281+lab+manual+2011+

Chem 281 Laboratory Manual

Fall 2011

Edited by R. A. Armitage and J.W Guthrie

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Table of Contents

Course Goals and Objectives ................................................................................................................................ 3 Guidelines for Laboratory Notebooks..................................................................................................................... 4 Review Criteria for CHEM 281 Laboratory Notebooks........................................................................................... 5 Chemical Laboratory Safety Rules......................................................................................................................... 6 EMU College of Arts & Sciences: Academic Honesty Policy ................................................................................. 7 Table of Atomic Masses ....................................................................................................................................... 12 Table of Formula Masses....................................................................................................................................... 9 Unknowns and Grading........................................................................................................................................ 12 Student Signature Form For Academic Honesty And Laboratory Safety ............................................................. 12 Waste Disposal .................................................................................................................................................... 12 Instructions for pH Electrode and Buck Flame Photometer ................................................................................. 12 Operation of the Sequoia-Turner Model 340 Spectrophotometer ....................................................................... 12 Experiments: ...................................................................................................................................................... 18

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Course Goals and Objectives

This course is your introduction to quantitative analytical chemistry, the measurement of chemical systems using classical wet chemistry and instruments. Throughout the course, the principles underlying common analytical methods will be discussed in lectures. Selected experiments that demonstrate the application of these principles will be performed in the laboratory. Data and calculations will be recorded in a scientific notebook. CHEM 281 is designed to give you practical experience using both classical and modern analytical methodologies and to provide you with the background theory and principles. As such, it has both a laboratory and lecture component. In the lectures, you will learn the chemical or physical principles exploited for making measurements, review how to calculate quantitative results from analytical data, and learn how to estimate the accuracy and precision of your measurements. In the laboratory, you will put this knowledge into practice by performing various experiments as examples of quantitative analysis. It is important to keep an accurate and reliable scientific notebook during any experiments you perform. In many industries, an individual is legally responsible for the accuracy of the data recorded in the notebook. Therefore, the notebook you keep as a record of your observations and measurements will be graded. In addition, you will use the fundamental information learned during lecture and practical experience gained in the laboratory to write coherent scientific reports. Technical writing, based on being logical, concise and precise, is desirable for any sort of formal paper that contains conclusions based on numerical data. To be successful in this class you must understand and be able to apply the lecture material to operation of the instruments. A common mistake for students is to concentrate on either the lecture or laboratory material exclusively. Remember, half of your grade is from each component.

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Guidelines for Laboratory Notebooks

You should be already familiar with the practices of keeping a scientific notebook. Those practices will be reinforced here. As in other classes, you will record the experimental results in ink in an orderly, neat manner in a hardbound notebook. Your laboratory notebook must be quadrille or ruled, with consecutively numbered pages. You may number the pages by hand prior to the start of the first laboratory session if necessary. There are several reasons why it is important to keep a thorough, accurate, and up-to-date notebook including:

1. Your notebook will help you to organize your laboratory work and will be a permanent record of your data.

2. Your notebook may help you or the instructor identify problems in your laboratory technique or calculations, and thereby save you from wasted time or a poor grade.

3. An accurate notebook is essential for many research and technical positions. Be as detailed as possible when writing information in the notebook: in general, a notebook cannot be too detailed. Write units for all numbers where appropriate. Write an explanatory note for any numbers – for example, “concentration of standardized NaOH solution = 0.1455 M” not just “NaOH = 0.1455” – the meaning of any information becomes less clear with time. Write all important numbers in your notebook. Do not try to remember numbers. It is good practice to date and initial each page as it is completed. Do not leave blank pages that you intend to fill in later but instead make a note in the notebook that refers to the relevant page. Tables, graphs, spectra, chromatograms and other papers should be permanently fixed into the notebook. A good way to do this is to copy and/or reduce a figure and glue it to the page. Loose papers, papers folded into the notebook or those that protrude from the edge tend to eventually become detached and lost. Do not remove any pages from your notebook. Mistakes should be crossed out with a single line: do not use correction fluid. On the cover of the notebook, record your name, address, course, and section number. You should reserve several pages at the beginning of your notebook for the table of contents. Include the title of each experiment, the date it was performed, and the page numbers where it can be found in the notebook. Be sure to keep the table of contents up to date throughout the semester. By the end of the class, your notebook will probably contain spills, splashes, crossed-out sections and revisions. This is normal. You will not be penalized as long as you have followed the other guidelines outlined here. The information in your notebook should be sufficient that you can reconstruct the whole experiment but it does not have to be as complete as your written report. Lab notebooks will be checked for proper formatting early in the semester. They will be collected approximately every two weeks. At least one experiment must be complete each time the notebook is graded.

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Review Criteria for CHEM 281 Laboratory Notebooks Laboratory notebook entries must be raw data; they are not typed. Only graphs and printouts of data (chromatograms, etc.) should be produced on the computer and pasted into the notebook. The following sections must appear in the lab notebook for each experiment:

• Title: The title of the experiment should be centered at the top of the first page of the report. (0.25 pts)

• Purpose: Succinctly state the purpose of the experiment. The word “analyze” is not appropriate here. “To analyze for chloride in an unknown” is wrong; “To determine the percent chloride in an unknown salt sample using gravimetry” is correct. Be sure to state the method. (0.5 pts)

• Procedure: Do NOT copy the procedure from the lab manual into your notebook! Cite the coursepack as the source for the basic procedure, including the page numbers. Note everything that you do, as you do it. (0.25 pts)

• Data: This section will include tables of all measurements and observations made during and for the experiment. List all sample masses, volumes used, etc. Identify each instrument used by letter where appropriate. (2 pts)

• Calculations/graphs: A sample calculation for each data manipulation must be shown with all units. Every calculation to determine appropriate masses or volumes of reagents must appear in the notebook. All graphs must be created by computer. Hand drawn graphs will not be graded. Graphs must be appropriately labeled and titled. Each axis must be labeled and include the units (for example, Concentration (ppm)). The title must be informative rather than simply restating the axis labels; “Calibration Curve for Fe” is appropriate, but “Concentration vs. Absorbance” is not. Calibration curves must include a linear regression best fit line and equation. This can be determined by graphing software (such as Excel) or using a graphing calculator. (4 or more pts, depending on experiment)

• Result: This consists of a single sentence reflecting the statement of purpose, such as “Unknown X contained 34.35% chloride.” The result is a single number, either the mean or median of your calculated values. It is this value alone that will be graded for accuracy and/or precision. (0.5 pts)

• Conclusion: Summarize what you learned from the experiment and answer any questions from either the coursepack or the handout. Answers must be in complete sentences that make clear what the original question was. The conclusion should also contain comments on your result: why you chose to report the median instead of the mean, why your standard deviation was so large, etc.(2 pts)

• Date: All labs must start with a date. If the lab is completed over several days, each day’s work should start with this entry. (0.5 pts)

Total: 10 points per notebook write-up + additional accuracy points

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Chemical Laboratory Safety Rules

1. MIOSHA approved eye protection (goggles) must be worn at all times when you are in the laboratory unless the instructor states otherwise. The wearing of contact lenses is

not permitted.

2. Never work alone in the laboratory.

3. Report all accidents and/or injuries to your instructor. Medical treatment/consultation is required (Snow Health Center-EMU Campus) for any injury. Tell your instructor you are going for treatment and give him/her your student number so Heath Services can be notified.

4. Never carry out unauthorized, unplanned, or non-scheduled experiments. Discuss any unusual work with your instructor prior to doing it.

5. Never eat, drink, or taste anything (food or chemicals) while you are in the laboratory. Don’t place fingers, pencils, pipettes, etc. in your mouth.

6. Confine long hair and sleeves when working. Wearing a lab coat or apron is the recommended protection. Wear closed shoes, not sandals in the lab.

7. Do not throw chemical waste in the sink or in the waste baskets. Always consult your instructor for the proper chemical waste disposal procedure.

8. Wear appropriate gloves when working with hazardous liquids, solids, or solutions.

9. Always use a suction bulb (never use your mouth) when filling a pipette.

10. Do not force glass tubing and/or thermometers into rubber stoppers – always lubricate the hole in the stopper and protect your hand with a towel when inserting the glass.

11. Never use an open flame near flammable solvents.

12. Clean up all spills immediately.

13. Do not test odors by direct inhalation from the container.

14. Wash chemical contacted areas of skin or eyes with water for 15 minutes.

15. During the first day of laboratory, locate all emergency and safety equipment that you may need to use. This includes drench showers, eye wash, fire extinguisher, and/or fire blanket. Locate the nearest emergency exit and telephone.

16. Be aware of your ‘right to know’ all safety information contained in the Material Safety Data Sheet (MSDS) for any chemical. These are available through the manager of Chemical Services Operations, in room B125 Mark Jefferson.

17. IN EMERGENCY: Telephone 911 for fire, ambulance, medical assistance, or police. Telephone 7-1222 for campus security. Telephone 7-0106 or 7-0107 for Chemistry Office

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EMU College of Arts & Sciences: Academic Honesty Policy

Education involves the search for the truth. Therefore, honesty and personal integrity are values highly esteemed by the academic community. They are ingredients essential to the cooperation and communication among students and faculty needed for progress. The following honesty policy is intended to clarify the College’s expectations for its students and to provide guidance in specific situations.

A. Definitions of Cheating

Cheating includes, but is not limited to the following:

1. Unless specifically told otherwise by the instructor, examinations, test papers, laboratory reports, computer programs, and graded homework assignments are to be completed independently by each student. Looking at another person’s paper or communicating with another person while working on an examination, test, laboratory report, computer program, or graded homework assignment is cheating.

2. Unless specifically told otherwise by the instructor, having books, notes, written material, or any means of accessing other than what is specified by the instructor readily available during an evaluation is considered cheating.

3. Obtaining or distributing exams in advance of their use is considered cheating.

4. It is cheating to represent as your own anything obtained from published materials or from another person. All source material must be appropriately acknowledged. Directions for proper acknowledgement of sources can be acquired in ENG 121 or found in the MLA Style Sheet, The Chicago Manual of Style book, and other style guides for specific disciplines. In chemistry, the ACS Style Guide is often used.

5. To allow another to represent your words or ideas as his/her own is cheating.

6. To use work from one class for another without prior approval of your instructor is cheating.

B. Penalties for Cheating

The cheating penalties will range from a minimum of receiving a zero grade on the experiment, computer program, report, paper, performance, project, examination, or test involved, to a maximum of both receiving an E in the course in which cheating occurred and reporting of the incident to the Department Head and Dean of Students for possible further disciplinary action including suspension or dismissal from the University.

C. Appeals

A student who has been found by his/her instructor to have engaged in cheating in a course may appeal the final grade received in that course. The appeal shall be made in accordance with Eastern Michigan University’s Grade Grievance Procedure, which appears in the Undergraduate Catalog and the Graduate School Catalog. In situations where cheating incidents are also referred to the Department Head and Dean of Students for possible disciplinary action, all proceedings and appeals shall be conducted in accordance with Eastern Michigan University’s Conduct Code and Judicial Structure for Students and Student Organizations, which appears in the Undergraduate Catalog and the Graduate School Catalog.

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International Atomic Masses Based on 12C=12 (Numbers in parenthesis indicate mass number of most stable known isotope)

Element

Symbol

Atomic No.

Atomic Wt.

Element

Symbol

Atomic No.

Atomic Wt.

Actinium Aluminum Americium Antimony

Argon Arsenic Astatine Barium

Berkelium Beryllim Bismuth Boron

Bromine Cadmium Calcium

Califonium Carbon Cerium Cesium Chlorine

Chromium Cobalt Copper Curium

Dysprosium Einsteinium

Erbium Europium Fermium Fluorine

Francium Gadolinium

Gallium Germanium

Gold Hafnium Helium

Holmium Hydrogen

Indium Iodine Iridium

Iron Krypton

Lanthanum Lawrencium

Lead Lithium

Lutetium Magnesium Manganese

Mendelevium

Ac Al

Am Sb Ar As At Ba Bk Be Bi B Br Cd Ca Cf C

Ce Cs Cl Cr Co Cu Cm Dy Es Er Eu Fm F Fr Gd Ga Ge Au Hf He Ho H In I Ir Fe Kr La Lw Pb Li Lu Mg Mn Md

89 13 95 51 18 33 85 56 97 4

83 5

35 48 20 98 6

58 55 17 24 27 29 96 66 99 68 63

100 9

87 64 31 32 79 72 2

67 1

49 53 77 26 36 57

103 82 3

71 12 25

101

(227) 26.9815

(243) 121.75 39.948

74.9216 (210)

137.34 (247)

9.0122 208.980 10.811 79.909 112.40 40.08 (251)

12.01115 140.12

132.905 35.453 51.996

58.9332 63.54 (247)

162.50 (254)

167.26 151.96 (257)

18.9984 (223)

157.25 69.72 72.59

196.967 178.49 4.0026

164.930 1.00797 114.82

126.9044 192.2

55.847 83.80

138.91 (256)

207.19 6.9417 174.97 24.312

54.9380 (258)

Mercury Molybdenum Neodymium

Neon Neptunium

Nickel Niobium Nitrogen Nobelium Osmium Oxygen

Palladium Phosphorus

Platinum Plutonium Polonium Potassium

Praseodymium Promethium Protactinium

Radium Radon

Rhenium Rhodium Rubidium

Ruthenium Samarium Scandium Selenium

Silicon Silver

Sodium Strontium

Sulfur Tantalum

Technetium Tellurium Terbium Thallium Thorium Thulium

Tin Titanium Tungsten Uranium

Vanadium Xenon

Ytterbium Yttrium

Zinc Zirconium

Hg Mo Nd Ne Np Ni Nb N

No Os O Pd P Pt Pu Po K Pr Pm Pa Ra Rn Re Rh Rb Ru Sm Sc Se Si Ag Na Sr S Ta Tc Te Tb Tl Th Tm Sn Ti W U V

Xe Yb Y

Zn Zr

80 42 60 10 93 28 41 7

102 76 8

46 15 78 94 84 9

59 61 91 88 86 75 45 37 44 62 21 34 14 47 11 38 16 73 43 52 65 81 90 69 50 22 74 92 23 54 70 39 30 40

200.59 95.94

144.24 20.183 (237) 58.71

92.906 14.0067

(255) 190.2

15.9994 106.4

30.9738 195.09 (244) (209)

39.102 140.907

(145) (231) (266) (222) 186.2

102.905 85.47

101.07 150.35 44.956 78.96

28.086 107.870 22.9898

87.62 32.064

180.948 (97)

127.60 158.924 204.37

232.038 163.934 118.69 47.90

183.85 238.03 50.942 131.30 173.04 88.905 65.37 91.22

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TABLE OF FORMULA MASSES

AgCl 143.32 AgNO3 169.87 CuSO4 159.60 Fe(NH4)2(SO4)2 ▪ 6 H2O 392.14 HC2H3O2 (acetic acid) 60.05 HCl (hydrochloric acid) 36.46 HClO4 100.46 HNO3 (nitric acid) 63.01 H3PO4 (phosphoric acid) 98.00 H2SO4 (sulfuric acid) 98.08 KCl 74.56 K2CrO4 194.20 K2Cr2O7 294.19 KHC8H4O4 ( KHP) 204.23 KMnO4 158.04 NaCl 58.44 Na2H2EDTA ▪2 H2O 372.23 NaOH 40.00 NH3 17.03 NH4Cl 53.49 (NH4)2SO4 132.14

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Unknowns and Grading

In each experiment you are given a solid or solution “unknown” and through experimentation are required to determine the quantity of a specific compound or element present in the “unknown” sample. Specific information for requesting the unknown for each experiment is found on page 11. Some solid unknowns must be dried in the oven prior to use. Those that have been dried should be stored in a desiccator. Store only solids in your desiccator, never liquids! Each “unknown” is numbered and randomly distributed by the teaching assistant who maintains a list of the distribution. Grading is based on a comparison with the “true” value for the sample as listed by the manufacturer with the experimentally determined one. It is essential, therefore, that you do not “swap” or mix your sample with anyone else’s as each sample is uniquely different in composition.

Deadlines: A schedule of due dates is included in the lab syllabus and is strictly adhered to. One laboratory point per laboratory period that a report is late will be lost, unless lateness is due to an excused absence (as determined by the instructor) and must be discussed with the instructor as soon as possible. No credit will be given for reports that are late more than three laboratory periods. Accuracy/Precision Grading: The accuracy of your results form the basis for the Accuracy/Precision grade for each experiment. Therefore, your experimental technique and data work-up is of significant importance. Each experiment is graded first on the requirements for the laboratory notebook described on page 5, and also, in some cases, on a basis of 10 or 15 points for accuracy and/or precision. The accuracy of the results, i.e., deviation from the true value, determines the score out of 10 possible points for accuracy. Repeating an Experiment: If you are dissatisfied with the grade you earned on an experiment you have the option of repeating the experiment with a different unknown. You may redo an experiment, but only for 80% of full credit. Thus, if on the KHP titration you receive zero points the first time, and twelve points on the second try, the score recorded for you would be (.8)(12) = 9.6 points. Of course, if your score on the redone experiment is lower than that on the original, we will just quietly forget that you redid the experiment, and let the original score stand.

Organization

You should always plan your work before coming to the laboratory. Time spent

in the laboratory puzzling over the laboratory instructions is wasted time. Thus, it seems clear that you must read your manual most carefully before coming into the laboratory. The analyses you will perform are not impossible, but can be botched most regally if you plunge into them without a clear notion of what you are doing. It is also useful to read ahead to the next experiment as you might find yourself with unexpected time to prepare for it.

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Unknown Requests:

Experiment: Submit the following and place in the prep room OR oven as

instructed. Include completed request form. Volumetric KHP Place 2 clean weighing bottles with caps (label one U and one S

with a Sharpie Marker) in a beaker, covered with a watch glass in the oven to dry. Samples will be added by GA. Dry samples for NO MORE than 2 hours. KHP starts decomposing with longer drying.

Macro Na Place a black-capped flint glass bottle, labeled on side using

marking pen, on the prep counter. A liquid unknown will be given. Also, place a clean weighing bottle, with cap, in a beaker, covered with a watch glass in the oven to dry. Solid KHP will be added by TA and returned to the oven for no more than 2 hours.)

Trace Na & K 1) For unknown: place a weighing bottle with cap in a beaker on the prep counter – a liquid sample will be given to you.

2) For standard Na and K: Place 2 weighing bottles with caps (label one Na & one K with Sharpie), in a beaker covered with a watch glass into oven to dry. Solids will be added by TA and returned to the oven

Back Titration of Aspirin tablets are located on the counter in the prep room. Aspirin Tablet There is no other unknown to request. Simultaneous Place a weighing bottle with cap in beaker and place on prep Co and Cr counter. The GA will provide a liquid unknown. Fe in Vitamin Place a clean 100-mL volumetric flask, , with cap, in a beaker on

the prep counter. GA will add exactly 10.00 mL of liquid unknown. Vitamin tablets are located on the counter in the prep room.

Water Hardness Bring at least 1L of tap water form home. Bring it in a plastic bottle.

(Place completed request form in beaker even if placed in oven)

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Student Signature Form For Academic Honesty and Laboratory Safety

I have read the ACADEMIC HONESTY POLICY of the Eastern Michigan University College of Arts and Sciences. I understand the policy, and will abide by it.

I have read the SAFETY RULES for Eastern Michigan University Chemistry Labs. I understand and will observe these rules. I understand that I am responsible for conducting myself in a safe manner, becoming aware of and informed about special hazards of technique, apparatus, or chemicals in the chemical laboratory, and will conform to any additional safety instructions presented orally or in writing by the instructor or contained in posted instructions or safety memoranda that are distributed.

I am aware that I have a ‘right to know’ all safety information contained in the manufacturer’s Material Safety Data Sheet (MSDS) for any chemical used in this course. I can obtain this information by requesting a copy of the MSDS from the main departmental office in room 225 Mark Jefferson.

Signature: ________________________________ Date: _______________________

Print Name: _____________________________________________________________

Student Number: _________________________________________________________

Course Number: CHEM 281

This form must be signed and returned to the instructor before you may begin work in the laboratory.

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Waste Disposal The proper disposal of waste is a responsibility of each student. Many solids and solutions are harmless and may be disposed of by washing down the sink. If discarding an acid or base solution down the drain it is essential to keep them separated. Make sure to flush each down with copious amounts of water. Some substances must be discarded into a properly labeled waste bottle found in one of the hoods. It is important to check the label BEFORE adding any new substances to the container. Do not overfill the containers. The table below shows the correct waste disposal method for substances encountered in the lab.

Graphing Technique

In the quantitative analysis laboratory, a great deal of emphasis is placed on proper graphing techniques. There are several graphing programs available (Excel, Graphical Analysis, etc.) which greatly simplify this process, however; the user must have a good understanding of the proper information to provide and the correct placement of said information. On any x-y graph, the x-axis represents the independent variable (the quantity you set) and the y-axis is used for the dependent variable (that which you have measured). Both axes must be properly labeled, using the appropriate units. (For example, “Concentration” is not a sufficient label. The units, such as molarity or ppm etc., must be included.) If the data points show a linear correlation, include the equation of the line and the correlation coefficient (calculated using linear regression) along with the graph. The equation should be used to draw the line that best fits the data points as well as to determine the x value given any measured y value and vice versa.

Sink

Labeled Waste Jars in Hood

dilute acids – lab sink concentrated H2SO4 - discard in a hood sink Concentrated HCl – discard in a hood sink

K2Cr2O7

NH3, NH4OH -discard in a hood sink Cr3+, Co2+

KHP CuSO4, Cu(EDTA)2

Fe(NH4)2(SO4)2·6 H2O

Fe(o-phen)3, and orthophenanthroline

KMnO4

NH2OH·HCl

CaC2O4•H2O, urea

NaCl, KCl

EDTA Aspirin

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Instructions for a pH Electrode

All pH electrodes are fragile and expensive. Do not bump or scrape them with your beaker. Do not rub them with a paper towel. They must be hydrated before using and gently dried with a Kim Wipe before inserting into a solution to be measured. All solutions measured must be at the same temperature as the buffer solutions used to calibrate the instrument and should be stirred. The electrodes should not be immersed so far that the level of the internal electrolyte is below the level of the external solution. If the electrodes are new or have dried out, several hours of soaking will be necessary before stable readings are possible. The pH meter must be calibrated with buffer solutions that bracket the range of the pH of the solution to be measured. Specific calibration instructions for each type of meter are located on the wall above the instruments.

Instructions for Use of the Buck Scientific Flame Photometer

The flame photometer measures light emitted (not absorbed) by excited electrons. The sample is sprayed (nebulized) into a flame with sufficient energy to excite electrons into a higher energy level. The light that is emitted when the electrons return to the ground state is selectively filtered and measured by the detector. Different filters are used for each element.

1. Switch on electrical power by depressing the "power" switch on the front of the photometer. The electronics should be permitted 15 minutes to warm up before using.

2. Ensure that the drain tap in the back of the instrument is pushed fully

down in its clip. Ensure that the drain trap has solution in it and that no air locks are present. If necessary, purge by adding deionized water and allowing surplus to drain away.

3. Fully open the "fuel adjust" valve by turning fully counterclockwise. To

avoid damaging the valve it should not be forced.

4. Open the air cylinder. Make sure there is more than 500 psi pressure in the tank and that the exit pressure is at the level marked. Ensure that air is present by listening for the hissing created as it passes through the nebulizer.

5. Open the natural gas valve located on the bench. Depress the ignition

switch and hold down several seconds, watching the "FLM" indicator in the display window. When this indicator is illuminated the flame is lit and the ignition switch can be released. This usually requires 3-4 attempts.

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CAUTION: Flames can extend over the top of the chimney. Do not

lean over it until you have adjusted the fuel. Make sure to tie back any long hair to prevent injury.

6. Turn the fuel level down until only a circular blue flame is visible at the

bottom of the chimney. Set the "filter select" control to the desired position.

7. Make sure liquid is aspirated into the flame at all times. Solution is

aspirated into the flame via narrow tubing. Wipe the tubing gently with a Kim Wipe before inserting into a solution. Do not touch the portion of the tubing that will enter the solution to avoid transferring ions (particularly sodium from your fingers) into the solution. Occasionally the tubing gets blocked, so routinely check to make sure solution is flowing through the tube. Aspirate deionized water and set the output to zero using the “Blank” knob.

8. Set the upper limit by aspirating the highest concentration solution and adjusting the “fine” and “coarse” control knobs until a reading of 700 is obtained.

9. Measure the emission of each solution, starting with the lowest

concentration. In between each solution, aspirate water into the flame and observe that the emission reverts to zero. Do not readjust the settings while measuring a series of solutions. Carefully wipe the tubing with a KimWipe in between each solution. Make sure the tip of the tubing is submerged into the solution to be measured and that the solution has been thoroughly mixed beforehand.

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Operation of the Sequoia-Turner Model 340 Spectrophotometer A. Turn on the instrument, allow the electronics a 15 minute warm-up. B. Turn mode switch to TRANS. Adjust instrument to desired wavelength. C. Insert appropriate filter into position into the slot on top of the instrument. D. Press and hold ZERO SET button while adjusting ZERO knob until display

indicates 0.0. Release ZERO SET. E. For absorbance measurements:

a. Insert a cuvette with water or other “blank” solution b. Set MODE switch to ABS. c. Adjust 100%T/OA COARSE knob (larger outer knob) to approximately 0. d. Adjust 100%T/OA FINE knob (inner knob) to exactly 0.000. e. Replace cuvette with sample cuvette and read absorbance from digital

display. How to Manipulate Cuvettes:

1. Cuvettes, or cells for the spectrophotometer, are rather like test-tubes in appearance. In reality, they are rather more carefully made.

2. It is necessary to insure that cuvettes are carefully used. Thus, the following rules must be diligently observed, if the student desires good analytical results. a. Never handle the lower portion of a cuvette, because light must pass

through it. The result of fingerprints on a cuvette where light is passing through is a scattering of the light resulting in a lower absorbance measurement.

b. Rinse the clean cuvette out with several portions of the solution of interest

before taking a measurement.

c. Gently wipe the cuvette with a Kim Wipe to remove any fingerprints or smudges. Be sure no bubbles cling to the inside walls (they will also scatter light), if they do gently tap the outside of the cuvette while hold it at an angle.

d. When inserting a cuvette into the sample holders always insert with the

index line facing the front of the instrument, to avoid scratching, then line up the index lines on the cuvette holder and cuvette exactly. Always insert the cuvette exactly the same way to avoid any discrepancies due to variability in the glass wall of the cuvette.

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EXPERIMENTS

Group 1 Experiments (door bench) Accuracy/ Precision Points

Notebook Points

Statistical Analysis (handout) 0 10 Determination of KHP 10 10 Potenetiometric Analysis of a Weak Acid 10 10 + 2 for graphs Determination of Macro Na in Water 10 10 Determination of Water Hardness 6 10 Determination of Trace Na by Flame Photometry 4 10 + 2 for graphs Fe in a Vitamin Tablet 10 10 + 2 for graphs Simultaneous Co and Cr 5 10 + 5 for graphs Back Titration of an Aspirin Tablet 10 10

Group 2 Experiments (middle bench) Accuracy/ Precision Points

Notebook Points

Statistical Analysis (handout) 0 10 Fe in a Vitamin Tablet 10 10 + 2 for graphs Simultaneous Co and Cr 5 10 + 5 for graphs Back Titration of an Aspirin Tablet 10 10 Determination of KHP 10 10 Potenetiometric Analysis of a Weak Acid 10 10 + 2 for graphs Determination of Macro Na in Water 10 10 Determination of Water Hardness 6 10 Determination of Trace Na by Flame Photometry 4 10 + 2 for graphs

Group 3 Experiments (window bench) Accuracy/ Precision Points

Notebook Points

Statistical Analysis (handout) 0 10 Determination of Water Hardness 6 10 Determination of Trace Na by Flame Photometry 4 10 + 2 for graphs Fe in a Vitamin Tablet 10 10 + 2 for graphs Simultaneous Co and Cr 5 10 + 5 for graphs Back Titration of an Aspirin Tablet 10 10 Determination of KHP 10 10 Potenetiometric Analysis of a Weak Acid 10 10 + 2 for graphs Determination of Macro Na in Water 10 10

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Experiment 1 Statistical Analysis of Data

Introduction There are few, if any, absolute values in science. Most values are stated with reference to a standard that has been arbitrarily selected. In quantitative chemistry, many results are expressed as percentages that are obtained by calculations using values that have been determined by physical measurements. However, such measurements involve many possible errors, which can be properly evaluated only by a statistical approach. In analytical chemistry, the term “precision” is used to describe the reproducibility within a set of values. It can be defined as the level of agreement among several measurements, observations, or other types of quantitative results obtained under identical operating conditions. “Accuracy” has been used to describe the degree of agreement between a particular experimental value and the correct or “true” value. Although there is no “true” value (as no measurement is completely accurate), reported results are judged with respect to those measured values which have been determined by skilled scientists using instruments of the best quality. To obtain a true value, it is necessary to have a number of observers and several instruments, both observers and instruments being interchanged to produce the final “average” or “true” value. In an elementary quantitative analysis course, a student’s chosen result is usually checked against a known value which has been previously established. In beginning courses, it is customary for the student to perform analyses in triplicate (or higher), using equal portions of the same sample. While the student cannot establish the accuracy of his results before reporting a value, he can determine the precision of his analytical observations. If the replicate results show good agreement, more reliance can be placed on the average value than on any one or two results. In an elementary course, a precision that is less than 2 parts per thousand (ppt) may be the result of a coincidence. On the other hand, a precision that is more than 5 ppt (involving the analysis of a homogeneous substance) indicates unreliable laboratory practice, and the values are subject to question. The task of an analytical chemist goes beyond that of correctly performing the manipulations and readings required in any given procedure. To make the work meaningful, one must also:

1) properly record and correctly calculate the results of each analysis 2) determine the best (most appropriate) value to report for a series of values 3) evaluate the results and establish the probable limits of error.

A simple statistical treatment of the data obtained in an experimental procedure can be carried out to complete the above requirements. In addition, the chemist must examine each measurement to determine the number of significant figures in each reading. The value having the smallest number of significant figures determines the number of significant figures in the final reported value.

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Although there is no doubt as to the mathematical importance of an infinite

population distribution of values, as far as real laboratory work is concerned, measurements are made on a finite number of samples. In most quantitative analytical procedures, practical considerations severely limit the number of trials which can be run.

Statistical treatment of experimental data employs a variety of terms, dealing with

both accuracy and precision parameters. The central tendency of a group of results is simply that value about which the individual results tend to “cluster.” For a finite number of measurements, it is the mean, designated as x, calculated as:

The median of an odd number of results is simply the middles value when the results are listed in order; for an even number of results, the median is the average of the two middle ones. In a truly symmetrical distribution, the mean and the median are identical. The average deviation, d, is often expressed as a measure of variability within a data set, although it is not very significant for a small number of observations. To calculate the average deviation, the differences between individual results and the mean, regardless of sign, is added up and divided by the number of values, as:

Often, the relative deviation is expressed relative to the mean, either as a percentage or in parts per thousand, as:

The standard deviation, s, is much more meaningful statistically than is the average deviation. The standard deviation, which may be thought of as a root mean square deviation of values from their average, is calculated as:

The spread or variability of a series of data values is often useful in evaluating an analytical technique. For a finite number of values, the simplest measure of variability is the range, which is the difference between the largest and smallest value. Unfortunately, one “wild” result often renders the range meaningless.

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Often, one result in a set of replicate measurements seems to be out of line with the others, and an analyst must decide whether to exclude this result from further consideration. Various criteria for rejection have been developed, including several rather simple and straightforward statistical tests. If one result in a small group (n = 4-8) is questionable, the “2.5 d” test may be applied. Application of this test is as follows:

1. Compute the mean and average deviation of the “good” results. 2. Find the deviation of the suspected result from the mean of the “good” ones

calculated in step 1. 3. If the deviation of the questionable result from the “good” mean is at least 2.5

times the average deviation of the “good” results, then reject the suspected result. Otherwise, retain it.

The “Grubbs” test also is applied to questionable values, but utilizes a lower level of confidence. For small data sets (n = 3-6), the Grubbs ttest allows rejection of data that deviate widely in one extreme. Application follows:

1. Calculate the average and standard deviation of the results, including the

questionable one. 2. Find the absolute value of the differnce between the questionable value and

the average and divide it by the standard deviation (page 92 in the text). 3. Consult a table of critical values of G for rejection of an outlier (p 92 of the

text) If the computed value of G is greater than the value in the table, the result can be discarded. If not, it must be retained.

If more universally applicable tests (or different confidence levels) are desired,

the student “t” test may be applied (p 86). This test states that the population mean, or true value lies within a certain region centered at the mean, as defined by the equation:

Where “t” is a constant corresponding to a certain probability level and the degrees of freedom associated with the statistical compilations. Thus, for a certain value to be a valid member of the data set, it must lie within the range about the mean given by the term “ts/n½”. If it lies outside this limit, it is not statistically valid at the confidence level used, and may be rejected. Irrespective of the statistical test used to check a questionable value, caution should be followed in rejecting values. If the number of replicate trials is large, rejection of one value is not critical. However, if the number of trials is small, a real dilemma can arise: the divergent value exerts a significant effect upon the mean of the complete set,

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while at the same time there are insufficient data to permit a real statistical analysis of the status of the suspected result. A sort of compromise between outright rejection and the retention of a suspected value is sometimes recommended. An analyst might prefer to report the median of all the results rather than a mean either with or without the deviant value. Some authorities recommend that the highest and lowest values both be rejected, and the mean of the remaining values reported as the most appropriate value. In any event, the ultimate decision on what value to report for an answer must be made by the analyst. Experimental Procedure Mass measurements, analytical balance Using the weighing procedures described by the instructor, accurately (to ± 0.1 mg) weigh each of the coins in one of the sets by the analytical balances. Record the balance number on the instrument used to make the measurements. Use forceps to handle the coins, remembering that fingerprints leave measurable traces on most substances handled. Place the weighing paper on the balance pan, and accurately weigh the paper. Record the mass of the paper, and then place each coin in turn on the paper and record the total mass. Calculate the mass of the coin by subtracting the paper mass from the combined mass (mass by difference). The paper is used to make it easier to place and remove coins from the balance pan. Label each coin/paper mass, using the mint date or other identifying marks on the individual coins. After all coin masses are calculated, list the masses by rank, smallest to largest, and label this table Data Set 1. If there is more than one outlier, find additional pennies in your “close” data range so you have at least 4 pennies with similar masses. Calculate the mean, median, and range of the data set. Round off each coin mass to three significant figures, and make a second table labeled Data Set 2. Calculate the average and standard deviation of this set. Using the “Grubbs” and “t” tests, determine the validity of any questionable member(s) of Data Set 2. Apply these tests (95% confidence level) only to those values which, upon preliminary inspection, appear suspect. Using the statistically analyzed values of Data Set 2, determine the most appropriate coin mass, and report the value in the appropriate units and tolerances. If there are one or more outliers, carefully look at the mint date and see of you can determine which year(s) have different masses. Consult the Wikipedia article on the Canadian penny (http://en.wikipedia.org/wiki/Penny_%28Canadian_coin%29) to see why and include this information in your notebook.

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Volume measurements Using a buret stand and clamp, set up a buret filled with deionized water. Collect

a set of 25.00-ml volumetric flasks either from the instructor or from your equipment drawer. Make sure the flasks are clean and dry before beginning this part of the experiment. Determine the volume (± 0.01 ml) of each of the volumetric flasks, by filling each in turn to the mark on the flask neck (exactly!). Record the initial and final volume readings on the buret to determine the volume of water used to fill the flask. Refill the buret after each measurement to insure the same region of the buret is used in filling each of the flasks.

List the volumes in order, smallest to largest. Calculate the mean, median, and

range of the data set. Calculate the average and standard deviation of the group of values. Using the “Grubbs” and “t” tests, determine the validity of any value that appears

questionable. Report the most appropriate value for the measured volume of the volumetric flask, using proper units and tolerances.

Calculate the relative error in the volumetric measurements, assuming the true

volume is 25.00 ml. Use the equation below to calculate relative error in parts per thousand, as:

When finished with a set of volumetric flasks, dry them (do not drain them by

standing them upside down!) and either return them to the instructor or replace them in your drawer.

Note: The two sections of this experiment can be performed in any

order, but both parts MUST be carried out.

Source: Quantitative Analysis Laboratory Manual, Chemistry Department, Thiel College, 1989.

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Experiment 2 Determination of Potassium Hydrogen Phthalate (KHP)

Introduction The best solutions for acid-base titrations contain strong acids and/or bases. In addition, such reagents must be fairly soluble, stable, relatively inexpensive, and not produce side reactions when used in acid-base neutralizations. Only a few acids and bases meet these criteria. The most satisfactory acids are hydrochloric acid (HCl) and sulfuric acid (H2SO4); the most frequently used bases are sodium hydroxide (NaOH) and potassium hydroxide (KOH). However, strong bases tend to react with glass containers and to absorb CO2 when exposed to the atmosphere. Consequently, certain precautions are necessary in the preparation, storage, and use of strong base solutions. In an acid-base neutralization of a weak acid, a strong base titrant must be used. The concentration of the base solution is usually determined experimentally by measuring the volume of base which is chemically equivalent to a known mass of a pure acidic substance. This process is called a direct standardization, and the pure analyte chemical is called a primary standard. A primary standard is usually a solid substance at least 99.95% pure, the composition of which is not affected by drying. It should be relatively soluble in water, and it must react quantitatively with the solution to be standardized. Ideally, its equivalent or molecular mass should be fairly high in order that a large sample may be used without requiring more than a buret full of titrant solution in the neutralization process. Potassium hydrogen phthalate or KHP (KHC8H4O4; 204.23 g/mol) is the most widely used primary standard for the standardization of base solutions. Its structure and reaction with a strong base such as NaOH is given in the equation:

The net reaction, of course, is: HP- + OH- = P2- + HOH. Substances that are used to a lesser extent include: potassium hydrogen tartrate (KHC4H4O6), potassium biiodate (KH(IO3)2), and sulfamic acid (HSO3 · NH2). By definition, the equivalent

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weight/mass of the acid may be calculated by dividing the formula/molecular weight of the acid by the number of protons (or H+) furnished in the reaction. Since KHP yields only one proton per molecule, the equivalent weight is equal to the formula/molecular weight (the hydrogens on the organic ring do not react). In the standardization of NaOH titrant solution, samples of potassium hydrogen phthalate are accurately weighed and dissolved in deionized water. In this medium, the KHP dissolves and dissociates to produce free potassium ion and hydrogen-phthalate, as: KHP → K+ + HP-. The HP- ion is then capable of acting as a weak acid, and, under appropriate conditions, may release the hydrogen, as: HP- ↔ P2- + H+ . In an acid-base neutralization reaction, one mole of acid will react with one mole of base. Similarly, the neutralization may involve millimoles of acid and base. A general relationship may be written as: # millimoles = grams/millimolecular weight = molarity x ml

You will be determining first the molarity of the NaOH solution- standardization of the base. Then, you will determine the amount of pure KHP present in an unknown which is a mixture of KHP and a salt. The general formula for calculating the % purity of an acid is given by:

OR

In the titration of impure KHP by NaOH, phenolphthalein indicator is used to determine the end point. Phenolphthalein (designated as HIn) is itself an organic acid, weaker than KHP (so it does not react with base titrant until the KHP has been quantitatively removed by neutralization). Reaction of excess base with phenolphthalein is shown as: HIn (colorless) + OH-

excess ↔ In- (pink) + HOH. This color change occurs only after all the stronger acids (KHP, etc.) have been neutralized by the added base titrant, and the color change occurs over a pH range from 8 to 10 (i.e. in slightly basic solution). If the solution is allowed to sit exposed to the atmosphere, CO2 may dissolve in the solution, forming carbonic acid, H2CO3, a stronger acid than phenolphthalein. This is why the color may fade and disappear after various periods of time. The phenolphthalein end point, then is seen at a point in the titration beyond the theoretical equivalence point for the neutralization (color change in the pH region 8-10, not at exactly pH 7). However, if the NaOH is standardized against pure KHP using the

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same visual indicator and operating conditions, the effective concentration of base determined experimentally will be such that the titration error will be canceled in the analysis of impure KHP samples. A note of caution about standardization procedures and analytical titrations with base solutions: bases tend to “layer” if stored for any period of time. To avoid concentration differences due to this non-homogeneous behavior, base solutions should be well shaken before portions are removed for use in titrations. Experimental Procedure Preparation of approximately 0.1 M NaOH solution Thoroughly clean a 1 liter plastic storage bottle and cap. Rinse the bottle and cap with several portions of deionized water. Using a 250-ml beaker, pour about 500 ml of deionized water into the clean bottle. Carefully obtain an appropriate amount (you will need to calculate this) of stock 19 M sodium hydroxide, NaOH (provided in a polyethylene bottle in the reagent hood), in a 25-ml graduated cylinder. Pour this solution into the water in the 1 liter bottle. Rinse the graduated cylinder several times with deionized water, adding the wash material to the large bottle (use a total of about 100 ml of water for rinses). Cap the bottle tightly, and mix the solution thoroughly by shaking/inverting the bottle. Use the large graduated cylinder to add another 400 ml deionized water to the solution, cap the bottle, and mix the solution well as described above. Label the bottle as 0.1 M NaOH, including your name and the date prepared. Standardization of the NaOH solution Reagent-grade potassium hydrogen phthalate, KHP, has been provided to you. This standard has been dried in an oven at 110° C for at least one hour. The material should have been cooled in a desiccator for at least 30 minutes before you weigh out the samples. Wash three 250-ml Erlenmeyer flasks, rinse the interiors with deionized water (do not dry!), and label for use. Accurately (to ± 0.1 mg) weigh by difference 0.5-0.7 g samples of pure KHP standard. Record the exact mass (to four decimal places) of KHP in each flask. To each sample, add about 75 ml deionized water and 2-3 drops of phenolphthalein indicator. Swirl the flask to aid in dissolving the KHP; all solid must dissolve and none must be stuck to the walls of the flask.

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Rinse the buret with deionized water. Make a final rinse (three preferably) with small (3-5 ml) portions of the 0.1 M NaOH solution. Fill the tip of the buret with the NaOH solution, making sure there are no air bubbles trapped in the stopcock. Fill the buret with the base solution so the upper level is between 0.00 and 1.00 ml in the calibrated region. Be sure to record the initial volume reading to two decimal places, and do not attempt to fill the buret exactly to the zero mark. Place one of the solutions of pure KHP under the buret, and begin adding the NaOH titrant, swirling the flask to thoroughly mix the solutions. Base may be added rapidly as long as the pink color fades quickly on mixing. When the color disappears more slowly, slow down the rate of base addition. Complete the titration by dropwise addition of titrant. Stop at the phenolphthalein end point: a single drop of added base produces a faint pink color throughout the solution, the color lasting for at least 15 seconds after thorough mixing. This pink color may fade upon sitting open to the air due to dissolving of CO2. Record the final buret volume, estimating the reading to the nearest ±0.01 ml. Discard the sample down the sink. Reset the buret with the NaOH solution from the storage bottle (mix the stock well before drawing out a portion), and repeat the procedure for the other KHP samples. Weigh out three additional samples of pure KHP, and carry out the titration process on a total of at least six KHP samples. Calculate the exact molarity of the base solution using data from each trial in the equation:

Calculate the relative deviation for all six trials. If this value exceeds 3 ppt, reject the most questionable value and recalculate the relative deviation of the five remaining runs. If the required precision level (3 ppt) is not attained, weigh out additional samples of pure KHP and repeat the standardization process until a minimum of five trials give molarities having 3 ppt or less deviation. Record the exact molarity of the base solution to four significant figures.

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Analysis of an Unknown Weak Acid (Impure KHP)

Unknowns have been dried at 110° C for 1 hour and cooled in a desiccator. Once cool, accurately weigh out three samples weighing 1.2-1.5 grams each into clean, labeled 250-ml Erlenmeyer flasks. Be sure to record all of the decimal places of the mass.

Proceed with the titration of the unknown weak acid, following the same

procedure as that described for the standardization using pure KHP samples. A minimum of four samples should be analyzed. Do more trials if time and material permit.

Calculate the percent KHP in the unknown, using the equation:

Save the remaining 0.1 M NaOH solution for use in the next experiment. Report the most appropriate value of the percent KHP in the original unknown sample including the standard deviation of your result. This will require that you propagate the error throughout the calculations. Assume that the error on reading the buret volume is ±0.02 mL. Adapted from Quantitative Analysis Laboratory Manual, Chemistry Department, Thiel College, 1989; and

Quantitative Analysis Chem 281 Coursepack, Revision 2, Chemistry Department, EMU, 2007.

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Experiment 3 Potentiometric Analysis of a Weak Acid

The potential of an electrode depends on the concentrations of one or more species present in the solution in which the electrode is immersed. This, coupled with knowledge of pertinent standard electrode potentials and a knowledge of the stoichiometry of a reaction, allows one to follow the course of a reaction and ultimately to determine the equivalence point. Thus it is possible, using a standard titrant, to determine the percentage composition of an unknown mixture. The potential of the glass electrode, E, depends upon two hydronium ion concentrations (activities):

or for dilute solutions

where E’ is a constant containing all contributing factors other than hydronium ion activities; n=1; [H+]1 is the concentration of H3O+ inside the glass electrode and [H+]2 is the concentration of H3O+ in the solution; R, T and F have their usually accepted values. Since E’, R, T, F, [H+]1 and n are all constants for a given system, the equation can be rewritten as:

E = K + k log [H+]2 Thus changes in [H+]2 will be directly reflected by changes in E. The change can be measured, and with appropriate instrumentation directly displayed as pH. Combination electrodes are available which combine the functions of the calomel reference electrode, the internal reference in the glass electrode, and the glass sensing portion which develops a changing potential across the glass membrane as the concentration of H3O+ in the solution changes. Weak acids can often be identified on the basis of their equivalent weight and the pH value of the half-titrated solution (pH at 50% titration = pKa for dissociation process). The equivalent weight may be experimentally determined by titration with a standard sodium hydroxide solution. The titration endpoint may be observed by use of a visual indicator such as phenolphthalein or by monitoring the pH changes with a pH meter. In either method, if the equivalent volume of base and the pH at 50% can be established, these values can be used to identify the acid. If the unknown acid is monoprotic, the interpretations of data is greatly simplified. If the acid is polyfunctional, analysis of the corresponding results is much more complicated, but may be successfully undertaken.

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Experimental Procedure

Titration procedure using pH meter Calibrate the pH meter as demonstrated by the instructor unless they have already been calibrated for you. Take care not to damage the glass combination electrode. The electrode should be rinsed thoroughly with deionized water before and after use, allowed to drain (touch the tip with a paper towel, but do not rub or wipe harshly), and stored in an appropriate storage solution when not in use. A 250-mL beaker – NOT an Erlenmeyer – is used as the titration vessel, and each sample is dissolved in 150 mL deionized water. Place a clean magnetic stir-bar in the beaker and adjust the stirring to a medium rate (do not generate a vortex). Do not allow the stir-bar to strike the electrode or beaker walls. Refer to page 12 for the proper use of the Vernier pH electrodes. Add 50-75 mL of water to the beaker, insert the pH electrode (and temp probe if present) and arrange as shown below: Place the buret above the titration beaker, and take an initial reading from the pH meter; note that the volume is 0.00 mL of added titrant. While the PC will take the pH readings, you will need to write in the volume of added titrant. Add the NaOH titrant incrementally, recording volume vs. pH at each addition. Plot the resulting titration curve. The equivalence point is the inflection point of the curve.

Magnetic Stirrer

Stirring Bar

pH Electrode

Buret with NaOH Solution

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Analysis of an unknown weak monoprotic acid Your unknown should have been dried at 110 °C for 1 hour; make sure the sample is cool before weighing it. Accurately weigh 0.2 – 0.3 grams each into clean, labeled 250-mL beakers. Dissolve each sample in 100 mL deionized water, heating if necessary on the hotplate in the hood. To the first sample ONLY, add 2-3 drops of phenolphthalein indicator. Run a very rapid potentiometric titration with the pH meter. Add the NaOH in 2 mL increments and record the volume and corresponding pH after each addition (follow the handout directions for using the Vernier pH electrodes and the computers). Wait no longer than a few seconds for the pH reading to stabilize. Note the pH and volume of titrant when the indicator changes color, and record this point for later use. Add a final total equal to twice the value at the color change (take the titration to 200%). When you are finished, copy the data into Microsoft Excel Repeat the procedure with the remaining two samples, leaving out the phenolphthalein. Add the titrant more slowly and in smaller portions. Add 1 mL increments to a point within 85% of the color change volume noted earlier, then change to 0.5 mL portions up to 115% titration, and then return to 1 mL portions until the titration is carried out to 200% completion. Using Excel, plot graphs of pH vs. mL titrant for each of the accurate (latter two) titrations. Because it is difficult to determine the equivalence point from this kind of curve, you will also need to plot the first derivative of that curve. To do this – approximately – you should subtract each volume measurement from the previous one. Set up your spreadsheet like this:

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The first two columns are the raw data, used to make your titration curve, which should look like this: The endpoint on this graph is the inflection point, indicated by the arrow, where the curve goes from being concave to convex. Because it’s hard to see this point, it’s easier to plot the first derivative of the curve, which looks like this:

Endpoint

Endpoint volume

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Even clearer is the plot of the second derivative, which has a zero crossing at the endpoint. The second derivative of the titration curve looks like this:

The equivalent weight is determined as:

Print out both of the good titration curves and tape them into your lab notebook. You must note on the graphs the end point volume and the ½ endpoint (where pH = pKa)

Endpoint volume

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Experiment 4 The Determination of Macro Concentrations of Sodium in Water by

Ion-Exchange and Titration The determination of macro concentrations (0.1 M) of Na+ has never been easy when classical means are employed. Gravimetric determination, though highly accurate is slow and tedious. Colorimetric techniques are most unhandy and also yield poor accuracy. The authors do not believe that a direct titrimetric method for sodium exists anywhere, and are not about to try to develop one. There is, however, a method employing both a cation-exchange resin and and acid-base titration which may be applied to the problem, provided that the Na+ is the only cation present in significant quantity in the solution to be analyzed. In this experiment, this method will be applied. What, first, is a cation exchange resin? Mostly, it appears to be made up of solid granules. These organic granules are polymeric in nature, and thus of high molecular weight, and are very, very insoluble in water. One sort of cation exchange resin has a polystyrene-like structure, a portion of which is shown below. Note the sulfonic acid groups (-SO3H) attached here and there. -------CH--------CH2--------CH--------CH2--------CH--------CH2--------CH--------CH2------ -------CH--------CH2--------CH--------CH2--------CH--------CH2--------CH--------CH2------ Even though the resin itself will not dissolve in water, it is quite capable of entering into chemical reactions with ions in aqueous solutions through the sulfonic acid groups. Sulfonic acids are quite strong, so that resins containing them are known as strong acid resins as well as cation exchange resins. The hydrogen ions on the resins may be exchanged for cations from the solutions the resins come in contact with. Thus, while the great, bulky, ugly and loathsome organic body of the resin is unaffected by an aqueous ionic solution, the attached ionic groups may carry on a lively and vigorous chemical traffic with the ions in solution. Such traffic may be described mathematically in a way that would cause economists to turn green, or at least chartreuse, with envy. Suppose that one has a cation-exchange resin, represented by RH+, and a solution containing Na+ ions. The reaction between the resin and Na+ may be given as

Na+ + RH RNa + H+

SO3H SO3H SO3H SO3H

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The equilibrium constant for this reaction is

In the case of the resin we are using (Amberlite IR-120(H) C.P.). Keq is sufficiently large so that if a neutral solution of NaCl is poured through it, all of the sodium ions will be grabbed by the resin. Each hydrogen ion will be replaced by a sodium ion. The hydrogen ions may then be washed from the resin and titrated with NaOH, and the amount of sodium originally present in the neutral solution of NaCl calculated. The resin may be regenerated (reconverted to its original form) by washing it with a concentrated (2M) solution of HCl. The magnitude of Keq is such that if [H+] = 2M, the reaction above is driven nearly completely to the left, and any residual sodium on the resin is replaced with hydrogen ion. In the experiment that you will do, you will be given roughly 100 mL of a NaCl solution of unknown concentration. You will pass a 25.00 mL aliquot of this solution through the resin, and collect it in a flask. You will wash the resin with deionized water, collecting the washings in the same flask. You will then titrate the contents of the flask with your standardized NaOH solution (from the KHP experiment), and calculate the molar concentration of NaCl in the solution you were issued. The exact procedure is given below. Procedure

READ THIS PROCEDURE CAREFULLY. ONCE YOU HAVE APPLIED YOUR SAMPLE TO THE COLUMN, STEPS 5-7 MUST BE COMPLETED DURING THE

SAME LAB PERIOD. 1. Request an aqueous solution of sodium chloride of unknown concentration and, if

you are out, another sample of standard KHP. 2. Obtain a column packed with cation

exchange resin from the prep room. The liquid level must never fall below the top of the glass wool or else air will enter the column and make portions of it inaccessible for ion exchange to take place and efficiency will be reduced.

3. Start with regeneration of the column

by passing 100 mL of 2M HCl through it. Slowly open the stopcock until liquid flows out at roughly 6 mL/min (flow rates are determined easily with a graduated cylinder).

Funnel

Glass Wool Resin

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Collect the effluent in a beaker. Wash the sides of the funnel with deionized water and allow the liquid to flow through the column, but never let the top level drop below the glass wool.

4. Fill the funnel with deionized water. Catch the water in the beaker. Run water

through the resin until the pH of the effluent is the same as that of the deionized water used (use pH paper). Discard the contents of the beaker.

5. Lower the water level in the column until it is just above the TOP of the glass

wool. Transfer 25.00 mL of your unknown solution to the funnel (measure with a transfer pipet). Pass it through the column at 6 mL/min, and collect it in a clean 250 mL Erlenmeyer flask. Wash the column with four 25 mL portions of deionized water (measure with a graduated cylinder), collecting in the same flask. Cover the flask, label it, and set it aside.

6. Repeat step 5 two additional times. 7. Close the stopcock on the column, leaving liquid in the funnel, wrap with

parafilm and return the column to the prep room. 8. Add two drops of phenolphthalein to each of the Erlenmeyer flasks. Titrate each

to the faint pink endpoint with the standardized NaOH solution from the KHP experiment. Check the concentration of the NaOH solution by titrating it again with standard KHP. If the molarity of the first titration is significantly differently than determined in the KHP experiment, titrate 2 more times and calculate the average molarity of the NaOH solution.

9. Save your remaining unknown for the next experiment.

Calculations If you restandardized your NaOH solution (or made more), calculate the new molarity of that solution as you did in the KHP titration experiment. Report the new molarity to four significant figures. Calculate the concentration of NaCl in your unknown in units of molarity. You have how many milliliters of NaOH were required to neutralize the H+ eluted from the column, the molarity of that NaOH, and the volume of salt solution you used. Refer to the reaction on page 35 for the stoichiometric ratio between Na+ that you put onto the column and H+ that you recovered from the column. Set up an example calculation in your notebook and have it checked by your instructor before continuing. Report your result as the molarity of NaCl with standard deviation. Also calculate your relative standard deviation, in ppt, and the 95% confidence limits on your result.

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Questions What would be different about the results of this experiment if your unknown were a solution of Na2CO3 or Na3PO4 rather than NaCl? Be sure to comment in your conclusion section how your titration technique has improved (or not) over the course of the first several experiments. Reference: R. Kunin, R. W. Percival, T. J. Kneip, and W. K. Dean, Experiments in Ion Exchange, 2nd Ed., Mallindrodt Chemical Works, St. Louis, MO., 1965.

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Experiment 5 Determination of Water Hardness

Introduction Ethylenediaminetetraacetic acid (EDTA) forms stable complexes with many metal ions, as shown in Figure 4. We will use a complexometric titration with EDTA to determine calcium concentration in YOUR hard water (bring about 1L from your tap). The titration must be carried out at pH 10 (in a NH3/NH4

+ buffer) using Eriochrome Black T indicator. Calcium does not form stable complexes with this indicator except in the presence of Mg2+ ion. The Mg2+ is added as a few drops of MgCl2 solution; the indicator changes from red to blue in the presence of Mg after all of the Ca has been complexed by EDTA. Retaining the form of the reactions shown in Experiment 4, the reaction here is: Ca2+ + H2B2- CaB2+ + 2H+ Because we will be working with ammonia buffer during this experiment, please inform the instructor if you are asthmatic, as the strong odor of the buffer can cause breathing difficulties if it is not handled properly. Summary

A variety of experimental procedures are used in applications of EDTA titrimetric analyses. The most common methods include: direct titration, back-titration, and displacement titration. A direct titration procedure is simplest and most straight-forward, both experimentally and theoretically. It involves the addition of EDTA titrant to a solution containing the metal ion(s) of interest, the titration carried out to a point at which an appropriate indicator changes color.

The calculations are quite simple, following the fundamental stoichiometry for

general titrimetric processes: mmol cation = mmol standard EDTA. The weight of a particular substance analyzed by complexometric titrimetry is given by:

milligrams of analyte = (mL EDTA) x (M EDTA) x (Formula weight analyte)

for any 1:1 EDTA-cation reaction. In many cases, the values are so small that the final result is best expressed in parts per thousand (ppt) or parts per million (ppm), rather than percentage.

If the disodium salt of EDTA is the titrating agent and Eriochrome Black T (EBT) is the indicator, the analytical and indicator reactions are as follows:

analytical: H2Y2- + M2+ MY2- + 2H+ indicator: H2Y2- + Min MY2- + HIn2- + H+

(wine-red) (blue)

40

where M2+ is any appropriate divalent metal cation. The EBT indicator does not give a satisfactory end point for the direct titration of a divalent cation, such as calcium, in the absence of magnesium ion. The calcium-EBT complex is weak; consequently, the end point is gradual and appears too soon. However, the magnesium-EBT complex is stronger, and the endpoint is sharper.

When calcium and magnesium ions are present in the same sample (as in hard water), the free calcium and magnesium ions are titrated together. Calcium forms a more stable EDTA complex than does magnesium, so calcium is titrated first, followed by magnesium. Since the strengths of the indicator complexes with calcium and magnesium are opposite to the EDTA complexes, the end point is signaled after all the calcium and magnesium ions have combined with EDTA. If the sample does not contain magnesium, the indicator solution (or EDTA titrant) is usually prepared with a small amount of magnesium ion to insure that the appropriate color change occurs properly.

Solid samples containing soluble metal cations can be analyzed by EDTA titration, but only in terms of total cation content (since EDTA complexes indiscriminately with nearly all cations). The analysis follows the general procedure used for any titrimetric determination of a solid unknown: a sample is weighed out, dissolved in appropriate solvent and indicator, and titrated to the correct end point.

Liquid samples can also be analyzed by EDTA titration, as in the case of so-called hard water (water containing dissolved materials such as compounds of calcium and magnesium). The materials responsible for the hardness of water may be removed through a process called water softening. The extent of the hardness determines the magnitude of the softening process, which is frequently a major industrial problem. The hardness of water is usually expressed in concentration units (ppm or mg/liter), defined in terms of CaCO3 content. This expression is not strictly true since hardness is due to both calcium and magnesium salts, but it serves as a convenient comparison method.

In the direct titration involving EDTA, it may be necessary to add an auxiliary complexing agent, such as citrate or tartrate ions, to prevent the precipitation of certain metals as insoluble hydroxides. Also, a buffer (NH4

+ + NH4Cl) is frequently used to supply a complexing agent for certain metals and to control the pH of the solution. At a pH of 10 (as provided by an ammonia buffer), many metals are held in solution as ammine complexes and can be titrated directly with EDTA (EDTA forms a more stable complex than ammonia).

As in the case of precipitation titrations, the disodium salt of EDTA is available in sufficient purity to use as a primary standard. However, the effective (or experimental) concentration of the EDTA solution may be significantly different from the theoretical value, so standardization of EDTA against primary standard magnesium solution is routinely performed. EDTA standardized in such a procedure can then be used for direct titrations of solid and/or liquid samples, as well as in back-titrations and displacement titrations of various metal cations.

41

Experimental Procedure

Preparation of glassware All glassware (volumetric flasks, Erlenmeyers, etc.) used in this

experiment must be cleaned thoroughly with soap and water. Use dilute nitric acid, HNO3, to remove traces and films. Rinse the glassware with tap water, and make several final rinses with deionized water. Turn the glassware upside down to drain dry, and do not wipe out the inside with towels. Preparation of approximately 0.01 M EDTA

The disodium salt of EDTA.dihydrate (F.W. 372.24) is provided in the balance room. Accurately weigh out about 2.5 grams of the solid into a clean 500.00-mL volumetric flask.

Add about 75 mL deionized water to dissolve the solid, mix well, and dilute to the mark with deionized water. Store the solution in a clean, dry bottle, and label as 0.025 M EDTA solution.

Calculate the theoretical or maximum EDTA concentration, using the

recorded exact weight and volume of solution. Record this value for later use. Preparation of primary standard magnesium solution

Obtain 1-2 grams of reagent-grade magnesium sulfate, MgSO4 (F.W. 120.37), from the TA, place the solid in a clean, dry, labeled weighing bottle or small beaker, and dry the container/contents in an oven at 110°C for 1-2 hours. Cool the material for 30 minutes in a desiccator, and accurately weigh out one sample of 0.5-0.6 grams into a clean 150 mL beaker. Add about 50 mL deionized water to the solid in the beaker, and mix well. Heat the beaker on a hot plate to hasten the dissolving process. Quantitatively transfer the solution to a clean 250.00 mL volumetric flask, rinsing the beaker with several portions of deionized water (adding the washing to the volumetric flask), and dilute to the mark on the flask with deionized water. Calculate the exact concentration of magnesium (or Mg2+ ) in the primary standard solution, using the equation:

Record this value for use in standardizing the EDTA solution. Standardization of EDTA solution

Use a clean, properly rinsed pipet to transfer a 25.00-mL aliquot of magnesium standard into a clean 250-mL Erlenmeyer flask. Add about 25 mL deionized water, and mix. Add 3 mL of pH 10 NH4Cl-NH4OH buffer (from the hood), and then add 5 drops of Eriochrome Black T indicator (from a dropper bottle on the bench tops). The sample solution should turn a wine-red color.

42

Titrate the magnesium solution with EDTA, adding the titrant slowly. Near the end point, each titrant addition must be added very slowly, allowing 4-5 seconds for complete mixing. Maintain a strong swirling action during the process. The correct end point is when one drop of EDTA titrant causes the solution color to change from light pink/violet to light blue/grey.

Complete chelometric titrations of at least four standard magnesium samples, and calculate the effective or experimental EDTA concentration for each trial, using:

Calculate the relative deviation for the standardization trials. If the value is 3 ppt or less, record the experimental concentration of EDTA for later use in determining the unknown percentages and/or concentrations. If the deviation is more than the allowed limit, drop the most questionable value and recalculate. If still too high, run additional standardizations until a minimum of three trials give a relative deviation of 3 ppt or less.

Determine the difference between the experimental and theoretical EDTA concentrations. Calculate the relative error in the titration process, assuming the theoretical EDTA value is correct. Offer a possible explanation for the difference (if any) in the laboratory report.

Analysis of Hard Water

Use a clean, properly rinsed 25.00-mL transfer pipet to obtain 100.00 mL aliquots of the hard water. Place an aliquot in a clean 250-mL Erlenmeyer flask. Add 3 mL of pH 10 buffer, and then 5 drops of EBT indicator. Titrate the aliquot of hard water with EDTA, following the same procedures as in the standardization process. Titrate four hard water samples and then calculate the concentration of CaCO3 in the hard water.

Report the most appropriate value for the hardness of the water sample. Question: If there are 7000 grains per pound and 946 mL per quart, what would be the grains of Ca2+ per gallon?

Source: Quantitative Analysis Laboratory Manual, Chemistry Department, Thiel College, 1989.

43

Experiment 6 The Determination of Sodium in Trace Amounts

by Flame Photometry The purpose of this experiment is to determine the sodium content in a water sample. On the face of it, this is routine. It is worth noting, though, that wet chemical methods for sodium are extremely slow and clumsy. Moreover, these methods are not very sensitive. There is a far better, far more sensitive means for determination of the alkali metals - flame photometry. Basically, the processes involved in the flame photometric analysis are shown in Figure 1:

Figure 1. Schematic representation of a flame photometer This configuration allows the measurement of light emitted from metal atoms as they return to their electronic ground state after excitation in a flame. The wavelength selector is set to a wavelength unique to the element of interest and thus, in accordance with Beer’s Law, the emission measured by the detector is directly proportional to the concentration.

For example, a sample, (here a solution containing sodium), is drawn up through an aspirator so that a fine mist is sprayed into the flame. Some sodium atoms become excited in the flame, and promptly emit light of wavelength 589.6 nm - the characteristic yellow radiation we so often see in flame tests. The light of this wavelength is isolated by the monochromator, either prism, grating, or filter as in our instruments, and then falls upon the detector. The detector, which may be a photocell, phototube, or photo-multiplier, gives an electric current which is proportional to the intensity of the light

Digital Display

0.000

44

falling on it. This current is monitored by some sort of readout device, which can range from a very simple to a very complicated piece of apparatus.

For analysis, one prepares a standard curve by plotting detector response versus corresponding known concentrations. This produces a straight line as they are linearly related [see Figure 2]. The concentration of an unknown solution can thus be calculated using the equation of the line obtained from a linear regression of the data points and substituting in the response measured from the unknown solution.

Figure 2. Standard Curve for Na Perhaps this is a good place to put in a word about the concentration units used in the experiment which are parts-per-million (ppm), or one part in a million parts:

1 ppm = 1 mg/one million mg When dealing with aqueous solutions in this concentration range it is convenient to use the fact that the mass of 1 liter is 1000 g or 106 mg (remember - the density of water is 1 g/ml). This permits the equivalent, convenient fraction to be used:

1 ppm = 1 mg / liter

This experiment may be divided into three parts.

1. Preparation of standard solution of sodium 2. Measurement of emission of unknown and standard solutions in order to

establish standard curves 3. Determination of sodium content of the unknown.

45

Procedure: 1. Preparation of standard solutions of sodium and potassium:

a. Request samples of pure NaCl along with your solution of unknown

concentration as instructed on page 11. Prepare 250.0 mL of a solution that is close to 1000 ppm in Na (not NaCl) Record the exact mass of the salt used as you will need it to calculate the exact concentration of the metal.

b. Glassware used to prepare the dilute standard solutions in this step should be

rinsed first with 1 or 2 M HCl and then with deionized water to remove Na+ adsorbed to the glass. Use deionized water to dissolve and prepare solutions. Beware of contaminating your solutions with sodium, which is present nearly everywhere, including your fingers.

c. Prepare 250.00 mL of a 100 ppm Na solution in Na by diluting an appropriate

amount (via a volumetric pipette) of your approximately 1000 ppm standard solution to 250.00 mL. Mix well. Prepare 100.00 mL of 1, 2, 3, 4 and 5 ppm solutions by transferring the appropriate amount of the 100 ppm solution using volumetric pipets and diluting to the mark.

2. Establish the sodium standard curve:

a. To start the instrument, follow the instructions for use of the BUCK flame

photometer given on page 15. b. Using the photometer fitted with the sodium filter,

establish the standard curve for sodium by recording the emission for the blank, the five standard solutions, and the unknown immediately after setting the zero and sensitivity of the instrument.

Calculations

First, calculate the actual concentration (in ppm) of sodium in your standard solutions. Then, use the data obtained in the experiment to prepare a standard curve of emission (y-axis) versus the actual concentration of Na (x-axis). Use linear regression to determine the equation that produces the best fit line to the Na data. Use the equation to determine the concentration of Na in your unknown solution.

46

Questions If we were to use the flame photometer to measure the amount of K and Ca in soils using an acid solution, would it be appropriate to use DI water as the blank? Explain and justify your answer. Be sure to include in your notebook for the Na by flame photometry experiment:

• A clearly labeled and written table of your flame photometer data. • All calculations for the Na concentration in your standard solutions. • Printouts of your graph and spreadsheet. • Exactly how you calculated the final concentration of your unknown. • The exact concentration in ppm of your unknown, as the result. • Conclusions that describe what you learned and what you should remember for

future use of the photometers. • Address the questions in the conclusions as well.

47

Experiment 7 Spectrophotometric Determination of Iron in Trace Amounts

The purpose of this experiment is to determine the concentration of iron present as a trace constituent in an aqueous sample by a spectrophotometric method. Generally speaking, spectrophotometric methods are far more sensitive than classical wet chemical methods, but are not as accurate. The chemistry of this system is not too terribly complex. Both Fe2+ and Fe3+ form quite stable complexes with 1,10 phenanthroline (or ortho-phenanthroline). The Fe2+ complex has an absorption maximum at 510 nm and the Fe3+ has an absorption maximum in the ultraviolet. Because the instrument we use is not designed for ultraviolet work, we have chosen to determine the iron present as Fe2+.

The Fe3+ present is reduced by the use of hydroxylamine hydrochloride, NH2OH •HCl, to Fe2+. The Fe2+ is then complexed with 1,10-phenanthroline:

With a Kf this large, the reaction is almost guaranteed to be quantitative. The pH of the solution can, however, play a role. Like ammonia, 1,10-phenanthroline (which we shall abbreviate as "o-phen") is a weak base and the reaction: o-phen + H+ ⇔ H(o-phen+) has a tendency to go very far to the right in a strongly acid medium. An acid medium must be used for the determination of iron; otherwise, a precipitate of Fe(OH)2 is apt to form. Thus, the reaction Fe2+ + 3H(o-phen)+ ⇔ Fe(o-phen)32+ + 3H+ occurs in acid solution. Too high an H+ concentration could retard the formation of the brightly-colored Fe(o-phen)32+ complex. Thus, the pH must be kept within limits; these are chosen as 2.5 to 4.5. The solution is buffered to accomplish this end. Beer's Law is obeyed rather faithfully by this system. At 510 nm a plot of absorbance versus concentration of [Fe(o-phen)3]2+ (hence CFe, or the total iron concentration) yields a beautiful straight line:

48

Thus, one can prepare a set of standard solutions, of known CFe, establish a straight line, and prepare the unknown solution in exactly the same way. The absorbance of the unknown solution is measured, and the unknown concentration determined from the equation of the line calculated using linear regression analysis. The stoichiometry of the Fe2+ ortho-phenanthroline complex may also be confirmed by performing a titration monitored spectrophotometrically. Consider the reaction:

Fe2+ + 3H(o-phen)+ ⇔ [Fe(o-phen)3]2+ + 3H+

Of all the reactants and products listed, only [Fe(o-phen)3]2+ , the complex, absorbs at 510 nm. Thus, if one began with a solution containing only Fe2+, and added ortho-phenanthroline to it, one would notice a deepening of the orange color as the complex [Fe(o-phen)3]2+ forms and an increase of the absorbance at 510 nm. The absorbance would continue to increase until all the Fe2+ had been complexed, after which addition of more ortho-phenanthroline would only dilute the solution and cause the absorbance to drop somewhat. So, although theoretically, absorbance remains a constant after the endpoint is reached there is a slight decrease observed experimentally. Hence, a plot of absorbance versus mL of ortho-phenanthroline added would look like this:

49

It is possible to accommodate for the dilution and flatten the horizontal portion, but here is not necessary. The stoichiometry of the complex, n, (the number of moles of o-phen required to complex one mole of Fe2+) is readily confirmed using the titrimetric endpoint which can be determined by calculating the point of intersection of the two linear portions of the graph. PROCEDURE: Note: The hydroxylamine hydrochloride and ortho-phenathroline solutions are located in burets in the the back corner of the lab 1. Prepare the standard and unknown iron solutions:

a. Request an unknown as per the instructions on page 11. b. Calculate and weigh out the amount of ferrous ammonium sulfate

hexahydrate, Fe(NH4)2(SO4)2•6 H2O (392.14 g/mole,found in the prep room) needed to obtain 0.0250 g Fe. Dissolve it in 50 mL of deionized water plus 2 mL concentrated HCl in a beaker, then quantitatively transfer (meaning rinse the beaker numerous times, adding the rinse to the flask) to a 250 mL volumetric flask and fill to the mark with water. This solution is 100 ppm in Fe. Loss of solution here can lead to grievous error later.

c. For the standard solutions:

Determine the amount of the 100 ppm iron solution needed to prepare 100 mL of a 1, 2, 3, and 4 ppm iron solution and 200-ml of a 0.5 ppm solution. To each flask, add the appropriate amount of 100-ppm iron solution. Then add 5 mL of hydroxylamine hydrochloride (NH2OH·HCl) solution. Place a small square (1/4" x 1/4") of Congo Red indicator paper in the solution. It will turn blue. Add 2 M sodium acetate solution dropwise until the paper turns red again. The pH will now be between 2.5 and 4.5. Add 5 mL of the 2.5 g/L ortho-phenanthroline solution. Then dilute to the mark with water. Let each standard solution stand for at least 10 minutes, so that its color can develop fully.

d. Now prepare the unknown you were given in a 100-mL volumetric flask.

Add 5 mL of hydroxylamine hydrochloride to the solution in the volumetric flask. Place a small square (1/4" x 1/4") of Congo Red indicator paper in the solution. It will turn blue. Add 2 M sodium acetate solution dropwise until the paper turns red again. The pH will now be between 2.5 and 4.5. Add 5 mL of the 2.5 g/L ortho-phenanthroline solution. Then dilute to the mark with water. Let the unknown solution stand for at least 10 minutes, so that its color can develop fully.

50

2. Dissolution of vitamin tablet: Select a vitamin tablet from the choices in the prep room. Record the brand you’ve chosen and the amount of iron reported in the tablet as per the label. IN THE HOOD place one tablet of the brand of your choice into a 100 mL beaker and heat slowly to a boil with 25 mL of 6M (not concentrated) HCl for 15 minutes. Dilute the mixture slightly with water and filter while hot through Whatman 114 filter paper (or other fast filtering paper) directly into a 100 mL volumetric flask. After washing the residue with hot water, allow the filtrate to cool, then dilute to the mark. Pippette out a 5.00 mL aliquot (10.00 mL if the tablet contains less than 15 mg Fe) of this solution and place into a 100 mL volumetric flask and dilute to the line. Next, transfer 10.00 mL of the diluted solution into a 100 mL volumetric flask and . Add 5 mL of hydroxylamine hydrochloride to the solution in the volumetric flask. Place a small square (1/4" x 1/4") of Congo Red indicator paper in the solution. It will turn blue. Add 2 M sodium acetate solution dropwise until the paper turns red again. The pH will now be between 2.5 and 4.5. Add 5 mL of the 2.5 g/L ortho-phenanthroline solution. Then dilute to the mark with water. Let the solution stand for at least 10 minutes, so that its color can develop fully.

3. Find the absorption maximum of the [Fe(o-phen)3]2+ complex: Fill a cuvette with your 4ppm solution. Plot the absorbance spectrum. From the

graph, identify the wavelength of maximum absorbance which should be around 510 nm.

4. Establish the standard curve or Beer's Law plot:

Set the wavelength dial to the absorbance peak you found. Then zero the spectrophotometer with distilled water. Read the absorbance, A, for each standard solution, the unknown and vitamin solution. Plot the A values on the ordinate (y-axis), and the corresponding concentration values of the standard solutions on the abscissa (x-axis). A straight line should result from the Beer's Law relationship, A = εbc. If the points are badly scattered, your procedure was most likely in error.

5. Determine the concentration of Fe in the two samples:

Calculate the equation of the standard curve using linear regression. Use this equation to calculate the concentration of iron in your unknown and vitamin tablet solutions. Remember to account for dilution: you were given a 10.00 mL sample of unknown which you subsequently diluted to 100.0 mL, and the vitamin solution has a number of dilutions that must also be accounted for in your calculations. Compare your value for the vitamin tablet to the value reported by the manufacturer. Calculate the percent relative error.

51

6. Confirm the stoichiometry, n, of the complex:

a. Set the wavelength dial of your spectrophotometer to the wavelength of maximum absorbance. Use water to set the instrument (blank the instrument) to zero absorbance. Place 5.00 mL of your 100 ppm Fe solution in a 250 mL beaker. Add 5.00 mL NH2OH·HCl solution and 90.00 mL of H2O. Use your buret to add the H2O. Then, add enough sodium acetate to adjust the pH to 3.5, again using the Congo Red paper.

(Note: The 5.00 mL aliquot contains 8.96 x 10-6 mole Fe if your original

solution is exactly 100.0 ppm.) b. The following operations should be done without losing a drop of solution.

Add 1.00 mL of the 1.0 g/L ortho-phenanthroline (NOT the 2.5 g/L solution) to the beaker without spilling any. The solution will get orangy. Mix well. Wait for 2 minutes. Now mix and transfer with an eyedropper some of this solution into a cuvette. Read its absorbance. Next, pour the contents of the cuvette back into the beaker. Add another 1.00 mL of the 1.000g/L ortho-phenanthroline to the beaker, mix well and wait for 2 minutes. Use your eyedropper and rinse the cuvette with the mixed solution, and read its absorbance. (Absorbance will rise as the solution gets more orange.) Pour the contents of the cuvette back into the beaker. Add another 1.00 mL of the 1.000g/L ortho-phenanthroline to the beaker, mix well and wait for 2 minutes. Use your eyedropper and rinse the cuvette with the mixed solution. Fill the cuvette with mixed solution and read its absorbance. You've probably got the idea now. Keep up the process until the absorbance stops climbing and drops, giving two nice lines. You won't need much over 10.00 mL of 1.000g/L ortho-phenanthroline for the entire process.

c. Use Graphical Analysis or some other spreadsheet program to plot mL

ortho-phenanthroline vs. corrected absorbance to confirm the stoichiometry of the complex. Note that the concentration of the 1.000 g/L solution of ortho-phenanthroline is 5.05 x 10-3 M. Carefully draw a straight line through the linear portions and extrapolate the lines until they intersect. Label the end point. Since it is not easy to read volumes accurately from graphs, it is better to calculate the end point volume. This can be done by doing a linear regression on each straight line portion of the corrected absorbance graph where:

52

Corrected A2 = m2VY2 + b2 Abs. A1 = m1VY1 + b1

endpoint volume ml o-phen added

At the end point:

A1 = A2 so:

m1VY1 + b1 = m2VY2 + b2

Also, at the intersection, VY1= VY2 so we can drop the subscripts and rewrite the equation: m1VY + b1 = m2VY + b2 and rearranging, solve for the endpoint volume:

This volume is used to calculate the amount of o-phen that corresponds to the amount of Fe used, thus verifying the stoichiometry of the complex where:

Be sure to include the following in your notebook:

Three graphs: (1) Plot of Absorbance vs. λ, showing λmax.

(2) Plot of Corrected Absorbance vs. mL o-phen added, showing endpoint volume. (3) Beer's Law plot of Absorbance vs. ppm Fe, including the equation of the best fitting line and the correlation coefficient.

53

Experiment 8 Simultaneous Spectrophotometric Determination of Two Components

in Solution (Co2+ and Cr3+)

A typical sample contains only one analyte (here, species X) and so the procedure is rather straightforward. First, one selects a wavelength that offers the most sensitivity to analyte concentration, usually found where absorbance is at a maximum. A spectrophotometer is used to gather data to make an absorbance versus wavelength plot, as shown:

Next, a working curve or Beer’s Law curve is plotted by measuring the absorbance of solutions of known analyte concentration at the selected wavelength:

Finally, the absorbance of the solution of unknown X concentration is measured and its concentration determined based on the equation of the standard curve obtained via linear regression.

From this you can see that the analysis of a sample with a single analyte is really quite simple. However, the procedure becomes slightly more complicated when there is

54

more than one analyte present. Separation of analytes is often time consuming, expensive, can result in a loss of sample, or sometimes very difficult. Therefore, it is often preferable to avoid separation procedures and instead develop a method that permits the quantitation of each analyte in the mixture simultaneously. It is quite possible to analyze simultaneously for two substances spectrophotometrically, if they do not interact chemically with each other. Many such pairs of substances exist: Cr3+ & Co2+, Cr2O7

2- & MnO4

-, and certain phenanthroline complexes of Fe2+ and Co2+ are examples. Suppose that we have two substances, X and Y, which do not interfere chemically with each other. Let us run an absorption spectrum for each species:

If we mix X and Y together, we get a curve which is a combination of the absorbance curves for X and Y. In short, at any wavelength, the absorbance of a solution of X and Y is the sum of the absorbances of X and Y at that wavelength. Let us proceed to the mathematical analysis of absorbances of mixtures of X and Y. Suppose that we wish to analyze an unknown solution. The concentrations of X and Y, CX and CY, are two unknowns. In order to determine two unknowns, we need two

For X alone

For Y alone

X Alone

X + Y

Y Alone

A1

55

equations. Let us, by way of developing these equations, measure the absorbance of the unknown solution at wavelengths λ1 and λ2. This is reasonable, for X has an absorbance peak at λ1, and Y a peak at λ2. Note that the absorbance of Y is near a minimum at λ1 as is the case at λ2 for species X. At λ1, the absorbance of the solution, A1, is given as the sum of the absorbances of X and Y at that wavelength: A1 = AX1 + AY1 But, by Beer's Law AX1 = ∈X1BCX

where εX1 is the molar absorptivity of species X at λ1, B is the cell length, and CX the concentration of species X. By the same sort of reasoning,

AY1 = ∈Y1BCY

Hence A1 = AX1 + AY1 = ∈X1BCX + ∈Y1BCY

At λ2, the absorbance of the solution, A2, is given as the sum of the absorbances of X and Y at that wavelength: A2 = AX2 + AY2 Using Beer's Law, we see that A2 = ∈X2BCX + ∈Y2BCY

Thus, we have two equations, A1 = ∈X1BCX + ∈Y1BC Y (Eqn. 1) A2 = ∈X2BCX + ∈Y2BC Y (Eqn.2) If we know A1, A2, ∈X1, ∈X2, ∈Y1, ∈Y2, and B, we may solve for CX and CY! There is only one problem with this approach. The alert student is no doubt curious about how all those ∈’s are to be determined. It’s really rather simple. One makes a standard curve for pure solutions of X and Y at varying concentrations at both

56

λ1 and λ2. For example, for species X:

at λ1,we see, while at λ2, we see

(Note the increased sensitivity to the concentration of M at λ1as opposed to λ2 ) The standard curves are linear in accordance with Beer’s Law, thus y = mx + b, where y = Absorbance, x = [X] or [Y], m = ∈XB. Ideally, b, according to Beer’s Law should be 0, however, experimentally this is not often the case and therefore the intercept, b, must be included in the calculations. Standard curves must be obtained and slopes and intercepts calculated for both components using linear regression. In this way the 4 molar absorptivity values, ∈X1, ∈X2, ∈Y1, ∈Y2, (B, the pathlength, equals 1) can be determined from the slopes of the calibration curves, ultimately leading to the calculation of CX and CY using equations 3 and 4 below:

A1 = ∈X1BCX + bX1 + ∈Y1BC Y + bY1 Eqn. 3

A2 = ∈X2BCX + bX2 + ∈Y2BC Y + bY2 Eqn. 4 The procedure contains 4 parts:

1. Solution preparation. 2. Determining λ1 and λ2 . 3. Preparing standard curves for both species, Cr3+ and Co2+ at λ1 and λ2 and

calculating molar absorptivity values at both wavelengths for each species. 4. Determining the concentration of Cr3+ and Co2+ in the mixture given.

Dispose of all Cr3+ and Co2+ in the appropriate WASTE bottles.

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PROCEDURE:

1. Solution Preparation: Stock solutions of Cr(NO3)3, at 0.1000 M, and of Co(NO3)2, at 0.300 M are available in the burets in the prep room. Prepare a series of solutions for each following the scheme below using six 50-mL volumetric flasks:

Stock solution: 0.1000 M Cr(NO3)3

50.00mL Volumetric Flasks

Flask 1 Flask 2 Flask 3 Cr 0.0500 M Cr(NO3)3 0.0250 M Cr(NO3)3 0.0125 M Cr(NO3)3 Co 0.1000 M Cr(NO3)2 0.0750 M Cr(NO3)2 0.0375 M Cr(NO3)2

2. Obtain the absorbance spectra of the 0.0250 M Cr3+, the 0.0750 M Co2+, and a solution containing the same concentrations made by mixing equal parts of the 0.0500M Cr3+ and the 0.1500 M Co2+ solutions. Plot all these spectra on one graph. Verify that the absorbance at each wavelength of the mixture is the sum of the absorbance at each wavelength of the individual solutions.

3. Select two wavelengths for analysis (which ones provide maximum sensitivity?).

Measure absorbance versus concentration at each wavelength for the series of

Add 25.00-mL stock solution

Fill to line with

deionized water,

mix well, then transfer

25.00 mL via a volumetric

pipet into the next flask

Fill to line with deionized water, mix well, then transfer 25.00 mL via a volumetric pipet into the last flask

Fill to line with deionized water and mix well

58

Cr3+ and Co2+ solutions. Measure the absorbance of the unknown mixture at the same wavelengths. If the absorbance of your unknown is greater than 0.8 at either wavelength, prepare a diluted solution using 4.00 mL of the unknown solution and 4.00 mL water (using pipets). Measure the absorbance of the diluted solution at both wavelengths.

4. Plot the absorbance versus concentration data for each element at each

wavelength – producing 4 calibration lines. Determine the equation for each line as well as the correlation coefficients. Using all the data acquired, calculate the concentration (molarity) of Cr3+ and Co2+ in your original unknown solution.

For example: If the following results were obtained:

Species Linear Regression Results λ 1

Linear Regression Results λ 2

X

slope =∈X1B =1.37 x 103

intercept = bX1 = 5.00 x 10-4 slope = ∈X2B = 2.65 x 102

intercept = bX2 = -5.00 x 10-4

Y

slope =∈Y1B =1.33 x 103

intercept = bY1 = - 5.00 x 10-4 slope =∈Y2B = 5.87 x 103

intercept = bY2 = 5.00 x 10-4

Unknown

Absorbance at λ 1 0.765

Absorbance at λ 2 0.414

Several approaches can be taken to solve for CX and CY. ( For this experiment, assign X and Y to [Cr3+] and [Co2+]). Shown here is the method of solving simultaneous equations which involves combining the equations, after manipulation, in order to solve for one term and then using that value, substituting and solving for the other.

Substituting the above information into equations 3 and 4 and forming 2

equations with 2 unknowns: (λ 1 ) 0.765 = (1.37 x 103)CX + (5.00 x 10-4) + (1.33 x 103)CY + (-5.00 x 10-4)

(λ 2 ) 0.414 = (2.65 x 102)CX + (-5.00 x 10-4) + (5.87 x 103 )CY + (5.00 x 10-4) In order to remove one variable and define the other, multiply by a factor that will produce the same coefficient for one of the terms and then subtract. To do this, the top equation is multiplied by 2.65 x 102 and the bottom equation by 1.37 x 103 producing: [2.027 x 102 = 3.631 x 105CX + 0.1325 + 3.525 x 105C Y - 0.1325] - [5.672 x 102 = 3.631 x 105CX – 0.685 + 8.042 x 106C Y + 0.685] Subtracting produces: -3.645 x 102 = -7.690 x 106C Y (Only one variable!)

59

so:

then, substituting the value for CY in one of the original equations allows solving for CX: [2.027 x 102 = (3.631 x 105CX) + 0.1325 + (3.525 x 105)(4.74 x 10-5) - 0.1325] so: This sheet may be used to help organize your data. Remember to account for dilution if you diluted your unknown mixture.

λ 1 =

λ 2 =

Concentration in original unknown:

Species at λ 1 at λ 2 Molarity Cr3+

slope = ∈Crλ1B = intercept = bCrλ1=

slope = ∈Crλ2B = intercept = bCrλ2 =

[Cr3+] =

Co2+

slope = ∈Coλ1B= intercept = bCoλ1 =

slope = ∈Coλ1B= intercept = bCoλ2 =

[Co2+] =

Absorbance of Unknown (Diluted or Undiluted?)

at λ 1

at λ 2

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Experiment 9 Determination of Aspirin Using a Back Titration

Many reactions are slow or present unfavorable equilibria for direct titration. Aspirin is acetylsalicylic acid, a weak acid, that also undergoes slow hydrolysis: Thus, each aspirin molecule ultimately reacts with two hydroxide ions, albeit at different rates. By adding hydroxide ion in excess, the reaction can be driven toward completion. The amount of aspirin present can then be calculated by determining the amount of hydroxide ion consumed. This can be accomplished by a back titration with HCl of the unused NaOH.

(fast) + OH -

+ H2O

+ OH - (slow)

+ CH3COO

OH

Procedure: 1. Preparation of approximately 0.10M HCl solution

Calculate the amount of concentrated (12 M) HCl you will need to prepare 250 mL of 0.10M HCl. Check your answer with your instructor before proceeding. Measure this volume of acid using a small graduated cylinder. Partially fill a 250 mL volumetric flask with water and gradually add the acid to it. Mix well, adding more water until the total volume is met. 2. Standardization of the HCl solution

Thoroughly mix the HCl solution. Clean, rinse, and fill a burette with the approximately 0.1M HCl solution. To each of three clean, labeled Erlenmeyer flasks, add approximately 50 mL distilled water and three drops of phenolphthalein indicator. Use your HCl burette to add approximately 25, 30, and 35 mL of acid to each Erlenmeyer flask, recording exactly how much you added – to the second decimal place. Mix each flask thoroughly. Clean, rinse and fill a second burette with the approximately 0.1M NaOH stock solution – recording the exact concentration of the solution. Titrate the three HCl solutions with the NaOH to the phenolphthalein endpoint. 3. Aspirin Samples

Accurately record the weight of a group of three aspirin tablets so that you can determine an average tablet weight. Use a mortar and pestle to crush enough tablets to produce about 1 gram of powder. Place approximately 0.30 g of the ground aspirin tablets into three different Erlenmeyer flasks, noting the exact mass for each. To each, add 20 mL of ethanol to help dissolve the sample. Add three drops of phenolphthalein indicator. Mix thoroughly. The solutions will probably remain cloudy with a few undissolved solids. 4. Aspirin Titration

Titrate one of the aspirin solutions with NaOH to the first permanent cloudy pink color. This endpoint indicates the completion of the acid-base reaction. The hydrolysis step also consumes one mole of hydroxide per mole of aspirin and so for a complete titration we will need to use a total of twice the amount of NaOH that you have already added, plus we will add some excess NaOH to ensure complete reaction. Double the amount of NaOH you already added plus an extra 10 mLs. Record exactly the amount of base you added. Add this same amount to each aspirin solution. 5. Heating the reaction to completion

Add two or three boiling chips to each flask and heat in a hot water bath to speed up the hydrolysis reaction. Avoid boiling, because the sample may decompose. While heating, swirl the flasks occasionally. After 15 minutes, remove samples from the water bath and cool for 5 minutes. If the solution is colorless, add a few more drops of phenolphthalein. If it remains colorless, add 10 mL more of NaOH and reheat. Record exactly the amount of extra base added.

2

6. Back titration with acid

The only base remaining in each flask will be the excess base that has not reacted with the aspirin. Using your burette with your HCl solution, titrate the excess base in each flask with HCl until the pink color just disappears, indicating the endpoint. 7. Waste Disposal

All solutions may be disposed of down the drain. Dispose of the HCl down the drain in the hood, the base in the sinks in the lab. Flush both with plenty of water. Calculations:

To determine the average mass of aspirin in each tablet the amount of hydroxide ion consumed can be determined by subtracting the remaining amount from the initial amount added. This can then be related stoichiometrically to the amount of aspirin present in each titration and eventually to the amount per tablet. Example: First, the molarity of the HCl solution must be determined. If, for example, it required 27.58 mLs of stock NaOH solution (0.091M) to reach the endpoint with a 25.00 mL portion of the HCl solution, then: (Repeat for each titration and determine the average HCl molarity.) Second, determine the total amount of NaOH added to the aspirin samples. If, for example, 44.00 mLs of 0.091 M NaOH were added to each sample, then: Third, determine the amount of NaOH remaining after both reactions with aspirin are complete. So, if for example, it required 9.00 mLs of HCl (0.1004M) to reach the endpoint with the NaOH remaining in the aspirin solution, then: Fourth, calculate the amount of NaOH that did react with the aspirin:

3

So: Fifth, determine the amount of aspirin per tablet by using the mass of the sample used and the average mass per tablet. If, for example, the above data was obtained using 0.3024g of the aspirin tablet and the average mass of a tablet was 0.3580g, then: The above calculations should be performed for each aspirin sample analyzed. The average amount of aspirin (mg) should be calculated and used to determine the relative percent error. Sixth, if the manufacturer reports that each aspirin tablet contains 325mg of aspirin, and the average amount of aspirin per tablet was determined to be 331mg, then the relative percent error is: