LAB Manual CHM256 - DETERMINATION THE PERCENT (W/W) OF THE ACTIVE INGREDIENT IN ASPIRIN TABLET PDF

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EXPERIMENT 3PREPARATION AND STANDARDIZATION OF HCl SOLUTION WITHPRIMARY STANDARD Na 2 CO 3 SOLUTIONPurposeTo learn the technique to prepare an acid solution from a concentrated HCl and to determine the accurate concentration of the HCl solution by standardization with standard solution.IntroductionH...

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ANALYTICAL CHEMISTRY

EXPERIMENT 3 PREPARATION AND STANDARDIZATION OF HCl SOLUTION WITH PRIMARY STANDARD Na2CO3 SOLUTION Purpose To learn the technique to prepare an acid solution from a concentrated HCl and to determine the accurate concentration of the HCl solution by standardization with standard solution.

Introduction Hydrochloric acid is a solution of hydrogen chloride (HCl) in water. It is a highly corrosive, strong mineral acid with many industrial uses. Commercially available con HCl is sold as 37-38 % w/w HCl. Since HCl is not a primary standard reagent, the solution of HCl prepared in the lab need to be standardized with a primary standard solution in order to determine its exact concentration. Apparatus 250 mL volumetric flask Retort stand White tile 250 mL conical flask (x3) 25 mL pipette burette

Chemical Reagents Na2CO3 solution (Exp. 2) Concentrated HCl Methyl orange – indicator

Procedure A. Preparation of HCl solution 1.

Read the label on the bottle of the concentrated HCl (should be located in the fume hood). From the assay (% w/w, specific gravity and molecular weight) of the concentrated HCl acid given, calculate the molarity of the concentrated HCl and the volume needed to prepare 250 mL 0.1 M of dilute HCl solution.

2.

Measure the calculated volume of the concentrated HCl (1) using a 10-mL measuring cylinder and transfer into a 250 mL volumetric flask (which has already contained some distilled water – to dilute conc. acid, add acid to water).

3.

Rinse the measuring cylinder and transfer it into the volumetric flask. Then, dilute the acid into the 250 mL of distilled water. Use a stopper to seal the volumetric flask and mix well by turning the flask upside-down a few times in order for the solution to be homogenous.

1

B. Standardization of HCl solution 1. 2. 3.

4. 5. 6.

Fill the burette with the prepared diluted HCl and record the initial burette reading. Pipette 25.0 mL of the standard Na2CO3 that you have prepared in Exp. 2 into 250 mL conical flask. Add 2–3 drops of methyl orange as indicator. Titrate the Na2CO3 in the conical flask with the HCl from the burette slowly while shaking the flask until the indicator color changes from yellow to red. Record your reading at the the end point of the titration. Repeat 2–3 times and record all your readings in a table. Calculate the exact molarity of your HCl solution. Transfer the remaining acid solution into a clean reagent bottle KEEP THIS SOLUTION FOR THE NEXT EXPERIMENT (i.e will be used to standardize NaOH solution)

ANALYTICAL CHEMISTRY

REPORT SHEET

Experiment No

3

Title

PREPARATION AND STANDARDIZATION OF HCl SOLUTION WITH PRIMARY STANDARD Na2CO3 SOLUTION Name Student ID Course Group Date of Experiment Date of Submission Lecturer’s Name

NUR AINA SAKINAH BT NORHISHAM 2019437562 AS120 4L

EXPERIMENT 3 PREPARATION AND STANDARDIZATION OF HCl SOLUTION WITH PRIMARY STANDARD Na₂CO₃ SOLUTION 1.

State the objective of the experiment. a. To learn the technique to prepare an acid solution from a solution from a concerntrated HCl. b. To determine the accurate concentration of HCl solution by standardization with standard solution.

2.

Describe the procedure (in brief).

A.

Preparation of HCl solution

B.

1.

Read the label on the bottle of the concentrated HCl (should be located in the fume hood). From the assay (% w/w, specific gravity and molecular weight) of the concentrated HCl acid given, calculate the molarity of the concentrated HCl and the volume needed to prepare 250 mL 0.1 M of dilute HCl solution.

2.

Measure the calculated volume of the concentrated HCl (1) using a 10-mL measuring cylinder and transfer into a 250 mL volumetric flask (which has already contained some distilled water – to dilute conc. acid, add acid to water).

3.

Rinse the measuring cylinder and transfer it into the volumetric flask. Then, dilute the acid into the 250 mL of distilled water. Use a stopper to seal the volumetric flask and mix well by turning the flask upside-down a few times in order for the solution to be homogenous.

Standardization of HCl solution 1.

Fill the burette with the prepared diluted HCl and record the initial burette reading.

2.

Pipette 25.0 mL of the standard Na2CO3 that you have prepared in Exp. 2 into 250 mL conical flask. Add 2–3 drops of methyl orange as indicator.

3.

Titrate the Na2CO3 in the conical flask with the HCl from the burette slowly while shaking the flask until the indicator color changes from yellow to red. Record your reading at the the end point of the titration.

4.

Repeat 2–3 times and record all your readings in a table.

5.

Calculate the exact molarity of your HCl solution.

6.

Transfer the remaining acid solution into a clean reagent bottle

3.

Results / Data Obtain the following information from the label on the bottle of the concentrated HCl. Table 1 Molecular Weight of HCl Density @ specific gravity Percent concentration (w/w)%

36.46 g/mol 1.18 37

Table 2: Titration of Na2CO3 with HCl Titration Final burette reading Initial burette reading Volume of HCl used

Rough

1

2

3

40.00 15.00 25.00

40.00 15.00 25.00

40.00 14.00 26.00

Average volume of HCl: 25.33mL Molarity of Na2CO3 (from Experiment 2): 0.05M

4.

Questions Based on the information in Table 2, calculate the molarity of the conc. HCl solution.

a) What volume of the concentrated HCl is needed to prepare 250 mL 0.1 M HCl solution?

b) Write a balanced equation for the reaction between HCl and Na2CO3.

c) Using the data above and the stoichiometric ratio from the equation, calculate the exact molarity of the HCl solution.

5.

Conclusion

EXPERIMENT 4 ACID BASED TITRATION I DETERMINATION OF PERCENT CONTENT OF ACETIC ACID IN VINEGAR

Purpose To determine the content percent of acetic acid in vinegar.

Introduction Vinegar is a diluted solution of acetic acid. Commercial vinegar should contain about 4-6 % of acetic acid (CH3COOH). The acetic acid content in vinegar can therefore be determined through titration method using a standard solution of a base for example the standard solution of NaOH. In this experiment, the acetic acid content in different brands of commercial vinegar is to be determined by titrating them with NaOH solution that has previously been standardized using a secondary standard solution of HCl (prepared in Experiment 2). Apparatus 250 mL volumetric flask Retort stand White tile 250 mL conical flask (x3) 20 mL pipette Burette

Chemical Reagents 3 brands of commercial Vinegar NaOH pellet Phenolphthalein – indicator

Procedure A. Preparation of NaOH solution and standardization of NaOH 1. 2. 3. 4. 5.

6. 7.

Calculate the mass needed to prepare 250 mL 0.1 M NaOH solution. Prepare the NaOH solution using the same procedure as in Experiment 2 (preparation of Na2CO3 solution). Fill the burette with NaOH solution. Pipette 20.0 mL of the HCl (from Exp. 3) into 250 mL conical flask. Add 2 – 3 drops of phenolphthalein as indicator. Titrate the HCl in the conical flask with the NaOH from the burette slowly while shaking the flask until the indicator color change. Record the color change and your reading at the end point of the titration. Repeat 2 - 3 times and record all your readings in a table. Calculate the exact molarity of the NaOH solution

B. Determination of acetic acid in vinegar 1. 2. 3. 4. 5. 6.

Pipette 10 mL of the given vinegar (record the brand name of the chosen vinegar) and dilute to mark with distilled water in a 100 mL volumetric flask. Then pipette 20 mL of this diluted vinegar solution into a 250 mL conical flask. Add 2 – 3 drops of phenolphthalein as an indicator. Titrate the vinegar in the conical flask with the NaOH from the burette slowly while shaking the flask until the indicator changes color. Repeat steps 2 and 3 two more times. Record your reading at the end point of the titration. Calculate the w/v % of acetic acid in the vinegar sample.

ANALYTICAL CHEMISTRY

REPORT SHEET

Experiment No Title

Name Student ID Course Group Date of Experiment Date of Submission Lecturer’s Name

EXPERIMENT 4 ACID BASED TITRATION I 1.

State the objective of the experiment.

2.

Describe the procedure to determine acetic acid in vinegar (in brief).

.

3.

Results / Data A. Standardization of NaOH solution Table 1: Titration of HCl with NaOH Titration Final burette reading Initial burette reading Volume of NaOH used

Rough

1

Average volume of NaOH: ......................... Exact molarity of HCl solution from Experiment 3: ………….. B. Determination of acetic acid in vinegar Brand of vinegar used:…………………………………………

2

3

Table 3.2: Titration of vinegar with NaOH Titration Final burette reading Initial burette reading Volume of NaOH used

1

2

3

4

Average volume of NaOH: .........................

4.

Questions a) Write a balanced equation for the reaction between NaOH and HCl.

b) Using the data above and the stoichiometric ratio from the equation, calculate the exact molarity of the NaOH solution.

c)

Write the chemical equation for the reaction between acetic acid (CH3COOH) and NaOH solution.

d) Determine the w/v % of acetic acid in the vinegar sample. i)

Using (MaVa / MbVb) = a/b, determine the molarity of the diluted acetic acid.

ii)

In the procedure 10 mL of the original vinegar has been diluted to 100 mL Calculate the molarity of the concentrated (original) vinegar. (M1V1=M2V2)

iii)

Calculate the mass of acetic acid (MW 60 g/mol) and then calculate % w/v of acetic acid in the vinegar sample. % w/v = mass of acetic acid x 100 volume of sample

e)

f)

Explain why you need to standardize NaOH solution after preparation.

In this experiment HCl solution from Experiment 3 is used as a secondary standard solution to standardize NaOH solution. Define secondary standard solution.

g)

5.

Draw a titration curve for the titration between acetic acid (CH3COOH) and NaOH solution.

Conclusion

EXPERIMENT 5 ACID – BASED TITRATION II DETERMINATION THE PERCENT (W/W) OF THE ACTIVE INGREDIENT IN ASPIRIN TABLET

Purpose To determine the percent (w/w) of the active ingredient, acetlysalicyclic acid (C9H8O4) in aspirin tablet. Introduction Acetylsalicylic acid (ASA) is an active ingredient in an aspirin tablet. Aspirin can be hydrolyzed by the excess amount of sodium hydroxide to give sodium salts of two weak acids, ethanoic acid and salicylic acid (equation 1). The unreacted NaOH can then be neutralized by HCl (equation 2). CH3COO-C6H4-COOH + 2NaOH  CH3COO- Na+ + HO-C6H4-COO- Na+ (excess) NaOH + (unreacted)

HCl 

(1)

NaCl + H2O (2)

In this experiment the principle of back titration is being used to determine the excess mole of NaOH and hence the mole of NaOH is consumed by acetylsalicylic acid in the reaction. Calculations Initial mole of NaOH used in hydrolysis = a Mole of unreacted NaOH (from Eqn. 2) = mole of HCl used in titration = b Mole of NaOH that reacted with acethylsalicyclic acid = a-b = c Mole of acethylsalicyclic acid (from Eqn. 1) = ½ c Mass of acethylsalicyclic acid = ½ c x MW % w/w = mass of acethylsalicyclic acid x 100 mass of aspirin Apparatus Volumetric flasks: 250 mL(1) 100 mL (2) Retort stand 250 mL conical flask (x3) 20 mL , 25 mL pipette Burette Weighing boat/bottle/paper Filter tunnel Dropper

Chemical Reagents 1 M NaOH – 50 mL or 100 mL (shared) Indicator : Phenol red or phenolphthalein Aspirin tablet (Brand:Disprin or Visprin, 300 mg ASA) 0.1M HCl (250 mL)

Procedure A. Preparation of solution Prepare 1 M NaOH solution (50 mL), 0.1 M HCl solution (250 mL) and 0.05 M Na2CO3 (100 mL) as you have learned in previous experiments. B. Hydrolysis of aspirin 1.

Collect one commercial aspirin tablet (300 mg ASA) from your lecturer. Record the brand name, the manufacturer’s name and % content (mass) of the active ingredient in the tablet (use 2 tablets if the ASA content is 100 mg). Weigh the aspirin tablet accurately in a weighing paper and transfer the tablets into a 250 mL conical flask. Add 25 mL (use pipette) 1.0 M NaOH solution and 25 mL of water to the tablet. Heat this mixture until it boils for about 10 minutes in order to hydrolyze the aspirin. Cool the solution at room temperature then transfer quantitatively into 250 mL volumetric flask. Dilute the solution to the mark carefully using dropper when the level is close to the calibration mark.

2. 3.

4.

C. Standardization of HCL solution Follow the procedure as in Experiment 3B to standardize your HCl solution. Record the data in Table 5.2 Equation: Na2CO3 + 2HCl  2NaCl + H2O + CO2

D. Back titration 1. 2. 3.

4. 5.

Pipette 20 mL of aspirin solution into a 250 mL conical flask . Add 1-2 drops of phenol red or phenolphthalein indicator Titrate the aspirin solution in the 250 mL conical flask with the diluted 0.1 M HCI solution from the burette slowly while shaking the flask until the indicator changes color (phenol red: light red to light yellow; phenolphthalein: pink to colorless). Record your reading at the end point of the titration. Repeat 2 – 3 times and record all your readings in a table. Calculate the % w/w of acetylsalicylic per tablet.

ANALYTICAL CHEMISTRY

REPORT SHEET

Experiment No Title

Name Student ID Course Group Date of Experiment Date of Submission Lecturer’s Name

EXPERIMENT 5 ACID – BASED TITRATION II 1.

State the objective of the experiment.

2.

Describe the procedure (use schematic diagram if necessary). Show all the calculations involved during the preparation of solutions. A. Hydrolysis of aspirin

B. Titration (back titration)

3.

Result/Data Table 5.1 Commercial/brand name of aspirin tablet Name of manufacturer Mass per tablet of active ingredient (acetylsalicylic acid) stated on the label Mass of aspirin tablet Calculated % w/w

Table 5.2 Standardization of HCl solution. Titration Final burette reading Initial burette reading Volume of HCl used

Rough

1

2

3

2

3

a) Average volume of HCl......................................mL b) Volume of Na2CO3 used:..............................mL c) Molarity of Na2CO3....................................M

5.3 Titration of aspirin solution with HCl Titration Final burette reading Initial burette reading Volume of HCl used

Rough

1

a) Average volume of HCl......................................mL b) Initial volume of NaOH used to hydrolyze aspirin........................mL

4. Calculation a) Based on the data obtained in Table 5.2 calculate the exact molarity of HCl.

b) Using back titration method, calculate the % (w/w) aspirin in the aspirin tablet. i) Calculate the initial mole of NaOH used for hydrolysis.

ii) Based on Equation (2), calculate the mole of excess (unreacted) NaOH.

iii) Calculate the mole of NaOH that has actually reacted with acetylsalicylic acid.

iv) Based on Equation (1), calculate the mole of acetylsalicylic acid in the solution.

v)

Calculate the mass then the % w/w of acetylsalicylic acid in the tablet.

vi)

5.

Conclusion

Compare the % w/w calculated in Table 5.1 and the result obtained in (e). Calculate the % error.

ANALYTICAL CHEMISTRY

EXPERIMENT 12 GRAVIMETRY II: GRAVIMETRIC ANALYSIS OF PHOSPHORUS IN PLANT FOOD Purpose To determine the content % of P in plant food through gravimetric analysis. Introduction Gravimetric analysis is a quantitative method that is based on determining the mass of pure compound to which the analyte is chemically related. In the method, the analyte is selectively converted to an insoluble precipitate; the precipitate is then dried or ignited and accurately weighed. Gravimetric analysis is considered as one of the most accurate method of chemical analysis. Usually, gravimetric analyses involve the following steps:      

Dissolving the sample Precipitating the analyte by adding precipitating reagent in excess Filtrating the separated precipitate from mother liquor Washing the precipitate (to avoid peptization) Drying or igniting the precipitate to constant weight Calculating the % of the desired analyte

Plant foods contain three essential nutrients; nitrogen, phosphorus and potassium. The labels on the plant food usually have a set of numbers such as 15-30-15 which indicate the content % of the nutrient (15% N, 30% P (expressed as P 2O5) and 15% K (expressed as K2O). The remainder is either anions or cations necessary to balance the charge in the chemical compounds. In this experiment, phosphorus will be determined through precipitation of the insoluble salt magnesium ammonium phosphate hexahydrate according to the reaction: 5H2O (l) + HPO 42- (aq) + NH 4 + (aq) + Mg2+ (aq) + OH- (aq)  MgNH4PO4.6H2O (s) The % P and % P2O5 in the initial sample can be calculated from the mass of MgNH4PO46H2O obtained using the following method:

73

Apparatus Weighing boat/bottle/paper 250 mL beaker Filter funnel Dropper Measuring cylinder 100 mL and 250 mL volumetric flask Retort stand, white tile, hotplate Porcelain evaporating dish

Chemical Reagent MgSO4.7H2O (10%) NH3(aq) 2 M (100 mL) 75% isopropyl alcohol

Procedure 1. Obtain your plant food sample and record the set of numbers written on the label. 2. Weigh the plant food sample between 3.0 - 3.5 g using a weighing paper and transfer the sample to a 250 mL beaker. 3. Add 35 - 40 mL of distilled water and stir the mixture with a glass rod to dissolve the sample. If your sample does not completely dissolve, remove the insoluble material through filtration. 4. To filtrate, add about 40 mL of 10% MgSO 4.7H2O. Then add about 100 mL 2 M NH 3 (aq) slowly while stirring. A white precipitate of MgNH4PO4.6H2O will form. 5. Allow the mixture to sit at room temperature for 10 minutes to complete the precipitation. 6. Weigh accurately a filter paper. Fold the paper and fit it into a glass funnel. Wet the paper with distilled water to hold it in place in the funnel. Transfer the precipitate carefully and all the solution from the beaker onto the filter paper. 7. Wash the precipitate by adding 2 - 3 times of 5 mL portions of distilled water to the beaker. Then rinse again the precipitate with two 10 mL portions of 75% isopropyl alcohol through the filter paper. 8. Remove the filter paper, place it on a numbered porcelain evaporating dish and store it to dry for 2 - 3 days on the bench or dry in the oven for 24 hrs at 60oC.

9. When the MgNH4PO4.6H2O is thoroughly dry, weigh the filter paper plus the MgNH4PO4.6H2O. Record the mass and calculate the % of phosphorus in your original sample.