experiment sodium thiosulfate titration

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Transition Metals - Iodine-Sodium Thiosulfate Titrations (A-Level Chemistry)

Carrying out a iodine-sodium thiosulfate titrations.

In order to find out the concentration of an oxidising agent , I odine-Sodium Thiosulfate titrations can be used.

Table of Contents

This involves adding an acidified solution of potassium iodid e (KI) to a solution of the oxidising agent under investigation. The iodide ions in solution will be oxidised to iodine :

I2 + 2e⁻ ⇌ 2I⁻

For example, if we were using potassium iodate (V) (KIO₃) as the oxidising agent, the reaction would be:

IO₃⁻ (aq) + 2I⁻ (aq) + 6H⁺ (aq) → 3I₂ (aq) + 3H₂O (l)

The higher the concentration of the oxidising agent, the more iodide ions will be oxidised to iodine.

In order to find out how many moles of iodine have been produced, the solution is titrated with a solution of sodium thiosulfate (Na₂S₂O₃) of known concentration . Iodine will react with the thiosulfate ions to form iodide ions once again, turning the solution from brown to colourless:

I₂ (aq) + 2S₂O₃²⁻ (aq) → 2I⁻ (aq) + 2S₄O₆²⁻ (aq)

As the thiosulfate solution is added from the burette drop by drop, the iodine solution in the conical flask will gradually become a very pale yellow as the end point is approached. The end point of the titration can therefore be difficult to see.

If we add 2cm³ of starch solution , the reaction mixture will turn dark blue to indicate that iodine is still present. Now you can continue to add sodium thiosulfate drop by drop until the blue colour disappears completely, indicating that all the iodine has just reacted .

Calculating the Concentration of the Oxidising Agent

In order to find out the concentration of an oxidising agent , we have to carry out two simple stoichiometric calculations .

Worked example: A student adds 25.0 cm³ of potassium iodate (V) solution to an excess of acidified potassium iodide solution. He then titres the resulting solution with 0.120 mol dm-³ sodium thiosulfate solution. It takes 11.0 cm³ of sodium thiosulfate solution to reach the end point in the titration. Calculate the concentration of potassium iodate. (4 marks)

Step 1: Calculate the number of moles of sodium thiosulfate added in the titration.

Number of moles = concentration x volume Number of moles = [0.120 mol dm⁻³ x 11.0 cm³]/1000 = 1.32 x 10⁻³ mol

Step 2: Calculate the number of moles of iodine that have reacted in the titration.

For this use the stoichiometry of the equation:

2 moles of thiosulfate ions are used per mole of iodine (Ratio 2:1)

Therefore, if moles of thiosulfate = 1.32 x 10⁻³ mol Then moles of iodine = 1.32 x 10⁻³ mol / 2 = 6.60 x 10⁻⁴ mol

Step 3: Calculate the number of moles of oxidising agent.

3 moles of iodine are produced for every mole of iodate ions (Ratio 3:1)

Therefore, if moles of iodine = 6.60 x 10⁻⁴ mol Then moles of iodate = 6.60 x 10⁻⁴ mol / 3 = 2.20 x 10⁻⁴ mol

Step 4: Calculate the concentration of oxidising agent.

Number of moles = concentration x volume

Concentration = number of moles / volume Concentration= (2.20 x 10⁻⁴ mol / 25.0cm³) x 1000 = 0.00880 mol dm⁻³

Iodine-Sodium Thiosulfate Titration with Copper (II) ions

An alloy is the combination of metals with other metals or elements.

An iodine-sodium thiosulfate titration can be used to calculate the percentage composition of copper metal in an alloy such as brass. The titration goes as follows:

1. Prepare a a solution of the alloy . A known mass of the alloy is first dissolved in concentrated nitric acid and the mixture made up to 250cm³ by adding deionised water. 25cm³ of the mixture is pipetted into a separate conical flask. Sodium carbonate solution is then slowly added until a white precipitate forms, indicating that any leftover acid has been neutralised. The precipitate can be removed by adding a bit of ethanoic acid .

2. Add an excess of potassium iodide solution . The iodide ions will reduce copper(II) ions in solution to copper (I) ions , forming a wash-off white precipitate of copper (I) iodide .

2Cu²⁺ (aq) + 4I⁻ (aq) → 2CuI (s) + I₂ (aq)

3. Titrate the resulting mixture with sodium thiosulfate solution. Fill a burette with sodium thiosulfate solution of known concentration and add it to the alloy mixture drop by drop until all of the iodine has reacted.

4. Calculate the percentage of copper in the alloy. The number of moles of copper can be calculated from the stoichiometric ratio of Cu²⁺ to I₂ derived from the reaction equation. This can then be used to calculate the mass of copper contained in the alloy sample used and hence its percentage composition.

Transition metals are elements in the periodic table that have partially filled d orbitals in their valence electron shells. They have unique physical and chemical properties that make them useful in various industries and applications.

An Iodine-Sodium Thiosulfate Titration is a laboratory experiment used to determine the amount of iodine present in a sample. Sodium Thiosulfate is used as the titrant, and iodine reacts with it to produce a yellow color. The reaction is monitored until the color disappears, which indicates the end point of the titration. The volume of Sodium Thiosulfate used is then used to calculate the amount of iodine in the sample.

The steps involved in an Iodine-Sodium Thiosulfate Titration are: 1. Preparation of the iodine solution: A known volume of iodine is dissolved in a solvent to make the solution to be titrated. 2. Preparation of the sodium thiosulfate solution: Sodium Thiosulfate is dissolved in water to make a solution that will be used as the titrant. 3. Titration of the iodine solution: A few drops of starch are added to the iodine solution. The sodium thiosulfate solution is then slowly added to the iodine solution while stirring. The reaction produces a yellow color, which disappears when the end point is reached. 4. Calculation of the amount of iodine: The volume of sodium thiosulfate used at the end point is recorded and used to calculate the amount of iodine in the sample.

Starch is used in an Iodine-Sodium Thiosulfate Titration as an indicator to indicate the end point of the reaction. When starch is added to the iodine solution, it reacts with iodine to form a blue-black complex. The appearance of the blue-black color indicates the end point of the titration.

The accuracy of an Iodine-Sodium Thiosulfate Titration can be determined by repeating the experiment several times and calculating the average value. The deviation of the values obtained from the average can be used to determine the accuracy of the experiment. Additionally, the use of a standardized sodium thiosulfate solution can also improve the accuracy of the experiment.

Iodine-Sodium Thiosulfate Titrations are commonly used in analytical chemistry to determine the amount of iodine in a sample. The method is widely used in various industries, such as water treatment, agriculture, and food science, to monitor the levels of iodine in water, soil, and food samples.

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Redox Titration -Thiosulfate & Iodine ( Edexcel International A Level Chemistry )

Revision note.

Core Practical 13b: Thiosulfate & Iodine Titration

Iodine-thiosulfate titrations.

• A redox reaction occurs between iodine and thiosulfate ions:

2S 2 O 3 2– (aq) + I 2 (aq) → 2I – (aq) + S 4 O 6 2– (aq)

• The light brown/yellow colour of the iodine turns paler as it is converted to colourless iodide ions
• When the solution is a straw colour, starch is added to clarify the end point
• The solution turns blue/black until all the iodine reacts, at which point the colour disappears.
• This titration can be used to determine the concentration of an oxidising agent , which oxidises iodide ions to iodine molecules
• The amount of iodine is determined from titration against a known quantity of sodium thiosulfate solution

Worked example

Analysis of household bleach

Chlorate(I) ions, ClO - , are the active ingredient in many household bleaches.

10.0 cm 3 of bleach was made up to 250.0 cm 3 . 25.0 cm 3 of this solution had 10.0 cm 3 of 1.0 mol dm -3 potassium iodide and then acidified with 1.0 mol dm -3 hydrochloric acid.

ClO - (aq) + 2I - (aq) + 2H + (aq) → Cl - (aq) + I 2 (aq) + H 2 O (l) 

This was titrated with 0.05 mol dm -3 sodium thiosulfate solution giving an average titre of 25.20 cm 3 .

2S 2 O 3 2- (aq) + I 2 (aq) → 2I - (aq) + S 4 O 6 2- (aq)

What is the concentration of chlorate(I) ions in the bleach?

    Answer:

• Therefore, 1 : 2 ratio of ClO - (aq) : S 2 O 3 2- (aq)
• Number of moles of ClO - (aq) in 250.0 cm 3 = 6.30 x 10 -4 x 10 = 6.30 x 10 -3  moles 
• 10 cm 3 bleach = 6.30 x 10 -3 moles of ClO - ions
• 1.0 dm 3 bleach = 0.630 moles of ClO - ions
• Therefore, the concentration of ClO - ions in the bleach is 0.630 mol dm -3  

General sequence for redox titration calculations

• Write down the half equations for the oxidant and reductant
• Deduce the overall equation
• Calculate the number of moles of manganate(VII) or dichromate(VI) used
• Calculate the ratio of moles of oxidant to moles of reductant from the overall redox equation
• Calculate the number of moles in the sample solution of the reductant
• Calculate the number of moles in the original solution of reductant
• Determine either the concentration of the original solution or the percentage of reductant in a known quantity of sample

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Preparation and Standardization of Sodium thiosulphate

Sodium thiosulphate solution is standardized against potassium dichromate in presence of hydrochloric acid and potassium iodide. Potassium dichromate oxidizes the iodide ion in acidic medium to equivalent amount of iodine. The iodine formed in the reaction oxidizes sodium thiosulphate giving sodium tetrathionate ion and the end point is detected by starch solution.

Hence, based on the above theory our aim is to prepare and standardize sodium thiosulphate using potassium dichromate and potassium iodide in presence of hydrochloric acid and starch solution as indicator. 1

REQUIREMENTS

Apparatus:     , chemicals:     , preparation of sodium thiosulphate solution.

Dissolve 25 g of sodium thiosulphate in CO 2 free water and make the volume upto 1000 ml. Keep the solution aside and filter to remove any cloudieness, if appears.

Preparation of starch solution

Standardization of sodium thiosulphate.

Dissolve 0.125 g of accurately weighed potassium dichromate in 25 ml of water present in a 250 ml erlenmeyer flask. Add 10 ml of hydrochloric acid and 2 g of potassium iodide, stopper, shake and keep in dark for 15 min. Add 100 ml of water to the above mixture and titrate with sodium thiosulphate using starch as the indicator. Near end point the color will be changed from dark blue to bottle green. Each ml of 0.1 M sodium thiosulphate is equivalent to 0.04904 g of K 2 Cr 2 O 7 .

Tabulation of standardization

Normality of sodium thiosulphate = (Weight of Pot.dichromate)/(Volume of sod.thiosulfate consumed)× 0.04904

From the above experiment it was evident that sodium thiosulpahte can be effectively standardized by using potassium dichromate and iodide with hydrochloric acid and starch as the visual indicator. After performing the calculations, strength of the prepared sodium thiosuplhate solution was found to be……..N.

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Standardization of Sodium Thiosulphate (Na2S2O3) solution with a Potassium Dichromate (K2Cr2O7) solution.

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Jeremiah Fernandex

Theory Aqueous iodine solutions normally contain potassium iodide (KI), which acts to keep the iodine in solution. This is due to the fact that an equilibrium is set up as follows: I 2 + I-I 3-I 3-is much more soluble than I 2 and it is as I 3-the iodine is kept in solution. In this experiment, a standard (0.06 M) solution of iodine is generated in the conical flask by reacting a standard (0.02 M) solution of potassium iodate, for each titration, with excess potassium iodide. Iodine is liberated from iodate and iodide according to the equation: IO 3-+ 5I-+ 6H + → 3I 2 + 3H 2 O The iodine solution, which is a golden-brown colour, can be titrated against sodium thiosulfate solution. The sodium thiosulfate solution is placed in the burette and, as it is added to the conical flask, it reacts with the iodine and the colour of the solution fades. When it reaches a pale yellow colour, a few drops of a freshly prepared starch solution are added. The solution becomes blue-black, and the titration is continued until it goes colourless. The titration reaction may be represented by the equation: I 2 + 2S 2 O 3 2-→ 2I-+ S 4 O 6 2-(Note that in this experiment a standard solution of iodine is used to standardise a sodium thiosulfate solution. But you also need to know that a standard solution of sodium thiosulfate can be used to standardise an iodine solution.)

Mohr Method

Nazira Mukhanbetova

Objective: Determination of chloride in solid and liquid samples by the Mohr Method Learning Outcome: • Students understand the terms volumetric analysis, morarity, molality normality and redox titration. • Students acquire the skill to prepare standard solutions of silver nitrate and sodium chloride. • Students understand the apparatus used for a titration. • Students acquire the skill to perform the precicpitation-titration in the real lab after understanding the different steps. Titration is a process by which the concentration of an unknown substance in solution is determined by adding measured amounts of a standard solution that reacts with the unknown. Then the concentration of the unknown can be calculated using the stoichiometry of the reaction and the number of moles of standard solution needed to reach the so called end point. Precipitation titrations are based upon reactions that yield ionic compounds of limited solubility. Classification of methods precipitation titration (on titrant): 1.

Mohd Zulhelmy Ahmad

dimas prasetyo

Novem Ylayron

Nurul Syafiqah

Sintia Lestari

Ajay Sharma

Determination of thiosulphate using photochemical exchange reaction of sodium nitroprusside has been investigated. It is an inexpensive, faster and convenient quantitative method. Sodium nitroprusside is a photolabile complex which undergoes photochemical ligand exchange reactions rapidly. Some recent efforts have been made to utilise such reactions for the estimation of some nitrogen containing anions and electron rich organic molecules. The progress of the reaction is observed spectrophotometrically. The effects of different parameters like pH, change of concentration of sodium nitroprusside, concentration of ligands, light intensity etc. on percentage error was investigated. The efforts were made to minimise the percentage error and some optimum conditions were obtained. Such reaction can be used for the determination of thiosulphate in the range of millimoles to micromoles; hence it is important to know whether such estimations can be done successfully and that too with the desi...

James Baker

Yana Mariana

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Master Chemistry

Iodometric Titration: Principle, Example, Advantages

Written By Adeel Abbas

Table of Contents

Iodometric titration is a method of quantitative analysis that involves the indirect determination of the concentration of an oxidizing agent in a sample solution.

This redox titration relies on the reaction between the oxidizing agent and iodide ions to produce iodine, which is then titrated using a standardized sodium thiosulfate solution.

The endpoint of the titration is indicated by the disappearance of the deep blue color of the starch-iodine complex, offering high precision in determining analyte concentration.

In our previous discussions, we have covered the broader concept of titration , which encompasses various types such as acid-base, redox, complexometric, and precipitation titrations. Each type of titration serves a specific purpose in quantitative analysis, allowing us to determine the concentration of analytes in a wide range of samples.

Moreover, we have explored the uses of titration in diverse fields, ranging from pharmaceuticals to environmental monitoring. This powerful analytical tool enables scientists to measure the concentration of substances with accuracy and reliability, providing valuable insights into chemical processes and ensuring quality control in various industries.

In the titration, the choice of indicators plays a crucial role in determining the endpoint of the titration reaction. We have delved into the fascinating world of indicators, both natural and synthetic, which serve as vital tools in visualizing the completion of titration reactions.

Natural indicators , such as litmus, turmeric, and red cabbage extract, harness the inherent color-changing properties of organic compounds. On the other hand, synthetic indicators , such as phenolphthalein and methyl orange, are carefully synthesized compounds specifically designed for titrimetric applications.

Now, let us delve into the intricacies of iodometric titration and discover its principles, procedure, advantages, and practical applications.

Principle of Iodometric Titration

Iodometric titration is a method used to measure the amount of an oxidizing agent in a solution. It works by mixing the oxidizing agent with iodide ions, which causes iodine to be released. To know when the reaction is complete, we use starch solution as an indicator. The starch forms a blue color when it reacts with iodine. We can tell that the reaction is finished when the blue color fades away.

Procedure of Iodometric Titration

To perform an iodometric titration, we need to follow a step-by-step process carefully. First, we dissolve an oxidizing substance in a suitable liquid. Then, we add sulfuric acid, hydrochloric acid, or acetic acid to make the solution acidic. This acid helps the next part of the reaction. Next, we introduce chlorine, which causes the release of iodide ions.

The freed iodide ions are then titrated using a standardized solution of sodium thiosulfate. As the sodium thiosulfate reacts with the iodine, the solution changes color from yellow to a lighter and more diluted shade. Finally, we add a starch indicator to the solution, and the color shifts dramatically from a deep blue to a pale yellow. This change in color tells us that the titration is complete.

Advantages of Iodometric Titration

Iodometric titration offers several noteworthy advantages, solidifying its position as a valuable analytical technique. Firstly, it allows for the precise determination of both reducing and oxidizing agents, expanding its analytical versatility. The visible color change associated with the formation of the starch-iodine complex provides a convenient visual indicator for identifying the endpoint of the titration.

Additionally, iodometric titration requires only small quantities of chemicals or substances, making it a cost-effective and efficient analytical method. The presence of iodine in starch imparts a vivid blue color change, facilitating the easy detection of the completion of the reaction. Moreover, iodometric titration offers the opportunity to visually observe reactivity at equilibrium points, fostering a deeper understanding of redox chemistry and its applications.

Applications of Iodometric Titration

Let us explore some practical applications of iodometric titration, demonstrating its relevance in analytical chemistry. One notable example involves the standardization of sodium thiosulfate (Na 2 S 2 O 3 ) using potassium dichromate (K 2 Cr 2 O 7 ). By accurately determining the concentration of the sodium thiosulfate solution, we can precisely quantify the concentration of the potassium dichromate, a powerful oxidizing agent.

Another application of iodometric titration involves the estimation of Cu(II) (copper(II) oxide) using a sodium thiosulfate solution. This method enables the determination of copper(II) oxide concentration in a given sample, contributing to industries reliant on copper-based materials.

Furthermore, iodometric titration finds application in the estimation of vitamin C, a potent reducing agent, through the iodometric method. This allows for the accurate quantification of vitamin C concentration in various biological samples, shedding light on its role in human health and nutrition.

The realm of iodometric titration unfolds as a captivating avenue in analytical chemistry, empowering scientists to unravel the mysteries of oxidizing agents’ concentrations. With a comprehensive understanding of its principles, procedure, advantages, and practical applications, we equip ourselves with the tools to navigate the intricate world of redox analysis.

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The chemistry of thiosulfate ions

This experiment will allow students to find out some interesting chemical reactions of sodium thiosulphate, record, observe, and understand this compound

Students will induce reactions between sodium thiosulfate and other chemicals. This practical takes place in three parts, with each part showing learners a new side of this complex substance.

This experiment should take 20 minutes.

Equipment 

• Eye protection
• Student worksheet
• Clear plastic sheet (eg OHP sheet)

Solutions should be contained in plastic pipettes. See the accompanying guidance on apparatus and techniques for microscale chemistry , which includes instructions for preparing a variety of solutions.

• Sodium thiosulphate, 0.1 mol dm –3
• Silver nitrate, 0.1 mol dm –3
• Sodium chloride, 0.1 mol dm –3
• Potassium bromide, 0.2 mol dm –3
• Potassium iodide, 0.2 mol dm –3
• Iron(III) nitrate, 0.1 mol dm –3
• Copper(II) sulfate, 0.2 mol dm –3
• Iodine solution, 0.05 mol dm –3 in 0.2 mol dm –3 KI

Health, safety and technical notes

• Read our standard health and safety guidance .
• Wear eye protection for part B and splash resistant goggles to BS EN166 3 for part C.
• Silver nitrate, AgNO 3  (aq), 0.1 mol dm –3 is an eye irritant (see CLEAPSS Hazcard HC087 ). Keep separate from organic waste containers.
• Copper(II) sulfate 0.2 mol dm –3 causes eye damage and is toxic to aquatic life (see CLEAPSS Hazcard HC027c ).
• Iron(III) nitrate, Fe(NO 3 ) 3 .9H 2 O(aq), 0.1 mol dm –3  (see CLEAPSS Hazcard HC055c ), potassium bromide, KBr(aq), 0.2 mol dm –3 , and potassium iodide, KI(aq), 0.2 mol dm –3 (see CLEAPSS Hazcard  HC047b ), are low hazard. As is Iodine solution 0.05 mol dm –3 , but this is also toxic to aquatic life (see CLEAPSS Hazard HC054 ).
• Sodium thiosulphate, 0.1 mol dm –3 is low hazard (see CLEAPSS RB087 for preparation and Hazcard HC9 5a ).

Part A — The reaction between thiosulfate ions and iodine solution:

• Cover the worksheets with a clear plastic sheet.
• Put one drop of iodine solution in the box provided on the worksheet.
• Add two drops of thiosulfate solution.
• Observe, comment and write an equation for the reaction.

What type of reaction are you observing?

Part B — The reaction between thiosulfate and silver halide:

• To form the silver halides, first put one drop of silver nitrate solution into each of the empty boxes provided on the worksheet, then add one drop of potassium bromide solution and potassium iodide solutions into the appropriate boxes. Observe and comment.
• Add three drops of sodium thiosulfate solution to each box and stir with the end of a pipette.
• Observe and comment.

What explanations can you give for your observations?

Part C — The reaction between thiosulfate ions and transition metal ions:

• Put two drops of iron(III) solution in the first box provided on the worksheet.
• Put two drops of iron(III) solution and one drop of copper(II) solution in the second box provided.
• Put two drops of copper(II) solution in the third box provided.
• Add one drop of thiosulfate solution to each box and observe carefully, especially the second box.

Observations

The brown colour of iodine is discharged as it is reduced by thiosulfate ions:

I 2 (aq) + S 2 O 3 2– (aq) → 2I – (aq) + S 4 O 6 2– (aq)

The addition of halide ions to the silver nitrate solution produces precipitates of the silver halides – pale yellow (silver bromide) and deeper yellow (silver iodide). Silver bromide dissolves readily in sodium thiosulfate solution, whereas silver iodide is less soluble. This could be used as a test to distinguish a bromide from an iodide.

The dissolution of silver bromide in thiosulfate solution is used in the fixing stage in photographic developing. Here, thiosulfate is used to dissolve unreacted silver bromide through the formation of soluble complexes such as Ag(S 2 O 3 ) 2 3– (aq).

The reaction of iron(III) with thiosulfate produces a deep violet complex anion, Fe(S 2 O 3 ) 2 – . This decomposes slowly with the fading of the violet colour:

Fe(S 2 O 3 ) 2 — (aq) + Fe 3+ (aq) → 2Fe 2+ (aq) + S 4 O 6 2– (aq)

The presence of copper(II) ions catalyses the decomposition reaction, and the violet colour fades more rapidly.

Thiosulfate reduces Cu(II) to Cu(I) and complexes the Cu(I):

2S 2 O 3 2– + 2Cu 2+ (aq) → 2Cu + (aq) + S 4 O 6 2– (aq)

2Cu + (aq) + 2S 2 O 3 2– → Cu 2 (S 2 O 3 ) 2 2– (aq)

The characteristic blue colour of copper(II) fades, leaving a colourless solution containing the complex ion Cu 2 (S 2 O 3 ) 2 2– (aq).

The chemistry of thiosulfate ions – teacher notes

The chemistry of thiosulfate ions – student sheet.

S. W. Breuer, Microscale practical organic chemistry . Lancaster: Lancaster University, 1991.

Additional information

This resource is part of our Microscale chemistry collection, which brings together smaller-scale experiments to engage your students and explore key chemical ideas. The resources originally appeared in the book Microscale chemistry: experiments in miniature , published by the Royal Society of Chemistry in 1998.

© Royal Society of Chemistry

Health and safety checked, 2018

• 16-18 years
• Practical experiments
• Redox chemistry
• Reactions and synthesis

Specification

• (c) redox reaction between Cu²⁺ and I⁻ and the determination of the liberated iodine with S₂O₃²⁻
• (o) reaction between aqueous Ag⁺ and halide ions followed by dilute aqueous NH₃

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