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 13 complexometric titration Experiment 3: Determination of active drug substance in a 18 pharmaceutical formulation using ultraviolet and visible spectrometry (uv/vis) Experiment 4: Determination of limonene concentration 21 by gas chromatography Experiment 5: interpretation of spectra of an unknown sample using fourier transform infra red (FT-IR) spectroscopy 25 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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 the set skills that are necessary to succeed in more advanced chemistry courses, such as physical chemistry and instrumental analysis. Additionally, these skills are essential to anyone working as a laboratory technician in industry and are essential to anyone that conducts research. ii Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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 iii Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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 experiments should be mentioned. 3. Objectives The objectives of the experiment are clearly stated. 4. Summarized experimental procedure Attached the summarized experimental procedure to the lab report. iv Analytical Chemistry (ERT 207) Sem I 2013/2014 5. Laboratory Module 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. 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. v Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module EXPERIMENT 1 POTENTIOMETRIC TITRATION 1.0 OBJECTIVE 1.1 To understand the principle of pH change in acid-base titrations using a pH meter. 1.2 To determine the concentrations of the acidic solutions as well as the acidionization constant of a weak acid. 2.0 COURSE OUTCOME CO3 : Ability to differentiate and to calculate concentration of analytes various titrimetric methods(acid-base,complexation,redox and precipitation) . 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: glass electrode combine with calomel electrode potentiometer temperature electrode Figure 1: A pH meter 6 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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 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 7 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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 on the next page). 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 8 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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. 9 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module The acid-ionization constant of acetic acid is: From the titration curve, the determination of acetic acid concentration in commercial vinegar can be achieved. 4.0 APPARATUS AND CHEMICALS 1. 2. 3. 4. 5. 6. 7. 8. 5.0 0.5 M HCl 0.5 M NaOH commercial vinegar phenolphthalein solution 150 mL beaker volumetric pipet buret pH meter 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. 10 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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. 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. 11 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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. 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? 12 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module EXPERIMENT 2 DETERMINATION OF WATER HARDNESS BY COMPLEXOMETRIC TITRATION 1.0 2.0 OBJECTIVE 1.1 To understand the principles of complexometric titration 1.2 To determine the total hardness of water samples COURSE OUTCOME CO1 : Ability to calculate and perform the correct stastitical method for data analytical and to remember the steps in quantitative analysis CO3 : Ability to differentiate and to calculate concentration of analytes various titrimetric methods(acid-base,complexation,redox and precipitation) . 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. 13 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module HO O HO O N N O O HO OH 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 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. 14 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 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. 1 mL 0.01 M EDTA = 1 mg CaCO3 4.0 5.0 APPARATUS AND CHEMICALS 4.1 50 mL buret 4.2 25 mL volumetric pipet 4.3 1 mL volumetric pipet 4.4 10 mL and 100 mL graduated cylinders 4.5 Four 250 mL Erlenmeyer flasks 4.6 Analytical balance 4.7 EDTA 4.8 pH 10 buffer 4.9 Eriochrome Black T indicator 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 5.2.1 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. 15 Analytical Chemistry (ERT 207) Sem I 2013/2014 5.2.2 Laboratory Module 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. 5.2.3 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. 5.2.4 Perform few blank titrations following steps 5.2.1 to 5.2.3 with 25.00 mL distilled water. 5.2.5 6.0 Subtract the average blank volume of sample titration value. RESULTS AND DISCUSSION Table 1: Titration of EDTA blank Trial1 Trial2 Trial3 Trial4 Trial5 Trial3 Trial4 Trial5 Final EDTA volume Initial EDTA volume Blank volume Average blank volume Table 2: Titration of hard water sample Trial1 Trial2 Final EDTA volume Initial EDTA volume Titration volume Average blank volume Net titration volume Average net titration volume 16 Analytical Chemistry (ERT 207) Sem I 2013/2014 7.0 Laboratory Module CALCULATIONS 1. Calculate the molarity of the EDTA titrating solution. 2. From the EDTA solution molarity and the average (of five 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 Actual amount of EDTA taken S1 = Wt. to be taken for 0.01 M EDTA solution X 0.01 Actual amount of EDTA taken = X 0.01 0.37224 b) Molarity of hard water sample V1 x S1 S2 = V2 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 100.09 x V1 x S1 x 1000 = ppm V2 17 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module EXPERIMENT 3 DETERMINATION OF ACTIVE DRUG SUBSTANCE IN A PHARMACEUTICAL FORMULATION USING ULTRAVIOLET AND VISIBLE SPECTROMETRY (UV/VIS) 1.0 OBJECTIVE 1.0 To determine the specific absorbance of the supplied drug standard using uv-vis spectrometry 1.1 2.0 To determine the content of drug in the supplied sample using uv-vis spectrometry. COURSE OUTCOME CO5 : Ability to understand the spectroscopic principles and to calculate concentration. 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. 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 (I 0), 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 18 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module Where, A = Absorbance; T = Transmittance; ε = molar absorptivity; and c = concentration of analyte. 4.0 5.0 EQUIPMENTS, CHEMICALS AND GLASSWARE: 4.0 UV-Vis spectrophotometer 4.1 Paracetamol tablet 4.2 NaOH 4.3 Distilled water 4.4 Beaker 4.5 Volumetric flask 4.6 Pipette etc. EXPERIMENTAL: DETERMINATION OF PARACETAMOL IN TABLETS 5.0 Take 2-3 tablets and measure the average weight and percentage deviation of the tablets. 5.1 Grind the tablets properly and transfer an amount of powder approximately equivalent to 150 mg active ingredient into a 200 ml volumetric flask. 5.2 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. 5.3 Dilute 10 ml of the solution to 100 ml with distilled water in a volumetric flask. 5.4 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. 5.5 For blank solution dilute 10 ml of 0.1M sodium hydroxide with distilled water in a 100 ml volumetric flask. 5.6 Scan the absorbance of the sample solution (step 5.5) with uv/vis spectrophotometer from 200 – 400 nm wavelength. What are the max value(s) of the sample? Include the chromatogram in the lab report. 19 Analytical Chemistry (ERT 207) Sem I 2013/2014 5.7 Laboratory Module Measure the absorbance of the resulting solution at the maximum at 257 nm at least five times each time changing the sample solution. 6.0 Calculation (a) Average wt. of tablets Weight of 10 tablets x= mg 10 (b) Calculate the content of paracetamol in each tablet. C A 200 x mg w where A= absorbance, = 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). (b) How many percentage of paracetamol content of the stated amount on the lebel? Content C 100 0 0 of the claimed amount 500 20 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module EXPERIMENT 4 DETERMINATION OF LIMONENE CONCENTRATION BY GAS CHROMATOGRAPHY 1.0 OBJECTIVE 1.1 To introduce the variety of laboratory techniques especially an extraction technique. 1.2 To familiarize the students with essential theory concerning gas chromatography. 2.0 COURSE OUTCOME CO2 : Ability to classify and use separation techniques and to use gravimetric methods for mass determination. CO4 : Ability to apply the chromatography principles and to interpret and calculate peak height for concentration determination. 3.0 INTRODUCTION 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. 21 Analytical Chemistry (ERT 207) Sem I 2013/2014 4.0 5.0 Laboratory Module MATERIALS AND EQUIPMENTS 4.1 Sample contains mixture of chemicals 4.2 Pentane 4.3 anhydrous sodium sulfate 4.4 separatory funnel 4.5 Aluminium foil 4.6 Water bath for evaporation 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 ml/L solution from the stock solution with series of dilution with pentane. 5.1.2 Measure the peak height of the standard samples using GC. 5.1.3 Draw the linear curve for the standard samples utilizing the peak height 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 2 ml of chemical mixtures in a 50 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 two times. 5.2.3 Combine the extracts in a 100 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 0.25 μl of the resulting solution into the GC. 22 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module Figure 1: A set up for liquid-liquid extraction using separatory funnel Figure 2: Technique to vent the separatory funnel 23 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module 6.0 RESULTS AND DISCUSSION 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? 24 Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module EXPERIMENT 5 INTERPRETATION OF SPECTRA OF AN UNKNOWN SAMPLE USING FOURIER TRANSFORM INFRA RED (FT-IR) SPECTROSCOPY 1.0 OBJECTIVE 1.1 To familiarize with the operation of Fourier Transform –Infra Red (FT-IR) spectrometer and solid sample handling techniques. 1.2 To learn how to interpret IR spectra. 2.0 COURSE OUTCOME CO5 : Ability to understand the spectroscopic principles and to calculate concentration. 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 H H C symmetric stretching H H H H C C asymmetric stretching bending or scissoring H H C rocking 25 H H H H C C twisting wagging Analytical Chemistry (ERT 207) Sem I 2013/2014 Laboratory Module Many localized vibrations are very useful for the identification of functional groups (Williams and Flemings, 1996). . 4.0 EQUIPMENTS, CHEMICALS AND GLASSWARE 4.1 4.2 4.3 4.4 5.0 KBr pellet die kit Mortar and pestle FT-IR spectrometer Uknown sample 1 PROCEDURE 5.1 OBTAIN IR SPECTRUM OF UNKNOWN SAMPLE VIA A KBR PELLET 5.1.1. Place about 100-200 mg of powdered KBr in the mortar. Add to this about 2 spatula tips (no more than 1-2 mg) of unknown sample 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 cavity. Screw second bolt on top and tighten both with the ratchets. 5.1.3. Then loosen and remove the bolts, leaving the pellet in place. Ideally it would be clear, uniform and unblemished; however, if light can shine through the pellet and it doesn’t have a hole then it may still be usable. If you want to make another pellet simply rinse the press with water and then follow with acetone and try again. 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. 6.0 RESULT & DISCUSSION 6.1 7.0 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. QUESTION 7.1. Why KBr is used to make the pellet? 26