Lab Manual 2009 by 27blgeE5


									          CHEM 230

 Quantitative Analysis

Laboratory Manual
            Fall 2009

Department of Chemistry and Biochemistry
       Southern Illinois University
         Carbondale, IL 62901
Table of Contents

Introduction ............................................................................................................................................... 3
Laboratory Grading Policies ................................................................................................................... 4
Experiment Schedule ................................................................................................................................ 5
Safety .......................................................................................................................................................... 7
Laboratory Rules ...................................................................................................................................... 9
Recommendations for Improving Laboratory Techniques and Data Quality .................................. 10
Experiment 1 ........................................................................................................................................... 12
Experiment 2 ........................................................................................................................................... 15
Experiment 3 ........................................................................................................................................... 17
Experiment 4 ........................................................................................................................................... 19
Experiment 5 ........................................................................................................................................... 21
Experiment 6 ........................................................................................................................................... 24
Experiment 7 ........................................................................................................................................... 28
Experiment 8 ........................................................................................................................................... 31
Experiment 9 ........................................................................................................................................... 41
Experiment 10 ......................................................................................................................................... 43
Experiment 11 ......................................................................................................................................... 45
Experiment 12 ......................................................................................................................................... 49
Experiment 13 (Optional Lab)............................................................................................................... 52


The purpose of the laboratory work in this course is to introduce you to the field known as analytical
chemistry. Specifically, you will learn in detail the wet chemical methods known as quantitative
analysis. This will acquaint you with the basic techniques and operations that are necessary to perform
precise analytical measurements. These operations will form the set skills that are necessary to succeed
in more advanced chemistry courses, such as physical chemistry and instrumental analysis. Additionally,
these skills are essential to anyone working as a laboratory technician in industry and are essential to
anyone that conducts research.

For most students, this will be the first time they will earn no credit for the act of completing an
experiment. Grades are determined by the accuracy of your analysis. The grade you receive for the
laboratory will be a reflection of the degree of skill you develop and the care in which you perform the
analyses. In many of the analyses that you perform, it will be possible to obtain answers that are within a
few parts per thousand of the correct answer. This is only possible, however, with careful and thoughtful

Each student will work independently in the laboratory and will be assigned a unique unknown sample.
For most experiments, we use Thorne-Smith unknowns that are commercially prepared and certified to a
high degree of certainty. For each experiment you will receive an unknown in glass vial based on your
student number. The TA will know which lot number corresponds to your student number. When you
submit you answer for grading, it will be compared to the certified answer. Your TA does not know the
value of your unknown.

Most students find the laboratory portion of the course relatively demanding in terms of time and effort.
Efficient use of the laboratory time requires that you be familiar with the principles underlying the
experiment; this allows you to know when measurements, calculations, and reagents can be made
approximately, and when you need to take time to be precise. This differentiation is key to a successful
experience in the course.

As you complete an experiment, you are responsible for learning the experiment in detail, including the
names and structures of all reagents, balanced chemical reactions, and calculations. Questions about
these experiments will be included in the periodic exams and in the final exam. While many students
tend to create and share formulas that make calculations easier, you must learn to derive the formula. It
is strongly recommended that you work out the calculations step-by-step and on your own. You will
remember how to do the calculations on a quiz or test if you do this. You might not remember the all-in-
one formula.

Sections of this manual are taken from a manual written by Billy Fairless with revisions by Matt McCarroll, Joe M. Davis, Robert B. Marquardt, Troy Seals,
Keith Smith & Albert White. Appreciation is also expressed to the students who have used this manual (or earlier variants) since 1993 and who have
provided many additional useful comments. Each year we continue to collect student comments and implement those we believe will improve this manual.

Laboratory Grading Policies
The following list includes common errors made in lab notebooks. This is not an exhaustive list.
-50 for each page torn out of the notebook
-20 (maximum) for missing or incorrect uncertainty calculations
-15 for missing data or data not entered directly in the notebook
-10 for missing or incomplete pre-lab summary at beginning of experiment
-10 for a new unknown sample or more unknown
-10 for missing or incorrect answers to pre-lab exercises.
-5 for omitted or illegible calculations
-3 for improper use of significant figures
-3 incorrect format for unknown value reporting
-3 table of contents not up to date in notebook.

Recalculation policy
If a calculation error is noted during grading, the student is allowed recalculate their answer
with a 5 point penalty. A second recalculation (on the same lab) is a 10 point penalty, and a
third is a 20 point penalty. The first recalculation is due within one week of the date of previous report
submission. Students should request due dates for the second and third recalculations, if needed, from Prof. Huff

Repeating Experiments
There is a major advantage of working ahead and turning lab reports as soon as you complete the experimental
work: there is ample time in the schedule to repeat an experiment if you are not happy with your experimental
results. There is a -10 point penalty for repeating your experiment (i.e., collecting additional data and
resubmitting your lab report for a grade). Please note: you must inform your TAs if you wish to repeat an
experiment at least 72 hours prior to the time when you plan to repeat the experiment. Failure to do this may
delay your work, as your TA will not have enough time to prepare stock solutions.

Lab Reports and Lab Notebooks
A lab notebook is any permanently-bound, sequentially-numbered, notebook with lined or quadrille pages.
-The first page must include a table of contents.
-Date each page.
-for mistakes—one line only
-must use black or blue ink (no pencil)
-right side of the pages is for text, data entry, and writeup
-the left side of the pages if for calculations

Each lab report (with the exception of experiment #1) is graded. Each lab report (with the exception of
experiment 8) is written directly in to your lab note book. Submit your lab notebook for grading by placing in the
box outside of Prof. Huff Hartz’s office (Neckers 291).

Lab reports should include the following sections:
1.Title of experiment
2. Purpose of experiment                               Sections 1 through 4 and calculations must
3. Experimental procedure                              be completed prior to beginning work on an
4. Answers to pre-lab question(s)                      experiment.
5. Data, observations, and calculations
6. Results and answers to questions (if applicable)

Due Dates for Lab Reports
Each lab report is due immediately after the experiment work and calculations are complete. Students are
discouraged from turning in several lab reports at the same time; this is a disadvantage because any instructor
feedback from prior lab reports cannot be incorporated into future lab reports. Lab reports that are turned in at or
just before the dates given below are at a significant disadvantage, and this practice is strongly discouraged.
Instead, students should turn in lab reports frequently, and as soon calculations are complete.

Important Dates
8/27        All students meet in the computer lab (Neckers 118) for the last hour of lab for a graded
10/8        After 5 PM, lab reports for Experiments 2, 3, and your first experiment assigned from
            Experiments 4 ,5, 6, and 7 will no longer be accepted for grading (this does not include
            repeated experiments or recalculations).
10/15       After 5 PM, the first recalculations for Experiments 2, 3, and your first experiment
            assigned Experiments from 4, 5, 6, and 7 will no longer be accepted for grading.
10/20       All students meet in the computer lab (Neckers 118) for the last hour of lab for a graded
            exercise (note: this date is subject to change, changes will be announced during lecture).
11/19       After 5 PM, lab reports from your second, third, and fourth experiments assigned from
            Experiments 4,5,6, and 7, plus Experiment 8, and your first assigned experiment from
            Experiments 9,10,11, and 12 will no longer be accepted for grading.
12/3        After 5 PM, the first recalculations your second, third, and fourth experiments assigned
            from Experiments 4,5,6, and 7, plus Experiment 8 and your first assigned experiment
            from Experiments 9,10,11, and 12 will no longer be accepted for grading
12/08       Last day for experimental work in lab.
12/10       Lab checkout and cleanup (attendance is mandatory)
12/10       After 5 PM, lab reports for Experiments 7, 8, 9, 10, 11,12, and 13 and all repeated
            experiment lab reports will no longer be accepted grading.
12/17       After 5 PM, recalculations for Experiments 7-13 will no longer be accepted for grading.

Experiment Schedule
In order to stay on schedule and complete all experiments, developing time-management skills is crucial.
The following table is a guide, and it indicates when you should begin each experiment. Students
should work to follow the schedule below. Students are encouraged to begin experiment work before
the start dates listed below and if equipment is available. However, students must begin the experiments
in the order given on page 6 of this manual unless consent of instructor has been obtained.
Experiment Number Start Date*                          Experiment Number Start Date*
1                          8/25                        8 (Take home)             10/13 (Nov. 19/ Dec. 3)
2                          9/3 (Oct. 8/15)             9, 10, 11, or 12          10/22 (Nov. 19/ Dec. 3)
3                          9/10 (Oct. 8/15)            9, 10, 11, or 12          10/29 (Dec. 10/17)
4, 5, 6, or 7              9/22 (Oct. 8/15)            9, 10, 11, or 12          11/5 (Dec. 10/17)
4, 5, 6, or 7              9/29 (Nov. 19/ Dec. 3)      9, 10, 11, or 12          11/12 (Dec. 10/17)
4, 5, 6, or 7              10/6 (Nov. 19/ Dec. 3)      13 (Bonus)**              11/19 (Dec. 10/17)
4, 5, 6, or 7              10/13 (Nov. 19/ Dec. 3)     -                         -
*The dates in parentheses are the last day that these labs will be accepted for grading/recalculations.
**The bonus lab (13) can be used to replace the lowest score from experiments 2-12. The bonus lab
should be attempted only after all other experimental work is complete. Students cannot use the bonus
lab to replace a lab report that was not completed.

Students will be assigned a student number by their TA on the first day of lab during check-in. Students
should find their student number in the left-hand column and plan to complete the labs in the order given

Number                                   Experiment Number
1             1     2     3      4     5     6      7    8       9     10    11     12
2             1     2     3      5     6     7      4    8      10     11    12      9
3             1     2     3      6     7     4      5    8      11     12     9     10
4             1     2     3      7     4     5      6    8      12      9    10     11
5             1     2     3      4     5     6      7    8       9     10    11     12
6             1     2     3      5     6     7      4    8      10     11    12      9
7             1     2     3      6     7     4      5    8      11     12     9     10
8             1     2     3      7     4     5      6    8      12      9    10     11
9             1     2     3      4     5     6      7    8       9     10    11     12
10            1     2     3      5     6     7      4    8      10     11    12      9
11            1     2     3      6     7     4      5    8      11     12     9     10
12            1     2     3      7     4     5      6    8      12      9    10     11
13            1     2     3      4     5     6      7    8       9     10    11     12
14            1     2     3      5     6     7      4    8      10     11    12      9
15            1     2     3      6     7     4      5    8      11     12     9     10
16            1     2     3      7     4     5      6    8      12      9    10     11
17            1     2     3      4     5     6      7    8       9     10    11     12
18            1     2     3      5     6     7      4    8      10     11    12      9

19            1     2     3      6     7      4     5      8    11     12     9     10
20            1     2     3      7     4      5     6      8    12      9    10     11
21            1     2     3      4     5      6     7      8     9     10    11     12
22            1     2     3      5     6      7     4      8    10     11    12      9
23            1     2     3      6     7      4     5      8    11     12     9     10
24            1     2     3      7     4      5     6      8    12      9    10     11
25            1     2     3      4     5      6     7      8     9     10    11     12
26            1     2     3      5     6      7     4      8    10     11    12      9
27            1     2     3      6     7      4     5      8    11     12     9     10
28            1     2     3      7     4      5     6      8    12      9    10     11
29            1     2     3      4     5      6     7      8     9     10    11     12
30            1     2     3      5     6      7     4      8    10     11    12      9
31            1     2     3      6     7      4     5      8    11     12     9     10
32            1     2     3      7     4      5     6      8    12      9    10     11
33            1     2     3      4     5      6     7      8     9     10    11     12
34            1     2     3      5     6      7     4      8    10     11    12      9
35            1     2     3      6     7      4     5      8    11     12     9     10
36            1     2     3      7     4      5     6      8    12      9    10     11

Red=spectrophotometer (experiment 12)
Blue=pH meter (experiments 6 and 11)


The Department of Chemistry and Biochemistry has prepared “Laboratory Directions for Chemistry
Students” that includes a section on safety and which may be obtained from Dan Parker in Room 225.
The general procedures in that document apply to this course as do the additional more specific safety
procedures listed below.

The EMERGENCY telephone numbers are:
      b) Urgent but not an emergency - Dial 453-4456 for the student health center care office.
      c) Other - Contact Mr. Dan Parker (453-6413) in Neckers, Room 225.

Additional Safety Rules for Chem-230
1. Always conduct yourself in a professional manner.
2. You must learn where the safety equipment is and how to use each item the first day in class. In the
     event of an emergency, you should use whatever you need to address the emergency. Again, you do
     not need to ask for permission to respond to an emergency. Usually, your response will be to advise
     your TA and then follow his/her instructions. As a general rule, and if time permits, students should
     not attempt to provide first aid but should concentrate on contacting a professional (phone numbers
     above) in that area.
3. No consumption of food (including gum) or beverages or application of cosmetics will be allowed.
4. You are not to perform any unassigned experiments.
5. Do not use your mouth to fill pipets.
6. If something is spilled on you, wash it off immediately with lots and lots and lots of water, and then
     report to the TA. Clean up the spill later according to instructions from the TA.
7. Uncontrolled long hair or clothing (loose sleeves, ties, jewelry) that might come in contact with a
     flame or become entangled in mechanical equipment will not be permitted. You will not be
     permitted to work in the lab without protection for your feet (no sandals, for example).
8. Never heat a closed system.
9. Never heat flammable materials with an open flame or near an ignition source.
10. Do not heat or mix anything near your face (or anyone’s face).
11. Review the hazards of all reagents for an experiment before you start, so you know how to respond
     to an emergency. The Materials Safety Data Sheets (MSDS) for each reagent we use are available in
     Dan Parker’s office on the second floor (Room 225 or ask at the Chemistry Office). You are
     encouraged to review any MSDS any time you have a question.
         You should also note that a considerable amount of safety information is on the reagent bottle
     labels. Read and understand them before you use the reagents. Refer any questions to your TA, your
     professor, or to Mr. Parker before you start. Safety related questions will be included on quizzes,
     tests and exams.
12. Do not rub your eyes or mouth with your hands. Your hands are frequently contaminated.
13. Protective clothing (lab coat) is not required, but highly recommended. Disposable aprons are
     available from the stockroom at no cost. You will not be permitted to work in the lab while wearing
     shorts without protection for your legs and feet.
14. You cannot tell when glass and other objects are hot by looking at them. Don’t get burned by trying
     to pick up something that is hot.

15. Do not store reagents near a sink or near the balance where they will be in the way and get knocked
    over. Return all reagents to their proper location as soon as possible after you are finished with them.
    Be sure everything is put up properly before you leave and that you have left nothing in the balance
    area, in a fume hood or at some other location.
16. Be sure you know where the safety equipment is located so you can find and use each item in an
    emergency (power off and the lab is dark, for example).
17. Be sure that, in an emergency, you know how to turn off all of the utilities (gas, water, electricity)
    you have been using.
18. Never attempt to identify an unknown by smelling or tasting it as recommended in some textbooks.
19. Use the appropriate safety equipment (safety shield, gloves, fume hood, shower, eye wash, etc.) and
    supplies as needed. Be sure any supplies you use are promptly replaced so they are available for the
    next emergency. It may be you again.
20. Label all containers containing chemicals with the name or formula of the material, the date and
    your initials.
21. Read all chemical labels prior to use. Be sure you know what you are using.
22. Do not store chemicals near non-compatible chemicals (acids with bases or oxidizers with fuels, for
    example) even for short periods of time.
23. Transport and dispose of all chemicals properly. If you are not sure how to do so, ask your TA.
24. Do not use chipped or broken glassware. Broken glassware will not be accepted at the end of the
    course and should be replaced during check-in or as soon as it is broken.
25. Do not operate electrical equipment with wet hands.
26. Do not block aisles or fire exits.

Laboratory Rules

All of the following rules will be followed at all times you are in the laboratory

1. Lab sections will meet only at the scheduled times. There will be no makeup sessions, so plan to use
     all of your available time effectively. If you do, you will have plenty of time to finish all of the
     required experiments, the optional experiments and to repeat some of the experiments you may not
     have done correctly the first time. If you do not come to the lab well prepared, and do not use your
     time effectively, you will have less time to do optional experiments or to repeat experiments and
     therefore, will be less likely to make a good grade. Planning and self-discipline are critical aspects of
     being a good chemist.
2. Most experiments should be done in triplicate. You must record all data directly into your lab
     notebook using the correct number of significant figures.
3. Use distilled water from your wash bottle for the final rinse of glassware that has previously been
     washed and rinsed with tap water. Use distilled water to prepare all solutions requiring water even
     when the directions only specify water. Do not rinse glassware at the distilled water tap; instead
     bring fill a squirt bottle with distilled water for rinsing at your bench.
4. Never assume glassware or other materials are clean.
5. Label all of your solutions and reagents to avoid mix up. Do not place or remove an unlabeled reagent
     or sample into/from the drying ovens. At a minimum, the label should have the name of the
     material in the container, the date, and your initials.
6. Store and use all concentrated acids and bases in the fume hoods. When you need to use dilute acids
     or bases, take a clean container to the hood, pour out the amount of the concentrated reagent you
     need and dilute it in the hood. Do not add water to concentrated acids, add the acid to water.
7. The equipment, glassware and supplies you use in this course are very expensive. You are
     responsible for all breakage of glassware and for any other damages above the normal wear.
8. It is essential that you keep all common work areas (such as fume hoods, balance areas, etc.) and your
     work area clean. All spills should be cleaned up and reported to your TA immediately. You should
     assume that all horizontal surfaces are contaminated with a corrosive and will destroy your clothing
     if you sit or rest your arms on them. Allow time to clean up at the end of each lab period.
9. You should be thoroughly familiar with each experiment before you begin labwork.
10. Unless directed by your teaching assistant or your professor, do not share your glassware, reagents,
     standards, or other chemicals with other students.

Recommendations for Improving Laboratory Techniques and Data Quality

1. Use of common reagents
a.) A common reagents is a reagent that is used by everyone. Do not contaminate them. To prevent
          -Don't stick a spatula into a common reagent. Instead, pour a small portion into a beaker.
          -Don't stick a pipette into a common liquid. Instead, pour a small portion into a beaker.
b.) It is better that a little goes to waste than contaminating an entire bottle of reagent.

2. Use of balances and massing solids.
a.) To begin, check that the balance surface is clean. Shut the sliding plastic doors to prevent air currents
from causing weight fluctuations. Zero the balance before you use it. Make sure you can get a steady
reading. Weigh by difference (see below). Leave the balance clean when you are done. Use the same
balance for all weighing of a single experiment.
b.) The method of weight-by-difference is best illustrated by an example. Place the material to be
weighed into a weighing bottle (small glass wide mouth bottle with glass top) and dry it if appropriate in
a drying oven with the top removed. Cool the container in a desiccator. When it is cool, zero the balance
and then place the weighing bottle with the material to be weighed on the pan of the balance. The weight
of the bottle when you start is the GROSS weight. After you have recorded the gross weight in your lab
notebook, pour some of the material directly (no spatula) into the receiving container. Weigh the
weighing bottle and the remaining contents again. The weight after you have dispensed an aliquot is the
TARE weight. Record the tare weight in your lab notebook. The sample weight is the difference
between the gross weight and the tare weight. If you are weighing out multiple aliquots, the tare weight
becomes the gross weight for the next aliquot.
c.) Always handle weighing bottles used in weight-by-difference measurements with tongs or gloves.
Fingers will leave oils and residue. Even a few mg can be enough to change an "A" grade to a "B" grade.
d.) When drying a sample or an unknown in the oven, prior to weighing it, allow the sample or unknown
to cool to room temperature in a dessicator. Otherwise, you'll simply rehydrate the sample or the
unknown and defeat the whole purpose of drying it (which is to remove adsorbed water than can bias the
weight). After it has cooled to room temperature, use your tongs to recap the weighing bottle, thus
preventing water from adsorbing onto the sample/unknown. Weigh it at your convenience and determine
the sample/unknown weight using weight-by-difference. The dessicator must contain anhydrous
Drierite to reduce water absorption. Some forms of Drierite contains an indicator. If so, blue =
anhydrous; red = hydrated.
e.) Weight-by-difference is not always essential, but is required for standards and samples. Sometimes, a
rough weight of a reagent suffices, when exact quantities aren't needed.

3. Dissolving solids/liquids.
a) Experiments sometimes fail because students don't make sure their reagents/samples are fully
dissolved or mixed. The manual will tell you if you need to heat something; if so, do so in the hood,
when required.
b) Otherwise, repeated inversions of a volumetric flask may be required to uniformly mix even easily
dissolved substances. When mixing two miscible liquids, be sure to invert a volumetric flask at least 3
times with shaking while it is inverted to obtain a homogeneous solution. If the liquids are slightly
insoluble, more inversions will be required. You can't mix a solution too much, only too little. Do not

put any solid into a volumetric flask. Dissolve all solids first and then transfer the liquid to a volumetric
flask. You never heat a volumetric flask.

4. Use of burets.
a) You can read a buret to ± 0.01 mL. If you only read to the nearest 0.1 mL, your data will not be as
precise as they could be.
b) If you have difficulty reading the buret, place a white piece of paper behind it. The contrast then is
sufficient to read it. You may find that a piece of white paper with a dark black line across it will
highlight the meniscus well also.
c) When you get near the equivalence point of a titration (and you may not know what volume this
corresponds to), open the stopcock slowly to allow only partial drops of titrant to emerge.
d) A clean buret dispenses liquid easily and crisply, and does not leave excessive "beads and residue" on
the walls above the meniscus.
e) Good analytical work requires rinsing a buret at least once with a small portion of the solution with
which it is to be filled. 10-15 mL is sufficient to rinse the walls; then discard the solution through the
stopcock and out the tip. Fill the buret with solution and drain it out the tip quickly to remove air
bubbles. There must be no air bubbles in the tip for good analytical work.

5. Volumetrics glassware and pipets in general.
a) All volumetric glassware is of two classes: To Contain (TC) or To Dispense (TD). These small
uppercase letters are on most volumetric glassware to identify them. Graduated cylinders and volumetric
flasks are TC items. Pipets are TD items.
b) Proper use of pipets is essential to their analytical utility. Fill a pipet with a pipet suction bulb after
the tip of the pipet is placed in the solution you need. Draw liquid up past the meniscus mark and set the
tip squarely and carefully on the bottom of the container. This will slow the rate that the liquid flows out
long enough for you to remove the pipet bulb and place your finger over the end where the bulb used to
be. By subtle movements of your finger, allow the liquid to drain out until the meniscus is on the scribed
mark of the pipet and completely seal the end with your finger to stop the flow out. Wipe the outside of
the pipet with a paper towel, move the pipet to the receiving container and allow the liquid to dispense
by gravity down the side of the receiving container. Hold the pipet in a vertical position with the tip
against the side of the receiving container. Do not try to get all the liquid out of the pipet. A TD pipet is
calibrated to deliver the appropriate volume and is made so some liquid is left in the tip of the pipet.

6. Quantitative Transfer.
a) Quantitative transfer of liquids requires pouring (decanting) down a stirring rod placed in the pour
spout depression of a beaker, or on the flat edge of any other glassware. This allows the liquid to run
down the side of the rod and not the side of the beaker. The inside of the beaker should be rinsed at least
three times with small portions of solvent (usually distilled water) and transferred in the same manner.
Successive, small portions are better than large, less numerous ones. The rod should be rinsed as well,
into the container, after the beaker is rinsed.
b) Quantitative transfer of solids can be tricky and usually requires more mechanical work like scraping
to be done. A rubber policeman works well for this. All other rules for liquids apply here as well. This
may take considerably longer than transferring of liquids only because of the sticky nature of some

Experiment 1: Preparation and Standardization of 0.10 Molar Solutions of NaOH and HCl

                 Hazardous Waste: No hazardous waste is produced in this experiment.

The purpose of this experiment is to prepare and standardize acid and base solutions that will be used for
subsequent experiments.

Prelab Exercise
1.) Write the balanced chemical equation for the reaction of a
    base (sodium hydroxide, NaOH) with an acid (potassium
    hydrogen phthalate, KC8H5O4, structure shown).

Preparation of KHP
To begin, fill a labeled weighing bottle about half full of potassium hydrogen phthalate (KHP) and dry it
for a minimum of one hour at 100°C and cool it in your desiccator (don’t close weighing bottle lid while

Preparation of 2.2 liters of 0.10 M NaOH
Prepare carbon dioxide free distilled water by boiling about 2300 mL of distilled water in a covered
beaker and then allowing it to cool. Put about 2200 mL of the boiled distilled water (not necessary to be
extremely accurate at this step) in a clean 2.5 L storage bottle. Add 13 mL of 50% NaOH solution. Cap
the bottle and mix thoroughly by shaking vigorously for one minute. Keep your storage bottles capped to
prevent contamination and evaporation. The solution will slowly react with CO2 and the glass, but
should remain stable for up to 6 weeks. A fluffy white precipitate may form in the bottom of your NaOH
solution. Do not to stir up the precipitate; it may adversely affect your measurements. If you later
question the accuracy of your standardized solution you can simply check it with one titration at any
time and recalibrate if it is not within the expected range.

Preparation of 2.0 L of 0.10 M HCl
Dilute 16 - 18 mL of concentrated HCl with 2.0 L of distilled water (not necessary to boil) in a clean 2.5
L storage bottle. Shake vigorously for one minute.

Approximate molarity of NaOH solution
The purpose of this step is to determine if the dilute NaOH solution prepared is within the desired range
of 0.08 and 0.12 M.

Weigh by difference into a clean 250 mL Erlenmeyer flask, 0.7 (±0.1) g of the dried KHP.
[0.7 g should require about 35 mL of a 0.10 M NaOH solution for complete titration.] You need not
worry about getting exactly 0.7 g, but you must know exactly how much you have. Dissolve and dilute
the KHP with about 75 mL of distilled water (the graduations on the flask are sufficient for this) and add
two drops of phenolphthalein indicator. The exact amount of indicator is not important, be consistent.

Exp. 1, Fall 2009                                                                                       12
Fill one of your clean burets with your 0.10 M NaOH solution and titrate to the end point. This should
require about 35 mL of solution. The purpose of this step is only to verify that your solution is close to
0.10 M, so you need not be extremely accurate (± 0.2 mL).

Calculate the approximate molarity of your solution (mol/L) of your NaOH. The molarity should be
0.10 ±0.02 M. If the molarity is not within this range, add an appropriate amount of either boiled
distilled water or 50% NaOH to the solution as necessary. Weigh out a new sample of KHP and titrate
with the new solution in order to determine the new approximate molarity. Repeat this process until the
NaOH molarity is within ±0.02 M of 0.10 M.

Determination of the approximate molarity of the HCl solution
The purpose of this step is to determine whether the HCl solution is between 0.12 and 0.08 M.

Fill one buret with the 0.1 M NaOH solution and the other with the ~0.10 M HCl solution. Take initial
readings on both burets. Deliver ~25 mL of the HCl solution into a clean 250 mL Erlenmeyer flask. Add
two drops of phenolphthalein and titrate with your NaOH solution until you reach the endpoint
(permanent pink color that persist at least 30s). It is not necessary to perform a perfect titration (±0.2 mL
is good enough), as your purpose at this point is to verify that the concentration is simply within the
desired range.

Use the approximate molarity of your standard base determined in the KHP titration to calculate the
approximate molarity of your HCl solution, knowing that at the end point, the number of moles of acid
is equivalent to the number of moles of base added. The molarity should be 0.10 ±0.02 M.

[If not in the desired range, adjust by adding distilled water or concentrated HCl. Determine the
concentration again and repeat the process until within the appropriate range.]

Once you are certain that your solutions are within the desired range of molarity, you will accurately
determine the concentration of the solutions by:

(1) determining the ratio of the NaOH concentration to the HCl concentration
(2) accurately determining the concentration of the NaOH solution by titrating samples of primary
standard KHP
(3) calculating the HCl concentration by using the ratio determined in (1) and the concentration
determined in (2).

Ratio of NaOH Concentration to HCl
Fill one buret with your NaOH solution and the other with your HCl solution. Take initial readings on
both burets to the nearest 0.01 mL. From the buret, dispense ~25 mL of HCl into a 250 mL Erlenmeyer,
add two drops of phenolphthalein indicator and titrate to a permanent endpoint with your base solution.
Now back titrate with the HCl until colorless and wash down the sides with distilled water from wash
bottle. Titrate with the base solution (probably only a few drops) until the solution is faintly pink. The
ideal endpoint is a faint pink that fades away slowly (~30 s) due to absorption of CO2 from the air.

Exp. 1, Fall 2009                                                                                         13
You should be able to determine the endpoint to within a fraction of a drop (partial drop titration). If the
endpoint is not satisfactory, you may back titrate acid and try again. As long as you don’t run out of
solution in your buret, the total volume doesn’t matter, it is the ratio of the acid and base added that
matters. Hence, you may back titrate and titrate until you are satisfied with the endpoint. Once you
reach a satisfactory endpoint, record the volumes delivered on the burets. Repeat this procedure until
you have at least three good ratios.

Calculate the volume ratios, (volume NaOH / volume HCl), the mean ratio and the percent relative
standard deviation (%RSD). The %RSD should be 0.5% or less. If your %RSD is greater than this, you
probably (1) did not adequately mix the solutions during preparation, (2) are using poor titration
technique, and/or (3) are not reading the buret closely enough. You must consistently estimate to ±0.01

Continue this process until you are able to get three consecutive titration ratios that have a percent
relative standard deviation of 0.5% or less. It may take several tries to achieve this, but don’t get
discouraged. You are developing good technique that you will need later to perform experiments

Standardization of NaOH versus KHP
Calculate the sample size of KHP that will require ~35 mL of your NaOH solution based on your
preliminary titration with KHP. Weigh (by difference) at least three samples of the dry KHP into clean
labeled 250 mL Erlenmeyer flask. The exact amount is not critical (within 0.1 g of target), but you must
know precisely the mass of the samples (± 0.0001 g). It is obvious, but important to carefully record the
mass and the sample number in your notebook.

Dissolve each sample in ~75 mL of water and add phenolphthalein. Titrate to a precise endpoint.
Calculate the molarity of the NaOH from each titration, the mean molarity and the RSD for the three
titrations. The %RSD should be 0.5% or less. If this is not achieved, you will need to titrate additional

Calculation of HCl molarity
Since you know precisely the ratio of the HCl and NaOH solutions, you can now calculate the HCl
molarity. You will need to know the relative uncertainty (i.e. reliability) of the HCl solution (recall that
RSD = %RSD/100). This is done by considering the error propagation from the two determinations used
to calculate the HCl molarity. The following equation can be used (eq 1):

For this experiment, you will not be graded since it does not involve determination of an unknown.
However, several labs are dependent on how precisely you determine the concentration of your standard
solutions. If this experiment is done improperly, it will be impossible to succeed on subsequent

Exp. 1, Fall 2009                                                                                           14
Experiment 2: Determination of Percent KHP in an Unknown by Titration with Standard NaOH

                 Hazardous Waste: No hazardous waste is produced in this experiment.

Prelab Exercise
1.) Why do you need to dry the KHP unknown before use?

You will receive your unknown in a vial that has your student number and the experiment number
printed on it. You should record these in your lab notebook. You will need to prepare your unknown for
analysis by drying it in a weighing bottle, just as you did with the pure KHP in the previous experiment.
[Hint: label your bottle since your unknown KHP and the pure KHP from previous experiment may look
the same, as well as most organic substances.]

You will be given 10 g of sample. This should be enough for at least five determinations. You will have
enough for a couple of overruns, etc. but not more. If you run out of sample, you may request more, but
it will be a 10 point penalty. In the “real world,” the sample is often a precious commodity that isn’t
unlimited. Realize that the number of replicates you are able to perform will depend on the
concentration of KHP in your sample. You will be able to run fewer replicates if you happen to have a
sample with a low concentration of KHP. The % KHP in your sample will be between 20-60%. You
will want to set up your analysis (sample size) so that the titration requires ~35 mL of titrant. This will
of course be dependent on the particular concentration of your sample, so you will want to estimate an
approximate sample size for the first titration, assuming your sample is somewhere in the middle (i.e.
40%). Using the standardized concentration of your base, calculate the necessary sample mass that will
require 35 mL of titrant (eqs 1-2).


You can then perform a single titration on a single sample to determine the approximate concentration of
your unknown. (Use proper technique, such as weighing by difference) If you happen to hit the end
point very precisely, you can use the answer in your final statistical analysis. You shouldn’t, however,
waste much time on this first titration as the main point is simply to get the approximate concentration in
order to set up the sample size for the rest of the titrations.

Once you know the approximate concentration of your sample, you can calculate the mass of sample
required to titrate to endpoint with 35 mL of your base. It is absolutely critical here that you properly
weigh by difference using your weighing bottle and that you accurately record to the full precision of the
balance the mass of each sample. These should all be within 0.1 g of your target mass. Since you are
unlikely to have each sample with exactly the same mass, you can use the ratio (volume of base / mass
of sample) to determine the approximate volume that will be required for each replicate titration.

You should accurately titrate at least three samples. If you have remaining sample and adequate time,
additional replicates can only help your determination (if done properly). [Hint: you can rapidly titrate

     Exp. 2, Fall 2009                                                                                    15
to within a milliliter or so of the expected endpoint and then do a careful drop-wise and partial drop
titration. In other words, use your time where it counts and be quick where it doesn’t.]

Calculate the percent KHP for each of your replicate samples. Using these values, calculate the mean
and relative standard deviation for the unknown. If you have a data point that is not known to contain a
gross error, but seems very much out of line with the other data, you can consider statistical justification
of omitting the data with the “Q-test”. You can find information on this is your text (Section 4-6). This
should be done with extreme caution, especially on very small data sets. It is very possible the datum
that looks out of line is actually balancing an opposite bias in the other data.

Once you are satisfied with your data, you will report a final answer for the determination with an
estimate of the uncertainty in the form of the absolute uncertainty and the percent relative uncertainty.
This should be written in your notebook and denoted such that the instructor can tell what to consider as
your reported value. An effective way of doing this is to draw a box around the reported answer. Make
sure your table of contents is up to date prior to turning in your notebook.


       Student #: 3
       %KHP = (42.3 ± 0.1)%

Exp. 2, Fall 2009                                                                                         16
Experiment 3: Volumetric Determination of the Percent Sodium Carbonate in an Unknown

                 Hazardous Waste: No hazardous waste is produced in this experiment.

Prelab Exercises
1.) Write a balanced chemical equation for the reaction that occurs in steps 1-3 of the procedure (p 18)
2.) Write a balanced chemical equation for the reaction that occurs in step 4 of the procedure (p. 18)

This method of carbonate analysis is a two step process, and your results in this analysis will depend on
the accuracy of your standardization procedure in Experiment 1. The first step is the neutralization of
the carbonate unknown with an excess of your standardized HCl solution. The second step is back-
titrating the excess acid with your standardized NaOH. Two moles of acid are required to neutralize
each mole of carbonate, which results in the production of carbon dioxide (eq 1). Carbon dioxide can be
removed by boiling the solution.


Note: unless mentioned otherwise, “water” always means distilled water.

Unknown Preparation and Estimation
You will be given 5 g of an unknown containing sodium carbonate. You should dry your unknown in a
weighing bottle for at least two hours at 140°C and cool in your desiccator.

It is necessary to know the approximate concentration in order to carry out the precise analysis, so you
will do an approximate titration with methyl orange indicator. Once the approximate concentration is
known, the more precise back titration method will be used for the actual analysis.

You unknown contains Na2CO3 and filler salts that do not react with acid. You should plan your
approximate concentration assuming 35% Na2CO3 by weight. Using the concentration of your HCl
solution and assuming a sample concentration of ~35%, calculate the amount of sample that will require
35 mL of your standard acid solution to reach the equivalence point.

Weigh (by difference, of course) one sample of your unknown that is close (± 0.1 g) to the calculated
mass into a 250 mL Erlenmeyer flask. Dissolve in ~70 mL of water and add 3 or 4 drops of methyl
orange indicator (pink in acid, yellow in base). The solution should be yellow at this point. Titrate with
your HCl solution to the first permanent color change to pink. The buffering action of the carbonate will
cause this result to be slightly inaccurate, but it is good enough to set up the back titration method. To
get the best endpoint, record the volume delivered at the first sign of a change to pink. Then add 0.5 mL,
and if the solution becomes “more pink,” record the second reading. Continue this process until the
solution doesn’t appear to become more pink with added acid.

     Exp. 3., Fall 2009                                                                                    17
Use this value to calculate the mass of sample that will require 30 mL of your HCl solution for
neutralization. (A ratio is sufficient for this, although it could be solved many ways.) Once you have
determined the desired sample size, weight out at least 3 samples into 250 mL Erlenmeyer flasks.

Unknown Determination
1. Add 70 mL of water and then add 30.00 mL of your HCl solution. If your sample mass
is slightly more or less than the calculated value, you should use proportionally more or less HCl. Then
add an additional 12 mL of the acid, which should result in ~12 mL of excess acid in the solution. You
can use your buret for this operation. As long as you have excess acid, it doesn’t matter exactly how
much add, as long as you know exactly how much you added.

2. Cover the flask with a watch glass and heat the solution gently on a hot plate until it boils, in order to
drive off the CO2. The solution should appear to effervesce before it actually reaches the boiling point
due to the release of the CO2. You should be very careful not to lose any of the material and should
probably add a glass bead or boiling chip to prevent bumping of the solution. [NOTE: be careful with
this hot, acidic solution as it will damage skin, eyes, etc…]

3. After the effervescent action has stopped and the solution has actually boiled, remove the flask from
the heat and allow it to cool. You should rinse any liquid on the watch glass back into the flask with
water. The solution needs to cool to room temperature. You can carefully run cool water on the side of
the flask, but care must be taken as any water spilled into the flask could compromise the analysis.

After the solution has cooled, add three or four drops of methyl red indicator (red in acid, yellow in
base). Since you added about 12 mL excess acid, the solution should be red. If it is yellow, then you did
not add enough acid. You can salvage the sample by adding additional acid until the solution is red, then
add 12 mL additional acid. You must keep track of the total amount of acid that was added. If you had to
add additional acid, you must boil and cool your solution again.

4. Once the solution is red at room temperature, you can titrate the excess HCl with your standardized
NaOH solution to the first permanent color change. (it will turn colorless then yellow) If you happen to
overrun this titration, you can add additional acid and then titrate again, noting the total volume of acid
added and the total amount of base used to titrate. If you do this, you should add enough acid to require
at least ~10 mL of titrant so the relative error will not be too large.

You can now proceed one at a time with your remaining samples. If you required an additional amount
of acid in the procedure, you can make adjustments based on this first accurate analysis so you will only
have to boil the sample once. You may wish to run as many replicates as sample and time allow.

Calculate and report the % Na2CO3 in each replicate of your unknown sample. Report the mean value,
the absolute uncertainty in the form the standard deviation, and the percent relative uncertainty in the
form of the percent relative standard deviation.

Exp. 3, Fall 2009                                                                                          18
Experiment 4: Gravimetric Analysis for Chloride

Hazardous Waste: The silver used in this experiment should be placed in a special container to be reclaimed. You
should have silver solutions in the filter flask from the precipitation and a silver chloride precipitate. Both go in the same
waste container.

Prelab Exercises
1.) Write the balanced chemical equation for the precipitation of silver chloride from silver ion and
    chloride ion.
2.) What errors in experimental procedure could cause systematic errors? List one error that leads to a
    higher than expected result and one error that leads to a lower than expected result.

This experiment is based on a gravimetric analysis method, which pertains to a method based on mass
measurement. Silver and chloride ions form a precipitate in solution that is relatively insoluble (~1.6 mg
AgCl per liter of water at 20°C). We can further decrease this by adding a slight excess of Ag due to the
common ion effect. High concentrations (~5 mM) should be avoided, however, due to the formation of
complex ions that are soluble (AgCl2- and AgCl32-). By determining directly the amount of AgCl
precipitate formed, the amount of soluble chloride can be determined.

If not already done, clean three sintered, medium porosity, glass crucibles by washing them with soap
and water and then, in the fume hood, soaking them in nitric acid for about five minutes. Draw the nitric
acid through the filter with a vacuum and then rinse both the filter and the rest of the crucible with tap
water, a dilute solution of base (use 0.1 M NaOH) and finally with repeated washings of distilled water.

Place the crucibles in a drying oven at about 160.0 °C for at least one hour or overnight. Cool the
crucibles in your desiccator for at least 30 minutes. Next, you must determine the “constant mass” of the
crucibles. [This part can be very time consuming, so you should not delay starting this experiment.] This
is an experiment where fingerprints really do count, as will rust and lint that might contaminate your
crucibles. Also be cautious of any powder from gloves that can leave residue. You should note the
amount of time that your crucibles cool in the desiccator and try to replicate it each time it is heated.
You should also use the same balance each time. You should work as quickly as feasible and mass the
crucibles in the same order each time.

Remove a crucible from the desiccator, replace the desiccator lid and place the crucible on the balance.
Record the first stable reading. It will probably slowly increase as it reaches equilibrium with the
atmosphere. You don’t want to let this happen, so record the first stable reading. Set the crucible aside
(not in the desiccator) and quickly mass the second crucible in the same manner as you did the first.
Repeat this with the third, etc. Dry the crucibles in the oven a second time for at least 30 minutes and
allow them to cool in the desiccator for the same amount of time as before. Mass the crucibles again, as
was done previously (same order). If the two mass measurements agree within ±0.0003 g, the crucibles
are at constant mass. If not, repeat the procedure until two consecutive mass readings are within the
required range.

While you are cleaning your crucibles, obtain and dry your unknown for at least two hours at 160 oC,
then cool it in a desiccator. Weigh out three portions of your unknown (you must calculate the amount

required to give you about 0.15 g of silver chloride precipitate per analysis) into 250-mL beakers. You
should assume your unknown contains ~55% soluble chlorides. Dissolve the portions completely in
~100 mL of 0.1 M nitric acid. Heat the solution to nearly boiling (but don’t boil) and keep the solution
hot during the addition of silver nitrate. Slowly add 0.2 M silver nitrate with stirring. (Caution - silver
nitrate reacts with protein and leaves a black residue in your skin if it comes in contact with your skin.)
Between small additions, allow the precipitate to settle from the surface and note when the addition of
silver no longer causes a precipitate. At this point add about 10% silver nitrate as excess. Continue
heating for about 10 minutes (do not allow to boil!) and allow the solution to cool for 15-20 minutes or
longer; a little longer is better so coagulation is more complete. Add a few additional drops of silver
nitrate to check for complete precipitation as soon as possible.

Filter each precipitate through one of the tared sintered glass filters as demonstrated by your TA. Do not
attempt to filter before the demonstration or until you are confident that you know how. Verify that you
are using a medium porosity filter, as fine porosity will take too much time and coarse porosity will
allow the precipitate to pass through. Wash the precipitate with 0.01 M nitric acid to remove all of the
silver nitrate.
[Note - silver chloride is more soluble in nitric acid than it is in water, so don’t use too much nitric acid (0.005 - 0.01 M is
recommended). However, the coagulated silver precipitate may go back to the colloid form and through the filter without the
nitric acid. Also, AgCl reacts with light to form silver metal (turning your precipitate a grayish-purple) and chlorine (which
may escape as a gas, giving low results, or react with any remaining silver ion, giving high results).]

If possible, first dry the crucibles containing the precipitate at ~110 oC for ~2 hours to prevent the
occlusion of water. Then dry for at least 30 minutes at 160 oC, cool in a desiccator, and weigh to a
constant weight as before using the same procedure and order (and balance). Repeat this procedure at
least twice (pre-drying step not necessary) until a constant mass is obtained.

Calculate and report the percent chloride in your unknown based on your replicates, including the
absolute uncertainty in the form the standard deviation, and the percent relative uncertainty in the form
of the percent relative standard deviation.

Exp. 4, Fall 2009                                                                                                            20
Experiment 5: Volumetric Determination of Arsenic Trioxide

Hazardous Waste: Aqueous arsenic solutions that are at a concentration less than 5 ppm
can go down the drain. Higher concentrations must be collected in hazardous waste containers.

This experiment is based on what is referred to as a redox (i.e. reduction-oxidation) titration. Rather than
being based on acid/base neutralization, this titration is based on electron transfer between chemical
species. If we haven’t yet covered this material in lecture, you will need to refer to your text in order to
effectively understand/perform this experiment. Make sure you understand the whole experiment and
have completed the following exercises prior to beginning the experiment.

Prelab Exercises
1. Write a balanced equation for the chemical reaction between arsenic trioxide and iodine.
2. What is the chemical structure of sodium thiosulfate? Is it an oxidizing agent or a reducing agent?

1. Determine the %composition of an unknown sample containing ~ 9% arsenic trioxide.
2. Learn about redox reactions and how to use then for quantitative analysis.
3. Develop better laboratory technique by completing a more challenging sequence of reactions.
4. Experience the challenges in dealing with the accumulation of errors from a multistep process.

Iodine titrations and reactions are historically very common. A three-step procedure described below is
widely used for standardization. In step 1, a known amount of iodine is produced by reacting potassium
iodate with potassium iodide. The resulting iodine is then titrated with sodium thiosulfate in order to
precisely determine the concentration of the thiosulfate solution (i.e. standardize). In step two, the
standardized thiosulfate is used to standardize a stock iodine solution. In step three, the standardized
iodine solution is then used to titrate the arsenic trioxide in an unknown sample.

Standard half cell reactions:

       IO3- + 5e- + 6H+ → ½I2 + 3H2O                         Eo = 1.196 V                           (1)

       I2 (aq) + 2e− → 2I-                                   Eo = 0.615 V                           (2)

       S4O62- + 2e- → 2S2O32-                                Eo = 0.080 V                           (3)

       AsO43- + 2e- + 2H+ → AsO33- + H2O                     Eo = 0.559 V                           (4)

Note: IO3- = iodate; S2O32- = thiosulfate; S4O62- = tetrathiosulfate; AsO43- = arsenate; AsO33- =

1. Standardization of sodium thiosulfate:
The half reaction with the higher reduction potential will go forward and will drive the half reaction with
the lower potential backwards (refer to text). From standard half cell reactions 1 and 2 above we see
that: a) reaction 1 will go forward, forcing reaction 2 backward (because 1.196 volts is larger than 0.615

volts) and b) reaction 1 must be multiplied by 2 and reaction 2 by 5 in order to balance the 10 electrons
lost and gained. Doing so and dropping the electrons gives eq 5:

        2IO3- + 12H+ + 10I- → 5I2 + I2 + 6H2O                                                       (5)

We combine the 5 iodine molecules from KI with the one from the KIO3, we get eq 6:

       2IO3- + 12H+ + 10I- → 6I2 + 6H2O                                                             (6)

Note: Equations 5 and 6 are identical; we have just combined all I2 sources.

Then dividing everything by 2 gives eq 7 which is the working equation for step 1 in the experiment.

       IO3- + 6H+ + 5I- → 3I2 + 3H2O                                                                (7)

The reaction in eq 7 shows that one mole of potassium iodate (KIO3), in an excess of potassium iodide
(KI) and in an acid medium, produces three moles of iodine. This iodine is then reduced to iodide by
thiosulfate in a subsequent reaction.

In triplicate, weigh out approximately 0.0500 - 0.1000 g of dried (100 oC for one hour) primary standard
grade potassium iodate and put it in a 500 mL Erlenmeyer flask with about 300 mL of distilled water
and an excess of potassium iodide (You should be able to calculate the amount. Do not use a large
excess but be sure that the potassium iodate is the limiting reagent.). Then add one mL of concentrated
HCl slowly, with a disposable pipet, drop by drop while mixing. Add a few drops of freshly prepared
starch indicator just before the endpoint as you titrate the solution with approximately 0.10 M sodium
thiosulfate to a colorless endpoint.

(1) The best results are obtained in this step by keeping the volume of iodine solution high (i.e. ~300
mL) and the weight of the potassium iodate lower than 0.12 g. Failure to do this may give incorrect
results since iodine is very slow to dissolve upon formation and may de-gas from the solution before you
titrate it, as evidenced by a purple vapor forming in the headspace above your solution in the flask. If
this occurs you must discard the sample and start over. Adding the HCl slowly forms iodine slowly,
reducing the risk that it will de-gas. The excess KI is also absolutely necessary as its presence
solubilizes the iodine.
(2) In excess oxidizing agent, e.g., IO3-, [I2]/[I-] is high and the solution is colored; in excess reducing
agent, the opposite is true. Furthermore, you should be able to approximate the endpoint from the
weight of KIO3, approximate molarity of Na2S2O3, and the stoichiometry.

Because it is more difficult to see the approach of the endpoint in the dark iodine solution, be careful not
to add too much titrant. Calculate the exact concentration of the sodium thiosulfate.

Exp. 5, Fall 2009                                                                                         22
2. Use of the standardized sodium thiosulfate to standardize a stock solution of iodine
From standard half reactions 2 and 3 above, we note that, in a reaction between iodine and thiosulfate,
reaction #2 will go forward, driving reaction #3 backward, and the electrons are balanced. The resulting
reaction, eq 8, is the working equation for step 2 in the experiment.

       I2 + 2S2O32 → S4O62- + 2I
                        −           −

In triplicate, pipet 25 mL of the stock iodine solution, previously prepared by your TA, into a 500 mL
Erlenmeyer flask. Add about 100 mL of distilled water, 1 mL of concentrated HCl, and titrate the
resulting solution almost to the endpoint. Add a few mL of starch indicator and complete the titration.
An iodine solution turns from a dark brown to a straw color just before you reach the endpoint. Note: I2
is the source of the color and when it is consumed in the reaction, the color fades. The approximate
molarity of the I2 solution is 0.05 M so you can approximate the endpoint. Calculate the concentration of
the stock iodine solution.

3. Use of the standardized stock iodine solution to titrate the arsenic in an unknown
In a reaction between iodine and arsenite ion (AsO33-), a review of standard half cell reduction reactions
2 and 4 above shows the electrons to be balanced and that reaction #2 will go forward. The resulting
equation is

       I2 + AsO33 + H2O → 2I- + AsO43- + 2H+
which is the working equation for step 3 in the experiment. In this reaction, one mole of iodine reacts
with one mole of arsenite ion to form iodide ion and arsenate ion (AsO43-).

In triplicate, weigh out approximately 0.900 gram unknown samples containing arsenic trioxide (dried
for one hour at 100 oC) and dissolve in 50 mL of 2% NaOH solution. Two % NaOH solution is easily
and approximately made by adding 1 mL of concentrated NaOH or 1 g of the solid to 50 mL of distilled
water. The samples should dissolve in less than 15 minutes without heat. Add 10 mL of 20% sulfuric
acid and cool. Adjust the pH to slightly acidic by using 5 drops of methyl red indicator and either more
acid or base until the desired pH is attained. Next add an excess of sodium bicarbonate powder, which is
effectively 5 to 7 grams. The solution will immediately effervesce (bubble). Dilute with distilled water
to approximately 200 mL and titrate immediately to “near” the endpoint and then add starch indicator.
Titrate until the first sign of any blue appears persistently (why is the endpoint indicated by a blue

In order to calculate the percentage of arsenic trioxide in your unknown, you will need to work though
the stoichiometry of the dissolution and titration reaction. It is necessary that you correctly determine
the concentration of the standardized solution that you used in the analysis. Report your unknown as a
percentage of arsenic trioxide, the absolute uncertainty in the form the standard deviation, and the
percent relative uncertainty in the form of the percent relative standard deviation.

Note: The effervescing of solution is due to formation of carbon dioxide, which is critical to the buffer
solution conditions created by the sodium bicarbonate. Unless you titrate immediately, place a stopper
on the Erlenmeyer to slow the loss of CO2. If the endpoint is still too difficult to see, increase the dilution
volume of the next sample to 300-350 mL.

Exp. 5, Fall 2009                                                                                            23
Experiment 6: Determination of Percent KHP by Potentiometric Titration

                 Hazardous Waste: No hazardous waste is produced in this experiment.

This experiment is similar to Experiment 2, in that you will be determining the amount of KHP in an
unknown sample. It differs in that you will use a pH meter to determine the endpoint of the titration,
rather than an indicator. When done carefully, this method will allow you to very accurately determine
the endpoint of a titration.

Prelab Exercise
1.) If we consider the procedures for experiment 2 and for experiment 6, which provides better accuracy
    and which provides better precision in terms of %KHP determination?

Obtain a KHP unknown from your TA and prepare it for analysis as you did in the previous KHP
experiment. You should have ~5 g.

Obtain a pH meter, electrode, stirring motor and electrode stand from your TA and set them up on your
bench. If your meter has a standby mode, leave it in this position when you plug in the meter and any
other time you are not actually recording pH measurements. Be careful with the electrode bulb, as it is
relatively fragile. Always store the electrode in a pH 4 or pH 7 buffer when not it use. You will need to
obtain about 2-3 cm of pH 7 buffer solution in a small beaker or flask (50 - 100 mL). Let the electrode
“condition” while you continue in the experiment. Electrodes are expensive and fragile, so don’t bump
them on the desk or bottom of the beaker.

Calculate the sample size necessary to use 35 mL of your standardized base, assuming the unknown
contains ~40% KHP. Weigh out one sample of approximately the calculated mass into a 250 mL beaker,
dissolve in ~80 mL of water and put a clean stirring bar in the solution.

Two Buffer Calibration Method
 The two buffer method of calibration adjusts the meter to fit the Nernst equation, which should give a
linear response over a wide pH range. The “Calibrate” knob or button will adjust the meter to “zero
potential” (i.e. intercept adjustment) and the “slope” knob or botton will adjust the slope. [note: when
making measurements of sample that differ in pH by more than one unit, it is advisable to give the
electrode about a minute to equilibrate in the new solution before making a reading from the meter. A
summary of the procedure is below. Ask your TA to show you how to use the pH meter.

Obtain 2-3 cm of pH 4.00 and 7.00 buffers in small beakers with stir bars. If the pH meter has a
temperate adjustment, the temperature knob is set to the solution temperature (room temp.) Place the
electrode the 7.00 buffer, using a stir plate and bar to stir the solution. Adjust the calibrate knob or push
the calibrate button so that the reading is pH = 7.00 ± 0.02. Rinse the electrode with a stream of distilled
water and gently blot (not wipe) to remove the excess water from the tip with a Kimwipe. Next, place
the electrode into the pH 4.00 buffer, allowing for equilibration (about one minute). Note the pH
reading. If it does not read pH 4.00 ± 0.02, adjust the slope knob (or push the calibrate button) to 4.00.
Double-check the calibration by switching the meter to standby, waiting a few seconds, and then
switching it back to take another meter reading to verify that it still reads 4.00. Adjust the slope knob
again if necessary. Once a stable and correct value is displayed, measure the pH 7.00 buffer again

(rinsing and blotting the electrode between buffers) to verify a correct reading of 7.00. If adjustment is
necessary to get the pH 7.00 reading, it is often necessary to adjust the pH 4.00 reading again as well.
The iterative procedure is continued until pH 4.00 and pH 7.00 readings are obtained without adjustment
of either knob.

Approximate Titration
It is best not to begin a titration unless there is ample time to finish it. This titration should not require
more than an hour.

Arrange the equipment so that the electrode and buret tip are in the beaker such that the electrode is
submerged and the buret tip is not submerged. The beaker should be sitting and stable on top of the stir
plate. Your buret should be filled with your base solution. It this experiment, it turned out to be very
convenient if you can start at exactly zero. You will also want to make sure that all of the buffer has
been rinsed off the electrode, as this would obviously affect your titration. Always rinse the electrode
with distilled water over a waste beaker. If necessary, you can blot the tip of the electrode with a
Kimwipe to remove excess water.

Once you are ready to begin the titration, immerse the electrode in the KHP solution and adjust the stir
plate such that the solution is mixed rapidly, but a vortex does not form. Make sure the stirbar does not
strike the electrode tip. At this point you can record the first pH reading that will correspond to 0.00 mL
of added titrant. Add the base at one mL increments, recording the pH and the precise buret reading. (Do
not worry if a drop is hanging on the tip. In this approximate titration it is not significant. Do not,
however, rinse the buret tip, as this would cause a dilution and significant error.) The solution and pH
meter should equilibrate very rapidly when you add titrant. Continue this process of added one mL and
recording the pH until the response is relatively constant at a value near pH = 11.

Plot the approximate titration curve. You can use computer software for this, or you could actually plot
it in your notebook by hand if you have graph paper pages. The purpose is to determine the approximate
endpoint and it doesn’t need to be extremely accurate. The pH values should be plotted on the ordinate
and the volume of titrant on the abscissa. Draw the best smooth curve through the points (don’t just
connect the dots) and identify the inflection point that corresponds to the endpoint. It will not be very
accurate, but this will tell you within ~1 mL of the endpoint. Calculate to the nearest percent the KHP in
the unknown sample.

Exact Titration
Based on information garnered from the approximate titration, calculate the sample size necessary to use
35 mL of your base solution at endpoint. Weigh out a sample based on this approximation. If, for some
reason the size deviates considerably from this goal, calculate where the endpoint will occur with your
specific sample size so you can anticipate the endpoint during the titration.

Set up the titration as you did for the approximate titration, however, add the same number of drops of
phenolphthalein indicator as you did in experiment 2. You will not use the phenolphthalein directly in
your determination, but you should make a note when the color change occurs during your titration.

Based on the approximate titration, you should be able to predict the general shape of the titration curve.
Therefore, you can spread out the data points in the areas that don’t matter (flat regions well before and

Exp. 6, Fall 2009                                                                                                25
after the endpoint) and take very closely spaced data when it does matter (inflection point). Specifically,
you will need to take very closely spaced data at the inflection point in order to precisely locate the end
point. A good way to this is described as follows:

1. When you have titrated to within 1 mL of the endpoint, take a precise buret reading. Then add one
drop and record the pH after the addition of the drop. Record this pH measurement as (1). Add another
drop and record the pH as (2). Continue this process keep a precise reading of the number of drops
delivered and the pH after each drop. Continue the drop wise titration until about 0.5 mL past the
endpoint (remember to note where the color changes).

2. After the last drop wise addition, take and record a precise buret reading. Then continue taking pH
measurements for a few mL at 0.5 mL increments, recording the buret reading as usual.

3. Calculate the volume of added titrant for the drop wise increments by considering the difference in
volume between the first and last drop wise addition and the total volume delivered based on the two
readings. In other words, determine the volume per drop for your buret. This turns out to be easier and
more accurate than taking buret readings at dropwise increments. An Excel spreadsheet is handy for this,
but not necessary.

34.15 mL - 32.58 mL = 1.57 mL

1.57 mL
          0.0785 mL/drop
20 drops

1 drop; 32.58 + 1(0.0785) = 32.659 = 32.66 mL
2 drops; 32.58 + 2(0.0785) = 32.737 = 32.74 mL
3 drops; 32.58 + 3(0.0785) = 32.816 = 32.82 mL
20 drops; 32.58 + 20(0.0785) = 34.15

Using some type of plotting software (Excel, etc.) and your data from the exact titration, produce the
following three graphs:
(1) All of the data on one graph
(2) Graph of expansion of endpoint region with scale division of the same precision as your buret
readings (e.g 0.10 mL divisions, 0.01 mL estimates). Indicate on this plot where the phenolphthalein
indicator changed color.
(3) Graph the region of the curve with drop wise additions in a derivative (slope) plot (see Section 11-5,
p 210 of your text)
For each addition, calculate pH/mL and plot these values versus the average of the two volumes. This
plot should have the same abscissa as the expansion plot in (2). Use this plot to accurately determine the
volume that corresponds precisely to the inflection point in the titration curve.

Using the endpoint determined in graph (3), calculate the percent KHP in your unknown. It should
correspond very closely (within ~1%) to the value determined in the approximate titration. If everything
seems good with your precise titration and is in general agreement with your approximate titration, you

Exp. 6, Fall 2009                                                                                        26
do not need to do a second precise titration. If you do not have confidence in your precise titration, you
may consider doing a second precise titration if time and sample allow.

Measure the volume difference between the pH meter endpoint and that observed by the indicator and
calculate the percent error for the indicator (assuming the meter is correct).
Report this error and state whether you feel phenolphthalein is a good indicator for KHP

Report the % KHP in you unknown based on the single precise titration, or the average of precise
titrations if more than one is performed. If you perform more than one titration, report the absolute
uncertainty in the form the standard deviation, and the percent relative uncertainty in the form of the
percent relative standard deviation.

Exp. 6, Fall 2009                                                                                         27
Experiment 7: Volumetric Determination of Percent Sodium Oxalate

Hazardous Waste: Solutions of permanganate at concentrations greater than 5 ppm must be collected as
waste. Solutions less than 5 ppm can go down the drain.

Prelab Exercise
1.) Why must potassium permanganate solutions be standardized by titration rather than by gravimetric

Potassium permanganate has been used in the analytical laboratory for more than a century. It is a strong
oxidizing agent and it is inexpensive. Furthermore, in most titrations the intense color of the
permanganate ion is sufficient to serve as the indicator. In strongly acidic media, permanganate
undergoes the following reaction (eq 1):

                               MnO4- + 8H+ + 5e- → Mn2+ + 4H2O                                         (1)

There are some disadvantageous of using permanganate. 1.) Permanganate will oxidize chloride ion to
chlorine, which limits the use of hydrochloric acid as a solvent and solutions of permanganate are not
stable. 2.) Potassium permanganate can be obtained in relatively high purity, but unfortunately it is not
feasible to prepare standard solutions by dissolving and dilute a massed quantity of KMnO4 due to the
fact that it will react with small amount of impurities present in the KMnO4, as well as traces of organics
in the water or on the surface of glassware. Results of such reaction produce manganese (IV) oxide,
which then acts as a catalyst for further oxidation. Therefore, it is necessary to standardize permanganate
solutions prior to use.

Preparation of 500 mL of 0.020 M KMnO4
1. Weigh ~ 1.65 g of KMnO4 and dissolve it in 550 mL of distilled water in an 800 mL beaker. Extra
care should be taken to use clean glassware and take precautions to exclude the introduction of any
organic matter. Do not take any measures to dry the KMnO4.
2. Heat the solution to boiling and keep it hot for ~1 hour and then allow it to cool. The final volume
should be ~500 mL. The solution should be covered and allowed to stand a couple hours, or overnight.
The process of heating hastens the oxidation of any organic matter contamination and enhances the
coagulation of manganese (IV) oxide.
3. Filter the solution through a clean sintered glass crucible with suction. Obviously, the filter flask will
need to be very clean since it will temporarily contain your standard solution. If there is question about
the cleanliness of the filter flask, you can rinse it will acidified permanganate (use about 1 mL of dilute
sulfuric acid and a few mLs of the permanganate solution). Wet all the surfaces with the solution that
will contact the permanganate and let it sit for a few minutes. Discard the solution and rinse with water.
If the filter becomes clogged with MnO2, just switch to a new crucible. The MnO2 can be removed later
with HCl.
4. Transfer the filtrate to a clean, 500 mL glass stopped bottle (You can use your 500 mL volumetric
flask). The standard permanganate solution should be stored in the dark. If you detect any solids that
form in the permanganate solution, you should filter again and standardize the solution again. If
carefully prepared and stored, the solution will be stable for at least a month.

Standardization of the permanganate solution
Anhydrous sodium oxalate (Na2C2O4) is available in high purity and is stable in air. For the
standardization of the permanganate solution, an accurately massed amount of the salt is dissolved in
sulfuric acid solution and titrated with the permanganate solution. The reaction between permanganate
and oxalic acid is (eq 2):

                       2MnO4- + 5H2C2O4 + 6H+ → 2Mn2+ + 10CO2 + 8H2O                                 (2)

When permanganate is first added to a solution containing oxalic acid, the solution immediately
becomes pink for some time, due to the slow rate of reaction. As the titration progresses the reaction rate
increases because the manganese (II) ion catalyzes the reaction. If the solution is heated the reaction
occurs rapidly near the endpoint of the reaction.

In this reaction, you will use an amount of a surfactant (cetyl trimethyl ammonium bromide, CTAB) to
catalyze the reaction at room temperature, making the analyses more convenient.

1.) Preparation of 0.050 M sodium oxalate solution
a. Weigh out 1.7 g of sodium oxalate and dry it at 100°C for one hour.
b. Cool and mass the weighing bottle plus contents
c. Transfer all of the sodium oxalate to a 250 mL volumetric flask and mass the weighing bottle a second
time to determine precisely the amount delivered.
d. Dilute to volume and use the mass delivered to calculate the molarity of the solution (moles of oxalate
per L of solution). Shake to mix well.

2.) Titration of oxalate solution with permanganate solution
a. Pipet a 25.00 mL aliquot of the standard oxalate solution into a 250 mL Erlenmeyer flask. Add ~100
mL of distilled water, 10 mL of 6 M H2SO4, And 10 drops of CTAB.
b. Assume your permanganate solution is 0.020 M and calculate the volume of permanganate solution
necessary to just react with the oxalate solution. Calculate 80% of this volume. (HINT: Remember that
the stoichiometry of the reaction of permanganate ion with oxalate ion is not 1:1.)
c. With continuous stirring, add the calculated volume of permanganate to the oxalate solution (using
you buret). The solution will probably undergo some interesting color changes and a temporary
precipitate may form, but after 5 minutes or so the solution should become colorless. If the solution is
brown instead of colorless, then you may have added more than the 80% or miscalculated. You will
need to discard this solution and begin again. Once the solution is colorless, continue the titration
(dropwise), adding additional titrant only after the solution is colorless. When a faint pink color persists
for 20 seconds or longer you have reached the endpoint. [Note: The dark colored solution will make
reading the buret meniscus very difficult. You will probably want to use the upper rim of the meniscus
in this case, which doesn’t affect the accuracy as long as you are consistent with your initial and final.]
d. Prepare a second 25.00 mL aliquot of oxalate solution for titration as in step (a). This time you can
add the permanganate rapidly to ~95% of the volume that was required in the previous titration. Once
the solution is clear continue the careful titration to endpoint. Repeat until you have at least 3-4
e. Because it takes a small amount of permanganate before your eye can detect the endpoint, you will
need to run a blank titration to determine the amount necessary to visualize the endpoint. Perform a

Exp. 7, Fall 2009                                                                                          29
titration as you did in the previous steps, except that you substitute water for the 25.00 mL aliquot of
oxalate solution. The titration volume will be very small since there is no oxalate in the solution. The
blank is the volume of permanganate that turns the solution pink, but did not react with the oxalate.
f. Take the blank into account (eq 3) and calculate the molarity and RSD of the permanganate solution.

                              mL of KMnO4 = titration volume – blank volume                       (3)

Determination of the percent sodium oxalate in an unknown
1.) Dry the unknown for one hour at ~100°C.
2.) Weigh out (precisely) 3 samples of ~1.0 g each into 250 mL Erlenmeyer flasks
3.) To each of the samples, add 125 mL of water (100 mL + 25 mL to mimic the standardization) and
prepare the solution for titration as you did in the standardization (sulfuric acid and CTAB).
4.) Calculate the volume of permanganate needed to reach the equivalence point if the unknown were
~20% sodium oxalate.
5.) Add the volume calculated in step 4 and then continue the titration as before.
6.) Using the results from the first unknown titration, you can calculate volume to titrate 95% of the
oxalate, which will allow faster titration of the remaining samples.
7.) Using the indicator blank determined previously, calculate the percent sodium oxalate in the samples
and calculate the mean and relative standard deviation (RSD) for the oxalate titrations.
8.) Calculate the relative uncertainty for the complete determination (eq 4), where RSDM is the RSD in
the permanganate concentration and RSD% is the RSD in the percent sodium oxalate.

                       relative uncertainty  RSDM   RSD% 
                                                                  2 12

Report the percent sodium oxalate in your unknown sample with the absolute uncertainty and the
percent relative uncertainty for the complete determination.

Note: Store your standardized permanganate solution in the dark if you cannot complete this experiment
in a single lab period.

Exp. 7, Fall 2009                                                                                       30
Experiment 8: Determination of the Equivalent Weight and Equilibrium Constant of a Known
Organic Acid by Potentiometric Titration

Note: this is a take-home exercise. No experimental lab data will be collected. Instead, you will
work through this exercise, answer the questions within this worksheet, and edit an Excel file.

Learning Objectives
 Analyze acid-base titration data using Microsoft Excel 2007.
 Identify whether an unknown organic acid is monoprotic, diprotic, or triprotic.
 Calculate the molecular weight of an organic acid.
 Determine the pKa values of an organic acid from experimental data.
 Identify an unknown organic acid from experimental data.

This experiment will acquaint you with a technique that is used to help identify organic acids. While it is
not possible to absolutely identify the substances with this technique, it will allow you to have a good
chance at identifying an organic acid from a relatively small list (25-50) of possibilities, by comparing
the measured molecular weight and the measured equilibrium constant(s) to the known values. You do
not collect any experimental data for this lab. However, you should read the appendix of these
instructions to review the procedure for data collection.

To begin this experiment, submit your email address, your name, and your student number to your
instructor at You will receive a Word version of this experiment and your
unknown as an Excel 2007 data file by return email as attached file. Make sure that your email account
can receive files 100 kB in size and that your filter settings do not reject your instructor’s return email.
Also, if you do not have MS Office 2007 installed on your computer, you should install a patch from the
Microsoft website or use the departmental computer lab.

The Excel data file that you receive shows the results of the data collection, and it includes the following
information: two columns of experimental data (the volume of NaOH titrated in first column and the pH
reads at each volume in the second column), the mass of the unknown organic acid titrated, and the
concentration of sodium hydroxide used to titrate the acid. You will use these data to determine the
whether the acid is monoprotic, diprotic, or triprotic, to determine the molecular weight of the unknown
acid, to estimate the acid dissociation constant(s) (pKa), and to determine the identity of the unknown

Graphing Procedure
Acid-Base Titration Curve:
Since you don’t know whether your sample is mono- di-, or tri-protic, you will need to plot the pH data
vs. NaOH volume data. In Excel 2007, highlight both columns of the titration data, click on the insert
tab, and select scatter with only markers from the chart menu. Next, select layout 1 from the chart
layouts menu and change the chart title to ‘Acid-Base Titration Curve’. Change the y-axis to ‘pH
Measured’ and the x-axis to ‘Volume of NaOH (mL)’. Cut and paste this graph into Sheet2 and rename
the worksheet ‘Graphs’ by double-clicking on the tab at the bottom of the spreadsheet. Change the name
of Sheet1 to ‘Data’ using the same procedure.

1. How many equivalence points do you observe?

First Derivative Plot:
First- and second-derivative plots are helpful for the determination of the equivalence point(s). A review
of section 11-5 in your textbook (pp. 208-211) is highly recommended.

To create the first derivative plot, you need to plot the slope of your data as function of the average
volume of NaOH delivered. To calculate the average volume of NaOH delivered, enter the following
formula into cell E7 in the Data worksheet:


Exp. 8, Fall 2009                                                                                         32
Verify that this gives you the average of the first two NaOH volume data points. Then, left-click on the
bottom right corner of cell E7 and drag your cursor down the column to copy this formula for all data

The slope of your data helps you identify the point(s) of maximum change in pH, which occur at
equivalence points. Remember that the slope is the change in y data divided by the change in x data. To
calculate slope, enter the following formula into cell F7:


This formula computes the change in pH over the change in NaOH volume for the first two data points.
Verify the calculation is correct and then copy the formula down the column for all data points. Now
that you have created two new columns of data, you are ready to make the first derivative plot. Highlight
the two columns and insert a scatter with smooth lines using the procedure you used to create the Acid-
Base Titration Curve. Change the title of this graph to ‘First Derivative Plot’, label the y-axis ‘First
Derivative’ and label the x-axis ‘Volume of NaOH (mL)’. Cut and paste this graph into the Graphs

2. What does a maximum indicate in your graph?

3. How many equivalence points do you observe?

Second Derivative Plot:
The second derivative plot calculates the slopes of the first derivative plot data, which provides
additional identification of your equivalence point(s). The second derivative plot is created in a similar
manner to the first derivative plot. Enter the following formula into cell H8 in the Data worksheet:


Verify the value is correct and then copy the formula down the column for all data points. Next enter the
following formula in cell I8:


Again, verify the value is correct and copy the formula down the column. Insert a scatter with smooth
lines plot and title the graph ‘Second Derivative Plot’. Label the y-axis ‘Second Derivative’ and the x-
axis ‘Volume of NaOH (mL)’. Cut and paste the graph into the Graphs worksheet.

4. What is indicated when your graph intersects the x-axis?

5. How many equivalence points do you observe?

Exp. 8, Fall 2009                                                                                          33
6. Do you think your organic acid is monoprotic, diprotic, or triprotic?

Second Derivative Expansion:
Once you have created the second derivative plot, copy it and paste a duplicate into the Graphs
worksheet. Right-click on the duplicate plot and select format axis from the menu. Adjust the maximum
and minimum values so that it expands the portion of the graph that intersects the x-axis (if your acid is
polyprotic, choose the final inflection point). Rename this title of this graph ‘Second Derivative

Determining the Molecular Weight of the Unknown Acid
7. What volume of NaOH was required to reach the end point based on the second derivative expansion

8. Using the equivalence point volume, and the sodium hydroxide concentration, calculate how many
   moles of NaOH were required to reach the end point.

9. Write a balanced chemical equation for the reaction of NaOH with…
      a. an organic acid, HA, assuming that HA is monoprotic.

       b. an organic acid, H2A, assuming that H2A is diprotic.

       c. an organic acid, H3A, assuming that H3A is triprotic.

10. How many moles of organic acid were present in the unknown solution assuming the acid is…
      a. Monoprotic?

       b. Diprotic?

       c. Triprotic?

Exp. 8, Fall 2009                                                                                       34
11. What is the molecular weight of the organic acid present in the unknown solution assuming the acid
       a. Monoprotic?

       b. Diprotic?

       c. Triprotic?

Determining Whether the Acid is Monoprotic, Diprotic, or Triprotic
Often, polyprotic acids will not give you clear breaks for each endpoint, and this can obscure the
identification of the acid. However, we can use our knowledge of acid/base equilibria to calculate the pH
at particular parts of the titration curve. For example, if we assume the acid to be monoprotic and
calculate the pH expected at various points we can compare the calculation with the measured value. If it
is considerably different, then our assumption is wrong and we can test another assumption.

Developing a Test for Monoprotic Acids:
Recall that a monoprotic acid (HA) exists in equilibrium with H+ and A-, and that this equilibrium can
be described by an equilibrium constant, Ka:

                                             HA  H+ + A-

                                                    [H  ][A  ]
                                             Ka 

The Henderson-Hasselbalch equation is merely a rearranged form of the equilibrium equation:

                                                             [A  ] 
                                                             [HA] 
                                            pH  pK a  log         
                                                                    
12. Compared to the total concentration of acid plus conjugate base, what is the fraction of the acid (HA)
    that remains in solution after it has been titrated halfway to its equivalence point with sodium

13. Compared to the total concentration of acid plus conjugate base, what is the fraction of the conjugate
    base (A-) that remains in solution after the acid has been titrated halfway to its equivalence point
    with sodium hydroxide?

14. Substitute your answers for #12 and #13 into the Henderson-Hasselbalch equation and solve for the
    pH in terms of pKa. Write the simplified equation.

15. Benzoic acid is a monoprotic acid with a pKa of 4.202. What is the pH of benzoic acid when it has
    been titrated halfway to its equivalence point with sodium hydroxide?

Exp. 8, Fall 2009                                                                                        35
16. Compared to the total concentration of acid plus conjugate base, what is the fraction of the acid (HA)
    that remains in solution after it has been titrated one-quarter of the way to its equivalence point with
    sodium hydroxide?

17. Compared to the total concentration of acid plus conjugate base, what is the fraction of the conjugate
    base (A-) that remains in solution at this point?

18. Substitute your answers for #16 and #17 into the Henderson-Hasselbalch equation and solve for the
    pH in terms of pKa. Write the simplified equation.

19. What is the pH of a benzoic acid solution that has been titrated a quarter of the way to its
    equivalence point with sodium hydroxide?

20. Compared to the total concentration of acid plus conjugate base, what is the fraction of the acid (HA)
    that remains in solution after it has been titrated three-quarters of the way to its equivalence point
    with sodium hydroxide?

21. Compared to the total concentration of acid plus conjugate base, what is the fraction of the conjugate
    base (A-) that remains in solution at this point?

22. Substitute your answers for #20 and #21 into the Henderson-Hasselbalch equation and solve for the
    pH in terms of pKa. Write the simplified equation.

23. What is the pH of a benzoic acid solution that has been titrated three quarters of the way to its
    equivalence point with sodium hydroxide?

Since you are attempting to identify an unknown acid, you do not yet know the correct pKa value(s) for
your particular acid. However, assuming your acid is monoprotic, you can be sure that it will have only
one pKa value. As a result you can rework these calculations (#12-23) by instead plugging in your
measured pH values and solving for the pKa. If your acid is, in fact, monoprotic, then the calculated pKa
value should be the same at all three volumes.

24. Based on the graph of your data, what is the equivalence volume (Vequiv.) of your acid assuming it is

Exp. 8, Fall 2009                                                                                        36
25. Fill out the following table based on the data for your unknown acid:

  Volume of NaOH Added                 Measured pH Value at              Calculated pKa Value from
                                             Volume                               pH value
          ¼ of Vequiv.

          ½ of Vequiv.

          ¾ of Vequiv.

26. Is your unknown acid monoprotic? Explain why or why not.

Developing a Test for Diprotic Acids:
If your acid is not monoprotic, the next step is to test if it is diprotic. A diprotic acid can exist as three
different species (H2A, HA-, and A2-) and is represented by the following two equilibria:

                                          (1) H2A  H+ + HA-          pKa1

                                          (2) HA-  H+ + A2-          pKa2

Furthermore, titrations of diprotic acids contain two different buffer regions that can be described by
simultaneous Henderson-Hasselbalch equations:

                                           [HA  ]                        [A 2 ] 
                                           [H A] 
                          pH  pK a1  log                                [HA  ] 
                                                           pH  pK a2  log         
                                           2                                      

It is helpful to think of titrating a diprotic acid as if it were two consecutive monoprotic titrations. For an
example, take a solution of a malic acid (pKa1 = 3.459, pKa2 = 5.097) that requires 10.00 mL to reach the
first equivalence point, and 20.00 mL to reach the second equivalence point. Therefore, the first buffer
region (0.00-10.00 mL) is described by the pKa1 equilibrium, and the second buffer region (10.00-20.00
mL) is described by the pKa2 equilibrium.

27. Consider the first buffer region for malic acid as described above:
       a. What is the pH of the solution when [H2A] = [HA-]?

        b. How many mL of NaOH are required to reach this point?

Exp. 8, Fall 2009                                                                                                37
28. Consider the second buffer region for malic acid as described above:
       a. What is the pH of the solution when [HA-] = [A2-]?

       b. How many mL of NaOH are required to reach this point?

To determine if your acid is diprotic, you will first calculate pKa1 and pKa2 from the Henderson-
Hasselbalch equations using your experimental pH data. With these two pKa values, you will then
calculate the expected pH at half of the endpoint volume using the simplified formula for a the pH of a
diprotic acid in its intermediate form: pH = ½(pKa1 + pKa2). This works because at half the endpoint
volume (ie, the first equivalence point), H2A has been converted to HA-. If the pH at half the endpoint
volume matches your experimentally measured pH at half of your endpoint volume, then your acid is
probably diprotic.

29. Assuming your acid is diprotic,
       a. What is the second equivalence point volume according to the graph of your data?

       b. What is the first equivalence point volume according to the graph of your data?

30. First, fill out the following table based on the data for your unknown acid:

  Volume of NaOH Added               Measured pH Value at             Calculated pKa Value at
                                           volume                             volume
         ¼ of Vendpoint.                                            pKa1 =

         ¾ of Vendpoint.                                            pKa2 =

31. Second, using the pKas that you calculated above and your graph, fill out the following table for your
    unknown acid:

       pH at ½ Vendpoint               pH = ½(pKa1 + pKa2

32. Do you think your unknown acid diprotic? Explain why or why not.

Exp. 8, Fall 2009                                                                                      38
Developing a Test for Triprotic Acids:
You can use similar logic as for diprotic acids above to test if your acid is triprotic. For example, the pH
should equal pK 1 at 1/6 the volume, at ½ volume pH should equal pK2, and pH equals pK3 at 5/6

33. Is your unknown acid triprotic? Explain why or why not.

34. Based on these three tests, do you think your acid is monoprotic, diprotic, or triprotic? Justify your

Identifying the Unknown Acid
Once you have determined whether the acid is mono-, di-, or triprotic, calculate the equilibrium
constant(s) and identify the molecular weight (as calculated above). The equilibrium constants that you
measure may differ from literature valued by a factor of 2 or more, due to activity effects. The molecular
weight may also vary some due to variable waters of hydration or water present as an impurity. Consult
your text book (table in appendix G) and attempt to identify your unknown based on the physical
appearance (see appendix at the end of this exercise), molecular weight and equilibrium constant(s). All
of the unknowns are relatively common organic acids that can be found in your textbook. If you do not
find a suitable candidate, you should seek the advice of the instructor. [Hint: In general, your MW
determination from the last endpoint will be more accurate than the pKa values.]

To be graded for this lab, complete the following table with the name of your organic acid, the
experimentally determined and literature values for pKa(s) and molecular weight.

Name of Organic Acid:

Parameter                                     pKa(s)                   Molecular weight, g/mol
Experimentally determined

Literature Value

35. Identify what you believe to be the most important concept(s) that you learned in this exercise.

36. What did you find to be the “muddiest point” (least clear) in this exercise?

Exp. 8, Fall 2009                                                                                            39
Submitting the Assignment
When you are satisfied with your work and results, clearly indicate the results to be graded in the table
above (molecular weight, equilibrium constant(s), and a reasonably justified “guess” as to the identity).

 To submit your lab for a grade, email your Excel file and Word file to your instructor for grading. You
must annotate your Excel file, label your graphs, and explain your calculations so that they can be
graded appropriately.

Appendix: Data collection procedure
An unknown sample that is a white crystalline solid is received in a numbered vial. The sample is not dried, because some
organic acids will decompose if heated. Also, some organic acids will have uncertain waters of hydration if allowed to
equilibrate with the desiccant in a desiccator. The unknown acid is stored in a stoppered weighing bottle in a desiccator. In
this procedure, the titration is performed with a pH meter similar to Experiment 6. Calibrate the electrode as follows.

Two Buffer Calibration Method The two buffer method of calibration adjust the meter to fit the Nernst equation, which
should give a linear response over a wider range than can be achieved with a one-buffer calibration procedure. The
“Calibrate” knob will adjust the meter to “zero potential” (i.e. intercept adjustment) and the “slope” knob will adjust the
slope. [note: when making measurements of sample that differ in pH by more than one unit, it is advisable to give the
electrode about a minute to equilibrate in the new solution before making a reading from the meter.

2-3 cm of pH 4.00 and 7.00 buffers is obtained in small beakers. The temperature knob is set to the solution temperature
(room temp.) The electrode is placed in the 7.00 buffer and adjust the calibrate knob until the reading is pH = 7.00 ± 0.02.
The electrode is rinsed with a stream of distilled water and gently blotted (not wipe) to remove the excess water from the tip
with a Kimwipe. Next, the electrode is placed into the pH 4.00 buffer, allowing for equilibration (about one minute) for it to
equilibrate and the pH reading is noted. If it does not read pH 4.00 ± 0.02, the slope knob is adjusted to 4.00. This is double-
checked by switching the meter to standby, waiting a few seconds, and then switching it back to take another meter reading
to verify that it still reads 4.00. The slope knob is adjusted again if necessary. Once a stable and correct value is displayed,
the pH 7.00 buffer is measured again (rinsing and blotting the electrode between buffers) to verify a correct reading of 7.00.
If adjustment is necessary to get the pH 7.00 reading, it is often necessary to adjust the pH 4.00 reading again as well. The
iterative procedure is continued until pH 4.00 and pH 7.00 readings are obtained without adjustment of either knob.

Titration Procedure Approximately 0.2 g of the unknown organic acid sample is massed, placed into a 250 mL beaker, and
dissolved using about 80 mL of distilled water. The solution of acid is titrated with with standard sodium hydroxide
solution, and a pH reading is taken every 0.50 mL. The volume data and pH data, along with the exact mass of the unknown
and the concentration of the sodium hydroxide are reported in the spreadsheet for analysis.

Exp. 8, Fall 2009                                                                                                               40
Experiment 9: Determination of Mg in a Solution by Ion-Exchange Chromatography

Hazardous Waste: No hazardous waste is produced in this experiment. The resin may be thrown in the
trash when the experiment has been completed.

Prelab Exercise
1.) The ion exchange resin used in this lab (Dowex 50X2-100) has a ion exchange capacity of 4 meq/g
dry weight, which means that for every gram of dry resin used, 4 milliequivalents of H+ can be
absorbed. How many meq of H+ can be absorbed by 5.0 g of this resin? How many meq of Mg2+ can be
absorbed by this resin?

Ion exchange is a process involving an interchange of ions of like sign between a solution and an
essentially insoluble solid in contact with the solution. The solid in this experiment is a synthetic ion
exchange resin in the form of small beads. Synthetic ion exchange resins are high molecular weight
polymeric materials containing a large number of ionic functional group (exchange sites) per unit
molecule. Synthetic ion exchange resins were first produced in 1935, and they have since then found
widespread laboratory and industrial application for water softening, water de-ionization, solution
purification, ion separation and quantitative determination.

The determination in this experiment is based on the exchange of Mg2+ ions for H+ ions. A cation
exchange process can be illustrated as follows (eq 1):

       xRSO3-H+ + Mx+ → (RSO3-)x + Mx+ + xH+                                                         (1)
Where xRSO3-H+ represents the number of exchange sites (x) within the resin occupied by protons and
M represents a cation of charge x (Mg2+ in this experiment). Mx+ will displace x protons from exchange
sites and will remain within, while the displaced protons will be eluted from the resin. The eluted
protons can then be titrated with your standard NaOH solution to determine the concentration of
displaced H+, which is then used to calculated the amount of Mx+ exchanged on the resin. The mobile
ions (H+ and Mg2+) occupying the exchange sites are known as counter ions. Generally, a more highly
charged ion will displace an ion of lesser charge from the exchange site, hence a Mg2+ ion will displace
two protons and electrical neutrality will be maintained.

Column preparation
You will use your buret as the ion exchange column. Insert a small wad of glass wool into the bottom of
the buret (above the stopper) to support the resin bed. Weigh 5.0 g of 50-100 mesh cation exchange
resin. A primary rule in pouring the column with ion exchange resin is never use dry resin.
Swelling of the resin upon wetting may create pressure that could break your glass column. [remember
that your buret is the most expensive glassware in your locker] Hydrate the resin by soaking it in a
beaker of distilled water for at least 10 minutes. After hydrating the resin, pour the resin-water slurry
into the column with the help of a stream of water from a wash bottle. The formation of air pockets must
be avoided. Wash the resin from the sides of your column down onto the bottom. The glass wool will
stop the resin and prevent it from entering the stopcock. Once the resin bed is in place, try not to stir it
up. In addition, a key point to remember is not to allow the liquid level to drain below the top of the
resin bed. The suggested flow rate from an ion exchange used to separate only a few components is from
1-3 mL per minute per square cm cross section of the resin bed. Measure the number of drops per mL
that come from the column you are using, count the number of drops necessary to fill a graduated
cylinder to 5 mL, and convert the rate flow rate limits from mL/min to drops/min; stay within these
limits when doing ion exchange with this column. Show these calculations in your lab notebook.

You must convert your resin to its fully protonated form by passing about one bed volume, a volume
approximately equal to the volume occupied by your resin, of 1.0 M HCl ( prepare 50.0 mLs of 1.0 M
solution) through the column. Rinse the column with two 10 mL portions of distilled water; collect the
third 10 mL portion of rinse and add one drop of methyl red; if the solution is yellow or turns yellow
upon addition of one drop of your stock NaOH solution then the column is sufficiently washed. If more
than, one drop is needed then continue rinsing with distilled water and repeat the process. Avoid stirring
up the resin bed when making additions. Drain each addition nearly to the level of the resin bed before
making the next addition.

You will be given an aqueous solution containing an amount of a magnesium salt. Quantitatively
transfer a 10.00 mL aliquot of the unknown solution to a 250 mL Erlenmeyer flask. Add 2-3 drops of
methyl red, and titrate with your base. Express the results as mmoles of HCl in the total sample. This
step is completed to determine the amount of titratable acid that may be present in your sample. You
may repeat replicate titrations if you wish, though you must make sure you retain enough solution for
the ion exchange portion of the experiment. Calculate the concentration of HCl in the solution (mg/mL).

Carefully place a 10.00 mL aliquot of your unknown solution on top of the prepared column. Start
collecting your aliquots in a clean Erlenmeyer flask. After the unknown has eluted nearly to the top of
the resin bed, wash with about one bed volume of distilled water. Repeat the washing two more times,
allowing each volume of wash to elute to the top of the resin bed before adding the next volume. All of
the effluent should be collected. Add 3-4 drops of methyl red to the effluent and titrate with your stock
base solution. Do not discard the solution. After the endpoint is attained, place the titration flask under
the column, and elute another one-half bed volume of distilled water into the solution you previously
titrated. If the indicator changes to red (acidic solution), continue the titration with your standardized
base to a yellow endpoint. Then repeat the process of eluting one-half bed volumes of distilled water and
titrate if necessary, until the eluted distilled water does not require more than one drop of the base to
change the indicator to yellow. The total amount of your base used in this process is then used to
         a) millimoles of H+ displaced from the resin.
         b) millimoles of Mg2+ remaining in the column on exchange sites.
         c) mg of MgSO4 in the total sample.

(HINT: remember that each Mg2+ displaces two H+ and that you must correct for the
amount of acid already present in your sample solution)

Repeat the procedure above for a total of three unknown determinations. Recharge the column between
each sample.

Finally, report the average concentration of Mg2+ in your unknown sample in mg/mL, the standard
deviation concentration, and the percent relative standard deviation.

Exp. 9, Fall 2009                                                                                       42
Experiment 10: Determination of Water Hardness

                  Hazardous Waste: No hazardous waste is produced in this experiment.

Pre-lab Exercises
The procedure described in Experiment 10 is optimized for mineral water samples containing
approximately 300 ppm total M2+ concentration, so that the endpoint volume is in the 25-mL range. The
endpoint volume gives a result with a percent relative error of order 0.1%.
1.) If distilled water, which has a total M2+ concentration that is 100 times lower than mineral water, is
analyzed by the same experimental procedure, what problem arises in the overall precision of the result?
2.) How should the experimental procedure be altered so that distilled water can be analyzed with the
same precision as mineral water?

The objective of this experiment is to determine the hardness of a mineral water sample, which we will
interpret as the total concentration of M2+. M2+ is defined as the sum of Ca2+ and Mg2+, since calcium
and magnesium are the typical dominant cations present in water.

Complexometric titration refers to methods of quantitative analysis utilizing titrations based on the
reaction of two species to form a soluble complex. The most common complexometric titrations are
those utilizing ethylenediaminetetraacetic acid (EDTA). EDTA forms stable 1:1 complexes with
divalent metal cations. While the range of metal complexes that form with EDTA is large (see Chapter
12 of Harris), under controlled pH conditions, the formation of EDTA complexes can be rapid and
complete, allowing for straightforward determination of cation concentrations.

EDTA is used as the disodium salt, which dissociates to give the H2Y2- form in aqueous solution.
Therefore the complexation reaction is given by equation 1.

        M2+ + H2Y2− → MY2− + 2H+                                                                        (1)
The metal ion competes with H+ for coordination sites on the EDTA molecule; hence the equilibrium is
controlled by the pH of the solution. Examination of equation 1 shows that basic conditions favor
complex formation. However, excessively basic conditions can result in the formation of metal
hydroxides or oxides, which can precipitate from solution. Furthermore, since the complexation reaction
results in the liberation of H+, the solution pH can change during the complexation titration. A good way
to address these problems is to use a buffer system. A good buffer to use in this type of analysis is an
ammonia/ammonium buffer system. The pH will be controlled in a range that favors EDTA
complexation and ammonia actually helps to stabilize metal ions so they will not precipitate.

In order for a titration to be useful, an indicator (In) is necessary that shows the endpoint of the titration.
Several indicators have been developed for use in EDTA titrations. The mechanism of action relies on
an indicator that has different colors in the free and complexed forms. If the formation constant of the
metal:In complex (MIn-) is significantly less than that of the MY2- formation constant, EDTA will
effectively compete for the metal species during the titration. Prior to the equivalence point, the
indicator will exist in the metal-bound state, because an excess of metal species exists. At the
equivalence point, however, the EDTA will effectively bind all the metal ions and the indicator will

exist in the free form, offering a visual signal of the endpoint. In general, the EDTA must bind the metal
10 times more strongly than the indicator for the system to work effectively.

In this experiment you will use the indicator, Eriochrome black T (EBT), which is red in the metal
complexed form and blue in the free form. Equation 2 shows the indicator reaction equilibrium.

       MIn (RED) + HY3- → HIn (BLUE) + MY2-                                                            (2)

Sample and Experiment Design
You will be given one sample that consists of mineral water that contains calcium and magnesium ions.
The mineral water sample has been degassed to remove dissolved carbon dioxide, which can alter the
pH of sample. You need to report the total concentration of divalent metals in each water sample. Since
you cannot differentiate between calcium and magnesium, you cannot report the concentration in terms
of mass. You can report, however, a molar concentration because you know the stoichiometry of the
complexation in 1:1. Therefore, you will report you findings in terms of molarity, [M2+].

1. Dry ~0.700 g of disodium EDTA for 2 hours at 80°C (EDTA may decompose if heated to a higher
temperature or a longer period of time) and cool in your dessicator.

2. Prepare an EDTA solution that is ~3 mM EDTA. This concentration should provide for an optimal
titration endpoint for samples that are ~200 ppm in M2+. You should prepare 500 mL of this solution,
which should be prepared quantitatively (i.e. weigh by difference and use a volumetric flask). If you
need to store the EDTA solution for more than a few hours, you should store it in a plastic bottle,
because the EDTA will leach metal ions from glass over time.

3. Prepare a pH 10 buffer solution immediately prior to use by dissolving 3.5 g of NH4Cl and 30 mL of
concentrated NH4OH in sufficient water to yield 50 mL of solution.

4. Quantitatively pipet a 10.00 mL aliquot of the water sample into a 250 mL Erlenmyer flask, add 2 mL
of the buffer solution, and add 2-3 drops of Eriochrome black T indicator. You want to add enough to be
able to visualize the red and blue forms, but too much will make for a broad color change at the
endpoint. This might take a bit of trial and error to get just right. You should practice using tap water
before you begin titrations on your unknown sample—titration error due to the indicator is the biggest
source of error for this experiment.

5. Titrate sample to the red to blue color change. If the endpoint is not satisfactory, consider the
        a. Adjust the amount of indicator.
        b. Check the pH of the solution (should be ~10)
        c. Warm the solution.

Report the average M2+ concentration in molarity, the standard deviation, and the percent relative
standard deviation.

Exp. 10, Fall 2009                                                                                           44
Experiment 11: Potentiometric Analysis for Iron

        Hazardous Waste: Solutions containing cerium must be collected as hazardous waste.

IMPORTANT SAFETY NOTE: ceric ion is a strong oxidizing agent, and it will oxidize chloride ion
(Cl-) to a highly toxic gas, chlorine (Cl2). For this reason, do NOT use any chloride containing-salts or
acids in this lab.

If this topic has not been covered in lecture yet, refer to Chapters 14 and 15 in the Harris text.

Prelab Exercises
1. What is a reference electrode?
2. What is an indicator electrode?
3. Since you don’t use a battery and nothing is plugged in to an electrical outlet, what causes the
electrons to flow? Can you think of any reason they might stop flowing?

1. Learn about the Nernst equation and how to use reference and working electrodes for quantitative
2. Successfully use the millivolt option on a pH meter
3. Learn more redox reactions and about different kinds of titration curves

In this lab, you will oxidize ferrous iron (Fe2+) to ferric iron (Fe3+) using a solution of ceric ions (Ce4+)
and will follow the change in potential of an electrochemical cell resulting from the reaction (eq 1):

                                  Fe2+ + Ce4+ → Fe3+ + Ce3+                                                (1)

This reaction is described by the standard half-cell reduction potentials (eqs 2-3):
                               Fe3+ + e- → Fe2+        Eo = +0.771 V                                       (2)
                                  4+           3+
                               Ce + e- → Ce            Eo = +1.44 V (in 1 M H2SO4)                         (3)

The higher reduction potential (1.44 V) dictates that reaction in eq 3 (cerium) will go forward as written.
The other half-cell reaction becomes an oxidation and is reversed. By recording and graphing the change
in potential during such a titration against the volume of titrant added, the endpoint can be precisely

The concentrations (activities) of reactants and products of a chemical reaction are related to the cell
potential by the Nernst equation (eq 4). The equation may be applied to each halfcell reaction or to the
entire cell reaction.

                                              -                                                      (4)

where E is the cell potential in volts at experimental conditions, Eo is the cell potential at standard
conditions, n is the number of electrons transferred during from the compound oxidized, and Q is the
reaction quotient (same form as an equilibrium constant).

For the iron half-cell reaction, the Nernst equation gives eq 5.


For example, if we were to titrate 50 mL of 0.1 M ferrous ammonium sulfate (Fe2+ ) with a solution
containing 0.1 M ceric ion to convert the ferrous ions to ferric ions, the equivalence point would occur
after the addition of 50 mL of the ceric solution. We would then have approximately (50 mL x 0.1 M =)
5 mmol of the ferric and the cerous (Ce3+) ions in 100 mL of solution less the small amount of the
products that reverse to ferrous and ceric ions. If we let x be the small amount of the ferrous and ceric
ion concentrations, then the concentrations of ferric and cerous ions will both be (5mmol – x) / 100 mL
or approximately 0.05 M.

At the equivalence point, the potential of each half-cell must be equal, therefore (eq 6):

       0.05916(log QFe + log QCe) = 1.44 - 0.771                                                    (6)

We can rearrange eq 6 as shown below (eq 7):


       Therefore Qreaction = 1011.3083 that Q = 2.0×1011 = [Fe3+] [Ce3+] / [Fe2+ ] [Ce4+ ]

Therefore, we can now define Keq as the following (eq 8):

If we assume that x is small relative to 0.05, then (eq 9):

       2.0 × 1011 x2 = 2.5 × 10-3                                                                   (9)

We can rearrange eq 9 and solve for x = x = 1.1 × 10–7 M = [Fe3+] = [Ce3+].

We can now use the iron half-cell reaction to calculate the cell potential at the equivalence point (eq 10).

       E = 0.771 + 0.05916 log (0.05 / 1.1 × 10–7) = 1.11 V                                         (10)

We can use the same iron half-cell Nernst equation to calculate the cell potential for volumes of titer
added prior to the equivalence point of 50 mL. For example, after adding 40 mL (40 mL × 0.1 M = 4
mmol) of the ceric solution, we would have produced 4 mmol of ferric
ions leaving 1 mmol of ferrous ions and both would be in (50 + 40 =) 90 mL of solution.

Exp. 11, Fall 2009                                                                                         46
Therefore, the concentrations would be 4/90 for ferric and 1/90 for ferrous or 0.04444 for
ferric and 0.01111 for ferrous. The cell potential would be (eq 11):

       E = 0.771 + 0.05916 log (0.04444 / 0.01111) = 0.771 + 0.05916 log 4 = 0.807 V                (11)

The potential at all other volumes between about 1 mL and 50 mL are calculated the same way. Side
reactions occur at 0 mL so the equation does not apply until some ceric ions have been added. A similar
problem is to be sure that all of the iron is reduced to the ferrous oxidation state and that none of the
ferrous ions are oxidized by other chemicals, such as the oxygen dissolved in the solution. In this
experiment, hydroxylamine hydrochloride is used to reduce the iron that has reacted with oxygen. The
titration with cerium should be done immediately after adding the reducing agent or some iron will be
lost to oxygen oxidation. Care should also be taken that the cerium does not react with an excess of the
reducing agent used to reduce the ferric ions to ferrous ions.

For volumes above the endpoint, it is more convenient to use the ceric half-cell reaction to calculate the
cell potential. Beyond the equivalence point, the concentration of Ce3+ will be (5 mmol / the total
volume of solution) because there was only 5 mmol of the ferrous ion available to convert ceric to
cerous ions. Any ceric ions added above that amount (50 mL of 0.1 M) will simply be diluted because
all the iron reacted with the first 50 mL. For those volumes above 50 mL, the Nernst equation is (eq 12)


At 50.01 mL, the voltage is up from 1.11 volts at the equivalence point to 1.28 volts and at 55 mL the
potential is 1.38 volts. The equivalence point is at the inflection point of the curve of the graph when the
voltage is plotted against the mL of ceric solution added.

Preparation of the ceric ion (Ce4+) solution
Ceric ion is available in the form a primary standard as the ceric ammonium nitrate salt
(Ce(NH4)2(NO3)6. Do NOT oven dry the ceric salt; it melts at 108 oC. You will prepare a 0.05 M
solution of ceric ion with a volume of 100 mL. Weigh out ~2.74 g (Ce(NH4)2(NO3) precisely and
transfer using good analytical techniques. The solvent for the ceric ion solution is 1.0 M H2SO4.
IMPORTANT: do NOT use HCl or HNO3 to dissolve the ceric salt.

Analysis of the unknown
Dissolve three aliquots of your unknown sample as received (do not dry) in a 250 mL beaker containing
about 150 mL of 1.0 M sulfuric acid. For initial calculation, you can assume that your sample is ~3%
Fe. You will need enough iron in each aliquot to react with about 20 mL of the 0.05 M cerium solution
you prepared according to the steps above.

Assemble the potentiometric system according to your TA’s demonstration. Titrate each solution by
adding an increment (about 1-2 mL) of cerium solution, reading the potential difference (i.e. millivolts
reading on the display), adding another increment and reading the voltage again, etc. The voltage will
not quickly reach an equilibrium value after each increment is added, so you must use judgment as to
how long to wait after adding an increment before you take a voltage reading. 2 seconds is usually
sufficient for a well-stirred solution. A suggested method is to quickly dispense an approximate amount,

Exp. 11, Fall 2009                                                                                          47
wait 2 to 3 seconds, and quickly look at the digital readout of the pH meter. Record the value and add
another increment. You should add smaller increments as you approach the endpoint but you do not
need to take additional time.

One of the advantages of a potentiometric titration is that you do not have to take the time to identify the
exact endpoint during the titration. Before the endpoint, the voltage will peak and then slowly decrease
after each addition of cerium. After the endpoint, the voltage will increase rapidly with the addition of
the increment of cerium as before and then drift higher rather than lower. Continue the titration for 5-10
mL past the endpoint.

After finishing all titrations, plot the voltages as a function of the volume of titrant added for your
unknown. Also construct a graph of the first and/or second derivative curve to better determine the
exact endpoint (see Harris text, 7th edition pp. 210-211.) Calculate the exact concentration of your
cerium and the percent of iron in your unknown. Report the %Fe in your unknown, the absolute
uncertainty and the percent relative uncertainty of your determination (propagate the uncertainty in the
ceric ion concentration and %Fe standard deviation). Include your graphs with your lab report.

Exp. 11, Fall 2009                                                                                       48
Experiment 12: Simultaneous Spectrophotometric Analysis of Caffeine and Benzoic Acid in a

                 Hazardous Waste: No hazardous waste is produced in this experiment.

[SAFETY NOTE: It is absolutely forbidden to taste anything in this laboratory. Even if the particular
item is not toxic, it is very possible that it has been contaminated with something else.]

Prelab Exercise
1.) Is an absorptivity value the same for all wavelengths?
2.) Why does the procedure for call for making the standard solutions and the unknown sample in acidic
    (0.01 M HCl) solution?
                                                                                    CH 3
                                                          H 3C

                                                            O        N

                                                                     CH 3

               benzoic acid, pKa = 4.202                           caffeine
                  122.12 g/mol                                   194.18 g/mol

Benzoic acid is added to many foods and beverages as bacteriostatic preservative, to inhibit bacterial
growth. The use of benzoic acid in beverages has been a concern recently, because benzene, a potent
carcinogen, was found in some citrus-flavored beverages that contain benzoic acid. Benzene formation
is attributed to reactions of ascorbic acid (vitamin C) and erythorbic acid (the d-isomer of ascorbic acid)
with benzoic acid.
(see and

Caffeine is a common central nervous system and metabolic stimulant that is added to many beverages.
The concentration of caffeine content varies in commercially-available drinks: 15-50 mg per 8 oz (tea),
20-80 mg per 12 oz (soda), 60-150 mg per 8 oz (coffee), and 100-400 mg/12 oz (energy drinks).

In this experiment, the concentration of benzoic acid and caffeine in beverages will be determined by
measuring the absorbance of these species.

According to Beer’s law, the absorbance (A) of a solution is equal to the absorptivity () times the cell
pathlength (b) times the concentration (c) at a given wavelength (eq 1):

       A=bc                                                                                        (1)

Since b is usually one cm, the working equation reduces to A = c, which is what we used in a previous
experiment (8). If a solution contains two (or more) absorbing species, then the measured absorbance at
a given wavelength will be the sum of all absorbances at that wavelength rather than just the absorbance

of a single analyte. By determining the absorptivities () of each species at each wavelength, we can
theoretically determine the concentration of each component in a mixture by measuring the total
absorbance at each wavelength and solving the resulting equations simultaneously. For the binary
mixture of this experiment, we have two equations (eqs 2-3):

        at 275 nm A275 = caffeine,275 [caffeine] + benzoic,275 [benzoic]                                     (2)


        at 230 nm A230 = caffeine,230 [caffeine] + benzoic,230 [benzoic]                                     (3)

where A275, A230 = total measured absorbances at 270 nm and 230 nm;
      caffeine,275, caffeine,230benzoic,275 benzoic,230 = absorptivities of caffeine and benzoic acid at 275 and
      230 nm;
      and [ ] = the unknown concentrations of the two species

If we measure A275 and A230 for our unknown and obtain caffeine,275, caffeine,230benzoic,275 and benzoic,230
from the calibration standards for the pure species, each measured at the two wavelengths, we can solve
the two equations for the concentrations of benzoic acid and caffeine.

You need to do four things.
1. Measure the absorptivity of caffeine at both 275 and 230 nm.
2. Measure the absorptivity of benzoic acid at both 275 and 230 nm.
3. Measure the absorbance of your unknown at the same two wavelengths.
4. Calculate the concentration of benzoic acid and caffeine in your unknown.

Prepare quantitatively (use volumetric flasks and pipets or burets; no graduated cylinders, beakers or
Erlenmeyer flasks) a set of at least five 100 mL calibration standard solutions (including a reagent
blank) (10 mL of each should be enough but you will probably need to prepare at least 50 mL since that
is the smallest volumetric flask you have in your desk) for caffeine and another set for benzoic acid.
The concentration of benzoic acid should be from about 0.5 mg/L up to approximately 20 mg/L. By now
you should know how to do this without detailed instructions.

Do not dry your benzoic acid or caffeine standard. They decompose at drying temperatures.

Put the benzoic acid and the caffeine standards in a solvent that is 0.01 M HCl. Use the 0.01 M HCl
solution for all of the calibration standards and unknown dilutions and to “zero” the spectrometer. You
will probably need about 2 L. Benzoic acid has a pKa of 4.202, and by making solutions in 0.01 M HCl
(~pH 2.0), we assure that 99% of the benzoic acid is in the protonated form.

You will obtain your unknown from your TA. The unknown is a commercially-available beverage
(either a soda or an energy drink) that contains caffeine and benzoic acid. To prepare your unknown for
analysis, dilute the unknown by a factor of 50, taking care to using the 0.01 M HCl as your solvent.

Exp. 12, Fall 2009                                                                                                   50
After you have been checked out on the UV/Vis spectrometer by your TA, measure the absorbance of
your benzoic acid calibration standards, your caffeine calibration standards, and your unknowns at 275
nm and 230 nm. Calculate the slope and intercept for the least squares best-fit lines and use the slopes to
first determine the four absorptivity values you will need and then the amounts of benzoic acid and
caffeine in your unknown. Be sure the concentrations of your calibration standards give absorbance
values within the 0.01 to 1.00 absorbance unit range (0.3-0.7 is better) and that the absorbance of your
unknown is in the linear part of your calibration curve. If not, adjust the concentrations accordingly.

Using the calibration standard data, graph the absorbance vs. concentration at 230 nm and 275 nm for
both benzoic acid and caffeine. Calculate the slope and intercept for the least squares best-fit lines for
the data at 230 nm and 275 nm. Use the four slopes that you obtain as the absorptivities shown in
equations 2 and 3 and calculate the concentrations of caffeine and benzoic acid in your sample. You
will need to solve two equations simultaneously since you have two unknowns - both species will absorb
to some extent at both wavelengths, so the absorbance you measure will be a sum at that wavelength.
You can solve these algebraically or using matrices as discussed in Chapter 19 of your textbook

Turn in your four graphs by attaching them to your lab notebook. Finally, report the concentration of
caffeine and the concentration of benzoic acid (in mg/L) in your unknown sample along with your
student number. (No uncertainty analysis is required for this lab).

Exp. 12, Fall 2009                                                                                       51
Experiment 13 (Bonus Lab): Determination of Hop Bittering Units in Beer

        Hazardous Waste: All organic solvents used in the extraction must be collected as waste.

[SAFETY NOTE: It is absolutely forbidden to taste anything in this laboratory. Even if the particular
item is not toxic, it is very possible that it has been contaminated with something else.]

Beer is a beverage that was likely first brewed as a method to preserve grains from spoilage, just as wine
was a method of preserving fruit. In the case of wine, the relatively high alcohol content (9-12%) is
typically sufficient to preserve the wine. Most beer, however, has a lower alcohol content (usually 3-6%)
which did not always adequately preserve the malt beverage. In the past, brewers relied on the addition
of various flowers and herbs to the beer, which acted as an antiseptic. This helped to preserve the
beverage by reducing the growth of bacteria and mold.

During the Middle Ages European monks perfected the brewing process and made nearly universal the
use of hops in beer as a preservative. Today virtually all beers are produced with hops, although most
commercial beers are now pasteurized and the hops are used primarily for the flavor they impart to the

Hops are the cone-like flower of humulus lupulus, a member of the cannabis family. Hops are used in
the brewing process and are responsible to two types of flavors in beer. Hops added during the early
stages of the boiling of the wort (unfermented beer) to release lupulin and various essential oils. A group
of compounds known as humulones are refered to as alpha acids. Boiling causes the alpha acids to
isomerize which makes them more soluble and allows them to impart bitterness to the beer. Hop flowers
can also be added at the end of the boiling process in order to impart a desired aroma and taste to the
finished beer.
                                                                                O          O
                     OH       O

                                  R                                                             R

          HO                  O
                                             boil                                     OH


Humulone: R = CH2CH(CH3)2                                    cis and trans isohumulones
Cohumulone: R = CH(CH3)2
Adhumulone: R = CH(CH3)CH2CH3

Because the amount of isohumulones is important to the taste and stability of the finished beer,
commercial brewers are interested in quantifying the amount of alpha acids in the beer. In this
experiment, we will develop a method for the determination of alpha acids in finished beer.

The International Bitterness Unit (IBU) is a tool used to describe the amount of alpha acid in beer, a
measure of how bitter the beer will taste. Specifically, the IBU describes the number of mg of alpha acid
per Liter of beer (mg/L). For example, it has been reported that Budweiser has an alpha acid content of
~10 IBUs, while Red Hook reports their ESB Amber Ale has a bitterness of 24.3 IBUs. Some styles of
beer, such as India pale ale have even higher IBU ratings. In this experiment, you will determine the
IBU value for 2 an unknown beer sample by using a method based on UV absorption. You will be
provided with 3 beers of known IBU. These beers will serve as standards to construct a calibration curve
for the method. Once you construct the calibration curve, you can determine the IBU for your unknown
beer. The method is based on the absorption of ultraviolet light by the isohumulones at a wavelength of
275 nm. A two phase liquid-liquid extraction is used to extract the alpha acids into an organic solvent,
which is then measured in the spectrophotometer. The absorption at 275 nm is proportional to the
concentration of the alpha acids in the beer sample.

Take your 10 mL pipette and rinse it with octyl alcohol. Then take a 10 mL aliquot of beer and place it
into a 50 mL volumetric flask. Add 1.0 mL of 3 M hydrochloric acid to the flask, and then using a 10
mL pipette add 20 mLs of isooctane. Place a stopper on the volumetric and shake for 8-10 minutes. Take
care to allow release of any gas evolved from the solution.

Allow the solution to separate into two phases, then, using a Pasteur pipette, separate the lighter
isooctane phase into a clean dry beaker being careful not to accidentally transfer any of the aqueous
phase into the beaker. The alpha acid will be in the isooctane phase. Next, take your samples to the UV-
Vis spectrophotometer and set the wavelength to 275 nm. (your TA will explain how to use this
instrument) Fill a cuvette with an isooctane/octyl alcohol solution. (Make sure there are no
fingerprints on the cuvette and that the cuvette is placed in the UV-Vis the same way each time.
The isooctane/octyl alcohol solution is to be used as the reference. Record the absorbance of the sample.

The procedure should be repeated for the three beer samples that serve as standards and for the one
unknown sample. A calibration plot should be constructed from the beers with known IBU values
(Absorbance versus IBU). Using least squares analysis, obtain the equation for the line that best fits the
data. Once the equation is known, calculate the IBU value for the unknown samples. Include an error
estimate based on the uncertainty in your least-squares fit. Your textbook has information about how to
do this.

Exp. 13, Fall 2009                                                                                       53

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