Experiment_613_Spectrophotometric determination of aspirin_1_2

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Work date:

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Experience number and title

Experiment613:Spectrophotometric determination of aspirin

Experiment613 Spectrophotometric determination of aspirin

Section1: Purpose and summary

It is said that most aspirin tablets contain 5 grains of the active ingredient acetylsalicylic acid. (A grain is a unit of mass and equals 0.0647989 grams.) How reliable is this value? This experience allows you to verify the manufacturer.

The analysis method we use is called "spectrometry" or "spectrophotometry". It depends on molecules being able to absorb electromagnetic radiation of certain wavelengths using the energy of the radiation to 'excite' the electrons in their atoms. (Therefore, they have absorption spectra, just like isolated gas atoms.) The higher the concentration of a given molecule in a sample, the more light of a given wavelength the sample can absorb. The absorption of light by

The sample increases in direct proportion to the concentration of the molecule present. We use electromagnetic radiation ("light") of wavelength (λ) 297 nm, which falls in the ultraviolet range. In this experiment, a 297 nm light is shined through the sample and the amount of light absorbed by the sample is measured.

The chemical name of the active ingredient in aspirin is acetylsalicylic acid. To completely dissolve aspirin in water, you must chemically convert it to the sodium salicylate salt before measuring absorbance. Note that 1.00 mol of sodium salicylate is produced for every 1.00 mol

Uses of acetylsalicylic acid (aspirin):

Sodium salicylate is then reacted with acidic Fe³⁺to form iron(III) salicylate complex, [FeSal]⁺. This complex shows maximum absorption at a wavelength of 525 nm and is purple in color. Absorption of [FeSal]⁺The complex is directly proportional to the concentration of salicylate in the sample but does not tell us the actual concentration in any sample. To convert an instrument reading to an actual salicylate concentration, we must first calibrate the instrument with solutions of known sodium salicylate concentration.

In general, this pilot project consists of:

By carefully preparing 5 solutions ofacquaintanceconcentrations
Measure the absorbance of eachacquaintancemiA strangerSolution

Carefully prepare an aspirin solution (fromA strangerConcentration)
(Depending on the laboratory) Creation of a calibration curve from the data of the 5 solutions ofacquaintanceconcentrations
(After the lab) Using the calibration curve to calculate the amount of aspirin in a tablet or pill.

Section 2: Safety Precautions and Disposal

Safety precautions:

All the materials in this experiment are relatively harmless. Laboratory-derived aspirin should not be consumed. Caustic soda is corrosive. The use of eye protection is recommended for all experimental procedures.

garbage deposit:

Discard solutions down sink drain with plenty of water. Dispose of any filter paper and undissolved residue from the aspirin tablet in the trash.

Section 3: Procedures

Part 1:Prepare solutions of known concentrations.

Ask your instructor to describe and demonstrate the use of volumetric pipets, volumetric flasks, and burettes. In your smallest beaker, obtain 10-15 mL of the standard solution of approximately 0.03 M sodium salicylate. Record the actual concentration (including significant numbers).	Concentration (molarity) of sodium salicylate standard solution:
Using a rubber bulb and a 5.00-mL graduated pipet, carefully measure 5.00 mL of the standard solution and place it in a 100.0-mL volumetric flask. at least 2 minutes. (Hold the glass or rubber stopper firmly with your thumb. Be careful not to spill any liquid before the solution is completely mixed.)fallback solution.
Get 6 clean, dry 100ml beakers. Label these beakers 'Stock', '1', '2', '3', '4' and '5'.
Empty the 100.0 mL volumetric flask into the beaker labeled "Stock." Rinse the volumetric flask with plenty of laboratory water.
Fill a 50 mL buret with your stock solution. Fill to near the 0.00 mL mark, making sure you can accurately read the volume at the meniscus (the lowest part of the solution surface in the buret). It is okay if the volume is not exactly 0.00 ml.	Initial burette value (ml): (Include 2 decimal places!)
Now prepare the first known solution,solution not. Place the 100 mL volumetric flask under the burette, then turn the burette stopcock to the appropriate position5 mlflows into the volumetric flask. Then read the buret.	Burette #1 reading (ml): (Include 2 decimal places!)
Quick check: Subtract theFirst burette readingVonread burette#1. This number should be close to 5 ml. You will then use this volume to calculate the concentration (molarity) of sodium salicylate in solution #1. Let's call this volume "Transfer Volume #1"	Calculate the difference between the initial value of the burette and the value of burette #1. (Include 2 decimal places!) “Transfer Belt #1”: ______________ml
Up to 5 mL contained in the volumetric flask, carefully fill the flask with 0.010 M FeCl.₃make up to the 100 mL mark in 0.1 M HCl solution and mix well. That's allsolution not. Empty the 100 mL volumetric flask into the beaker marked "1". Rinse the volumetric flask with plenty of laboratory water.
Now prepare the second known solution,Solution #2. Place the 100 mL volumetric flask under the burette, then turn the burette stopcock to the appropriate position10mlflows into the volumetric flask. Then read the buret.	Buret #2 reading (mL): (Include 2 decimal places!)
Subtractread burette#1Vonread burette#2. This number should be close to 10 ml. RecordTransfer volume #2. Carefully fill the volumetric flask with 0.010 M FeCl.₃add exactly to the 100 mL mark in 0.1 M HCl solution and mix well. That's allSolution#2. Empty the 100 mL volumetric flask into the beaker marked "2". Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between the #1 and #2 buret readings. (Include 2 decimal places!) “Drive Belt #2”: ______________ml
Now prepare the third known solution,Solution #3. Place the 100 mL volumetric flask under the burette, then turn the burette stopcock to the appropriate position15mlflows into the volumetric flask. Then read the buret.	Burette #3 Reading (mL): (Include 2 decimal places!)
Subtractread burette#2Vonread burette#3. This number should be close to 15 ml. RecordTransfer Volume #3. Carefully fill the volumetric flask with 0.010 M FeCl.₃add exactly to the 100 mL mark in 0.1 M HCl solution and mix well. That's allSolution#3. Empty the 100 mL volumetric flask into the beaker marked "3". Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between the #2 and #3 buret readings. (Include 2 decimal places!) „Transfer belt no. 3": ______________ml
Refill the burette with your stock solution (use the beaker marked "Stock"). Read the buret and record this volume (it should be close to the 0.00 mL mark). It is okay if the volume is not exactly 0.00 ml.	Initial buret value after RELOAD (mL): (Include 2 decimal places!)
Now prepare the fourth known solution,Solution #4. Place the 100 mL volumetric flask under the burette, then turn the burette stopcock to the appropriate position20mlflows into the volumetric flask. Then read the buret.	Reading burette #4 (ml): (Include 2 decimal places!)
subtract theFirst burette readingAfter LOADINGVonread burette#4. This number should be close to 20 ml. RecordTransfer tape #4. Carefully fill the volumetric flask with 0.010 M FeCl.₃add exactly to the 100 mL mark in 0.1 M HCl solution and mix well. That's allSolution#4. Empty the 100 mL volumetric flask into the beaker marked “4”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between the #4 burette reading and the initial burette reading after FILL. (Include 2 decimal places!) „Transfer belt no. 4": ______________ml
Finally, prepare the fifth known solution,Solution #5. Place the 100 mL volumetric flask under the burette, then turn the burette stopcock to the appropriate position25mlflows into the volumetric flask. Then read the buret.	Burette reading #5 (mL): (Include 2 decimal places!)
Subtractread burette#5Vonread burette#4. This number should be close to 25 ml. RecordTransfer Volume #5. Carefully fill the volumetric flask with 0.010 M FeCl.₃add exactly to the 100 mL mark in 0.1 M HCl solution and mix well. That's allSolution#5. Empty the 100 mL volumetric flask into the beaker marked “5”. Rinse the volumetric flask with plenty of laboratory water.	Calculate the difference between the #4 and #5 buret readings. (Include 2 decimal places!) "Transfer Tape #5": ______________ml
You should now have 5 beakers, labeled 1-5, each containing 100 mL of sodium salicylate at various concentrations.

paper2:Using an aspirin tablet to prepare a solution of unknown concentration

Crush an aspirin tablet in a clean porcelain mortar. Read the label on the aspirin bottle to find out how many grains are in each pill. If grains are not reported, record the number of grams.	On the aspirin bottle, indicate the number of GRAINS (if you do not have them, indicate grams) that each tablet should contain:
Transfer all of the solid to a beaker (any size larger than 100 mL). Use several portions of 0.1 M NaOH (about 50 mL total) to rinse out all the solid in the beaker. Be sure to rinse the mortar and pestle.
Gravity Filtration: Prepare a folded filter paper (coffee filter paper works!) in a funnel and insert the funnel into the neck of a clean 250.0 mL volumetric flask. (The vial should be clean, but it doesn't have to be dry.) Be careful not to get dust, paper, or other contaminants inside the vial. Pour the aspirin-containing liquid from the beaker onto the filter paper in the funnel. The clear solution flowing through the filter contains your sample; finally the solid that remains on the filter paper is discarded. Rinse the beaker three times with 10 mL of 0.1 M NaOH (about 30 mL total) and also pour the rinsings into the filter funnel.
After filtering the sample, remove the funnel and dispose of the filter paper and any undissolved residue from the aspirin tablet with your household waste.
Carefully add enough laboratory grade water to bring the solution level to just above the 250.0 mL line on the neck of the bottle.
Mix the sample well by gently inverting the capped vial for at least 2 minutes. (Hold the glass or rubber stopper firmly with your thumb. Be careful not to spill any liquid before the solution is completely mixed.)
The solution is too concentrated to use at this point. Prepare a dilution of this by transferring exactly 2.00 mL to a clean 100.0 mL volumetric flask using a volumetric pipet (flask can be wet while clean). Carefully add 0.010 M FeCl₃in a 0.1 MHCl solution to bring the solution level to the 100.0 mL mark on the vial. Homogenize.
Label this 100 mL volumetric flask "unknown."

paper3:Measurement of ultraviolet light absorption by sodium salicylate solutions

1. See Appendix: Technique I Use of Spec 20/Spec 200 for proper use and operation of the spectrophotometer.

2. Set up the spectrophotometer to measure absorbance at 525 nm.

3. Use 0.010 M FeCl₃in 0.1 M HCl solution as the reference liquid, ie, set the instrument to zero absorbance when the light path passes through this ferric iron(III) solution.

Measure the absorbance of each of your solutions with known concentrations (Beckers 1-5) and of your unknown solution (the 100 mL volumetric flask labeled "Unknown").

Rinse the cuvette (sample holder) well with laboratory grade water and then with each solution between measurements.

Enter your data in the following table.

Investigation	absorption in525New Mexico
Solution #1
Solution #2
Solution #3
Solution #4
Solution #5
A stranger

Quick check:

The absorbance should increase from solution #1 (lowest) to solution #5 (highest).
The absorbance of the unknown should be between the absorbance of solution #1 and solution #5.

Section 4:calculations

(ºto becan be completed after the internship. However, if time is available, it is recommended that you work on these calculations before you leave the lab, because if an error is found in your data, you may have time to get better data in the lab.)

Calculate the concentration of your known solutions using the dilution equation: M₁v₁= METRO₂v₂.

solution name	METRO₁= concentration of the solution to be transferred (the initial solution)	v₁= Volume transferred	METRO₂= final concentration (Calculate this numberwith correct SIGNIFICANT NUMBERS)	v₂= final volume of solution after dilutionEsfinished
fallback solution	Concentration (molarity) of sodium salicylate standard solution (on the standard vial label)=	5,00ml		100,0ml
solution not	Concentration of your stock solution =	Transfer volume Nr. 1 =		100,0ml
solution not	Concentration of your stock solution =	Transfer volume Nr. 2 =		100,0ml
solution not	Concentration of your stock solution =	Transfer volume Nr. 3 =		100,0ml
solution not	Concentration of your stock solution =	Transmission Band #4 =		100,0ml
solution not	Concentration of your stock solution =	Transmission Band #5 =		100,0ml

2. Create the calibration curve

Using Microsoft Excel or another spreadsheet program, plot the final concentration (y values) against the absorbance (x values) for solutions 1-5. Use Microsoft Excel to find the line of best fit (also known as linear regression).

Basic instructions for creating charts with Microsoft Excel^®

Enter the data into the worksheet with the x-coordinate data in the first column and the y-coordinate data in the second column.

Select data columns with the mouse.

Choose Chart Wizard on the toolbar or from the Insert menu.

Select the type of chart you want. Typically xy scatterplot.

choosenextGusto.

If necessary, select the data range (normally not required).

choosenextGusto.

Enter chart titles and data labels.

choosenextGusto.

Choose Insert Chart Asnew page.

Select data points with the mouse.

Add trend line from chart menu or right click.

Choose the type of trend line (linear).

Use the Options tab to include linear equations and r²not graphic value.

Make the chart pretty using the format plot area.

Double click on the axes to adjust the limits.

Write the equation of the line obtained from linear regression(step 14 above):

______________________________________________________________________

Ask your instructor for a graph of your datahe mustto be thereRlaboratory report.

Using the equation written above, calculate the concentration of the unknown solution:

Investigation	absorption in525New Mexico	concentration
A stranger

4. From the concentration of the unknown solution, calculate the amount of aspirin in the tablet.

Copy the concentration of the unknown solution (from the table above) that was placed in the spectrophotometer.	Concentration of unknown solution:
In the laboratory, you prepared a dilution of the unknown solution by using a volumetric pipet to transfer exactly 2.00 mL of it to a clean 100.0-mL volumetric flask. Let's calculate the concentration of this 2.00 mL solution using M₁v₁= METRO₂v₂. METRO₁= (calculate) v₁= 2,00ml METRO₂= concentration of the unknown solution (from the top) v₂= 100,0ml	Concentration of unknown solution before dilution:
This 2.00 mL solution comes from the 250.0 mL volumetric flask (this volume corresponds to 0.2500 L). Calculate the number of moles of sodium salicylate in the 250.0 mL volumetric flask by multiplying 0.2500 L by the concentration of the unknown solution before dilution.	Mole sodium salicylate:
How many moles of aspirin (acetylsalicylic acid) were in the tablet for each mole of sodium salicylate in the original solution? (See the balanced equation at the beginning of this experiment)	Mol Aspirin:
Calculate the number of grams from the number of moles using the molar mass of aspirin (180.158 g/mol).	Aspirina Gramm:
Convert grams to grains (15.43 grains = 1,000 grams).	Aspirin grains:
Calculate the percent error using grains or grams of aspirin: (Your aspirin grains - aspirin grains from the bottle) * 100 Cereal aspirin bottle	Error rate:

PAGost-Lab Questions:

Discuss your test results: Do your results agree with the aspirin manufacturer's claim regarding grains (or grams) of aspirin in each pill? Discuss any of your experimental errors that might affect your conclusions.

Discuss how the structure of aspirin becomes more soluble in water after reacting with sodium hydroxide.

In another experiment, known ibuprofen solutions were used to create a calibration curve. The best line of fit for this calibration curve was

y=0,000924x+ 0,0000343

where x = absorbance and y = concentration (molarity) of ibuprofen in solution.

If absorbance = 0.207, what is the concentration of ibuprofen in the solution?

Predict absorbance for 5.62 x 10^-4ibuprofen solution M.

Assume that the anibuprofen tablet is crushed, dissolved in sodium hydroxide, and filtered. This solution is diluted to a total volume of 5.00 L and mixed well. If this solution has an absorbance of 0.163, how many milligrams of ibuprofen were in the tablet? The molar mass of ibuprofen is 206.29 g/mol.

Experiment_613_Spectrophotometric determination of aspirin_1_2_2 (2023)