Revisions 1/18/2005 1/11/2006 4/15/2007
UNIVERSITY OF HAWAI’I AT MANOA CHEMISTRY DEPARTMENT
CHEMISTRY 274L LABORATORY EXPERIMENTS
T.A.: OFFICE: OFFICE HOURS:
2 Introduction to the Analytical Chemical Laboratory The purpose of this laboratory is to provide an opportunity for the student to become familiar with a number of analytical techniques that most chemists use today. In addition to the more classical methods of gravimetry and titrimetry, the experiments will introduce you to the more sophisticated methods of spectrophotometry, potentiometry, and chromatography. The analytical laboratory differs in many ways from other undergraduate laboratory courses. In order to successfully analyze the samples, the student must be very conscientious and meticulous. The procedures should be followed exactly if you expect to obtain reliable results – accuracy and precision are the goals. PLAN AHEAD. You will often encounter delays while waiting for things to heat or cool. Organize your work so that these delay times are not wasted. Keep your work area clean – sloppy work is rarely accurate nor precise. In some respects the analytical laboratory appears more difficult and demanding. However, this is compensated for by the rewarding experience of obtaining good results. Success is practically guaranteed if you know what you are doing and work carefully. The Laboratory Notebook The laboratory notebook is an important part of this course because it is the complete record of everything you do. Your notebook should meet the basic requirements of any research or control laboratory. The notebook should be bound and each page numbered. Square rule paper is convenient for tables and graphs. It must begin with a Table of Contents, listing each experiment and corresponding page numbers. All entries must be made in ink, with errors or data to be ignored neatly lined out, never erased. A neat, well-organized, and easily read notebook will make your TA happy, but the primary purpose is to compile a complete record of your original data and calculations. Enter all data immediately into your notebook—not on loose pieces of paper which you may lose or the TA may confiscate. A uniform format is essential in your laboratory notebook. This is important because it will increase the organization of your reports and your lab work. Use the following format:
Title: ____________________ Date Performed: ____________________ Date Submitted: ____________________ Partner’s Name (if any): ____________________ I. Purpose – One or two sentences which state the essence of the experiment: include the method and species that is being determined. Theory – In this section you should include the following: What are the chemical reactions pertinent to this experiment? Introduce all equations and formulas and define all variables. What is measured and how is it related to the results? What instruments are used and how do they work? What equations are used in the calculations?
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3 III. Procedure – Reference the source of the procedure. For instance, ―The procedure followed for experiment number 2 was found in the Chem 274L laboratory handout from the Chemistry Department, University of Hawaii, Spring 1984.‖ Note any and all exceptions to the referenced procedure. List the unknown identification number, the starting amounts of materials, and any dilutions. Include a physical description of the unknown. Data and Calculations – Tabulate all original data, including weighing data. Following the table, include one sample calculation of each type necessary for the data analysis. Results Use the following format to report results: The results of unknown (ID number) are: Trial # 1 2 3 ¯ X S.D. Rel. S.D. VI. Concentration ____________________ ____________________ ____________________ ____________________ ____________________ ____________________
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Conclusion – Look over your data and explain any trends or anomalies. Support the discarding of any results by performing a Q-test on the data (see pages 167-168 of Skoog, West, Holler, and Crouch, 8th edition). Analyze your experimental errors. In particular, list the sources of error and the magnitude of the error generated in each case. Suggestions for modification of the lab procedure and handout should be included in this section. List observations of unexpected or inconsistent results.
The mechanics of the lab write-up are important. The following guidelines for equations, tables, and figures are recommended. 1. All tables and figures must be numbered and titled. Titles should be clear and descriptive. 2. All equations must be numbered and set off by spaces. 3. Specific titles are required for all columns in tables. Use asterisks and footnotes for long titles and supplemental information. 4. All pieces of graph paper and chart paper must be neatly trimmed and attached to the lab notebook by tape, glue, or staples. 5. Use the ―RIGHT HAND‖ pages of the notebook for format-graded report. 6. Use the ―LEFT HAND‖ pages of the notebook for class notes, exp’t. notes and calculations.
4 Due Dates The lab reports are due one week from the day the experiment is completed. The penalty for late lab reports is two points for each weekday after the due date. A 3‖ x 5‖ card with a copy of your unknown results is due at the next class meeting after the experiment is completed. Evaluation of Performance Your performance in the determination of ―unknowns‖ will be a major part of the evaluation of your final grade. One ―unknown‖ sample per student will be dispensed on the day it is required. The results will be graded on scale of 65 points. Both accuracy and precision count. If a grade of 40 or less is received, the work may be repeated at the discretion of the T.A. However, repeated work is subject to a 20% grade penalty. A short quiz will be given at the beginning of each new experiment. These quizzes will cover material given in the pre-lab lectures and instructions. In general the student is expected to know the underlying principles of the methods used—both the chemistry and the instrumentation. The final grade will be based on the following scheme: Unknown Analysis Lab Notebook Quizzes T.A. Evaluation Total 65 10 20 5 100
Rewards for accurate and precise analytical work will be provided for analyses that are reported to within 4 parts per thousand of the known value with precisions of less than 1.5%. Regardless of scores on notebooks or quizzes for such a lab, students obtaining such results will be given 95 points. (Note that these rewards WILL NOT be provided in the case of students having to go back track and recalculate their results after the fact. Normal grading will apply in such cases.) Subsequent TA evaluations will likely push the total grade for such an experiment to 100 points.
Time Constraints Time is of the essence. You must stay on schedule. The instruments and some reagents will only be available for a limited time. No one is permitted to work in the laboratory unless a TA is present. To facilitate the completion of your lab work within the allotted time, it is highly recommended that you read the lab handout and create a flow chart of the experiment in your notebook before coming to lab. Read the appropriate chapter in the Chem 274 text and take notes during the pre-lab lectures.
5 Experiment 1. Weighing and Statistical Evaluation of Data: Is a Penny Worth Its Weight in Copper? The Way to Weigh The measurement of the mass of an object is perhaps the oldest, most common, and simplest of all the measurements you will make in the laboratory. It is also one of the more precise measurements that can be made. With a high quality balance and careful technique, the mass of a 100 g object can be accurately determined to the nearest 0.1 mg, or 1 part per 1 million. In this course, 1 part per 1000 will normally be adequate in the final result, but even this level of precision requires care and attention to detail. A spring scale determines weight, which depends on both the mass and the gravitational force. A balance compares the mass of an object to the mass of standard masses (weights) which is independent of gravity. Thus, an object weighs less on the moon than on the earth, but its rest mass is constant everywhere. When you ―weigh‖ an object on a balance, you are really ―massing‖ it: however, as long as gravity is constant, we use ―weight‖ and ―mass‖ interchangeably without confusion. Our electronic analytical balances are top-loading and have a capacity of 100 g and a precision of 0.1 mg, or 1 part in 1.0 million. They are delicate instruments and would require ~$3000 to replace. Common sense dictates that they must be treated with respect or they will give you erroneous results. Generally, the first operation in any analysis is to weigh a sample. This should be one of the most accurate and precise steps. With a little practice, you can do this in a minute or so. For your own benefit, and that of other users, there are some precautions you must observe: 1. Absolute cleanliness is imperative. Avoid spills in the balance and the area around it. If you do have an accident, clean it up immediately. Never place a wet or hot object on the balance pan. Liquids and solid chemicals must be weighed in appropriate containers. Use only the balance assigned to you. Report malfunctions to the TA.
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You will be given a demonstration and instructions on the Way to Weigh. Briefly: 1. 2. Be sure all doors are closed. Be sure that doors are in a clean condition.
6 You should get a correct result every time, but here are some potential causes of trouble: 1. 2. 3. 4. 5. Room doors open (air currents). Balance is out of adjustment, call TA. Object should be at room temperature. Static charge on surface of object should be avoided. Do not record your weight on scrap paper.
The Correction for Buoyancy All objects experience a buoyant force equal to the weight of the displaced fluid—―Archimedes Principle.‖ Two objects of the same mass will displace different amounts of fluid (air or water) if they have different densities. Similarly, an object will have a different weight in air than in water (but its mass is constant!). Thus, the buoyant force correction on mass, fb, is
Mo f b = o x df V o lu m e o f = O b je ct X D e n sity o f d ispla ce d fluid
and the true mass, Mt, is
M M t = Mo + o df
o
When you weigh an object, the object and the weights in the balance experience a different buoyant force because of their different densities. You can easily show that
M t = M wt + M wt a ir obj _ wt a ir
where Mt = true mass or weight in vacuum Mo or Mwt = apparent mass of object or mass of weights air = density of air = 0.0012 g/ml wt = density of weights = 7.8 g/ml
7 The buoyancy correction for solids is usually negligible, but not for liquids. Find out for yourself by calculating buoyancy correction for weighing cork, = 0.20 g/mL; water, = 1.00 g/mL; sodium chloride, = 2.17 g/mL; and copper metal = 8.94 g/mL. Although buoyancy corrections are not required in Experiment #1 (why?), you will find that the buoyancy corrections required in Experiment #2 (why?) have been conveniently included in your tabulated reference data used in your calculations. Experimental Procedure: In this experiment you will test your ability to weigh accurately and quickly, use statistical methods to evaluate your data, and determine the market value of the copper in a penny. You will use ten pennies which constitute a representative sample of a large population of pennies. 1. 2. 3. 4. Study the material on statistical evaluation and in the Chem 274 text. Familiarize yourself with the proper operation of the balance. Obtain ten clean pennies from your T.A. Weigh each penny to the nearest 0.1 mg. It should not be necessary to wear gloves if your hands are clean and dry. Record the individual weights and minting dates in your notebook. Weigh the ten pennies at one time and compare the sum of the individual weights with the weight of all ten at once. Apply the Q test and reject any suspect values. Calculate the following: (a) the mean (b) the median (c) the average deviation (d) the standard deviation, s (e) the relative standard deviation (f) the student t test at the 90% and 99% probability levels. Calculate the value of the copper in 100 pennies. On 1/18/2007, the market price of copper metal was $2.50/lb. Tabulate and summarize your results.
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8 Experiment 2. Calibration of Volumetric Glassware Volumetric Measurements Volumetric flasks, pipets, burets, and calibrated syringes are used to measure accurate volumes of liquids. With care, you should be able to obtain routinely an accuracy of 0.1%. The principal sources of error are (a) dirty glass surfaces, (b) errors in reading the miniscus (parallax), and (c) temperature variations. It is absolutely essential that the inside surface be scrupulously clean and free from the inevitable, invisible greasy film that causes droplets to hang up. If the glassware does not drain completely without visible droplets, it is too dirty to use. Usually, a brushing with detergent solution will suffice. If not, a more drastic soaking in ―cleaning solution‖ is called for. CAUTION: Cleaning solution (dichromate/sulfuric acid) is treacherous—it eats practically everything except glass. Consult your TA for directions and precautions. Finally, rinse well, first with tap water, then with de-ionized water. Fill it, drain it, and look for droplets. Once you have it clean, keep it that way by storing it filled with de-ionized water. Volumetric flasks are less critical except for the neck in the region of the calibration mark. Volumetric glassware is calibrated by the manufacturer and is accurate enough for routine work. For highest accuracy, you must calibrate it yourself, or pay many dollars for NIST-certified quality. Certain tolerances are allowed, and if your calibration differs from the nominal value, you should use the value you measure as the true volume. As an exercise, you will calibrate a 25 mL pipet and a 50 mL buret. Your TA will demonstrate the proper technique for using volumetric glassware. CAUTION: Never suck on a pipet—it makes a very good straw and many solutions are poisonous or corrosive to your mouth/teeth. Experimental Procedure: Part A. Calibration of a 25 mL pipet 1. 2. Record the temperature of the water to be used. Weigh a clean dry 50-mL Erlenmeyer flask with a rubber stopper. Weighing to the nearest 0.001 g is adequate. Fill the pipet to slightly above the mark, wipe the outside dry and let it drain to the mark. Touch the tip of the pipet to the side of a beaker to remove any hanging drops. Transfer the water from the pipet to the weighed flask. Let it drain freely for 10-15 sec after it is empty. Touch the tip to the side of the flask and replace the stopper. Reweigh the flask and calculate the weight of water delivered by the pipet. Repeat the calibration at least 3 times or more until you obtain agreement within 0.005 g. It is not necessary to dry the inside of the flask if you reweigh the empty/wet flask each time. Convert the mass of water to volume using the data in Table 2-1. These values correct for the density of water at the given temperature, the buoyancy effect, and the expansion of glassware with temperature.
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9 Part B: Calibration of 50 mL Buret In this procedure, we will construct a graph like Figure 2-1 that can be used to convert the measured volume delivered by a buret to the true volume delivered at 20C. Procedure 1. Fill the buret with distilled water and force any air bubbles out the tip. See that the buret drains without leaving drops on the walls. If drops are left, clean the buret with soap and water or have your TA soak it with chromic acid cleaning solution. Adjust the meniscus to be at or slightly below 0.00 mL, and touch the buret tip to a beaker to remove the suspended drop of water. Allow the buret to stand for 5 min while you weigh a 125 mL Erlenmeyer flask fitted with a rubber stopper. (Do not touch the flask with your hands, to avoid changing its mass with fingerprints.) If the level of the liquid in the buret has changed, tighten the stopcock and repeat the procedure. Drain approximately 10 mL of water (at a rate of 99.99% pure), (b) be stable on storage, (c) be easily weighable in ordinary air without picking up water or carbon dioxide, (d) react stoichiometrically with the analyte, (e) be soluble (preferably in water), and (f) have a high equivalent weight to reduce weighing errors. Examples of good primary standards are: potassium dichromate, oxalic acid, potassium hydrogen phthalate, and silver nitrate. Primary standard grade chemicals are expensive and it would be wasteful to use them in preparing large volumes of standard solutions. Instead, such solutions are made directly by the analyst when needed. Secondary Standard Solutions Most titrants are made up to have an approximate concentration, and then standardized by titration against a small amount of a primary standard. For example, a sodium hydroxide solution is usually standardized with pure potassium hydrogen phthalate, and potassium permanganate solutions with sodium oxalate. Once standardized, the solution should be protected from contamination, evaporation, and/or dilution. Preliminary Adjustment of Oxidation State In redox titrations it is frequently necessary to convert the analyte quantitatively to a known oxidation state, either because the sample contains more than one oxidation state, or because the available titrant requires that the analyte be in a higher or lower oxidation state. This is especially true for iron samples which often contain both Fe(II) and Fe(III), and for which either oxidizing or reducing titrants can be used. The preliminary adjustment is made by adding an excess of an appropriate oxidizing or reducing agent. These reagents must (a) react rapidly and quantitatively, (b) not introduce interferences, and (c) have properties such that the excess can be easily removed or rendered harmless. Examples of good preliminary oxidants are: ozone, bromine, sodium bismuthate, perchloric acid, and sodium peroxide. Examples of good preliminary reducing agents: sulfur dioxide, stannous chloride, chromous chloride, amalgamated zinc metal, and hydrazine. Electrochemical methods are also useful. Determination of Iron There are numerous ways to determine iron. The dichromate titration is one of the easiest to perform and introduces some interesting chemistry. The iron in most samples is found in the +3 oxidation state. It could be titrated with a standard reducing titrant to the +2 oxidation state. However, standard solutions of reducing agents are subject to air oxidation on storage, thus less convenient to use than oxidizing titrants. In this experiment we will first reduce the Fe(III) to Fe(II) with a slight excess of stannous chloride. 2Fe3+ + Sn2+ 2Fe2+ + Sn4+
25 The excess SnCl2 is then destroyed with mercuric chloride. Sn2+ + 2Hg2+ + 2Cl- Sn4+ + Hg2Cl2 Finally, the solution is titrated with the standard potassium dichromate solution. 6Fe2+ + Cr2O72- + 14H+ 6Fe3+ + 2Cr3+ + 7H2O Potassium dichromate is a moderately strong oxidant (E = 1.33 V) and is an excellent primary standard. You might think that its orange color would serve to signal the end point after the first excess has been added. However, the product ion, Cr3+, is green and the Fe3+ in the solution may also interfere with seeing the color change. Therefore, an intensely purple-colored redox indicator, sodium diphenylamine sulfonate, must be used. Two titration curves are sketched in the Figure 6-1 below. The dashed curve shows how the theoretical potential would change using the E values for the Fe3+/Fe2+ and Cr2O72-/Cr3+ couples. Obviously the end point (EP) is not sharp and the color change occurs gradually, well before the equivalence point. This would not be a satisfactory titration. Figure 6-1.
The remedy to refining the endpoint location is to load the solution with phosphoric and sulfuric acids. Phosphate forms very strong colorless complexes with both ferric ion and chromic ion. Fe3+ + e- Fe2+ Fe(PO4)x + e- Fe2+ + H3PO4 E = 0.77 V E’ = 0.43 V (in 2 M H3PO4)
26 Thus, Fe(II) is a stronger reducing agent in the presence of phosphate. Also the high concentration of H+ from sulfuric acid increases the oxidizing power of dichromate: Cr2O72- + 14H+ + 6e- 2Cr+3 + 7H2O (excess H+) H3PO4 complex E’ = 1.40 V (in 2 M H2SO4)
These double-barreled effects sharpen the end point and bring it more precisely in range of the indicator change. It is a beautiful example of controlling the conditions to fit the chemistry. The prereduction of Fe(III) is carried out with SnCl2 in hot HCl. The disappearance of the yellow color of Fe3+Clx- complexes indicate when the reduction is complete. The solution is cooled and the small excess of Sn(II) is destroyed by rapid addition of HgCl2. If too large an excess of Sn(II) is added, it will reduce the HgCl2 to finely-divided Hg metal instead of Hg2Cl2. Mercury metal will react with K2Cr2O7 and cause high results. Experimental Procedure: 1. Weigh accurately (to the nearest 0.1 mg) ~1.3 g K2Cr2O7, transfer to a 250-mL volumetric flask, dissolve in water, dilute to the mark, and mix well by inverting and shaking the stoppered flask a few times. Calculate the concentration of this primary standard solution. 2. Obtain your unknown ore sample from your T.A. and dry it for one hour at 110oC. 3. After drying and cooling, weigh accurately (to the nearest 0.1 mg) three 0.6-0.9-g samples of unknown iron ore, and quantitatively transfer to labeled 500-mL Erlenmeyer flasks. 4. Add 20 mL of conc. HCl to each of three labeled 500-mL Erlenmeyer flasks., cover with parafilm, and give them to your TA to store until the next lab period. 5. The remaining should be done during the second lab period. 6. Treat each sample separately through the remaining steps. Do not delay once you have started. 7. Heat the sample nearly to boiling and remove from the hot plate. Some residue may remain (probably SiO2) which can be ignored. Recording the total drops used, slowly add SnCl2 solution dropwise until the solution is almost colorless. Rinse the inside of the flask with a squirt of water, reheat, and continue adding drops of SnCl2 until the color disappears—but not more than 5 drops excess. (If you do add too much SnCl2, add a little dilute KMnO4 solution until the color returns and again decolorize with SnCl2.) 8. Cool immediately by holding the flask in a stream of cold tap water. Add ~100 mL freshly boiled and cooled water. Add rapidly 10 mL saturated HgCl2 solution. CAUTION: HgCl2 is poisonous! A white, silky precipitate of Hg2Cl2 should form. If you get no precipitate within a minute, or a black one, too bad—discard the sample in the appropriate waste container and be more careful next time!
27 9. Allow the solution to stand 3-4 min, then add in order 10 mL 3 M H2SO4, 20 mL 6 M H3PO4 and 5 drops sodium diphenylamine sulfonate indicator solution. 10. Titrate the unknown solution with the K2Cr2O7 solution, swirling constantly, adding dropwise near the endpoint when you begin to see flashes of purple. Stop when you see the first permanent purple color, i.e., the point at which dichromate also oxidizes the diphenylamine sulfonate indicator. This is a sharp color change. 11. Repeat steps 7-10 for each of the replicate samples. 12. There is a significant indicator blank for this titration. To run a blank, all the reagents except for the iron unknown should be added to a 500-mL Erlenmeyer flask and the volume should then be brought up to the approximate volume at the end point of the earlier titrations using distilled water. Hence, 20 mL conc. HCl, 2 drops of the SnCl2 solution (note only the potential excess is added here), ~100 mL freshly boiled and cooled water, 10 mL saturated HgCl2 solution, 10 mL 3 M H2SO4, 20 mL 6 M H3PO4 , and finally dilute with deionized water to the approximate endpoint volume. Once cooled, add 5 drops sodium diphenylamine sulfonate indicator solution and titrate with potassium dichromate solution through the grayish tinge to a permanent violet. Remember the actual end point change is slow. The blank should be no more than 1 to 4 drops (about 0.20 mL max.) and its value needs to be subtracted from the volumes of dichromate standard used in the preceding titrations. The blank should be done twice. 13. Calculate the % Fe in the unknown ore sample. CAUTION: Do not discard solutions containing mercury salts in the sink. Consult your TA for instructions.
28 Experiment 7. Determination of Ca and Mg with EDTA Compleximetric Determinations Most metal ions form stable but soluble complexes with a number of aminocarboxylic acids or their ions. The most important of these complexing agents is ethylenediaminetetraacetic acid (EDTA or H4Y), introduced by Schwarzenbach in 1945. Before beginning the next experiment, the student should review the material on EDTA in the Chem 274 text. Water Hardness Most ground waters throughout the world contain calcium and magnesium as major contaminants. Such waters are known as ―hard water‖ because the metal ions form precipitates with soaps (scum or bathtub ring) and when the water is heated, calcium and magnesium carbonates precipitate and foul boilers, valves, and other machinery. Thus, the determination of calcium and magnesium (―hardness‖) in water is an important procedure. The EDTA titration is the simplest, fastest, and now the most popular way to do this. Reaction Equilibria The equilibria involved in this titration are rather complicated. If you do not understand what is going on in the titration beaker and carefully control the conditions, poor end points will frustrate you. The right combination of conditional formation constants must prevail in order to obtain a sharp endpoint. Calmagite, a triprotic acid here represented as H3In, is used as an indicator. It dissociates the first proton as a strong acid, but the dissociation of the other two protons are as weak acids. H2In(red) HIn2- + H+ (blue) pK1 = 8.1 pK2 = 12.4
HIn2- In3- + H+ (blue) (orange)
Obviously, Calmagite could function as a pH indicator, but that is not desired here. Calmagite also forms a weak complex with calcium (too weak to be useful) and a stronger complex with magnesium. Ca2+ + HIn2- CaIn- + H+ Mg2+ + HIn2- MgIn- + H+ (blue) (red) log Kf’ = 3.7 (pH 10) log Kf’ = 5.7 (pH 10) weak complex stronger complex
In the main titration reaction, EDTA reacts preferentially with calcium, but the end point will be poor because the CaIn- complex is too weak to be seen. Therefore a small amount of magnesium must be present. The magnesium can be added to the EDTA titrant before standardization, thus canceling out its presence. Ca2+ + EDTA CaY2log Kf’ = 10.2 (pH 10) Mg2+ + EDTA MgY2log Kf’ = 8.2 (pH 10)
29 To summarize the several steps of this procedure: Initially, titrant contains EDTA and MgY2- at pH 10 beaker contains Ca2+ and HIn2- at pH 10 During the titration, Y4- + Ca2+ CaY2MgY2- + HIn2- HY3- + MgIn(titrant) (beaker) At the end point, when all Ca2+ has reacted with the titrant MgIn- + excess EDTA MgY2- + HIn(blue) (red)
The color change should be sharp from red to blue. Note that data given for all conditional formation constants are for pH 10, where the relative stabilities increase in the order: CaIn- < MgIn- < MgY2- < CaY2-. The effect of pH is shown in Figure 7-1 below. Figure 7-1.
E D T A titra tio n o f C a
10 8
pCa
2+
a s fc t(p H )
6 4 2 0 0 .0 0 .5 1 .0 1 .5 2 .0 S to ic h io m e tric R a tio (1 :1 b a s is ) pH 10 pH 8 pH 6 pH 4
2+
30 Experimental Procedure: 1. It is important to use good deionized water for this procedure – traces of heavy metal ions will form irreversible complexes with Calmagite and interfere with end points.
2. Weigh 3.0 g Na2H2Y2H2O, transfer to a 1-liter plastic bottle and dissolve in 500 mL water. Add 15 mL 1% MgCl2 solution and 2 mL 6 M NH3. Shake gently and warm if necessary. 3. Weigh accurately (to the nearest 0.1 mg) 0.5 g of previously dried (110oC for one hour) primary standard CaCO3. Transfer to a 125-mL Erlenmeyer flask and add 20 mL water and 5 mL conc. HCl dropwise. Warm the flask gently until the CaCO3 is dissolved, then boil gently for a few minutes to remove all CO2. Transfer to a 100-mL volumetric flask and dilute to the mark. 4. Transfer a 10-mL aliquot of the standard Ca solution from Step 3 to a 250-mL Erlenmeyer flask, add 10 mL pH 10 buffer (32 g NH4Cl + 285 mL conc. NH3 per 500 mL of solution). Add 4 drops of fresh 10% Na2S solution and 5 drops of Calmagite solution. 5. Titrate immediately with EDTA solution (from Step 2) until the color changes from red to blue. If the color appears to fade before the end point, add a few more drops of indicator. 6. Repeat Steps 4-5 two more times. With the data from the three titrations, calculate the EDTA concentration in the titrant. 7. The calcium unknowns will be issued as solutions in a 100-mL volumetric flask which you will have turned in to your TA ahead of time. Add 8 mL conc. HCl. Dilute to the mark and mix well. 8. Transfer a 10-mL aliquot of the unknown solution to a 250-mL Erlenmeyer flask and repeat the titration procedure of Steps 4-5. 9. Repeat Step 8 two more times. 10. Calculate the concentration of Ca in the unknown after it is diluted to 100 mL. Report your result in ppm Ca.
31 Experiment 8. Properties of an Ion-Exchange Resin Introduction: In this experiment, we explore the properties of a cation-exchange resin, which is an organic polymer containing many sulfonic acid groups ( – SO3H). When a cation, such as Na+ or Ca2+, flows into the resin, the cation becomes tightly bound by sulfonate groups, and one H+ is released for each positive charge bound to the resin. The general exchange reaction that takes place with cation-exchange resins is n R – SO3H (s) + Mn+ (aq) (R – SO3)nM(s) + n H+ (aq)
where Mn+ is any metal cation, and R represents the polymeric organic portion of the resin sometimes coated onto inert plastic beads. Schematically the exchange process takes place as follows:
R e sin C u 2+ O S OH O O S OH O O S OH O in O S O O O Cu S O O O S OH O out R e s in 2 H 3O +
Having previous knowledge of the nature of the retained cation, titration of the acidic eluate with a standardized base allows the analyst to determine the amount of cation retained on the resin. Analytical laboratories often use deionized water which is tap water that has been passed through two separate resin columns, one being a cation-exchanger (H+ generating) and the other an anionexchanger (OH- generating), or through a mixed resin bed containing both types of resin. The net effect is the removal of contaminating cations and anions, and the production of non-electrolytic water molecules. Prior to passing such ―dirty‖ water through ion exchange columns the presence of significant ion concentrations imparts a relatively high conductivity to the water source. Removal of such ions by replacement with water molecules (i.e., H+ + OH- H2O, K~1014) reduces the conductivity to much lower levels. When new resin columns are placed into service output conductivities are quite low inferring the generation of ―clean‖ water. In time, output conductivities rise because the capacity of the columns are eventually exceeded and the columns either have to be regenerated or replaced. Cation-exchange resins can be regenerated by treatment with large excesses of H+ or by any excess of any other cation for which the resin might have some affinity. Likewise, anion-exchange resins can be regenerated by treatment with large excesses of OH- or by any excess of any other anion for which the resin might have some affinity. Removal of non-ionic substances from water sources usually requires some other process such as distillation, reverse osmosis, and/or passing the water source through an activated charcoal bed. These procedures are not relevant to the activities of this laboratory experiment. The analytical goal of this experiment is to determine the concentration of positive charge present in a water source, calculated as ppm NaCl, by using ion exchange chromatography. To do this, you will pass an accurately measured volume of an unknown water sample through a prepared column and then titrate the collected eluate with a standardized base solution. The number of millimoles of H+ found in the
32 eluate is directly related to the number of millimoles of positive charge that must have been present in the unknown sample. Reagents 0.02 M NaOH: Each pair of students should have an accurate 1/5 dilution of standard 0.1 M NaOH. This is prepared by standardization against KHP as done in an earlier 274L lab. 1 M HCl Bio-Rad Dowex 50W-X2 (100/200 mesh) cation-exchange resin Unknown saline solution. Experimental Procedure Tips: It is important not to disturb the resin once it has swollen. This can be done by adding liquid towards the walls of the column with an eyedropper instead of directly to the resin. 1. Wash four 0.7-cm diameter x 15-cm length columns with tap water. As a resin retainer, insert a small plug of glass wool towards the column bottoms being careful not to pack too tightly and close with caps. 2. Check whether water flows through the columns without the cap. If no water flows, you packed the glass wool too tightly and need to start over. 3. Check for column leaks with caps attached. 4. Drain the water until 2 cm remains and close the columns again. 5. Prepare a slurry of 1.1 g of Bio-Rad Dowex 50W-X2 (100/200 mesh) cation-exchange resin in 5 mL of water for each column. For size refinement, let settle twice and discard the supernatant. 6. As a fluid-like slurry, slowly pour the resin into the columns at ~30 degree angle to avoid creation of air bubbles within the resin that can affect fluid migration once in operation. 7. Clamp the columns upright and allow the swollen resin to settle for ~15 minutes. As settling occurs you should see a water layer forming above the resin layers. 8. Using an eyedropper, remove the upper water layers without disturbing the resin layers. 9. Trying to minimize resin disturbance, add ~10-mL of 1 M HCl slowly into each column using an eyedropper. 10. Allow the acid to soak the resin beds for ~15 minutes. Some shrinkage might be noticed but that is OK.
33 11. Remove the column caps and allow the acid solutions to drain to the point where the resins are still covered by liquid. 12. Replace the caps and add ~10 mL of deionized water slowly to each column. 13. Allow the resins to soak in the water for ~10 minutes. 14. Remove the column caps and allow the water to drain to the point where the resins are still covered by liquid. 15. Check the pH of one drop of eluate from each column with blue litmus paper. If necessary, repeat the washing process until the eluate is no longer acidic (i.e., litmus turns red). 16. With the columns now closed, accurately and slowly transfer 5.000 mL of unknown solution onto each column. Cover and let soak for ~30 minutes. 17. Place a clean, not necessarily dry, Erlenmeyer flask below each column. 18. Remove the column caps and collect the eluates until the liquid level just covers the resin. Recap each column, add ~10-mL of water to each resin and let soak for ~15 minutes. 19. Remove the column caps and collect the eluate in the same Erlenmeyer flasks. When the water has drained just to the top of the resin, check the pH of the eluate and, if necessary, repeat the washing process until the eluate is no longer acidic. 20. Add 3-4 drops of phenolphthalein and titrate the contents in each Erlenmeyer flask with standardized NaOH until a faint pink endpoint is reached. 21. When done, clean the columns and return back to the zip-lock bag. 22. Assuming the source of the ions in your unknown saline solution is NaCl, report the concentration of NaCl in your unknown solution.
34 Experiment 9. Spectrophotometric Determination of Iron Spectrophotometric Methods The absorption of ultraviolet and visible radiation is the basis of a wide variety of analyses, both organic and inorganic. If the analyte itself does not absorb, it can often be converted to a related substance that does. There are numerous color-producing and color-intensifying reagents available. These methods are rapid, sensitive, and selective. Dilute solutions are preferred and only a small amount of sample is required. The student should review the material on spectrophotometry in the CHEM 274 text before attempting the experimental work of this section. It is important that you understand the use and limitations of Beer’s law of absorption, and the principles/construction of a single-beam spectrophotometer. Operation of the Turner Model SP-830 Spectrophotometer The instrument and schematic diagram of the optical system are shown in Figure 9-1 below. Figure 9.1.
Optical System of a Single Beam Spectrophotometer
35 The fixed-width exit slit determines a nominal spectral bandwidth of ~8 nm at all wavelengths. Study the diagram and locate each control knob and function buttons on the actual instrument. The general directions for operating the spectrophotometer follows: 1. Turn the power on, and allow the spectrophotometer to warm up for at least 20 minutes. Display should come on. The power should remain on at all times during a lab period to avoid having to repeat the war-up period. 2. Set the stray light filter select wheel in the sample compartment to the visible range (~482-736 nm, position #3). 3. Place a cuvette containing a blank solution, i.e., a solution containing all solvents and reagents but no analyte, in the sample compartment and close the compartment cover. 4. Select the desired wavelength (510 nm) with the large knob on the right-hand side of the instrument. 5. Select Absorbance mode by pressing the TRANS/ABS button. The appropriate light will be illuminated under the display, indicating Transmittance or Absorbance mode. 6. Using the blank solution, zero the instrument by pressing the 100% T/0A button. The display will show 100.0 in the transmittance mode (really %T mode), or 0.000 when in the absorbance mode. 7. Carefully open the sample compartment and remove the blank cuvette. 8. Insert a cuvette with the sample (standard or unknown) to be analyzed into the sample compartment. The cuvette should be oriented the same as the cuvette used to zero the blank solution. 9. Being this is an electronic instrument, you can read the absorbance directly and record it. Older analog instruments had absorbance/%transmittance scales that were difficult to reproducibly interpolate in terms of absorbance. More accurate readings were obtained by using the linear %T scale and converting to A = - log T, or A = 2 - log %T. An accuracy/reproducibility of 1% at A=1 (%T= 10%) is the best one can expect with this instrument. Such photometric errors are shown in Figure 9-2 on the next page for four different systems. Clearly, errors are more significant at very low %T (high absorbance) and very high %T (low absorbance) readings.
36 Figure 9.2.
E rrors in C oncentration V alues D erived F rom T ransm ittance M easurem ents as fct(read ing error in % T )
1%
C on cen t ration , % R elative Error in
0 .5 0 %
0 .2 0 %
0 .1 0 %
7 6 5 4 3 2 1 0 0 20 40 %T 60 80 100
Sample Tubes Sample tubes, sometimes referred to as cuvettes, come in many sizes, shapes, materials, and tolerances. In our instruments we will use selected test tubes. If you are lucky, you can select a set of ―matched‖ tubes; that is, they will give the same reading for the same solution. In any event, you should check all tubes to be used for agreement – or correct your readings if necessary. 1. Pick one tube arbitrarily as a reference and half-fill it with a portion of solution #4 described below. Use the index mark near the top of the tube to reproducibly insert the tube in the instrument in the same position. 2. Without a tube in the sample compartment set the zero adjustment to 100%T. 3. Place your reference tube in the sample compartment and obtain a %T reading at a wavelength setting of 510-nm. 4. Replace the first reference tube with another one containing the same solution. If the reading is not the same as the reading obtained in Step 3, try rotating the tube in the sample holder. If necessary, try additional tubes. Reminder: All readings must be made with the sample compartment closed. 5. You will use the two most nearly matched sample tubes in this and the next experiment. Some Iron Complexation Chemistry Both Fe2+ and Fe3+ ions are colorless unless complexed. Fe3+ usually appears yellowish because the [Fe(H2O)6]3+ ion partially hydrolyzes to [Fe(H2O)5OH]2+. However, there are many complexing agents that form colored ions with either Fe2+ or Fe3+. In this experiment we will use bipyridine, bipy, which forms an intense red complex with Fe2+.
37 Fe2+ + 3 bipy [Fe(bipy)3]2+
2,2’ bipyridine Iron(III), if present, must be reduced to iron(II). The complex forms rapidly and is stable to oxidation over the pH range 3 to 9. It is useful in determining iron concentrations in the range of 0.5 to 10 ppm. The low concentration range requires several successive dilution steps that can complicate the final calculations. Since you will be using many volumetric flasks, be sure to mark them well, or you will end up in total confusion. Draw a flow chart in your notebook to help you keep track of the dilutions. Experimental Procedure: 1. Prepare a standard iron solution. Weigh accurately enough FeSO4(NH4)2SO46H2O to make 250.0 mL of 0.002000 M Fe(II) solution. Transfer the salt to a 250-mL volumetric flask, add water to dissolve and 8 mL 3 M H2SO4. Dilute to the mark with de-ionized water and mix well. 2. Using a pipet, transfer a 10-mL aliquot of the 0.002000 M solution from Step 1 into a 100-mL volumetric flask, add 4 mL 3 M H2SO4, dilute to the mark, and mix. This is a more dilute standard solution of approximately 0.0002000 M. The concentrations should be known to within about 0.5%. 3. Add 4 mL of 3 M H2SO4 to the unknown, dilute to the mark with de-ionized water, and mix well. 4. Using a pipet properly cleaned and rinsed with the solution from step 3, transfer a 10-mL aliquot of the unknown solution of step 3 into a 100-mL volumetric flask. Add 4-mL of 3 M H2SO4, dilute it to the mark, and mix well. 5. Obtain seven 50-mL volumetric flasks; number them 1 through 7. Add the following to each flask: #1 – 0 mL of standard Fe2+ solution (blank) #2 – 5 mL of standard Fe2+ solution from step 2 #3 – 10 mL of standard Fe2+ solution from step 2 #4 – 15 mL of standard Fe2+ solution from step 2 #5 – 10 mL of unknown Fe2+ solution from step 4 #6 – 10 mL of unknown Fe2+ solution from step 4 #7 – 10 mL of unknown Fe2+ solution from step 4 6. To each of the volumetric flasks add the following: 1 mL of 10% hydroxylamine hydrochloride 10 mL of 10% sodium acetate 10 mL of 0.1% bipyridine Dilute each flask to the mark, and mix well.
38 7. Take a complete visible spectrum (400 to 700 nm) of the solution in flask #3, using the blank solution (#1) as a reference (i.e., for setting 100% T). Measure the absorbance at 10-nm intervals. Select the wavelength of maximum absorbance and use this wavelength in Step 8. 8. Measure the absorbance of each of the solutions prepared in Steps 5-6 at the wavelength of maximum absorbance. 9. Plot absorbance versus wavelength for the solution in flask #3. Plot absorbance versus [Fe2+] for the standard iron solutions (#2, 3, and 4). Calculate the concentration of Fe2+ in the original unknown sample. 10. Keep the sample tubes in your drawer until you finish the next unknown experiment (Spectrophotometric Determination of P).
39 Experiment 10. Spectrophotometric Determination of Inorganic Phosphorus Principle The inorganic phosphorus in a protein-free filtrate is reacted with ammonium molybdate [Mo(VI)] to form ammonium phosphomolybdate. This is reduced with a mild reducing agent to produce ―molybdenum blue,‖ a heteropoly molybdenum(V) species. Molybdates are not further reduced under these conditions and the blue color of the solution is measured spectrophotometrically. Chemical Equations 7PO43- + 12(NH4)6Mo7O24 + 36H2O 7(NH4)3PO4 12MoO3 + 51NH4+ + 72OH(NH4)3PO4 12MoO3 + mild reducing agent Mo(V) species (blue) [Although normal ammonium molybdate (NH4)2MoO4 can be crystallized, the common crystalline form is (NH4)6Mo7O24 4H2O or 3(NH4)2O 7MoO3 4H2O.] Reagents You are provided with: Aminonaphtholsulfonic acid reducing agent Solution of unknown phosphorus conc. (issued in 100-mL volumetric flasks) You have to prepare: Phosphorus standards of 100, 10, 5 and 2 mg/L P Ammonium molybdate solution Experimental Procedures 1. Dissolve 0.439 g KH2PO4 in water and dilute to 1000 mL in a volumetric flask with DI water. This stock solution contains 100 mg/L P. Pipet 10 mL of the stock solution (from step 1) into a 100-mL volumetric flask and dilute to the mark with DI water, resulting in a working standard which contains 10 mg/L P. Pipet 25 mL of the 10-mg/L P solution (from step 2) into a 50-mL volumetric flask and dilute to mark with DI water, resulting in a working standard that contains 5 mg/L P. Repeat step 3 with 10 mL of the 10-mg L P solution (step 2) into a 50-mL volumetric flask and dilute to mark with DI water, resulting in a working standard that contains 2 mg/L P.
2.
3.
4.
You now have three working standards: 10, 5 and 2 mg/L P solutions. 5. Dilute your unknown to the mark with DI water.
40 6. Pipet 10 mL of your unknown (from step 5) into a 50-mL volumetric flask and dilute to the mark with DI water. Repeat this step to prepare two more solutions (three unknown replicate samples needed). (*) Pipet 25 mL of DI water (to serve as the blank) and the three working standards into four small beakers (or Erlenmeyer flasks). Add the following accurately to each of these containers: 5 mL molybdate solution (**) 2 mL reducing agent Obtain a calibration curve of Absorbance versus [P] by measuring the absorbance of the four solutions in step 7 at 690 nm (filter wheel position #3), using the blank solution for setting 100% T. Repeat step 7 with your three unknown solutions and measure their absorbances. Report your results in mg/L P.
7.
8.
9.
* Be sure to complete steps 7-9 during one class period as the solutions will air oxidize over long periods. ** Ammonium molybdate solution. Dissolve 2.5 g ammonium molybdate. (NH4)6Mo7O244H2O, in 80 mL water and add 30 mL of 5 M H2SO4. The solution should be stable indefinitely. Discard if blanks show a blue color.