CHEMISTRY 247 QUANTITATIVE CHEMICAL ANALYSIS LABORATORY … · Free silver (black stain): Use conc....

Illinois Institute of Technology Chemistry 247 Laboratory Manual, Edition: January 2003

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ILLINOIS INSTITUTE OF TECHNOLOGY

CHEMISTRY 247 QUANTITATIVE CHEMICAL ANALYSIS

LABORATORY MANUAL

Revised by William R. Penrose January, 2003

to conform to the text

D.C. Harris Quantitative Chemical Analysis, Fifth Edition

IN CASE OF EMERGENCY From Campus Phone, call 8-6363 for Public Safety Call 9+911 for Chicago Emergency Services From cell or other phones, call (312)-808-6363 for Public Safety Call 911 for Chicago Emergency Services Be prepared to give location of the laboratory: Wishnik Hall, Corner of 33rd and State Streets, Second Floor, South end.


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TABLE OF CONTENTS General Instructions Including SAFETY – READ THIS SECTION!! Labeling Chemicals Removing Chemical Stains Common Errors in Technique Laboratory Reports Use of the Balance Polymer Body, Sealed Reference Combination Electrode Use of the Spectronic 20 Visible Spectrophotometer Expt. 1 - Safety – Checkin - Calibration Of A Buret Expt. 2 - Titrimetric Determination Of Khp Expt. 3 - Titration Curve Of A Weak Monoprotic Acid Expt. 4 - Titration Curve Of Polyprotic Acids Expt. 5 - Gravimetric Chloride Expt. 6 - First Nitrite Lab: Calibration Expt. 7 - Second Nitrite Lab: Complex Sample Expt. 8 - Chloride By Ion-Selective Electrode Expt. 9 - Gas Chromatographic Separation Of An Alkane Mixture Expt. 10 - Experimental Design (Ascorbic Acid)


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INTRODUCTION

GENERAL INSTRUCTIONS Chapters 1 and 2 of D.C. Harris, Fifth Edition, provide an introduction to general techniques and common equipment used in analytical chemistry. Specific readings will be assigned with each experiment.

The first five sections in this laboratory manual give specific instructions as to the proper protocol, safety precautions, and conduct expected in this course. If these are followed, everyone will benefit from a safer, cleaner, more productive, and even enjoyable learning experience. Please read the following sections carefully. PLANNING To save time in the laboratory, plan ahead for each experiment. For example, drying and cooling of samples usually takes over an hour, and other tasks such as reagent preparation or cleaning can be done concurrently. NOTEBOOK AND LABORATORY REPORT POLICIES To record experimental data use the standard bound laboratory notebook available at the lIT bookstore (soft cover with carbon pages, National No. 43-647 or 43-469). Write your name on the front cover and along the side page edge. Leave 2-3 pages at the beginning for a table of contents. Enter all of your laboratory work and related activities directly at the time of performance. Identify all entries by name of experiment and by date. Keep the table of contents up-ta-date. Begin each experiment with a written section that was completed before coming to the lab. This section must contain an introduction which includes the objectives of the experiment. List the chemical reactions and equations involved as well as molecular and equivalent weights. The introduction must contain the procedure that will be followed. It is often helpful to rearrange instructions published in the text syllabus into a flow-chart or ordered program. Include personal notes on how to effectively save time. In the beginning of the semester, most instructions will be explicit. However later experiments will give general directions only and may require preliminary calculations of amounts, quantities, proportions, etc., in this section. The general ground rule for this section is simple. You must be able to perform the lab experiment by reference only to what has been written in your lab notebook. If the section has not been completed prior to the lab period, you will not be allowed to proceed and will not receive any credit for the experiment. Preparing such a write-up will improve your understanding and increase the probability that you will perform a successful experiment.

Further instructions regarding the content and format for lab reports are in Section 5. Sample lab reports are available on request.


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SAFETY, HOUSEKEEPING, AND SOME COMMON LABORATORY ERRORS. All students must behave in a mature and responsible manner. Careless behavior could result in serious injury to yourself, others, and to property. See Sections 2, 3 and 4 of this manual regarding safety, housekeeping and common lab errors. SAFETY GLASSES Safety glasses, or prescription glasses with safety lenses, must be worn at all times in the laboratory to protect your eyes from injury. Those who do not have approved glasses will not be allowed to work in the laboratory. FIRST AID In cases of severe injury, call IIT Public Safety Immediately or Chicago emergency services. From a cell phone, IIT is 312-808-6363, Chicago is 911. Report the location as Wishnick Hall, 33rd and State Streets. Someone should wait outside to guide the paramedics, fire or police to the scene of the event. Water is to be used first on all chemical burns. Immediate flooding with water of the affected area is more effective than all the specific remedies that can be applied later. A first aid kit is located near the main door and a larger first aid kit is kept in the storeroom. Always consult the instructor before using the first aid kits. Burns. Surfacaine is recommended for small burns. Larger burns or burns producing immediate blistering should be treated by the school nurse or doctor or at Mercy Hospital. In hospital cases no preliminary treatment should be given, if possible. Bromine Burns. Avoid the use of bromine when possible. Bromine burns are treated by immediately flooding with water, followed by a liberal application of glycerine. Hospital examination is necessary within the hour. If a serious burn is suspected, medical treatment is advised. Many chemical burns develop slowly. Spilled Acids or bases. For major emergencies that involve acids or bases spilled on your body or clothes, proceed to the emergency shower and pull the cord. DO NOT hesitate. Burns may be delayed, especially with strong alkalis. Small acid spills on clothes should be immediately sponged with dilute (approx. 1:3) ammonium hydroxide. Even the vapors will neutralize areas not directly sponged. Consult the instructor or TA for benchtop or floor spills. These spills may require neutralization before cleanup. Treatment All injuries requiring treatment should be reported immediately to the instructor or TA. Arrangements for emergency treatment by the school nurse or doctor or for transportation to Mercy Hospital are made through an instructor or the Chemistry Department office.

Reporting Injuries

All injuries, whether requiring medical treatment or not, must be reported to the instructor. Students


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working in the laboratory are insured by the school, but adequate records must be kept. FIRE Shout “FIRE” several times if a fire occurs, to alert people near you. A dry powder fire extinguisher is located in each corner of the laboratory. Familiarize yourself with their location and method of operation. A dry powder extinguisher is good for all types of fires. Remove the locking pin, squeeze the handle, and direct the blast of powder at the BASE of the fire. DO NOT ATTEMPT TO FIGHT A FIRE IF: YOU ARE ALONE THERE ARE OPEN OR UNOPENED BOTTLES OF CHEMICALS OR SOLVENTS IN THE

FIRE. YOU CANNOT SEE THE ENTIRE FIRE THE FIRE IS BETWEEN YOU AND THE DOOR THE FIRE IS MORE THAN THREE FEET IN DIAMETER YOU ARE UNSURE OF YOUR ABILITY OF CONTROL IT A fire blanket is located just below the light switches near the main door of the laboratory. An emergency shower is located just inside the main door. The large ring can be easily located in an emergency, if you observe and remember its location. Either the shower or the fire blanket can be used when a person's clothes have caught fire. The shower can also be used for accidents with acids, alkalies, or other dangerous chemicals where the coverage is too great to be treated at the sink. (If the shower is used for an emergency, others should immediately begin mopping water to prevent damage to scientific instruments on the first floor.) A fire alarm box which sets off the alarm in the building only is located above the light switches in the laboratory. To call the fire department, the Institute switchboard should be notified by dialing 8-6363. If the switchboard is closed, dia1 911 on the public phone located on the first floor at the foot of the south stairs. 911 can be reached from any Institute phone after dialing "9" and waiting for the second dial tone.

Exits. In case of fire, exit Room 208 in an orderly fashion down the stairs to the front of Wishnick Hall. If this area is not available, proceed out the east door to the area in front of Perlstein Hall. If the main door of the laboratory is blocked, the second exit through the storeroom and freshman lab, should be used. If the glass doors are locked, break them from a safe distance.


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LABELING CHEMICALS

If you encounter unlabeled chemicals, assume the worst: Assume they are poisonous, volatile, and flammable. Failure to label a dangerous chemical or to treat the chemical with caution is a major source of laboratory accidents AND a cause for dismissal at many companies. Strong acid and pure water look exactly alike. Use a marker or a gummed label to identify all your solutions, reagents, etc. This is an excellent habit to cultivate for all chemicals, even if they are not dangerous. Many items of glassware have etched spots which can be written on with a pencil and erased with an eraser. This provides a permanent marking which is easy to remove.

All objects placed in the oven or in common areas must labeled with the name of the chemical and your initials. Beakers should be covered with watch glasses. Only air dry objects may be placed in the ovens. REMOVING CHEMICAL STAINS Do all of these operations in the hood. Almost all involve noxious gases.

Silver chloride ( in Gooch crucibles): Use conc. ammonium hydroxide followed by water, in the hood.

Free silver (black stain): Use conc. HNO3 followed by water, in the hood. In this case aqua regia is ineffective since silver chloride is precipitated. A mixture of silver chloride and free silver: Dissolve free silver in conc. HNO3, wash thoroughly with water, then dissolve the residue in conc. NH4OH and wash with water. Repeat the cycle if necessary. Potassium permanganate stains - MnO2 (violet or brown): Dissolve in an acidic solution of sodium sulfite (Na2S03 or NaHSO3) acidified with dilute sulfuric acid. Other acidic reducing agents can be used. Concentrated hydrochloric acid can be used, but chlorine is evolved. Work in the hood. Rust or iron: Digestion with conc. HCI. Iron stains in porcelain crucibles usually cannot be removed. If the glaze is still smooth, the stain usually does no harm. Silicone type greases: These are occasionally used for stopcock greases. Because of the tenacity with which they adhere to glass, they are never to be used in the lab. Glassware treated with silicone type compounds (e.g. "Drifilm") is not wet by water and will show a flat or even convex meniscus. Prolonged action of CHCl3 or other chlorinated solvents is the usual method of removal Iodine stains on hands or clothing: sodium thiosulfate solution. Not usually effective. Dust and dirt in sintered glass crucibles that have been left on suction for long periods of time is almost impossible to remove completely. Mechanical removal with soap, brush, a stream of water, and backwashing are most effective. Concentrated HNO3, a HNO3-H2S04 mixture, or warm cleaning solution will be moderately effective. Residues that withstand such treatment will usually come to constant weight and do no harm in an analysis.


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COMPUTING THE ERROR IN A STANDARDIZATION (E.G., LAB 2) Suppose you have accurately weighed a 0.7 g sample of potassium hydrogen phthalate (KHP), and its weight (by amazing coincidence) is exactly 0.7000 +/- 0.0002 g. The error shown is an estimate based on the stability of the balance while you are weighing the salt, and the known four-decimal-place accuracy of the balance. The formula weight of KHP is 204.233. Because there is just one dissociable proton, the equivalent wt is also 204.233. In titrating the sample with a nominal 0.1 N NaOH, the initial buret reading is 0.05 +/- 0.02 mL and the final reading is 40.05 +/- 0.02. The total titrant is therefore 40.00 mL. The no. of equivalents of KHP is 0.7000/204.233 = 3.427 x 10^-3 eq. For convenience, this is expressed at 3.427 milliequivalents, or meq. When titrating, eq = N x V or meq = N x V (mL) The normality N = 3.427/40.00 = 0.085675 N (keep all the significant figures until the end). Error: The volume error is the propagated error of two buret readings: Error = sqrt ( err1^2 + err2^2) = sqrt ( 0.02^2 + 0.02^2) = 0.028 (note: 2 or 3 sig figs is enough in error calculations, because you are only going to keep one at the end. Volume is therefore 40.00 +/- 0.028 mL. When combining with the weighing error, a division operation is used, so the proportional errors are used. Volume %error = 0.028/40.00 x 100% = 0.070% Weight %error = 0.0002/0.7000 x 100% = 0.0286 % NOTE THAT, ALTHOUGH THE WEIGHT WAS DIVIDED BY THE EQUIVALENT WEIGHT OF KHP, THERE IS NO ERROR PROPAGATION SINCE THE EQUIVALENT WEIGHT IS CONSIDERED TO BE EXACT. THE PROPORTIONAL ERROR THEREFORE IS THE SAME FOR THE NO OF MEQ OF KHP AS FOR THE ACTUAL WEIGHT. Combining proportional errors: % error = sqrt ( (0.028)^2 + (0.070)^2 ) = sqrt (0.005716) = 0.0755% The absolute error is 0.0755% of 0.085675 N = 0.00065. The concentration of the standard should be 0.08568 +/- 0.00007 N ! If replicate determinations are made, and the results reasonably close, the propagated errors of each will not be much different, so this calculation is normally done once. To confirm this, the mean and standard deviation of at least three replicate titrations should be calculated (a posteriori error) and compared with the propagated (a priori) error.


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COMMON ERRORS IN LABORATORY TECHNIOUE Good technique does not consist of simply not making mistakes - it requires the use of the best techniques you know at all times.

The following list contains some of the more common errors. Most Serious Errors 1. Improper use and care of the balance: e.g., using or leaving a dirty balance, forgetting to tare, or misreading the weights. Spilled chemicals on the balance pan will be weighed along with those in the weigh bottle. 2. Improper note-taking procedure: e.g., notebook record not made at time of observation; data recorded anywhere but in the correct place in the notebook; use of both a "good" and a "rough" notebook; notes taken in pencil (borrow a pen if necessary); corrections made by any method other than drawing a line through the unwanted result; pages torn from notebook

3. Returning chemicals to bottles.

4. Any act endangering others in the laboratory.

5. Leaving spilled water on the floor.

Errors Which Contribute to Poor Results

1. Using dirty volumetric equipment.

2. Banging a stirring rod against the side of a beaker (producing chips that will affect the weight of a precipitate).

3. Spilling chemicals or not cleaning up spilled reagents.

4. Using tubing or rod without fire-polishing .

5. Beakers and Erlenmeyer flasks left uncovered, especially when on the hot plate or in the ovens. 6. "Drying" solids in deep layers. Always turn and redry the sample for a second period.

7. Weighing samples before they have cooled to room temperature.

8. Losing sample by boiling solutions.


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9. Using reagents, solid or liquid, without checking name and quality (e.g., confusing sulfates and

sulfites). The solids on the shelf may be a different hydrate than that called for and a different weight will

be required. Technical grade reagents occasionally get on the shelves by mistake. A liquid may contain a

sediment, particularly when solvents are refilled from barrels or drums. In all cases, the responsibility of

observation lies with the user.

Discard weighing papers after use in trash receptacles (not in sink).

LABORATORY REPORTS ALL LABORATORY WRITEUPS MUST BE SUBMITTED WITH A PREPRINTED COVER PAGE WHICH SUMMARIZES YOUR RESULTS. The writeup of an analytical chemistry experiment can be divided into 3 phases, preparation (pre-lab), experimental. and analysis (post-lab). Pre-Lab: Title and Your Name (use both experiment number and name)

Objective of the experiment Background and introduction Procedure and materials (some students have used flowchart and checklist formats)

Experimental: Pertinent data (weights, unknown numbers, optical absorbances, times, etc. Deviations from the prescribed procedure, and reasons for these Observations of unexpected events Potential sources of error List shared data and sources of shared data Clean your equipment and work area. Post-Lab: Calculate the results asked by the lab instructions Prepare any required plots and graphs, either computer or neatly hand-drawn Explain your results Explain sources of error and any mistakes that were made in the procedure Fill in the prepared cover sheet. Staple or clip cover sheet, carbon sheets from notebook, and copies of graphical data together and submit on the required date. The Pre-Lab sections of the report must be entered in the lab notebook before the lab work begins. You will not be allowed into the lab without these sections completed. The content of preliminary sections has been discussed in section 1.5; sample lab reports should also be consulted. Introduction: This section (about 1/2 page) will contain a brief description of the experimental technique to be used. The appropriate principles of physical chemistry should be introduced, and the equations used to relate the data to the results derived. Balanced chemical equations should be used throughout. Indicate the limits


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of the technique and the special precautions needed to obtain good results in the experiment. Procedure: Present, in summary, outline, or flow chart form, all the information you need to perform the experiment in the lab. You must be able to do the experiment on the basis of what is written here. Include approximate quantities of reagents to be used, the temperature and duration of dryings required for samples and unknowns, comments on endpoints and tests of reagent purity. For example, if the quantity of a particular reagent must be calculated on the basis of the weight and approximate composition of the unknown, then the calculation must be shown at the proper place in the procedure. .

Data Pages:

An integral part of the procedure is establishing the sequence of measurements and their precision. The layout of the data page should reflect this chronological order. All data and measurements must be recorded here. Never record a measurement, even preliminary ones, on scrap paper. Neatness on the data page is not as important as obtaining a complete record. Calculations: The calculations section should show a sample calculation that contains the sequence of steps used in reducing the measurements to find the required results. Show all units and how they cancel to give the final answer. A clear, informative calculation section is the best defense against turning an accurate experiment into a reported failure. The sample calculation should show real data from one trial in the experiment. Results: Present the results of the experiment in tabular or graphic form. For a report on a gravimetric experiment, for example, include in the table the weight of each unknown sample, the weight of AgCl ppt, the calculated %Cl, the mean value, and the standard deviation. Clearly label all table entries with appropriate headings. For experiments involving multiple determinations, it is permissible to reject a single determination if it is inconsistent with other results. You can reject a trial only if you have sufficient other data to show that the suspect trial is indeed atypical. For example, suppose you have standardized a NaOH solution five times to give the results as normality (N):

#1 0.1009 #2 0.0875 #3 0.1021 #4 0.1011 #5 0.1018

Trial #2 could be legitimately excluded from a calculation of the average normality using a ‘Q’ test. If so, you must clearly label this as a bad run and, if you can, explain why it may not have been valid. Average results as well as standard deviations for an unknown must be calculated and displayed clearly. Be sure to report values for unknowns in the correct units. The grader will not recalculate or change units for you. Discussion: This section has two parts. First, critically discuss the experiment in terms of the inherent strengths and weaknesses of the techniques used. Second, make a careful and quantitative analysis of all


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errors. Part 1. Carefully examine the experiment for its unavoidable or theoretical limitations. For example, consider a gravimetric precipitation experiment. Formation and analysis of a precipitate assumes that the salt is "insoluble". What if it has some finite solubility? How would that affect results? Make them too high or low? How might you correct for this?

Many of these limitations have been discussed in your text. Occasionally your discussion will repeat some of the same points; strive to understand these points and discuss them in your own words. The text does not necessarily cover all points, so see if you can contribute original criticism.

Next you may comment on blunders that may have occurred. "Trial #4 is too low because half the solution was spilled. Weights for run #2 are too high because there was insufficient time to dry the soaking wet precipitate. No endpoint was found for titration number 7 because the indicator was omitted." Avoid using personal pronouns in this discussion. Experiments are graded on results, and blunders will cost you in terms of your grade. However, you can salvage some credit and prove that you understood the experiment by mentioning blunders that occurred. Part 2. Here discuss the quantitative analysis of errors. Most of the errors to be considered are indeterminate errors associated with many quantitative measurement made during the course of the experiment, including weights, volumes, pH meter readings, absorbance readings, etc. You must make a reasonable estimate of all indeterminate errors. For example, suppose you analytically weigh a sample of a salt and obtain 0.4869 g. The balance instrument has uncertainty in the last decimal point, so the "true" weight might be 0.4870 or 0.4868 g or 0.4968 +/- 0.0001 g. You will see by movement or drift of the scale that the error might be even larger. Many of your indeterminate error estimates can come from your analysis of the "readability" of an instrument. However, the error estimate should be larger because of chemical considerations. For example, suppose you are weighing a precipitate and you note that the weight is drifting high by a few mg per minute. This could be the result of water adsorbed from the air. Since you don't know how much water was adsorbed, this is an indeterminate error. You might estimate that your weight was correct to +/- 0.002 g. Or you might consider a pH meter reading. The scale gives numbers to +/- 0.01 units, but you notice a drift, perhaps due to CO2 adsorption, temperature changes, or electronic "noise". Therefore, you might choose to revise your error estimate to +/- 0.1 units. After all errors have been estimated, you must show how they can affect your results by a propagation of errors analysis (Harris, Ch. 3). Often, it will be possible to see that one or a few sources of error are more important than all others. This should be mentioned. This error analysis will provide an expected amount of error for your results, considering all the error sources. Now compare this to standard deviations obtained statistically from your final data. Be careful in the comparison of units. Propagation of errors gives a relative (fractional) error estimate, while calculated standard deviations are usually in absolute units. Comment on observed standard deviations that are much larger than the expected error.

Some experiments will require that you answer some direct discussion questions. These must be answered as part of the lab report. They are intended to aid you in a thoughtful evaluation of the experiment.


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USE OF THE BALANCE Older weighing instruments were slow and required great patience to use. The balances in the lab are faster "single pan" balances. The ease of operation may disguise the fact that it is a delicate (and expensive) instrument and should, therefore, be used with great care.

Never place a chemical directly on the balance pan. Weighing bottles, weighing papers, plastic weigh boats, dry beakers, or watch glasses must always be used. The weight of container and sample together should not exceed the maximum load of the balance. Avoid touching any interior part of the balance. All analytical weighings should be done by difference, as described in the textbook. One exception is for volatile liquids, which are difficult to quantitatively remove from a bottle. In this case, you should use a pipet or dropper to add liquid to a tared (weighed empty) weigh bottle, and the glass cover replaced. After each addition, reweigh the vessel until the desired weight is attained. . Don't use the analytical balances unless they are required. Approximate chemical weights can be obtained with the coarse top-loading balances or triple beam balances. No chemical should touch the pan. A recurrent problem for beginning students is misinterpretation of weighing instructions. “Accurately weigh 0.5 g of X” means that you should tare the weigh bottle to four decimal places, add X to the weigh bottle on the analytical balance until the net weight is between 0.45 and 0.55 g, then read the accurate weight to four decimal places. “Add 2 g of X to your sample” means to weigh 1.9 to 2.1 g of X on the top-loading balance, and add to the sample without further reading.


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USE OF THE POLYMER BODY, SEALED REFERENCE, COMBINATION pH ELECTRODE Polymer-body, combination electrodes are easy to. Because the pH bulb is recessed inside the polymer body, the electrode can rest against the bottom of a beaker without damaging the glass bulb. In many measurements this recessed bulb design eliminates the need for electrode holders. The sealed reference design eliminates the need to add filling solutions, minimizes reference dryout, and allows the electrode to be used in systems up to 100 psi without the need for external pressurization. Helpful Operating Techniques The electrode tip is covered by a protective cap which serves both to keep the reference from drying and to prevent breakage. This cap fits snugly and it contains a pressure relief hole to facilitate removal and installation. This hole is covered by vinyl tape to retain moisture inside the cap. Before removing or reinstalling this cap, the tape must be removed to expose the pressure relief hole. During shipment the air pocket in the electrodes stem may move into the bulb area. If bubbles are seen in the bulb area, hold the electrode by its cap and shake downwards as is done with a clinical thermometer. Vigorously stir the electrode in the sample, buffer or rinse solution. This action brings solution to the electrodes surface quickly and improves the speed of response.

After exposure to a sample, buffer or rinse solution, shake the electrode with a snap motion to remove residual drops of solution thus minimizing contamination from carryover. Blot the electrode with tissue paper, but do not rub. Rubbing produces static electricity. As a rinse solution, use a part of the next sample or buffer to be measured. This action also reduces contamination from carryover. When calibrating, use a buffer close to that expected from the sample to control span errors. Keep buffers and samples at the same temperature to eliminate the need to correct values for temperature effects. pH readings stabilize faster in some solutions that others; allow time for the reading to stabilize. In general, buffers provide stable readings in several seconds (tris buffers take somewhat longer) while samples usually take longer. Keep in mind that all pH electrodes 'age' with time. Aging is characterized by shortened span and slower response speed. If the pH meter has a 'slope' (span) control, the control can be adjusted to compensate for electrode span errors but will not affect speed of response. Aging is best detected by calibrating the electrode in, for example, pH 7 buffer and then rinsing and placing the electrode in pH 4.01 buffer. As a rule, if the span is 10% or more in error (a reading of 4.3 or higher for this example) the electrode should be cleaned and retested. If reconditioning does not restore performance the electrode should be replaced.


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USE OF THE ‘SPECTRONIC 20’ VISIBLE SPECTROPHOTOMETER Theory of Spectroscopy When monochromatic light of a particular wavelength A and initial intensity I is passed through an absorbing substance, the exiting light is decreased to an intensity I, according to a relationship known as BEER's LAW:

I = I. 10 -aBC

where C is the concentration of the absorbing species in a light path length B, and a is a constant known as the absorption coeffient of the particular species. a is a function only of the species and the wavelength,

not of the species' concentration. Note that the factor 10 -aBC

is always less than one for positive a, B, C. The ratio I / I is called the transmittance (T) of the sample, and is sometimes expressed as a percentage. The absorbance (A) is, by definition, A = - log T = - log (I / I) The absorbance is sometimes called optical density. The convenience of this definition is clear when T and A are expressed in terms of equation (1):

A = - log T = - log (I / I) = -log10 (10 -aBC

), or A = aBC (3) In other words, the absorbance of a sample at a particular wavelength (and thus constant a) and for a particular path length B, varies directly as its concentration C. Study the above relationships until you are satisfied that decreasing T means increasing A, that T = 1 means A = 0, that T decreases with increasing C while A increases with C. As stated above, the absorption coefficient varies with the wavelength of light used. A complete plot of a versus wavelength is known as the absorption spectrum of the species studied. The Bausch and Lomb Colorimeter-Spectrophotometer (Similar instructions apply to the Spectronic 20D) A narrow range of wavelengths is selected from the continuum emitted by the tungsten lamp source by appropriate orientation of a diffraction grating. The grating is turned mechanically by an arm and earn that is joined to a knob on the instrument top. The knob simultaneously turns a precalibrated wavelength scale from which wavelength readings are obtained. There may be some error in this scale. After light passes through the sample, it strikes a photocell through which electronic amplification of the photocurrent causes a meter to deflect to a degree proportional to the intensity of the transmitted light. The needle scale is calibrated in % T (the linear scale) and optical density (A; the logarithmic scale). Digital models have a readout that can be in units of T or of A. In these labs, we recommend that transmittance readings be taken and converted to absorbance afterward.


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Figure 1. Schematic Diagram of Spectrophotometers, showing both analog and digital readouts.


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When the sample tube is removed from the sample compartment, a mechanical shutter falls into the path of the light beam, so that no light reaches the photocell. The left hand knob is then adjusted until the meter reads 0% T or infinite absorbance. When a tube containing deionized water, or a reagent blank, is placed in the sample compartment, the needle should be adjusted to read 100% T or A = 0.000 by turning the right front knob which adjusts the beam size incident on the test tube and photocell (see light control in Figure 1). Finally, with no change in instrument settings, the sample to be measured is inserted into the light path and its absorbance read directly. Operating Instructions 1. Turn on the spectrometer. Allow the spectrometer to warm up for about 20 minutes. 2. Set the wavelength selector to the desired wavelength (see Figure 1). With digital instruments, press the MODE button until the ‘T’ light is on, meaning that the display will show transmittance. 3. With the cell compartment empty and the lid closed, adjust the left knob until the meter reads 0% T. Watch it carefully for drift. The instrument may not yet be fully stable. 4. Place a cuvette (a preselected test tube sold for the purpose) filled with deionized water in the cell compartment. Align the cuvette so that the vertical mark is toward you. Close the cover, and adjust the right front knob until T is 100%. Depending on the experiment, a blank solution may be used in place of water. 5. When using the Spec 20 for the first time that day, repeat steps 3 and 4, making fine adjustments if necessary, then go on to step 5. 5. Replace the standard with sample, align the marks close the lid and read absorbance with no knob adjustments. Record the transmittance or absorbance. 6. Rinse the cuvettes sample tubes thoroughly and allow to drain. DO NOT SCRUB TUBES WITH A TEST TUBE BRUSH. It spoils the calibration. Return sample tubes to instructor. EXPERIMENTS Instructions for individual experiments are given in the following sections. Remember, the key to a successful experiment is a thorough understanding of the principle as well as the meaning and purpose of each step in the procedure.


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LABORATORY 1

CALIBRATION OF A BURET Chem 247 - Analytical Chemistry Although many laboratory supplies and items come from the manufacturer with calibration marks on them, it is not wise to trust these marks, particularly when high analytical accuracy is necessary. In this course, you are expected to calibrate some of your glassware. The buret is a device used to deliver precise volumes of a reagent. You will find in this class that it is also a delicate instrument that requires as much care skill in its use and storage as the most expensive analytical instrumentation.

We will calibrate our burets using a gravimetric standard based on the density of water. Keep in mind that this exercise assumes that the balances in the laboratory have themselves been properly calibrated! Preparation Read Chapter 2 in your textbook on use and care of glassware, and of making corrections for calibrations. Look up “buret, calibration” in the index. You may want to check out the following web sites: http://www.dartmouth.edu/~chemlab/techniques/buret.html http://www.cbu.edu/~mcondren/c214/titration/titraon.htm http://onsager.bd.psu.edu/%7Espudich/Quant.html Calibration and (Important!) a graph of density of water corrected for expansion of glass buret: http://chemistry.csudh.edu/oliver/che230/labmanual/calbur.htm

Buret stand or retort stand

Make sure stand grips buret firmly

Stopcock of Teflon (white) or glass (clear)

Waste flask in position whenever buret is unused

Check that tip is not chipped or broken


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Standards used by Kodak for analytical chemistry work: http://www.kodak.com/US/en/motion/support/processing/h243/h2403ulm0005-1.shtml You need the following supplies to do this experiment:

buret, 50 mL buret stand or retort stand with appropriate clamp beaker, 100 mL waste flask, 100 – 500 mL, left under buret tip whenever another receiver is not there narrow-stem funnel (for filling buret) square of Parafilm, 4” x 4” (10 cm x 10 cm) with a 1/4” (6 mm) hole punched 1 “ from one edge transfer pipet or dropper plain white card to use as a background while using and reading the buret

These supplies should be available in the lab:

stopcock grease either: aqueous Alconox detergent, acid alcohol, or cleaning solution gloves (to wear while cleaning and using the buret) buret brush deionized or distilled water Kleenex or Kim-Wipes thermometer ( 0 – 110 C)

Step 1: Preparing and Cleaning the Buret Keep in mind that a buret is a $100 item.

1. If the buret is visibly dirty, scrub it out using a buret brush and Alconox. Rinse with ten changes of tap water and two changes of distilled water (Alconox sticks tightly to glass).

2. Remove the stopcock by unscrewing the nut (Teflon valve) or removing the steel clip (glass). Examine the valve. The hole should be clear and there should be no residue of grease on the ends of the valve. Clean it with a paper clip and paper tissue if necessary.

3. Replace Teflon valves without greasing, and tighten the nut snugly. Lightly grease glass valves and insert. Check that the valve turns freely without popping out.

4. Mount the buret securely in the stand. It should not slide up or down, and you should be able to read the entire calibrated scale.

5. Test the buret for cleanliness by filling it with deionized water and draining slowly to the 50 mL mark. Watch it drain--If the meniscus distorts from a clean concave shape, or if drops of water cling to the inside of the barrel, it is not clean.

6. Clean the buret by filling with acid alcohol (preferred), Alconox, or cleaning solution. Leave a few minutes and drain. If necessary, demount the buret and use a brush. Rinse 10 times with tap water and twice with deionized water.

7. If the buret is clean, measure the time needed for the film of water to stop draining, i.e., the time in seconds for the meniscus to stop rising. You will need to wait this time before reading every time you read the buret. In no case should it be less than 15 seconds.

8. Practice reading the meniscus before going on to the next step. Hold a white card behind the buret and look at the meniscus with your eye at the same level as the meniscus. Sometimes it is necessary to lower the buret to do this. Always read the bottom of the meniscus. You should be able to read the buret to 0.02 mL (one fifth of the smallest graduation mark) without eyestrain. See if your colleagues agree with you!!!


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The Right Way to Hold a Buret

Two of the websites listed above show the buret being used improperly! This is how to hold a buret (for right-handed persons): Turn the buret so the valve is to the right. With the left hand, loosely grip the entire valve body, and operate the valve with the fingertips. With the right hand, hold the receiver, tilted so that the tip of the buret touches its inner wall. Did you know that it is normal technique to titrate to the nearest tenth of a drop? Small increments of a drop are deposited on the inside wall of a receiver and rinsed into the bulk solution with a stream from a wash bottle. (Do not do this in today’s experiment, however.) Step 2: Calibration

1. Prepare a 100 mL clean beaker by covering it with Parafilm so that the small hole is next to the wall of the beaker. (What is the purpose of the Parafilm?)

2. Obtain some deionized water. Measure its temperature. 3. With the waste flask in place. Fill the buret with deionized water. Open the valve briefly so that

the tip fills with water, with no bubbles. A single bubble can defeat your calibration efforts. Adjust the meniscus as closely as possible to 0.00 mL (the topmost mark) by adding water with a transfer pipet, and draining water slowly from the tip, using the tilted waste flask shown above.

4. Note the time and the exact reading of the meniscus (it may not actually be 0.00). 5. Set the balance to zero with an empty pan. Weigh the empty beaker and Parafilm to an accuracy

of about 5 mg. (Do not tare the balance for your beaker. The experiment will take some time, and others may come and change the tare setting in the meantime.)

6. Hold the beaker so that the buret tip passes through the hole in the Parafilm and touches the inside wall of the beaker. Drain 5.00 mL as accurately as possible into the beaker.

7. Weigh the beaker again while waiting for the meniscus to drain. 8. Read the meniscus accurately. 9. Repeat the above two steps until the bottom of the scale is reached. Do not empty the water from

the beaker between steps. 10. If the buret can be stored vertically between classes, fill it with water and cover the top with

Parafilm.


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Step 3: Calculations If you have been successful, you will have a temperature and a list of buret readings and beaker weights.

Calculate the net weight of water added to the beaker by subtracting the initial weight of the beaker before water is added from the buret.

Convert the weight of the water to volume by looking up the density of water at the temperature of the experiment. Divide that density into the net weight of water to get the actual volume of water dispensed from the buret. (Note: the attached table reads in mL/g, the opposite of density!!!)

Subtract the buret reading from the actual volume of water to get the correction factor. The correction factor (F) should be plotted on a graph as the dependent variable against the

volume dispensed (B). Using the Correction Plot or Table When doing actual titrations, the raw buret reading should be looked up on the abscissa of the error plot. The correction factor should be estimated and added to the raw number to get the actual volume.

Buret Calibration Chart

0

0.01

0.02

0.03

0 10 20 30 40 50 60

Scale Reading (mL)

Co

rrec

tio

n (

mL

) (add correction to reading)


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Chem 247 - Buret Calibration Data Sheet (REPRODUCE IN YOUR NOTEBOOK) Date Name Temperature = Density of water at temperature = g/mL Buret Reading (mL)

Total weight of beaker, water and Parafilm(g)

Total weight of water added to beaker (g)

Exact volume of water (mL)

Exact volume minus buret reading (correction)

B observed C observed D = C – tare wt E = D/density F = E - B


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Density of Water at Various Temperatures and After Changing to 20 C Corrections for the thermal expansion of the glass buret have been applied. Temp Volume of Volume of 1 gram (oC) 1 gram water at 20 oC at temp. 10 1.0013 1.0016 11 1.0014 1.0016 12 1.0015 1.0017 13 1.0016 1.0018 14 1.0018 1.0019 15 1.0019 1.0020 16 1.0021 1.0022 17 1.0022 1.0023 18 1.0024 1.0025 19 1.0026 1.0026 20 1.0028 1.0028 21 1.0030 1.0030 22 1.0033 1.0032 23 1.0035 1.0034 24 1.0037 1.0036 25 1.0040 1.0037 26 1.0043 1.0041 27 1.0045 1.0043 28 1.0048 1.0046 29 1.0051 1.0048 30 1.0054 1.0052 Questions 1. Your weight of water is converted to the true volume using data from the table above. What are the three corrections that are embodied in those values? 2. Explain why it is not necessary to weigh the water samples to the maximum available accuracy. 3. When the glass of a buret expands due to an increase in temperature does the diameter of the bore increase or decrease? 4. Most volumetric glassware is calibrated at what temperature? How can you find out? 5. What do the letters T.D. and T.C., found on various types of volumetric glassware, signify?


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EXPERIMENT 2: ACID-BASE TITRATION: DETERMINATION OF THE PERCENTAGE OF KHP IN AN UNKNOWN PRELIMINARY READING: Chapter 2 of Harris 5. These sections discuss the operation of the analytical balance, the proper use of burets and volumetric flasks, and the method of weighing by difference. Lab notebooks should be prepared as instructed in your textbooks and in this lab manual. INTRODUCTION: This experiment is similar to those you may have carried out in General Chemistry or other classes. However, in this course, the emphasis is on technique and accuracy. The experiment consists of two parts. Using primary standard grade potassium hydrogen phthalate (KHP), the concentration of a standard sodium hydroxide solution will be determined to at least three significant figures by manual titration. This standardized NaOH will then be used to determine the percentage of potassium biphthalate (KHP) in an unknown sample. PROCEDURE: Preparation of carbonate-free NaOH solution, approximately 0.1 N. You will be supplied with water that has been sparged with nitrogen to remove most of the carbon dioxide dissolved in it. (Alternately, boil 1 liter of deionized water in a one liter beaker for 3-4 minutes. After boiling the water, COVER it with a watch glass and allow it to cool until it can be touched comfortably with your bare hand.) Use the water to half-fill a one liter polyethylene (plastic) bottle. OBTAIN 6 mL of 50% NaOH solution in your 10 mL graduated cylinder, and ADD it to the bottle water. MIX thoroughly by inverting the bottle several times. Finally, fill the bottle to just below the top with degassed water. The result is an approximately 0.1 N NaOH solution, which will be standardized. Do not keep NaOH solutions in glass-stoppered bottles. The stoppers will stick permanently in place. NOTE: You will use the NaOH solution prepared above for your first three experiments. Keep the bottle tightly capped except when dispensing solution (WHY?). Drying KHP and 'UNKNOWN" samples. Obtain a 5-6 g portion of primary standard grade KHP and an individual "UNKNOWN" sample from the Teaching Assistant. Be sure to keep a record of the Unknown number in your notebook and one the lab form. PUT the samples in labeled weighing bottles. Place the labeled weighing bottles into a beaker (to avoid tipping and confusion in the oven) and COVER with a watch glass. DRY in an oven at 120 C for one hour. STORE in a desiccator until needed. Standardization of NaOH Solution. Accurately WEIGH out 0.7 to 0.9 g dried reagent grade KHP into each of three labeled 250 mL conical fIasks. ADD 50 mL of CO2 free deionized water to each flask, SWIRL gently until the solid has


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dissolved, and COOL to room temperature. ADD 4 to 5 drops of 0.1% phenolphthalein indicator solution to each flask and MEASURE the volume of carbonate-free NaOH solution necessary to reach a stable end point by recording the initial and final buret readings. You should be able to do this to a precision of 0.02 mL. When titrating, be cautious as you come near the endpoint. The phenolphthalein color will persist longer as you get close. Fractional drops from the buret can be made by allowing a partial drop to form at the tip, touching on the side of the flask, and rinsing it into the solution with a brief stream from your wash bottle. Aim to get a light pink color that persists for a few seconds. Performance is the key in this exercise. If you miss an endpoint, note it in your notebook and prepare another sample. You should have three reliable titrations to fix the concentration of the standard NaOH.

Measurement of KHP in the Unknown Accurately weight three 0.7 to 0.9 g samples of the dried unknown into three labeled 250 mL flasks and dissolve in water as above. Titrate in a similar way. Don’t forget the indicator. Calculations To compute the normality of the NaOH solution, you can take advantage of the two main equations that rule volumetric analysis:

Eq. (equivalent weights) = N1V1 and N1V1 = N2V2 The equivalents (Eq) of a sample is the weight in grams divided by the equivalent weight (grams per equivalent). The EW of KHP is the same as the molecular weight. It has just one acidic proton. From your accurate weight of each KHP sample, you can calculate the equivalents of KHP in each. Since you know the volume of titrant for that sample (in liters), you can calculate the normality of the titrant. The calculation is performed in reverse when titrating the unknown samples. The Eq of KHP can be computed from the titration volume and the newly-determined normality of the NaOH titrant. Computation of Error Compute the a priori error in the final results by propagating the estimated errors in buret readings and weights. Compute the a posteriori estimate of error by calculating the mean and standard deviations in the final results. Obtain the results from another person who is analyzing the same sample number, if possible. Pool their results with yours. If they are greatly different, try to explain why that might be so. If the error limits by the two methods are not similar, explain why this might be so.


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EXPERIMENT 3: DETERMINATION OF THE PKa OF A WEAK MONOPROTIC ACID; PRELIMINARY READING: Harris 5, Chapter 12.1, 12.2; Ch. 15.4 INTRODUCTION: In this experiment, you will derive a titration curve for potassium hydrogen phthalate (KHP), and then determine the pKa and the equivalent weight of a unknown weak monoprotic acid. A pH meter will be used to follow the change in pH of the solution during the course of the titration. The use of the pH meter will be demonstrated in class. A shortage of pH meters may require that this lab be done with a partner. If so, each partner must do at least one titration curve, and the unknown should be titrated twice. PROCEDURE: Standardization of the pH meter. OBTAIN a pH electrode and pH meter (millivoltmeter). Immediately PLUG in the meter and set the switch to STANDBY to allow the circuit to warm up. POUR 10 to 15 mL of standard pH buffer solutions, pH 4.0 and pH 10.0, into two small, labeled beakers. RINSE the electrode with deionized water (DIW) and blot it carefully with a tissue. IMMERSE the electrode in a standard calibration solution, making sure the liquid covers the glass bulb on the bottom, as well as the reference junction. Examine the pH electrode to determine where these features are located. Consult the instructor if you cannot determine that depth. ADJUST the meter to read the known pH of the solution. Several seconds may be required for the meter to reach a stable reading. RETURN the meter to its STANDBY position. The meter should always be in this position when the electrode is removed from a solution. RINSE and BLOT the electrode and immerse it into the second calibration solution of a different known pH. READ the pH. If it is within 0.3 units of the second known pH value, everything is functioning properly. Consult the instructor if there are any problems. Titration Curve for KHP WEIGH accurately 0.2 to 0.3 g dry potassium hydrogen phthalate and ADD to a 100 mL beaker. Add about 35 mL DIW and swirl to dissolve. ADD a few drops of phenolphthalein indicator. Insert the pH electrode to a depth of at least one inch. A stir bar may be added, but do not allow it to strike the electrode. Stir slowly without forming a vortex. RECORD the inital pH of the sample solution. Fill the buret with standardized NaOH and RECORD the initial volume. ADD enough base from the buret to cause the pH to change by about 0.3 units. RECORD pH and NaOH buret reading. Initially, larger volumes of base may be necessary to change the pH by 0.3 units, but near the endpoint, only a fraction of a drop will cause large pH jumps. Go slowly in this region. Good technique is required to produce a smoothly progressing pH curve. NOW OBSERVE AND RECORD the pH at which the indicator first turns pink. Continue past the endpoint until 10 mL of added base causes no further pH change.


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Titration of Unknown. OBTAIN a sample of an unknown weak acid. UNKNOWNS are provided in numbered containers. Select one of these unknowns and record the number. The source bottles for these unknowns, with vital statistics on the label, are in the hood. Take about 2 g of the acid in a weighing bottle and dry for 30 min if a solid. NOTE: Some of the unknowns will be liquids. To weigh liquids, dispense drops of the liquid into a tared beaker and take the gross weight quickly. Immediately rinse the weighed acid into the titration beaker, taking care to washdown the sides so that all of the acid is transferred. Make the final volume to approximately 50 mL with DIW. When not using unknown, make sure it is capped since these acids are also volatile. SAFETY CAUTION: Although these unknowns are weak: acids, they may be in concentrated form. Therefore, be very careful when handling them. If you do spill some on yourself, quickly rinse with copious amounts of water. Derive a titration curve using the same technique as for the titration of KHP. Keep in mind that the endpoint will occur at a different place. LAB REPORT: TITRATION CURVES should be included and plotted with the volume of NaOH added along the x-axis (abscissa) and the pH along the y-axis (ordinate). DRAW a smooth curve through the data points. To determine the EQUIVALENCE POINT, find the midpoint of the steep portion of the titration curve. Its corresponding volume on the abscissa is the volume of NaOH required to reach the equivalence point (Veq). Determine the pKa of the acid by graphically finding the pH corresponding to half the volume of base necessary to reach the equivalence point. At this halfway point, the solution is a buffer with equal concentrations of HA and A- and its pH is equal to the pKa of the unknown acid. CALCULATE from the equivalence point, the known normality of the base, and the weight of the unknown, its EQUIVALENT WEIGHT. Include answers to DISCUSSION QUESTIONS. NOTE: Although the value for pKa should be reproducible, it may be as much as 1 pH unit different from the “true” pKa that is usually found in tables and is sometimes called the thermodynamic pKa. The difference is partly due to ion activity effects. DISCUSSION QUESTIONS: 1. Show by appropriate equations why the graphical data analysis gives the equivalence point and pKa correctly. 2. Calculate pKa values from five data points in the buffering region (one curve only). In this region

pH = pKa + log10 [conjugate base A-]/[weak acid AH]) OR

pKa = pH + log10 [(Veq - V)/V] where pH is the measurement corresponding to the volume of added titrant V.


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EXPERIMENT 4: EQUIVALENT WEIGHT OF POLYPROTIC ACIDS. PRELIMINARY READING: Harris 5, Chapter 12.1, 12.2; Ch. 15.4 (same as last week)


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29


30

INTRODUCTION:


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This experiment introduces you to more complex acids having multiple dissociating protons. The goal of this experiment is to determine the pKa's and equivalent weight of an unknown polyprotic acid. Two different titrations will be done using the pH meter. One will be on maleic acid, which has well-separated pKa’s and gives a good, classical curve. The other will be on a more difficult unknown acid.


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34


35

PROCEDURE:


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37

OBTAIN an unknown sample from the series of acids on the side table. Record the number.


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39


40

For the first titration, WEIGH accurately 0.3 to 0.4 g of maleic acid, and add to 50 mL water in a conical flask or beaker. Rinse down the sides of the flask.


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NOTE: Some of the solid acids require much stirring and perhaps gentle heating to dissolve completely. Complete dissolution is critical to an accurate equivalent weight determination.


42


43


44


45

Set up the pH meter as in the last lab, and TITRATE with your standard NaOH until the pH reaches about 9 or 10. Record the pH and volume of NaOH as you go. WEIGH a 0.4 to 05 g sample of your UNKNOWN and repeat the titration. Multiple "jumps" or steps should be seen (two or more). Each step corresponds to the titration of one dissociating proton. There is one equivalence point per step. If time allows, repeat this last titration. LAB REPORT:


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47


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PLOT titration data (pH vs. volume) and graphically determine the equivalence points, the volume of base required to reach them, and the corresponding pKa’s. In the case of maleic acid, the two equivalence points should be clearly seen. With other polyprotic acids, the first equivalence point may be indistinct, but the second will be visible. CALCULATE the equivalent weight of your unknown from the titration curves. REPORT pKa's, equivalent weights, and appropriate statistics.


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EXPERIMENT 5: GRAVIMETRIC CHLORIDE PRELIMINARY READING: Harris 5, Ch. 27 INTRODUCTION: The chloride in a water soluble unknown sample will be determined by precipitation by silver nitrate. The silver chloride precipitate will be filtered, dried, and weighed. The results of this determination can be compared with the results of the ion-selective electrode method to be done later in this series. PROCEDURE: Clean three Gooch crucibles if necessary by placing them in a beaker and adding 2 or 3 mL concentrated ammonium hydroxide in them. DO THIS IN THE FUME HOOD. After five minutes, add about 100 mL water to dilute the ammonium hydroxide and wash it down the drain. Dry the Gooch crucible in the oven for 30 minutes and store in the desiccator to cool. When cool, number the crucibles with pencil and take the TARE weights. OBTAIN a numbered unknown from the teaching assistant. Find someone in the lab who is using the same unknown. KEEP THIS UNKNOWN FOR THE ION-SELECTIVE ELECTRODE EXPERIMENT. Dry your sample for two hours. Do this during the preceding laboratory if possible. WEIGH accurately two 0.5 g samples. DISSOLVE in 100 mL 0.1 N nitric acid. Keeping the light dim, ADD 50 mL 0.2 N silver nitrate solution slowly with stirring. Rinse the stirring rod into the solution. DIGEST the precipitate on a hot plate at 60 – 70 C, without stirring, for about 30 minutes. OBSERVE the changes in appearance as digestion proceeds. Set up the Gooch crucibles as shown by the instructor, and FILTER the solution through it carefully. Use the stirring rod to guide the flow of solution from the beaker. Use your wash bottle to rinse down and transfer all the precipitate to the crucible. RINSE the precipitate with 5 mL lots of 0.1 N nitric acid. DRY the precipitate at 160 C for two hours and cool the crucible. WEIGH the crucible. Replace in the oven until the next lab period and weigh again. (This is called ‘drying to constant weight.’) LAB REPORT: Decide which weight to use for your calculation. Calculate the weight of the precipitate and the amount of chloride in the samples. Express it as a percentage of the total weight of the sample. Share your results with another person who has used the same unknown. Calculate the mean and standard deviation of the percent chloride. DISCUSSION QUESTION: Explain what happened during the digestion process and why. Name possible sources of error in this lab. If your precipitates turned gray or violet on heating, what happened?


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EXPERIMENT 6 NITRITE ANALYSIS METHOD READING: This lab manual, section on spectrophotometry and Spectronic 20. INTRODUCTION: The determination of nitrite ion in water or in food samples is an important and typical analytical problem. Nitrites can contribute to pollution in water supplies and have been implicated as causing the formation of carcinogens in certain food products. A highly sensitive, spectroscopic technique uses the diazotization coupling of nitrite with amines. The compound 4-amino-benzenesulfonamide (sulfanilamide) reacts with nitrites as follows:

This diazonium salt is then coupled to another amine to give a strongly red colored product.

This experiment introduces the use of standards, and the construction of a calibration curve. Absorbance, the spectroscopic signal, varies linearly with concentration as given in Beer's law, A = .b.c. The absorbance of the unknown nitrite solution is then measured and its concentration determined from the calibration curve. In this first laboratory using this method, we will prepare our own reagents, construct a calibration curve, examine the acid-base properties of the colored product, and measure an unknown nitrite solution. In the next laboratory, we will extract the nitrite from a commercial meat sample, ‘clean up’ the sample in preparation for analysis, and finally measure the amount of nitrite in the sample. PROCEDURE: NOTE to TA: The reagents used in this experiment can be unstable, and the experiment should be checked beforehand. Preparation of stock reagent solutions. Accurate Nitrite Stock: PLACE 2 to 3 g of NaNO2 into a weighing bottle and dry for 1 hour. Cool in your desiccator. Accurately WEIGH about 0.25 g (record exact weight) of the dried NaNO2 into a conical flask or beaker. DISSOLVE in deionized water and TRANSFER quantitatively to a 500 mL volumetric flask. ADD water to the 500 mL mark, MIX thoroughly, and TRANSFER to a labeled bottle or flask. This solution is still too concentrated to use. RINSE out the 500 mL volumetric flask twice. PIPETTE exactly 5.00 mL of the concentrated solution into the 500 mL volumetric and dilute to 500 mL. This more dilute solution is the stock standard nitrite solution and has to be made up fresh. Keep the more concentrated stock for next week.


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Sulfanilamide Reagent: WEIGH about 5 g (4.9 to 5.0 g) of sulfanilamide into a 400 mL beaker and ADD 188 mL of water. ADD 62 mL of concentrated hydrochloric acid to the sulfanilamide solution in a FUME HOOD (conc HCl emits irritating fumes). STIR until solid is dissolved (which may take a while, then TRANSFER to a storage. This solution is UNSTABLE and should be discarded after 8 days (both nitrite experiments). NED Reagent: DISSOLVE approximately 50 mg of N-naphthylethylenediamine (NED) in 250 mL of water and STORE in a sealed vessel. This solution will keep for 2 or 3 weeks. Preparation of calibration solutions. The final amount of colored product in the above reaction is subject to many complex variables that are difficult to control. The most accurate results are obtained if a calibration curve is prepared using samples of known nitrite concentration, AND the same reagents that will be used for later analysis of the samples. This curve should consist of a plot of observed absorbance (y-axis) versus nitrite in moles or micromoles (x-axis) for a series of solutions where the nitrite level is increased in a systematic way. Then an unknown nitrite sample can be subjected to the same reaction, and from the observed absorbance of the unknown, its concentration can be graphically determined. PREPARE six calibration solutions in 200 mL Erlenmeyer flasks by adding 0, 3, 6, 12, 21, and 30 mL aliquots of the stock standard nitrite solution. Use a buret to add the appropriate aliquot of stock nitrite to the flask. ADD approximately 50 mL of water, and then ADD 10 mL of sulfanilamide reagent. ALLOW solution to stand for at least 10 minutes. ADD 10 mL NED reagent and let stand another 10 minutes. Quantitatively TRANSFER the solution to a 100 mL volumetric flask. ADD water to 100 mL mark, and MIX completely. Use several small portions of this solution to RINSE the flask from which the solution originally came. Transfer the accurately diluted solution back to the flask, so that the volumetric flask can be used again for the other calibration samples. Let solutions stand for 15 minutes before measuring absorbance. Determination of the wavelength of maximum absorbance of the final complex. Get two cuvettes from the TA. One of these can be used for the reference (zero) sample, and the other for the solution being measured. SET wavelength of spectrophotometer to 570 nm. USE the "0.0 mL" (zero) calibration solution as a reference to zero the spectrophotometer. ZERO the spectrophotometer as described in the lab manual or as demonstrated. Keep the cuvette filled with the zero solution. You will need it throughout the measurements. Using another cuvette, CHOOSE one of the other solutions that has approximately median color intensity and MEASURE its absorbance in 20 nm intervals between 400 and 600 nm, except where a peak is detected. When in the peak absorbance region, use 5 nm increments. REMEMBER TO RE-ZERO THE INSTRUMENT WITH YOUR REFERENCE SOLUTION EACH TIME THE WAVELENGTH IS CHANGED. PLOT the absorbance observed versus wavelength. DETERMINE the wavelength of maximum absorbance.


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Measure the absorbance of the calibration solutions with the spectrophotometer set at the wavelength of maximum absorbance. MEASURE the absorbance of each of the nitrite calibration solutions. SAVE the solution used to determine the peak: absorbance wavelength of the final complex. Analysis of the Unknown OBTAIN a nitrite unknown solution and RECORD the unknown number. The unknown is prepared and measured in the same lab period as the calibration solutions. No dilution may be necessary. REACT the unknown in the same fashion as the calibration solutions. Use your own judgement as to the size of the aliquot to add and attempt more than one determination. IF THE ABSORBANCE IS TOO HIGH (A >1.00) OR TOO LOW (< 0.04), do the unknown again, adjusting the amount of sample used appropriately. Effect of pH. pH is a critical variable in this experiment. TAKE a portion of the same calibration solution used to determine the peak absorbance wavelength and ADD concentrated ammonia using a dropper. Continue additions until the color change is complete (dye changes from pink to orange), but avoid excessive base. DETERMINE the absorption spectrum of the basic solution by measuring absorbance between 400 and 600 nm as before. PLOT absorbance versus wavelength on the same grapb as the previously measured absorption spectrum. SAVE YOUR SULFANILAMIDE, NED AND NITRITE STANDARD SOLUTIONS. LAB REPORT: The report must include (a) the absorption curves vs wavelength under acidic and basic conditions, (b) the Beer's law calibration curve of absorption versus nitrite concentration (in MOLAR UNITS), and (c) the concentration of the unknown in molar units. ANSWER discussion questions. DISCUSSION QUESTIONS: 1. Which form of the final product, in acid or in base, is better for a spectrophotometric determination? Why? 2. When the stock nitrite was prepared, a solution of 0.25 g/500 mL was first made and then diluted 100x. Why not just prepare a 0.0025 g/500 mL solution without dilution? Compute and compare the expected relative error by the first technique, called serial dilution, with that of the second. Assume that the weights are accurate to 0.0002 g and volume to 0.01 mL


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EXPERIMENT 7 NITRITE ANALYSIS IN A COMPLEX SAMPLE READING: INTRODUCTION: One of the central problems of real-world analytical chemistry is that the anaIyte is often a minor part of the sample. The bulk material, or matrix, not only dilutes the desired species, but may interfere with analytical chemical reactions or with the instrumental technique of analysis. The purpose of this experiment is to gain some experience with analyzing a trace constituent of a typical complex sample. Sodium nitrite is added to cured meat products such as hot dogs, salami, ham, and bacon because it preserves the red color of meat and it has strong antibiotic activity against microorganisms, including Clostridium botulinum, which causes a deadly form of food poisoning. Nitrite will be extracted from cured meat in this experiment and analyzed by the diazonium dye reaction practiced in the last laboratory. Three cleanup steps are involved: Water extraction of nitrite from the meat sample. Removal of fat from the water extract. Ion exchange chromatography to concentrate the nitrite and separate it from interfering species. This experiment uses a strong Type 1 anion exchange resin, such as Dowex 1 or AG 1, which bears fixed quaternary amine groups. This resin is provided as a suspension in deionized water in the chloride form

R-NH4+Cl- It will be used to adsorb NO2- from the cured meat according to the exchange reaction

R-NH4+Cl- + NO2- --> R-NH4+NO2- + Cl- Nitrite is adsorbed by displacing chloride ion from the resin. A large volume of sample, containing a trace of nitrite, can be passed over a short column of ion exchange resin. The nitrite is in turn displaced from the resin by washing with 1 M NaCl. At this high concentration of salt, nitrite is not able to compete with chloride for binding sites and is released into solution for subsequent analysis. PROCEDURE: Several of the following steps can be performed simultaneously. Organizing and scheduling your work to finish within the lab period is one of the goals of this laboratory. Preparation of meat sample. OBTAIN a cured meat sample from the instructor and WEIGH 30 g on a top-loading pan balance, recording the weight to the nearest 0.1 g. Cut it into a small pieces and TRANSFER to a beaker. ADD 50 mL cold deionized water. With a spatula, stir the mix every few minutes. ALLOW water to perfuse the sample for at least 10 minutes, then DECANT suspension through a Kim-Wipe tissue stretched over the top of a glass funnel. Only pour the liquid into filter initially, and try to retain the solid pieces in the beaker. Add 10 mL water to the solid


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meat and leach for another 10 minutes. Pour this fluid through the filter. This may be the most time-consuming step. If the filtrate appears excessively turbid, REPEAT the filtration with new filter paper. If there are globules of fat floating on the filtrate, however small, be sure to filter it again. If filtration is slow, it can usually be improved by mixing one gram of Celite filter air with your sample. Celite prevents small particles from closing off the flow paths through your sample. Preparation of the ion exchange column. PRETREAT a sample of anion exchange resin. Put 3 to 4 g Dowex 1 or AG-1 in a 50 mL beaker and wash with at least three changes of DIW. Add DIW until the beaker is half full above the resin. Stir or swirl, let the resin settle, and DECANT the overlying liquid. (Decantation means pouring the liquid off without losing any of the solid resin.) Then add about 20 mL 1 N NaOH, decant, and wash three times with DIW. Finally, cover the resin with 20 mL 1 M NaCl, decant and wash three times with DIW. The resin is now ready for use. PREPARE a chromatography column by attaching a Pasteur pipet to a glass funnel with a short piece of latex tubing. Slide a piece of thin vinyl tubing over the narrow end of the pipet and close the tubing with a pinch clamp. MOUNT the column vertically on a stand making sure the bottom valve stays closed. Using a transfer pipet and your wash bottle, add resin to the column until the resin plug is about one inch long. Trapped air can make this procedure quite difficult. Ask the instructor to help if necessary. DON'T LET THE RESIN GO DRY, or the flow rate of the column will be much slower. Use a long Pasteur pipet to remove bubbles from the column or to resuspend the resin to remove air. Always leave a little water standing above the resin. Loading of filtered sample and elution of nitrite. ADD the meat extract to the column and start the column flow. COLLECT this effluent in a beaker and save (in case of error). You can continually ADD extract to the top of the column during this loading process, but the liquid level must return to 1 - 2 cm above the resin after all the filtrate has passed through. WASH the column with 15 mL of deionized water using a dropper. Maintain the liquid level at least 1 cm above the resin bed. ELUTE the column with 10 mL of 1 N NaCl solution and COLLECT the effluent in a 25 mL graduated cylinder. RECORD the collected volume. Spectroscopic analvsis for nitrite. Solutions of sulfanilamide, NED, and stock nitrite from the previous experiment are used here. The nitrite stock has to be diluted to working strength by pipetting 5.00 ml into a 500.0 mL volumetric flask and diluting with water. One nitrite determination should be done on the stock aliquot and compared with previous results to insure reagent quality. If in doubt, do two or three calibration points. Since you don’t know how much nitrite is present, a “range-finding” analysis is done first to estimate the correct amount of sample to use to yield a color within the colorimetric range. More than one range-finding test may be needed. ANALYZE the final effluent for nitrite using the dye reaction. Measure 1.00 mL of the eluate into a flask, and add water, sulfanilamide, and NED as in the last experiment. Make the sample to 100 mL. If the color is very pale or very dark, change the dilution appropriately.


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ADJUST the aliquot size of effluent to obtain reliable absorbance readings. As much as a 1;100 dilution of the eluent may be needed to bring the sample into range. Once it is diluted, samples of 1, 5, and 15 mL might be needed to get a reasonable absorbance. Keep track of your serial dilutions. From the absorbance readings, the previously prepared calibration curve, and known dilution factors, the concentration of nitrite in the unknown can be determined. TEST an aliquot of the effluent from applying the sample to the column for nitrite by running the dye reaction. DO this in a small beaker, since it is just a test of the effectiveness of the ion exchange method. Visually estimate whether there is any color apparent. PLACE all meat scraps and filter paper in the trash area recommended by the instructor. WASH glassware thoroughly to remove grease. LEAVING MEAT SCRAPS IN THE SINK WILL ANNOY THE INSTRUCTOR!!! LAB REPORT: Sketch the chain of dilutions and extractions from the original meat sample to the final dilution used for the analysis. Use this to guide you calculations to compute the amount of NaNO2 in the sample. Report the concentration of NaNO2 found in the meat in units of both micrograms NaNO2 per g of meat and micromoles per gram. Discuss the practical problems of the procedure and sources of error.


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EXPERIMENT 8: DETERMINATION OF CHLORIDE BY ION SELECTIVE ELECTRODES. PRELIMINARY READING: Harris 5, Chapter 15, especially 15-1, 15-5, 15-6. INTRODUCTION: Recent technology has greatly simplified the analysis of certain ions in solution by the development of ion-specific (selective) electrodes (ISE). These electrodes give electrical indication of concentrations. The familiar pH electrode is an ISE. ISE’s can be used for fast, routine analyses on complex mixtures such as foodstuffs, blood, sewage, etc. Commercial electrodes are now available for many common ions such as Br-, Cl-, F-, CN-, K+, Na+, and NO3-. This experiment introduces you to the chloride specific electrode, which is typical of most ISE’s. You will examine typical responses, prepare calibration curves, and investigate limitations and interferences. You will also become acquainted with serial dilution techniques, because ISE’s typically work across many order of magnitudes of concentration. The proper protocols for preparing the calibration solutions should be computed before the lab begins. SIMPLIFIED THEORY: The Cl- electrode consists of finely dispersed AgCl and Ag2S mixed in an epoxy plastic and cast into a membrane or pellet. When the electrode is immersed in a Cl- solution, Cl- ions in the solution compete with Cl- ions in the plastic matrix for the available silver ions.

AgCl(s) + Cl-(aq) AgO(s) + Cl-(solid fixed in matrix) The loss of Cl- from the solid matrix is prevented by the materials selected for the construction of the electrode and the sulfide. The reaction is confined to the surface of the electrode and occurs only to an extremely small extent. By conventional techniques, the AgCl formed in solution or the additional Cl- in the solid could never be detected. Only the outer surface in contact with the solution has excess Cl-. Note that all species above are solids and hence have unit activities, except for Cl- in solution. Therefore, the extent of reaction depends solely on the solution concentration of Cl-. The charge separation across the membrane, caused by the imbalance of Cl- at the solution-electrode interface, establishes a potential. This potential of a single half-cell cannot be measured alone, but must be coupled to a second half-cell so that the potential difference can be electrically measured. The combination ion-selective Cl- electrode has another internal cell containing solid AgCl, solid Ag, and Cl- in solution that forms the second or reference potential. The potential difference is measured by a potentiometer that has a very high impedance (> 10^12 ohms) to prevent current from passing during the measurement. The actual potential developed across the solid membrane is impossible to predict a priori because it depends strongly on properties of the cast membrane, surface, scratches, etc. Therefore, an electrode must be calibrated against known concentrations of Cl- and recalibrated frequently. The exact theory of electrode response demonstrates that the developed potential depends on the thermodynamic activity of Cl- ions. Because total ionic strength affects ion activity, it is important to insure total ionic strength remains constant, although the Cl- concentration may change.


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Other ions, especially anions, may interfere by causing membrane potentials to develop, although they are usually smaller in magnitude. On the other hand, bromide and iodide ions may produce a stronger response than chloride itself. Hence, these electrodes are more accurately called ion-selective rather than ion specific electrodes. PROCEDURE: Preparation of calibration solutions. PREPARE 100 mL of a 1 M KNO3 solution. The nitrate ion does not react with the chloride electrode, but can be used to establish a constant ionic strength. PREPARE 500 mL of an accurate 0.100 M solution of NaCl (i.e., the concentration may not be exactly 0.1 M, but is accurately known). (Read the formula weight from the sodium chloride bottle and divide by 20 to get the approximate amount of NaCl to weigh.) By SERIAL DILUTION, make a series of six calibration solutions containing 0.1, 0.05, 0.01, 0.001, 0.0001, and 0.00001 M NaCl and sufficient KNO3 solution to make the final total ionic strength (NaCl + KNO3) equal to 0.1 M. Approximately 50 mL of each calibration solution is needed, but it may be most convenient to use your 100 mL volumentric flask. For example, to prepare the 0.05 M NaCl solution, take 50 mL of 0.1 M NaCl solution and 5 mL of 1 M KNO3 solution and dilute to 100 mL in a volumetric flask. Final concentrations of NaCl and KNO3 equal 0.05 M each, or a total salt concentration of 0.1 M. To prepare the next NaCl solution, take an appropriate aliquot of this newly made solution and adjust salt concentration with an appropriate KNO3 aliquot. Continue likewise for rest of solutions. Before the laboratory, prepare a table showing the amounts of NaCl and KNO3 solutions to be used to make each calibration solution. For the most dilute NaCl concentrations, consider that you may have to dilute your 0.1 M NaCl standard solution by 1:100 or 1:1000. Preparation of chloride unknown solution. By using the percent of chloride in your unknown (as you determined it in the gravimetric chloride experiment), calculate the amount of solid unknown necessary to prepare a 0.05 M solution of Cl- (100 mL). To prepare this solution, WEIGH out the solid, DISSOLVE in 20 mL of water, ADD enough KNO3 to make its final ionic strength 0.05 M, and ADD water to bring it to a final volume of 100 mL. Measurement of potentials. Use the chloride-selective electrode as demonstrated by the instructor. CHECK that the KCl electrolyte in the electrode shell is full nearly to the fill hole. Set the pH meter to “millivolts”. CAUTION: Since the millivolt scale is arbitrary, be sure not to reset the “calibration” setting once the potential measurements have begun, for the entire laboratory session! MEASURE the potential difference in millivolts for each of your calibration solutions and your unknown solution.


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Measurement of MnCl+ formation constant. Note the formula weight of the manganese sulfate provided. ADD 2.0 g of MnSO4 to 50 mL of the 0.01 M NaCl calibration solution (with KNO3). DETERMINE the potential of the resulting mixture and COMPARE with the potential of the original 0.01 M NaCl solution. Assuming Mn+2 + Cl- MnCl+ and Kf = [MnCl+]/([Mn2+].[Cl-]) calculate the equilibrium constant from your data on the mixed MnSO4 - NaCl solution. Selectivitv of electrode. PREPARE an accurate 0.100 M solution of NaBr in 500 mL deionized water. Make dilutions of 0.010 M and 0.0010 M and add KNO3 solution to make the ionic strength equal, using the table you generated for chloride solutions. Measure the potential of the three solutions. Plot the concentrations and potentials on the same plot as your chloride curve. The horizontal displacement of the straight part of the line, translated into concentration terms, is your selectivity factor. For example, if 0.010 M NaBr gives the same potential as 0.0010 M NaCl, then the electrode is 10 times less selective for bromide than chloride. Optionally, you may carry out the same experiment using 50 mL of 1 mg/mL KI or NaI, as provided. LAB REPORT: Plot the recorded potential (ordinate) versus log10[Cl-] (abscissa) for the calibration solutions. Using the recorded potential for the unknown, graphically determine its Cl- concentration. Calculate and report the weight % chloride in unknown. Calculate the MnCl+ equilibrium constant. Include error propagation and results analysis. Calculate the relative sensitivity of the electrode to bromide ion and chloride ion. Report as the molarity of bromide ion needed to give the same signal as 1 M chloride ion. DISCUSSION QUESTION: 1. Is the potential of the NaCl + NaBr solution proportional to the NaBr concentration? CALCULATE the error (in millivolts) as the change in potential upon the addition of 1 mL of NaBr solution. DIVIDE this error by the final NaBr concentration in the measured solution and REPORT the error per M NaBr. This electrode is commercially sold as "chloride specific". Is it able to distinguish between Cl- and Br-? 2. What is the slope of the your calibration curve? What is the theoretical value of the slope and how did you derive it? 3. Why must the ionic strength of the solutions be kept constant by the addition of KNO3 ?


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EXPERIMENT 9: GAS CHROMATOGRAPHIC SEPARATION OF AN ALKANE MIXTURE PRELIMINARY READING: Harris 5; Chapt 23. INTRODUCTION:

In this experiment, a mixture of alkanes will be analyzed and the weight percent of one of the components will be determined by quantitative gas chromatography. The goals of this experiment will be to understand the fundamental principles of gas chromatography, to develop skill in instrument use, and to explore the use of an internal standard in quantitative analysis.

Because of limited instrument time, this experiment will be conducted in groups of five to eight persons. The instructor will demonstrate the use of the instrument and the sample injection technique.

You will be provided with a number of alkane (straight-chain) hydrocarbons plus toluene. Some designated person in the team should make a mixture of the alkanes by weight. Each student in a team should make at least one injection into the GC, under varying GC conditions. The Varian gas chromatograph (GC) uses He as the carrier gas and employs a flame-ionization detector. The column may be packed with one of several possible stationary phases. Be sure to make a note of the vital statistics of the column (length, diameter, stationary phase, etc). These are embossed on a metal tag attached to the column.

The purpose of this experiment will be to measure retention times and adjusted retention times. You will relate the retention times to column temperature, measure peak widths, and estimate the number of theoretical plates in the column. NOTE: ALL COMPOUNDS IN THIS EXPERIMENT ARE FLAMMABLE. DO NOT EXPOSE TO OPEN FLAMES.

The most sensitive manipulation in this experiment is the sample injection. Both the GC inlet and the microliter syringe used to inject the sample are fragile. Care is required when making the injection, but at the same time, a smooth, rapid injection is essential for good chromatograms. Correct sample injection technique will be demonstrated by the instructor. If possible, practice your injection technique on unused sample inlets before attempting a chromatogram. A poor injection will require a wait time for the sample to clear from the column before another run can be made. PROCEDURE: Preparation of unknown and reference mixtures. PREPARE a mixture of the alkanes provided (probably dodecane, undecane, dodecane, and perhaps others) and toluene in hexane. They should be no more than 1% by weight in acetone or hexane.

Using the analytical balance and disposable pipettes, WEIGH about 0.2 g of each of the above compounds into a stoppered vial. ADD about 20 g hexane solvent. Mix the solution carefully, if possible without wetting the inside of the vial cap. Keep the vial STOPPERED when not in use.

Data will be produced either with a chart recorder or a computerized data system. Make copies of

the data for all members of your group.


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Data

Individually, measure retention times (in seconds) from the time of injection and peak widths (in seconds) of each peak in your chromatogram. Do this wherever possible (since it may not be possible with some data). Whether you measure half-widths or baseline widths will depend on the graphs you get. Measure peak heights in millimeters. Also measure the time of the leading edge of the solvent peak. Collate these data with the temperature information.

Calculations

Plot the retention times as a function of column temperature for the isothermal runs. Plot retention time as a function of carbon number in the alkanes for all chromatograms.

Calculate the number of theoretical plates of the column using at least two well-formed peaks

from one chromatogram. Calculate the area of each peak (height x base width x 0.5). How does this compare with the mole

fraction of each component in the mixture? In an appropriate literature source, or from the reagent bottles, find the boiling points of the

alkanes in the unknown. Plot the retention time versus boiling point. DISCUSSION QUESTIONS 1. What happens to peak height as the retention time increases? To the peak area? Would you have detected a component which has a boiling point of 300 C? 2. Compute the resolution of two selected peaks from your own data. Use the most clearly defined component and its nearest neighbor. If your own data cannot be used, then use that from another member of your group.


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EXPERIMENT 10.

EXPERIMENTAL DESIGN

PRELIMINARY READING: Harris, 5th Edition, page 856, Section 16-7, and any other pages referenced by the experiment. INTRODUCTION THIS LAB COUNTS DOUBLE. You have to apply what you have learned so far in modifying the experiment as needed. Work in pairs. This lab will take at least two lab periods to complete. You should expect to repeat operations as necessary to get good and reliable results. In this experiment, the assigned problem is the iodometric determination of ascorbic acid (vitamin C; formula weight 176.13) in commercial vitamin tablets by back-titration. (If you normally take vitamin pills, bring 10 or 12 of them to this lab, and you can determine whether there really is as much ascorbic acid in them as it says on the label.) Commercial vitamin pills will also be provided, which have been deliberately selected to introduce analytical difficulties. One mole ascorbic acid reacts stoichiometrically with one mole elemental iodine. The iodine (actually

water-soluble triiodide ion, I3-) is generated by reacting an accurate amount of potassium iodate (KIO3)

with an excess amount of potassium iodide (KI).

KIO3 + 8 KI + 6 H+ 3 I3

- + 3 H2O + 6 K

+

After the iodine has reacted with the ascorbic acid, the excess elemental iodine is titrated with standard thiosulfate solution. This is therefore a back titration. The principle of the back titration is to consume the iodine that is not destroyed by the vitamin C.

I3- + 2 S2O3

-2 3 I

- + S4O6

-2

The indicator in this experiment is aqueous starch solution. Iodine reacts with starch to form a deep violet color. As the iodine is titrated, the color will be consumed and will disappear at the endpoint. The blank in this experiment, remember, is a flask to which iodate, sulfuric acid, and iodide are added, but no ascorbate. It would be the same as your titrations to standardize the thiosulfate solution. The following products are expected from this laboratory: A preliminary determination of the ascorbic acid content of the selected tablets by direct analysis. A more accurate determination using the method of standard additions. Among the analytical difficulties you may encounter in this experiment are: A color from the orange-flavored tablets that will interfere with the iodine endpoint. Insoluble material in the tablets that may or may not have to be removed before titration. Although a complete prelab writeup is required before the first lab session, keep in mind that the unknown problems may necessitate changes in the method as you go.


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PROCEDURE Standard 0.02 M potassium iodate solution.

Do not pre-dry the KIO3 before use. Prepare ~0.02 M KIO3 by accurately weighing ~2 g of the solid reagent and carefully dissolving it in a 500-mL volumetric flask. Starch solution May be provided. If not, prepare the indicator by making a paste of 5 g of soluble starch and 5 mg of HgI2 in 50 mL of water. Pour the paste into 500 mL of boiling water and boil until it is clear or faintly opalescent. Standard thiosulfate solution (back-titrant).

Prepare 0.07 M Na2S2O3 by dissolving ~8.7 g of Na2S2O3.5H2O in 500 mL of freshly

deionized water containing 0.05 g of Na2CO3. Store this solution in a tightly capped amber bottle.

Standardize the solution against your accurate KIO3 solution. Pipet 25.00 mL of standard KIO3 solution

into a flask. Add 10 mL of 0.5 M H2SO4 and 2 g of solid KI and immediately titrate with thiosulfate until the solution has lost almost all its color (pale yellow). Then add a few drops of starch indicator (enough to give a medium violet color), and complete the titration. Do three *good* titrations, ie, if you miss an endpoint, or if the titrations do not agree, then do another. If you add the starch solution too soon, the starch will precipitate and the endpoint will be slow and icky. The yellow color of the iodine should be almost gone before you add the starch indicator. Ascorbic acid analysis Your sample should be enough multivitamin or vitamin C tablets to provide 65 to 130 mg of the vitamin. You can estimate this from the label on the bottle. Tablets and capsules typically contain starch, sugar, flavoring, chalk, and other additives. You may have to estimate the amount of sample to use by weighing an intact tablet and comparing with the listed ascorbic acid content (i.e., if a “250 mg” flavored tablet actually weighs 750 mg, then you would need 300 mg of the ground tablet per sample to get about 100 mg ascorbic acid). Crush the tablet and dissolve in water (try 25 mL to start). This process may not be easy. You may have to use a mortar and pestle, which will be provided, and weigh the resulting powder, rather than the pills. Perhaps simply breaking a tablet into halves or quarters will suffice. Bring enough tablets for some trial and error in the sample preparation process. Heating should be avoided where possible, as ascorbic acid is somewhat unstable.

Add 25.00 mL of the standard KIO3 to the sample of ascorbic acid, plus 10 mL 0.5 M H2SO4 and 2 g of solid KI, and titrate. With some Vitamin C preparations, the sample color may obscure the iodine color, so it may be better to add the starch sooner rather than later.


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It is possible that the yellow iodine color will not persist. This means that there is enough (or more than

enough) ascorbic acid to consume all the iodine. Try adding a second 25.00 mL aliquot of KIO3 before giving up. Be sure to record when you deviate from the standard procedure. Method of standard additions Over the following week, use the method of standard additions to determine interferences in the ascorbic acid analysis. The design of this part of the experiment is left as an exercise for you. A standard solution of pure ascorbic acid must be weighed and dissolved, added in increments to identical samples of the dissolved vitamin pills, and titrated. At least six points should be generated on the standard addition curve. The correct amount of vitamin pill sample and the amounts of ascorbic acid to use should be determined after you have done the first part of the experiment. Ideally, you should judge your increments of added ascorbic acid (FW 176.13) to be 33 to 50% of the amount of ascorbic acid in the sample. For example, if your vitamin pill samples contain 100 mg ascorbic acid each, you should add, say, 0, 50, 100, 150, 200, and 250 mg of pure ascorbic acid. A bonus from standard additions is that calculations are minimized. If the signal is plotted against the amount of added ascorbic acid, the answer may be read directly from the graph or determined using a bet-fit line without having to compute moles of iodate, iodine, or thiosulfate. In this case, signal would be: Signal = (vol of thiosulfate to titrate iodine blank) – (vol of thiosulfate to titrate a sample) CALCULATIONS The calculations are easier if done in equivalents (normalities) rather than moles. You may use either. For this experiment,

1 M I3- = 2 N

1 M ascorbic acid = 2 N 1 M thiosulfate = 1 N Avoid assigning a normality to the iodate solution, as its reaction with iodide is a rare example of a “reverse disproportionation”. Calculate the equivalents of iodine (as triiodide) generated by the 25.00 mL standard iodate solution. After the standardization titration of the thiosulfate solution, calculate the normality of the thiosulfate. When titrating samples, the difference in volume of thiosulfate needed to reach the endpoint is used to calculate the equivalents of ascorbic acid in the sample. Draw a graph, as demonstrated in class, of the results of the standard additions experiment. Present your results as milligrams of ascorbic acid. DISCUSSION QUESTIONS 1. What types of interferences might affect the results? Think in terms of each step of the process, and the fact that iodine is reactive with many classes of compound. What effect would each type of interference have? 2. Can you explain the difference, if any, between the results from direct analysis and standard additions?

CHEMISTRY 247 QUANTITATIVE CHEMICAL ANALYSIS LABORATORY … · Free silver (black stain): Use conc....

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Transcript of CHEMISTRY 247 QUANTITATIVE CHEMICAL ANALYSIS LABORATORY … · Free silver (black stain): Use conc....