Lab Manual Quantitative Analytical Method, Deepak Pant [epub ebook reader .TXT] 📗

Cu Ore Unknown
1. Dry ~ 3.5 g unknown at 110 ºC for one hour. Weigh (to ± 0.1 mg) three 1-g samples and transfer each to a 250 mL beaker. Proceed separately with each sample as follows.
2. Add 5 mL of concentrated HNO3 (Caution: Handle concentrated acids carefully in the fume hood!) and heat below boiling in the hood to dissolve the unknown. Cool and add 10 mL of concentrated H2SO4 and heat until yellow SO3 gas is evolved, then cool. Cautiously add 30 mL of DI water, boil the solution for 1-2 minutes, then cool. Add concentrated ammonia dropwise until the deep blue color of the Cu(NH3)42+ complex ion appears. Then add 3 M H2SO4 to make the deep blue color just disappear. Finally, add 2 mL of 85% phosphoric acid and cool to room temperature. (This procedure adjusts the pH to a proper range.)
3. Each of the Cu2+ solutions should be completely analyzed before the next one. Add ~ 4 g of KI (to precipitate CuI) and titrate with standard Na2S2O3 until the solution is pale yellow. Add ~ 2 mL of starch solution and continue the titration to a faint blue color. Add ~ 2 g of KSCN, stir for about 30 seconds, and then finish the titration to the disappearance of the blue color. (KSCN is used because I2 and I3- tend to adsorb on the surface of solid CuI, thus becoming less available for reduction by the thiosulfate. Thus, iodometric titrations involving reduced copper tend to yield low results unless the adsorbed iodine is liberated by adding SCN- which competes with the adsorbed iodine on the surface of CuI particles.)
4. Determine the percent Cu for each trial. Report the individual results, the mean, and the standard deviation.


Iodometric Titrations
Molecular iodine (I2) is only slightly soluble in water but adding iodide, I-, produces the "triiodide" ion (I3-) in solution. Thus, KI is almost always added when redox reactions of I2 are involved in quantitative analysis.
I2(s) + I-(aq)  I3-(aq)
iodine iodide triiodide
Starch is used as the indicator in most iodometric titrations because iodine (i.e., I3-) forms an intense blue colored "starch-iodine complex."


In "iodometric" titrations, the analyte is first reduced with an excess of I-, producing I2 (actually, I3-) which turns blue in the presence of starch. Back titration of the I2 with a reducing agent like thiosulfate (S2O32-) permits the determination of the analyte. The end-point is the disappearance of the blue starch-iodine color.


I3- + 2 S2O32-  3 I- + S4O62-


Experiment: To determine ascorbic acid (vitamin C) concentration by a redox titration with potassium iodate
Object: To determine vitamin C (C6H8O6) by potassium iodate titration and to master iodometric titrations.
Introduction
A suitable method for the determination of vitamin C (C6H8O6) quantities is a titration with potassium iodate (KIO3). Potassium iodate is used as a titrant and it is added to an ascorbic acid solution that contains strong acid and potassium iodide (KI). Potassium iodate reacts with potassium iodide, liberating molecular iodine (I2):
KIO3 + 5KI + 6H+ → 3I2 + 6K+ + 3H2O (1)
As long as the solution contains ascorbic acid, the iodine produced in (1) is used up in a rapid reaction with ascorbic acid, during which dehydroascorbic acid and iodide ion are formed:
C6H8O6 + I2 → C6H6O6 + 2I- + 2H+ (2)
Potassium iodide must be added in excess to keep iodine dissolved. Once all the ascorbic acid has been consumed, any excess iodine will remain in solution. Since aqueous iodine solutions are brown in colour, iodine can act as its own indicator. However, it is quite difficult to detect endpoints using iodine coloration alone, and it is more usual to add starch, which forms an intensely blue coloured complex with iodine but not with the iodide ion.
According to the above reactions, each mole of potassium iodate added corresponds to 3 moles of ascorbic acid dehydrogenated in the sample.
Procedure
1. Pipette 25 cm3 of the provided ascorbic acid solution into a 250 cm3 conical flask, add  4 cm3 of 2M HCl,  5 cm3 of potassium iodide (KI) solution and 3 cm3 starch solution. Then titrate with the standard potassium iodate (KIO3) solution (? M) until the solution changes from brown to intense blue. Write down the standard potassium iodate (KIO3) solution volume.
2. Pipette 25 cm3 of unknown ascorbic acid sample, a kind of juice, into a 250 cm3 conical flask, then follow the same procedure of step 1 and write down the volume of the standard KIO3 solution.
Repeat the steps 1 and 2 to check the precision of the experiment. Use the mean result of step 2 to determine the concentration (mol/cm3) of ascorbic acid in the selected sample.
Note: The end point is reached when the solution turns a permanent, dark blue colour, due to the complex formed between starch and iodine. During an iodometric titration an intermediate dark blue iodine-starch complex may form momentarily, before the iodine reacts with ascorbic acid. However, if the colour disappears upon mixing, the end point has not yet been reached. Thus, magnetic stirrers or glass rod are employed in the titration to ensure proper mixing and to facilitate the reaction of iodine with ascorbic acid.

Report
1. Write up this experiment results in full (including the discussion) and hand in at the beginning of the next lab session. Marks awarded are shown in brackets.
2. Tabulate your results as follows:
Standardization Unknown sample
1st Titration 2nd Titration 1st Titration 2nd Titration
Final reading/ cm3
Initial reading/ cm3
Titre/ cm3
Mean titre ± / cm3
3. Comment upon the precision of the titrations.
4. Calculate the ascorbic acid concentration of the juice sample.
5. Briefly describe factors which influence the results from this titration.
6. What are the limitations of this method of determination of ascorbic acid?
7. Describe one other method for determining ascorbic acid in juice and decide whether it is more sensitive and more accurate than the method you have employed.
Chemicals employed in the experiment:
1). The standard KIO3 solution with a suitable concentration of-- according to the concentration of ascorbic acid in the samples.
2). Ascorbic acid solution, 0.01136M. This concentration has been chosen since it gives you 10 mg ascorbic acid in each 5.00 cm3 of the solution, which was taken as much of this solution as necessary to simulate the amount of ascorbic acid expected in the sample.
3). Potassium iodide is provided as a concentrated solution of KI in water. Use 5 cm3 of KI solution for each titration.
4). 2M HCl Use 4 cm3 of the HCl solution in each titration. The addition of acid is necessary to provide the acidic conditions required in reaction (1) above.
5). Starch solution. Use 3 cm3 of the starch solution in each titration.


3. Precipitation Titration
Precipitation titrimetry, which is one of the oldest analytical techniques, is based on reaction that yield ionic compounds of limited solubility. The number of precipitating agents that can be used is limited because of the slow rate of formation of most precipitates. The most widely used precipitation reagent is silver nitrate, and precipitation titrimetry based on silver nitrate are known as argentometric methods.

Titration curves for precipitation reactions are very similar to those involving strong acids and strong bases. Consequently, they can be derived in the exact same way. One need only substitute the solubility product of the precipitate for the ion-product constant for the water. The following diagram illustrates such a titration curve.


Clearly, one can see the resemblance of this curve to that of an acid/base titration shown below.


Both have a sharp inflection point. Furthermore, it is evident that at pre-equivalence and post-equivalence points, the p-function does not chance very much. At the equivalence point in contrast, a small volume of titrant corresponds to a large p-function change.

Figure 3 below shows the effect of product solubility on the sharpness of the endpoint. The change in the p-function at the equivalence point becomes greater as the solubility


product becomes smaller. In other words p-function increases as the reaction between the analyte and the silver nitrate becomes more complete.


The endpoint produced by a chemical indicator usually consists of a color change. Once again, we see the relation between argentometric titrations and acid/base titrations, as the requirements for an indicator for the titrations are analogous. These requirements are namely: (1) the color change should occur over a limited range in the p-function of the reagent or the analyte, and (2) the color change should take place within the steep portion of the titration curve for the analyte. Three indicators that have widespread use in argentometric titrations are the Chromate ion, Flourescein and Iron(III) ion.
The chromate ion is employed as an indicator in the Mohr method. Sodium chromate can serve as an indicator for the argentometric titrations of chlorine, bromine and cyanide ions, by reacting with silver to form silver chromate (Ag2CrO4). The brick red Ag2CrO4 precipitates at the equivalence point region and thus is a useful indicator. The Mohr method has an inherent systematic error that is correctable. At reagent concentrations lower than .1M, the chromate ion imparts an intense yellow color to the solution that hinders the identification of the red silver chromate. As a consequence, excess silver nitrate is required before precipitation can be detected and as a result, the volume is skewed. The correction that is often made to account for this error is making a blank titration of a chloride free suspension of calcium carbonate. This volume can then be used to adjust the experimental results. The Mohr method will follow these reactions:
Ag+ + Cl- AgCl (s) white
2Ag+ + CrO4-2 Ag2CrO4 (s) red

The second popular indicator is Flourescein, an adsorption indicator used in the Fajans method. An absorption indicator is an organic compound that tends to be absorbed onto the surface of the solid in a precipitation titration. In the ideal scenario, the absorption occurs near the equivalence point and results in a color change. Flourescein is a widely used absorption indicator.


In aqueous solution partial dissociation into H3O and the negatively charged Flourescein ion. The Flourescein forms a bright red silver salt. These types of indicators give a rapid, accurate and reliable titration endpoint.
The third indicator used is Iron(III) ion, used in the Volhard method. In this method, silver ions are titrated with a standard solution of thiocyanate ion, where Fe3+ serves as the indicator imparting a red color to the solution. This titration must be carried out in an acidic solution lest the Fe(III) precipitates out as a hydrated oxide. The reaction that accompany the Volhard method