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CHEMISTR CHEMISTRY Y LABORA LABORAT TOR ORY Y [CHEM 102]

NAME: NO. : 05 EXPERIMENT

DATE:

ID: NAME OF THE ST STAND AND ANDARDIZA ARDIZA ARDIZATION TION EXPERIMENT SECTION (GR (GROUP): OUP):

OF

SODIUM

THIOSULPHA THIOSULPHATE TE SOLUTION WITH ST STAND AND ANDARD ARD

PO POT TASSIUM

DICHR DICHROMA OMA OMATE TE

SOL SOLUTION UTION

BY

IODOMETRIC METHOD THEOR THEORY Y The process by which estimation of oxidant or reductant is made by the use of liberated iodine is called iodometric titration. In this experiment, a standard solution of potassium dichromate is treated with excess potassium iodide in acidic solution (hydrochloric acid) and the liberated iodine is titrated with sodium thiosulphate solution which is to be standardized. The reactions are:

=

K2Cr2 O7 + 14 H+ + 6 e 2 KI

=

2Cr3+ + 7 H2O + 2K+

[reduction half reaction]

I2 + 2 e - + 2K+ …………× 3 [oxidation half reaction]

K2 Cr2 O7 + 6 KI + 14 H+ = 2Cr3+ + 3I2 + 7 H2O + 8K+ Or, K2 Cr2 O7 + 6 KI + 14 HCl = 2CrCl3 + 3I2 + 7 H2O + 8KCl 2 Na2S2O3 I2 + 2e -

=

Na2S4O6 + 2Na+ + 2 e -

=

2I-

2Na2S2O3 + I2 = Na2S4O6 + 2NaI Now, from the above reactions we get, 1 mole K2Cr2O7 ≡ 3 mole I2 ≡ 6 mole Na2S2O3

PREPARATION OF 100ml 0.01M POTASSIUM DICHROMATE SOLUTION Molecular weight of K2Cr2O7 = 294 gm 1000ml 1M K2Cr2O7 solution contains = 294gm K2Cr2O7 294 × 0.01 ×100 Chem-102

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100ml 0.01M K2Cr2O7 solution contains =

gm K2Cr2O7

1000 = 0.29gm K2Cr2O7 Amount of K2Cr2O7 taken = . . . . . . . . gm Weight taken (gm) Strength of the prepared K2Cr2O7 solution =

× 0.01M Weight to be taken (0.29gm)

=

× 0.01M 0.29gm = ………… M

Transfer . . . . . . . gm of pure K2Cr2O7 crystals in a 100ml volumetric flask, then dissolve it with distilled water and make it upto the mark.

PRO PROCEDURE CEDURE 1. Dissolve 1gm of pure sodium bicarbonate (NaHCO 3) in about 50ml distilled water in a conical flask.

2. Add 4ml of 12% solution of iodate-free potassium iodide (KI) and shake well for thorough mixing. 3. Now add about 4 ml of concentrated HCl acid slowly while rotating the flask in order to mix the liquids well.

4. Pipette out 10ml of standard dichromate (K2Cr2O7) solution and pour into the same flask. Shake gently for thorough mixing. Cover the flask with a watch glass and allow standing in the dark for 5 minutes for completion of the reaction. The solution will be deep brown. In the mean time, fill the burette with the supplied thiosulphate solution in the appropriate manner. Rinse the watch glass collecting the rinse water in the conical flask and dilute with about 100ml distilled water.

5. Titrate the liberated iodine with the thiosulphate solution while shaking the flask until the brown color fades (light yellow).

6. Add 2ml of starch solution. At the end-point, the deep blue color of the starch-iodine complex disappears leaving the light green color of the chromic ion.

7. Calculate the strength of the supplied thiosulphate solution.

Table: Standar Standardization dization of Na2S2O3 solution against K2Cr2O7 solution

Chem-102

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No. of Observation

Volume of K2Cr2O7 solution (ml)

Burette Reading Initial (ml)

Final (ml)

Mean (ml)

Difference (ml)

1 2 3

CALCUL CALCULA ATION We know, V Na2S2O3 ×

S Na S O = 6 ×V K Cr O × S K Cr O 2

2

3

2

2

7

2

2

RESUL RESULT T The strength of supplied thiosulphate (Na2S2O3) solution is . . . . . . . . . M.

Chem-102

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7

ERROR The accurate result is …………M. Known Value ~ Observed Value Percentage of error =

× 100 Known Value

=

× 100

= …………

CAUTION! There are a few significant sources of error in carrying out the experiment. So the following precautions should be taken: 1. The solution should be strongly acidic. Insufficient acidity causes incomplete reduction of dichromate by iodide ion. 2. In order to avoid oxidation of hydrochloric acid by air, a reducing atmosphere should be provided. This may conveniently be done by adding some solid sodium bicarbonate or sodium carbonate to the solution. 3. Iodine is very weakly soluble in the water, and can be easily lost from the solution due to its volatility. The flask should, therefore, be kept closed except during the titration and vigorous shaking should be avoided. 4. As the reaction between dichromate and iodide is not instantaneous, after adding the reagents the solution should be allowed to stand for a while. The solution should not be allowed to stand for a long time as air oxidation of the iodide may take place. 5. KI is generally used to produce I2 to run the reaction through iodometric method. 6. 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. If we add starch, 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. This is why for the end point detection starch solutions are used. 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. Thus at the end point some absorbed iodine may remain un-titrated giving erroneous end point. This is why starch is added at the last moment when iodine concentration is very small. Chem-102

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7. At the elevated temperatures adsorption of the iodine on the starch surface decreases, so titration should be carried out at the temperature that does not exceed 300C- 350C.

DISCUSSION

[Write, in brief, about the precautions taken, causes for deviation of the experimental result from the expected value on the basis of the following major matters. i) washing and rinsing all apparatus ii) weight measurement iii) Solution preparation iv) maintaining lower meniscus of the surface level of liquid on the mark in the volumetric flask and eye level at the same plane v) maintaining lower meniscus of the surface level of liquid on the mark in the burette and eye level at the same plane vi) rate of dropping vii) detection of color change. These are just examples. Each titration has its own quirks. They are usually related to chemical characteristics of titrant and other substances involved. Discussion should be written in past-passive form.]

Chem-102

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…………………………… Signature of the faculty SOME POINTS ON SOURCES OF ERRORS TITRATION ERROR After the reaction between the substance and the standard solution is practically complete, the indicator should give a clear visual change (either a color change of the formation of turbidity) in the liquid being titrated. The point at which this occurs is called the end point of the titration. In the ideal titration the visible end point will coincide with the stoichiometric or theoretical end point. In practice, however, a very small difference usually occurs; this represents the titration error. There are thousands of possible random errors that can't be adjusted for. Some of them are typical human errors that can be limited by sticking to lab procedures, but as long as there is a human operator involved, they will be never completely eliminated. Some of possible cases are: 1.

Misjudging the color of the indicator near the end point - this is probably the most common one. Not only color change is sometimes very delicate and slow, but different people have different sensitivity to colors. This is not the same as being color blind, although these things are related.

2.

Misreading the volume - at any moment, and due to whatever reason. This can be for example a parallax problem (when someone reads the volume looking at an angle), or error in counting unmarked graduation marks. When reading the volume on the burette scale it is not uncommon to read both upper and lower value in different lighting conditions, which can make a difference.

3.

Using contaminated solutions - for example when two different solutions are transferred using the same pipette and pipette is not rinsed with distilled water in between.

4.

Using diluted titrant and diluted titrated solution - if the burette and/or pipette was not rinsed with transferred solution after being rinsed with distilled water. In effect titrant (or tittrated substance) is slightly diluted.

5.

Using solutions of wrong concentration - titrant we use may have different concentration then expected. This can be due to incorrect standardization, error in copying the concentration, contamination of the bottle content, titrant decomposition, solution being kept in open bottle and partially evaporated and so on.

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6.

Using wrong amount of indicator - in the case of single color indicators amount added can shift end point.

7.

Using dirty glass - if glass was not properly cleaned before use it may be contaminated with old reagents, which can react with new ones, changing their concentration. Also, dirty glass is not properly wetted by the solutions and they can form droplets on the glass surface making exact volume measuring impossible.

8.

Rinsing burette and/or pipette with wrong solution - if the burette or pipette is not dry before use, it has to be rinsed with the solution that will be transferred. Using just distilled water for rinsing will mean transferred solution is slightly diluted. Obviosuly it is important only when transferring sample, titrant or stoichiometric reagents used for back titration. Small errors in amounts of other substances (buffers, acids used to lower pH in redox titrations, solutions masking presence of inteferring substances and so on) are not that important.

9.

Not filling burette properly - if there is an air lock in the burette stopcock it can block the flow of the titrant, but it can also at some moment flow with the titrant; after that we have no idea what was the real volume of solution used.

10. Not transferring all solid/liquid when preparing samples - it may happen that part of the

solid was left in the funnel during transferring it into flask, or it was simply lost. It is also not uncommon to forget to rinse walls of the glassware after solution was transferred - it may happen both to solution pipetted to some vessel, or to titrant that formed droplet on the flask wall and was not rinsed with distilled water. If the pipette is not clean, some of the solution can be left inside in form of drops on the glass. 11. Transferring excess volume of liquid - by blowing pipette for example, or by incorrectly

leveling meniscus with the mark on the single volume pipette. 12. Not transferring all the volume - shaken pipette may lose a drop of the solution when it is

being moved between flasks, one may also fill the single volume pipette leveling not the meniscus, but the upper edge of the solution with pipette mark. 13. Using wrong reagents - sounds stupid, but happens now and then. Too many possibilities to

list, but we have to remember - if the reaction doesn't proceed as expected, it won't hurt to check if burette is not filled with something different than expected. Or perhaps there is no indicator in the solution? 14. Titrating at wrong temperature (other then glassware was calibrated for)- This is a very

common problem. Quite often we have no choice other, then to calibrate the glass once again. This is time consuming and - especially in the student lab - almost impossible without additional arrangements. 15. Titrating at wrong temperature (other then the method was designed for)- Some

indicators are sensitive to temperature changes. Some reactions need correct temperature range to keep stoichiometry (avoid side reactions). 16. Losing solution - too vigorous swirling can end in liquid splashing from the titration flask

before the end point had been reached. It may also happen that some titrant lands on the table instead of inside the flask. Chem-102

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17. Leaking burette - sometimes burettes leak slowly enough to allow titration, but will lose

several tenths of milliliter if left for several minutes after titrant level has been set to zero and before titration started. These are just examples. Every day in every lab in the world old mistakes are repeated and new cases are recorded. Finally, each titration has its own quirks. They are usually related to chemical characteristics of titrant and other substances involved.

Now, solve the following problems:

1. State the fundamental principle of iodimetric and iodometric titration. 2. Why is starch used as an indicator in titrations involving iodine? 3. Why is it added at the last stage of titration? 4. Why does deep-blue coloration appear after adding starch? 5. Why does light green color appear at end point? 6. Why light green color is not appeared before the end point? 7. Why is the end point detected by the appearance of light green color of chromic ion (Cr3+)?

8. What are the sources of error and steps of elimination against each of the errors? 9. Why do you add) i) sodium bicarbonate ii) conc. HCl iii) KI? 10. Write down the reactions. Which is oxidant and reductant and why? 11. Why the titration’s temperature does not exceed 300C- 350C?

Chem-102

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