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

until the entire contents turn color. For accurate work, rinse down the inside surface of the flask with a stream of distilled. water near the endpoint. Half-drops of solution may be added by touching the side of the flask to the buret and rinsing down the flask.

Procedure
Overview:
A. Dissolving the antacid tablet. Since some brands of tablet do not dissolve in water, you will use a measured amount of "stomach acid" (actually about 0.2 M HCl) to dissolve the tablet. You will use more of this acid than can possibly be neutralized by the tablet, resulting in a mixture containing excess acid.
B. Neutralizing the excess acid. You will add base to the above mixture and measure the amount of base required to finish neutralizing the excess acid. The more base required in this step, the less effective was the antacid tablet at stomach acid neutralization.
Detailed Procedure:
A. Dissolving the antacid tablet.
1. Ground one antacid tablets with Porcelain Mortars & Pestles. Transfer all the grounded antacids into a weighed 250-mL flask. (If tablet is of extra strength, transfer half the amount.) Weigh the flask with the antacid and calculate the antacid weight.
2. Fill one 50 mL buret of them with 0.2 M HCl (Do follow the directions for buret use above first).
3. After recording the initial volume, allow approx. 45 mL of HCl to run into your 250 mL Erlenmeyer flask containing the grounded antacid. Do not allow the liquid level to drop below the 50 mL mark. Leave the solution in the buret!
4. Set the flask on a wire gauze supported by a ring, and heat until the solution boils. Continue boiling the solution for 5 minutes to remove any CO2 gas that may be liberated from carbonates in the antacid. Cool the flask.
B. Neutralizing the excess acid.
5. Fill a second buret with 0.2 M NaOH. Record the initial volume.
6. Add 5 drops methyl red indicator to the solution in the flask, and titrate it with the NaOH solution. Methy red is yellow if it is dissolved in a basic solution, and red in an acidic one. As the excess stomach acid is neutralized by the NaOH, the solution changes from red to yellow; the exact neutralization point is orange. If you add too much NaOH solution, you can add slightly more HCl from the first buret, and continue with NaOH addition until the endpoint is reached.
7. Read the final volumes on both burets. Repeat steps 1~7 one more time.


7. Redox Titration
Redox Indicator is a chemical compound that undergoes a color change as it goes from its oxidized form to its reduced form. Color Change for a Redox Indicator occurs over a potential range.Similar to acid-base indicators that change color with a change in protonation state
Color Change for a Redox Indicator occurs mostly over the range:


Where Eo is the standard reduction potential for the indicator and n is the number of electrons involved in the reduction


Experiment : STANDARDISING SODIUM THIOSULPHATE SOLUTON
Iodine-thiosulphate titrations are examples of redox titrations. Iodine is reduced to the iodide ion by the thiosulphate ion. The half equations for this reaction are:
I2(aq) + 2 e- 2 I-(aq) 2 S2O32-(aq) --> S4O62-(aq) + 2 e-
Overall reaction
I2(aq) + 2 S2O32-(aq) --> 2 I-(aq) + S4O62-(aq)
yellow colourless colourless colourless
Iodine-thiosulphate titrations are often used to determine the concentration of oxidising agents. The oxidising agent is reacted with an excess of potassium iodide, producing iodine. The iodine that is produced is then titrated against sodium thiosulphate, to find how much iodine was produced by the reaction of the oxidising agent with potassium iodide. Once the amount of iodine has been found, the amount of the original oxidising agent can be calculated.
In iodine-thiosulphate titrations, a standard solution of Na2S2O3 is added to the iodine solution from a burette. In these circumstances the end-point would be a change from a pale yellow solution of iodine to a colourless solution of iodide ions, which does not give a clear end-point. In order to obtain a clear end-point, as it is approached (a pale yellow colour), starch solution is added. This forms an intense blue-black colour in the presence of iodine, even in low concentrations. When all the iodine has been reduced at the end-point, the solution turns colourless (although it may change back to blue due to the re-oxidation of iodide to iodine by air - the end-point is taken to be when the solution remains colourless for at least 30 seconds). The starch should not be added until the iodine solution is pale yellow, as the iodine would become strongly adsorbed onto the starch making the titration less accurate.
Method The equation for the reaction between iodate (V) and iodide is:
IO3-(aq) + 6 H+(aq) + 5 I-(aq) --> 3 I2(aq) + 3 H2O(l)
1) The sodium thiosulphate has a concentration of about 0.1 M. Calculate the mass of potassium iodate needed to be dissolved in a 250 cm3 standard solution to give a titre of around 25 cm3. Check this value with your teacher.
2) Make up a 250 cm3 standard solution of potassium iodate.
3) Pipette 25 cm3 of the solution into a conical flask and add about 20 cm3 of 1 M sulphuric acid and about 10 cm3 of 0.5 M potassium iodide solution (both are large excesses).
4) Titrate the mixture against the sodium thiosulphate solution, using as starch as indicator near the end point.
5) Repeat until you have concordant results.
Analysis 6) Use your results to find the concentration of the sodium thiosulphate solution.
7) Calculate the maximum percentage apparatus error in the final result. Standard errors in apparatus are as follows:
balance  0.001 g
volumetric flask  0.1 cm3
25 cm3 pipette  0.1 cm3
burette (start & end readings and end point )  0.15 cm3
Evaluation 8) Comment on the reliability of your titration results.
9) Write down the mean class value for the concentration of the sodium thiosulphate (we shall assume that the class mean is the correct value).
10) Find the percentage difference between your value and the class mean value.
11) Comment on the size of your experimental error compared to the maximum apparatus error.
12) If your value is outside the apparatus error, suggest reasons why it is too high or too low (be specific, give reasons why it is too high or too low, not just why it is outside apparatus error). If your value is within apparatus error, give reasons why somebody else’s value may be too high.
13) Show that the volumes of sulphuric acid and potassium iodide added are excesses.


Iodometric Determination of Copper
INTRODUCTION
The object of this analysis is to determine the amount of copper in an ore sample (reported as the mass percentage of Cu) by an "iodometric" titration method. Like most analytical procedures based on redox chemistry, this experiment involves a lot of steps and some complex reactions that must be fully understood ahead of time.
The ore sample is first dissolved in hot, concentrated HNO3 which oxidizes Cu, yielding Cu2+ in solution. After some steps to make the solution less acidic, an excess of KI is added. This reduces Cu2+ to Cu+ in the form of insoluble CuI and, more importantly, produces I3- (which is just I2 + I-) in solution.
2 Cu2+ + 5 I-  2 CuI(s) + I3-
The liberated I3- is then titrated with a standard sodium thiosulfate (Na2S2O3) solution, reducing it back to I-. Starch is used as an indicator because it forms an intense blue color with I2 (actually, I3-). The end-point is the disappearance of the blue starch-iodine color.
I3- + 2 S2O32-  3 I- + S4O62-
The amount of Cu in the ore sample is thus determined by the overall stoichiometry of these two reactions: moles Cu2+ = moles S2O32-.
The same titration procedure, with the starch-iodine indicator, is also used to standardize the Na2S2O3. Primary standard KIO3 is pre-treated with an excess of KI which again yields I3- in solution for subsequent titration with thiosulfate.
IO3- + 8 I- + 6 H+  3 I3- + 3 H2O
I3- + 2 S2O32-  3 I- + S4O62-
Thus, the overall stoichiometry for the standardization procedure is that one mole of IO3- requires 6 moles of S2O32-.
UNKNOWN
The "copper ore" unknown will be a powdery solid containing Cu in various ionic forms, mainly CuO, along with other oxides that do not interfere in this Cu analysis.
REAGENTS and STANDARD SOLUTIONS
Preparation of 0.10 M sodium thiosulfate (Na2S2O3). Boil 500 mL of DI water for 10-15 minutes. Cool to room temperature and add 12.5 g Na2S2O3.5H2O and ~ 0.05 g Na2CO3. Stir to dissolve the salts. Transfer to a clean bottle and store in a dark place. (Boiling the water is necessary to destroy micro-organisms which can metabolize the thiosulfate ion. A small amount of Na2CO3 is added in order to bring the pH to about 9.)
Preparation of standard 0.01 M KIO3 solution. Dry ~ 1.2 g of pure KIO3 at 110 ºC for about one hour. Weigh (to ± 0.1 mg) ~ 1.1 g of KIO3 and quantitatively transfer it to a 500 mL volumetric flask. Dilute to the mark with DI water.
Preparation of starch indicator solution. Starch solutions should be freshly prepared and will be provided by the lab TA. Dissolve ~ 2 g of "soluble starch" per 100 mL of DI water. Boil the mixture gently, if necessary, to obtain a clear solution.
Standardization of 0.1 M sodium thiosulfate (Na2S2O3). Pipet four 50.00 mL aliquots of the standard KIO3 solution into 250 mL conical flasks. Proceed with each flask to completion of the titration before starting the next one. (Here the concern is to avoid error due to air oxidation of the I- ion.) Add ~ 2 g of KI and stir to dissolve. Add ~ 2 mL of 6 M HCl and titrate at once with thiosulfate until the solution becomes pale yellow. Then add ~ 2 mL of starch indicator and titrate until the blue color disappears. (If the starch indicator is added too early, the I3- tends to form a kinetically stable complex with starch from which the I3- is released too slowly, resulting in endpoint detection problems.) Repeat the titration with each of the other KIO3 samples. Determine the molarity of Na2S2O3 to four significant figures, paying careful attention to the overall stoichiometry.


PROCEDURE -- Analysis of