Iodometry is one of the most important redox titration methods. Iodine reacts directly, fast and quantitively with many organic and inorganic substances. Thanks to its relatively low, pH independent redox potential, and reversibility of the iodine/iodide reaction, iodometry can be used both to determine amount of reducing agents (by direct titration with iodine) and 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 mentioned above is
2I- ↔ I2 + 2e-
and obviously whether it should be treated as oxidation with iodine or reduction with iodides depends on the other redox system involved.
Second important reaction used excesivelly in iodometry is reduction of iodine with thiosulfate:
2S2O32- + I2 → S4O62- + 2I-
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.
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. Still, 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). Iodine solutions are prepared dissolving elemental iodine directly in the iodides solution. Elemental iodine can be prepared very pure through sublimation, but because of its high volatility it is difficult to weight. Thus use of iodine as a standard substance, although possible, is not easy nor recommended. 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 practice it is much better to use iodate as a primary substance to standardize thiosulfate, and then standardize iodine solution against thiosulfate.
Reversible iodine/iodide reaction mentioned above is
2I- ↔ I2 + 2e-
and obviously whether it should be treated as oxidation with iodine or reduction with iodides depends on the other redox system involved.
Second important reaction used excesivelly in iodometry is reduction of iodine with thiosulfate:
2S2O32- + I2 → S4O62- + 2I-
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.
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. Still, 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). Iodine solutions are prepared dissolving elemental iodine directly in the iodides solution. Elemental iodine can be prepared very pure through sublimation, but because of its high volatility it is difficult to weight. Thus use of iodine as a standard substance, although possible, is not easy nor recommended. 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 practice it is much better to use iodate as a primary substance to standardize thiosulfate, and then standardize iodine solution against thiosulfate.
end point detection with starch
Iodine in water solutions is usually colored strong enough so that its presence can be detected visually. However, close to the end point, when the iodine concentration is very low, its yellowish color is very pale and can be easily overlooked. Thus for the end point detection starch solutions are used.
Iodine gets adsorbed on the starch molecule surface and product of adsortion has strong, blue color. Exact mechanism behind adsorption and color change is not known, see for example this explanation of starch as an indicator usage.
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. In the case of titration with iodine solution we can add starch at the very beginning, as high iodine concentrations are not possible before we are long past the end point.
At the elevated temperatures adsorption of the iodine on the starch surface decreases, so titrations should be done in the cold.
Finally, it is worth of noting that starch solutions, containing natural carbohydrate, have to be either prepared fresh, or conserved with antibacterial agent like mercuric iodide HgI2.
Iodine gets adsorbed on the starch molecule surface and product of adsortion has strong, blue color. Exact mechanism behind adsorption and color change is not known, see for example this explanation of starch as an indicator usage.
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. In the case of titration with iodine solution we can add starch at the very beginning, as high iodine concentrations are not possible before we are long past the end point.
At the elevated temperatures adsorption of the iodine on the starch surface decreases, so titrations should be done in the cold.
Finally, it is worth of noting that starch solutions, containing natural carbohydrate, have to be either prepared fresh, or conserved with antibacterial agent like mercuric iodide HgI2.
Two most important solutions used in iodometric titrations are solution of iodine and solution of sodium thiosulfate. Both substances can be easily obtained in a pure form, but their other characteristics (volatility, hard to control amount of water of crystallization) make them difficult to use as a primary standards.
It is also worth of mentioning that both solutions are not quite stable and they can not be stored for a prolonged period of time. Iodine can be lost from the solution due to its volatility, while thiosulfate slowly decomposes giving off elemental sulfur. The latter process is easily visible, as thiosulfate solutions become slightly cloudy with time.
Iodine solution
It is not difficult to prepare high purity iodine through sublimation, but - due to its volatility - iodine is difficult to weight accurately, as it tends to run away. To minimize losses it should be weight in closed weighing bottle. Iodine should be kept in a closed bottles also because it is highly corrosive and it vapor can damage delicate mechanism of analytical balance.
Commonly used solutions are 0.05M (0.1 normal).
To find out amounts of substances required to prepare the solution for a needed volume use ChemBuddy concentration calculator. Download the iodine solution preparation file. Open it with the free trial version of the concentration calculator. After opening the file enter solution volume and click on the Show recipe button. Read amounts of the substances, but don't follow the general directions. It is better to use as small initial volume of the solution as possible, that is, dissolve potassium iodide in about 1/100th of the final volume of water, before adding iodine.
To minimalize losses it is important to transfer iodine to the solution as fast as possible, or even to weight a 1% excess. Solution should be kept in dark glass bottle with grinded glass stopper and standardized every few weeks or before use.
Sodium thiosulfate solution
Sodium thiosulafte can be realtively easily obtained in a pure form, but it is quite difficult to obtain samples with known amount of water of crystallization, as the exact composition of the solid is very temperature and humidity dependent. Thus solution has to be standardized against potassium iodate KIO3 or potassium dichromate.
Commonly used solutions are 0.1M (0.1 normal).
To prepare the recipe for a needed volume of the solution use ChemBuddy concentration calculator. Download the sodium thiosulfate solution preparation file. Open it with the free trial version of the concentration calculator. After opening the file enter solution volume and click on the Show recipe button.
Small amount of carbonate added helps keep solution pH above 7, which slows down thiosulfate decomposition. Some sources also call for addition of 0.5 mL chloform per liter of the solution, to stop possible growth of bacteria that can speed up decomposition process.
Starch solution
Starch solution is used for end point detection in iodometric titration.
To prepare starch indicator solution, add 1 gram of starch (either corn or potato) into 10 mL of distilled water, shake well, and pour into 100 mL of boiling, distilled water. Stir thoroughly and boil for a 1 minute. Leave to cool down. If the precipitate forms, decant the supernatant and use as the indicator solution. To make solution long lasting add a pinch of mercury iodide or salicylic acid, otherwise it can spoil after a few days.
2% sodium bicarbonate
This solution is used for neutralization of sodium arsenite, before it is titrated with iodine solution during iodine solution standardization.
It is also worth of mentioning that both solutions are not quite stable and they can not be stored for a prolonged period of time. Iodine can be lost from the solution due to its volatility, while thiosulfate slowly decomposes giving off elemental sulfur. The latter process is easily visible, as thiosulfate solutions become slightly cloudy with time.
Iodine solution
It is not difficult to prepare high purity iodine through sublimation, but - due to its volatility - iodine is difficult to weight accurately, as it tends to run away. To minimize losses it should be weight in closed weighing bottle. Iodine should be kept in a closed bottles also because it is highly corrosive and it vapor can damage delicate mechanism of analytical balance.
Commonly used solutions are 0.05M (0.1 normal).
To find out amounts of substances required to prepare the solution for a needed volume use ChemBuddy concentration calculator. Download the iodine solution preparation file. Open it with the free trial version of the concentration calculator. After opening the file enter solution volume and click on the Show recipe button. Read amounts of the substances, but don't follow the general directions. It is better to use as small initial volume of the solution as possible, that is, dissolve potassium iodide in about 1/100th of the final volume of water, before adding iodine.
To minimalize losses it is important to transfer iodine to the solution as fast as possible, or even to weight a 1% excess. Solution should be kept in dark glass bottle with grinded glass stopper and standardized every few weeks or before use.
Sodium thiosulfate solution
Sodium thiosulafte can be realtively easily obtained in a pure form, but it is quite difficult to obtain samples with known amount of water of crystallization, as the exact composition of the solid is very temperature and humidity dependent. Thus solution has to be standardized against potassium iodate KIO3 or potassium dichromate.
Commonly used solutions are 0.1M (0.1 normal).
To prepare the recipe for a needed volume of the solution use ChemBuddy concentration calculator. Download the sodium thiosulfate solution preparation file. Open it with the free trial version of the concentration calculator. After opening the file enter solution volume and click on the Show recipe button.
Small amount of carbonate added helps keep solution pH above 7, which slows down thiosulfate decomposition. Some sources also call for addition of 0.5 mL chloform per liter of the solution, to stop possible growth of bacteria that can speed up decomposition process.
Starch solution
Starch solution is used for end point detection in iodometric titration.
To prepare starch indicator solution, add 1 gram of starch (either corn or potato) into 10 mL of distilled water, shake well, and pour into 100 mL of boiling, distilled water. Stir thoroughly and boil for a 1 minute. Leave to cool down. If the precipitate forms, decant the supernatant and use as the indicator solution. To make solution long lasting add a pinch of mercury iodide or salicylic acid, otherwise it can spoil after a few days.
2% sodium bicarbonate
This solution is used for neutralization of sodium arsenite, before it is titrated with iodine solution during iodine solution standardization.
0.05M iodine standardization against arsenic trioxide
Chemical characteristics of the arsenic trioxide As2O3 make it a good candidate for a standard substance in many potentiometric methods, however, because of its toxicity it is used less and less frequently.
Arsenic oxide is dissolved in sodium hydroxide, producing sodium arsenite, which is a good reducing agent. In iodometry it is quantitatively oxidized by iodine to arsenate:
Na3AsO3 + I2 + H2O → Na3AsO4 + 2I- + 2H+
Direction of this reaction depends on pH - in acidic solutions As(V) is able to oxidize iodides to iodine. To guarantee correct pH of the solution we will add solution of sodium bicarbonate NaHCO3.
Interestingly, when using As2O3 as a standard substance in other types of redox titrations, we often add small amount of iodide or iodate to speed up the reaction. For obvious reasons in the case of iodometric titration we don't have to.
Procedure to follow:
Weight exactly about 0.15-0.20g of dry arsenic trioxide and transfer it to Erlenmayer flask.
Add 10 mL of 1M sodium hydroxide solution and dissolve solid.
Add a drop of phenolphthalein solution.
Neutralize with 0.5M sulfuric acid, adding several drops of excess acid after solution loses its color.
Add slowly (to not cause the solution to foam up) 50 mL of 2% NaHCO3 solution.
Add 5 mL of the starch solution.
Titrate swirling the flask, until a blue color persists for 20 seconds.
To calculate iodine solution concentration use EBAS - stoichiometry calculator. Download iodine standardization against arsenic trioxide reaction file, open it with the free trial version of the stoichiometry calculator.
Note, that to be consistent with the use of arsenic trioxide and its molar mass, reaction equation is not the one shown above, but
As2O3 + 2I2 + 5H2O → 2AsO43- + 4I- + 10H+
These are equivalent. Enter arsenic troxide mass in the upper (input) frame in the mass edit field above As2O3 formula. Click n=CV button below iodine in the output frame, enter volume of the solution used, read solution concentration
Chemical characteristics of the arsenic trioxide As2O3 make it a good candidate for a standard substance in many potentiometric methods, however, because of its toxicity it is used less and less frequently.
Arsenic oxide is dissolved in sodium hydroxide, producing sodium arsenite, which is a good reducing agent. In iodometry it is quantitatively oxidized by iodine to arsenate:
Na3AsO3 + I2 + H2O → Na3AsO4 + 2I- + 2H+
Direction of this reaction depends on pH - in acidic solutions As(V) is able to oxidize iodides to iodine. To guarantee correct pH of the solution we will add solution of sodium bicarbonate NaHCO3.
Interestingly, when using As2O3 as a standard substance in other types of redox titrations, we often add small amount of iodide or iodate to speed up the reaction. For obvious reasons in the case of iodometric titration we don't have to.
Procedure to follow:
Weight exactly about 0.15-0.20g of dry arsenic trioxide and transfer it to Erlenmayer flask.
Add 10 mL of 1M sodium hydroxide solution and dissolve solid.
Add a drop of phenolphthalein solution.
Neutralize with 0.5M sulfuric acid, adding several drops of excess acid after solution loses its color.
Add slowly (to not cause the solution to foam up) 50 mL of 2% NaHCO3 solution.
Add 5 mL of the starch solution.
Titrate swirling the flask, until a blue color persists for 20 seconds.
To calculate iodine solution concentration use EBAS - stoichiometry calculator. Download iodine standardization against arsenic trioxide reaction file, open it with the free trial version of the stoichiometry calculator.
Note, that to be consistent with the use of arsenic trioxide and its molar mass, reaction equation is not the one shown above, but
As2O3 + 2I2 + 5H2O → 2AsO43- + 4I- + 10H+
These are equivalent. Enter arsenic troxide mass in the upper (input) frame in the mass edit field above As2O3 formula. Click n=CV button below iodine in the output frame, enter volume of the solution used, read solution concentration
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