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CONTENT

CONTENT i

PREFACE ii

SAFETY IN THE LABORATORY iii

FORMAT OF LABORATORY REPORT iv

Experiment 1: Potentiometric Titration 6

Experiment 2: Determination of water hardness by

complexometric titration

13

Experiment 3: Determination of limonene concentration by gas

chromatography

17

Experiment 4: Determination of active drug substance in a

pharmaceutical formulation using ultraviolet and

Visible Spectrometry (UV/VIS)

20

Experiment 5: Interpretation of spectra of an unknown sample

using fourier transform infra red (FT-IR)

spectroscopy

23

Analytical Chemistry (ERT 207) Laboratory Module

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PREFACE

The purpose of the laboratory work in this course is to introduce you to the field known as

analytical chemistry. Specifically, you will learn in detail the wet chemical methods known

as quantitative analysis. This will acquaint you with the basic techniques and operations

that are necessary to perform precise analytical measurements. These operations will form

a set of skills that are necessary to succeed in more advanced chemistry course like

instrumental analysis. In addition, these skills are essential to the students who will be

working in laboratories in industry and research instituition.


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SAFETY IN THE LABORATORY

ALWAYS:

Wear safety glasses

Wear protective clothing

Know the location and use of all safety equipment

Use proper techniques and procedures

Add acid to water

Be very cautious when testing for odors

Use hoods whenever poisonous or irritating fumes are evolved

Discard wastes properly - flush liquids down the sink with a large excess of water

Report any accident, however minor, to the instructor at once

All the times think about what you are doing

Be alert, serious, and responsible

NEVER:

Eat or drink in the lab

Perform unauthorized experiments

Leave anything unattended while it is being heated or is reacting rapidly

Aim the opening of a test tube or flask at yourself or at anyone else

Add water to acid

Insert droppers, pipettes and other laboratory equipment into reagent bottles – this

is a sure way of contaminating the contents

Return unused reagents to stock bottles

Clutter your work area

Take unnecessary risks

Enter chemical storage area


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FORMAT OF LABORATORY REPORT

The purpose of the laboratory report is to provide information on the measurement

procedure, obtained results, analysis and interpretation and discussion of the results. The

discussion and conclusions are definitely significant in a report because these sections

deliver the knowledge you gained upon doing the experiments.

For this particular laboratory, the following format is suggested:

1.Cover page

2.Introduction

3. Objectives

4.Summarized experimental procedure

5.Experimental data/Results

6. Discussion

7.Conclusion

Detailed descriptions of every item are given below:

1. Cover page

It should have the individual name, matrix number, course name and academic session,

the number and title of the experiment, group number, names of the team members and

the date of the report delivery.

2. Introduction

Brief introduction about the theory of the experiments should be mentioned.

3. Objectives

The objectives of the experiment are clearly stated.

4. Summarized experimental procedure

Attached the approved summarized experimental procedure to the lab report.

5. Experimental data/Results

This section deals with the management of data obtained after experiment. Data can be

presented in many forms such as tables, graphs, and some calculations or data

analysis. The best presentation of some data is graphical. Figures should be numbered.

Each figure must have a caption following the number. All graphs, beside captions, should

have clearly labeled axes.

6. Discussion

The discussion can be written in two ways;

a) Compare the expected outcome of the experiment with theory or

b) Make an appropriate graph on which the theory is represented and experimental data

by points. For overall, second method is the best way to present the discussion.


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A critical part of discussion is error analysis. In comparison of theory and experiment you

may not get a perfect agreement. It does not necessarily mean that your experiment was

failed. The results will be accepted, provided that you can account for discrepancy.

Precision and accuracy of the instrument or your ability to read the scales may be one

limitation. However, the reason for the difference between the expected and measured

values lies in the experimental procedures or in not taking into account all factors that

enter into the analysis. Apart from this, data analysis requires you to open your mind and

critical approach to your work and that routine methods may not be sufficient.

7. Conclusion

The conclusion should contain several shorts statements closing the report. They should

inform the reader if the experiments agreed with the theory. If there were differences

between measured and expected results, explain possible reasons for these differences.

You may also say what could have been done differently, how experiments may be

improved, or make other comments on the laboratory.


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EXPERIMENT 1

POTENTIOMETRIC TITRATION

1.0 OBJECTIVE

1. To understand the principle of pH change in acid-base titrations using a pH meter.

2. To determine the concentrations of the unknown acidic solutions as well as the acid-

ionization constant of a weak acid.

iii. To differentiate acid strength by the shape of a titration curve.

2.0 LEARNING OUTCOME

1. Ability to solve complex problem by using techniques, skills and modern engineering

tools to analyze the concentration of analytes of various classical titrimetric and gravimetric

methods for mass determination.

3.0 INTRODUCTION

Acid-base titrations have been performed in the past to determine the concentration of

an acidic or basic solution using a colored indicator. However, there are times when an

appropriate indicator does not exist, or where the color of the solution would obscure any color

change associated with the endpoint. In such cases, a pH meter can be used to monitor the

acidity of the solution throughout the titration. Recall the definition of pH:

Figure 1: pH meter

A pH meter consists of two electrodes: a glass electrode, which is sensitive to the

concentration of hydronium ions in solution, and a reference electrode. The reference

glass electrode combine with calomel electrode

temperature electrode

potentiometer


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electrode is often a calomel electrode, which supplies a constant potential (E° = +0.24 V

versus the standard hydrogen electrode) as determined by the half-reaction

Calomel is the trivial name for the compound Hg2Cl2. When both the reference and glass

electrodes are contained in a single unit, it is referred to as a combination electrode.

The potential of the glass electrode is proportional to the logarithm of the ratio of

[H3O+] inside and outside the electrode. The pH meter measures the total potential across the

two electrodes and displays this measurement on a scale calibrated in pH units. The pH meter

is an accurate and easy-to-use device for determining the pH of a solution.

Figure 2 on the next page shows a plot of pH versus volume of base added for the titration of

a strong acid with a strong base. There is very little change in pH when the base is initially

added. Below the equivalence point, the pH is a function of the amount of excess acid

present. Above the equivalence point, the pH is a function of the amount of excess base

present. The equivalence point for the titration of a strong acid with a strong base occurs

when [OH–] exactly equals [H3O+] in the solution; pH = 7.0. The situation in the case of the

titration of a weak acid with a strong base is somewhat different due to the fact that a weak

acid is only partially ionized in aqueous solution. A dynamic equilibrium exists which is

represented by the following equation:

Figure 2: Titration curve for the titration of a strong acid with a strong base


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The equilibrium expression for this reaction is:

where Ka is the acid-ionization constant for the weak acid. Let us assume that the initial

dissociation of the weak acid is negligible. The progressive addition of NaOH during the

titration decreases the concentration of HA and increases the concentration of its salt, NaA:

The presence of both HA and its salt, NaA, creates a buffer system which resists a large

change in pH. The ratio of [HA]/[A–] changes only slightly; therefore, according to Eq. 1, the

change in [H3O+] (or pH) must also be small. The pH increases slowly until the equivalence

point is approached (see Figure 3).

At the halfway point in the titration, exactly half of the HA originally present will have been

neutralized, and therefore the concentrations of HA and A– will be equal.

Substituting this information into Eq. 1, we obtain:

Figure 3: Titration curve for the titration of a weak acid with a strong base


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Thus, the ionization constant of a weak acid is equal to the hydronium ion concentration at the

halfway point in the titration; pKa = pH1/2. This relationship is valid only if the initial dissociation

of the acid is negligible. When the degree of dissociation is appreciable, as in the case of a

very dilute solution, the pH at the midpoint of the titration bears no relation to the value of Ka.

The subsequent rapid increase in pH and the inflection in the titration curve at the equivalence

point can be accounted for. As the equivalence point is approached, the concentration of

unreacted HA becomes progressively smaller so that successive increments of NaOH

neutralize a greater fraction of the HA remaining. This produces a large change in the [HA]/[A–

] ratio and, therefore, in the pH of the solution. At the equivalence point, the acid and base

have reacted completely to yield the salt, NaA. The pH at the equivalence point is determined

by the strength of the base, A–. The conjugate base of a weak acid is a strong base. It will

react with water to produce hydroxide ions (hydrolysis):

For this reason, it is not surprising to see a pH which is greater than 7 at the equivalence

point. Beyond the equivalence point, the pH is determined by the ion product for water:

The first small excess of NaOH greatly increases the concentration of OH–, concomitantly

decreasing the H3O+ concentration, and causing the pH to continue to increase. Well past the

equivalence point, the concentration of OH– becomes so large that only slight changes in pH

are produced.

This experiment involved titration a solution of HCl with a standardized solution of NaOH while

measuring the pH throughout the course of the titration. From the titration curve, the

concentration of the HCl solution can be determined. The second part would be titration a

sample of a commercial vinegar using a standard solution of NaOH. The active ingredient in

vinegar is acetic acid, which is a weak acid.

The acid-ionization constant of acetic acid is:


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From the titration curve, the determination of acetic acid concentration in commercial vinegar

can be achieved.

4.0 CHEMICALS AND EQUIPMENTS

1. 0.5 M HCl 2. 0.5 M NaOH

3. commercial vinegar 4. phenolphthalein solution

5. 150 mL beaker 6. volumetric pipet

7. Buret 8. pH meter

9. Stirring rod

5.0 PROCEDURES

5.1. Titration of a Strong Acid

5.1.1 Dispense 10.00 mL aliquot of the 0.5 M HCl solution (do not pipet directly from

the bottle) into a clean and dry 150 mL beaker using a carefully rinsed volumetric

pipet. Add exactly 75.0 mL of distilled water and 10 drops of phenolphthalein solution.

5.1.2 Fill a clean and carefully rinsed buret with the standardized 0.5 M NaOH solution

(record the exact molarity from the label). Record the initial buret reading in your

notebook.

5.1.3 Remove the pH electrode from the buffer solution. Thoroughly rinse the electrode

with distilled water, wipe the drops of water and place it in the acid solution such that

the tip is immersed. Stir the acid solution with the glassrod. Now arrange the buret

over the beaker so that the NaOH can be dispensed directly into the acid solution.

5.1.4 Record the pH reading when the pH is stable. This would be the reading of initial

pH when 0 mL of NaOH added. Begin the titration by adding, with stirring, about 1.0

mL of NaOH. Be careful not to splash any liquid out of the beaker. When the pH

reading is stable, stop stirring, then record the pH of the solution. Note the total volume

of NaOH that was added (the volume reading on the buret minus the initial volume

reading that you recorded in your notebook).

5.1.5 Continue to add base, record the volume of NaOH and pH of the solution

throughout the titration. Slow down as you approach the equivalence point (as

indicated by the appearance of the pink color)! As the pH readings approach 2,

reduce the amount of base added to 0.5 mL increments. Record in your notebook the

pH reading when the pink phenolphthalein endpoint color persists for 30 seconds. Add

increments of 0.5 ml until the pH reaches 12.


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5.2. Titration of vinegar

5.2.1 Refill the buret with the standardized 0.5 M NaOH solution provided. Record the

initial volume of NaOH in your notebook. Dispense a 10.00 mL aliquot of vinegar into a

clean, dry 150 mL beaker (do not pipet directly from the bottle). Dilute this aliquot with

exactly 75.0 mL of distilled water. Add 10 drops of phenolphthalein solution.

5.2.2 Titrate the vinegar as you did the HCl solution above.

6.0 RESULT AND CALCULATIONS

1. a) On the plot resulting from the titration of HCl with NaOH, draw the best curve

through the data points by hand. Locate the equivalence point and determine the

pH and volume at this point. Label the point where the phenolphthalein endpoint

became visible.

b) Using the molarity and volume of base solution required to reach the

equivalence point, calculate the concentration of the HCl solution which was in the

bottle in molarity.

2. a) On the plot resulting from the titration of vinegar with NaOH, draw the best curve

through the data points by hand. Locate the equivalence point and note the pH and

volume at this point. Label the point where the phenolphthalein endpoint became

visible.

b) Using the molarity and volume of base solution required to reach the equivalence

point, calculate the concentration of acetic acid in the vinegar bottle in molarity.

3. To find the pH at the halfway point of the curve, divide the volume of base

needed to reach the equivalence point by 2, and read off the corresponding pH from

the titration curve. Determine the value of Ka for acetic acid.

4. Using the concentration of acetic acid that you calculated in question 2.b)

above along with the literature value for Ka given in the introduction (1.74 x 10–5 M),

calculate the initial pH of the vinegar sample (after diluting itwith 75.0 mL of water

but before any base has been added).

5. Using your calculated concentration of acetic acid and the literature value for Ka,

calculate the pH of the vinegar sample at the equivalence point. For the volume of

base added at the equivalence point, use the amount determined from your titration

curve.


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6. Using your calculated concentration of acetic acid, calculate the pH of the

vinegar sample after 25.00 mL of standardized NaOH solution has been added (when

the solution is well past the equivalence point). At this point in the titration, you may

assume that the pH of the solution depends upon the

concentration of OH– ions.

7.0 QUESTIONS

1. Differentiate the equivalence point of the first titration (hydrochloric acid) with the

endpoint determined by phenolphthalein indicator.

2. Do you find the volume of NaOH needed to reach the equivalence point for strong

acid is less or more compare to weak acid. Explain your answer.

3. What is indicator and why it is important in titration procedure?


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EXPERIMENT 2

DETERMINATION OF WATER HARDNESS BY COMPLEXOMETRIC TITRATION

1.0 OBJECTIVE

1. To understand the principles of complexometric titration

2. To determine the total hardness of water samples


1. Ability to apply knowledge of mathematics to analyze the correct statistical

method for data analytical and to remember the steps in quantitative analysis.

2. Ability to solve complex problem by using techniques, skills and modern

engineering tools to analyze the concentration of analytes of various classical

titrimetric and gravimetric methods for mass determination.

3.0 INTRODUCTION

Water hardness is an expression for the sum of the calcium and magnesium cation

concentration in a water sample. These cations form insoluble salts with soap, decreasing

soaps cleaning effectiveness. They also form hard water deposits in hot water heaters. The

standard way to express water hardness is in ppm CaCO3 which has the formula weight of

100.1 g/mole.

Water hardness can be readily determined by titration with the chelating agent EDTA

(ethylenediaminetetraacetic acid). This reagent is a weak acid that can lose four protons on

complete neutralization; its structural formula is below. Each nitrogen atom has one unshared

electron pair and each doubly-bonded oxygen has two lone pairs that form coordinate

covalent bonds to metal ions.

N

N

O

HO

O

HO O

HO

O

OH

ethylenediamminetetraacetic acid Figure 1 : Ethylenediamminetetraacetic acid

The four acid (COOH) sites and the two nitrogen atoms all contain unshared electron pairs, so

that a single EDTA molecule can bind up to six sites on a given cation, utilizing the cation’s

vacant d orbitals. The complex is typically very stable, and the conditions of its formation can


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ordinarily be controlled so that it contains EDTA and the metal ion in a 1:1 mole ratio. In a

titration to establish the concentration of a metal ion, the EDTA that is added combines

quantitatively with the cation to form the complex. The end point occurs when essentially all of

the cation has reacted.

In this experiment you will standardize a solution of EDTA by titration against a standard

calcium solution made from calcium carbonate. You will then use the EDTA solution to

determine the hardness of an unknown water sample. Since both EDTA and Ca2+ are

colorless, it is necessary to use a special indicator to detect the titration end point. The

indicator, Eriochrome Black T, forms a stable wine-red complex, MgIn-, with magnesium ion. A

tiny amount of this coordination complex will be added to the solution before titration. As

EDTA is added, it will react with free Ca2+ and Mg2+ ions leaving the MgIn- complex alone until

essentially all of the calcium and magnesium have been converted to chelates. At this point,

the EDTA concentration will increase sufficiently to remove Mg2+ from the indicator complex.

The indicator reverts to its acid form, which is sky blue, establishing the end point of the

titration.

The titration is carried out at pH 10 NH3-NH4+ buffer, which keeps the EDTA (H4Y) mainly in

the half-neutralized form, H2Y2-, where it complexes the Group IIA ions very well but does not

react readily with cations such as Fe3+ that may be present as impurities. Taking H4Y and H3In

as the formulas for EDTA and Eriochrome Black T respectively, the equations for the

reactions that occur during the titration are as follows.

Titration:

Ca2+ + H2Y2- � CaY2- + 2H+

Ca2+ + MgY2- � CaY2- + Mg2+

End point:

Mg2+ + HIn2- � MgIn- + H+

MgIn- + H2Y2- � MgY2- + HIn2- + H+

Since the indicator requires a trace of Mg+2 to operate properly, you will add a little

magnesium ion to each solution and titrate it as a blank to adjust for the volume of EDTA

solution required to react with the added magnesium.

4.0 CHEMICALS AND EQUIPMENTS

1. 50 mL buret 6. 25 mL volumetric pipet

2. 1 mL volumetric pipet 7. 10 mL and 100 mL graduated cylinders

3. 250 mL Erlenmeyer flasks 8. Analytical balance

4. EDTA 9. pH 10 buffer

5. Eriochrome Black T indicator

1 mL 0.01 M EDTA = 1 mg CaCO3


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5.0 PROCEDURE

5.1 Preparation of 0.01 M EDTA solution

Dissolve 0.360 - 0.380 g disodium ethylene diamine tetra-acetate dihydrate (EDTA)

(analytical reagent grade) in distilled water in a 100 mL volumetric flask. Calculate the

actual concentration of the solution.

5.2 Determination of Ca2+ and Mg2+ in Water sample

5.1.1 Pipet 25.0 mL of water into a 250 mL Erlenmeyer flask (conical flask). Add 15

mL of distilled water to the flask. Add 10 mL of pH 10 buffer and five drops of

Eriochrome Black T indicator to the flask.

5.1.2 Titrate with EDTA from a 50 mL buret until the indicator color changes from

wine-red to blue. Titrate until every trace of purple has just disappeared. Save

a completed titration for a color comparison.

5.1.3 Repeat the titration at least two more times. The %RSD for three trials should

be less than 2.0%. If not, perform additional trials until this criterion is satisfied.

5.1.4 Perform few blank titrations following steps 5.2.1 to 5.2.3 with 25.00 mL

distilled water.

5.1.5 Subtract the average blank volume of sample titration value.

6.0 RESULTS AND DISCUSSION

Table 1: Titration of EDTA blank

Trial1 Trial2 Trial3

Final EDTA volume

Initial EDTA volume

Blank volume

Average blank volume

Table 2: Titration of hard water sample

Trial1 Trial2 Trial3

Final EDTA volume

Initial EDTA volume

Titration volume

Average blank volume

Net titration volume

Average net titration volume


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7.0 CALCULATIONS

1. Calculate the molarity of the EDTA titrating solution.

2. From the EDTA solution molarity and the average (of three trials) volume of EDTA

needed to titrate each aliquot of your unknown hard water sample, calculate the

molarity of calcium in your unknown.

3. From the molarity of the unknown hard water sample, calculate the mass of calcium

carbonate per liter in the hard water.

4. Calculate the hardness of your unknown hard water sample in ppm CaCO3.

a) Molarity of EDTA soln

S1 = X 0.01

= X 0.01

b) Molarity of hard water sample

S2 =

V1 is the volume of EDTA required to titrate V2 ml (25 ml) of hard water.

c) Content of CaCO3

= 100.09 x S2 x 1000 ppm

= ppm

Actual amount of EDTA taken

Wt. to be taken for 0.01 M EDTA solution

Actual amount of EDTA taken

0.37224

V1 x S1

V2

100.09 x V1 x S1 x 1000

V2


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EXPERIMENT 3

DETERMINATION OF LIMONENE CONCENTRATION BY GAS CHROMATOGRAPHY

1.0 OBJECTIVE

1.1 To familiarize the students with essential theory concerning gas chromatography.

1.2 To be able to prepare standard limonene at different concentrations and establish standard curve.

1.3 To quantify the concentration of limonene in the citrus peel extract by using gas

chromatography.


1. Ability to evaluate the sample concentration and analyze data of complex

chromatography and spectroscopic problem in order to design solution for the

problems.

3.0 INTRODUCTION

Gas chromatography is a powerful separation technique for detection of volatile

organic compounds. Combining separation and on-line detection allows accurate quantitative

determination of complex mixtures, including traces of compounds down to parts per trillions

in some specific cases. Gas chromatography has found in a variety of analytical uses, which

include qualitative analyses of illicit drug samples and forensics evidence, trace analyses of

pesticides and other toxic residues present in soil and ground water samples, and performing

quality control analyses in both the pharmaceutical and food industries. Thus, it has become

essential for students to know these basic analytical techniques to a better prepare for future

careers in industrial settings.

The essential oils, derived from natural sources, are found in many common

household products, including fruit juices, spices, flavor components in beverages and bakery

products, and fragrances in incense and many household cleaning products. Many of these

essential oils belong to a family compounds known as terpenes and terpenoids. Terpenes are

small organic hydrocarbon molecules; they may be cyclic or acyclic, saturated or unsaturated.

Terpenoids are oxygenated derivatives of terpenes, which may contain hydroxyl groups or

carbonyl groups. Terpenes such as limonene may be found in abundance in oil sacs located

in the outer, colored or flavedo portion of the rinds of many common citrus fruits. Limonene

may be readily isolated through a variety of methods, such as cold pressing, steam distillation,

or extraction. The concentration of limonene in the extracts can be determined using gas

chromatography by comparing of the retention time with that of a standard of the pure

substance.


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4.0 MATERIALS AND EQUIPMENT

1. Citrus peel 5. Pentane

2. Anhydrous sodium sulfate 6. 100 ml Separatory funnel

3. 50 ml Beaker 7. 50 ml measuring cylinder

4. GC sample vial 8. Micropipette

5. Water bath

5.0 PROCEDURES

5.1 PREPARATION OF THE STANDARD SOLUTION

5.1.1 Prepare the standard concentration of limonene for 0.1, 0.2, 0.3 and 0.4 ppm

solution from the stock solution with series of dilution with pentane.

5.1.2 Measure the peak area of the standard samples using GC.

5.1.3 Draw the linear curve for the standard samples utilizing the peak area reading

against each concentration.

5.1.4 Calculate the linear equation from the plot. Y = mX +c.

5.1.5 Calculate the unknown concentration of limonene utilizing the linear equation of the

standard samples.

5.2 PREPARATION OF SAMPLE

5.2.1 Place approximately 5g of grinded citrus peel in a 100 ml separatory funnel.

5.2.2 Add 10 ml of pentane in the separatory funnel and shake for 10 minutes. Remember

to release the pressure by frequently vent the separatory funnel (see Figure 1).

Drain of the lower layer of extract into a beaker and repeat the procedure.

5.2.3 Combine the extracts in a 50 ml beaker. Add approximately 1g of anhydrous sodium

sulfate and leave for 15 minutes. Filter the solution through a filter paper on a glass

funnel.

5.2.4 Transfer the extract to a 50 ml beaker wraped with Aluminium foil and remove all

pentene using water bath at 35oC in a fume hood for 15 minutes.

5.2.5 For the chromatographic analysis, dilute the citrus oil with 1.0 ml of pentane and

inject 1.00 µl of the resulting solution into the GC.


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Figure 1: A set up for liquid-liquid extraction using separatory funnel

Figure 2: Technique to vent the separatory funnel


6.1 Plot a standard curve of limonene concentration (mg/l) versus the peak area for

limonene using the obtained data.

6.2 Using the standard curve, calculate the concentration of limonene in the sample.

7.0 QUESTIONS

(1) Why should you vent the separatory frequently during the extraction?

(2) What is the purpose of using anhydrous sodium sulfate in the experiment?

(3) The extracts should immediately injected into the GC injection port. Give one

appropriate reason.

(4) Why there is the need to compare the observed retention time values to the standard

values?


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EXPERIMENT 4

DETERMINATION OF ACTIVE DRUG SUBSTANCE IN A PHARMACEUTICAL

FORMULATION USING ULTRAVIOLET AND VISIBLE SPECTROMETRY (UV/VIS)

1.0 OBJECTIVE

1. To determine the specific absorbance of the supplied drug standard using uv-vis

spectrometry

2. To determine the content of drug in the supplied sample using uv-vis spectrometry.

2.0 LESSON OUTCOME

1. Ability to evaluate the sample concentration and analyze data of complex

chromatography and spectroscopic problem in order to design solution for the

problems.

3.0 INTRODUCTION

UV-vis spectroscopy is the measurement of the wavelength and intensity of absorption

of near-ultraviolet and visible light by a sample. Ultraviolet and visible light are energetic

enough to promote outer electrons to higher energy levels. UV-vis spectroscopy is usually

applied to molecules and inorganic ions or complexes in solution. The uv-vis spectra have

broad features that are of limited use for sample identification but are very useful for

quantitative measurements. The concentration of an analyte in solution can be determined by

measuring the absorbance at some wavelength and applying the Beer-Lambert Law.

Spectrophotometry is the quantitative measurement of the reflection or transmission

properties of a material as a function of wavelength. The advantages of these methods are

low time and labor consumption. The precision of these methods is also excellent. The use of

UV–Vis spectrophotometry especially applied in the analysis of pharmaceutical dosage form

has increased rapidly over the last few years.

The extent of absorption of radiation by an absorbing system at a given

monochromatic wavelength is described by the two classical laws which relate the

intensity of radiation incident on the absorbing system (I0), to the transmitted intensity

(I). Lambert's (or Bouguer's) Law that at a given concentration (c) of a homogeneous

absorbing system, the transmitted intensity (I) decreases exponentially with increase in

path length (b). The complementary Beer's Law states that for a defined path length (b),

the transmitted intensity (I) decreases exponentially with increase in concentration (c) of a

homogeneous absorbing system. Combination of these observations gives the familiar

Beer-Lambert Law.

A= log1/T = log I0/I = εbc Eq. 1

where, A = Absorbance; T = Transmittance; ε = molar absorptivity; and c = concentration of

analyte.


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Figure 1 : Schematic diagram of UV-VIS spectrometry system

4.0 MATERIALS AND EQUIPMENT

1. UV-Vis spectrophotometer

2. Paracetamol tablet

3. 0.1M NaOH

4. Distilled water

5. Beaker

6. 200ml and 100ml volumetric flask

7. Pipette

8. Peslte and mortar

9. Sonicator

5.0 EXPERIMENTAL: DETERMINATION OF PARACETAMOL IN TABLETS

5.1 Take 2-3 paracetamol tablets and measure the average weight and percentage deviation

of the tablets.

5.2 Grind the tablets properly and transfer an amount of powder approximately equivalent to

150 mg active ingredient into a 200 ml volumetric flask.

5.3 Add 50 ml of 0.1 M sodium hydroxide, shake for 5 minutes. Fill 2/3 of the volume with

distilled water and put in sonicator for 5 minutes. Add sufficient distilled water to make up

to the mark.

5.4 Dilute 10 ml of the solution to 100 ml with distilled water in a volumetric flask.

5.5 Again take 10 ml of the resulting in 100 ml volumetric flask. Add 10 ml 0.1M sodium

hydroxide distilled water to make up to the mark.

5.6 For blank solution dilute 10 ml of 0.1M sodium hydroxide with distilled water in a 100 ml

volumetric flask.

5.7 Scan the absorbance of the sample solution with uv/vis spectrophotometer from 200 – 400

nm wavelength. What are the λmax value(s) of the sample? Include the chromatogram in

the lab report.

5.8 Measure the absorbance of the resulting solution at the maximum at 257 nm at least three

times each time changing the sample solution.


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6.1 Please show detail calculation.

(a) Average wt. of tablets

x = mg

(b) Calculate the content of paracetamol in each tablet.

w

xAC

×

××=

ε

200 mg

where A= absorbance, ε = specific absorbance [A(1%, 1 cm) at 257 nm] = 715, w =

weight of powder taken for sample preparation (mg) and x = average weight of the

tablets (mg).

(c) How many percentage of paracetamol content of the stated amount on the lebel?

Content 00100

500×=

C of the claimed amount

6.2 Compare the concentration of paracetamol obtained in the experiment to the stated

amount of the commericalized paracetamol.

6.3 Discuss any discrepency and give suggestion for improvement.

Weight of 10 tablets

10


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EXPERIMENT 5

INTERPRETATION OF SPECTRA OF AN UNKNOWN SAMPLE USING FOURIER TRANSFORM INFRA RED (FT-IR) SPECTROSCOPY

1.0 OBJECTIVES

1.0 To familiarize with the operation of Fourier Transform –Infra Red (FT-IR) spectrometer and solid sample handling techniques.

1.1 To learn how to interpret IR spectra.


1. Ability to evaluate the sample concentration and analyze data of complex chromatography and spectroscopic problem in order to design solution for the problems.

3.0 INTRODUCTION

The energy of most molecular vibrations corresponds to that of the infrared region of

the electromagnetic spectrum.Molecular vibration maybe detected and measured either in an

infrared spectrum or indirectly in a Raman spectrum. The most useful vibrations occur in the

narrower range of 2.5-1.6 µm which most infrared spectrometers covers. The position of an

absorption band in the spectrum in microns or in reciprocal of wavelength, cm-1. The usual

range of an infrared spectrum is therefore between 4000 cm-1 at the high frequency end and

625 cm-1 at the low frequency end (Williams and Flemings, 1996).

A complex molecule has a large number of vibrational modes which involve the whole

molecule.To a good approximation, however some of these molecular vibrations are

associated with the vibrations of individual bonds or functional groups (localized vibrations)

while others must be considered as vibrations of the whole molecule. The localized vibrations

are either stretching, bending, rocking, twisting or wagging. For example, the localized

vibrations of the methylene group are


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H

C

H

symmetricstretching

H

C

H

asymmetricstretching

H

C

H

bendingor scissoring

H

C

H

rocking

H

C

H

twisting

H

C

H

wagging

Many localized vibrations are very useful for the identification of functional groups (Williams

and Flemings, 1996).

4.0 EQUIPMENT, CHEMICALS AND GLASSWARE

4.1 KBr pellet die kit 4.2 Mortar and pestle 4.3 FT-IR spectrometer 4.4 Uknown sample 1

5.0 PROCEDURE

5.1 OBTAIN IR SPECTRUM OF UNKNOWN SAMPLE VIA A KBr PELLET 5.1.1. Place 200 mg of powdered Pottasium Bromide (KBr) in the mortar. Add about 2

mg of unknown sample 1 (100:1). Mix well and grind together till uniform. Do all this quickly, as the KBr will absorb water from the atmosphere, and this makes it difficult to press a good pellet.

5.1.2. Put 1 bolt on the bottom of the pellet holder – note the shiny surface of the bolt,

this will form the surface of the pellet so take care and do not scratch. Add about 30-50 mg of the mixture into the evacuable die. Screw second bolt on top and tighten both with the ratchets.

5.1.3. Place the evacuable die on a hydraulic press 10 ton. Close the pressure. Press

at 10 ton of pressure for 3 minutes. Release the pressure, remove the die from press, disassemble the die and remove the KBr pellet.

5.1.4. Record a spectrum (after having already collected a background) by placing

pellet holder with pellet in place into the transmission cell holder in sample chamber. Your instructor will show you how to navigate the software.

5.2 ACQUIRE SPECTRA OF THE UNKNOWNS 5.2.1. Record IR spectra of the unknowns.


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6.0 RESULTS & DISCUSSION

6.1 For the unknown, tabulate all the wavelengths of significant absorption (select the wavelengths of maximum absorption in the 2- to7 µm region). Deduce the possible chemical structure or functional group for each absorption.

7.0 QUESTION

7.1. Why KBr is used to make the pellet?

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