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The purity of a sample of picric acid, C6H3N3O7, isdetermined by controlled-potential coulometry, convertingthe picric acid to triaminophenol, C6H9N3O. A 0.2917-gsample of picric acid is placed in a 1000-mL volumetric flask and diluted to volume. A 10.00-mL portion of this solution istransferred to a coulometric cell and diluted till the Pt cathodeis immersed. The exhaustive electrolysis of the samplerequires 21.67 C of charge. Report the purity of the picricacid.

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I am unable to solve these questions due to a time constraint. I received help on part 1 of the unknown acid. Need help with the rest! I really appreciate it!

Lab: ACID-BASE TITRATION

Objective: To use titration to determine in Part 1, the concentration (molarity) of an unknown acid solution and in Part 2, the purity of a sample of KHP acid. Background: Titration is a volumetric technique used to determine the concentrations of solutions, molar masses of solids, and purity of samples. A titration inolves the addition of a titrant (a solution of known concentration) to an analyte ( a solution of unknown concentration), or vice versa, using a piece of glassware called a burette. The titration is carried out until it reaches an equivalence point (the exact point at which the reaction between the two solutions is complete). A chemical indicator is often used to aid in the identification of the equivalence point. An indicator changes color upon reaching the equivalence point of endpoint. Since we will be carrying out an acid-base titration, the indicator must change color upon reaching the endpoint at a specific pH. An example of such an indicator is phenolphthalein, which changes from colorless to pink near the equivalence point when pH approaches 7. It should be noted that the indicator selected is dependent upon a given titration, so that the observed color change is close to the ideal equivalence point. Once the titration is completed, we can use the volumes measured for each solution, as well as the concentration of the titrant, to determine the concentration, molar mass, or purity of the unknown analyte. Procedure Note: During lab, record all measurements and data in a clearly labeled table. Part 1: 1. Obtain ~80mL of ~0.1M of base (sodium hydroxide, NaOH) into a labeled beaker. Make sure to report the actual concentration from the bottle. 2. Obtain ~25mL of unknown acid into a second labeled beaker. Record whether your acid is monoprotic, diprotic, or triprotic. 3. Using a funnel, or carefully pouring, fill the burette with sodium hydroxide NaOH. Make sure a labeled waste beaker is under the buretter to catch any drippings. Rinse your burette by filling it with ~5-10mL of NaOH. Turn the knob to let ~2 mL out the tip into the waste beaker. Make sure your waste beaker is under the burette to catch any drips. Next, turn the Burette upside down and pour the remaining NaOH into your waste beaker while twisting the burette to coat the inside with NaOH. 4. Fill your burette with sodium hydroxide. Turn the knob to fill the tip of the burette with NaOH. Record the initial volume to the correct number of significant figures. 5. Using a volumentric pipette, transfer 10.00 mL of your unknown acid into the Erlenmeyer flask. 6. Add 3 drops of pH indicator and magnetic stirrer to the Erlenmeyer. 7. Under the burette, place the Erlenmeyer on a hot plate and turn on the stir knob. Heat should be off. 8. Titrate the unknown acid until the solution turns pink and remains pink. 9. Record the final volume on the burette to the correct number of significant figures. 10. Repeat (4) to (7). Pre-plan the whole lab and who will do which task. MY RESULTS: PART 1 The actual concentration of NAOH is 0.5939, 80 mL Initial Final 24.9 mL 44.5mL 26.1 mL 45.57mL PART 2 Impure KHP (SAMPLE B ) Only Paper: 0.280 g Impure KHP Mass: 0.526 g Initial volume: 22.30 mL + water +26.10 mL Final Volume: 26.80 mL QUESTIONS: 1. Calculate the concentration (molarity) of your unknown acid. 2. Calculate how many grams in your impure KHP sample was indeed KHP. 3. Calculate the percent (by mass) KHP in the sample. Part 2: 1. What would happen to the final result( state what the final result is ) if the tip of the burette was not filled when reading the initial volume? 2. What would happen to the final answer if you used Ca(OH)2 instead of NaOH and DID NOT know it? 3. What would happen to the final result if the Erlenmeyer flask was contaminated with other types of acid? 4. Recalculate the concentration of your unknown acid if it was given as a diprotic acid. 5. Draw a molecular level picture (showing Na+, H+, OH-, H2O, etc...) of what is in your beaker before, at half point, at equivalent point, and at end point of the titration. Part 3: 1. What would have happened to the final result if the impurity in the sample was also acidic? 2. What would have happend to your final result if you used ~25mL of water to dissolve the solid acid instead of 20mL? Summary data: Create a clearly labeled table of all your final data results for each run, and athe average of multiple runs. Part 1: Unknown #, Vol base, M base, Vol acid, M acid. Part 2: Sample #, % KHP (purity)

24.

Kagel and Farwell report a reversed-phase HPLC method for determining the concentration of acetylsalicylic acid (ASA) and caffeine (CAF)

in analgesic tablets using salicylic acid (SA) as an internal standard.

series of standards was prepared by adding known amounts of acetyl

salicylic acid and caffeine to 250-mL Erlenmeyer flasks and adding 100

mL of methanol. A 10.00-mL aliquot of a standard solution of salicylic

acid was then added to each. The following results are obtained for a

typical set of standard solutions.

milligrams of

peak height ratios for

standard

mg ratio of peaks

ASA CAF ASA/SA CAF/SA

1 200.0 20.0 20.5 10.6

2 250.0 40.0 25.1 23.0

3 300.0 60.0 30.9 36.8

A sample of an analgesic tablet was placed in a 250-mL Erlenmeyer

flask and dissolved in 100 mL of methanol. After adding a 10.00-mL

portion of the internal standard, the solution was filtered. Analysis of

the sample gave a peak height ratio of 23.2 for ASA and of 17.9 for

CAF.

(a)

Determine the milligrams of ASA and CAF in the tablet.

(b)

Why was it necessary to filter the sample?

(c)

The directions indicate that approximately 100 mL of methanol is

used to dissolve the standards and samples. Why is it not necessary

to measure this volume more precisely?

(d)

In the presence of moisture, ASA decomposes to SA and acetic acid.

What complication might this present for this analysis? How might

you evaluate whether this is a problem?

Equipment Materials

50.00 ml buret, 250.ml volumetric flask, 1.00 ml volumetric pipet 10.0 mL volumetric pipet, 25.00 mL volumetric pipet, pipettor, 250 mL Erlenmeyer flask, 10.0 mL graduated cylinder, several beakers, several watchglasses, unkown bleach solution, 10% potassium iodide solution, 2M hydrochloric acid solution, 0.150 M sodium thiosulfate solution, 1% starch solution.

Background information regarding this experiment is as follows:

An aqueous solution of sodium hypochlorite (NaOCI) is a clear, slightly yellow liquid, and is commonly known as bleach. Aside from its uses as a bleaching agent, sodium hypochlorite solutions are also used as sterlizing agents and in water treatment. Industrial uses include agricultural, food handling, paper production, and textile production. Sodium hypochlorite is also added to waste water to reduce odors.

The concentration of sodium hypochlorite in bleach solutions can be determined by titration. Therefore, we must use a two-step method to titrate sodium hypochlorite. In the first step, sodium hypochlortie, hydrochloric acid, iodide ion, and starch are combined to form a starch-triiodide complex. In this step there are 3 reactions that take place:

The result of these 3 reactions is that when sodium hypochlorite is present the starch-triiodide complex is produced.

In the second step the starch-triiodide product is titrated by sodium thiosulfate to form a colorless solution of iodide, dithionate.

I do not know how to answer the following on the data sheet: (? means I do not know how to answer)

Table 1:

Mass of 1.00 ml of unknown (by tare): 10.552 g

Density of bleach: (g/ml) : 1.055 g/ml

Volume of diluted bleach used in each titration: 10.00

Concentration of sodium thiosulfate: ? (M)

Titration Data:

Initial buret reading (ml): trial1: 0.000 ml Trial 2: 0.00 ml Trial 3: 0.000ml

Final buret reading (ml): trial 1: 13.50ml Trial 2: 13.30ml Trial 3: 13.40ml

Volume delivered(ml: trial 1: ? Trial 2: ? trial 3:?

Calculations: (show clearly)

For trial 1,2,3 ???

1.Moles of Na₂S₂O₃ in titration: ??? trial 1,2,3

2.Moles of NaOCl: ???trial 1, 2,3

3.Average moles of diluted NaOCl: (one answer) ?

4.Average moles of undiluted NaOCl: (one answer.. mol) ?

5.Standard Deviation: (one answer) (mol)

6.Average mass of NaOCl in 25.00mL of undiluted bleach: ? (g)

7.mass of 25.00 mL of undiluted bleach: ? (g)

8. Mass % of NaOCL from experiment: ? (%)

9.Actual mass % of NaOCL (from prof; student does not make an entry):? (%)

More questions all need to be answered please:

1. Calculate the volume of the reagent thiosulafte solution which would be required if you titrated a 25ml sample of the original commercial bleach instead of 25ml of the diluted solution.

2. Why would we choose a graduated cylinder to measure the HCl and KI solutions instead of a volumetric pipette? When would we use a volumetric pippette?

3. You calculated the mass percent of sodium hypochlorite in the bleach( dilute and undiluted). Now calculate the molarity of the sodium hypochlorite in the diulte and undiluted solution/

4. Use oxidation states to show that equations (2) and (4) in the introduction are redox reactions and that equation 1 is not a redox reaction. :

-(1) NaOCl(aq) + HCl(aq) yields HOCl(aq)+NaCl(aq)

-(2) HOCL (aq) + HCl(aq) + 3I(aq) yields I3(aq) + 2CL(aq) + H₂O(l)

-(4) (I₃) (Starch) + 2S₂O₃ yields 3I + S₄O₆ + starch

5. combine equation 1-3 to show the overall reaction for the products of the starch-triiodide complex.

-(1) NaOCl(aq) + HCl(aq) yields HOCl(aq)+NaCl(aq)

-(2) HOCL (aq) + HCl(aq) + 3I(aq) yields I3(aq) + 2CL(aq) + H₂O(l)

-(3) I₃ + Starch yields (I₃) (Starch)

6. What standard was used in this experiment?

7. Describe the naming of NaClO and HClO. What are the names of the following compounds? NaClO₂ , NaClO_3, NaClO₄ , HClO₂, HClO₃, HClO₄

8 Submit a one paged (typed double spaced) report that discusses the following topics:

a. the goals or objectives of the experiment

b.your results as they relate to the goals and objectives

c. the quality of your results.

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