COLLEGE CHEMISTRY I
PHS 1025
Laboratory Manual
INDEEPENDENCE COMMUNITY COLLEGE
Department of Chemistry
Fall 2014
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Table of Contents
Introduction ....................................................................................................................................2
Laboratory Philosophy .................................................................................................................2
Laboratory ...................................................................................................................................4
Laboratory Notebook ...................................................................................................................6
Comprehensive Laboratory Report ..............................................................................................9
Laboratory Grading Policies ......................................................................................................10
Laboratory Techniques ...............................................................................................................22
Experiment 1 & 2: Statistics & Calibrations ............................................................................22
Experiment 3: Gravimetric Anaylsis .........................................................................................33
Synthesis of Alum ......................................................................................................................33
Analysis of Alum .......................................................................................................................39
Experiment 4: Titration-Acid & Base ........................................................................................46
Standardization of NaOH ...........................................................................................................46
Determination of Fruit Juice ..........................................................................................................
Determination of KHP ...............................................................................................................57
Experiment 5: Titration-Oxidation & Reduction .....................................................................58
Standardization of Thiosulfate ...................................................................................................58
Determination of Vitamin C Content ........................................................................................ 58
Experiment 6: Beer’s Law...........................................................................................................63
Determining Iron Content ..........................................................................................................68
Determining Aspirin Content .....................................................................................................74
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Introduction
Lab Philosophy
The fundamental assumption made of you as you begin College Chemistry I PHS 1025 is that
you are an adult person with a scientific curiosity ready to learn first year chemistry. This may
not be a valid assumption in all cases, but if this is not your approach to this course, then try to
pretend that it is. There is much work to be done. This may be one of the most demanding
courses you will take at Independence Community College. However, it is also one of the most
useful courses for anyone planning to do science at any level.
What you can learn from College Chemistry I & II will be extraordinary. College Chemistry I,
and College Chemistry II that follows are ultimately practical to all aspects of scientific careers.
The lab skills which you will learn and refine are the most "marketable" part of your degree. The
informed skepticism that you should learn to apply to the data you obtain is an essential approach
for all scientists. At the completion of College Chemistry I and College Chemistry II you will
have hands-on experience to all but a small fraction of the instruments and methods used in the
various aspects of chemistry. The laboratory portion of College Chemistry I will introduce the
most commonly used methods and instruments and prepare you for work, summer jobs, and
research in chemical laboratories.
Laboratory (READ THIS PART TWICE)
You will find that detailed instructions for experiments will become limited to specific
techniques as we get further along in the class. Many details are left to you to figure out once we
have conducted a few experiments. All of these tasks are within your ability with a little thought.
This is done on purpose to force you to spend time thinking about the labs and to learn to plan
experiments. The goal is to give you a real life experience and force you to do most of this
preparation before you arrive at the lab. The Quizzes are provided to assess your preparation.
You will not be allowed into the lab without completing the assigned Quizzes. If you take
shortcuts and do not adequately prepare for lab, I can almost guarantee you will not have
sufficient time to complete the experiments and no additional time will be allowed. If you come
to lab well prepared, you can expect to leave by the end of the class period. Class will end
promptly at 2:50 since I have a class coming in at 3:00 pm. If you come to class prepared to do
your experiment, you will find you have more than enough time to make the usual mistakes.
A proper approach to the lab is to read the experiment and references before you come to lab.
You should then reread the experiment and establish a plan of approach to the experiment in
detail. Such a plan will detail each step to be taken in the order you feel is the most efficient
Frequently, the most efficient or even necessary order of the experiment will be different than the
order of topics in this manual. Your plan of approach must include the weights and volumes of
reagents and solutions to be used at each step in the procedure. You know what glassware you
have and we will attempt to inform you of the specific reagents available for the experiments.
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The plan of approach should also include preparation of data tables for the data you will be
taking. This will lead to better organization of your notebook and faster work in the lab.
As you progress through the semester, you will find the labs are more sophisticated (and the data
you get will in general be worse, so don't panic) and require more preparation. The reporting
requirements also become more demanding. You will have several opportunities this term to
design parts of your own experiments, each time offering more options and greater
sophistication.
I hope that, by being forewarned of the expectations of the lab you will not panic when things go
wrong in the lab. Changes in procedure are much more easily managed when you understand
(translated as not just memorized or transcribed from handout to notebook) what is really going
on. Such changes will occur due to a variety of reason ranging from the perversity of the
instructor to mistakes in the preparation of solutions. Stay loose!! As you will learn, even
procedures printed in the chemical literature do not always work for everyone. This is where you
must learn to be skeptical and think about what you are doing. If it doesn't work, figure out why
and fix it. Double check your own calculations and those presented in the manual.
Do not hesitate to ask for help at any point in the process. If you don't understand what you are
doing before lab, make a point of asking questions in class before the lab or find the instructor
before lab. This will insure that you are ready for the lab and will not have to wait until the
instructor has gotten things going for the prepared students. When you arrive unprepared in the
lab, you go to the bottom of the priority list for help. When you ask for help before lab, you are
my top priority.
Finally, a few comments on cooperation between scientists (and students) are in order. No person
in science can work in a vacuum (figuratively speaking). The reason for scientific meetings and
the journal literature is to provide a means for the sharing of ideas and results. Out of this
interaction, new ideas are born. The same can be said for your work in this course. The
interactions you have with other students can be very valuable for learning the material.
However, the same rules apply to these interactions that apply to those of any group of scientists.
ALL SCIENTIFIC INTERACTIONS MUST BE PROPERLY REFERENCED. If you work with
someone as a lab partner on an experiment, you are both expected to keep a lab notebook
detailing what you did. The notebooks will not be identical as you will certainly have different
styles of presenting what you have done. The Student Handbook has some examples of proper
practice under the section on plagiarism. READ IT! These rules apply to everything you do as a
scientist or student.
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Laboratory Notebook
A scientific notebook should hold a permanent record of the experimental work. Consequently,
the pages should be securely bound (not a spiral binding) and entries made in ink at the time
work is done. The pages are numbered and the entries are dated. The format of the entries must
be such that the book could be read by any scientist who is familiar with quantitative chemical
work. Readability of the original data is essential. Neatness is convenient but not central.
Neatness is not a grading criterion for the notebook. Too much neatness is often a sign that the
information has not been actually written during the lab.
It is common practice to record data on the right hand page and to use the left hand page for
preliminary readings, notes and calculations. The left hand page is the only place you are
allowed to write scratch notes. Any evidence of recording data on scratch paper, lab manual
pages, paper towels, etc. will result in a laboratory notebook grade of 0 for the experiment.
Errors are a part of scientific work. Data is frequently invalid - experimental conditions may not
be adequately controlled, reagents may be contaminated, and instruments may not be operating
correctly, scales of instruments may be incorrectly read, etc. If results have more than the
anticipated variation, the first reaction may be to discard the whole thing and start over. Later
more significance may be read into the data. All data - including data that is known to be invalid
- remains a part of the laboratory record. To correct an entry, draw a single line through it and
enter the correct value above the original value which remains readable. If the reason for the
change is not obvious, an explanation should be recorded. If a large section of the work is
considered invalid, a single line is drawn diagonally across the page and the reason for
discarding the work stated. All entries remain readable. Under no circumstances is a page
removed from the book. To erase or block out data, or to remove a page from a notebook is
considered a violation of scientific integrity. In research, valid notebooks are the basis of
determining priority of scientific discoveries and the granting of patent rights.
At the end of a determination, the work is summarized in a table which includes the data and
calculated values for all trials. The pages on which the data, the calculations and the summary
occur are cross referenced if the pages are not consecutive. With experience, it is possible to
record essentially all data directly in this summary table. These tables, particularly if they show
intermediate values in the calculations, are helpful in discovering trends and locating errors in
calculations. One of the chief sources of errors in quantitative determinations is the incorrect
treatment of the data. Calculations are as much a part of a determination as any measurement.
One example of each calculation with all units must be included in the notebook. Replicate
calculations should not be included.
The correct use of significant figures requires alertness and judgment. The most common error is
the copying of all of the digits on the computer or calculator printout. The use of computers and
calculators require careful attention to the correct entering of data followed by careful
consideration of the number of figures to be retained in the final answer. One additional figure
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should be retained in all intermediate values of a calculation and the final answer reduced to the
correct significance. It takes hard work to get good data. Don't invalidate the results with sloppy
calculations.
1. Anyone can keep an accurate notebook if they put their mind to it. Here are some
suggestions which may help organize the notebook so it is easier to read.
2. Plan what you are going to do before you start writing.
3. Don't try to put too much on one page. Leave plenty of space.
4. Use one end of your bench space for your notebook and the other for wet work - which
depends on whether you are right or left handed and also the position of the sink.
5. When not in use, keep the notebook closed and off of the desk top. If a disaster occurs,
data may be transferred. Draw a single line across the first page and cross reference the
two pages.
6. Run through calculations on scrap paper first. This is quite permissible since calculations
can always be repeated as long as the data are available. Show only one example
calculation in your notebook to allow later verification of the method used.
7. Data and calculated values should be presented in tables. Even though instrument
recording and computer printouts are attached to the reports, significant values should be
included in the tables.
Tables of results are possibly the most crucial part of the notebook. A well-organized student can
often prepare the table prior to beginning the experiment. This saves laboratory time and
indicates a thoughtful, orderly approach to the experiment. A well thought out table will simplify
recording and subsequent calculation of results.
The importance of establishing a pattern of careful notebook keeping is difficult to
overemphasize. Most industrial laboratories engaged in analytical chemistry are required to
follow Good Laboratory Practice standards (GLP's) established by professional groups or
government agencies. In some laboratories, for example, every weighing, buret reading, etc.,
must be witnessed and countersigned. This would be a little extreme in this course. An almost
universal practice is to have each page witnessed at the end of each day’s work.
Although no grade will be associated with maintaining the laboratory notebook, the experimental
data and information contained in the notebook will be invaluable to your laboratory report. If
you maintain a poor notebook, your report will show this because of the lack of information and
the inaccurate results. Approximately 60% of your laboratory report will come from the
information you keep in you laboratory notebook.
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Comprehensive Laboratory Report
Brief reports of analysis will be submitted for each experiment. For a limited number of
experiments a comprehensive, technical journal style report will also be required. All reports
will be submitted electronically by email attachment.
The nature of the report required will be specified for each experiment at the end of the section
for the experiment. Comprehensive reports will be due one week after the completion of the
specific experiment. Those dates are specified in the syllabus. You will never prepare a "perfect"
report and the attempt to do so can result in procrastination and the absence of all productive
activity.
There will be no late laboratory reports accepted but there is always a possibility that due to
illness or other factors you will need an extension on your lab report. Extensions will always be
granted for any reasonable cause. Extensions will only be granted through an e-mail request to
the professor. If you informally ask for and are granted an extension, this is only my way of
saying, yes I will grant an extension if you ask through the formal procedure. When you request
an extension, you must state your general reason for the extension and give a specific new due
date when you will hand in the report. I can accept the proposed due date or change it. In any
case, you will receive an e-mail reply stating the new due date for the lab. You must copy that
message and include it at the end of your lab report when you submit it for evaluation. If
there is no written extension document submitted, the full late penalty will be assessed. No
open ended extensions will be granted.
Although you are not English majors, it is still very important to be able to communicate your
science to others. Writing poor enough to be unclear or confusing will affect your grade, as will
lack of regard for grammar, punctuation, and spelling. Learn to use the spell-check facility that is
built into your word processor.
Here are some general comments to help you with your scientific writing. For more detail, we
recommend the first chapter of the ACS Style Guide.
Scientific writing is not literary writing. You should aim to be brief, precise, and unambiguous.
The reader should clearly understand what you are trying to say. Try to keep your verb tense
consistent and appropriate. You may use either passive or active voice, but try to be consistent.
Avoid using jargon or slang and use full sentences. It is rare that you would need to use first
person; i.e., try not to refer to “I”, “we”, “our”, “us”, nor should you speak about yourself, e.g.,
“the student”.
Title
The Title should describe the experiment in an appropriate, adequate and concise manner.
Abstract
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Is a summary of the goals, objectives and the results of the experiment. It should convey a sense
of the full report in a concise and effective manner. Remember, a reader should be able to read
the abstract and get a full sense of what was accomplished in an experiment.
Introduction
Connects the basic scientific concepts being tested, or utilized in an experiment with the
objectives and purpose of the experiment. Hypotheses and logical reasoning should be given.
Data and Observations
Report all data, including appropriate units, preferably in tabular form. Any data that was
collected digitally (i.e., using computerized data collection) should be printed out graphically as
described in the data analysis section of your experiment and included in your report. There is no
need to print out the large data tables collected by the computer. Any data that was collected on
chart paper should be photocopied onto regular US letter size paper, reduced or enlarged if
necessary, labeled appropriately, and included in your report. Do not attach long rolls of chart
paper to your report. Include any relevant observations (e.g., color changes, unintended spills,
unexpected results, etc.) and note anything you did differently from the procedure given in this
manual.
Data Analysis
See the Useful Analysis Techniques section of this manual for important information, relevant
to your data analysis. Your sample calculations may be hand-written, but must be clear and
legible. Follow the analysis steps in this manual, labeling them by number so we can follow
what you are doing. You must clearly show each step of your analysis, by writing out each
equation used (the formula in words, or symbols with an explanatory key), and followed by the
calculation plugging in the relevant data, including the appropriate units, even if the equation is
explicitly stated in the lab manual. Put a box around, underline, or in some other way highlight
the important answers.
Unless you are explicitly told otherwise in the data analysis instructions, if you studied multiple
samples, carry out all calculations for all replicates; i.e., do not average them together until the
very end. Do not show working for all replicate calculations, simply show one complete sample
calculation for the first replicate. You may then simply summarize the calculation results for the
other replicates in a table or spreadsheet, making sure you label the columns well with
informative names and units. It is not necessary to show standard deviation or line of best fit
calculations, simply provide the answer from your spreadsheet/calculator. In addition to your
calculations, your data analysis section must communicate to the reader what you are doing.
There needs to be enough detail and explanation that a reader could follow your steps and
reasoning and reproduce your results. State any reasoning processes that you have used that are
not described by equations, such as anywhere you used common sense, logic, bizarre thought
processes to decide something (like not dropping a data point, choosing which equation best
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represents your data, etc). Each step of your analysis should be clear and obvious to anyone
reading your report.
Results
This section should report the results of the experiment separately from your calculations in a
clear and concise format. Where multiple data sets are used, report each individual
determination, the average with appropriate significant figures, the standard deviation, the
percent relative standard deviation (%RSD), and the 90% confidence interval for your data.
When using a line-of-best fit to data, report your correlation coefficient. If you Q-test away one
of your three replicates, then you cannot report a standard deviation. That’s ok, just talk about it
in your conclusion.
Conclusions
The most scientifically important part of an experiment is your interpretation of your
observations. For an analytical chemist, the accuracy and precision of your results is also
important. These priorities are reflected in the allocation of some of your lab report grade to
these topics. It may seem difficult for you to imagine what more there is to discuss once you
have determined something like the % content of ascorbic acid in a Vitamin C tablet; however,
here are some guidelines to help you:
Summarize the point of the experiment. (What did you do? Why did you do it?) Discuss your
results. (Can you compare to literature values or values given by a manufacturer? Do they seem
reasonable? Why or why not?) In the cases where values are available for comparison, you
should quantify the difference between your experimental values and the expected values. Rather
than simply saying you were “way off”, a comment such as your value was 15% larger than the
literature value would be more appropriate. List the possible sources of error in your experiment,
being as specific as you can in your error descriptions; e.g., rather than saying “operator error”,
describe exactly what aspect of the experiment you may have done incorrectly. Discuss any
problems encountered and suggest ways around them if the experiment was repeated. Evaluate
possible sources of error (operator, random, systematic, etc.) in the experiment. Discuss the most
significant contributions to the error in this experiment and describe specifically how each of
these significant errors would have affected your results (i.e., increased or decreased your
answers). Give reasonable methods to eliminate or reduce these errors. Consider the experiment
in terms of accuracy, precision, reproducibility, selectivity, and analysis time. Suggest ways to
improve your procedure if you were to perform it again.
Questions
Write numbered, coherent answers to all questions found in the lab manual. Use complete
sentences and be sure your answer demonstrates understanding of the material. Like the rest of
your report, the discussion section must be in your own words. Plagiarism will not be tolerated.
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Laboratory Grading Policies
The laboratory section of the course contains two (2) laboratory activities worth 25 points each,
four (4) comprehensive laboratory reports turned in each worth 100 points, and a laboratory
practical worth 50 points. The breakdown for grading the laboratory activity will be as follows:
Laboratory Activity:
Quiz
= 5 points
Experimental Procedure
= 5 points
Results & Discussion, Conclusion, Figures & Tables
= 10 points
Questions
= 5 points
Total
= 25 points
Comprehensive Laboratory Report:
Notebook
= 10 points
Quiz
= 10 points
Title, Abstract, Experimental Procedure
= 10 points
Results & Discussion, Figures & Tables
= 30 points
Conclusion
= 20 points
Questions
= 20 points
Total
= 100 points
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Laboratory Techniques
There are essentially two criteria for judging a technique –
 Does it work?
 Is there a less laborious technique which works equally well?
In a few cases, the scientific world frowns on perfectly adequate procedures. An analogous
situation is eating mashed potatoes with a knife. This section will provide some guidance to help
you develop habits of work which are both simple and effective.
Reagents
In so far as practical, reagents will be supplied in the manufacturer's bottles. Note the analysis
reported on the label. Record the manufacturer and grade of the reagent used in your lab
notebook. This information is always required in writing a technical paper and you should get in
the habit of recording it.
The purity of the reagents used is a determining factor in the results obtainable in analytical
work. Consequently, every effort must be made to keep stock bottles free from contamination.
Under no circumstances should material be returned to a stock bottle in a general laboratory. A
large stainless steel scoopula with a handle or a porcelain spatula may be used to break up caked
solids. Use some judgment as you would not want to introduce the stainless steel scoopula into a
reagent used for trace metal analysis. An example from recent history was a geological analysis
which was ruined because of contamination from a platinum wedding ring worn by the
technician doing the sample preparation. This error received world-wide publicity due to the
nature and importance of the erroneous findings.
In general, solids are best poured from bottles. Rotating the bottle back and forth helps to control
the rate of flow. Droppers and pipettes should never be dipped into stock bottles. Droppers from
dropper bottles should not come in contact with any surface outside of the dropper bottle itself.
Primary Standards
Substances prepared for use as primary standards are so labeled. These materials are expensive
and must not be used for routine procedures.
Drying at Elevated Temperatures
Drying ovens for this course are thermostated and are usually operated at 120-125oC. As the
name implies they are used for drying at an elevated temperature. The efficiency of the drying
process also depends upon the pressure of water vapor in the immediate atmosphere.
Consequently a very wet object introduced into an oven may actually cause another object in
the same oven to pick up water. Consequently, one should not place wet glassware in an oven
being used to dry analytical samples or standards. Before using the ovens, note the temperature
at which they are operating. Do not change the setting on an oven without consulting the
instructor.
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Space is always at a premium in a drying oven. Distribute objects around the sides and back of
the shelves so that all objects can be reached with tongs. Crucibles and weighing bottles should
be dried in small labeled beakers covered with small ribbed watch glasses. Any object left in an
oven overnight may be confiscated between 8:30 and 9:00 A.M. unless specific permission was
obtained.
Two common problems encountered in "community" ovens are the loss of samples (someone
else takes yours) and knocking over someone else's samples. The former problem is avoided by
clearly labeling the samples. If you place them in a beaker, you can easily place a large paper
sign with your name on it in the beaker with the sample. When manipulating samples in the
oven, be careful not to knock over others.
Microwave ovens are replacing conventional ovens in the laboratory just as they are in the
kitchen. Typically, samples can be dried in about 25% of the time it would require in a
conventional oven. However, there are some significant exceptions. Some samples melt, burn,
or decompose when placed in a microwave oven for an extended time. Check with the
instructor before drying a sample in the microwave oven. It is always advisable to
experiment with a small quantity of the material before entrusting your entire sample to the
microwave.
CAUTION: MATERIALS REMOVED FROM BOTH MICROWAVE OVENS AND
CONVENTIONAL OVENS ARE HOT. USE TONGS OR GLOVES WHEN HANDLING
THESE OBJECTS TO PREVENT BURNS.
Cleaning Glassware
The cleaning process should be as simple as possible. Rinsing with tap water and several small
portions of distilled water may be adequate. For more dirty glassware, scrubbing with detergent
and water should precede the rinsing. From a chemical point of view soap, detergents, etc. are
dirt. Rinse very thoroughly with tap water and then at least three times with small volumes of
distilled water. In a few cases special chemicals may be needed to dissolve solids or oils. The
inside of a container which may come in contact with chemicals is not dried with a towel since
this introduces lint. In many cases it is not necessary to dry glassware. Simply rinse with the
solution to be used unless this would invalidate measurements. Beware of drying glassware with
compressed air. This may introduce oil vapor from the pump. I am not aware of any
instance where this is appropriate in this lab. The best strategy is to keep all your glassware
clean so that you have clean and dry equipment available at the start of each lab. Before you dry
glassware, ask yourself if there is another approach. If you are going to immediately add water to
the container, why does it need to be dry?
Cleaning Volumetric Glassware
Volumetric glassware must be clean so that water drains from the surface without leaving
droplets. It will not do this if there is the least bit of oil on the surface. Once "clean" glassware
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becomes dry it will usually not drain properly when water is again added. Volumetric glassware
is cleaned just before use or cleaned and stored full of distilled water or other solvent. The 50 ml.
burets can be scrubbed with lab soap - 1 teaspoon in a large beaker of warm water - and a buret
brush. In doing this, scrub the buret in sections - about 10 cm. at a time. The 10 ml. buret is
fragile and difficult to clean. It is too small for a brush and is cleaned in the same fashion as the
pipet. The tip is very small and the cleaning solution should not be taken through it since it often
contains undissolved pieces of soap.
Working Surfaces
Use paper towels to wipe up all spilled materials. Repeatedly wash the surface with a wet towel
to remove water soluble materials including acids and bases. The working surface should be kept
clean and dry. Spilled material is then quite evident and contamination can be kept at a
minimum. If working materials are arranged in an organized manner, there are fewer
opportunities for confusion and there is a higher probability that a determination will be carried
through to completion without error.
Quantitative Transfer
Quantitative transfer is the complete transfer of a sample without loss of any kind. The
techniques used are a matter of common sense - do not spill, splash, drool or abandon. Dry solids
are poured or transferred with a spatula. If the surface tension of a liquid is high it should be
transferred by pouring down a stirring rod. This prevents the liquid from running down the
outside of the original container and also prevents splashing as the liquid enters the second
container. Last traces are transferred by washing the original container and transfer equipment
such as spatula, stirring rod and funnel with a miscible liquid. This can be done by a batch
method using repeated small volumes of the wash liquid or by a continuous flow method using a
stream of wash liquid from a wash bottle. A rubber policeman is used to facilitate the process
and minimize the amount of wash solution necessary. Transfer of material from a weighing
bottle to a flask should always be done by pouring and not with a spatula.
Reading Instrument Scales
The advent of digital readouts has reduced the opportunities to read analog and vernier scales.
Therefore, when you encounter the need to read instrument scales or calibrations on burets and
pipets, it is especially important that you pay close attention to this process. Even with digital
readouts, it is possible to make serious errors if the instrument has several modes of display (eg.
Transmittance and Absorbance). Become familiar with the scale. Does it read from left to right
or right to left? Top to bottom or bottom to top? Determine the significance in both units and
magnitude of the largest divisions then determine the significance of the smallest divisions. The
number and magnitude of the small divisions may not be the same in all ranges of the scale.
With most analog scales the orientation of the operator's eye and the instrument determine the
reading. If all readings occur at a fixed position on the instrument, it is only necessary that the
position of the eye be the same for a set of measurements.
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To read a variety of positions on a horizontal scale, the most reproducible orientation is to have
the eye directly above the pointer. For a variety of positions on a vertical scale the most
reproducible orientation is to have the eye at the same level as the pointer. In both cases the line
of vision is at right angles to the surface bearing the scale.
If the surface of a liquid is to be read in place of a pointer, a reproducible position on the surface
must be chosen. This is usually taken to be at the bottom of the meniscus if the liquid wets the
glass and at the top if the liquid does not wet the glass Calibrations which cover at least half of
the circumference of the tube serve as a check on the correct eye level. The angle of reflection of
light makes a significant difference in the appearance of the meniscus. The lighting can be
controlled by holding a card against the back of the tube: For clear liquids a white card
containing a broad dark line serves as the best means of reading the meniscus. The dark line is
raised until it just touches the bottom of the meniscus. The top of the line is then compared with
the graduations on the device to determine the value.
Readings are in general made to 0.1 of the smallest calibration division. For example, the
readings with a scale calibrated to 0.10 cm. are estimated to the nearest 0.01 cm. Since
calibration lines have width, some convention must be established for the use of the line. If the
line width itself is 0.02 cm., the top of the 3.10 line could be read as 3.09, the center as 3.10 and
the bottom as 3.11. A pointer also has width. Choose a point of reference - the center, one edge,
some irregularity on the pointer, etc. To record a reading as 3.1 cm. states that the value is
thought to be closer to 3.1 cm. than it is to 3.0 cm. or 3.2 cm. To record a reading as 3.10 cm.
indicates that the value is thought to be closer to 3.10 than to 3.09 or 3.11 cm. To record 3.1 cm.
implies one of two things. It is impossible to determine the value more carefully or the operator
simply chose not to read the value more carefully. In the first case, 3.1 is the correct reading. In
the second, 3.1 is an approximate reading. To make an approximate reading is stupidity when the
more exact value is needed. To record 3.1 when the value 3.10 has been read is unscientific - or
to be more brutal, just plain sloppy.
Weighing Out Samples and Using Weighing Bottles
Your equipment includes three weighing bottles. These are small glass bottles with ground glass
tops. Weighing bottles are to be used only for drying, storing, and weighing solid standards and
unknowns. Weighing bottles should be numbered in pencil on the ground glass surface.
Samples to be dried are placed in the weighing bottle without the stopper and placed in a beaker
with a watch glass cover and a piece of paper with your name. This entire apparatus is then
placed in the oven for the specified time. Upon removal from the oven, the weighing bottle is
allowed to cool until it can be easily handled and then transferred to the desiccator. The
weighing bottle should not be inserted until the bottle has come to room temperature in the
desiccator.
The preferred method is known as weighing by difference, is to weigh the weighing bottle
containing the dry sample, transfer the sample to the flask in which it is to be used, and reweigh
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the weighing bottle. This last weight becomes the first weight for a second sample if multiple
samples are being prepared. This method assumes that the receiver flask has a large enough
opening that there is minimal risk for spillage. The receiver flask need not be dry, so
considerable time can be saved through not needing to dry glassware. When transferring from a
weighing bottle to a volumetric flas, a powder funnel should be used to facilitate transfer.
Samples can be added to a clean dry container, often a weighing boat, which has been previously
weighed or tared on the balance. Extreme care must be taken when samples are transferred from
this container to insure that no material is lost. Normally, the solvent should be used to wash any
residue from the boat into the container at the end of the transfer. Caution must be exercised with
the common plastic boats as they can accumulate static electricity which either attracts or repels
the particles of sample. This method is not recommended for the most precise quantitative
work.
The best way to manipulate the weighing bottle is to use a band of dry paper pulled firmly
around the bottle. Do not use your fingers directly on the weighing bottle as the moisture from
your fingers will affect the weight. If the weighing bottle stands for several hours in the
desiccator before taking the next sample, its weight should be rechecked.
Rules for Use of Analytical Balances
Weighing performed on the analytical balances shall never be done on weighing paper (or
filter paper, paper towels, etc.) If sufficient precision is demanded to require the analytical
balance it also requires the use of procedures which do not involve such high risk of loss during
transfer. The second problem with use of paper is that it leads to dirty, and subsequently,
damaged balances. Acceptable weighing containers are weighing bottles, plastic weighing boats,
and glassware with sides to contain the material.
No reagent shall be added to or subtracted from a container while in the analytical balance.
Remove the container to the bench top, make the addition and return to the balance. This is the
most fundamental rule of the use of the balance and students found violating this rule will be
disciplined.
Heating and Concentrating Solutions
Aqueous solutions may be heated either on a Bunsen burner with wire gauze or on the hot plates.
The burner is often a quicker method for rapid heating while the hot plate will provide a constant
level of heat for a long time. Solutions other than water or dilute aqueous salts should be
heated in the hoods. A watch glass should always be used during the heating of solutions both
to prevent entry of extraneous material from your neighbors sample, the paint on the ceiling, etc.,
and to prevent loss of sample due to splattering. Ordinary watch glasses restrict the loss of vapor
and are used to maintain the value of the solution. When evaporation is desirable the watch glass
may be supported with three glass hooks or ribbed watch glasses may be used. These allow
escape of vapors. In either case, some sample will collect on the watch glass and must be rinsed
14 | P a g e
back into the container with a small volume of the solvent whenever the watch glass is removed.
Boiling is generally to be avoided with samples due to the high risk of mechanical loss. When
boiling is desired, boiling chips should be used whenever possible. Selection of an appropriate
boiling chip requires knowledge of what may be added without causing contamination of the
material to be boiled. Glass beads or chips, marble chips, silicon carbide, and many other
substances have been used as boiling chips.
Handling Stock Solutions
A uniformly mixed solution may develop a concentration gradient on standing even in a closed
bottle. Evaporation occurs from the surface of the liquid. Condensation takes place on the wall of
the container above the surface and the condensate flows down the wall into the solution again.
Shake or mix stock solutions before use.
Control Charts
Control Charts may be used to evaluate the consistency of the instruments you are using. This is
a procedure used in most routine labs. The charts will be posted in the lab and you are
responsible for entering your data before you leave the lab at the end of each day.
The left side of the chart indicates the sequence of lab days numbered sequentially. There is
also a place for you to enter the actual date. The right side of the chart has a column for you to
enter the names of the members of your group and the numerical value to be recorded. The
graph in the middle shows the trends in analytical results. The first three groups who enter data
are only to fill in the values on the left and right side. After this much data is obtained, the
average and standard deviation of the values will be determined and the scale for the chart will
be established.
You will find the Control Chart very useful in your lab work after the first few weeks. It will
provide an indication of the validity of your analytical results and will help the staff to detect any
problems with standard solutions, LC columns, electrodes, etc.
Volumetric Glassware
Flasks, Burets and Pipets
The specifications used by most manufacturers of volumetric glassware meet and usually exceed
the National Institute of Standards and Technology (NIST) recommendations. Measurements of
volumes larger than 10 ml. can easily be made with an accuracy of 1 or 2 parts per thousand.
Smaller volumes, which are so much a part of present day chemistry, may present special
problems.
Since the volume of a container depends upon the temperature, volumetric equipment is
calibrated for a specified temperature - usually 20oC. The coefficient of expansion of glass is so
small that calibrations for 20oC are valid over the usual range of laboratory temperatures. Even at
30oC the \error is less than 0.3 ppt. Tolerances for various pieces of equipment have been
recommended by NIST. For example, the tolerance for a 50 ml. flask or a 50 ml. transfer pipet is
15 | P a g e
0.05 mL properly used the volume is 50.00 +_ 0.05 mL This is a maximum error of 1 ppt. The
tolerance for a 10 ml. transfer pipet is 0.02 mL, 2 ppt. Note that for the smaller volume the
tolerance is proportionally larger although smaller in absolute magnitude.
Graduated Cylinders
Graduate cylinders should not be used for any quantitative measurement which is part of
an analytical determination in this course. They are useful for preparing stock solutions,
HPLC mobile phases, and other places where the accuracy of the measurement does not have a
direct contribution to the quantitative calculation. Graduated Cylinders are very crude pieces of
volumetric equipment. Non-uniformity of glass at the base of the cylinder makes the
measurement of small volumes of liquid in the bottom of the cylinder particularly unreliable.
Small volumes can be measured with more precision as the difference between two larger
volumes.
Graduated Measuring Pipets
Measuring pipets look like a buret without a stopcock. They are frequently convenient to use but
should in no sense be considered as precision equipment - largely due to the difficulty in
controlling the level of the liquid. In normal use they are no better than graduate cylinders
and should not be used quantitatively unless errors greater than 10% relative are expected
in the results.
Volumetric Transfer Pipets
A transfer pipet has a single calibration line on tubing of small diameter and is capable of high
precision. It is, however, frequently used incorrectly and becomes a serious source of error. For
this reason an operator should run enough calibration checks to attain self confidence in their
technique as well as confidence in the equipment.
All transfer pipets in this lab are marked TD (To Deliver). The proper use of these pipets is as
follows:
Fill the pipet above the calibration line using a bulb. Mouth pipeting is not allowed and will
result in expulsion from the lab.
1.
2.
3.
4.
Tip the pipet to an angle to prevent leakage
Wipe excess liquid from the outer surface of the pipet with a clean towel or wipe.
Drain the pipet until the liquid level reaches the calibration line.
Touch the tip of the pipet to a glass surface to remove the attached drop which is
probably present.
5. Tilt the pipet to carry it to the receiving vessel. This will prevent loss of sample.
6. Drain the contents into the receiving vessel. Do not force the liquid out. Let it take its time.
7. After the pipet has stopped draining for 10-20 seconds (during which time the film of
liquid in the pipet will continue to drain down), touch the tip of the pipet to the edge of
the receiving vessel. Do not blow out the liquid remaining in the pipet after this
16 | P a g e
procedure.
8. You will rarely encounter a blow out pipet in a chemistry lab, but they are still common
in biology labs. The blow out pipet will be labeled TC rather than TD.
Finally, it is never proper to place your mouth on a pipet. Many solvents and chemicals are
toxic or carcinogenic. In a biology lab one must also contend with pathogenic substances.
Always imagine you are pipetting a sample of the AIDS virus or a terrible toxin.
Absolute Calibration of a Pipet
Determine the weight of distilled water of known temperature transferred by a TD pipet. The
degree to which the weights obtained in several trials agree is a check of your technique in the
use of the pipet. Using the average of the weights and the temperature of the water calculate the
absolute volume of the pipet. The volume delivered depends both on the surface tension and the
viscosity of the liquid. Consequently the calibration value is only valid for water or dilute
aqueous solutions.
Use of a Volumetric Flask for Solution Preparation
The interior wall of the flask - particularly the neck - should be checked for uniform drainage
with a small volume of the solvent. If necessary, re-clean the flask. A funnel is usually used to
obtain quantitative transfer of the sample. If the space between the stem of the funnel and the
neck of the flask is small, air may be trapped in the flask when liquid is added. The trapped air
may in turn force liquid back up the neck of the flask. This can be avoided by using a funnel
with a stem that extends into the bulb of the flask or by tipping a shorter stemmed funnel so
that the end of the stem touches the wall of the flask at one point. A small piece of paper
between the funnel and the top of the flask will hold the funnel in this cocked position. Solids
which dissolve readily in the solvent at room temperature may be added through the funnel if
care is taken to wash the solid down with solvent. Solvent is added until the bulb of the flask is
about 9/10 filled and a uniform solution is obtained by swinging the flask in a small circle to
promote swirling of the liquid without bringing it into the neck. Mixing at this time allows
volume changes which accompany dilution to take place before the solution is made up to
volume. Solvent is now added to bring the solution to the calibration mark. The last few drops
may be added with a dropper. The stoppered flask is repeatedly inverted to obtain uniform
mixing - at least l5 inversions, more if the solution is viscous. Volume changes which
accompany this mixing are included in experimental error. Since a uniform solution has been
prepared, solution that is now removed on the stopper in no way changes the concentration of
the solution.
Volumetric flasks should never be placed on a flame or hot plate to dissolve a difficult solute.
This can cause permanent changes in the volume of the flask. If a solute is difficult to dissolve,
carry out the dissolution in a beaker or flask and then quantitatively transfer the solution to the
volumetric flask and bring to the final volume. Ultrasonic bath cleaners can be used to dissolve
solutes in volumetric flasks.
17 | P a g e
The solution is prepared at some temperature - usually room temperature. The volume of the
solution and consequently the concentration of the solution is dependent on the coefficient of
expansion of the solution. (If all volumes are measured in glass equipment, then it is the
difference between the coefficient for the solution and the coefficient for glass that is
significant.) Around room temperature the coefficient of expansion for dilute aqueous solutions
is about 0.024% per oC. A change of 4oC, therefore, corresponds to a change in concentration of
about 1 ppt.
Burets
The calibration lines on a 50 ml. buret are at 0.10 ml. intervals. To obtain maximum precision,
volumes are estimated to 0.01 ml. The calibrations on the 10 ml. micro-burets are at 0.020 ml.
intervals and volumes are estimated to 0.002 ml.
When using a 50 ml buret, the standard deviation of a single buret reading is assumed to be ±
0.02 ml. The two readings required in measuring the volume of reagent transferred introduces an
uncertainty of [(0.02)2 + (0.02)2]1/2 = 0.03 ml. Thus, volumes smaller than 20 ml, even under
ideal conditions, would have an uncertainty greater than 1 ppt. Most operators prefer to work in
the 40 ml. range.
Use of a Buret to Carry Out a Titration
Teflon stopcocks are commonly used in burets. The Teflon stopcock require no lubricant. Tension
on the stopcock is simply increased until the fit is sufficiently secure to prevent leakage but not
interfere with the easy rotation of the stopcock. A Teflon stopcock should be stored free from
tension- loosen the nut. This applies to those on burets and on other glassware such as
separatory funnels. The correct assembly of the Teflon stopcock has the white Teflon ring next
to the glass and the black O-ring next to the Teflon nut which tightens the assembly. Check that
assembly is correct before you attempt to use a Teflon stopcock.
All stopcock burets are designed to be operated with the left hand so that the right hand is free to
agitate the reaction mixture. With the scale of the buret facing the operator, the handle of the
stopcock is on the operator's right. With the base of the left hand to the left of the buret, the
thumb and first two fingers encircle the buret to control the handle of the plug, the last two
fingers against the left of the tip. This braced position of the hand leads to maximum control of
the stopcock. It also makes it possible to keep constant pull on the plug into a secure position in
the seat. This is essential with glass stopcocks to avoid leakage. If initially the position of the left
hand seems awkward, make a conscious effort to develop skill. The instructor will be glad to
give advice and commiserate with your difficulties. (This technique was developed for glass
stopcocks to keep the stopcock from being pulled out. It is not an issue with Teflon stopcocks,
but the design has not changed)
To fill the 50 ml buret rinse three times with 3-4 ml portions of the liquid to be used. Use a buret
funnel so this liquid can be directed to flow over the entire interior surface. Do this in the buret
18 | P a g e
stand to prevent any titrant which may spill from running down to your hand. Allow time for
each portion to drain from the buret before the next is added. Fill the buret, including the tip, and
replace the buret funnel with a buret cap. Never leave the buret funnel in the buret during a
titration since it may add a drop of titrant during the titration. Just exactly what you do with the
buret funnel is a problem. The next time it is used it can be a source of contamination due to
evaporation or reaction of the solution remaining on it or to the solvent remaining in it if it has
been washed. One procedure that is adequate for most reagents is to hang the funnel in the neck
of the appropriate reagent bottle and cover the funnel with a watch glass.
To fill the tip of the buret, fully open the stopcock to allow the liquid to flow rapidly through the
tip. If necessary, use a rubber bulb to exert pressure on the liquid and increase the rate of flow.
Filling the 10 ml buret presents special problems since the diameter of the barrel is too small to
allow liquids to be poured into the top. Consequently liquids are brought into the buret from the
bottom. These burets may also be filled through the tip in exactly the same manner that a pipet is
filled. This is the procedure most frequently used unless a large number of titrations are to be
carried out. The tips of these burets are extremely fine and care should be taken not to pull solid
particles into them.
CAUTION: If titrant is spilled on the outside of the buret, it must be cleaned up and the
waste neutralized and disposed of in a proper manner.
Between laboratory periods, the burets are filled to the top with the solution being used or with
distilled water and left in position on the buret stand in your lab bench.
Logically one should proceed to carry out an absolute calibration of the buret. In practice the
quality of the burets manufactured today is so high and the technique is so straightforward this is
not necessary unless unusual precision is necessary.
To proceed with the titration bring the level of the liquid of the filled buret onto the scale and
read the position after allowing a short time for the film of liquid to drain down the walls. The
reading is more objective if an effort is not made to set the level at zero on the scale. You will be
chastised by the instructor if you have any initial buret readings of 0.00 ml. Trying to set
the buret at exactly 0.00 has been shown to increase errors. Excess liquid on the tip of the
buret is removed by washing with a stream of solvent from the wash bottle and touching the tip
to a glass surface.
Using the right hand to swirl the flask and the left hand to control the stopcock, add liquid at a
rapid and uniform rate. Reaction in the localized region of mixing produces an indicator change.
The addition of the titrant is periodically stopped and the rapidity with which the indicator
returns to its color in the first solution is observed. Using this as a guide, the addition of the
titrant is continued at a gradually decreasing rate. The tip of the buret and the walls of the flask
are washed down with a small volume of solvent from the wash bottle. The process of addition
and rinsing is continued until the end has been located within a drop or within a fraction of a
drop. After a suitable drainage period the buret is read. Fractional drops are obtained by
19 | P a g e
stopping the addition before a full drop has formed. This fractional drop is then washed into the
reaction mixture. The volume of solvent added must be adequate to bring all of the reagents into
the reaction mixture at the end point. Premature washing with large volumes of solvent may
reduce the precision of the work. This is particularly true when the concentration of the titrant is
small. "Small" cannot be specified since it depends upon the properties of the reactants and the
indicator. The time that should be allowed for drainage depends upon the volume of liquid
withdrawn, the rate of withdrawal and the dimensions of the buret. The fine tip has been
designed to place a maximum limit on the second. The micro burets are particularly
troublesome since the surface area is large in comparison to the volume. A check on the reading
after a short time indicates whether adequate drainage time had been allowed.
An overrun end point is an overrun end point. In spite of this the initial addition should be
continuous and reasonably rapid. There is a limit to how long attention can be focused on drop by
drop addition. If the effort invested in a sample is not large, it is good to use one sample to find
the approximate volume of titrant per unit of sample taken. The estimated volume of titrant for the
following samples can then be approached rapidly with confidence. Very close to the end point it
is good to keep a running record of buret readings (use the back side of the notebook page). This
locates the end point between two successive readings and places a limit on the maximum error
that could be involved in locating the end point.
Visual Indicators
The selection of an indicator may be a tricky business which depends upon knowing a great deal
about the chemical reactions involved. Once an indicator has been selected some familiarity
with that indicator should be acquired before making a serious attempt to carry out a titration.
This can be done in a very qualitative manner - even to adding the titrant with a dropper to an
approximate small sample in a beaker. Knowledge of the colors involved leads to confidence
and greatly reduces the time consumed in the first titrations. If a back titration is feasible, the
preliminary investigation provides an opportunity to decide which color change is preferred color A to color B or color B to color A. It is easier to titrate to a definite color change but in
some cases a suitable indicator is not available and it is necessary to match a color standard.
The preliminary investigation sets up this color standard and gives experience in judging how
rapidly the color changes. Work with an indicator until you are confident you can judge its
behavior.
Determination of an Indicator Blank
The indicator gives the "end point," the point at which the titration is ended. Ideally this would
be at the "equivalence point," the point at which chemically equivalent quantities of reagents
have been brought together. In practice the end point and the equivalence point may not
coincide. The determination of an indicator blank gives some information on this point. Even
when the indicator has been correctly chosen, a significant quantity of the titrant may be required
to produce the indicator change or to react with contaminants in the reagents. In some cases it is
20 | P a g e
possible to determine the quantity so used by running a blank - a titration that is equivalent in
every way with the exception that the substance to be determined is not included in the reaction
mixture. In order to determine an indicator blank that has meaning, the operator must be secure
in their knowledge of the behavior of the indicator.
21 | P a g e
Experiment 1 & 2: Statistics & Calibration
Penny Statistics & Calibration Laboratory
Introduction
There are three terms that are used by scientists in relation to their data’s reliability. They are
accuracy, precision and error. Accuracy is how close a measured value is to the true, or accepted,
value, while precision is how carefully a single measurement was made or how reproducible
measurements in a series are. The terms accuracy and precision are not synonymous, but they are
related, as we will see. Error is anything that lessens a measurement’s accuracy or its precision.
To beginning science students the scientific meaning of “error” is very confusing, because it
does not exactly match the common usage. In everyday usage “error” means a mistake, but in
science an “error” is anything that contributes to a measured value being different than the “true”
value. The term “error” in science is synonymous with “mistake” when we speak of gross errors
(also known as illegitimate errors). Gross errors are easy to deal with, once they are found. Some
gross errors are correctable (a mistake in a calculation, for example), while some are not (using
the wrong amount of a reactant in a chemical reaction). When met with uncorrectable gross
errors, it is usually best to discard that result and start again.
The other types of “errors” that are encountered in science might be better referred to as
uncertainties. They are not necessarily mistakes, but they place limits on our ability to be
perfectly quantitative in our measurements because they result from the extension of a
measurement tool to its maximum limits. These uncertainties fall into two groups: systematic
errors (or determinate errors) and random errors (or indeterminate errors).
A systematic error is a non-random bias in the data and its greatest impact is on a measurement’s
accuracy. A systematic error can be recognized from multiple measurements of the same
quantity, if the true value is known. For example, if you made three measurements of copper’s
density and got values of 9.54, 9.55 and 9.56 g/cm3, you would not be able to determine whether
a systematic error was present, unless you knew that the accepted value of copper’s density is
8.96 g/cm3. You might then suspect a systematic error because all of the measured values are
consistently too high (although the closeness of the data to each other implies some level of
confidence). Often in science one needs to assess the accuracy of a measurement without prior
knowledge of the true value. In this case the same experiment is performed with samples where
the quantity to be measured is known. These standards, or known’s, can reveal systematic errors
in a procedure before measurements are made on unknowns, and give the experimenter
confidence that they are getting accurate results.
The last type of uncertainty is random error. As the name suggests, these uncertainties arise from
random events that are not necessarily under the control of the experimentalist. Random errors
can be thought of as background noise. The noise restricts our ability to make an exact
22 | P a g e
measurement by limiting the precision of the measurement. Because indeterminate errors are
random, they may be treated statistically.
Assessing Accuracy
Accuracy can be expressed as a percent error, defined by Eqn. 1, if the true value is known. Note
that the percent error has a sign associated with it (‘+’ if the measured value is larger than the
true value, and ‘-’ if it is less than the true value). Using the copper density data
(1)
from above and Eqn. 1, we can calculate a percent error for each data point of approximately
+6.5%. This suggests the presence of a systematic error because, if there were no systematic
error, we would expect the percent error for each member of the data set to be very small and
that there would be both positive and negative values. When the true value is not known, no
conclusion about accuracy may be made using a percent error. In this case, standards must be run
or other statistical methods based on the precision can be used. However, the latter can be used
only to assess the accuracy of a group of measurements.
In the absence of systematic errors, the average of a set of measurements (Eqn. 2) should
approximate the true value, as the number of measurements, N, becomes very large (i. e., there
are many individual data points, xi). But if a systematic error is present, then making more
(2)
measurements will not make the average approach the true value (as is the case for the copper
data we have been discussing). So to make the most accurate measurements (smallest percent
error), all systematic errors must be eliminated. Note that the percent error for a set of
measurements can be made using the average. The average value of copper’s density, using the
data that we have been discussing, is 9.55 g/cm3, which has +6.6% error.
Assessing Precision
The range is the simplest, and crudest, measure of the precision for a set or measurements. The
range is simply the highest value minus the lowest value, and can be used to get a rough idea of
the spread in the data, but not much more. Sometimes you will see a range reported in the form
± (range/2), which should not be confused with the confidence limits discussed below.
A better measurement of precision for a data set is the standard deviation (σ) which may be
calculated using Eqn. 3 for data sets that have more than about 20 points.
(3)
23 | P a g e
In Eqn. 3 μ is the true mean (what the average becomes when N is large). Since it is rare in
chemistry to have more than three to five replicate experiments, the estimated standard deviation,
S, is used instead (Eqn. 4). In either case, a smaller S or σ indicates higher precision.
(4)
Note the dependence of both S and σ on the number of data points. If the difference terms are all
about the same, then the precision should increase (S and σ decrease) as N increases. So, it is
statistically advantageous to make more measurements, although this must be balanced with
practical considerations. No one wants to do a ten-day experiment 30 times just to get better
statistics!
The standard deviation is related to another estimate of precision known as the confidence limit
or the confidence interval. The confidence interval is a range of values, based on the mean and
the standard deviation of the data set, where there is a known probability of finding the “true”
value. A confidence limit is written as
± Δ at the given confidence level. For example, a
3
volume expressed as 2.16 ± 0.05 cm at the 95% confidence level means that there is at least a
95% probability of finding the “true” value in the range 2.11 cm3 to 2.21 cm3 (in other words,
within ± 0.05 cm3 of the average, 2.16 cm3). It does not mean that only 95% percent of the time
we are confident of the result! To some extent precision is separate from accuracy. However, if
enough precise measurements are made in the absence of systematic error, we have increased
confidence that our average is a good approximation to the true value, even though we do not
know the true value. So, a confidence limit also expresses a level of certainty that the true value
lies within ±Δ of the average, in the absence of systematic error.
To determine a confidence limit, the uncertainty, Δ, must first be calculated from the estimated
standard deviation using Eqn. 5. The value of t in Eqn. 5 may be calculated in Excel using the
TINV function, or may be taken from a table such as Table 1, which gives the value of t for
various degrees of freedom (usually the number of data points minus one, i.e., N - 1) at the 95%
confidence level. Note that as the precision of a set of measurements increases, Δ will decrease at
a set confidence level. Higher confidence levels also reflect higher precision in the data set.
(5)
Degrees
of
Freedom
t
1
2
3
4
5
6
7
8
9
10
15
∞
12.7 4.30 3.18 2.78 2.57 2.45 2.36 2.31 2.26 2.23 2.13 1.96
Table 1. Values of t at the 95% confidence level for various degrees of freedom.
24 | P a g e
Precision and Significant Figures
In lecture and on exams and quizzes when we write a number, we assume that the precision is ±1
in the last number written (for example, the number 31.778 would be assumed to have a
precision of ±0.001). We do this for simplicity. Because when we make this assumption we only
need to concern ourselves with significant figures and we can ignore statistics and the
propagation of error. In real life we are not so lucky and we must worry about significant
figures, statistics and the propagation of error. However, significant figures are always our first
step in analyzing our data.
The uncertainty in a number tells us directly how many significant figures our result has. This is
because the uncertainty tells us in what place the first uncertain digit is (or you could say it is the
first digit where certainty ends). For example, if you had a result 15.678±0.035 kJ/mole at the
95% confidence level, then you could tell from the uncertainty that the first digit that has any
uncertainty in it is the tenths place. We know the 1, the 5 and the 6 (and are confident that we
know them), but the 7 we have some doubt about. We only really know this digit to ±3 at 95%
confidence and the hundredths place is not known with any certainty. How we show this is
discussed below.
Reporting Results
There are three ways in which the statistical information that accompanies a measurement
(average, standard deviation, and confidence limit) can be stated. If, for example, five replicate
measurements of a solid’s density were made, and the average was 1.015 g/cm3 with an
estimated standard deviation of 0.006, then the results of this experiment could be reported in
any of the following ways:
The average density is 1.015 g/cm3 with an estimated standard deviation of 0.006 g/cm3.
The density is 1.015(6) g/cm3.
The density is 1.015 ± 0.006 g/cm3 at the 95% confidence limit.
In this example the density has four significant figures, and the uncertainty is in the last decimal
place. Sometimes the uncertainty and the number of significant figures in the measurement do
not match. This means that each individual measurement was measured more exactly than the
reproducibility within the group. If the standard deviation in the density experiment had instead
been 0.010 g/cm3, then the results might be reported as:
The average density is 1.02 g/cm3 with an estimated standard deviation of 0.01.
The density is 1.02(1) g/cm3.
The density is 1.02 ± 0.01 g/cm3 at the 95% confidence limit.
The results have been rounded off because the number of significant figures does not reflect the
precision of the data set. In other words, the statistical analysis shows us that the first digit where
uncertainty begins is the 1/100ths place, even though each measurement was made to the
25 | P a g e
1/1000ths place. The last significant figure is in the 1/100ths place, so this is where rounding
occurs. Sometimes the average and the uncertainties are quoted to the maximum number of
significant figures (i. e., 1.015(10) g/cm3). In this way the precision of each individual
measurement and the precision of the set of measurements are shown.
Using Statistics to Identify Hidden Gross Error
Another way in which statistics can be used is in the evaluation of suspect data by the Q-test.
The Q-test is used to identify outlying (“bad”) data points in a data set for which there is no
obvious gross error. The Q-test involves applying statistics to examine the overall scatter of the
data. This is accomplished by comparing the gap between the suspect point (outlier) and its
nearest neighbor with the range, as shown in Eqn. 6. The calculated Q is then compared to the
critical Q values, Qc, at given confidence level, like those in Table 2. If the measured Q is greater
than Qc, then that data point can be excluded on the basis of the Q-test.
(6)
N
3
4
5
6
7
8
9
10
Qc
0.94
0.76
0.64
0.56
0.51
0.47
0.44
0.41
Table 2. Critical Q (Qc) values at the 90% confidence limit for a small number of data points, N.
For large data sets (N > 10) a data point that lies more than 2.6 times S (or σ) from the average
may be excluded. Although for medium-sized data sets (between 11 and 15 data points), there is
an alternative treatment that is usually sufficient. In these cases, we can use Qc for N = 10, but in
doing so, a higher criterion is placed on the data for exclusion of a point than is required by
statistics. So, an outlying point that could have been discarded is retained and the precision is
quoted as being less than it actually is. But again, it is better to err on the side of caution in our
data treatment.
In any case, only one data point per data set may be excluded on the basis of the Q-test. More
than one point may be tested, but only one may be discarded. For example, you have measured
the density of copper as 9.43, 8.95, 8.97, 8.96 and 8.93 g/cm3; can any of these points be
excluded?
First, we must remember that the Q-test is only valid at the extremes, not in the middle of the
data set. So before performing a Q-test, it is best to sort the data (as already been done with the
data that we are considering). Now look at the extremes and see whether either of the points look
odd. In this case, the low value (8.93 g/cm3) is not that much different than the values in the
middle of the set, while the high value (9.43 g/cm3) looks to be suspect.
Having decided that the 9.43 g/cm3 value is suspect, we can calculate Q using Eqn. 6, (suspect
value = 9.43, closest value = 8.97, highest value = 9.43 and lowest value = 8.93). This gives Q =
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0.92 for this point. Since this exceeds Qc for five data points (for N = 5, Qc = 0.64 in Table 2),
this point may be excluded on the basis of the Q-test. The Q-test may not be repeated on the
remaining data to exclude more points.
One last important thing about the Q-test is that it cannot be performed on identical data points.
For example, if our data set had been 9.43, 9.43, 8.95, 8.97, 8.96 and 8.93 g/cm3, we would not
have been able to use the Q-test on the 9.43 g/cm3 values.
Propagation of Uncertainty
So, now we have an average and an associated uncertainty at given confidence level for a data
set. What happens if we use this result in a calculation? The simple answer is that the uncertainty
carries through the calculation and affects the uncertainty of the final answer. This carrying over
of uncertainty is called propagation of error, or propagation of uncertainty, and it represents the
minimum uncertainty in the calculated value due entirely to the uncertainty in the original
measurement(s). The following example demonstrates how a propagation of uncertainty analysis
is done.
The dimensions of a regular rectangular wood block are 15.12 cm, 3.14 cm and 1.01 cm, all
measured to the nearest 0.01 cm. What is the volume and the confidence limits on the volume
based on this single measurement? The equation for the uncertainty in the volume is given in
Eqn. 7, where ΔV, Δx, Δy and Δz are the uncertainties in the volume and the x, y and z
dimensions, respectively. Do not be confused by the notation! The Δ represents the uncertainty,
not a change, in these parameters. Since each measurement was made to the nearest 0.01 cm, Δx
= Δy = Δz = ±0.01 cm. First we calculate the volume, being careful with our significant figures
(note the extra “insignificant” figures from the calculator output, shown as subscripts, carried
along in the calculation for rounding purposes).
(7)
Substituting the known values of V, x, y, z, Δx, Δy and Δz into Eqn. 7 gives
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So, the volume would be reported as 48.0 ± 0.5 cm3 for the single measurement, and this
represents a minimum uncertainty in the volume based on the uncertainties in the block’s
dimensions. Note that the propagated uncertainty usually has only one significant figure.
To see how the propagated uncertainty differs from an uncertainty for a population (data set),
imagine that we did this measurement three times and got volumes of 48.1, 47.8 and 48.3 cm3.
Each individual measurement has an uncertainty of ±0.5 cm3, from the propagation of
uncertainty analysis, but the uncertainty for the set of measurements is ±0.7 cm3. This was
calculated with S = 0.3 cm3 (determined using Eqn. 4) and the value of t taken from Table 1 (for
N - 1 = 2) by substitution into Eqn. 5. Thus, the volume would be reported as 48.1 ± 0.7 cm3 at
the 95% confidence limit. Notice that the uncertainty in the population is not the same as the
uncertainty in each individual measurement. They are not required to be the same, nor are they
often the same. In this example, the propagated uncertainty is less than that for a series of
volume measurements, indicating another source of uncertainty besides that arising from the
uncertainty in the block’s dimensions. This is often the case, and in your conclusions to an
exercise or experiment you should try to identify its source and discuss its impact on your result.
Regression Analysis
Once we have data from an experiment, the challenge is to determine the mathematical
expression that relates one measured quantity to another. The problems that confront us when we
attempt to mathematically describe our data are 1) how to establish the mathematical formula
that connects the measured quantities and 2) how to determine the other parameters in the
equation. The process by which a mathematical formula is extracted from a data set is called
fitting, or regressing, the data.
A linear relationship is the simplest, and most useful, mathematical formula relating two
measured quantities, x (the independent variable) and y (the dependent variable). This means that
the equation takes the form y = m•x + b, where m is the slope of the line and b is the intercept. It
is possible to relate two quantities with other equations, but unless there is a good theoretical
basis for using another function, a line is always your best initial choice. For a linear relationship
the values of m and b must be found from the data (x and y values), which is done through a
linear least squares regression (or fit). The mathematics behind the fitting algorithm is not
relevant at this time, but it is important to know that the least-squares procedure assumes that the
uncertainty in the x values is less than the uncertainty in the y values. This means that, if we
want to get a meaningful slope and intercept from our fit, we must make the measured quantity
with the smallest uncertainty be the independent variable.
The analyses to be run include calculating the mean, the standard deviation, and
confidence intervals; running t tests and least-squares analyses; and identifying
discrepant data. A computer spreadsheet (Excel) will be used for all the mathematical
analyses.
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One of the most useful aspects of spreadsheets, and thus one of the more important things to
learn, is the ability to reference cells in equations. Every cell is identified by its column letter
and row number (e.g., A1, B10, C55). When you write an equation in a cell, you may simple
enter the letter and number of the cell to be referenced, or you may used your mouse to select
the cell. Often, you will need to reference a range of cells, which is done by giving the first and
last cell, separated by a colon (e.g., A1:A100 for the range of data in the cells A1 to A100), or
you may use your mouse to select the cells. If you cut and paste a formula, the pasted formula
will reference the cells relative to the cells referenced in the original formula (e.g., if the
formula in B1 is "A1*2" and this is copied to B2, the formula in B2 will be "A2*2"). Often this
is exactly what you want, but sometimes you want a certain reference to always refer to the
same cell. In order to do this, use the "$" (e.g., if the formula in B1 is "$A$1*2" and this is
copied to B2, the formula in B2 will be "$A$1*2").
The data set to be collected consists of the masses of a large number of pennies. U.S. pennies
minted after 1982 have a Zn core with a Cu over-layer. Prior to 1982, pennies were made of
brass, with a uniform composition (95 wt% Cu / 5 wt% Zn). In 1982, both the heavier brass
coins and the lighter zinc coins were made. We will, therefore, only consider pennies made
after 1982. In this experiment, your class will weigh many coins and pool the data in order to
determine whether the averages mass of pennies each year has changed.
Procedure:
Part 1: Statistics
1. Create a table with the headings of the columns being the years of the pennies you are
measuring the mass.
2. Use an analytical balance to find the masses of about 50 pennies. Record these in your
notebook according to their year of minting.
3. Using Excel, prepare a new spreadsheet. Enter the years of the pennies you measured.
4. Input the masses of your pennies into the document. Save this spreadsheet! Record the
name of the spreadsheet in your notebook.
5. Cut and paste the masses from your spreadsheet into the Google Docs spreadsheet.
6. Once the class has completed entering their data, download the spreadsheet to your
computer and open it with Excel. Save the class data spreadsheet with a different name
(do no overwrite your original data!) Record the name of the file in your notebook.
Part 2: Calibration
1. Determine the density of water based on the temperature of the DI water in the containers
at the end of the aisles and the data in the table at the end of this write-up.
2. Calibrate the volume of a 25 mL glass graduated cylinder using the mass of water (do not
use tap water. Use the distilled water) needed to fill it exactly to the 25 mL mark.
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Weigh the empty graduated cylinder. Then carefully fill it so the meniscus is exactly at
the 25 mL mark using a Pasteur pipet so you can be careful not to get any drops of water
above the 25 mL mark. Weigh the full cylinder.
3. Determine the density of a set of 5, 10, 15, 20 pennies that pre-date 1983 by using the
method of water displacement. Weigh the 15 pre-1982 pennies (combined mass) directly
on the balance. Then record the mass of the dry graduated cylinder plus the 15 pennies.
Carefully remove some water from you filled graduate cylinder using the Pasteur pipet.
Carefully slip each of the 15 pennies into the graduate cylinder to avoid splashing. Once
again, fill it so the meniscus is exactly at the 25 mL mark using a Pasteur pipet, being
careful not to get any drops of water above the 25 mL mark. Weigh the cylinder with the
15 pennies plus the water. Record the mass of the cylinder plus the 15 pennies plus the
water. Calculate the average mass and volume of the pennies. Use the mass of water
needed to fill the 25 mL graduated cylinder (step 2) to calculate the volume of the pennies.
Add this data to spreadsheet 2 on the laptop.
4. Repeat step 3 using the 5, 10, 15, 20 pennies that are dated post-1983.
5. Each pair of student will enter their data onto a spreadsheet so you can use the class data
for your analysis. This data will be used to complete a lab report for this experiment.
Analysis
A. Setting up the Spreadsheet
1. Be sure that the first two rows and first column are blank. If you have already written
in the first column, click on the heading for Column A to highlight the entire column,
right click on any highlighted cell, and hit “Insert Column.” Insert rows in a similar
manner.
2. Write your name, the date, and a title for the spreadsheet in cells in the first row.
3. Put the following row headings in the first column: Mean, StdDev, Mean+4s, Mean-4s.
Sort each individual column of data from lightest to heaviest. Highlight the data in just
one column, go to the DATA menu, select SORT, and follow the directions that come up.
Be sure the year is NOT sorted and that Excel does not select and sort the neighboring
data.
4. Set the number format for all the cells that will have data and calculations to have 4
decimal places. Highlight all of these cells, right click, select “Format Cells,” be sure
“number” is selected, and increase the number of decimals to 4.
B. Discrepant Data
1. At the bottom of each column in the appropriately labeled rows, compute the mean and
standard deviation.
2. The Grubb's Test for an Outlier is an excellent way to determine whether any of your data
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are discrepant and should be thrown out. It should be clear that only the maximum and/or
the minimum value in a list would possibly be discrepant. In a cell just to the right of the
year 2000 data, find the maximum and minimum masses of the pennies made in 2000 and
then calculate G for these values.
3. Although the Grubb's test is a very rigorous test for outliers, it can be time consuming,
because the mean and standard deviation must be recalculated after each discrepant
datum is thrown out. Notice that the Grubb's test essentially looks at the number of
standard deviations a value is from the mean, and then sets a minimum number of
standard deviations that defines a discrepant datum for a given number of observations.
To quickly analyze all our sets of data for outliers, we will perform a crude Grubb's test,
erring far on the side of caution by choosing our Critical Value to be 2.5 standard
deviations. Calculate the range of acceptable data (those within 4 standard deviations)
under the data for each year using the formulas:
 average + (2.5 × standard deviation)
 average - (2.5 × standard deviation)
Analyze your data for grossly discrepant masses (those lying > 2.5 standard deviations
from the mean) in any one year. (For example, if one column has an average mass of
3.000 g and a standard deviation of 0.030 g, the 2.5-standard-deviation limit is ± (2.5 ×
0.030) = ±0.075 g. A mass that is  2.925 or > 3.0750 g should be discarded.) Copy and
paste the data from each column that is within the acceptable limits to a series of cells
somewhat below the original data.
4. Calculate the mean and standard deviation of the new set of data just below it. Be sure to
create new row labels for these calculations.
C. Confidence Intervals
1. Find the year with the highest and the year with the lowest average masses.
2. Label the next row “n” (the number of data points). Calculate “n” for the data resulting in
the highest and lowest average mass.
3. Label the next row “t”. Use the TINV (0.05, n-1) function to find the value of Student's t
at the 95% confidence level for the data resulting in the highest and lowest average mass.
Remember the degrees of freedom = n – 1. Record these data in your notebook.
4. Calculate the 95% confidence (μ95%) interval for the highest average mass by hand. Show
this work in your notebook.
5. Have Excel calculate the 95% confidence intervals (μ 95%) for the highest and lowest
average masses in the cells below the calculation for t. Double check that it found the
same interval for the highest mass that you did! Record these intervals in your notebook.
6. Run the t Test for Comparison of Means, note the values of t stat and P in your notebook.
D. Least-Squares Analysis: Do Pennies Have the Same Mass Each Year?
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1. Copy all of the data into a single column several to the right of the data. As you copy the
data, put the year that corresponds to each mass in the column to the left.
2. At the top and to the right of the data, calculate the number of pennies, average and standard
deviation for all the penny masses. Calculate the slope, intercept, standard error of linear
regression, and R-sqaured for the best-fit line, using the years as the x-data and masses as ydata. Record these values in your notebook.
3. Use the worksheet 8-StatisticalFunctions.xls to find the standard deviation for the slope.
Record this number in your notebook.
4. Calculate the 95% Confidence interval for the slope of the best-fit line. Record this in your
notebook.
5. Insert an X-Y scatter plot. The x-data are years, and the y-data are the masses of the
pennies. Give the graph a title, and label the axes. You do not need to show the legend.
6. While still in the mode to edit the graph, go to "Insert -> Trendline" and insert a linear
trendline on the graph.
Conclusion
1. Compare the calculated values of G (step B2) with the Critical Value of G for the given
number of observations. By the Grubb's test, are these values outliers?
2. Compare the result of the t Test for Comparison of Means with the value of Student's t
you calculated for the largest and smallest average masses of pennies (C4 and C6). Is the
difference in these two averages statistically significant (i.e. is the result of the t test
larger than the larger of the two calculated t's)?
3. The P you calculated in C6 is a statistical function that tells us the chance of the
differences between the means of two groups being due to random chance. Say you get a
P value of 0.10 (or 10%). This means that there is a 10% chance that the differences
between your two groups are due to random chance alone. Another way to say this is that
there is a 90% chance that the differences between these two groups is significant.
Normally will say that a P value of .05 or less is significant. What does the P value you
calculated tell you about the significance of the difference between the largest and
smalled average masses?
4. Does the trendline you calculated (D2) and drew on the graph (D6) indicate whether the
masses of pennies has increased, decreased, or remained relatively the same over time?
5. If the masses of pennies had remained relatively constant over time, or varied randomly,
the slope of the trendline would be 0. Is "0" in the 95% confidence interval you calculated
for the slope of the trendline (D4)? What does this tell you about the general trend in
penny masses over time?
6. Print out two copies of the portion of the spreadsheet containing the acceptable data and
the calculations done immediately under these. Do NOT print out the list of 300+ masses
in one column. Format your print area so that all values are printed on a single sheet of
paper. Attach one copy in your notebook and hand the other in with your report.
7. Print out two copies of the graph you made. Attach one copy in your notebook and hand
the other in with your report.
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Experiment 3: Gravimetric Analysis
The Synthesis of Alum
The term alum is a general family name for a crystalline substance composed of cations with 1+
and 3+ charges. In this experiment, you will synthesize a type of alum called potassium
aluminum sulfate dodecahydrate, KAl(SO4)2•12H2O. You will synthesize this compound by
placing the appropriate ions in one container in aqueous solution and then evaporate the water to
form the alum crystals.
2 Al(s) + 2 KOH(aq) + 6 H2O(l)  2 KAl(OH)4(aq) + 3 H2(g)
This particular compound has been chosen because it is relatively simple to prepare a pure
sample. The process of synthesizing this compound is interesting in that it involves both
chemical and physical reactions. Chemically, aluminum is oxidized from aluminum foil to
prepare the Al3+ ions. Physically, as the solution that contains the mixture of ions evaporates,
crystals will form which contain six waters of hydration bonded to the aluminum ion and six
waters bonded to the potassium ion.
Aluminum is considered a reactive metal, but because its surface is usually protected by a thin
film of aluminum oxide, it reacts slowly with acids. It does, however, dissolve quickly in basic
solutions. Excess hydroxide ion converts the aluminum to the tetrahydroxoaluminate (Al(OH)3)
precipitates. Continued addition of acid causes the hydroxide ions to be completely neutralized,
and the aluminum exists in solution as the hydrated ion [Al(H2O)6]3+. Aluminum hydroxide is
considered to be an amphoteric hydroxide because it dissolves in both acids and bases.
OBJECTIVES
In this experiment, you will

Synthesize a sample of potassium aluminum sulfate dodecahydrate (alum).

Observe and record the process of synthesizing a compound.

Calculate the percent yield of your synthesis.
PROCEDURE
1. Obtain and wear goggles.
2. Obtain a piece of aluminum foil and measure its mass. For best results, you should have
about 1.00 g of aluminum. Tear the foil into small pieces and place the pieces in a 250 mL
beaker.
3. Set up a Büchner funnel and filter flask so that you are ready to filter the reaction mixture
that will be produced in Step 4.
4. Conduct the first part of the synthesis. CAUTION: Potassium hydroxide solution is caustic.
Avoid spilling it on your skin or clothing.
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a. Use a graduated cylinder to measure out 25 mL of 3 M KOH solution.
b. Slowly add the KOH solution to the beaker of aluminum pieces. Notice that the reaction is
exothermic. Allow the reaction to proceed until all of the foil is dissolved.
c. Carefully pour the reaction mixture through your Büchner funnel and filter flask setup,
and rinse the filter paper with a small amount of distilled water. Note: The reaction
mixture contains three ions: K+, [Al(OH)4–], and excess OH–.
d. Rinse the beaker with distilled water, and pour the filtered liquid back into the beaker.
5. Allow the solution to cool to near room temperature. If you are pressed for time, you may
cover the beaker with plastic wrap or Parafilm, and store the liquid overnight.
6. Clean the Büchner funnel and filter flask, and prepare it for more filtering that you may need
to do in Step 7 or Step 10.
7. Complete the synthesis.
a. Use a graduated cylinder to measure out 35 mL of 3 M H2SO4 solution. CAUTION: The
reaction mixture must be cooled to room temperature before proceeding. Handle the
H2SO4 solution with care. It can cause painful burns if it comes in contact with the skin.
b. After the reaction mixture has cooled, slowly add the sulfuric acid solution to the beaker
of liquid. Stir the mixture constantly. The reaction is strongly exothermic, so be careful as
you stir the mixture. Note that aluminum hydroxide will precipitate initially. It will
dissolve as more sulfuric acid is added.
c. If there is some solid remaining in the beaker after the 35 mL of sulfuric acid has been
added, pour the mixture through the Büchner funnel and filter flask to separate the
undissolved solid from the mixture.
8. Gently boil your mixture until you have about 50 mL of liquid in the beaker.
9. Cool the beaker of solution. Choose one of the two methods listed below.
a. Allow the solution to cool overnight. In most cases, this gradual cooling forms a good
crop of alum crystals.
b. Prepare an ice bath for the 250 mL beaker. Place your beaker of solution, uncovered, in
the ice bath. Do not move the ice bath or the beaker. After about fifteen minutes, crystals
of alum will appear in the beaker. If there are no crystals after fifteen minutes, scrape the
bottom of the beaker with a glass stirring rod to create a rough spot for crystal growth.
You may also heat the solution to evaporate more water and cool the solution again.
10. Collect your alum crystals by pouring them onto the Büchner funnel and filter-flask setup.
Use vacuum filtration to wash the crystals on the filter paper with 50 mL of an aqueous
ethanol solution (50%). The crystals will not dissolve in this solution.
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11. Remove the filter and crystals from the Büchner funnel and allow the crystals to dry at room
temperature. Measure and record the mass of your sample of alum. Store the crystals for
further analysis.
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Name: ________________________________
Group Members: _______________________
Date: ______________________________
Class & Section: ____________________
PRE-LAB QUESTIONS
1. What are the two cautions you need to take when conducting the synthesis of alum experiment?
2. Fill in the blanks for the physical properties of alum:
a. Molecular Weight: _____________g/mol
b. Melting Point: _________________ oC
3. A student conducting the synthesis of alum experiment started with 0.9156 g of aluminum foil.
a. What is the theoretical yield of alum? (Show your work)
b. The student isolated 8.3181 g of alum. What was the percent yield the student obtained? (Show
your work)
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA TABLE
Mass of Aluminum Foil used: ___________________g
Moles of Aluminum Foil used: __________________mol
Mole Ratio of Aluminum Foil to Alum: ___________ratio
Moles of Alum theoretically obtained: ____________mol
Theoretical Mass of Alum: _____________________g
Actual Mass of Alum obtained: _________________g
DATA ANALYSIS
1. Determine the theoretical yield of the alum. Use the aluminum foil as the limiting reagent
and presume that the foil was pure aluminum.
2. Calculate the percent yield of your alum crystals.
3. Discuss the factors that affected the percent yield.
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4. Write the balanced equations for the following: (a) aluminum and potassium hydroxide,
yielding [Al(OH)4] – and hydrogen gas; (b) hydrogen ions and [Al(OH)4] –, yielding
aluminum hydroxide; (c) aluminum hydroxide and hydrogen ions, yielding [Al(H2O)6]3+; and
(d) the formation of alum from potassium ions, sulfate ions, [Al(H2O)6]3+, and water.
38 | P a g e
The Analysis of Alum
After a compound has been synthesized, tests should be carried out to verify that the compound
formed is indeed the compound desired. There are a number of tests that can be performed to
verify that the compound is the one desired. In the previous experiment, you prepared alum
crystals, KAl(SO4)2•12H2O. In this experiment, you will conduct a series of tests to determine if
your crystals are really alum.
The first test is to find the melting temperature of the compound and compare this value with the
accepted (published) value for alum (92.5°C). The second test determines the water of hydration
present in the alum crystals. The third test is a chemical test to determine the percent sulfate in
your sample of alum.
Objectives
In this experiment, you will

Determine the melting temperature of a sample of alum.

Determine the water of hydration of a sample of alum.
Procedure
Part I: Determine the Melting Temperature
1. Obtain and wear goggles.
2. Connect the Temperature Probe to LabQuest.
3. Take a piece of weighing paper folded in half and using a micro-spatula remove a small
amount of alum from the sample you prepared in the past experiment. Place the alum into
the fold of the weighing paper, fold the weighing paper over, and using the edge of the
micro-spatula pulverize the alum sample. Use the micro-spatula to pile the alum in the
weighing paper. Push the open end of a capillary tube into the pile of the alum powder. Pack
alum into the capillary tube to a depth of about 0.5 cm by tapping the tube lightly on the table
top.
4. Use a rubber band to fasten the capillary tube to the Temperature Probe. The tip of the tube
should be even with the tip of the probe. Use a utility clamp to connect the Temperature
Probe to a ring stand (see Figure 1).
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Figure 1
5. Prepare a water bath to be heated by a hot plate.
6. Monitor the temperature readings on the Main screen. Immerse the capillary tube and
Temperature Probe in the water bath. Warm the alum sample at a gradual rate so that you can
precisely determine the melting temperature. The white powder will become clear when it is
melting. Observe the temperature readings and record the precise melting temperature when
the substance is completely clear.
7. Conduct a second test with a new sample of alum in a new capillary tube.
Part II Determine the Water of Hydration
8. Heat a crucible with cover over a burner flame until it is red hot. Allow the crucible to cool,
and then measure the total mass of crucible and cover. Handle the crucible with tongs or
forceps to avoid getting fingerprints on it.
9. Place about 2 g of your alum crystals in the crucible, and then measure the mass of the
crucible, cover, and alum. Record this measurement in the data table.
10. Set up a ring, ring stand and triangle over a lab burner. Use tongs or forceps to set the
crucible at an angle on the triangle and place the cover loosely on the crucible. Use a lab
burner to very gently heat the crucible of alum until you can see no vapor escaping from the
crucible. It is important that the vapor does not carry any alum with it. After the vapor is
gone, heat the crucible more strongly for five minutes, and then cool the crucible.
11. Measure and record the mass of crucible, cover, and alum after drying. This would be the
mass after the first heating.
12. Reheat the crucible and alum sample for five additional minutes. Cool and measure the mass
of the crucible again. This would be the mass after the second heating. If the two masses are
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the same (or very nearly so), the test is done. If not, repeat the heating and weighing until a
constant mass is obtained.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
PRE-LAB QUESTIONS
1. Write the definition/explanation of each of the following terms:
a. Hydrated Salt
b. Anhydrous Salt
c. Hygroscopic substance
d. Desiccant substance
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2. A student was asked to identify an unknown hydrate by following the following the
procedure in this experiment to determine the percent water of alum. A 2.752 g sample
of the unknown sample is heated and weighted after cooling to constant weight of 1.941
g. The unknown is believed to be one of the following compounds: LiNO3.3H2O,
Ca(NO3)2.4H2O, or Sr(NO3)2.4H2O.
a. Calculate the percent water in the unknown sample.
b. Calculate the percent water in the three known samples.
c. What is the identity of the unknown sample? Explain your reasoning.
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Name: ________________________________
Date: ______________________________
Group Members: _______________________
Class & Section: ____________________
DATA TABLE
Part I Melting Temperature Test Results
Trial 1
Trial 2
Trial 1
Trial 2
Melting Temperature (°C)
Part II Water of Hydration Test Results
Mass of crucible and cover (g)
Mass of crucible, cover, and alum before heating (g)
Mass of Alum before heating (g)
Mass of crucible, cover, and alum after 1st heating (g)
Mass of anhydrous Alum after 1st heating (g)
Mass of crucible, cover, and alum after 2nd heating (g)
Mass of anhydrous Alum after 2nd heating (g)
Mass of crucible, cover, and alum after final heating (g)
Mass of anhydrous Alum after final heating (g)
Average mass of anhydrous Alum (g)
Moles of anhydrous Alum (mol)
Average mass of water lost (g)
Moles of water (mol)
Ratio of moles Alum : moles water
Chemical Formula of Alum
44 | P a g e
DATA ANALYSIS
1. Is your sample alum? Use the results of the three tests to support your answer. Discuss the
accuracy of your tests and possible sources of experimental error.
2. Suggest other tests that could be conducted to verify the composition of your alum.
3. If the melting temperature test was the only test that you conducted, how confident would
you be in the identification of your sample? Explain.
45 | P a g e
Experiment 4: Titration-Acid/Base
Standardizing a Solution
of Sodium Hydroxide
It is often necessary to test a solution of unknown concentration with a solution of a known,
precise concentration. The process of determining the unknown’s concentration is called
standardization.
Solutions of sodium hydroxide are virtually impossible to prepare to a precise molar
concentration because the substance is hygroscopic. In fact, solid NaOH absorbs so much
moisture from the air that a measured sample of the compound is never 100% NaOH. On the
other hand, the acid salt potassium hydrogen phthalate, KHC8H4O4, can be measured out in
precise mass amounts. It reacts with NaOH in a simple 1:1 stoichiometric ratio, thus making it an
ideal substance to use to standardize a solution of NaOH.
Objectives
In this experiment, you will

Prepare an aqueous solution of sodium hydroxide to a target molar concentration.

Determine the concentration of your NaOH solution by titrating it with a solution of
potassium hydrogen phthalate, abbreviated KHP, with an exact molar concentration.
Figure 1
Choosing a method
If you choose Method1, you will conduct the titration in a conventional manner. You will
deliver volumes of NaOH titrant from a buret. You will determine the buret readings manually
and monitor the change in pH by acid-base indicator (phenolphthalein).
46 | P a g e
If you choose Method 2, you will conduct the titration in a conventional manner. You will
deliver volumes of NaOH titrant from a buret. You will enter the buret readings manually to
store and graph each pH-volume data pair.
If you choose Method 3, you will use a Vernier Drop Counter to conduct the titration. NaOH
titrant is delivered drop by drop from the reagent reservoir through the Drop Counter slot. After
the drop reacts with the reagent in the beaker, the volume of the drop is calculated and a
pH-volume data pair is stored.
Materials
Materials for both Method 2 (buret) and Method 3 (Drop Counter)
LabQuest
magnetic stirrer
LabQuest App
stirring bar or Microstirrer
Vernier pH Sensor
wash bottle
weighing dish or weighing paper
distilled water
solid potassium hydrogen phthalate
ring stand
solid sodium hydroxide
utility clamp
pipet bulb or pump
250 mL beaker
250 mL Erlenmeyer flask
100 mL graduated cylinder
balance (± 0.0001 g)
Materials required only for Method 1 & 2 (buret)
50 mL buret
buret clamp
Materials required only for Method 3 (Drop Counter)
Vernier Drop Counter
100 mL beaker
60 mL reagent reservoir
10 mL graduated cylinder
a second 250 mL beaker
Preparation of stock sodium hydroxide solution
1. Obtain and wear goggles.
2. Boil about 1L of distilled water for about 10 minutes. Let water cool
3. Measure out the 4g of NaOH and add it to the 1L volumetric flask. Add enough water to fill
about half the flask with distilled water. Swirl the flask to dissolve the solid. CAUTION:
Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.
4. Once the solid has dissolved, add enough water to prepare 1L of solution.
47 | P a g e
Method 1: measuring volume using a buret
1. Measure out 0.2g to 0.3g to the closes 0.1 mg of KHP. Dissolve the KHP in about 50 mL of
distilled water in a 125 mL Erlenmeyer flask. Add 2-3 drops of Phenolphthalein indicator.
2. Set up a ring stand, buret clamp, and buret to conduct a titration (see Figure 1). Equilibrate
the buret with the NaOH solution.
3. Fill the buret with the NaOH solution. Determine the initial volume of the NaOH solution.
4. While swirling the flask, slowly add the NaOH solution to the flask containing the KHP.
Watch for the phenolphthalein to change color from colorless to pink. The pink color should
persist for about 1 minute before returning to the colorless state. Record the final volume of
NaOH. Determine the volume of NaOH used by subtracting the initial volume from final
volume.
Method 2: measuring volume using a buret
1. Measure out 0.2g to 0.3g to the closes 0.1 mg of KHP. Dissolve the KHP in about 50 mL of
distilled water in a 250 mL beaker. Place the beaker of KHP solution on a magnetic stirrer
and add a stirring bar. If no magnetic stirrer is available, stir the reaction mixture with a
stirring rod during the titration.
2. Connect the pH Sensor to LabQuest and choose New from the File menu. If you have an
older sensor that does not auto-ID, manually set up the sensor.
3. Set up a ring stand, buret clamp, and buret to conduct a titration (see Figure 1). Equilibrate
the buret with the NaOH solution.
4. Use a utility clamp to suspend the pH Sensor on a ring stand as shown in Figure 1. Position
the pH Sensor in the KHP solution and adjust its position so that it is not struck by the
stirring bar. Gently stir the beaker of solution.
5. On the Meter screen, tap Mode. Change the data-collection mode to Events with Entry. Enter
the Name (Volume) and Unit (mL).Collect titration data. Conduct the titration carefully, as
described below.
a. Start data collection.
b. Before you have added any of the NaOH solution, tap Keep and enter 0 as the buret
volume in mL. Select OK to store the first data pair.
c. Add the next increment of NaOH titrant (enough to raise the pH about 0.15 units). When
the pH stabilizes, tap Keep, and enter the current buret reading as precisely as possible.
Select OK to store the second data pair.
d. When a pH value of approximately 6.0 is reached, change to 1–3 drop increments. Enter a
new buret reading after each increment. At about pH 6.7, add NaOH one drop at a time.
48 | P a g e
e. After a pH value of approximately 10 is reached, again add larger increments that raise the
pH by about 0.15 pH units, and enter the buret reading after each increment.
f. Continue adding NaOH solution until the pH value remains constant.
6. Stop data collection to view a graph of pH vs. volume. Dispose of the reaction mixture as
directed. Rinse the pH Sensor with distilled water in preparation for a second titration.
7. Examine your titration data to identify the region where the pH made the greatest increase.
The equivalence point is in this region.
a. To examine the data pairs on the displayed graph, select any data point.
b. As you move the examine line, the pH and volume values of each data point are displayed
to the right of the graph.
c. Identify the equivalence point as precisely as possible and record this information.
d. Store the data from the first run by tapping the File Cabinet icon.
8. An alternate way of determining the precise equivalence point of the titration is to take the
first and second derivatives of the pH-volume data.
Determine the peak value on the first derivative vs. volume plot.
a. Tap the Table tab and choose New Calculated Column from the Table menu.
b. Enter d1 as the Calculated Column Name. Select the equation 1st Derivative (Y,X). Use
Volume as the Column for X and pH as the Column for Y. Select OK.
c. On the displayed plot of d1 vs. volume, examine the graph to determine the volume at the
peak value of the first derivative.
9. Determine the zero value on the second derivative vs. volume plot.
d. Tap Table and choose New Calculated Column from the Table menu.
e. Enter d2 as the Calculated Column Name. Select the equation 2nd Derivative (Y,X). Use
Volume as the Column for X and pH as the Column for Y. Select OK.
f. On the displayed plot of d2 vs. volume, examine the graph to determine the volume when
the 2nd derivative equals approximately zero.
10. Repeat the titration with a second KHP solution. Analyze the titration results in a manner
similar to your first trial and record the equivalence point.
11. Print the graph directly from LabQuest, if possible. Alternately, transfer the data to a
computer, using Logger Pro software. Print a copy of the graph of each titration.
Method 3: measuring volume using a drop counter
1. Obtain and wear goggles.
49 | P a g e
2. Measure out 100 mL of distilled water into a 250 mL Erlenmeyer flask.
3. Measure out 0.2g to 0.3g to the closes 0.1 mg of KHP. Dissolve the KHP in about 50 mL of
distilled water in a 250 mL beaker. Place the beaker of KHP solution on a magnetic stirrer
and add a stirring bar. Swirl the flask to dissolve the solid. CAUTION: Sodium hydroxide
solution is caustic. Avoid spilling it on your skin or clothing.
4. Measure out the mass of KHP that will completely neutralize 10 mL of 0.10 M NaOH
solution. Dissolve the KHP in about 40 mL of distilled water in a 100 mL beaker.
Figure 2
5. Connect the pH Sensor to LabQuest. Lower the Drop Counter onto a ring stand and connect
it to DIG 1. Choose New from the File menu. If you have older sensors that do not auto-ID,
manually set up your sensors.
6. Obtain the plastic 60 mL reagent reservoir. Close both valves by turning the handles to a
horizontal position. Follow the steps below to set up the reagent reservoir for the titration.
a. Rinse the reagent reservoir with a few mL of the 0.10 M NaOH solution and pour the
NaOH into an empty 250 mL beaker.
b. Use a utility clamp to attach the reservoir to the ring stand.
c. Fill the reagent reservoir with slightly more than 60 mL of the 0.10 M NaOH solution.
d. Place the 250 mL beaker, which contains the rinse NaOH, beneath the tip of the reservoir.
e. Drain a small amount of NaOH solution into the 250 mL beaker so that it fills the
reservoir’s tip. To do this, turn both valve handles to the vertical position for a moment,
then turn them both back to horizontal.
f. Discard the drained NaOH solution in the 250 mL beaker as directed.
50 | P a g e
7. Calibrate the Drop Counter so that a precise volume of titrant is recorded in units of
milliliters.
a. Choose Calibrate from the Sensors menu and select Drop Counter.
b. If you have previously calibrated the drop size of your reagent reservoir and want to
continue with the same drop size, select Equation. Enter the value for the Drops/mL and
select Apply. Select OK and proceed directly to Step 8.
c. If you want to perform a new calibration continue with this step.
d. Select Start.
e. Place a 10 mL graduated cylinder directly below the slot on the Drop Counter, lining it up
with the tip of the reagent reservoir.
f. Open the bottom valve on the reagent reservoir (vertical). Keep the top valve closed
(horizontal).
g. Slowly open the top valve of the reagent reservoir so that drops are released at a slow rate
(~1 drop every two seconds). You should see the drops being counted on the screen.
h. When the volume of NaOH solution in the graduated cylinder is between 9 and 10 mL,
close the bottom valve of the reagent reservoir.
i. Enter the precise volume of NaOH and select Stop. Record the number of drops/mL
displayed on the screen for possible future use. Select OK.
j. Discard the NaOH solution in the graduated cylinder as directed, and set the graduated
cylinder aside.
8. Assemble the apparatus.
a. Place the magnetic stirrer on the base of the ring stand.
b. Insert the pH Sensor through the large hole in the Drop Counter.
c. Attach the Microstirrer to the bottom of the pH Sensor. Rotate the paddle wheel of the
Microstirrer, and make sure that it does not touch the bulb of the pH Sensor.
d. Adjust the positions of the Drop Counter and reagent reservoir so they are both lined up
near the center of the magnetic stirrer.
e. Lift up the pH Sensor, and slide the 100 mL beaker containing the KHP solution onto the
magnetic stirrer. Lower the pH Sensor into the beaker.
f. Adjust the position of the Drop Counter so that the Microstirrer on the pH Sensor is just
touching the bottom of the beaker.
g. Adjust the reagent reservoir so its tip is just above the Drop Counter slot.
h. Turn on the magnetic stirrer so that the Microstirrer is stirring at a fast rate.
51 | P a g e
9. You are now ready to perform the titration.
a. Start data collection. No data will be collected until the first drop goes through the Drop
Counter slot.
b. Fully open the bottom valve. The top valve should still be adjusted so drops are released at
a rate of about 1 drop every 2 seconds. When the first drop passes through the Drop
Counter slot, check the graph to see that the first data pair was recorded.
c. Continue watching your graph to see when a large increase in pH takes place—this will be
the equivalence point of the reaction. When this jump in pH occurs, let the titration
proceed for a few more milliliters of titrant.
d. Stop data collection to view a graph of pH vs. volume.
e. Turn the bottom valve of the reagent reservoir to a closed (horizontal) position.
10.
Dispose of the reaction mixture as directed. Rinse the pH Sensor with distilled water in
preparation for a second titration.
11.
Examine your titration data to identify the region where the pH made the greatest increase.
The equivalence point is in this region.
a. To examine the data pairs on the displayed graph, select any data point.
b. As you move the examine line, the pH and volume values of each data point are displayed
to the right of the graph.
c. Identify the equivalence point as precisely as possible and record this information.
d. Store the data from the first run by tapping the File Cabinet icon.
12.
An alternate way of determining the precise equivalence point of the titration is to take the
first and second derivatives of the pH-volume data.
Determine the peak value on the first derivative vs. volume plot.
a. Tap the Table tab and choose New Calculated Column from the Table menu.
b. Enter d1 as the Calculated Column Name. Select the equation 1st Derivative (Y,X). Use
Volume as the Column for X and pH as the Column for Y. Select OK.
c. On the displayed plot of d1 vs. volume, examine the graph to determine the volume at the
peak value of the first derivative.
Determine the zero value on the second derivative vs. volume plot.
d. Tap Table and choose New Calculated Column from the Table menu.
e. Enter d2 as the Calculated Column Name. Select the equation 2nd Derivative (Y,X). Use
Volume as the Column for X and pH as the Column for Y. Select OK.
52 | P a g e
f. On the displayed plot of d2 vs. volume, examine the graph to determine the volume when
the 2nd derivative equals approximately zero.
13.
Conduct another titration with a second KHP solution. Analyze the titration results in a
manner similar to your first trial and record the equivalence point.
14.
Print the graph directly from LabQuest, if possible. Alternately, transfer the data to a
computer, using Logger Pro software. Print a copy of the graph of each titration.
53 | P a g e
Name: _______________________________ Date: ______________________________
Group Members: _______________________ Class & Section: ____________________
Pre-lab exercise
1. Calculate the mass of sodium hydroxide needed to prepare 1000 mL of a 0.10 M solution.
2. Calculate the mass of KHP needed to react completely with 25 mL of a 0.10 M NaOH
solution. Consider the reaction equation to be as shown below.
HP– (aq) + OH– (aq) → H2O(l) + P2– (aq)
54 | P a g e
Data table
Trial
KHP
Sample
(g)
Final
Volume
(mL)
Initial
Volume
(mL)
Equivalence
point
(mL)
1 (Method 1)
---
2 (Method 1)
---
3 (Method 1)
---
4 (Method 1)
---
Molarity
NaOH
(M)
5 (Method 2)
6 (Method 2)
7 (Method 2)
8 (Method 2)
9 (Method 3)
10 (Method 3)
11 (Method 3)
12 (Method 3)
55 | P a g e
Data analysis
1. Calculate the molar concentration of the NaOH solution that you prepared for each trial using
Method 1, Method 2 and Method 3.
2. Determine all the relevant statistical analysis for the molarity of NaOH for each method.
56 | P a g e
Determination of purity of unknown KHP
The unknown should already have been dried at 110-120oC for 1 hour. Use 0.6 to 0.8 g of
unknown for each titration. Choose the method you think will give you the best result for
determining the amount of KHP in your unknown sample. Obtain at least 3 titrations. (You should
receive enough unknown for at least 6 trials.)
Evaluation - the accuracy and precision of your result will be graded on a non-linear sliding scale
with a grade of 100% being awarded for results within 1% error. Penalties will be assessed for the
following:
1. Need more unknown due to running out.
2. Calculation error in result
Data table
Trial
KHP
Sample
(g)
Final
Volume
(mL)
Initial
Volume
(mL)
Equivalence
point
(mL)
Percent KHP
1
2
3
4
5
6
7
8
Data Analysis
1. Calculate the percent KHP for each of your trials.
57 | P a g e
2. Determine all relevant statistical analysis for the percent KHP
58 | P a g e
Experiment 5: Titration-Redox
Standardization of Sodium Thiosulfate &
Analysis of Vitamin C
Introduction
Oxidation is defined as a loss of electrons, and reduction is a gain of electrons. Sodium
thiosulfate is a very common reducing agent. Sodium thiosulfate solutions must be standardized
to determine their accurate molarity. In this experiment we will standardize a sodium thiosulfate
solution by doing a redox titration with a solution of primary standard potassium iodate. Since
all reactants and products are colorless, we must perform the titration indirectly using an
intermediate reaction. The titration endpoint is determined using a starch-iodine color change
from blue to colorless. Once the Sodium Thiosulfate has been standardized, the vitamin C
content of commercial tablets is determined titrimetrically and compared with the manufacturers’
specifications.
Vitamin C is ascorbic acid which is rapidly and quantitatively oxidized by iodine in acid solution
according to the following equation:
The standard method for determination of ascorbic acid involves the direct titration of the acidified
sample with a standard iodine solution. But the low solubility of iodine makes this procedure less
than ideal.
The proposed experiment avoids these difficulties by generating in situ a known excess quantity
of iodine by the reaction between iodate and iodide:
IO3–(aq) + 5I–(aq) + 6H+(aq)
3I2(aq) + 3H2O(l)
After reacting with ascorbic acid, the remaining iodine may then be titrated against standard
sodium thiosulfate solution.
2 S2O32-(aq)
S4O62- (aq) + 2e-
59 | P a g e
2e- + I2
2I-
Procedure
[Hazard Warning: Potassium iodate is oxidizing, and 0.5 M sulfuric acid is
irritant.]
Part A: Standardize Thiosulfate Solution
1. Accurately weigh out 0.6 g to 0.7 g of potassium iodate. Record the mass.
2. Dissolve this in deionized water and make up to 250 mL in a volumetric flask.
3. Use part of this iodate solution to standardize the given thiosulphate solution as follows:
a) Pipette 25.0 mL of the potassium iodate solution into a conical flask, add to it about 5 mL
of 1 M potassium iodide solution followed by about 10 mL of 0.5 M sulfuric acid, then
immediately titrate with sodium thiosulphate solution.
b) Add a few drops of freshly prepared starch solution when the reaction mixture turns pale
yellow and continue to titrate to the end point (colorless). Record your results in Table 1.
Part B: Determine the Amount of Vitamin C in One Tablet
4. Dissolve ONE Vitamin C tablet provided in about 150 mL of 0.5 M sulfuric acid.
5. Transfer the resulting solution to a clean volumetric flask and make up to 250 mL using
deionized water.
6. Pipette 10.0 mL of the vitamin C solution into a conical flask and add to it 5 mL of 1 M
potassium iodide solution.
7. Pipette 25.0 mL of the standard potassium iodate solution into the flask containing vitamin C
and potassium iodide.
8. Titrate the iodine content immediately with the standardized sodium thiosulfate solution as in
(3). Record your results in Table 2.
Results
Mass of weighing bottle and potassium iodate:
_______________ g
Mass of weighing bottle:
_______________ g
Mass of potassium iodate:
_______________ g
60 | P a g e
Table 1
Trial
First
Second
Third
Trial
First
Second
Third
Final burette reading (mL)
Initial burette reading (mL)
Volume of Na2S2O3 (mL)
Mean titration (mL)
Table 2
Final burette reading (mL)
Initial burette reading (mL)
Volume of Na2S2O3 (mL)
Mean titration (mL)
Molarity of standard KIO3 solution
Molarity of sodium thiosulfate solution
Trial 1 __________M
Trial 2 __________M
61 | P a g e
Trial 3 __________M
Average (mean) molarity of sodium thiosulfate solution
______________M
Questions
1. What is the function of the starch solution? Why should it be added only when the reaction
mixture becomes pale yellow?
2. Why is it preferable to carry out the titration immediately after the addition of sulfuric acid
in step 3(b)?
3. It is well known that ascorbic acid (vitamin C) deteriorates on heating and on exposure to
air. Describe, giving experimental details, how you would investigate these two factors.
62 | P a g e
4. What is the mass of the ascorbic acid per tablet? How does the result compare with the
manufacturer’s specifications?
5. The net ionic equation for the half-reaction for the conversion of potassium iodate to iodine
in acidic solution is
12H+(aq) + 2IO3- (aq) + 10 e- -----> 3I2(s) + 6H2O(l)
Why is this considered a half-equation? Does this half-reaction represent an oxidation or
a reduction?
63 | P a g e
Experiment 6: Beer’s Law
Determining the Concentration
of a Solution: Beer’s Law
The primary objective of this experiment is to determine the concentration of an unknown
copper (II) sulfate solution. The CuSO4 solution used in this experiment has a blue color, so
Colorimeter users will be instructed to use the red LED. Spectrometer users will determine an
appropriate wavelength based on the absorbance spectrum of the solution. A higher
concentration of the colored solution absorbs more light (and
transmits less) than a solution of lower concentration.
You will prepare five copper (II) sulfate solutions of known
concentration (standard solutions). Each solution is transferred to
a small, rectangular cuvette that is placed into the Colorimeter or
Spectrometer. The amount of light that penetrates the solution
and strikes the photocell is used to compute the absorbance of
each solution. When you graph absorbance vs. concentration for
Figure 1
the standard solutions, a direct relationship should result. The
direct relationship between absorbance and concentration for a solution is known as Beer’s law.
You will determine the concentration of an unknown CuSO4 solution by measuring its
absorbance. By locating the absorbance of the unknown on the vertical axis of the graph, the
corresponding concentration can be found on the horizontal axis. The concentration of the
unknown can also be found using the slope of the Beer’s law curve.
Objectives
In this experiment, you will

Prepare and test the absorbance of five standard copper (II) sulfate solutions.

Calculate a standard curve from the test results of the standard solutions.

Test the absorbance of a copper (II) sulfate solution of unknown molar concentration.

Calculate the molar concentration of the unknown CuSO4 solution.
Materials
LabQuest
0.40 M copper (II) sulfate, CuSO4, solution
LabQuest App
copper (II) sulfate, CuSO4, unknown solution
Vernier Colorimeter or Spectrometer
pipet pump or pipet bulb
one cuvette
distilled water
64 | P a g e
five 20 × 150 mm test tubes
test tube rack
two 10 mL pipets or graduated cylinders
stirring rod
two 100 mL beakers
tissues (preferably lint-free)
Procedure
Both Colorimeter and Spectrometer Users
1. Obtain and wear goggles.
2. Obtain small volumes of 0.40 M CuSO4 solution and distilled water in separate beakers.
3. Label five clean, dry, test tubes 1–5. Use pipets to prepare five standard solutions according
to the chart below. Thoroughly mix each solution with a stirring rod. Clean and dry the
stirring rod between uses.
Test
Tube
1
2
3
4
5
0.40 M CuSO4
(mL)
2
4
6
8
~10
Distilled H2O
(mL)
8
6
4
2
0
Concentration
(M)
0.080
0.16
0.24
0.32
0.40
4. Prepare a blank by filling a cuvette 3/4 full with distilled water. To correctly use cuvettes,
remember:

Wipe the outside of each cuvette with a lint-free tissue.

Handle cuvettes only by the top edge of the ribbed sides.

Dislodge any bubbles by gently tapping the cuvette on a hard surface.

Always position the cuvette so the light passes through the clear sides.
Spectrometer Users Only
5. Connect the Spectrometer to LabQuest and choose New from the File menu.
6. Calibrate the Spectrometer.
a. Place the blank cuvette in the Spectrometer.
b. Choose Calibrate from the Sensors menu. The following message is displayed: “Waiting
60 seconds for lamp to warm up.” After 60 seconds, the message will change to “Warmup
complete.”
c. Select Finish Calibration. When the message “Calibration completed” appears, select OK.
7. Determine the optimal wavelength for creating this standard curve and set up the
data-collection mode.
a. Remove the blank cuvette, and place the 0.40 M standard into the cuvette slot.
65 | P a g e
b. Start data collection. A full spectrum graph of the solution will be displayed. Stop data
collection. The wavelength of maximum absorbance ( max) is automatically identified.
c. Tap the Meter tab. On the Meter screen, tap Mode. Change the mode to Events with
Entry.
d. Enter the Name (Concentration) and Units (mol/L). Select OK.
Colorimeter Users Only (Spectrometer users proceed to the Spectrometer section)
5. Connect the Colorimeter to LabQuest and choose New from the File menu.
6. Calibrate the Colorimeter.
a. Place the blank in the cuvette slot of the Colorimeter and close the lid.
b. Press the < or > button on the Colorimeter to select the wavelength of 635 nm (Red). Then
calibrate by pressing the CAL button on the Colorimeter. When the LED stops flashing,
the calibration is complete.
7. Set up the data-collection mode.
a. On the Meter screen, tap Mode. Change the mode to Events with Entry.
b. Enter the Name (Concentration) and Units (mol/L). Select OK.
c. Proceed to Step 8.
Both Colorimeter and Spectrometer Users
8. You are now ready to collect absorbance-concentration data for the five standard solutions.
a. Start data collection.
b. Using the solution in Test Tube 1, rinse the cuvette twice with ~1 mL amounts and then
fill it 3/4 full. Wipe the outside with a tissue and place it in the device (Colorimeter or
Spectrometer). Close the lid of the Colorimeter.
c. When the value displayed on the screen has stabilized, tap Keep and enter 0.080 as the
concentration in mol/L. Select OK. The absorbance and concentration values have now
been saved for the first solution.
d. Discard the cuvette contents as directed. Using the solution in Test Tube 2, rinse and fill
the cuvette 3/4 full. Wipe the outside and place the cuvette in the device (close the lid of
the Colorimeter). Wait for the value displayed on the screen to stabilize, and tap Keep.
Enter 0.16 as the concentration in mol/L.
e. Repeat the procedure for Test Tubes 3 and 4. Trial 5 is the original 0.40 M CuSO4
solution. Note: Do not test the unknown solution until Step 11.
f. When you have finished testing the standard solutions, stop data collection.
66 | P a g e
g. To examine the data pairs on the displayed graph, tap any data point. As you tap each data
point, the absorbance and concentration values are displayed to the right of the graph.
9. Write down the absorbance values, for each of the five trials, in your data table.
10. Display a graph of absorbance vs. concentration with a linear regression curve.
a. Choose Graph Options from the Graph menu.
b. Select Autoscale from 0 and select OK.
c. Choose Curve Fit from the Analyze menu.
d. Select Linear as the Fit Equation. The linear-regression statistics for these two data
columns are displayed for the equation in the form: y = mx + b where x is concentration,
y is absorbance, a is the slope, and b is the y-intercept. Note: One indicator of the quality
of your data is the size of b. It is a very small value if the regression line passes through or
near the origin. The correlation coefficient, r, indicates how closely the data points match
up with (or fit) the regression line. A value of 1.00 indicates a nearly perfect fit.
e. Select OK. The graph should indicate a direct relationship between absorbance and
concentration, a relationship known as Beer’s law. The regression line should closely fit
the five data points and pass through (or near) the origin of the graph.
11. Determine the absorbance and concentration values of the unknown CuSO4 solution.
a. Tap the Meter tab.
b. Obtain about 5 mL of the unknown CuSO4 in another clean, dry, test tube. Record the
number of the unknown in your data table.
c. Rinse the cuvette twice with the unknown solution and fill it about 3/4 full. Wipe the
outside of the cuvette and place it into the device (close the lid of the Colorimeter).
d. Monitor the absorbance value. When this value has stabilized, record it in your data table.
e. Tap the Graph tab.
f. On the Graph screen, choose Interpolate from the Analyze menu. Tap any point on the
regression curve (or use the ◄ or ► keys on LabQuest) to find the absorbance value that
is closest to the absorbance reading you obtained in Step 11d. Determine the concentration
of your unknown CuSO4 solution and record the concentration in your data table.
67 | P a g e
Data table
Trial
Concentration (mol/L)
1
0.080
2
0.16
3
0.24
4
0.32
5
0.40
6
Unknown number ____
Absorbance
Data analysis
1. Describe an alternate method for determining the molar concentration of your unknown
sample of copper (II) sulfate solution, using the standard data.
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Determining the Quantity
of Iron in a Vitamin Tablet
As biochemical research becomes more sophisticated, we are learning more about the role of
metallic elements in the human body. For example, copper and zinc are present in enzymes, and
trace amounts of molybdenum and selenium are vital in regulating internal oxidation-reduction
reactions. Iron is necessary for oxygen transport in the bloodstream. Many people gain these
essential elements through their diets, or by taking multivitamin tablets.
The iron that is present in these tablets is in the form of water-soluble Fe2+ ions. In this
experiment, you will use the Colorimeter or Spectrometer to determine the quantity of iron in a
vitamin tablet. You will prepare four solutions of known Fe2+ concentration. The Colorimeter or
Spectrometer will be used to determine the absorbance of each solution at a specific wavelength
of light. When a graph of absorbance vs. concentration is plotted for these solutions, a direct
relationship should result (Beer’s law).
You will also use a Colorimeter or Spectrometer to measure the absorbance of a solution
prepared from a multivitamin tablet. This absorbance value is converted to concentration using
the Beer’s law curve. From this concentration, the total mass of iron in the original tablet can be
calculated.
The iron in each tablet tested will be reacted with 1,10 phenanthroline so that it develops a
detectable color. Sodium acetate and hydroxylamine hydrochloride are added to control the pH
of the mixtures and keep the iron in the +2 oxidation state.
Objectives
In this experiment, you will

Prepare Fe2+ standard solutions.

Measure the absorbance of each standard solution.

Plot a graph of absorbance of Fe2+ vs. concentration.

Use your Beer’s law plot to determine the quantity of iron in a vitamin tablet.
Materials
LabQuest
20 mg/L iron (II) solution
LabQuest App
1% hydroxylamine hydrochloride
Vernier Colorimeter or Spectrometer
1 M sodium acetate
one cuvette
1% 1,10 phenanthroline
four 20 X 150 mm test tubes
0.010 M HCl
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test tube rack
distilled water
mortar and pestle
50 mL graduated cylinder
hot plate
two 100 mL beakers
filter funnel
two 10 mL graduated cylinders
filter paper
pipet pump or pipet bulb
stirring rod
two 10 mL pipets
multivitamin tablet with iron
Procedure
Both Colorimeter and Spectrometer Users
1. Obtain and wear goggles.
2. Weigh one multivitamin tablet and record its mass to the nearest 0.001 g. Use a mortar and
pestle to grind the tablet to a fine powder. Also record the mass of iron, in mg, listed on the
label of the vitamin container.
3. Measure out 0.200 g of the powdered tablet and transfer it to a 100 mL beaker. Note: Try to
avoid including bits of the tablet coating in your 0.200 g sample. Add 20 mL of 0.010 M HCl
and 10 mL of distilled water. Mix thoroughly. CAUTION: Handle the hydrochloric acid
with care. It can cause painful burns if it comes in contact with the skin.
4. Use a hot plate to heat the mixture in the beaker to a gentle boil. Boil the mixture for
5–7 minutes. Frequently stir the mixture with a stirring rod. After boiling, allow the mixture
to cool to nearly room temperature. Some particles may remain undissolved, which is normal
and will not affect your experiment.
5. Set up a funnel and filter paper. Pour the mixture through the filter, collecting the liquid
filtrate in a 50 mL graduated cylinder. Rinse the 100 mL beaker with 10 mL of distilled water
and pour the water through the filter. When the filter has drained, fill the graduated cylinder
to the 50 mL mark with distilled water. Transfer this solution to a clean, dry, 100 mL beaker.
6. Label five clean, dry, test tubes 1–5. Prepare a set of four standards and one blank solution in
the test tubes using the following solutions: CAUTION: Handle the solutions with care.
A. 20 mg/L iron (II) solution
B. 1% hydroxylamine hydrochloride
C. 1 M sodium acetate
D. 1% 1,10 phenanthroline
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Note: Solutions B, C, D, and the distilled water may be measured successively into one
10 mL graduated cylinder. Then add solution A, using a pipet for greater precision. Transfer
the solution into a test tube. Important: Mix each solution thoroughly with a stirring rod.
Test Tube
number
1
2
3
4
5 (blank)
A
(mL)
1.00
2.00
3.00
4.00
0.00
B
(mL)
1.00
1.00
1.00
1.00
1.00
C
(mL)
1.00
1.00
1.00
1.00
1.00
D
(mL)
1.00
1.00
1.00
1.00
1.00
Distilled water
(mL)
6.00
5.00
4.00
3.00
7.00
7. Prepare a blank by filling an empty cuvette 3/4 full with distilled water. Seal the cuvette with
a lid. To correctly use a Colorimeter cuvette, remember:

All cuvettes should be wiped clean and dry on the outside with a tissue.

Handle cuvettes only by the top edge of the ribbed sides.

All solutions should be free of bubbles.

Always position the cuvette so the light passes through the clear sides.
Colorimeter Users Only (Spectrometer users proceed to the Spectrometer section)
8. Connect the Colorimeter to LabQuest and choose New from the File menu.
9. Calibrate the Colorimeter.
a. Place the blank in the cuvette slot of the Colorimeter and close the lid.
b. Press the < or > button on the Colorimeter to set the wavelength to 565 nm (Green). Then
calibrate by pressing the CAL button on the Colorimeter. When the LED stops flashing,
the calibration is complete.
10. Set up the data-collection mode.
a. On the Meter screen, tap Mode. Change the data-collection mode to Events with Entry.
b. Enter the Name (Concentration) and Units (mg/L). Select OK.
c. Proceed to Step 11.
Spectrometer Users Only
8. Connect the Spectrometer to LabQuest and choose New from the File menu.
9. Calibrate the Spectrometer.
a. Place the blank cuvette in the Spectrometer.
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b. Choose Calibrate from the Sensors menu. The following message is displayed: “Waiting
60 seconds for lamp to warm up.” After 60 seconds, the message will change to “Warmup
complete.”
c. Select Finish Calibration. When the message “Calibration completed” appears, select OK.
10. Determine the optimal wavelength for creating this standard curve and set up the
data-collection mode.
a. Remove the blank cuvette and pour out the water. Using the solution in Test Tube 1, rinse
the cuvette twice with ~1-mL amounts and then fill it 3/4 full. Wipe the outside with a
tissue and place it in the Spectrometer.
b. Start data collection. A full spectrum graph of the solution will be displayed. Stop data
collection. The wavelength of maximum absorbance ( max) is automatically identified.
c. Tap the Meter tab. On the Meter screen, tap Mode. Change the mode to Events with
Entry.
d. Enter the Name (Concentration) and Units (mg/L). Select OK.
Both Colorimeter and Spectrometer Users
11. You are now ready to collect absorbance-concentration data for the five standard solutions.
a. Leave the cuvette, containing the Test Tube 1 mixture, in the device (Colorimeter or
Spectrometer). Close the lid on the Colorimeter. Start the data collection.
b. When the value displayed on the monitor has stabilized, tap Keep and enter 2.00 as the
concentration in mg/L. Select OK. The absorbance and concentration values have now
been saved for the first solution.
c. Discard the cuvette contents as directed by your instructor. Using the solution in Test
Tube 2, rinse the cuvette twice with ~1 mL amounts and then fill it 3/4 full. Place the
cuvette in the device. When the value displayed on the screen has stabilized, tap Keep and
enter 4.00 as the iron concentration in mg/L. Select OK. The absorbance and
concentration values have now been saved for the first solution.
d. Repeat the procedure for Test Tube 3 (6.00 mg/L) and Test Tube 4 (8.00 mg/L).
e. Stop data collection.
f. Examine the data pairs on the graph of absorbance vs. concentration. As you tap each data
point, the absorbance and concentration values are displayed to the right of the graph.
Record the absorbance values in your data table.
12. Prepare the tablet solution in the 100 mL beaker for testing.
a. Measure out 5.00 mL of the tablet solution into a second clean, dry, 100 mL beaker.
b. Add 5.00 mL each of solutions B, C, and D, plus 30.0 mL of distilled water.
72 | P a g e
c. Mix the solution thoroughly.
13. Find the absorbance of the tablet solution.
a. Rinse the cuvette twice with the tablet solution you prepared in Step 12, and fill it about
3/4 full.
b. Wipe the outside of the cuvette and place it in the device. CAUTION: Handle the
solutions with care.
c. Tap the Meter tab and monitor the absorbance value. When this value has stabilized,
record the absorbance value for the tablet solution in your data table. Note: Do not exit the
LabQuest App; you will analyze the data and the graph momentarily. Proceed to Step 1 of
Processing the Data.
d. Dispose of the remaining solution in the flask as directed by your instructor.
14.
Discard the solutions as directed by your instructor. Proceed directly to Steps 1–4 of
Processing the Data.
Processing the data
1. To determine the iron concentration in tablet solution, you can interpolate along the
regression line on your Beer’s law curve. Use the following method to convert the
absorbance value of the unknown to concentration, in mg/L.
a. Tap the Graph tab and choose Curve Fit from the Analyze menu.
b. Select Linear as the Fit Equation. The linear-regression statistics for these two data
columns are displayed for the equation in the form
y = ax + b
where x is concentration, y is absorbance, a is the slope, and b is the y-intercept.
Note: One indicator of the quality of your data is the size of b. It is a very small value if
the regression line passes through or near the origin. The correlation coefficient, r,
indicates how closely the data points match up with (or fit) the regression line. A value of
1.00 indicates a nearly perfect fit.
c. Select OK. The graph should indicate a direct relationship between absorbance and
concentration, a relationship known as Beer’s law. The regression line should closely fit
the five data points and pass through (or near) the origin of the graph.
d. Choose Interpolate from the Analyze menu. Interpolate along the regression curve to
determine the concentration of the unknown solution. Tap any point on the regression
curve (or use the ◄ or ► key of LabQuest) to find to the absorbance value that is closest
to the absorbance reading you obtained in Step 13. The corresponding iron concentration
of your tablet solution, in mg/L, will be displayed to the right of the graph. Record this
value in your data table.
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2. Print a graph of absorbance vs. concentration, with a regression line and interpolated
unknown concentration displayed.
3. Determine the mass of iron in your vitamin tablet.
a. Because the tablet solution in Step 12 was diluted by a factor of 1/10, multiply your final
iron concentration (in mg/L) by 10 to find the pre-dilution iron concentration, in mg/L.
b. A portion of the tablet was dissolved in 50 mL of solution during the procedure. Use this
volume and the answer in the previous step to calculate the mass of iron, in mg, in the
tablet solution.
c. Since the analysis used only a fraction of the total mass of the tablet (see Step 3), multiply
the mg of iron in the previous step by the ratio of the total tablet mass to the 0.200 g
portion:
mg of iron X
total tablet mass
= mg of iron in the vitamin tablet
0.200 g
4. Compare your final answer in the previous step with the mg of iron shown on the label of
vitamin tablets you tested. Use these values to calculate the percent error.
Data and calculations
Mass of tablet
g
Mass of iron on label
mg
Test Tube
Number
Concentration
(mg/L)
1
2.00
2
4.00
3
6.00
4
8.00
Absorbance
tablet solution
Iron concentration in the pre-dilution
tablet solution (use a factor of 10)
Mass of iron in the original tablet solution
mg/L
mg
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Calculated mass of iron in vitamin tablet
Percent error
mg
%
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Analysis of Aspirin
Aspirin, the ubiquitous pain reliever, goes by the chemical name acetylsalicylic acid. One of the
compounds used in the synthesis of aspirin is salicylic acid, which is itself a pain reliever that
was known to many ancient cultures, including the Native Americans who extracted it from
willow tree bark. Salicylic acid is extremely bitter tasting, and frequent use can cause severe
stomach irritation. The search for a milder form of this pain reliever led to the successful
synthesis of acetylsalicylic acid by the German chemist Felix Hoffmann in 1893.
Objectives
In this experiment, you will

Conduct a colorimetric analysis of your aspirin sample.
Materials
Colorimeter Test Materials
LabQuest
solid salicylic acid, C6H4(OH)CO2H
LabQuest App
aspirin crystals
Vernier Colorimeter or Spectrometer
95% ethanol, CH3CH2OH
50 mL graduated cylinder
0.025 M iron (III) nitrate solution, Fe(NO3)3
plastic cuvette with lid
distilled water
250 mL beaker
100 mL volumetric flask
100 mL beaker
250 mL volumetric flask
Procedure
Test the Colorimetric Absorbance of an Aspirin Sample
Both Colorimeter and Spectrometer Users
1. Quantitatively prepare the stock salicylic, acid solution.
a. Measure out about 0.20 g of salicylic acid. Record the precise mass that you use.
b. Transfer the salicylic acid to a 250 mL beaker and add 10 mL of 95% ethanol. Swirl the
beaker to dissolve the solid.
c. Add 150 mL of distilled water to the beaker. Mix the solution.
d. Quantitatively transfer the solution from the beaker to a 250 mL volumetric flask.
Thoroughly rinse the beaker with several portions of distilled water, and transfer the rinse
water to the volumetric flask. Add distilled water, as needed, to fill the flask to the
250 mL mark. Mix the solution thoroughly. Calculate the precise molar concentration of
your stock solution and record it in your data table.
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2. Prepare four standard solutions of varying concentrations of salicylic acid.
a. To prepare 100 mL of your standard solution (the solution you will use for Trial 1),
quantitatively transfer 10 mL of the stock salicylic acid solution you prepared in Step 14
to a 100 mL volumetric flask.
b. Add 0.025 M Fe(NO3)3 solution to the flask to make precisely 100 mL.
c. Prepare the remaining three salicylic acid standard solutions according to the table below,
diluting the standard solution in the 100 mL flask with distilled water. Mix thoroughly.
Trial
1
2
3
4
Standard salicylic
acid solution
10.0
7.5
5.0
2.5
Water
(mL)
0
2.5
5.0
7.5
Calculate the precise molar concentrations of the four standard solutions in the table above
and record them in your data table.
Colorimeter Users Only (Spectrometer users proceed to the Spectrometer section)
3. Disconnect the Temperature Probe and connect the Colorimeter. Choose New from the File
menu.
4. Calibrate the Colorimeter.
a. Prepare a blank by filling an empty cuvette 3/4 full with distilled water.
b. Place the blank in the cuvette slot of the Colorimeter and close the lid
c. Press the < or > buttons on the Colorimeter to set the wavelength to 565 nm (Green). Then
calibrate by pressing the CAL button on the Colorimeter. When the LED stops flashing,
the calibration is complete.
5. Set up the data-collection mode.
a. On the Meter screen, tap Mode. Change the mode to Events with Entry.
b. Enter the Name (Concentration) and Units (mol/L). Select OK.
c. Proceed to Step 19.
Spectrometer Users Only
6. Disconnect the Temperature Probe and connect the Spectrometer. Choose New from the File
menu.
7. Calibrate the Spectrometer.
a. Place the blank cuvette in the Spectrometer.
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b. Choose Calibrate from the Sensors menu. The following message is displayed: “Waiting
60 seconds for lamp to warm up.” After 60 seconds, the message will change to “Warmup
complete.”
c. Select Finish Calibration. When the message “Calibration completed” appears, select OK.
8. Determine the optimal wavelength for creating the standard curve.
a. Remove the cuvette from your Spectrometer and pour out the water. Using the solution in
the first 100 mL volumetric flask of salicylic acid, rinse the cuvette twice with ~1 mL
amounts and then fill it 3/4 full. Wipe the outside with a tissue, place it in the
Spectrometer.
b. Start data collection. A full spectrum graph of the solution will be displayed. Stop data
collection. The wavelength of maximum absorbance ( max) is automatically identified.
c. Tap the Meter tab. On the Meter screen, tap Mode. Change the mode to Events with
Entry.
d. Enter the Name (Concentration) and Units (mol/L). Select OK.
e. Proceed to Step 19.
Both Colorimeter and Spectrometer Users
9. You are now ready to collect data for the four standard solutions.
a. Remove the cuvette from your device (Colorimeter or Spectrometer) and pour out the
water. Using the solution in the first 100 mL volumetric flask of salicylic acid, rinse the
cuvette twice with ~1 mL amounts and then fill it 3/4 full. Wipe the outside with a tissue
and place it in the device. (Close the lid on the Colorimeter.)
b. When the value displayed on the screen has stabilized, tap Keep and enter the molar
concentration. Select OK. The absorbance and concentration values have now been saved
for the first solution.
c. Discard the cuvette contents as directed by your instructor. Using the solution in the
second 100 mL volumetric flask, rinse the cuvette twice with ~1 mL amounts, and then
fill it 3/4 full. Wipe the outside, place it in the device, and close the lid. After closing the
lid, wait for the absorbance to stabilize and tap Keep. Enter the molar concentration, and
select OK.
d. Repeat the procedure for the remaining salicylic acid solutions that you prepared.
e. Stop data collection to view a graph of absorbance vs. concentration.
f. To examine the data pairs on the displayed graph, select any data point. As you tap each
point, the absorbance and concentration values of each data point are displayed to the right
of the graph. Record the absorbance and concentration values in your data table.
78 | P a g e
g. Choose Curve Fit from the Analyze menu.
h. Select Linear as the Fit Equation. Record the best-fit line equation in Part III of your data
table and select OK.
10. Prepare the aspirin sample for testing. Complete this step quickly.
a. Measure out about 0.40 gram of aspirin and transfer it to the 250 mL beaker. Record the
precise mass of aspirin that you use.
b. Add 10 mL of 95% ethanol to the beaker of aspirin sample. Swirl the mixture to dissolve
the solid.
c. Add 150 mL of distilled water to the beaker. Mix the solution.
d. Quantitatively transfer the solution from the beaker to a 250 mL volumetric flask.
Thoroughly rinse the beaker with several portions of distilled water, and transfer the rinse
water to the volumetric flask. Add distilled water, as needed, to fill the flask to the
250 mL mark. Mix the solution thoroughly.
e. Transfer 5 mL of the aspirin solution from the 250 mL volumetric flask to a clean and dry
100 mL volumetric flask. Add 0.025 M Fe(NO3)3 solution to the flask to make precisely
100 mL. Mix the solution thoroughly.
11. Measure and record the absorbance of the treated aspirin sample. This must be done within
5 minutes of completing Step 20.
a. Rinse and fill the cuvette 3/4 full with the sample. Cap the cuvette and place it in the
device (Close the lid of the Colorimeter).
b. Tap the Meter tab. If the absorbance value falls within the range of the salicylic acid
standard solutions, record it in your data table. If it does not, repeat Step 20d with a more
dilute or more concentrated sample, depending on the absorbance value from your first
test.
c. Repeat Parts a–c of this Step two times with new aliquots of the treated aspirin sample.
12. Discard all solutions as directed.
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Data table
Salicylic Acid Standard Stock Solution
Initial mass of salicylic acid (g)
Moles of salicylic acid (mol)
Initial molarity of salicylic acid (M)
Beer’s Law Data for Salicylic Acid Standard Solutions
Trial
Concentration (M)
Absorbance
1
2
3
4
Best-fit line equation for the salicylic acid standards
Determine the amount of Aspirin in Tablet
Initial mass of aspirin sample (g)
Absorbance of aspirin sample
Moles of salicylic acid in aspirin sample (mol)
Mass of salicylic acid in aspirin sample (g)
Mass of aspirin in sample (g)
Percent error in sample (%)
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