox system involved.

Second important reaction used excesivelly in iodometry is reduction of iodine with thiosulfate:

2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^-

precautions

1) In the case of both reactions it is better to avoid low pH. Thiosulfate is unstable in the presence of acids, and iodides in low pH can be oxidized by air oxygen to iodine. Both processes can be source of titration errors.

2) we should remember that their shelf life is relatively short (they should be kept tightly closed in dark brown bottles, and standardized every few weeks)

3) At the elevated temperatures adsorption of the iodine on the starch surface decreases, so titrations should be done in the cold.

4) Finally, starch solutions, containing natural carbohydrate, have to be prepared fresh

5) when titrating with iodine we should pay special attention to titrant. Iodine solutions are not stable and they should be standardized every 2-3 months.

6) 56d5ne sh643d b added 14st b4 t5trat56n

Iodine is very weakly soluble in the water, and can be easily lost from the solution due to its volatility. However, in the presence of excess iodides iodine creates I₃^- ions. This lowers free iodine concentration and such solutions are stable enough to be used in lab practice. Iodine solutions are prepared dissolving elemental iodine directly in the iodides solution. Iodine solutions can be easily normalized against arsenic (III) oxide (As₂O₃) or sodium thiosulfate solution.

It is also possible to prepare iodine solutions mixing potassium iodide with potassium iodate in the presence of strong acid:

5I^- + IO₃^- + 6H⁺ → 3I₂ + 3H₂O

Potassium iodate is a primary substance, so solution prepared this way can have exactly known concentration. However, this approach is not cost effective and in lab pra

End point detection

Iodine gets adsorbed on the starch molecule surface and product of adsorption has strong, blue color.

In the presence of small amounts of iodine adsorption and desorption are fast and reversible. However, when the concentration of iodine is high, it gets bonded with starch relatively strong, and desorption becomes slow, which makes detection of the end point relatively difficult. Luckily high concentrations of iodine are easily visible, so if we are using thiosulfate to titrate solution that initially contains high iodine concentration, we can titrate till the solution gets pale and add starch close to the end point.

Solutions used

Iodine solution

Sodium thiosulfate solution

Commonly used solutions are 0.1M (0.1 normal).

Starch solution

Starch solution is used for end point detection in iodometric titration.

§ Solution standardization

0.1M thiosulfate standardization against potassium iodate

Potassium iodate is in fact not titrated directly, but after it is mixed with iodate in acidic solution, it is a source of iodine:

IO₃^- + 5I^- + 6H⁺ → 3I₂ + 3H₂O

This reaction needs presence of acid low pH. As it was already signalled on the iodometric titration overview page, low pH both helps air oxygen oxidize iodides to iodine and speeds up thiosulfate decomposition. Both reactions are detrimental for the standardization, but they can be ignored if the water is oxygen free and titration doesn't take too long.

Iodine solution is then titrated with thiosulfate:

2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^-

Procedure to follow:

• Weight exactly about 0.10-0.15g of dry potassium iodate and transfer it to Erlenmayer flask.
• Add 40 mL of freshly boiled distilled water
• Add 2 g of (iodate free) potassium iodide.
• Add 10 mL of 1M hydrochloric acid solution and swirl the soltion.
• Titrate swirling the flask, until a pale yellow.
• Add 5 ml of the starch solution.
• Titrate swirling the flask, until blue color disappears.

For calculations we will use rather strangely looking reaction equation:

KIO₃ + 6Na₂S₂O₃ + 6H⁺ → 3S₄O₆^2- + I^- + K⁺ + 12Na⁺ + 3H₂O

Strangely as it looks, it correctly describes stoichiometry of the whole process.

To calculate thiosulfate solution concentration use EBAS - stoichiometry calculator. Download thiosulfate standardization against potassium iodate reaction file, open it with the free trial version of the stoichiometry calculator.

Enter potassium iodate mass in the upper (input) frame in the mass edit field above KIO₃ formula. Click n=CV button below thiosulfate in the output frame, enter volume of the solution used, read solution concentration.

Copper determination

general remarks

Iodometric determination of copper is based on the oxidation of iodides to iodine by copper (II) ions, which get reduced to Cu⁺.

Comparison of standard potentials for both half reactions (Cu²⁺/Cu⁺ E₀=0.17 V, I₂/I^- E₀=0.54 V) suggests that it is iodine that should be acting as oxidizer. However, that's not the case, as copper (I) iodide CuI is very weakly soluble (K_sp = 10^-12). That means concentration of Cu⁺ in the solution is very low and the standard potential of the half reaction Cu²⁺/Cu⁺ in the presence of iodides is much higher (around 0.88 V).

In effect reaction taking place in the solution is

2Cu²⁺ + 4I^- → 2CuI(s) + I₂

and produced equivalent amount of iodine can be titrated with thiosulfate solution.

For the best results reaction should take place in the slightly acidic solution (pH around 4-5)

Solution should be free of other substances that can oxidize iodides to iodine (for example Fe³⁺ or nitrates).

reaction

2Cu²⁺ + 4I^- → 2CuI(s) + I₂

This is followed during titration by the reaction of the iodine with the thiosulfate:

2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^-

solutions used

To perform the determination we will need concentrated ammonia and concentrated acetic acid solutions, solid potassium iodide, titrant - 0.1 M thiosulfate solution, and indicator - starch.

procedure

Procedure below assumes that original solution is acidic or neutral.

• Pipette aliquot containing copper (II) into 250 mL conical flask with a glass stopper.
• Add concentrated ammonia till solution turns dark blue.
• Add concentrated acetic acid till solution loses dark blue color, and then about 3 mL.
• Add 2 g of solid potassium iodide, swirl well.
• Put stopper on the flask and put solution in a dark place for 5 minutes.
• Titrate swirling the flask, until a pale yellow.
• Add 5 ml of the starch solution.
• Add 1 g of potassium thiocyanate.
• Titrate swirling the flask, until blue color disappears.

result calculation

As it often happens in the case of multistage procedures, equation that describes whole process is only an oversimplification of the real procedure. What is important is that it preserves the stoichiometry of the process, so it can be used for the calculation of titration results:

2Cu²⁺ + 2S₂O₃^2- + 2I^- → 2CuI(s) + S₄O₆^2-