1. technology and computing
2. computer certification

CHML 1045
General Chemistry I:
Laboratory Manual
Fall 2014
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Table of Contents
Table of Contents .......................................................................................................................................................2
Table of Figures ..........................................................................................................................................................3
General Chemistry Lab: Your Mindset........................................................................................................................4
Chemistry Lab Software .............................................................................................................................................5
Important Instructions for Lab Success ......................................................................................................................7
Important Lab Protocols/Techniques .........................................................................................................................8
The Laboratory Notebook ..........................................................................................................................................9
Chemistry Lab Notebook and Participation Grading Rubric ................................................................................... 14
Laboratory Techniques ............................................................................................................................................ 16
Laboratory Safety Markings .................................................................................................................................... 22
Lab Safety Definitions .............................................................................................................................................. 23
Eastern Florida State College - Chemistry Lab Safety Rules .................................................................................... 25
Eastern Florida State College - Chemistry Lab Safety Rules .................................................................................... 29
Experiment 1 - Measurement and Common Laboratory Techniques ..................................................................... 33
Experiment 2 - Chemical Analysis, Determination of solution concentration by stepwise reactions .................... 51
Experiment 3 – Emission Spectroscopy ................................................................................................................... 55
Experiment 4 – Spectrophotometry and Molecular Structure ............................................................................... 63
Experiment 5 – Chemical Nomenclature................................................................................................................. 69
Experiment 6 - Chemical Stoichiometry: Determination of the formula of a complex salt .................................... 82
Experiment 7 – Covalent Bonding and Molecular Geometry.................................................................................. 90
Experiment 8 - Endothermic and Exothermic Reactions ......................................................................................... 99
Experiment 9 - Determination of the Number of Moles in a Chemical Reaction ................................................. 102
Experiment 10 – Reactions of Copper ................................................................................................................... 105
Experiment 11 - Aqueous Reaction Stoichiometry: Analysis of Vinegar by acid-base titration............................ 110
Experiment 12 - The Potentiometric Titration of Hydrogen Peroxide .................................................................. 116
Experiment 13 - The Synthesis and Analysis of Aspirin ......................................................................................... 121
Experiment 14 - Additivity of Heats of Reaction: Hess’s Law ................................................................................ 125
Appendix A – Common Glassware ........................................................................................................................ 131
Appendix B – Using a Buret ................................................................................................................................... 133
Appendix C – Conducting a Titration ..................................................................................................................... 137
Appendix D – Using the Balances .......................................................................................................................... 142
Appendix E – Handling Chemicals.......................................................................................................................... 146
Appendix F – Filtration........................................................................................................................................... 148
Appendix G – Calorimetry ...................................................................................................................................... 154
Appendix H – Using a Centrifuge ........................................................................................................................... 156
Appendix I – Heating ............................................................................................................................................. 157
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Table of Figures
Figure 1: These dartboard targets illustrate the difference between accuracy and precision. .............................. 34
Figure 2: Top Loading Analytical Balance ................................................................................................................ 36
Figure 3 - Reading the meniscus in a graduated cylinder properly ......................................................................... 38
Figure 4: Visible Spectrum of Chlorophyll-a ............................................................................................................ 66
Figure 5: Ions of Some Main-Group Metals (Groups IA - IIIA) ................................................................................ 70
Figure 6: Ions of Some Transition Metals and Post-Transition Metals (Groups IVA and VA) ................................. 70
Figure 7: Ions of Some Nonmetals (Groups IVA - VIIA) ........................................................................................... 71
Figure 8: Formulas and Names of Some Polyatomic Ions ....................................................................................... 71
Figure 9: Typical Buret Setup................................................................................................................................. 112
Figure 10: Reading a buret meniscus with a reading card. ................................................................................... 133
Figure 11: Using Wax Paper to Transfer a Solid .................................................................................................... 146
Figure 12: Proper technique for transferring a liquid. .......................................................................................... 147
Figure 13: Covering a flask for temporary storage using parafilm ........................................................................ 147
Figure 14: Fluting filter paper for gravity filtration. .............................................................................................. 149
Figure 15: Parts and proper use of a Bunsen burner. ........................................................................................... 157
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General Chemistry Lab: Your Mindset
I am making a few serious assumptions about you as you enroll in this course. First, I’m making the
assumption you are an adult who is serious about their academic career. If this is wrong, I’d suggest you either
act like it is or drop the course now. This class has a lot of work that must be done, and much of it is to make
sure you and those around you are safe. This lab is a place to explore the chemistry you are learning in lecture.
To see that all of the equations and “stuff” I tell you really does work.
This manual is written to give you the detailed set of instructions you need to conduct the labs
contained within. It does however, assume you will learn the technique when first exposed to it, and retain that
knowledge for later use. What this means is you will get detailed instructions on how to use a most devises
once…after that, you’ll just be told to use it and I expect you’ll use it correctly. My goal with this manual above
all else, is to force you to prepare for the labs; to think about them and plan your work out. If you take shortcuts
and don’t come to lab with a plan, you will not finish on time…and I will not give you extra time. Assignments
are due when the due date and time are listed. It is your job to get it done, just like you would have to do in the
real world. Mistakes happen, and a safety margin of time is built into each lab, so that will not be an acceptable
excuse.
The best way to approach the lab is to read the experiment before you come to lab. Then, reread the
experiment and establish a plan of approach to the experiment in detail. This is done in your lab notebook, in a
way you feel will allow you to get the work completed on time. Your plan should include weights, volumes and
conversions as well as relevant physical data such as density and melting points. You should also have a list of
needed glassware and equipment for the announced reagents you will have available. Your preparation should
also include developing a data table for the measurements you will collect and results you will calculate.
Working this out in a consistent manner will lead to a more organized notebook, and more efficient progress in
the lab. I WILL BE COLLECTING YOUR PREP WORK AND RESERVE THE RIGHT TO REFUSE YOUR PARTICIPATION
IN LAB IF IT IS INSUFFICIENT.
In General Chemistry labs, you usually learning a new technique each week or at the very least, some
new chemistry to observe. But there are only so many techniques and so much chemistry at this level.
Consequently, as the semester moves forward the labs often become more and more sophisticated. But, you
will be building upon the work you did previously, and hence it shouldn’t be that much harder; only what you
prep for will change. Getting into a habit of good preparation will make this easier. But no matter how much
you prepare, when experiments begin to get sophisticated, “bad” data becomes a possibility. That’s part of the
process, and you should not “freak” out. How you deal with bad data in science is just as important as when it
comes out perfect.
So what am I trying to do here? Scare you into dropping the course? Prove what a jerk I am? No. I am
trying to forewarn you of what lies ahead so you are prepared for the semester and don’t panic when things go
wrong. Changes in the procedures are much more manageable when you UNDERSTAND (Read: did not
memorize or just scribble in your notebook) what is going on. You need to start thinking about what went
wrong and how to fix it. Start with your calculations…redo them without looking at the old set. That’s where
the most common preparation mistakes are made.
If you don’t understand; if you have tried and cannot figure it out; by all means ask me! But do so
before the lab begins, and be ready to prove to me you tried….really hard. If I feel you are fishing, I will let you
know how disappointed I am (read: I will treat you like an employee not doing their job). If you arrive
unprepared, you are not a priority for me. I will get the prepared students started and get to you when and if I
have time.
In this, the age of information it is necessary for me to say a few words about how science collaborates.
If you look up the definition of that word, collaborate, you will find something along the lines of ‘to work with
another person or group in order to achieve something’. The key word there is to “work” with those people.
Copying the work of others and claiming it as your own is not collaboration, but cheating or plagiarism. For the
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latter, you will fail the course and possibly be asked to leave the school or fired if it is a job. For the former, you
will be rewarded and praised for your actions. Science cannot work in a vacuum. It thrives on sharing data and
deriving conclusions independently to further our understanding. This is why we have conferences, peer
reviewed journals and approved databases. The price for this is a small one; you must reference the source of
your new information. Anything you lookup must be referenced to the source. If you have a lab partner for a
portion of the lab, then BOTH of you should keep records of what was done, and both names should appear in
the final report (be it as authors if the paper is co-written or an acknowledgement if only data sharing occurred).
The Student handbook addressed plagiarism in detail…you should go read it before copying anything in lab.
While I’m on it, a word about sources seems in order. The internet has made of all knowledge of the
universe available to you in the palm of your hand. This is great, but much of that knowledge is unregulated and
thrown out without relevant references. For this reason, you should be referencing ONLY primary, scientific
sources. Wikipedia is often a tertiary source, meaning it is a summary of a summary that is summarized. You
should be referencing Peer Reviewed journals primarily, textbooks at a minimum and no tertiary sources.
You did re-read this section twice didn’t you? This essay could be entitled “How to pass this class.”
Chemistry Lab Software
The science lab at EFSC-Titusville is equipped with several pieces of computer equipment to help in your
investigation of chemistry.
All the computers have a default login account. The login for all machines is Student and the password is
student. Most require that you type T01215-##\Student (where the ## is the number found on the front of the
computer you are using)
PLEASE DO NOT TRY TO LOG IN USING YOUR STUDENT ACCOUNT, IT DOESN’T WORK
There are 10 computers in the lab equipped with several pieces of software you might find helpful.
Among these software installations are:
1. MS Office – Word, Excel and Powerpoint.
2. Spartan 2008 v 1.2 – Additional molecular modeling software you might be interested in. Only installed on
the front 6 Computers.
3. LoggerPro – Used to collect data during your experiments.
4. ChemOffice (Front 3 machines only)– ChemDraw and Chem3D, used to help in drawing molecular
structures, model chemicals and examine theoretical properties of chemicals.
5. Gaussian 03W – Only on the win 7 machines. Includes mathematical functions needed for advanced
modeling work in Organic Chemistry.
6. ChemSketch – Is installed on some computers. Similar to ChemOffice but less powerful. This one is also
free if you would like to have a copy. They are at www.acdlabs.com
7. Internet Explorer/Chrome Browsers – Should you need to access CANVAS, Mastering, my web site or
chemical information. It is not for surfing the web or reading your blog pages instead of doing your work.
You can also use it to access chemical information such as Material Safety Data Sheets (MSDS). The
department also keeps a set of MSDS in the lab as required by law.
Several reference books are also available should you need to use them. PLEASE DO NOT REMOVE THEM FROM
THE LAB!
Reference Sections:
• The Merck Index is an excellent reference book for over 10,000 important organic substances. It has a handy
cross index and molecular formula index that you will find useful.
• The CRC Handbook is another reference book that provides some physical and spectral information on a
wealth of substances. Overall, however, I find the Merck Index to be both easier to use and more relevant.
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Several textbooks (out of date and not worth any money) are also on the shelf…the book is out of date but the
chemistry hasn’t changed.
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Important Instructions for Lab Success
1. Read the assignment before coming to lab.
2. Plan your experiment BEFORE you come to lab. Often simply reading through and thinking about/visualizing
what you will be doing will help you work out the difficult portions of an experiment. This has the added
benefit of helping you think of questions you can ask during lecture instead of during the experiment when
time is more critical.
3. There are picture guides on ANGEL and in the laboratory. If you do not know what something is, consult the
picture guides.
4. Read the technique guides in the lab manual introduction or appendices prior to conducting the lab.
5. Complete any prelab assignments WELL before arriving at the lab.
6. Setup any math you may have to conduct during the lab. This will save you time during lab.
7. Work in the pairs unless instructed otherwise.
8. Think about what you are about to do before you do it. Consciously make an effort to not cause accidents and
mistakes.
9. Record data DIRECTLY onto your data sheets or into your lab notebook. NEVER COPY FROM “SCRAP” PIECES
OF PAPER.
10. Keep your lab bench clean and uncluttered. Wash dishes during downtime.
11. Do not borrow equipment from other drawers. If you are missing equipment, ask for a replacement.
12. No matter how cautious you are, there is no guarantee your lab mates will be. Be aware of your surroundings!
If someone around you is doing something unsafe consult your instructor.
13. Science thrives on collaboration…not copying. Discussing things with the instructor or your lab partners may
provide useful help, but you should not copy from them. You learn nothing from this behavior.
14. Complete the post-lab assignment neatly and place it in the return box before leaving.
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Important Lab Protocols/Techniques
It is REQUIRED and expected you will have these guidelines memorized prior to conducting
experiments in this lab.
1. Always pour acids into water, never in the reverse. When acids dissolve they release large amounts of heat
and may boil and spatter.
2. DO NOT PUT CHEMICALS BACK INTO A REAGENT BOTTLE!!
3. If you must insert glass into a hole in a stopper use glycerol to lubricate it. Wrap the rod in a thick cloth as
well to reduce the risk of stab wounds from breaking glass. Twist the rod or tube into the stopper, never
push. Keep your hands close together to reduce leverage force.
4. Never point a test tube at someone. If they flash boil, they will often shoot chemicals out like a cannon.
5. When dealing with toxic fumes, explosive chemicals or flammables, the experiment should be conducted in
the hood.
6. Never use wet glassware in an organic lab. If you need to wash your glassware (and you should) use soap
and water, dry it on the outside with a paper towel (NOT INSIDE!) and then use acetone or ethanol to
remove the water. Finish it with a heat gun (or hair dryer) to evaporate the solvent. Make sure the
glassware is not hot when you go to use it.
7. DO NOT USE OPEN FLAMES AROUND VOLATILE SOLVENTS! If you are not certain if it is volatile…assume it is!
8. Do not place a reagent stopper on the desk, keep it in your hand (chemical side away from your palm).
9. Put reagents and equipment back from where you found them.
10. Try to take only the chemicals you need, avoid excess.
11. Use Distilled (or Deionized) water unless specifically told otherwise.
12. Always use wax paper or a weighing boat when weighing chemicals on the balance.
13. When weighing objects they must be at room temperature.
14. Never place hot glassware or objects directly on the bench. Use a wire gauze or heat pad.
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The Laboratory Notebook
Your lab notebook should be a permanent record of your experimental work. Because of this, it should
be perfect bound (not spiral or ring bound). If it is a single page notebook, meaning no carbon copies are made
you should make all entries in pen at the time the experiment is conducted. If it is a carbon style notebook, then
pencil is ok. However, most industrial and research labs prefer all work be done in pen. I prefer pencil myself
because it is resistant to spills. For this lab, you may use either. Just remember you are using a carbon copy
notebook, erasing is not an option. Readability however, is. If I cannot read it easily, then it is wrong. Neatness
is beneficial, but I am not expecting perfection. After all, excessive neatness is a sign the work was not written
during the lab. You wouldn’t do that would you?
It is common practice for you to use the right hand page to record data, and the left hand side (the back
of the previous carbon page) to make notes, calculations and rough readings of data. This is recommended, but
not required. However, any evidence of recording data on scratch paper, paper towels, the back of your hand,
etc…anywhere not directly into your lab notebook will result in an automatic zero for the experiment.
Errors happen. They are frustrating and common unfortunately. You may misread a balance, an
instrument might not be working correctly, or a reagent may be contaminated. If your results have more than
the anticipated variation, you’d be tempted to toss everything out and rip the pages from your manual. This
would be BAD! Not only is “bad” data often found to be “good” data later on, but all data, be it invalid or
perfect is part of the laboratory record. If you need to make a correction, draw a single line through the entry
and enter the correct value above the original. If the reason(s) for the change is not obvious, make a note next
to the entry to explain it. Any pages missing (i.e. removed) from a lab notebook, any data blocked out
completely is a violation of scientific integrity and will result in a zero for the lab it occurs in. Go on, ask me
what happens if it is between two labs. In science, lab notebooks are sacred. They are used to determine
patent rights, scientific priority and the validity of results.
When you are done with your experiment, all relevant data and results1 should be organized into a data
table for all trials and the statistical results. Tabular data fields are useful for detecting trends and patterns in
the data, especially if intermediate values form calculations are included. It also makes it easier to spot errors.
Calculations are one of the greatest source of error in the post analysis of an experiment. Usually, it is the
incorrect treatment of data that leads to the error. Because of this, calculations are an important part of the
determination of a result as the data itself. For this reason, one example of each calculation should be
included in the notebook with all units and numbers included.
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Data is a measurement taken directly. A result is a value calculated from data points. Know the difference!
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The most important criterion for an acceptable lab notebook is that the record be complete enough for
a second person to be able to follow your experimental work, thus being able to repeat the experiment and to
obtain the same results. The record must be written while you do your work. Your records should reflect your
work so well that any odd results can be interpreted later, even if you don't notice them at the time. Later, you
can do calculations, etc. If you find yourself unable to figure out where a mistake or error has occurred, you are
likely NOT keeping an effective notebook.
How to Keep a Laboratory Notebook:
As the term progresses, you will learn more and more of the details concerning how to keep a professional
style notebook. This includes using clear handwriting in your notebook; illegible labs will not be graded!
Lab notebooks should be written in third person…if you don’t know what this means seek help from a
Communications professor.
The lab notebook is not hard to keep properly. You just have to break all those bad habits we humans
acquire over the early years of education. Here’s a list of the things you should be doing:
1. Prepare ahead of time and PLAN.
2. Don’t try to cram it all on one page.
3. Keep your lab notebook at one end of the bench, and wet work at the other. Think about which
hand you write with when working this out. Also, your lab notebook should be the only paperwork
you have in your work area. Anyone with a lab manual will be asked to leave the lab and receive a
zero for the experiment.
4. Keep a notebook closed when it is not in use. Data can be transferred in the case of a disaster, but
not if the pages are destroyed.
5. Use the first page for a table of contents or to cross reference data.
6. It is ok to do calculations on scrap paper first before transferring them permanently into the lab
notebook. Just not data!
7. Put all data in tables. Even if instrument printouts are attached, summarize the result in the tables.
8. Have someone sign your data page during lab to prove you were there and completed the work.
Keep in mind signors, you are saying they did the work…if they didn’t, that is scientific dishonesty
and will be dealt with accordingly.
Tables are probably of utmost importance to developing a better lab practice and efficiency. A well
prepared student can prepare a table beforehand that saves time in the lab and makes calculations much
simpler to perform.
So what should you expect upon turning in your work? Here you go, a data table of grading!
Prelab Preparation
Labs will be graded out of 100 points.
Lab Section:
First page of experiment: Prelab
-Date, Name and Title on every page
Objective/purpose (usually 1-2 sentences)
Pre-Lab info, Data and Calculations
-A summary of the important theory
-balanced reaction equation
-formula weights of all reagents and products
-amounts of reagents to be used
-incomplete references
Experimental procedure
-written so as to show understanding of what is to
be done. Need not and should not be written
word for word.
Points off
-5 points per missing page or incomplete section
- 5 point (too long, wrong or unclear)
- 5 points each item missing
- zero for lab prep for that experiment if the procedure
could not be given to another student and be done with
only your lab notebook. Remember, you can be asked to
leave if you are unprepared.
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Postlab Report
Data and Observations
-30 points for Unknown Label not given
-30 points Data missing
-10 points Data not in Table form
-10 points Observations missing
Waste Disposal
-10 points for no information in report regarding disposal
of waste products
General Comments
-Illegible reports will not be graded
-10 points for unsigned data pages
-10 points for missing units
-10 points for significant figure errors
-10 points for blotted out non-data errors (remember data
blocked out is a zero for lab)
-in addition to the report grade you will receive a grade for
how accurate/precise your result is. For example, if a good
yield is 75% and you get a 45% yield…points will be lost
(generally 1 point per % error over 5%)
Result
Final Result
You will be submitting the above information for each experiment. They are always due the first lab meeting of
the week after the lab is completed.
There are no late lab reports. There are no late lab reports. I will drop two prelabs and two postlabs for
free no questions asked. However, you cannot drop a prelab for a lab you intend to complete. No prelab
preparation turned in and you will be asked to leave the lab.
Here’s a Little More Detail
1. You will need a carbon-copy style lab notebook.
2. You should organize your lab notebook in the following way:
a. Put your NAME and course number on the cover.
b. Page 1 is the table of contents of your lab notebook; fill in the date and title as you go through the
semester. Make sure to add the page numbers as you go.
c. All entries should be made in Pencil. Pencil will not run if it gets wet unlike ink. However, do not erase
errors, do NOT use white-out on errors, do NOT scribble out errors, do not make your lab notebook look
like a coloring book from a 3 year old. Simply cross off the mistake with a single line. Besides, you are
using carbon paper…it doesn’t erase.
i. Example: Place 3 grams of NaCl in a beaker and dissolve in 200 250 ml of water.
d. Record all data in your notebook at the time that it was performed. Take your lab notebook with you to
the balance room. Do not record information on paper towels (they get thrown away or spilled upon) or
scrap pieces of paper (that get thrown away or spilled upon). Do not rely on your memory. Simple
transposition of numbers can throw off calculations and the error will propagate itself throughout the lab.
e. Provide as much detail as you can. What color are the reagents? What crystal type are the solids. Baby
powder and salt are both white solids, but they look nothing alike – describe describe describe!!!! What
happens when you mix reagents together? Does the product form immediately? Where any gases
emitted? What does the product look like?? Be thorough. It is not uncommon for a seemingly trivial
observation to become important later in the experiment.
f. Organize your data in a straight-forward manner. Make sure that all sections in your lab notebook are
clearly labeled. Use tables where applicable to better present your results. Draw schematics or flow
diagrams to simplify a long procedure. Be neat!!!! This requires planning planning planning and more
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planning BEFORE you come to lab. The shell of your notebook should be filled out before walking in the
door. Leave enough space for observations and calculations that you might have to perform during the
lab.
Your lab notebook will constructed using two main submissions for each lab per week.
1. Due Before Class Starts, Prelab Entries: Part of your grade for the lab portion of this class will be based
upon your pre-lab preparation. Depending on the experiment, this may or may not include a separate
online or in class pre-lab quiz. Regardless, you will be expected to come to lab with your notebook
prepared for the day’s investigation. After the Table of Contents, you can sequentially start keeping
information from each lab. See below for the recommended format for keeping information in lab. The
organization of the pages for a particular lab should contain the following:
a. Title block – The Full experiment title and Lab number, Your Name, the date, page numbers of the
experiment and course number MUST be on every page of the notebook. (-2 points per block not
filled out)
b. Objective Statement – A restatement of what CHEMICAL PRINCIPLE is being demonstrated.
c. Theory Section – You should restate the theory for each lab as well as add additional information from
your textbook. A chemical reaction should be included as well as a description of the theoretical
concepts addressed in the lab. (-10 if missing or under developed)
d. Equipment – What materials are needed to conduct the lab? You should draw pictures of apparatuses
and look up those items you do not know. (may be part of the procedure)
e. Procedure – This is a full step by step re-writing of the procedure from the lab manual. It is designed
for you to think about what you will be doing in the lab each week. It should be detailed enough that
you can conduct the lab WITHOUT THE LAB MANUAL, as that is exactly what you will be required to
do. Do not include data tables in your prelab write-up as you need them in the POST LAB. (you will
receive a zero if this is not complete and detailed)
f. Calculations – The calculations should be laid out prior to the start of lab. This will include any
calculations you need to conduct as well as calculation of the theoretical yield (-10 if missing)
g. Prelab Questions. Worked out with work shown. (-3 for each question not answered)
CUTTING AND PASTING FROM THE LAB MANUAL IS NOT ACCEPTABLE AND WILL RESULT IN A FAILING GRADE
FOR THE LAB.
2.
Due ONE week after the lab is completed (before class begins):
a. Observations – What did you see happen? What color, smell, shape, texture, etc did the chemicals
have? Did you notice bubbling, color changes, etc. This is not a section of the manual that has
numbers and data. (-5 points if missing or incomplete)
b. Data/Results – Data is any number you take from a measurement device (i.e. 12 cm, 14.5 g, 30oC). A
result is a number you must mathematically manipulate to achieve. For the most part data and results
may be combined, and a data table format is provided in the lab manual. This must be copied into the
lab notebook each week. (you will receive a Zero for the lab if this is missing)
c. Calculations - A section devoted to the mathematical manipulations you conduct. There should be a
section in the post lab where you complete the math based on your data. Including a % yield or error
calculation ALWAYS! (-10 points if missing)
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d. Analysis/Interpretation/Questions - Answer any questions posed by the lab in the data analysis
section. (-5 points for each missing)
e. Conclusion – Summarize the results of your experiment, and summarize the general chemical theory
addressed by the lab.
i. In this section you should describe any conclusions that you have found based on the analysis of
your data.
ii. What did you learn from this experiment, what went wrong, how could you have fixed it?
iii. Summarize the reasons for your error in complete sentences! Human error is unacceptable…it can
be prevented. “good” errors are those you can’t anticipate or adjust for (i.e. instrument error,
weather conditions, etc.)
iv. You and your lab partner will have the same results, and the same reasons for error etc. . . You did
the lab together so that is to be expected. You and your lab partner WILL NOT copy your reasons
from one another verbatim. You will come up with your own way of expressing the data,
interpretation, and conclusion.
v. A common and HORRENDOUSLY unacceptable conclusion goes something as follows: “Overall our
experiment seemed to go just fine.” You will lose points for this. No thinking here, this is about
summing up based on facts. A better conclusion would be, “The yield of 78% indicated a successful
precipitation of silver chloride from silver nitrate.”
Additional notebook requirements:

NO LOOSE PAGES. All lost pages or data are the students fault. If you give me your notebook
and I lose it that is my responsibility…if something falls out of that note book and is lost, that is
YOUR responsibility.
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Chemistry Lab Notebook and Participation Grading Rubric
Students will receive a successful grade on a lab for attempting the lab in earnest. This means coming to
class, trying the experiment and making an effort at success. Those wishing for more than just a passing grade
will need to make sure a few criteria are met:
Lab completion checklist
Prelab (Must be completed on carbon pages for credit):
1.
ONLY CARBON COPIES TURNED IN, Originals will not be accepted for a grade)
2.
Title block completed. (-2 points/box not filled on EACH page, -5/page if all boxes not filled in)
a. Your name
b. Complete title of lab
c. Experiment number
d. Date experiment conducted
e. Course information (e.g. CHML 1045) and section (e.g. 01T).
3. Theory section based on the lab introduction sections and/or appropriate textbook sections (-10 if not
included or insufficient detail given).
4. All chemical reactions shown and/or chemicals used listed out (-5 if not included).
5. Procedure written out; detailed enough that the experiment is done (and will be) without your lab manual
(automatically given a Zero on the lab if not included or insufficiently detailed).
6. All calculations laid out showing you have thought about the math and studied it before coming to class (10 if not included or insufficiently detailed).
7. Prelab questions written out and completed showing all work and/or complete sentences used. (-3 per
question not completed).
Postlab (must include carbon copies from lab experiment. Failure to do so once will result in a Zero:
1.
Cover page from lab manual and filled out completely (-10 if not included or completed).
2.
Title block completed. (-2 points/box not filled on EACH page, -5/page if all boxes not filled in)
a. Your name
b. Complete title of lab
c. Experiment number
d. Date experiment conducted
e. Course information (e.g. CHML 1045) and section (e.g. 01T).
3. Data table clearly shown with all data filled in; units included for each measurement (automatic Zero if not
completed or filled out).
4. Observations written out in detail (-5 if not included or detailed).
a. Step information included
b. Information from all senses included
c. Information about chemicals and your results …not what the graph or computer is doing.
5. Calculations written out in detail (-10 if not included or detailed).
6. Conclusion (-5 if not included or sufficient)
a. Includes basic statement of the chemistry included in the experiment
b. Includes a summary of your results.
c. Includes an error discussion (i.e. what went wrong and how could you fix it)
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illegible labs will not be graded!
15
Laboratory Techniques
Lab techniques come with two criteria:
1. Does it work?
2. Is there and easier/more efficient way to get the same answer?
Science likes appropriate techniques. Adequate is rarely acceptable…would you eat mashed potatoes with a
steak knife?
These next few sections along with your text book will help you develop good lab techniques and avoid the
potato embarrassment.
Reagents
In most cases I will give you the manufacturer supplied bottle for all reagents. Check the labels, an
analysis is usually done to assay the purity and grade of the reagent. You should record this assay information
in your lab notebook. It is mandatory information in technical papers (and comprehensive reports), so get into
the habit of recording it.
Your results are dependent upon this purity measurement, and you should make all efforts to preserve
the purity of the reagent. Under no circumstances should you ever return material to a stock bottle. A large
metal scoopula or porcelain spatula can be used to break up caked solids. Keep in mind what the reagent is
used for…if it is metal analysis, you don’t want to add stainless steel to the bottle.
Droppers and pipettes should never be dipped into a bottle. Pour it out into a small beaker first.
Additionally, droppers from dropper bottles should never come into contact with any surface outside of the
bottle itself.
Drying at Elevated temperatures
We have several drying ovens used by several courses. They are generally kept at 125oC or 250 oC. How
fast they dry substances out depends on the water content of the atmosphere in the oven. Let’s face it, we live
in Florida and the humidity is insane usually. Because of this, drying is hard to do. Add to the process, the
addition of a wet beaker or hydrated sample can quickly make everything in a drying over “wet”. Because of
this, you should never put wet glassware in a drying oven when analytical samples are already in the oven.
Before using an oven check the temperature and contents of the oven. Never change the settings!
Oven space is at a premium most days. To aid in this, place objects along the back and sides whenever
possible. Use tongs to place and remove items. Sample bottles and crucibles should be dried in small beakers
with ribbed watch glasses over the beaker. Your name and lab should be on everything in the oven; any
unlabeled items will be removed. This will also help solve one of the other big problems with drying
ovens…sample loss. Community ovens are often prone to someone taking your sample or knocking over
someone else’s sample. Labelling your samples clearly will solve the first problem quickly. By placing your
bottle, crucible, etc in a beaker, you can add a piece of paper with your name and lab information on it. The
latter problem is solved by using tongs and being careful when moving items in the oven.
One last thought on ovens. As discussed before, a plan is necessary. Standing at the oven with the door
open lowers the temperature and lets water creep back in. Go to the oven knowing what you will do once
there, and have all of the tools you’ll need to do it. Constantly opening and closing the door or standing with the
door open ruins the lab experience for everyone!
In the late-90’s, I was an overwhelmed graduate student teaching microbiology. I got behind on lab
prep one day and cut a corner by rapidly heating test tubes in a microwave. My thought was the higher
temperature would drive the water out of the glass faster than the drying oven. I was right (though not the first
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to think of this path), and the microwave became a much bigger part of the microbiology lab2. These days, the
microwave oven is becoming a much bigger part of labs everywhere just as they did in kitchens. You can dry
glassware and samples in about 1/4th the time you can in the drying oven. However, some things don’t take to
the microwave very well. Some samples will melt, burn, or decompose when in the microwave for long
periods of time. Check with your instructor before drying a sample in the microwave. You can also test a
small quantity of your sample before entrusting the entire sample.
WARNING! OBJECTS COMING OUT OF THE OVENS AND MICROWAVES ARE HOT!!!! USE TONGS AND GLOVES
WHEN HANDLING THESE OBJECTS.
Once your samples and glassware come out of the oven, they will usually need to cool to room
temperature before use. This is accomplished by placing the items in a desiccator while they cool.
Cleaning Glassware
There are several levels of cleaning when it comes to glassware:
1. Dirty glassware needing general cleaning. This couldn’t be simpler. Rise with tap water several times,
use soap available in the lab in squeeze bottles and a scrub brush available by the sinks if necessary.
Rinse a few times with tap water, and several times with distilled water. This will then be clean enough
to use for most aqueous based experiments. Do not use paper towels to dry the inside of the glassware
out. It leaves lint which is dirty in our world. If other chemicals are to be used besides water, rinse a
few times with the chemical to be used.
2. Dirty glassware needing chemical cleaning. If basic cleaning doesn’t remove the chemicals, the next
step is to try cleaning with acetone or ethanol. This is usually enough to remove fats, oils and organic
chemicals that soap leaves behind.
3. Glassware for use in non-aqueous organic reactions. Many reactions in organic chemistry require the
presence of water be minimized. To do this, we wash as if it needs general cleaning, then rinse the
inside with acetone. Once this is done, a hot gun (e.g. hair dryer) is used to dry the outside and heat the
glass through conduction to evaporate the acetone on the inside. Never use the hair dryer or any
airstream (i.e. compressed air) to dry glassware. This can introduce oil from the compressor or other
contaminants.
4. Glassware for use with HPLC. HPLC is very susceptible to hard water salts as it can clog the column. For
this reason, glassware is thoroughly cleaned with soap, water, rinsed with tap water, and then DI water
many times. Last the glassware is rinsed inside several times with 95% ethanol and allowed to dry. If
the glassware has spots, it must be rewashed.
5. Last ditch nothing else worked. If the above doesn’t work, the last two options are horribly toxic and
relatively unsafe. You can use a solution of sulfuric acid/chromic acid, commonly known as “cleaning
solution”, to rinse the inside out and clean some of the toughest dirt. This is a very toxic mess, and very
bad for you. If needed, see your instructor for assistance in this method. You should wear thick rubber
gloves and minimize contact with the solution. Another less toxic, but equally nasty solution is sodium
persulfate in sulfuric acid. This is a good general solution that is almost as effective as the cleaning
solution. Otherwise, it is used the same way the cleaning solution is.
6. Cleaning volumetric Glassware. Volumetric glassware is clean when water drains from the surface
without leaving droplets. If any oil is on the surface, droplets will form. Once it dries, it will usually also
not form droplets and should be re-washed. If you must clean and store volumetric glassware, it should
2
This story has an unhappy ending however. Though the glass was dry, it was not sterile. Every sample we
made with those test tubes was contaminated. I had to do it all over again, losing a weekend in the process.
Cutting corners doesn’t usually work out in your favor!
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be filled with deionized water or other nonabrasive solvent. Burets can be cleaned with soap, water and
a buret brush. Rinse with distilled water when done. You should clean the buret in 10 cm sections.
Pipets should be cleaned with soap and water, then rinsed with distilled. Do not drain the soap through
the tip as it often has undissolved salts that can clog it.
One thing to be certain of, is to clean glassware the way it needs to be cleaned. Don’t do more than you have
to; it is time consuming and can be expensive as ethanol and acetone are not cheap. Ask yourself a simple
question before you begin: “Is there another approach?” If you are about to add water to the glassware, why
does it need to be dry!?
Cleaning Bench Tops
Use paper towels to clean up most simple spills. If it is a concentrated acid or base, neutralize the spill
first. Use a wet paper towel repeatedly to clean the surface of water soluble compounds. Keep your bench top
clean and dry at all times!
Quantitative Transfer
Quantitative transfer is the complete transfer of a sample without loss of any kind. The techniques are
usually a simple matter of common sense – don’t spill, splash, dribble or drool. Dry solids should be poured or
transferred by spatula. If the surface tension of a liquid is very high, it should be transferred by pouring down a
glass rod. This prevents the liquid from running down the side of the container and reduces splashing. Last
traces are transferred by washing the original container and transfer equipment such as a spatula, glass rod,
funnel, with a miscible liquid. Constant stream methods using a squeeze bottle or batch rinsing with small
volumes are both acceptable methods.
Transferring material from a weighing bottle should always be done by pouring and not spatula.
Reading Calibrated Scales
The digital revolution has reduced the opportunity to read an analog or Vernier scale. It is a very
necessary skill however as so much of the glassware, such as burets and pipets, you’ll encounter possesses
calibration marks. Familiarize yourself with the various scales; do they read top to bottom or reversed, left to
right or right to left, etc. Determine the significance in both units and magnitude of the largest divisions as well
as the smallest divisions. With instruments, this can change over various ranges of the readout, look carefully
for these.
With a fixed readout such as an instrument dial, what matters is consistency of your eye location above
the scale. It is best to have the eye at the same level as the pointer. In any case, having your eye perpendicular
to the surface of the scale is of the utmost importance.
If the surface to be read is a liquid, a reproducible point must be selected. In most cases this is lowest
point of the meniscus (curved down) or top point (curved up). Most calibrated glassware will have a line that
circumscribes at least half the circumference of the device. You can check for level by adjusting the angle of the
glassware until this circle becomes a line. At this point the glassware is level. It is best that you bring the
glassware up to eye level. Do not kneel or crouch down to bench level. Hold the glassware up to your eye level,
level it and read it. Reading it is best accomplished using a white card with a straight dark line on it. The dark
line is raised up to the point of the meniscus and used to read the calibration marks. This technique also solves
another problem of reading a meniscus, that of the angle the light hits it.
Correctly Reading the Divisions
Readings are generally made to one tenth of the smallest calibration division. For example, if the
readings of a scale are calibrated (the smallest division lines) to 0.1 cm, you would estimate the reading to 0.01
cm. If a reading is right on the line, then you must keep the size of the line in mind while estimating the reading.
Most calibration lines are 0.02 cm wide. Therefore, the top of a line is one reading, and the lower half is
another. For example, if the top of a line is 7.11 cm, the middle would be 7.10 cm and the bottom of the line is
18
7.09 cm. How you take your measurement is important. Reporting that a reading is 7.2 indicates it is close to
7.2, and further away from 7.1 or 7.3. Reporting 7.10 indicates it is closer to 7.10 than 7.09 or 7.11. Compare
the percent difference in those readings:
%D 
0.1
x100  1.4%
7.1
%D 
0.01
x100  0.14%
7.10
That is quite a difference in error! If you report 7.1 when 7.10 was possible, that is stupidity! It is unscientific
and unacceptable.
Weighing out Samples
The preferred method of weighing is the by difference method. You weigh the wax paper or weighing
boat, transfer the sample to the device and weigh it again. The difference in mass is how much sample is
added..
Samples can be added to a clean dry container, often a weighing boat which has been previously tared
on the balance. You must take extreme caution to ensure no material is lost however. Normally, solvent is used
to wash any residue from the boat into the container after transfer.
Rules for using an Analytical Balance
Balances are dreadfully expensive and fragile devices. Because of this, there are several rules that must be
followed when using them.
1. No reagent should be added or subtracted from a container while in the balance. Remove the
container to the bench, make the addition or subtraction and return it to the balance. If you are found
violating this rule, you will receive a zero on the lab you commit this most heinous violation.
Heating Solutions
Aqueous solutions may be heated on a Bunsen burner with wire gauze or a hot plate. The Bunsen
burner gets hotter and heats quicker than the hot plate and will provide a more consistent level of heating
(prevents bumping). Solutions other than water or aqueous salts should be heated in the hood. You should
always place a ribbed watch glass on the flask or beaker when heating it to both prevent your sample from
finding its way to your neighbor’s sample and to prevent the building (paint, ceiling times, etc.) from falling into
your sample. Ordinary watch glasses prevent evaporation, and therefore should have a glass hook on the rim of
the beaker to facilitate evaporative loss. Remember to rinse any sample that splashes onto the watch glass back
into the beaker. Boiling is generally avoided as it raises the risk of sample loss. If you must boil, it is also
advisable to use boiling stones when possible. This too however increases the chance of introducing a
contaminant. Glass beads, marble chips, and silicon carbide are commonly used materials for boiling chips.
Stock Solutions
Even uniformly mixed solutions may form concentration gradients upon standing for long periods of
time. Evaporation occurs at the surface for instance. For this reason they should be mixed before being used.
Volumetric glassware
Flasks, burets and Pipets
Volumetric measurements are temperature dependent. Because of this, most glassware indicates what
temperature the equipment is calibrated to; usually this is 20oC. Luckily for us, the coefficient of expansion for
glassware is so small that even at temperatures up to 30oC the error is less than 0.3 ppt. Larger volumes are
actually easier to measure with accuracy. I know that sounds backwards, but take the following examples: a
50.00 mL volumetric flask has a tolerance of 0.05mL or 1 ppt. A 10.00 mL pipet has a tolerance of 0.02 mL or 2
ppt. As microscale chemistry begins to pervade research, an entirely new set of equipment has been developed
to address this problem.
Graduated cylinders
19
Graduated cylinders should not be used for any quantitative measurement in this course. They are
useful for making solutions where concentration is not required to be highly accurate or does not contribute to
the overall quantitative calculations. Graduated cylinders are very crude pieces of equipment whose glass is
known to vary greatly especially at the bottom and above the top calibration mark.
Use of Volumetric Flask for Solution Preparation
The flask should be cleaned and checked for uniform drainage in the neck in particular using a small
sample of the solvent to be used. Use a funnel to facilitate a complete transfer of solids and liquids. It is a good
idea to use a long stem funnel to prevent air from being trapped in the neck and being forced back through the
top. Use solvent to wash any stuck solids into the bulb of the flask. Fill the flask until it is 80% full, and mix by
swirling (not inversion!) until the solid is dissolved. Do not hold the flask by its neck, this can cause breakage.
Mixing the solution in this manner allows for the volume change that occurs when solids are placed into liquids.
You may place difficult to dissolve solids in an ultrasonic bath. You should never heat a volumetric flask on a
Bunsen burner or hot plate. This can damage the flask and change its volume. Add solvent to the flask to bring
the solution to the calibration mark. Use a dropper to add the last few drops. Stopper the flask and invert it a
minimum of 15 times or more if the solution is highly concentrated or viscous.
The solution is prepared at room temperature usually. At this temperature, the glassware has a
coefficient of expansion of about 0.024%. +/-4oC this results in a 1 ppt or less error.
Burets
The calibration marks on our 50 mL burets are at 0.1 mL intervals. To obtain maximum precision,
volumes are estimated at 0.01 mL. This leads to a standard deviation of 0.02 mL. A volume change reading
would require a minimum of two readings which introduces an uncertainty of [(0.02)2+(0.02)2 ]1/2 = 0.03 mL or
greater than 1 ppt. In a 20 mL delivery, this would yield uncertainty greater than 1 ppt which is not acceptable
for our work. However, if we increase the delivery to 40 mL, the uncertainty would decrease to acceptable
levels. We therefore usually plan our experiments to have endpoints in this range.
Use of a Buret in Titrations
Most burets have a Teflon stopcock to control flow. This stopcock does not require lubricant, and is
tightened using a nut until leaks are prevented but the stopcock turns freely. You should store the Teflon
stopcock separate from the buret to prevent warping; this includes those found on any glassware such as
separatory funnels. When reassembling the stopcock, insert the stopcock, then add the white plate against the
glass, the black O-ring and finally tighten the nut in place. You should always test for leakage before starting
your titration. Permanent stopcock bases are all designed to be left-handed to accommodate right handed
people. This leaves the predominant hand free to agitate, add or otherwise manipulate the experiment. With
Teflon stopcock mounts, this is not a problem and you can adjust until you are comfortable.
To fill a 50 mL buret, rinse the buret 3-4 times with small volumes of the solution you will use. This is a
process called conditioning. You should use a buret funnel to ensure all surfaces are covered, and additionally
roll the barrel when pouring out the final 2 washings. Make sure to clear the stopcock and tip with solution as
well. You should fill the buret while it is in the buret clamp to prevent pouring solution on your hands. Fill the
buret to over the top calibration mark and remove the funnel. Never leave a funnel in the buret while titrating
as this can lead to drops from the funnel adding volume to the titrant. You can place the funnel in the neck of
the solvent container to prevent contamination. Cover the funnel with a watch glass to prevent evaporation.
To begin the titration, drain the meniscus to a level below the 0.00 mL level. You will be chastised for
and have points removed from every lab containing this reading. Research has shown it actually increases errors
in titrations. Remove drops from the tip that fail to fall with a small amount of solvent from a squeeze bottle or
by touching the tip to the side of the receiving flask.
Add liquid to the flask at a rapid but uniform rate with continuous stirring. Some labs require manual
stirring to prevent splashing, other are fine with magnetic stirring. Stop the flow intermittently to assess the
20
rate of indicator color change. You can add one drop at these stoppages to assess that rate. Usually, as the
volume increases, the rate of addition will decrease until you reach the end point and a drop or fraction thereof
will indicate the end. Fractional drops are achieved by turning the stopcock slowly until the drop begins to form,
but doesn’t drop off. The partial drop is washed into the flask with solvent.
If you over run an end point, it is over run, be it by a drop or 10 mL. But your time is important, and you
don’t have enough of it to invest in a drop by drop addition approach. You can make the process much easier if
you have enough sample to perform a test titration. In this method you setup a titration, and just add large
volumes and estimate the end point color change.
Many times, you will be using a probe to collect data on the state of the titration. This can be pH, ORP,
or ISE probes. These will allow you to follow a reaction instantly, and very accurately determine the location of
the end point. When using these, it is good to keep track of the points either in an Exceltm spreadsheet or using
the data collection software of the probeware.
Indicators
Indicators still have a significant role in the chemistry lab. Even with modern probes, indicators are
often used as a check to ensure the probe is functioning within acceptable limits. Selection of an indicator is an
art and science in itself. Usually, you need to know a great deal of information regarding the reaction involved
to make this choice. Once you have selected an indicator, you should familiarize yourself with its behavior and
color transitions before employing it. Knowing this information will help you confidently make decisions when
quantitatively using the indicator. This is especially true when a definite color change does not occur, but a
color standard is used to determine the progress of a reaction.
One of the dilemmas of using an indicator is the indicator gives you the “end point” of the reaction. This
is the point when the titration ended (i.e. the indicator color changed to its final color state). In a perfect world,
this would also be the “equivalence point” as well. The equivalence point is the point at which chemically
equivalent quantities of reagents have been brought together. It most reactions, the end point and equivalence
point do not coincide.
21
Laboratory Safety Markings
You should know the properties of the chemicals you are working with. While many chemicals in the
general chemistry lab are harmless, several are not. As you advance in both General and later into Organic
Chemistry the chemicals become more dangerous. Some are toxic and irritating, others caustic or corrosive;
many seem innocuous enough, but cause internal damage on a genetic level not detectable until much later in
life. Understanding the chemicals around you in lab is made easier by a few standard practices implemented
by scientists and workplace hygienists.
NFPA Labels
Each chemical possesses a diamond with four small squares inside it. Each square tells you important
information regarding its use and safety. The top (red) square is for flammability, The left middle (blue)
square reveals Health effects, the right middle (yellow) square tells about reactivity and the bottom (white)
square tells about special handling or storage (often OX for oxidizer, TER for teratogen, MUT for mutagen or
COR for corrosive). The top 3 quadrants are given a number between 0 and 4. A chemical with a 0 is no
hazard, while a 4 denotes extremely hazardous properties. An example of this label is shown below.
MSDS Sheets
The federal government requires that chemical manufacturers provide a series of papers describing
the properties of every chemical sold and used in the workplace. The workplace has post these sheets for all
to consult. These sheets are called Material Safety and Data Sheets, or MSDS sheets. Under the Right-toKnow act, these must be available upon request. To make this easier, these are generally kept in a binder and
left for all to access.
The format can vary by chemical manufacturer, but the content is generally the same. You can find
information regarding the lethal dose (LD50) for an average human (for example 260 mg/kg is the lethal dose
of diethyl ether that will kill a human when ingested)…Organic Chemistry uses a lot of ether, and they share a
lab with General Chemistry! So follow the rules! You can also learn that diethyl ether is flammable, causes
headaches, nausea and must be kept away from flames. The safe disposal of chemicals is important as well.
The MSDS often provides information on this as well. We have reduced our disposal load by conducting more
“green” experiments and using a microscale approach, however there is still waste generated.
22
Lab Safety Definitions
Autoignition
temperature
The temperature at which the substance will spontaneously burst into
flame. Given a supply of oxygen. Such a fire can occur if a lid is
removed from an overheated vessel, as may happen when oil is
heated in a covered frying pan.
Carcinogen
The substance has been found to cause cancer in annuals or humans.
We will not use known carcinogens.
Caustic
The substance is a strong base. Caustic substances cause burns to
human flesh and eat holes in clothing.
Corrosive
The substance is a strong acid Corrosive substances cause burns to
human flesh. They also break down certain fibers (including cellulose,
the polymer found in cotton).
Doseresponse
curve
When plotting close versus physiological effect, the result is normally
not a straight line. Instead, at low closes there is often no effect at all
(or a therapeutic effect), while at sufficiently high doses the effect is
typically toxic. See LD.
Flammable
Inflammable3
Flash point
The temperature at which the vapor above a liquid forms an explosive
mixture with air. Ignition results in a flash (or explosion) rather than a
flame. Many organic liquids (such as gasoline) are used at
temperatures well above the flash point this is relatively safe because
the liquid’s vapor pressure is so high that oxygen is displaced from the
air above the liquid4. And it still takes a spark to start things off. So
don’t smoke around gas pumps!
“Flush down
the sink”
Pour down the drain, rinsing thoroughly with water. Run water after
the material for 5 minutes or more, depending on how much reagent
was discarded and how concentrated it was.
Inflammable
The substance will burn. See also flash point and autoignition
temperature.
3
“Inflammable” is an older term meaning the substance or material can become “inflamed”. A more modern
adaptation is the word “flammable”, which means the same thing.
4
Things like car engines, fuel-air explosives and any internal combustion engine are carefully calibrated to run at
peak efficiency. Too much fuel and the bang is less impressive.
23
Irritant
The substance will irritate skin, eyes or mucous membranes. The
fumes may cause you to sneeze.
LD50
The close that caused death in 50% of a test group (usually of rats).
The dose is expressed in mg/kg, milligrams of substance per kilogram
of test animal. The dose-response curve typically rises sharply at this
point, so that amounts smaller than the LD50 are often innocuous, and
larger amounts are definitely toxic.
Mutagen
The substance causes genetic damage (mutations) to cells with which
it comes in contact. Mutagens in the bloodstream can adversely affect
germ cells, causing genetic defects in offspring. Normally we will not
use known mutagens.
Pyrophoric
Pyrophoric materials ignite spontaneously in air below about 45oC.
They will also ignite nearby inflammable materials.
Smelly
The substance has an objectionable odor. Smelly substances should
be handled only in the hood: this includes cleaning glassware
contaminated with them. If they need to be weighed. They must be
weighed in closed containers.
Teratogen
The substance causes developmental abnormalities in unborn babies:
therefore pregnant women should not be exposed to it. We do use
some suspected teratogens: if you think you might be pregnant you
should inform the instructor immediately.
Toxic
The substance is poisonous. Normally we will not use “highly toxic”
substances LD50 < 50 mg/kg). Mildly toxic” is used to describe
substances with LD50> 500 mg/kg.
Reference Sources
Physical data of organic compounds may be found in the following sources. Most of these are available
in the lab should you wish to consult them.
• The CRC Handbook of Chemistry and Physics
• The Merck Index
• The Aldrich Catalog Handbook of Fine Chemicals
Chemical safety data may be found in the following sources.
• The Merck Index
• MSDS online database which is accessible from the course ANGEL page. These are among the
few reliable online data: they are reprints of published MSDS forms. Information here may be
out of date, so be sure to check the date of the MSDS you are reading!
24
• The BLACK MSDS binder in the Chemistry Lab contains an MSDS for every substance used in
the labs. It is highly recommended that you use these references before coming to lab!
Eastern Florida State College - Chemistry Lab Safety Rules
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1. Safety goggles are mandatory in all labs when chemicals are in use.
This means you need to have them covering your eyes at all times when ANYONE in the room
is using chemicals, not just your group. Goggles are available for use your use in the green
cabinet in the back of the room. The goggle will fit (un)comfortably over your prescription
eyeglasses. Contact lenses are not recommended as chemicals can stick to the lenses causing
damage to your eyes.
We reserve the right to refuse service to students improperly dressed for lab.
No Shirt, No pants, no shoes, no safety glasses…no lab!
Safety: In the event a foreign object or chemical enters your eye, have your lab partner lead
you to the eye wash. Wash your eye in eyewash with flowing water for 10-15 minutes.
Additionally, portable eye wash bottles are available along the wall.
2. Chemistry involves some potentially hazardous chemicals. Therefore, students should
some dressed in long pants and shirts with at least short sleeves.
I know Florida is hot, but long term safety is more important than short term comfort. It is
better to lose a pair of pants than to have a caustic chemical burn your skin. You will be asked
to leave the lab if you come inappropriately dressed, makeup of any missed work will not be
allowed. Additionally, you should try to wear non-synthetic fibers in a chemistry lab (or any
lab that uses chemicals). Many chemical solvents will easily dissolve synthetic fibers, often
creating toxic byproducts or adhesives. NO BULKY OR LOOSE SLEEVES IN LAB! They are
perfect adapted to dipping into chemicals or knocking things off the bench. You can bypass
the shirt rule by wearing a long sleeve, below the knee lab coat. You will still need to wear
long pants however.
Safety: In the event you become covered with a chemical, there is a safety shower in the back
of the lab. You should run water over yourself for at least 10 minutes. In the event of a toxic
chemical spill on your clothes, remove the clothes while under the shower…once again, your
long term safety is more important than modesty.
3. Eating drinking and smoking are not permitted in the lab.
Many chemicals can have toxic results when ingested even in small amounts.
4. You must wear shoes that are closed toed and closed heeled in the lab.
Once again, your safety is at issue here. Dropped glassware and chemicals are a real hazard in
the lab. Open shoes allow glass or chemicals to easily reach your skin.
5. Long hair should be tied back during experiments.
Bunsen burners and beakers of chemicals on the bench can easily cause damage to hair, or
sweeping hair may knock chemicals off the bench.
6. You should immediately clean up spills or notify the instructor of a spill.
Simple spills can often be cleaned up with paper towels. However, toxic chemicals and strong
acid/base spills should be cleaned up with the proper tools. Acid/base cleanup kits are
available on each side of the room. In the event of mercury or any other toxic spill, notify
your instructor immediately.
25
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Initial
7. Chemical data regarding use, hazards and cleanup can be found in the MSDS sheets
located in the front of the room.
The MSDS sheets are there for your knowledge, and can be reviewed at any given time.
Additionally, you can find chemical data online at www.fishersci.com . You should review this
data anytime you are concerned about the effects of a chemical on you.
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Initial
8. Never inhale chemicals.
Your instructor will show you the proper way to smell chemicals should it be necessary. As
many chemicals produce toxic fumes you should not inhale anything without permission from
your instructor.
9. Many chemicals require safe disposal.
You should dispose of chemicals in the proper waste containers. Never pour chemicals down
the drain without express permission of your instructor. Many chemicals pose serious
environmental threats and should not be allowed to enter the sewer system.
10. You will perform no unauthorized experiments.
You should carefully read all labels to ensure you are using the proper chemicals. Never mix
chemicals together in any way other than that described by your lab manual. Failure to do so
may lead to fire, explosion or other health risks. Remember your safety is not the only one to
consider as others are performing experiments near you.
11. Broken glass should be disposed of in the labeled glass containers around the lab.
Never throw broken glass in the trash as it can pose a cut hazard. Additionally, be careful
when cleaning up broken glass. Use the appropriate tools to ensure your safety.
12. Report all injuries to your instructor no matter how insignificant they seem.
Even small cuts may pose health risks when chemicals are involved. Your instructor has safety
and first aid equipment to help with treating these injuries.
13. Never work in the lab alone.
If you are injured, it is possible no one will find you in time to render assistance.
14. Keep drawers closed and Aisles clear or obstructions.
Blacking the walkway is a sure way to make any emergency worse. For this reason all bags
should be left at the desks; take with you only what you need to do the experiment.
Additionally, DO NOT TAKE THE STOOLS WITH YOU! They are major obstructions and trip
hazards. You can stand for a 2 hours; I have faith in your ability.
15. Gloves are available if needed.
Should you need gloves, they are available in the lab. They are useful for protecting you from
corrosive liquids as well as mutagenic compounds. Please only use them when necessary, and
keep them one until you are done. Taking them off and getting a new pair every 2 minutes is
expensive and WASTEFUL! Making balloon animals that look like cow udders even more so. If
you develop a redness or irritation, you may be allergic to latex. Let us know and we will
order latex free gloves for you to use.
16. No excessive Jewelry Should be worn. Do not wear rings, watches, or other types of
jewelry in the laboratory if at all possible. You will be asked to remove what we might
consider excessive amounts of hand and wrist jewelry. Such items can catch on protruding
objects, serve as nice traps for spilled chemicals keeping them in contact with the body,
function as nice “wires” to lead current into the body from electrical equipment resulting in
shocks and/or electrocution. In addition, jewelry can be seriously damaged by chemicals. We
reserve the right to require you to remove any object that we consider to be dangerous.
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26
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17. Visitors are not permitted in lab.
I am responsible for your safety and yours alone. Guests are not permitted. Additionally,
going out in the hallway for a chat is equally as bad. You have either abandoned your
experiment or you have abandoned your lab partner…both cases are not acceptable.
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18. CELL PHONES, LAPTOPS AND TABLETS ARE NOT TO BE USED IN THE LAB.
First and foremost they are a distraction. You are not paying attention to your experiment, or
your surroundings or me. Any of which can lead to safety hazards. Additionally, some phones
have been known to build up static charge leading to sparking. As flammables are often
present in the lab this can lead to serious safety hazards.
19. Safety Equipment is available in the lab. You should take note of all the safety
equipment available in the lab and have a basic knowledge of its function. This includes but is
not limited to: All three Emergency Exits, Fire Extinguishers, Blankets, First Aid Kit, Spill
Cleanup kits, Safety Shower, Eye Wash and Safety Goggles.
20. Underwear. Always wear clean underwear to lab. That way, if you have an accident and
have to go to the emergency room, the medical personnel can see that you wore clean
underwear, and will no doubt comment that you are wearing such nice clean underwear. In
addition, your mother will be very proud of you. All her years of nurturing and nagging have
paid off.
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I have fully read and understand the safety rules listed above as well as the safety lecture given by the
instructor. I have familiarized myself with the location of the safety equipment in the laboratory we are
using.
_________________________ ___________________________
Sign your name
Print Name
_________________________
Date
27
28
Eastern Florida State College - Chemistry Lab Safety Rules
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1. Safety goggles are mandatory in all labs when chemicals are in use.
This means you need to have them covering your eyes at all times when ANYONE in the room
is using chemicals, not just your group. Goggles are available for use your use in the green
cabinet in the back of the room. The goggle will fit (un)comfortably over your prescription
eyeglasses. Contact lenses are not recommended as chemicals can stick to the lenses causing
damage to your eyes.
We reserve the right to refuse service to students improperly dressed for lab.
No Shirt, No pants, no shoes, no safety glasses…no lab!
Safety: In the event a foreign object or chemical enters your eye, have your lab partner lead
you to the eye wash. Wash your eye in eyewash with flowing water for 10-15 minutes.
Additionally, portable eye wash bottles are available along the wall.
2. Chemistry involves some potentially hazardous chemicals. Therefore, students should
some dressed in long pants and shirts with at least short sleeves.
I know Florida is hot, but long term safety is more important than short term comfort. It is
better to lose a pair of pants than to have a caustic chemical burn your skin. You will be asked
to leave the lab if you come inappropriately dressed, makeup of any missed work will not be
allowed. Additionally, you should try to wear non-synthetic fibers in a chemistry lab (or any
lab that uses chemicals). Many chemical solvents will easily dissolve synthetic fibers, often
creating toxic byproducts or adhesives. NO BULKY OR LOOSE SLEEVES IN LAB! They are
perfect adapted to dipping into chemicals or knocking things off the bench. You can bypass
the shirt rule by wearing a long sleeve, below the knee lab coat. You will still need to wear
long pants however.
Safety: In the event you become covered with a chemical, there is a safety shower in the back
of the lab. You should run water over yourself for at least 10 minutes. In the event of a toxic
chemical spill on your clothes, remove the clothes while under the shower…once again, your
long term safety is more important than modesty.
3. Eating drinking and smoking are not permitted in the lab.
Many chemicals can have toxic results when ingested even in small amounts.
4. You must wear shoes that are closed toed and closed heeled in the lab.
Once again, your safety is at issue here. Dropped glassware and chemicals are a real hazard in
the lab. Open shoes allow glass or chemicals to easily reach your skin.
5. Long hair should be tied back during experiments.
Bunsen burners and beakers of chemicals on the bench can easily cause damage to hair, or
sweeping hair may knock chemicals off the bench.
6. You should immediately clean up spills or notify the instructor of a spill.
Simple spills can often be cleaned up with paper towels. However, toxic chemicals and strong
acid/base spills should be cleaned up with the proper tools. Acid/base cleanup kits are
available on each side of the room. In the event of mercury or any other toxic spill, notify
your instructor immediately.
7. Chemical data regarding use, hazards and cleanup can be found in the MSDS sheets
located in the front of the room.
The MSDS sheets are there for your knowledge, and can be reviewed at any given time.
Additionally, you can find chemical data online at www.fishersci.com . You should review this
data anytime you are concerned about the effects of a chemical on you.
29
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8. Never inhale chemicals.
Your instructor will show you the proper way to smell chemicals should it be necessary. As
many chemicals produce toxic fumes you should not inhale anything without permission from
your instructor.
9. Many chemicals require safe disposal.
You should dispose of chemicals in the proper waste containers. Never pour chemicals down
the drain without express permission of your instructor. Many chemicals pose serious
environmental threats and should not be allowed to enter the sewer system.
10. You will perform no unauthorized experiments.
You should carefully read all labels to ensure you are using the proper chemicals. Never mix
chemicals together in any way other than that described by your lab manual. Failure to do so
may lead to fire, explosion or other health risks. Remember your safety is not the only one to
consider as others are performing experiments near you.
11. Broken glass should be disposed of in the labeled glass containers around the lab.
Never throw broken glass in the trash as it can pose a cut hazard. Additionally, be careful
when cleaning up broken glass. Use the appropriate tools to ensure your safety.
12. Report all injuries to your instructor no matter how insignificant they seem.
Even small cuts may pose health risks when chemicals are involved. Your instructor has safety
and first aid equipment to help with treating these injuries.
13. Never work in the lab alone.
If you are injured, it is possible no one will find you in time to render assistance.
14. Keep drawers closed and Aisles clear or obstructions.
Blacking the walkway is a sure way to make any emergency worse. For this reason all bags
should be left at the desks; take with you only what you need to do the experiment.
Additionally, DO NOT TAKE THE STOOLS WITH YOU! They are major obstructions and trip
hazards. You can stand for a 2 hours; I have faith in your ability.
15. Gloves are available if needed.
Should you need gloves, they are available in the lab. They are useful for protecting you from
corrosive liquids as well as mutagenic compounds. Please only use them when necessary, and
keep them one until you are done. Taking them off and getting a new pair every 2 minutes is
expensive and WASTEFUL! Making balloon animals that look like cow udders even more so. If
you develop a redness or irritation, you may be allergic to latex. Let us know and we will
order latex free gloves for you to use.
16. No excessive Jewelry Should be worn. Do not wear rings, watches, or other types of
jewelry in the laboratory if at all possible. You will be asked to remove what we might
consider excessive amounts of hand and wrist jewelry. Such items can catch on protruding
objects, serve as nice traps for spilled chemicals keeping them in contact with the body,
function as nice “wires” to lead current into the body from electrical equipment resulting in
shocks and/or electrocution. In addition, jewelry can be seriously damaged by chemicals. We
reserve the right to require you to remove any object that we consider to be dangerous.
17. Visitors are not permitted in lab.
I am responsible for your safety and yours alone. Guests are not permitted. Additionally,
going out in the hallway for a chat is equally as bad. You have either abandoned your
experiment or you have abandoned your lab partner…both cases are not acceptable.
30
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Initial
_____
Initial
_____
Initial
18. CELL PHONES, LAPTOPS AND TABLETS ARE NOT TO BE USED IN THE LAB.
First and foremost they are a distraction. You are not paying attention to your experiment, or
your surroundings or me. Any of which can lead to safety hazards. Additionally, some phones
have been known to build up static charge leading to sparking. As flammables are often
present in the lab this can lead to serious safety hazards.
19. Safety Equipment is available in the lab. You should take note of all the safety
equipment available in the lab and have a basic knowledge of its function. This includes but is
not limited to: All three Emergency Exits, Fire Extinguishers, Blankets, First Aid Kit, Spill
Cleanup kits, Safety Shower, Eye Wash and Safety Goggles.
20. Underwear. Always wear clean underwear to lab. That way, if you have an accident and
have to go to the emergency room, the medical personnel can see that you wore clean
underwear, and will no doubt comment that you are wearing such nice clean underwear. In
addition, your mother will be very proud of you. All her years of nurturing and nagging have
paid off.
I have fully read and understand the safety rules listed above as well as the safety lecture given by the
instructor. I have familiarized myself with the location of the safety equipment in the laboratory we are
using.
_________________________ ___________________________
Sign your name
Print Name
_________________________
Date
31
32
Experiment 1 - Measurement and Common Laboratory Techniques
Objectives
 To practice weighing objects using the top-loader balances
 familiarize yourself with the concept of density
 perform some basic statistical analyses and basic calculations
 data manipulation using Microsoft Excel
 learn to use common laboratory glassware
Relevant Readings
 Kotz et al Chapter 1,
33

Appendix D – Using the Balances.
Note: you need not write up my notes on Excel, as they are just there for informational purposes.
Theory and Introduction
Experimental observations often include measurements of mass, length, volume, temperature, and
time. There are three parts to any measurement:
• its numerical value
• the unit of measurement that denotes the scale
• an estimate of the uncertainty of the measurement.
The numerical value of a laboratory measurement should always be recorded with the proper number
of significant figures. The number of significant figures depends on the instrument or measuring device used and
is equal to the digits definitely known from the scale divisions marked on the instrument plus one estimated or
"doubtful" digit. The last, estimated, digit represents the uncertainty in the measurement and indicates the
precision of the instrument.
Measurements made with rulers and graduated cylinders should always be estimated to one place
beyond the smallest scale division that is marked. If the smallest scale division on a ruler is centimeters,
measurements of length should be estimated to the nearest 0.1 cm. If a ruler is marked in millimeters, readings
are usually estimated to the nearest 0.1 or 0.2 mm. The same reasoning applies to volume measurements made
using a graduated cylinder. A 10-mL graduated cylinder has major scale divisions every 1 mL and minor scale
divisions every 0.1 mL. It is therefore possible to "read" the volume of a liquid in a 10-mL graduated cylinder to,
the nearest 0.02 or 0.05 mL. Three observers might estimate the volume of liquid in the 10-mL graduated
cylinder shown at the right as 8.32, 8.30, or 8.33 mL. These are all valid readings. It would NOT be correct to
record this volume of liquid as simply 8.3 mL. Likewise, a reading of 8.325 mL would be too precise.
Some instruments, such as electronic balances, give a direct reading-there are no obvious or marked
scale divisions. This does NOT mean that there is no uncertainty in an electronic balance measurement; it means
that the estimation has been carried out internally (by electronic means) and the result is being reported
digitally. There is still uncertainty in the last digit. On an electronic centigram balance, for example, the mass of a
rubber stopper might be measured as 5.67 g. If three observers measured the mass of the same rubber stopper,
they might obtain readings of 5.65, 5.67, and 5.68 g. The uncertainty of an electronic balance measurement is
usually one unit in the smallest scale division that is reported-on a centigram balance this would be ± 0.01 g.
Accuracy and precision are two different ways to describe the error associated with measurement.
Accuracy describes how "correct" a measured or calculated value is, that is, how close the measured value is to
an actual or accepted value. The only way to determine the accuracy of an experimental measurement is to
compare it to a "true" value-if one is known! Precision describes the closeness with which several
measurements of the same quantity agree. The precision of a measurement is limited by the uncertainty of the
measuring device. Uncertainty is often represented by the symbol ± ("plus or minus"), followed by an amount.
Thus, if the measured length of an object is 24.72 cm and the estimated uncertainty is 0.05 cm, the length would
be reported as 24.72 ± 0.05 cm.
34
Figure 1: These dartboard targets illustrate the difference between accuracy and precision.
Variations among measured results that do not result from carelessness, mistakes, or incorrect
procedure are called experimental errors. Experimental error is unavoidable. The magnitude and sources of
experimental error should always be considered when evaluating the results of an experiment.
Units
There are two common sets of units used worldwide when taking measurements: the English system
and the Metric system. While the English system is common in the United States the rest of the world has
largely moved on to the metric system more appropriately named the SI system. You are probably very familiar
with the English system and how it is used, having heard of inches, yards, feet, pounds etc. The metric system
however is used by many countries and almost exclusively in science. The metric system consists of units based
on powers of 10 including things like centimeters, meters, kilograms, etc. there are some measurements in the
SI system that are not based on powers of 10, with specific values, such as the mole, the Kelvin and the second.
In this lab today, we will deal mostly with the meter, the kilogram, and the Kelvin. We will directly measure
several of these quantities and will create more complex units using these three simple units.
The first unit we will consider is the kilogram, a unit of mass. In most chemical labs we deal not with
kilograms but with a smaller unit the gram, as it is much more convenient on the scale of our lab. Next we will
look at the meter, a distance measurement which is slightly longer than a yard. The yard is 36 inches long, and
the meter 39.37 inches long. Last we will consider the measurement of temperature. In science we utilize the
Celsius scale almost exclusively despite the official SI unit of Kelvin. The two units have identical sizes however
the Kelvin is the Celsius temperature plus 273.15 K. The Kelvin system is an absolute scale, meaning there is
theoretically no lower value than 0K.
Using these three units, it is possible to create more complex units to measure other physical properties
of a chemical. This is achieved by mathematically manipulating values through multiplication or division creating
complex units. Examples of such manipulation would be mass in kilograms divided by volume in liters to
calculate density and length x width x height in meters for volume in cubic meters.
Statistics
35
If you were to type random numbers in your calculator, and conduct random arithmetic operations you
would end up with a number that might look reasonable as an answer. However if you would repeat this process
randomly several times, each time you would have which seemingly ‘made sense’. In order to determine
whether a number is meaningful we must perform some basic statistics. Do not be afraid. I repeat do not be
afraid… And above all don't panic. Statistics is really just a fancy name for more algebra. And not even the hard
algebra just some adding subtracting dividing and multiplying.
The quality of your data is what statistics will give you. In all of science the quality of one's data is one of
the most important aspects of conducting research. One of the simplest ways to measure the quality of data is
known as standard deviation. Standard deviation of a set of measurements gives you an idea of the variation of
your data from the average (which is better known as the mean or expected value). If the standard deviation is
very low your measurements are closely centered on the mean value, however standard deviation is high your
measurements are spread out over a much larger range of values. For why the latter is bad, see the dartboard
above in Figure 1 labeled not accurate nor precise.
Standard deviation for a set of measurements is often represented by the Greek letter sigma (). It is
calculated manually using the following formula,

d
2
1
 d 22  d32  ...  d n2 
n 1
2
where each d represents the square of the difference from a measurement to the mean, and n is the total
number of measurement you taken. All the squares do is get rid of negative differences from the mean. There
are several other the statistical analyses one can take, the standard deviation will suffice for most of your chem
lab experience. Additionally well he can be calculated by hand, it is much easier to calculate standard deviation
using a spreadsheet program such as Microsoft Exceltm.
Length Measurements
You will be given several measuring devices allowing you to measure the length of objects in the lab.
Remember the rules for decimal places when taking measurements from ruled measurement devices.
Mass Measurements
One of the most common and ultimately most quantitative measurements you can calculate the
chemistry lab is mass. In many instances students beginning in the sciences confuse the concept of mass and
weight. Mass is the quantity of matter contained in an object and is a constant value no matter where you are in
the universe. In contrast to this weight is a measurement of force caused by gravitational attraction between the
Earth and an object. As this value of gravity changes across the surface of the Earth, weight is often a poor
measurement of the quantity of matter in an object. While I have made an important distinction between these
two values you will often hear them used interchangeably in this lab and most of undergraduate science. This is
largely because variations in the value of gravity across the surface of the earth are very very small often outside
the realm of error in introductory labs. However if one is launching rockets and planning celestial navigation, the
differences in gravity may be extremely large and must be considered when planning calculations. As we will not
be launching any rockets in this lab sadly, will assume weight and mass are the same for now. Additionally we
will be comparing weights of several objects under identical gravitational conditions, and therefore the values of
mass and weight are essentially identical.
The balances we use in the general chemistry laboratory are top loader analytical balances. They are
analytical devices with protective shields to prevent air currents from causing fluctuations on the balance
surface. The balances are precise enough to read +/- 0.001 g. A typical picture of a balance utilized in lab is
shown in Figure 2.
36
Figure 2: Top Loading Analytical Balance
Density Measurements
Another very helpful property that is often determined in chemistry is density. In chemistry, density is
defined as mass per unit of volume. Mathematically, density is shown as: , where d is density, m is mass (in units
of grams, kilograms, etc.), and v is volume (in units of liters, milliliters, etc.). Typically, density is expressed in
units of grams per milliliter (g/mL). So, if a 15.0-gram solid displaces 5.0-mL of water in a graduated cylinder,
what is its density? It’s mass (m) is 15.0 grams, and its volume (v) is 5.0 mL, so we can simply use the density
equation: d = 15.0 grams/5.0 mL = 3.0 g/mL.
Density can be very useful when comparing unknown objects. For instance, if we have two pieces of a pure
metal and want to know whether or not they are the same, we can determine their densities. If the two
densities are the same (or very close), we can be somewhat sure that the two objects are the same metal.
**One very interesting note about density: Water has a density equal to 1.00 g/mL.
A Note on Numbers in Science
Scientific notation
One of the things you will be doing quite routinely in this lab is collecting measurements that represent
extremely large or extremely small numbers. It is quite unwieldy and can be extremely time-consuming to write
these numbers over and over again. To simplify this we often use a method of representing numbers called
scientific notation. In scientific notation quantities are reduced to a number between one and 10 multiplied by a
power of 10 that represents the number of decimal movements that have occurred; where the power is positive
for left-handed movements and negative for right hand movements. In this style the number between one and
10 is called coefficient and the number of places the decimal is moved is called exponent. You should use this
type of notation any time your answer has more than two decimal places or is larger than 1000 (in many cases
scientific notation is always used no matter what the magnitude of your number).
As an example of this, take for instance a drop of water. One drop of water occupies a volume of
approximately 0.0001 liters. While this number is not unreasonably small it can be difficult to grasp how much
volume this occupies and can be unwieldy to keep writing all those zeros after the decimal. To simplify this
number we move the decimal for places to the right. Our new number becomes 1x10-4 liters. The number one is
a little easier for minds to comprehend.
In the opposite direction, consider how many water molecules are in that same drop. There are
approximately 3,344,444,444,444,444,444,444 of them! Imagine having to write that number over and over
again as part of a series of calculations; you would run out of ink quickly. Again we can simplify this number
significantly by reducing it using scientific notation. There is a decimal not shown after the last four, and we
37
move this decimal place 21 places to the left. Our new number becomes 3.34x1021 (we've taken the liberty here
of reducing the number of decimal places as well to simplify the exponential notation. You should be aware of
significant figures in your calculations to know how many of those decimal places you can go without).
Scientific notation is handy in simplifying math as well. If all of your values have the same exponent, you
may multiply divide add and subtract only the coefficients and the answer will have the same exponent in the
end.
Significant figures
It is good practice in science to only use the digits you are certain of plus one digit of uncertainty. While
it is tempting to write down every number your calculator gives you, at some point your calculator begins
guessing those numbers. It is your job to know which numbers are meaningful and which ones have been
estimated by the calculator. Your final answer should use the number of significant digits of the least accurate
measurement. Imagine you are weighing a piece of metal and find it weighs 10.1238 g. The balance that has
been utilized is giving you for decimal places worth of accuracy, and six total significant digits. If you then take
that piece of metal and determine its volume using a graduated cylinder that indicates 9.91 mL your final answer
can only have three significant digits, the number found in the graduated cylinder measurement. When doing
the calculations you would use all of the numbers from the mass and all of the numbers from the volume, but
the answer would be reported to only three significant figures. You should always determine significant figures
at the end of the series of calculations.
Glassware
Many glassware items have volume marks printed on them. Before using a piece of glassware to make a
volume measurement, you should take a moment to study its calibrations to insure that you know how to read
them properly. A beaker or Erlenmeyer flask can be used for rather rough measurements, while a graduated
cylinder of the appropriate size can be used for measurements of moderate accuracy. In making any volume
measurement, the liquid level should always be the same as your eye level. Erlenmeyer flasks and graduated
cylinders are usually filled and read by raising them to your eye rather than by squatting down to bring your eye
level to the bench top.
Graduated Cylinders
The volume of liquids is often taken in chemistry labs using a graduated cylinder. This volume is a
measurement of the amount of space a given amount of matter takes up. In the case of liquids this volume is
measured directly by reading the marks on the side of graduated cylinder as compared to the fluid level. In the
case of solids however if a graduated cylinder must be used volume is determined by displacing liquid, and
measuring the difference between the liquid level before and after the solid object is introduced. In either case
most graduated cylinders within the scope of introductory science labs measure in milliliters (mL). You will find
that most graduated cylinders made of plastic or glass, and come in many shapes and sizes.
Reading a graduated cylinder requires some practice and patience. When liquid is placed into a
graduated cylinder, a meniscus is formed. This is a curved layer of liquid (either curved up or down, as in the
case of water). A graduated cylinder is read at the very bottom (or very top) of the meniscus, as can be seen in
Figure 3. Properly reading the meniscus will provide you with the volume of the liquid contained in the
graduated cylinder.
38
Figure 3 - Reading the meniscus in a graduated cylinder properly
Though we will rarely need to determine the volume of a solid object in this lab, there will be times
when this type of measurement is necessary. If the object is a simple geometric object such as a cue cylinder or
smear we can take a few simple measurements and determine the volume. There are times however when we
use objects that do not have a regular shape. If we need to determine the volume of in a regular object this is
easily done using Archimedes concept of displacement. Legend has it while taking a bath Archimedes noted that
the water level in his bath rose when he entered it. He made a mark on the tub for this high water level and then
stepped out of the tub. By filling the bathtub with water between the high mark and the low level and keeping
track of how much water he added he was able to determine his volume. Repeating this type of measurement
with objects of known volume he determined objects displace an amount of water equal to their own volume.
What this means is if we measure of volume of water before adding an object and the volume of water after
adding the object the difference between those two measurements is the volume of the object.
As an example imagine filling a graduated cylinder with 15.00 mL of water. We then place any regular
solid object (e.g. a chunk of metal) into the graduated cylinder and note that the level of water rises to 25.00 mL.
25.00 mL - 15.00 mL equals 10.00 mL of displacement, and therefore are object has a volume of 10.00 mL. It
should also be noted that other units of volume are often used in some fields. The medical industry in particular
often refers to volume in terms of “cc”. The cc refers to the concept of a cubic centimeter (or cm 3). Interestingly,
1 mL of water is actually the same as 1 mL of water. This can be a handy conversion in many instances.
Beakers
Beakers are used for mixing chemicals and as containers for samples during testing. The lip at the top
makes beakers easy to use for pouring and the flat bottom and straight sides makes boiling or heating
easy. When used for mixing, the straight sides make the use of some type of stirring equipment necessary. As
with Erlenmeyer flasks, the graduation marks on beakers are only approximate volumes. Beakers should not be
used when accurate volume measurements are required.
Flasks
Erlenmeyer Flasks
The primary use for Erlenmeyer flasks is in mixing chemicals. Because the sides of Erlenmeyer flasks are
slanted and the mouth is narrow, mixing reagent liquids can be accomplished by swirling without fear of spilling
the contents. The advantage of this type of flask is that special stirring equipment (such as a magnetic stirrer) is
not needed. Note: Although many Erlenmeyer flasks have volume markings, they are approximate volumes
only. The graduations on Erlenmeyer flasks should not be used when accurate measurements are required.
Interestingly, the Erlenmeyer flask is named for Albrecht Erlenmeyer, a friend of Sigmund Freud. He persuaded
Freud not to use cocaine as a "cure" for his psychiatric patients.
Filter Flasks
Filter flasks are Erlenmeyer flasks with an adapter (called a side-arm) near the top. A rubber hose is
attached to the side-arm and connected to a vacuum pump or aspirator. When suction is applied, air is drawn
out through the side-arm and a vacuum is created inside the flask.
Volumetric Flasks
39
Volumetric flasks are specially designed containers for very accurate mixing and diluting. Like
volumetric pipettes, volumetric flasks are designed to measure one volume only. The long neck makes it easy to
determine when the final volume has been reached.
Pre-Lab Quiz Questions
To be answered in your notebook before the beginning of the Lab Session.
1. You are determining the density of an irregularly shaped solid that has a piece of dirt on it. You submerse
it in water in a graduated cylinder and measure a displacement of 15.00-mL. Is the actual displacement
more or less than 15.00-mL? Explain your answer.
2. a) A student weighs a piece of solid to be 7.418 g and it displaces 2.5mL of water. What is the density of
that solid?
b) If a student weighs out 6.155 g of the same type of solid as the student in part a, use the density from
part (a) to determine how much water will be displaced.
3. How many milliliters (mL) of water are in 426 cubic centimeters (cm3) of water?
4. Draw a picture of a graduated cylinder containing water. Include the meniscus and label where on the
meniscus you should take a reading.
5. You take the mass of 5 objects and find them to be: 5.01g, 5.06g, 4.97g, 5.01g and 4.98g. (a) What is the
average mass of one object? (b) what is the standard deviation for these measurements?
Procedure
1. Obtain goggles and wear them over your eyes.
Length measurements
2. Your kit should include a small metal block and metal cylinder. Remove the metal cylinder and describe it in
your observations section. Make sure to include the ID letter of the cylinder.
3. Measure the length of the cylinder using a ruler. Make sure you use the appropriate number of decimal
places in your measurements.
4. Rotate the object around in your hand a few times and take another measurement.
5. Repeat step 4 one more time.
6. This time using a meter stick, measure the cylinder 3 times as you did with the ruler.
7. Determine the average for each measurement.
8. Determine the percent difference between the two averages using the following formula:
%D 
Avg ruler  Avg meter stick
Avg ruler
x100
Mass Measurements
9. Your kit should have a bag full of ten nuts. Make sure to note the ID on your bag in your data table.
10. Take any one of the nuts and determine its mass on each of the three balances.
11. Take the mass of each nut using the Analytical Top loading balances. Note each mass in your lab notebook.
The use of the draft shield will help minimize the fluctuation of mass during measurement.
40
The top loading analytical balance is used in the chemistry lab are sensitive instruments that
can be easily damaged if used improperly (which by the way will incur the wrath of your
chemistry professor). You should make sure to read
Appendix D – Using the Balances before coming to lab.
12. Manually determine the average and standard deviation for your measurements in the table below.
a. Add the masses together and divide by 10 to get the average mass.
b. For the column marked “d”, subtract the average from the mass.
c. For the column marked “d2”, multiply d by itself for each entry.
d
d
For the block labeled “
d. For the block labeled “
e.
f.
2
”, sum together all of the d2 values.
2
n 1
For the block labeled “
d
”, divide by the number of measurements minus 1 (nine in this case).
2
n 1
”, take the square-root of the value you calculated in step e. This is the
standard deviation ().
Temperature Measurements
13. Start LoggerPro and open “Lab Experiment 1 – Measurement” from the desktop folder marked – CHM 1045
Experiments; connect the stainless steel temperature probe. The temperature should be displayed on the
computer screen.
14. Fill a 100mL beaker with distilled water.
15. Measure the temperature of the water with both the thermometer and stainless steel temperature probe.
16. Empty and fill the beaker with ice cubes and distilled water.
17. Allow the beaker to sit for 2 minutes and take the water’s temperature with a thermometer and the
stainless steel temperature probe.
18. Pour the cold water down the drain, and refill it with about 50mL tap water.
19. Heat this water on a hotplate set at 350o until it begins to boil.
20. Take the temperature with both the thermometer and stainless steel temperature robe.
Volume Measurements
1. Obtain a 400 mL beaker and fill it to the 100 mL mark with water. Pour this water into a 250 mL beaker.
Note the volume indicated by the beaker; is it still 100 mL exactly?
2. Pour the water from the 250 mL beaker into 150 mL beaker. Note the volume of the water again, how close
is it to 100 mL?
3. Now transfer this water into a 100 mL graduated cylinder, if it reaches the 100 mL mark stop pouring. Note
the volume as measured with the graduated cylinder, the most accurate device of the four. How close is it to
100 mL? In all probability you will be well over or well under 100 mL. The point of this is to show you that
beaker's are not used to measure accurately. If a lab asks for “about” a particular volume then a beaker is
41
4.
5.
6.
7.
fine, however if the lab indicates “exactly” a particular volume then you must use graduated cylinders or
pipets if available.
Look at the markings on the 100 mL graduated cylinder. What is the smallest volume increment indicated I
the marks? How many decimal places should you use when taking measurements with this graduated
cylinder?
Note the level of the water, and the shape of the water surface. This curved feature is called a meniscus, and
you should read the volume at the bottom of the meniscus always.
Pour out approximately half of the water, and take a measurement of the volume.
Fill a 10 mL graduated cylinder with water and repeat the process for taking a volume measurement.
Density Measurements- Water
1. In this first part you will determine the density of deionized water.
2. Take your 10 mL graduated cylinder to a balance and determine its mass when empty. You should make
sure it is dry as possible for doing so.
3. Fill your 10 mL graduated cylinder approximately 5 to 7 mL of water.
4. Return to the same balance (remember the mass of your washer) and determine the mass of your
graduated cylinder and the water?
5. Accurately read the volume of water in your graduated cylinder.
6. Take the mass of both cylinder and water and subtract the mass of the empty cylinder. This is the mass
of water in the graduated cylinder.
7. Divide the mass of water by the volume in milliliters. This is the density of water.
8. Empty your graduated cylinder out and fill it with a different volume of water and repeat the steps for
determining the density.
9. Repeat the process for a third volume of water.
10. Determine the average for your density measurements.
11. Then determine percent difference as you did with the ruler, by using the following formula
%D 
Avg Exp.  1.0
1.0
g
mL
g
mL x100
Density Measurements – Unknown object
1. Remove the small metal block from your kit and note its ID letter.
2. Make sure to describe it in detail in your observation section.
3. Take the cube to the balance and determine its mass.
4. Fill a 100 mL graduated cylinder about half full with DI water and determine an accurate volume for the
water.
5. CAREFULLY, drop the cube in the cylinder and accurately determine the volume of water with the cube.
6. Pour the water out and dry the cube.
7. Using a ruler, determine the volume of you cube (V =l x w x h) in cubic centimeters (cm3)
8. Determine the density of the cube using the following formulas (based on your data):
42
Density 
mass( g )
mass( g )
and Density 
volume(cm3 )
volume(mL)
9. Using the following list, determine the identity of the metal cube.
Metal
Copper (Cu)
Aluminium (Al)
Silver (Ag)
Iron (Fe)
Brass (Alloy of 2 metals)
Nickel (Ni)
Zinc (Zn)
Density (g/cm3 or g/mL)
8.96
2.70
10.49
7.87
8.40
8.91
7.13
You can now clean up and return your equipment. The rest of the lab can be done at home or anywhere you
have access to a computer.
Data Tables(Put these as part of your postlab write up!)
Data Table
Measurement
1
2
3
Average
%D
Balance
Number
1
2
3
Meter Stick
Cylinder ID
_______________
Ruler
Mass (g)
43
44
Data Table
Measurement
1
2
3
4
5
6
7
8
9
10
Average Mass
Nut Set ID _______________
Mass (g)
d=(Mass-Avg)
d2
d
2
d
=
2
n 1

d
=
2
=
n 1
Data Table
Measurement
Cold
Room Temp.
Hot
Trial
1
2
3
Thermometer
Volume (mL)
Mass (g)
SS Temp. Probe
Density (g/mL)
Average Density =
%Diff. =
Grad. Cylinder
100 mL
10 mL
Density Data
Mass of Cube (g)
Volume of Cube (mL)
Liquid Based Density (g/mL)
No. of Decimal Places
Measurement (mL)
Cube ID Letter ___________
Mass of Cube (g)
Volume of Cube (cm3)
Solid Based Density (g/mL)
45
Using Microsoft Excel
All of the computers in the lab have Microsoft Excel installed on them, you're welcome to use this
program in the lab or conduct your last analysis somewhere else. You will be using Microsoft Excel to do the
standard deviation calculations for the washers you did before by hand. I should note that you can use any
spreadsheet program (such as open office or whatever it is that Mac users use) to complete this portion of the
lab; however since Microsoft Excel is the most popular spreadsheet (and the most powerful… And the one I
know best) I will be using it to demonstrate how it can be used to determine the statistics.
You should open Microsoft Excel and have a default blank spreadsheet displayed.
1. In the box that corresponds to column A, row 1, enter “mass (g)”.
2. Starting in the box that corresponds to A2, enter the masses of your washers. Do not put units with your
numbers. You can see how this should look in the figure below.
3. In the box corresponding to A12, type the word "average =".
4. In the box corresponding to B12, type the following exactly: =AVERAGE(A2:A11) this will determine the
average for all your measurements. Note: the equals sign tells Excel that you are about to enter some
sort of formula that will result in the calculation on the numbers in the blocks shown in parentheses.
5. Once you have this done, it is time to calculate the standard deviation. In this cell corresponding to B1,
type “ss2”
6. In the cell corresponding to B2, type the following exactly: =($B$12-A2)^2 this will calculate the square
of the difference between the average and the value in cell A2. This of course is the first step in filling in
the standard deviation formula as you did in the data table in the mass measurement section. Note: the
$ in front of the letter and number are location locks. This tells Microsoft Excel always use that cell even
if the formula location changes. For example say you wanted to move the ss2 column over to column D.
The formula would still look to B12 for the average, but it would now start looking for your
46
measurement in cell C2. Micro soft Excel follows the basic rules of operation order, conducting the
calculation inside the parentheses first. The ^2 symbol tells Excel to multiply the value inside the
parentheses by itself, essentially square that value.
7. There are two ways you can complete the ss2 column for the rest of the values.
a. You can enter the formula in part six for each mass value in column B or
b. You can copy cell B2 and paste it into the next nine cells. I like this method because it's quick.
8. Now we will determine the sum of all the squared differences. In cell A13, type the following exactly:
=SUM(B2:B11) this tells Excel to add all the numbers between B2 and B11 together.
9. The last thing we need to complete his taking the square root of the sum of the squares (cell A13). In
cell A 14 type the following exactly =SQRT(A13/9) this tells Microsoft Excel to take the value in cell A
13 and divide by nine. We have chosen the number nine because there are 10 measurements, and this
represents the sample size. Remember however the standard deviation formula is n -1, where n is the
sample size.
If you have done all of this correct, your spreadsheet should look much like the following figure.
What you have just done is calculate standard deviation… The long way a.k.a. the hard way. Microsoft Excel is an
amazing program that has a way of simplifying this quite a bit. I do this because it is important for you to
understand that Excel is doing exactly the same thing you did in the measurement section above. Similarly,
when using your calculator it is important that you don't just look at a number and hope it is right. By
understanding how we number is calculated is easier for you to figure out when the number is wrong.
So let's go about doing this the easy way.
1. In cell A16, type the following “STDEV”.
47
2. In cell B16 type the following exactly: =STDEV(A2:A11) and hit enter. That is it. What you should have
now is a value that's identical to the one you calculated in B14. Your spreadsheet should look similar to
the one in the figure below.
I have uploaded a copy of the spreadsheet to the class webpage. At the bottom are two tabs one that says raw,
and another that says finished. The raw contains just my data and the finished includes the document as you see
it in my sample. This way you can check how I've done everything and how I've entered everything when you're
creating your own.
Graphs and Excel
You will be expected to create graphs throughout the semester in most cases simply sketching them into
your lab notebook will be appropriate. However occasionally you will be asked to create a graph, print it out and
turn it in. In this case Microsoft Excel and again help you make a neat professional graph.
To do this with our washer data, will create a bar graph based on the number of measurements in
particular ranges. Selection of the range size can be done in many different ways. For the purposes of this
demonstration I'll simply create bins that are 20% of the difference between the largest and smallest washers
(Note: you can use Excel again to find the maximum and minimum size. This is especially helpful if your data set
is too large to skim through. If you type =MAX(data range) in one cell, and =MIN(Data range), where data range
is the start and end cells as we did before, you'll get the maximum value and the minimum value from the data
set).
So let's get to setting up this graph.
1. In cell E4, type “Max”
2. In cell F4, type exactly =MAX(A2:A11)
3. In cell E5, type “Min”
4. In cell F5, type exactly =MIN(A2:A11)
5. In cell E6, type "bin size”
48
6. In cell F6, type exactly =(F4-F5)/5 this tells Microsoft Excel to find the difference between the
maximum and minimum values for your masses and then divide by four. I divide by five, the number of
bins that I wish to set up.
7. In cell E8, type in the label “bin range”.
8. In cell E9, type in the minimum value - minimum value plus bin size (e.g 3.997 – 4.020 g). Note: if you
type numbers into a cell with text the numbers become text as well. This is a way of showing
calculations without having Excel actually do them. Without the = to start off Excel will not perform a
calculation. However typing a number by itself in a cell will leave it as a numerical value.
9. In cell E10 through E13, type in the remaining bin sizes. I recommend you add 0.001 to the starting value
for each successive bin, and 0 .0242 each successive ending range. See the sample spreadsheet for
clarification of this. This ensures that you have no overlap in values, the graph will be skewed if a
particular measurement fell into two bins.
10. In cell G9, enter the number of measurements that fall in the range of the first bin. Excel does have a
way of setting up the bins automatically. However this process is beyond the scope of this simplistic lab.
Because our data set is small enough you can analyze it to determine how many measurements fit into
each bin.
11. In cells G10 through G13, repeat this process. Enter zero if there are no measurements that fall in this
bin range. Note: remember the number of measurements shown in the bin ranges must be the same as
the number of measurements in your sample. In our case there should be 10 measurements and the
number of measurements in the bins should also equal 10.
What you should have now is a spreadsheet that looks like the following figure.
Next, we will set up a bar graph based on our bins.
12. Highlight the bin ranges making sure that all five are selected.
49
13. While holding down the ctrl key, select the five bin counts as well. Note: holding down the control key
allows you to select individual data points. In our case we have a space between the bin ranges and the
bin counts this allows us to not select the empty cells.
14. At the top of the page is a tab labeled insert (See next photos). Click on it and then click on the icon that
says column. In this drop-down list is a series of chart types, select the one that says 2D column. This
should give you a simple bar graph.
15. You should never turn in a graph that has no title and no axes labels. Click once on the graph so that is
highlighted. At the top of the page should be a section labeled “chart tools”. Is a tab labeled “layout”.
Click on this tab. You should now see icons that allow you to change the axes title as well as the chart
title.
16. Click on the chart title icon and label your chart “mass distribution”.
17. Click on the axis title icon and change the primary horizontal axis to “mass ranges”. Note: with the title
box that appears highlighted, you will type in your text in the entry line at the top of the page next to
where it says fx. Remember to not use “=” when typing your title.
50
18. Click on the axis title icon and change the primary vertical axis to a rotated title and type in “bin count”.
What you should have is a spreadsheet and chart that looks like the following figure.
19. Print out your graph and spreadsheet and attach them to your report page with your carbons.
Discussion Questions
1. Comment on the differences in measurements obtained when measuring the cylinder with two different
measurement devices. How close were the measurements? Which gave you better results in your
opinion?
2. Comment on the variation in masses of the nuts in your sample (i.e. are they close or very spread out
based on your observations). Is this reflected in the standard deviation?
3. Comment on the differences found between the different volumetric measuring devices.
a. How does the volume of the beaker compare to that obtained by the graduated cylinder?
b. Which measuring device provides the most consistency between measurements?
c. Which of these measuring devices would you want to use when a very certain volume of liquid is
needed and why?
51
Name ____________________________________
Lab Section_________
Experiment 1: Measurement and Common Laboratory Techniques
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Nut Set ID ____________
Average Mass of a single nut
________________
Standard Deviation
________________
Remember units are important in your answers!
1. What is your unknown metal cylinder ID letter?
_________________
2. What is the average length of your ruler measurements for the cylinder?
_________________
3. Assuming conditions were perfect in the lab, water would have boiled at 100oC. Based on
your experiment, which temperature device gave you the best (closest) measurement of
the boiling temperature of your water sample?
______________________________
4. What is the density of water as determined by your results?
_________________
5. What is your unknown cube ID Letter?
_________________
6. What was its density?
_________________
7. Identify the metal of your unknown cube.
_________________
Make sure to attach your spreadsheet and graph along with the carbon pages for this
experiment to the back of this sheet.
52
Experiment 2 - Chemical Analysis, Determination of solution
concentration by stepwise reactions
Objectives





Learn to measure a specified amount of liquid using a graduated cylinder.
Learn to make observations regarding completion of a reaction.
Learn to separate by filtration.
Learn to minimize errors due to loss.
Practice utilizing dimensional analysis in solving chemical problems.
Relevant Readings
 Kotz et al Lets Review: The Tools of Quantitative Chemistry, Chapter 4 section 5
You should read
53

Appendix E – Handling Chemicals and
54

Appendix F – Filtration.
Theory and Introduction
Early in your chemistry career, you will conduct the majority of your chemical study in an aqueous
environment. What that means in studying reactions and chemical behavior in a solution that is made by
dissolving a chemical in water. In studying this chemistry, it is often essential that you know how much of a
chemical is dissolved into the solution. This last measurement is known as concentration, the amount of a
solute (dissolved chemical) in a given volume of solvent (the stuff doing the dissolving). The most common
concentration unit used in the chemistry you are studying is called molarity (M or Cx). Molarity is defined
mathematically as
molarity ( M ) 
moles of solute
L of solution
This is the focus of the lab you are about to conduct.
By doing a series of reactions, a dissolved substance can be rendered insoluble and separated from the
solution. Knowing that all the dissolved substance came from the original solution, you can back calculate the
mass of dissolved chemical in the given volume of solution you are studying.
Today we will be studying various concentrations of copper (II) sulfate, CuSO4. The copper ions in your
sample will react with excess Zn to form solid copper which we can filter out and determine the mass of. The
reaction for this process is a type of reaction called a Redox reaction, and occurs as follows,
CuSO4 (aq) + Zn (s)  Cu (s) + ZnSO4 (aq)
We then add hydrochloric acid (HCl) to our solution to remove any remaining zinc. Copper does not react with
HCl, however zinc dissolves into the solution releasing hydrogen gas. This reaction is also a redox reaction and
occurs as follows,
Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2 (g)
The copper produced by the reaction is isolated and weighed. Knowing that all the copper came from the
original sample solution, we can relate the amount of copper to the amount of copper (II) sulfate in that
solution. Copper (II) sulfate contains one copper atom for every molecule of the salt. Knowing this, we can
determine how many copper atoms we precipitate and know that each one came from a copper (II) sulfate. To
make this connection however we need to relate mass in grams used in the lab to the number of atoms in a
sample. This is accomplished using the mole concept. A mole is defined as 6.022x1023 atoms per mole, and a
mole is the number of grams one mole of the substance weighs. So this allows the following conversions,
moles CuSO4  gCu x
1 mole Cu 1 mole CuSO4
x
63.546 g Cu
1 mole Cu
This value is the top portion of the molarity equation above. The bottom number of the molarity equation is of
course the amount of sample you originally started with.
Prelab Questions
To be completed as part of your prelab write-up.
1. If 2.075g Co are produced from a 40.0 ml sample of cobalt (II) chloride, what is the molarity of the solution?
2. What is the role of acetone in the final washing?
Procedure
1. Clean all the glassware with soap and tap water and then rinse them twice with distilled water.
55
2. In a clean 250-mL beaker, accurately measure with a 50-mL graduated cylinder 25.0 mL of CuSO4 solution
with the use of a clean graduated cylinder and dilute it with about 15 mL of distilled water.
Remember that you don’t have to use exactly 25.0-mL of solution, you only need to know
EXACTLY how much you use. 24.6 mL or 25.2 mL is all the same experiment wise…let the
math work!
3. Warm this solution gently (set the hotplate to about 175oC).
4. Add about 0.6 g of powdered Zn to the solution and allow the reaction to proceed with occasional stirring
until the blue color of the solution is gone. If needed, add a small amount of additional Zn grains.
5. Add approximately 2-3 mL of conc. HCl and stir well to remove the remaining excess Zn. This reaction is
evidenced by the evolution of H2 gas.
6. Continue heating the reaction mixture on your hot plate to speed up this rather slow reaction. If you still see
Zinc metal at the bottom, add a few more drops of HCl.
7. After reaction of the excess Zn is complete, add about 25 mL of distilled water to dilute the mixture.
8. Decant (means to pour carefully so as not to pour out the solid) the solution into a second beaker, making
sure to leave the copper in the first beaker. If you don’t lose any copper you can pour the liquid down the
drain.
9. Fill the beaker containing the copper with about 200-mL deionized water and stir gently.
10. Repeat the decanting process.
11. Repeat this wash-decant process for a total of four times. This ensures that most of the hydrochloric acid is
removed from the copper and doesn’t add mass to your final copper yield.
12. Using as little deionized water as possible, transfer the copper to a dry, pre-weighed recrystallizing dish.
13. Heat this recrystallizing dish on a hotplate gently until all the water evaporates away.
14. Once completely dried, cool the filter paper completely, and weigh it accurately on a balance. To ensure the
accurate data, place it under the heat lamp for another 5 minutes, cool and weigh. Continue this until you
obtain consistent masses (within +/- 0.001) after successive dryings. Remember, hot objects weigh more
than cooler objects. Before weighing, the filter paper must be cooled to room temperature.
15. Transfer the dry copper in the copper waste bottle and discard the filter paper in trash. Clean all the
glassware with water.
Calculations
The molarity of CuSO4 equals the number of moles CuSO4 per liter of solution.
𝑀 𝐶𝑢𝑆𝑂4 =
𝑀𝑜𝑙 𝐶𝑢𝑆𝑂4
𝐿 𝑆𝑜𝑙𝑢𝑡𝑖𝑜𝑛
To solve for Molarity, we must first determine the number of moles of CuSO4.
𝑚𝑜𝑙𝑒𝑠 CuSO4 = g Cu 𝑥
1 𝑚𝑜𝑙 𝐶𝑢
1 𝑚𝑜𝑙 𝐶𝑢𝑆𝑂4
𝑥
63.546 𝑔 𝐶𝑢
1 𝑚𝑜𝑙 𝐶𝑢
Data and Results
Trial 1
Trial 2
Volume of solution
Weight of filter paper
56
Weight of filter paper & Cu
Weight of Cu produced
Moles of Cu produced
Molarity of solution
Average Molarity
Questions
These questions should be answered COMPLETELY in your lab notebook. You should show all your calculations and
explain your answer
3. Why is it unnecessary to know the exact weight of Zn used?
4. If the HCl does not dissolve all of the excess Zn, how will this affect your calculated molarity of CuSO4?
57
Name__________________________________
Section___________
Experiment 3 Post Lab Sheet - Chemical Analysis, Determination of solution
concentration by stepwise reactions
CuSO4 Molarity __________________Unknown Code __________________
Show the calculation of molarity for your sample from the beginning of your
experiment in the space below (NEATLY!).
58
This lab is homework and should be completed on your own time. It is due at the same time as the
Spectroscopy and Electronic configuration Labs
Experiment 3 – Emission Spectroscopy
Objectives




Gain experience using the Rydberg equation to calculate electronic transitions in the emission
spectrum of hydrogen.
Use emission data to calculate the value of the Rydberg constant.
Confirm the emission properties light generated from elemental discharge sources.
Gain experience using an atomic emission spectrometer.
Relevant Readings

Kotz et al Chapter 6
Theory and Introduction
One the most important advances in experimental chemistry during the 19th century was the
observation that light emitted from elements as they were vaporized into flame contained only wavelengths of a
few specific colors. Later, advances in electrochemistry lead to the ability to repeat these tests bypassing electric
current through gases of these elements, resulting in the same colored light emissions. For example mercury
atoms in fluorescent lamps will glow bluish white, however this light contains only for specific wavelengths; one
blue, one green, one yellow and one red. The light emitted from an elemental source is characteristic of that
particular element. The particular wavelengths detected are called elements spectrum, and in particular its
emission spectrum. This spectrum is the same no matter the source of where the element comes from. For
example the light released by elemental sodium will be the same as that emitted from vaporizing sodium
chloride or any sodium compound.
This property is useful as it allows substances to be analyzed by heating them in flame and collecting
spectrum. This technique is known as atomic emission spectroscopy or AES. One of the benefits of this
technology is it is one of the most sensitive methods available to analyze substance no matter what its
concentration is. This technique is commonly used by astrophysicists to study the gases found in interstellar
space for this exact reason. We however we used the technique to examine a much more fundamental principle
of the element; in particular why do atoms release only certain colors of light?
In the late 1800s, a Swiss high school mathematics teacher by the name of Johann Jakob Balmer was
asking this very question. In analyzing the spectra emitted from hydrogen atoms, he was able to develop a
relationship between the wavelength of the light produced and the physical position of the emitted light in the
spectrum.

hm2
(m2  n 2 )
Where m and n were integers, and h=3654.6 x 10-8 cm. If the equation is solved using n=2 and m=3,4,5,
and 6, the calculated wavelengths are very close to the hydrogen emission wavelengths in the visible range;
While Balmer’s equation was based on experimental observations, it was purely theoretical and not supported
by any physical evidence; Balmer had solved his equation by trial and error. However, it could hypothetically
predict the emission spectrum in the visible, infrared and the ultraviolet ranges. Unfortunately for Balmer, he
would not live to see Niels Bohr and Johannes Rydberg prove his equation valid.
A few years after Balmer published his formula, another mathematician extended his work and
expanded the applicability of the equation. Johannes Rydberg was a high school mathematics teacher like
Balmer. He studied spectroscopy in a much more detailed but less specific (he studied other elements besides
hydrogen) manner. His work allowed him to discover that Balmer's equation was a very specific example of a
59
much more general principle. Rydberg used his impure cult evidence to derived a more general version of
Balmer's equation in which he substituted the quantity he called the wave number (1/) And combined the
resulting constants from his derivation into one constant he called the Rydberg constant. In his notes Redbird
hinted at the idea that light was behaving as both a particle and wave, a concept later demonstrated by Albert
Einstein in the photoelectric effect. Ryberg’s formula exactly matched the light emitted from hydrogen atoms in
the visible region, and is published formula was as follows:
1

 Rm(
1 1
 )
n2f ni2
Here,  is the wavelength of the given emitted light, ni and nf are integers known as principal quantum numbers,
and R is a constant known as the Rydberg constant whose value is 1.0967758x107 m-1. Unfortunately, it took
about 40 years before the power of this equation was really understood. Balmer and Rydberg had discovered
without realizing it, that energy at the atomic level does not obey the rules set forth by Newton. As we now
know from our studies class, the cause of this shocking deviation from ‘normality’ is the idea put forth by Max
Planck; the idea of quantization of energy levels. Thus, the hydrogen emission spectrum opened the door to a
much greater understanding of the chemical and physical world around us.
In 1913, Niels Bohr postulated that electrons orbited the nucleus, and applied work from Max Planck’s
black body investigations and Albert Einstein’s photon energy calculations. His application of these giants of
physics proved Balmer’s equation valid. Using a simple substitution of Planck's law into Rydberg equation he was
able to derive the following formula.
E  hcR(
1 1
 )
n 2f ni2
This can be simplified as h, c and R are all constants.
E  2.18 x1018 J (
1 1
 )
n2f ni2
In this experiment, we will use a spectrophotometer to capture emission spectra from various light
sources. Spectrophotometer function by allowing electromagnetic radiation emitted from the light source
answer through a small slit to creating narrow beam. The light is then passed through a grating that bends the
light according to its wavelength. Each wavelength is reflected off a mirror into a photo diode detector. A photo
diode detector is essentially Einstein's photoelectric device. As light of different wavelengths reaches the
detector, energy levels rise creating a positive signal that is sent to the computer. By adjusting the angle of the
mirror, and passing the light through a second slit, we can to the detector to read only certain wavelengths than
accuracy of +/- 0.5 nm. Luckily for us spectrophotometer functions mostly in an autonomous fashion, we only
need to examine the emission spectra displayed on the computer. Using the wavelengths of the largest peaks
spectrum, we can determine the final principal quantum number each peak.
60
Prelab Questions – To be completed as part of your prelab writeup.
1. Suppose we have hydrogen emission lines originating from an excited energy level ni and ending in the
ground state nf as indicated in the figure below.
a. Calculate the energy released in the emitted radiation by each transition shown in the figure. Calculate
the value in J and J/mol.
b. What type of radiation is emitted by these transitions?
Prelab
Transition
Wavelength
71
91.2 nm
93.8 nm
95.0 nm
97.3 nm
102.6 nm
121.6 nm
61
51
41
31
21
Energy of one
Photon (J)
Energy of a Mole
of Photons (J)
Procedure
Part I – Hydrogen Emission Spectrum
1. Using the emission spectra provided, calculate the wavelengths for the 6, 5, 4 and 3  2 transitions.
2. Label the spectrum with the appropriate transition.
3. Using the wavelengths you calculated, calculate the energy of one photon of visible light at each
frequency (do so in J).
4. Using the value of one photon, calculate the value of a mole of photons.
5. Calculate the wavelength for the 72 transition. Why is this peak not on the spectrum?
61
Part II – Helium Emission Spectrum
6. Calculate the wavelengths for the 6, 5, 4 and 3  2.
7. If possible, Label the spectrum with the appropriate transitions.
Data Table
Hydrogen
Transition
Wavelength
Energy of one Photon
62
52
42
32
Helium
Transition
Wavelength
Energy of one Photon
62
52
42
32
Energy of One photon of light
___________________________
Energy of One mole of Photons
___________________________
62
63
64
Name________________________________________ Section _________
Experiment 3: Emission Spectroscopy
This sheet is due before the start of lab one week after the lab is completed. No Carbon pages need be
attached to this page for this lab only. Everything must be completed on this page in detail for full
credit.
Data Table
Hydrogen
Transition
Wavelength
Energy of one Photon
62
52
42
32
Helium
Transition
Wavelength
Energy of one Photon
62
52
42
32
Questions
1. Show the calculation for a transition in the Hydrogen Spectrum for any one peak in the spectrum. Show ALL
your work in the space below.
2. For the transition you selected in Q1, show the calculation of the energy of one photon at that wavelength.
3. What is the relationship between wavelength and energy based on your calculations? Explain why this
trend occurs from the perspective of the photon
65
4. Why is the 72 transition not found on the spectrum?
5. Theoretically (i.e. according to your calculations) the wavelengths for the hydrogen and helium spectrum
should be the same. As evidenced by the two spectra being different, the wavelengths are not the same.
Suggest TWO reasons why this is so.
6. Think about the aurora borealis. If you have never had the fortune to see one, you’re missing one of
nature’s greatest events…look up a video or images of the aurora. Based on what you learned in this lab,
and what you see in the imagery, explain how the aurora occurs from a chemistry standpoint.
66
Experiment 4 – Spectrophotometry and Molecular Structure
Objectives
 Measure and analyze the visible light absorbance spectra of three standard olive oils: extra virgin, regular,




and light.
Measure the absorbance spectrum of an “unknown” olive oil sample and identify its concentration.
Analyze food dye spectra.
Relate the absorbance pattern of a chemical to its molecular structure.
Gain experience with spectrophotometry.
Theory and Introduction
Olive oil is made by pressing or extracting the rich oil from the olive fruit. It seems like a simple matter
to press the olives and collect the oil, but many oil extraction processes exist for the many different types of
olives grown around the world. To complicate things further, there are also various grades of olive oil, and
carefully selected groups of officials meet to define and redefine the grading of olive oil. To help make our
experiment a more scientific and less political exercise, we will winnow our investigation of olive oil down to a
manageable few variables.
After processing, olive oil comes in three common grades: extra virgin, regular, and light. Extra virgin
olive oil is considered the highest quality. It is the first pressing from freshly prepared olives. It has a greenishyellow tint and a distinctively fruity aroma because of the high levels of volatile materials extracted from the
fruit. Regular olive oil is collected with the help of a warm water slurry to increase yield, squeezing every last
drop of oil out of the olives. It is pale yellow in color, with a slight aroma, because it contains fewer volatile
compounds. Light olive oil is very light in color and has virtually no aroma because it has been processed under
pressure. This removes most of the chlorophyll and volatile compounds. Light olive oil is commonly used for
frying because it does not affect the taste of fried foods, and it is relatively inexpensive.
The visible light absorbance spectrum of chlorophyll gives interesting results. The chemistry of
chlorophyll (some references site four types: a, b, c, and d) creates absorbance peaks in the
400–500 nm range and in the 600–700 nm range. The combination of visible light that is not absorbed appears
green to the human eye, but different sources of chlorophylls will have different ratios of these peaks, which
create various shades of green. The ability of chlorophyll to soak up light energy across a wide swath of the
visible range helps power photosynthesis at optimum efficiency in plants.
In this experiment, you will have two primary goals. First, you will analyze the various grades of olive oil
to determine the absorbance peaks that are present and the relative amount of chlorophyll found in each grade.
You will use a spectrophotometer to measure the absorbance of the olive oil samples over the visible light
spectrum. You will then test an unknown sample of olive oil and grade it as extra virgin, regular, or light.
Procedure
1. Obtain and wear goggles.
2. Connect the Spectrometer to the USB port of your computer.
3. Start Logger Pro. If it is already running, choose New from the File menu.
2. Obtain the 3 small cuvettes containing the known oil concentrations; light, regular and Extra Virgin.
3. Calibrate the Spectrometer.
a) Prepare a blank by filling an empty cuvette 3/4 full with distilled water.
b) Choose Calibrate  Spectrometer from the Experiment menu.
67
c) When the warmup period is complete, place the blank in the spectrometer. Make sure to align the
cuvette so that the clear sides are facing the light source of the spectrometer.
d) Click Finish Calibration, and then click
.
Part I Comparing Three Grades Of Olive Oil and Identifying an Unknown
For Part I of this experiment, you will calibrate the spectrometer with distilled water. Your goals are: (1) to
compare the absorbance spectra of the different grades of olive oil; and (2) to identify the grade of an unknown
sample of olive oil.
4. Conduct a full spectrum analysis of an olive oil sample.
a) Clean one of the oil sample cuvettes by wiping it with a kim wipe.
b) Place the olive oil sample in the spectrometer.
c) Click
. A full spectrum graph of the olive oil will be displayed. Review the graph to identify the peak
absorbance values. Click
to complete the analysis.
5. To save your data for reference later, choose Store Latest Run from the Experiment menu.
6. Repeat Steps 6–7 with the remaining olive oil standard samples.
7. When all 3 samples are completed, return your known samples to your instructor or Lab TA and acquire an
unknown. Note its letter in your notebook and on the data sheet.
8. Repeat Step 6 with the unknown. Note: Do not store the last run.
9. Examine the plots of the olive oil samples and compare your unknow to the known curves. Determine from
these curves which oil type you had.
10. Return your unknown to the lab TA or instructor and get the known samples again.
Part II Comparing the Chlorophyll Concentration of Regular and Extra Virgin Olive Oil
In Part II, you will use the light grade of olive oil to calibrate the spectrometer and presume that light
olive oil contains no chlorophyll. Next, you will compare the chlorophyll content of the regular grade with the
extra virgin grade.
1. Set up a new file and calibrate the spectrometer using light olive oil.
a. Choose New from the File menu.
b. Use the light olive oil known sample as your blank.
c. Choose Calibrate  Spectrometer from the Experiment menu.
d. When the warmup period is complete, place the light olive oil blank in the spectrometer. Make sure to
align the cuvette so that the clear sides are facing the light source of the spectrometer
e. Click “Finish Calibration”, and then click
.
68
2.Measure the absorbance spectrum of regular and extra virgin olive oil.
a) Remove the cuvette of light olive oil from the spectrometer and replace it with the cuvette of regular
olive oil.
b) Click
. A full spectrum graph of the regular olive oil will be displayed. Note the slight difference in
the plot as a result of using the light olive oil as the calibration blank. Click
.
c) To save your data, choose Store Latest Run from the Experiment menu.
d) Measure the absorbance spectrum of the extra virgin grade in the same way.
Part III – Analysis of Vegetable Dyes
14. Calibrate the Spectrometer.
a) Prepare a blank by filling an empty cuvette ¾ full with distilled water.
b) Open the Experiment menu and select Calibrate → (Spectrometer). The following message appears in the
Calibrate dialog box: .Waiting . seconds for the device to warm up.. After 60 seconds, the message
changes to: .Warmup complete..
c) Place the blank in the cuvette holder of the Spectrometer. Align the cuvette so that the clear sides are
facing the light source of the Spectrometer. Click .Finish Calibration., and then click .
15. Conduct a full spectrum analysis of a food dye sample.
a) Empty the blank cuvette and rinse it twice with small amounts of a food dye mixture. Fill the cuvette ¾
full with the food dye mixture and place it in the spectrometer. Align the cuvette so that the clear sides
are facing the light source of the spectrometer.
b) Click . A full spectrum graph of the food dye sample will be displayed.
c) Examine the graph, noting the peak or peaks of very high absorbance or other distinguishing features.
Save and/or print a copy of the graph.
16. Repeat Step 15 with the remaining food dye samples. Remember to keep a copy of each graph.
Part III Data table
Dye Color
Absorbance Peaks (nm)
69
Data Analysis
Part I Comparing Three Grades Of Olive Oil and Identifying an Unknown
1. Describe the graph of each of the standard olive oil solutions. Emphasize the differences between each grade
of olive oil, identifying the absorbance peaks and other distinguishing features.
2. Compare the absorbance spectra of the three grades of olive oil with the sample graph in Figure 4: Visible
Spectrum of Chlorophyll-a. What evidence is there that regular and extra virgin olive oil contains
chlorophyll while the light grade of olive oil does not?
Figure
4:
Visible
Spectrum of Chlorophyll-a
3. Identify your unknown olive oil as extra virgin, regular, or light. Explain your choice.
Part II Comparing the Chlorophyll Concentration of Regular and Extra Virgin Olive Oil
4. Which grade of olive oil, regular or extra virgin, contains the greater amount of chlorophyll? Use your
absorbance spectrum graphs to speculate about how much more chlorophyll one grade contains compared
to the other. You can do this simply by dividing the larger absorbance by the smaller one. This is how many
more times concentrated the chlorophyll is.
Part II Comparing the Chlorophyll Concentration of Regular and Extra Virgin Olive Oil
5. You can see from your data that each dye has a characteristic absorbtion peak. For each dye, explain how
it has the color you saw in the cuvette. BASE YOUR ANSWER ON THE SPECTRAL DATA FOR EACH DYE.
70
6. Based on your unknown sample’s spectrum, which of the dyes do you think are contained in the sample?
Justify your answer using your dye spectra and comparing it to the unknown’s spectra.
7. Below are the structures of two common food dyes most likely found in your diet. Lookat each structure
and describe what structural components lead to the color of the dye. (Hint: I am not asking you to tell me
how they get their specific color, but how they have any color at all based on their structure.)
71
Name________________________________________ Section _________
Experiment 4: Spectroscopy and Molecular Structure
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Unknown Oil Sample ID __________
Did your unknown sample have Chl-a? __________ (Y/N)
What grade of oil was it? ______________
Which grade had the most chlorophyll? __________How many times more Chl-a? __________
72
Experiment 5 – Chemical Nomenclature
Objectives
 To learn the basic types of compounds.
 To gain experience with the rules of naming ionic compounds.
 To gain experience with the rules of naming molecular compounds.
 To gain experience with cations, anions, and polyatomic ions.
 To gain familiarity with the periodic table and elemental symbols.
Relevant Readings
 Kotz et al Chapter 2.
Note: this write up is mainly a theory section that is needed. There are no calculations and you need not
write up the practice sections of this lab. You will be turning in the actual pages from this lab manual as your
post lab.
Theory and Introduction
One of the most important skills a chemist must develop is the ability to name compounds as well as
generate names for compounds. This requires that you learn the language of chemistry. It is the same in any
field where you are just starting out, the first thing you must do is learn to speak in terms of the subject.
Essentially must learn the alphabet, then learn words and then you can learn to write sentences. In our case
that means learning the periodic table, then learn to form compounds and then we can begin studying
reactions.
There are four basic protocols for naming molecules in chemistry; ionic nomenclature, molecular
nomenclature, organic nomenclature, and biochemical nomenclature. While these are not the only sets of
nomenclature rules they are the four most common. For the purposes of this class as well as this lab, we will
focus only on ionic and molecular nomenclature. While there is some overlap between the different types of
nomenclature rules, there are extreme differences amongst them as well. Because of this it is important that
you recognize not only what kind of compound you are working with, but also which rules apply to naming that
compound.
Types of Compounds
For the purposes of this lab, there are two types of compounds we will deal with. The difference
between these two compounds at the molecular level is largely the type of bonding that is involved in joining
the atoms together. The regardless of the type of compound, there are two basic methods for connecting
atoms together in a compound or molecule.
The first is ionic bonding where one atom completely removes an electron from another leaving the first
positively charged and the second negatively charged. These opposite charges will now attract creating a strong
electrostatic attraction between the two atoms called an ionic bond. This type of bonding normally occurs
between a metal and a nonmetal or polyatomic anion. With rare exception the metal is usually positively
charged having lost an electron and is henceforth called a cation. The negatively charged ion is called an anion.
Examples of some of these metals becoming ions are shown in the table below (Figure 5).
73
Group
IA
IIA
IIIA
Element
H
Li
Na
K
Cs
Mg
Ca
Sr
Ba
Al
Cation
H+
Li+
Na+
K+
Cs+
Mg2+
Ca2+
Sr2+
Ba2+
Al3+
Ion name
hydrogen ion
lithium ion
sodium ion
potassium ion
cesium ion
magnesium ion
calcium ion
strontium ion
barium ion
aluminum ion
Figure 5: Ions of Some Main-Group Metals (Groups IA - IIIA)
When the metal involved is a transition metal the process can be more complex. The transition metals in
many cases can form more than one type of cation. In older nomenclature systems these cations often had
trivial names or common names that ended with –ic or –ous (e.g. ferrous or ferric). While these names still
persist today, they are not universally utilized. You should be aware however when a transition metal is part of
a compound you are naming the process is slightly different and the charge on the cation is important.
Examples of transition metals and their cationic charges and names are shown in the table below (Figure 6).
Metal
Ion Systematic name Common name
Cadmium Cd2+ cadmium ion
Chromium Cr2+ chromium(II) ion chromous ion
Cr3+ chromium(III) ion chromic ion
Cobalt
Co2+ cobalt(II) ion
cobaltous ion
3+
Co
cobalt(III) ion
cobaltic ion
+
Copper
Cu
copper(I) ion
cuprous ion
2+
Cu
copper(II) ion
cupric ion
2+
Iron
Fe
iron(II) ion
ferrous ion
3+
Fe
iron(III) ion
ferric ion
Manganese Mn2+ manganese(II) ion manganous ion
Mercury
Hg22+ mercury(I) ion
mercurous ion
2+
Hg
mercury(II) ion
mercuric ion
2+
Nickel
Ni
nickel(II) ion
nickelous ion
+
Silver
Ag
silver ion
2+
Zinc
Zn
zinc ion
2+
Tin
Sn
tin(II) ion
stannous ion
4+
Sn
tin(IV) ion
stannic ion
2+
Lead
Pb
lead(II) ion
plumbous ion
4+
Pb
lead(IV) ion
plumbic ion
3+
Bismuth
Bi
bismuth(III) ion
Figure 6: Ions of Some Transition Metals and Post-Transition Metals (Groups IVA and VA)
74
The nonmetal portion of ionic compound is formed by electrons being added to an atom. If the ion is
monoatomic (a single atom) the charge is fairly consistent; group IV has a -4 charge, group V has a -3 charge,
group VI has a -2 charge and groups seven a -1 charge. However the nonmetals will often form complexes of
two or more atoms with an overall charge. This is called a polyatomic ion. Examples of monoatomic ions are
shown in, and polyatomic ions in.
Group
Element
Anion
Ion name
IVA
C
Si
N
P
As
O
S
Se
Te
F
Cl
Br
I
H
C4Si4N3P3As3O2S2Se2Te2FClBrIH-
carbide ion
silicide ion
nitride ion
phosphide ion
arsenide ion
oxide ion
sulfide ion
selenide ion
telluride ion
fluoride ion
chloride ion
bromide ion
iodide ion
hydride ion
VA
VIA
VIA
VIIA
IA
Figure 7: Ions of Some Nonmetals (Groups IVA - VIIA)
Formula
NH4+
H3O+
OHCNO22N3NO2NO3ClOClO2ClO3ClO4AsO43SeO42MnO4-
Name
ammonium
hydronium
hydroxide
cyanide
peroxide
azide
nitrite
nitrate
hypochlorite
chlorite
chlorate
perchlorate
Arsenate
Selenate
permanganate
Formula
C2H3O2C2O42CO32OCNSCNS2O32CrO42Cr2O72SO42SO32PO43PO43PO43HCO3HSO4HSO3-
Name
acetate (OAc-)
oxalate
carbonate
cyanate
thiocyanate
thiosulfate
chromate
dichromate
sulfate
sulfite
phosphate
monohydrogen phosphate
dihydrogen phosphate
hydrogen carbonate (bicarbonate)
hydrogen sulfate (bisulfate)
hydrogen sulfite (bisulfite)
Figure 8: Formulas and Names of Some Polyatomic Ions
75
When the cation and an anion combine, they will do so with a ratio that results in a neutral or
uncharged molecule. The neutral form of this molecule is often called a salt.
The second type of compound is the molecular compound. This type of compound involves two atoms
sharing a pair of electrons between the two in order to create a lower energy state. This lower energy state and
sharing of a pair of electrons yields in a strong attraction between the two atoms but does not involve any full
charges. This type of bonding is almost exclusive to nonmetal elements.
Nomenclature Rules- Ionic Compounds
Naming ionic compounds is the simplest of the rules we will examine. The process involves naming the
metal first, and then the anion.
1. To name the metal you simply write the name of the metal directly from the periodic table. For
example sodium in NaCl would be written as “sodium” with no modification.
a. The exceptions to this are the transition metals that are able to form multiple charge states.
For example iron can form two salts with chlorine. FeCl2 and FeCl3. To differentiate these
two we would need to add the charge state to the name of the cation. This is done by
adding a after the name of the metal. In our example the first would be iron (II) and the
second would be iron (III). This is only done with transition metals that have more than one
charge state. Metals such as zinc and silver are transition metals but only have one charge
state and would therefore not need Roman numerals in their name.
b. If there is more than one metal simply write their names in the order they are listed and the
formula. For example, NaKSO4 would be named sodium potassium sulfate.
2. If the and ion is monoatomic, we would write the stem name of the element and change the ending
to –ide. For example chlorine would become chloride. In the example above NaCl would be called
sodium chloride. Other examples of anion names Nitrogen – Nitride, Sulfur – Sulfide, Oxygen –
Oxide, Selenium – Selenide.
3. If the ion is polyatomic you must know the name of the complex that is formed. For the most part
this will require simple memorization of the most common ions. There are however some rules that
will help you when you must turn in name into a formula. If the name of the polyatomic ion ends in
–ate, this denotes that it is the most common form of a polyatomic ion with oxygen in it. If the
polyatomic ion ends in –ite, this denotes that it is a form of the polyatomic ion with less oxygen than
the –ate form. If the polyatomic ion has a prefix such as per-, this indicates that is a form of oxygen
containing polyatomic ion with more oxygen than the most common form. Conversely a polyatomic
ion with the prefix hypo- indicates a form of the polyatomic ion with less oxygen than the -ite form.
An example of this is shown in the table below.
Polyatomic Ion
ClOClO2ClO3ClO4-
Name
Hypochlorite
Chlorite
Chlorate
Perchlorate
4. If it is necessary to add more of a cation or anion to create a neutral compound, this is indicated
using subscripts after the cation or anion. For instance if you were to form a compound from Ca2+
and Br-, you would need to bromine (two negatives) to neutralize the +2 charge of calcium. When
you write the formula you would write CaBr2. Subscript to after bromine tells you that there are two
76
of them in the compound. Similarly if the aluminum ion, Al3+, and the nitrite and ion, NO2- form a
compound, you would need three nitrite and ions to balance the +3 charge of the aluminum. To
illustrate this in the formula you would place the polyatomic anion in parentheses and uses
subscript to tell how many nitrates are attached to aluminum. This would be shown as the formula
Al(NO2)3.
a. Do not use parentheses around a single polyatomic ion. For example, iron (III) phosphate
would be written as FePO4.
b. You also do not use parentheses around cations. For example calcium phosphate would be
written as Ca3(PO4)2.
Examples of ions forming neutral salt compounds are shown in the table below.
Cation
Anion
Formula
Na+
Cl-
2+
-
Ca
Br
+
NaCl
CaBr2
Na
2-
S
Na2S
Mg2+
O2-
MgO
3+
O
+
SO42-
Na2SO4
2+
Mg
NO3-
Mg(NO3)2
NH4+
SO42-
(NH4)2SO4
Fe
Na
2-
Fe2O3
Nomenclature Rules- Acid Compounds
The naming rules for acids are very similar to those of the ionic compounds. While acids are not actually
ionic compounds the cation in acids is essentially the hydrogen ion H+. These compounds can be named in the
same fashion as the ionic compounds, but they are often given special names to denote that they are acids. This
is especially true when they are dissolved in water where they're acidic nature is most prevalent. If they are in
gaseous form it would be more appropriate to use the ionic naming rules.
In the acid naming scheme, hydrogen is generally omitted from acids made with polyatomic ions, and
the word Hydro is used in monoatomic acids. For example an acid made with hydrogen and chlorine would be
called hydrochloric acid, while an acid made from hydrogen and sulfate (SO42-) would be called sulfuric acid. The
polyatomic ions containing oxygen that end in -ate would have the ending change to -ic and word acid added to
the end; the polyatomic ions containing oxygen that end in –ite would have the ending changed to –ous and the
word acid added to the end. The use of per- and hypo-would still be used as it was above in the polyatomic ions
naming section. Examples of some acids are shown in the table below.
Example
Compound Name
Acid name
HClO3
hydrogen chlorate
chloric acid
H2SO4
hydrogen sulfate
sulfuric acid
HClO2
hydrogen chlorite
chlorous acid
HCl
hydrogen chloride
hydrochloric acid
77
Nomenclature Rules- Molecular Compounds
When to nonmetals form they combine using covalent bonds to form a molecular compound. In many
cases it is possible for the scene to atoms to combine in different ratios. For example carbon and oxygen form
two very common compounds carbon monoxide (CO) and carbon dioxide (CO2). If we simply applied the ionic
nomenclature rules both compounds would have the same name carbon oxide. This would make distinguishing
between the different chemicals difficult. Because of this molecular compounds use slightly different
nomenclature rules.
1. In general molecular compounds have their formulas written based on the order the atoms appear on
the periodic table. Atoms that are furthest to the left are written first, and those on the right are written
second and so on. If both atoms are in the same group the atom lowest in the column is written first.
This order is important as you will name the atoms in the order they appear in the formula.
2. The first element in the formula is given the same name that appears in the periodic table. For example
in the formula NCl3, nitrogen is furthest left and would be named first. We would simply write nitrogen
without any modification.
a. If more than two elements are in the formula each subsequent elements is named in the same
fashion. For example the chemical NOCl would be named nitrogen oxygen chloride. I should
note that this chemical has an actual name that is not based on these rules but it is beyond the
scope of the course.
3. The last element in the formula is written the same way monoatomic ions were written in the ionic rule
system. We simply change the ending from the stem name to –ide.
4. In order to differentiate different ratios in compounds that are made from the same elements the
molecular naming system uses prefixes that didn't know how many of each element are in the formula.
The table below sums up the first 10 prefixes, and using the only ones you're responsible for. For
example in the chemical P4O10, we would need prefixes on both the first and second elements. As there
are for phosphorus atoms, we would need to use the prefix tetra; and with 10 oxygen atoms, we would
need to use the prefix deca-. That would make the name of this compound tetraphosphorous decoxide.
1
mono-
6
hexa-
2
di-
7
hepta-
3
tri-
8
octa-
4
tetra-
9
nona-
5
penta10
decaa. An exception to the use of prefixes is on the first element in the formula if it is a single atom we
omit the use of mono-. However mono would be used if there is a single atom of a second
element. For example, NO2 is nitrogen dioxide, and the compound N2O would be called
dinitrogen monoxide. Note that we have not use Mono in the first example for nitrogen.
78
This lab will be completed on the post lab data sheets that follow, not in your notebook.
Name________________________________________ Section _________
Experiment 5 Post Lab Sheet – Nomenclature
Monoatomic Ion Containing Salts (-ide endings) : Give the name of the following compounds:
1. NaCl
____Sodium chloride______
2. BaBr2
_______________________
3. K2S
_______________________
4. Al2O3
_______________________
5. LiI
_______________________
6. CaS
_______________________
7. BeF2
_______________________
8. Zn3N2
_______________________
9. Na2S
_______________________
10. Rb2Se
_______________________
11. Mg3P2
_______________________
12. KBr
_______________________
13. Na2O
_______________________
In in the following problems convert the name first into ions and then use the ions to create a neutral
compound. Note these are all monoatomic anions.
Name
Cation
Anion
Compound
1. cesium oxide
Cs+ _
O2-
2. potassium sulfide
______
______
__________
3. barium iodide
______
______
__________
4. zinc chloride
______
______
__________
5. sodium hydride
______
______
__________
6. calcium nitride
______
______
__________
7. gallium oxide
______
______
__________
_Cs2O ___
8. hydrogen bromide
______
______
__________
9. lithium phosphide
______
______
__________
10. beryllium fluoride
______
_______
__________
11. magnesium nitride
12. lithium sulfide
13. aluminum iodide
______
______
______
_______
_______
_______
__________
__________
__________
79
Polyatomic Ions Containing Salts that end in -ATE
Give the name of the following compounds:
1. NaNO3
_______________________
2. CaCO3
_______________________
3. AlPO4
_______________________
4. CsSO4
_______________________
5. Sc(ClO3)3
_______________________
6. AgIO3
_______________________
7. Ga2(SO4)3
_______________________
8. Mg(NO3)2
_______________________
9. Be3(AsO4)2
_______________________
10. Sr(BrO3)3
_______________________
11. BeCO3
_______________________
12. LiClO3
_______________________
13. Ag2SO4
_______________________
14. Zn3(PO4)2
_______________________
In in the following problems convert the name first into ions and then use the ions to create a
neutral compound.
Name
Cation
Anion
Compound
1. zinc chlorate
______
_______
__________
2. aluminum nitrate
______
_______
__________
3. silver sulfate
______
_______
__________
4. gallium chlorate
______
_______
__________
5. rubidium bromate
______
_______
__________
6. beryllium nitrate
______
_______
__________
7. cadmium iodate
______
_______
__________
8. lithium carbonate
______
_______
__________
9. calcium phosphate
______
_______
__________
10. rubidium sulfate
______
_______
__________
80
Polyatomic Ion Containing Salts – Mixed Endings
Give the name of the following compounds:
1. NaNO3
_______________________
2. Ba(NO2)2
_______________________
3. CaSO4
_______________________
4. SrSO3
_______________________
5. GaAsO3
_______________________
6. RbClO3
_______________________
7. CsClO2
_______________________
8. Zn(IO2)2
_______________________
9. KNO2
_______________________
10. BeCO3
_______________________
11. KBrO3
_______________________
12. LiClO2
_______________________
13. AgIO3
_______________________
14. AlPO3
_______________________
In in the following problems convert the name first into ions and then use the ions to create a
neutral compound.
Name
1. aluminum sulfite
2. potassium carbonate
Cation
______
Anion
______
Compound
____
__________
3. barium nitrite
______
______
__________
4. rubidium chlorite
______
______
__________
5. scandium iodite
______
______
__________
6. cesium phosphite
______
______
__________
_______
__________
7. calcium bromite
8. zinc sulfate
9. cadmium phosphate
10. magnesium chlorate
11. silver bromate
______
______
______
______
______
_______
_______
_______
_______
__________
__________
__________
__________
81
MIXED ANIONS - Be careful, some of these are polyatomic and some are monoatomic.
Give the name of the following compounds:
1. TiCl3
_______________________
2. ZnCO3
_______________________
3. FeSO3
_______________________
4. MnF2
_______________________
5. Li3P
_______________________
6. CaCrO4
_______________________
87 KIO2
_______________________
8. Ga2O3
_______________________
9. V(NO3)5
_______________________
10. CrBr3
_______________________
11. NaBrO3
_______________________
12. FePO4
_______________________
13. Cu2O
_______________________
In in the following problems convert the name first into ions and then use the ions to create a
neutral compound.
Name
1. cesium nitrite
Cation
Anion
Compound
____
2. nickel(II) chromate
______
______
__________
3. lithium nitride
______
______
__________
4. chromium(III) sulfate
______
______
__________
5. mercury(II) sulfide
6. iron(III) chlorate
______
______
_______
_______
__________
__________
7. manganese(III) oxide
______
_______
__________
8. vanadium(III) fluoride
______
_______
__________
9. copper(I) carbonate
______
_______
__________
10. scandium nitrate
______
_______
__________
82
Complex Salts and hydrates
Give the name of the following compounds:
1. (NH4)2S
_______________________
2. Cu(NO3)2 ∙3H2O
_______________________
3. Fe(ClO4)2 ∙6H2O
_______________________
4. KMnO4
_______________________
5. Na2SO4 ∙7H2O
_______________________
6. Mn(C2H3O2)2 ∙4H2O
_______________________
7. NaOCl
_______________________
8. (NH4)2CrO4
_______________________
9. K2S ∙5H2O
_______________________
10. Ca(OH)2
_______________________
11. AgCN
_______________________
Give the formulas of the following compounds:
Name
Compound
1. sulfur trioxide
__________
2. tetraphosphorus decoxide
__________
3. iodine pentafluoride
__________
4. ammonia
__________
5. dinitrogen pentoxide
__________
6. ammonium phosphate trihydrate
__________
7. calcium hypochlorite
__________
8. strontium cyanide
__________
9. calcium sulfate heptahydrate
__________
10. aluminum chloride hexahydrate
__________
11. magnesium permanganate
__________
12. cobalt(II) perchlorate pentahydrate
__________
13. barium hydroxide octahydrate
__________
14. lithium acetate
15. ammonium perchlorate
__________
__________
83
ACIDS
Name the following compounds in aqueous solution:
1. HCl(aq)
_______________________
2. HNO3(aq)
_______________________
3. H2SO4(aq)
_______________________
4. HI(aq)
_______________________
5. HF(aq)
_______________________
6. HClO4(aq)
_______________________
7. HClO3(aq)
_______________________
8. HNO2(aq)
_______________________
9. H3PO4(aq)
_______________________
Molecular Compounds
1. NO3
_______________________
2. H2O
_______________________
3. Cl2O7
_______________________
4. IF7
_______________________
5. HCl (gas)
_______________________
6. SO2
_______________________
7. OF
_______________________
8. COCl2
_______________________
9. S2O8
_______________________
10. As3O6
_______________________
11. N2Cl4
_______________________
12. NO
_______________________
13. NH3
_______________________
84
PUTTING IT ALL TOGETHER
Give the name of the following compounds:
1. (NH4)3PO3
_______________________
2. HF(aq)
_______________________
3. Ba3P2
_______________________
4. Fe2(SO3)3 ∙9H2O
_______________________
5. LiClO2
_______________________
6. CdSeO4 ∙2H2O
_______________________
7. Ba(HSO4)2
_______________________
8. MnSO4 ∙7H2O
_______________________
9. Be(MnO4)2
_______________________
10. KH2PO4
_______________________
11. Sr(OH)2
_______________________
12. CaSiO3
_______________________
13. PbHPO4
_______________________
Give the formulas of the following compounds:
Name
1. chromium(III) phosphate hexahydrate
Compound
____
2. potassium bisulfate
__________
3. lithium hydrogen carbonate
__________
4. tin(IV) fluoride
__________
5. zinc hydroxide
__________
6. strontium dihydrogen phosphate
__________
7. barium hydride
__________
8. potassium bisulfate
__________
9. potassium oxalate
__________
10. gold(III) cyanide
__________
11. hydrosulfuric acid
__________
12. scandium iodide
__________
13. dichlorine heptoxide
__________
85
Experiment 6 - Chemical Stoichiometry: Determination of the
formula of a complex salt
Objectives
1. Use stoichiometry to determine the formula of a salt hydrate
2. Practice vacuum filtration
3. Utilize dehydration to determine water content
4. Practice crucible burner use
5. Practice analytical balance use
6. Safely use Bunsen burners
Relevant Readings
 Kotz et al Chapter 2.7-2.11,
86

Appendix D – Using the Balances,
87
Appendix E – Handling Chemicals and
88

Appendix I – Heating.
Theory and Introduction
In this experiment you will be determining the formula of a hydrated metal sulfate salt. The basic formula
is either MSO4. xH2O or M2SO4. xH2O over the course of 2 lab periods.
Part I Determination of the mass percent of sulfate
According to Dalton, mass must be conserved and the formula of a compound is consistent. Because of
this, we can transfer part of a molecule from one form to another and know the mass of what is transferred is
still the same and now contained in the new chemical. Here sulfate in the sample is precipitated out by adding
a solution of barium chloride, quantitatively collected (filtered), dried and weighed to determine the weight
percent of sulfate. Barium sulfate is an insoluble solid precipitates from the solution according to the net ionic
equation:
Ba2+(aq) + SO42-(aq)  BaSO4(s)
By knowing the mass of precipitate formed (g BaSO4) and the molecular weight of BaSO4 you can
calculate the moles of barium sulfate by dividing by the molecular mass of barium sulfate, and according to
Dalton one mole of barium sulfate contains one mole of sulfate; which all came from our unknown metal salt.
Therefore,
moles BaSO4 
gBaSO4
MW BaSO4
And,
moles SO42- = moles BaSO4 x
1 mol SO42
1 mol BaSO4
Using the number of moles and the formula weight of sulfate ion you can then calculate the grams of
sulfate in your original sample.
g SO42- = moles SO42- x
1 mol SO42 96.0 g SO42
x
1 mol BaSO4 1 mol SO42
and finally the mass percent of sulfate ion in your original sample
mass % SO42-=
g SO42
x100
g sample
89
Part II
Determination of the mass percent water
The second week, you will determine the weight percent of water in your salt. From this data and the
weight percent of sulfate ion from week one (if both values are accurate!), you should be able to determine the
empirical formula of your unknown salt.
The water of hydration of these salts is rather loosely bound and can be driven off by heating a sample
in a crucible over a Bunsen burner flame. The equation for the dehydration is
MSO4. x H2O MSO4(s) + x H2O (g)
The difference in the weight of the crucible and sample before and after heating will give the grams of
water lost. From this weight and the original sample weight, you can determine the weight percent of water in
the salt.
mass % H2O = g H2O lost x 100
g sample
Part III
Determination of the empirical formula
You can determine the mass percent of metal from the mass percent of sulfate and mass percent of
water in your unknown salt:
mass % of metal = 100 – [(mass % of SO42-) + (mass % of H2O) ]
Since you now have the equivalent of an "elemental" analysis you should be able to calculate the
empirical formula.
If you assume you have a 100-g sample. This unknown then would contain (use the decimal version of
the percent here):
(% SO42- x 100 g) / (MW of SO42-)= moles of SO42(% H2O x 100 g) / (MW of H2O)= moles of H2O
(% metal x 100g) / (MW of metal) = moles of metal
Examination of the last equation shows us that we cannot calculate the moles of metal since we don't
know the MW of the metal (indeed we want to determine this MW). Since we do know the metal is
monovalent or divalent only, and that the stoichiometry of the compound requires that we have X moles of
H2O for each mole of SO42- there must be either one or two moles of metal for each mole of SO42- (think of it
this way…if it is a +1 metal, you’ll need two of them to cancel sulfate; and if it is a +2 metal you only need one.
Therefore the +1 metal is a 2:1 ratio and the +2 metal is a 1:1 ratio).
Note that the ratio (moles of H2O /moles of sulfate) must be an integer. Similarly (moles of metal/
moles of SO42- ) must be equal to 1.00 or 2.00. Using these facts we calculate two possible molar masses for the
metal.
Assuming MSO4 . xH2O (divalent cation or +2 cation)
moles of metal = moles of SO42-= (% metal x 100 g) / (MW of metal)
MW of metal = (% metal x 100 g)/ moles of SO42Assuming M2SO4 . xH2O (monovalent cation)
90
moles of metal = 2 x moles of SO42-= (% metal x 100g )/ (MW of metal)
MW of metal = (% metal x 100g / (2 x moles of SO42-)
By referring to the periodic table and considering the possible charges and your two calculated
molecular weights, you should be able to identify the metal and therefore the empirical formula for your
hydrated metal sulfate. Also, the color of the sample will give some clue as to which metal is most probable.
Prelab Questions
To be completed as part of your prelab write-up.
Use the following reaction to answer the questions below:
CaSO4·8 H2O+ HEAT CaSO4 + 8 H2O
a. What is the name of the hydrated substance?
b. What is the molar mass of the hydrated substance?
c. What is the molar mass of the anhydrous substance?
d. The molar ratio of hydrated substance:anhydrous substance is 1 to 1. What is the theoretical yield of
anhydrous substance if you start with 500.00 grams of the hydrated substance (remember to convert
to moles before answering this question)?
e. If 275.00 grams of anhydrous substance is recovered, what is the percent yield of this reaction?
f. Would you consider this percent yield to be good or bad? Explain your answer!
Procedure
Part 1 - Sulfate (SO42-) determination
1.
2.
3.
4.
5.
6.
7.
8.
In gravimetric procedures you should weigh your samples by difference. This is done by weighing the
container you will use when empty, and again with the substance to be measured inside. The difference
between these two values is the amount of added substance. Accurately weigh out a sample (0.2 to 0.3
g) of the unknown hydrated metal sulfate on the analytical balance directly in a 400-mL beaker.
Dissolve the sample in approximately 100mL of distilled water. You need not measured this volume
accurately it is acceptable to fill up the beaker to the 100 mL mark.
Add 2mL of concentrated HCl to the sample and heat the solution to near boiling.
Obtain approximately 10mL of 0.25 M BaCl2 solution. Dilute it to about 25mL, and heat to near boiling.
Slowly add the hot BaCI2 solution to the sample from step (1) while stirring down a glass rod.
Allow the precipitate to settle and test for completeness of precipitation by adding a drop of BaCI2
solution. (If precipitation is complete, you should not see any more solid form when you add additional
BaCl2. If it is not complete, add an additional 1 mL portion of hot BaCl2 solution, allow settling and testing
again.)
After complete precipitation, heat the sample for about 20 minutes just below the boiling point. During
the heating, place a watch glass over the beaker. This will prevent splattering in case your sample begins
to boil. Cool to near room temperature.
Meanwhile, obtain a piece of filter paper and flute (fold for filtration) it. Weigh the filter paper on an
analytical balance. Using this weighed filter paper, assemble a gravity filtration apparatus as demonstrated
in
91
9. Appendix F – Filtration. Wet it with distilled water so that the filter paper is seated in the funnel.
10. Without disturbing the precipitate on the bottom of the beaker, pour as much of the clear solution
through the filter paper as you can. Never pour more than enough to fill the half the cone.
11. After the complete transfer of the supernatant solution add enough distilled water to wash all the sides of
the beaker and stir the precipitate with a glass rod and let it settle.
12. Once again transfer almost all the supernatant solution and then stir the remaining solution and the
precipitate to make a suspension, then carefully transfer this suspension to the filter paper. Use a stream
of distilled water and a rubber policeman to facilitate a complete transfer of the solid to the filter paper.
(This is to insure that no solid is lost)
13. When the solid has been completely transferred to the filter paper, rinse it with three separate 10-mL
portions of distilled water. The water washes off excess H+, Cl-, Ba2+ and metal ions.
14. Follow these water runs with three separate acetone rinsings. Acetone displaces off water and aids the
drying process.
15. Remove the paper from the funnel and place it on a watch glass.
16. Place to watch glass and filter paper in the drying oven for approximately 10 to 15 min. or until the salt
and paper is dry.
17. Once everything is dry and cool, weigh the sample and determine the amount of barium sulfate you
collected. Remember, never weigh when objects are hot since heated objects often weigh less than when
they are cooled.
18. Scrape the precipitate in “solid sulfate waste” container. Do not place the filter paper in the waste
container, that may be placed in the trash
19. From the weight of the BaSO4 precipitates, calculate the weight percent of SO42- in your salt.
20. Clean up your bench and glassware, and calculate the weight percent of sulfate in your salt. Make sure
you turn in the sulfate portion of the post lab BEFORE you leave the lab. You will receive this back prior to
the start of part two of the lab in order to ensure you use the correct percentages. However you will be
graded based on your results for this portion of the lab.
Part 2 - Dehydration (% water determination)
Procedure
3.
4.
5.
6.
7.
8.
Your crucible and cover have been set up on a Bunsen burner flame to heat during the pre-lab lecture. This
will ensure they are clean, dry and ready to use once you arrive at your lab station. Take note of the
position of the lid before beginning the lab. This is the position you should use whenever heating sample in
crucible. Remove the flame from under the crucible, and carefully place the lid on your crucible.
Carefully move the crucible to the wire gauze for cooling. Make sure the lid is completely closed well your
crucible was cooling. This will ensure water does not enter the crucible as it cools.
Once the crucible is cooled to room temperature, you should take it to a balance and determine its mass
with the cover on.
Add approximately 0.2 to 0.3 g of your unknown salt hydrate to the crucible and replace the cover.
Determine the total mass of crucible, cover and salt hydrate. The difference between these two
measurements (step four minus step three) is your actual sample size. It is very important that you
determine this mass to at least three decimal places of precision.
Make sure that the cover stays on the crucible and you do not touch the crucible or cover with your fingers
as fingerprints will add mass. Return the crucible to the clay triangle at your station, and vent the lid as you
saw in step one.
Start with a meek flame in the beginning and then increase the intensity (crucible bottom will glow). You
should take observations by carefully removing the cover and observing the chemical process inside the
crucible. You may also occasionally carefully mixed the salts to ensure even heating and complete
dehydration. After 5-7 minutes of strong heating, remove the flame and reposition the crucible lid on
straight.
92
9.
10.
11.
12.
13.
Allow the crucible to cool to room temperature right on the clay triangle (Handle the crucible carefully with
clean tongs.) Weigh the crucible and sample again on the analytical balance.
Repeat the heating, cooling & weighing process until consecutive weighings agree to ±0.001 g.
From the experimental weight loss, determine the weight percent of water in your sample.
From the weight percents of SO42- and H2O, calculate the percent of metal ion in the unknown.
Determine the metal and the formula of the hydrated metal sulfate.
93
Sulfate analysis
Unknown Sample Label
Describe the sample
(Color, texture, etc.)
Data
Mass of sample bottle
Mass of bottle after transfer
Mass of the sample
Mass of filter paper
Mass of filter paper & BaSO4
Mass of BaSO4
Percent sulfate in the sample
Water analysis
Unknown Sample Label
Data
Mass of crucible & cover
Mass of crucible, cover & sample
Mass of the sample
Mass of crucible, cover & sample after heating
Mass of heated sample
Mass of H2O Lost
Total Analysis
Data
2-
Percent SO4 in sample
Percent H2O in sample
Percent metal in sample
Molar mass of metal if M+ (monovalent)
Molar mass of metal if M2+ (divalent)
Identity of the metal
Empirical formula
Questions to be completed at the end of your postlab carbon sheets.
1.
2.
3.
4.
Why must the initial heating of the original unknown sample be done gently?
How can you tell when the sample has been completely dehydrated?
What is the purpose of washing the BaSO4 precipitate with distilled water? With Acetone?
Why must the liquid level remain below the top of the filter paper during filtration?
94
Name________________________________________ Section _________
Experiment 6 – Salt Hydrate Analysis Data Sheet
Must be included with your final lab, carbon copies attached to the BACK of this sheet.
Unknown Sample Label
______________________
% SO42-
__________________
% H2O
__________________
% Metal
__________________
Molecular Formula of your Salt hydrate ______________________
Show the calculations for determining the identity of your sample for all steps. (i.e.
determining % H2O, %SO42-, %metal and elemental analysis to molecular formula.
95
Lab 7 Prelab - Covalent Bonding and Molecular Geometry
You should complete MOST of this experiment before you get to the lab. This will allow you to focus on
the experiment instead of trying to catch up. This lab is about comparing predictions to calculations. If you
haven’t made a prediction the calculation is meaningless.
These are the steps you should follow for all molecules you are assigned:
1. Draw a Lewis Diagram for each compound in the table below, showing all unshared electrons, multiple
bonds and formal charges.
a. Follow the rules on determining the central atom (least electronegative is central) unless otherwise
noted.
b. Remember in molecular Lewis Diagrams ALL electrons must be paired (except rarely on nitrogen.
c. Follow the octet rule unless it is an atom that can deviate from it (e.g. I, S, P, Be or B)
d. Use the formal charges to determine the most likely structure if more than one can be constructed.
2. Determine whether or not resonance structures can be generated and would be necessary to represent the
structure correctly (just say ‘yes’ to the resonance question if it exists. You need not show all resonance
forms).
3. Name the Central atom and how many central atoms there are.
4. For each central atom determine the electronic geometry and the hybridization of the atom.
5. The first one is done for you to serve as an example.
Lewis
Diagram
Nitrogen
trifluoride, NF3
Resonance
No
Number of
Central
Atoms
1
Molecular
Hybridization Geometry
sp
3
Tetrahedral
(trigonal
pyramidal also
correct)
Methyl Amine
CH3-NH2
Tetrachloroiodate
ion, ICl4-
Nitrate ion, NO2-
96
Experiment 7 – Covalent Bonding and Molecular Geometry
Assignment
In your classroom text, review the following topics thoroughly:
 writing Lewis diagrams of molecular structures,
 computing formal charges on atoms in a Lewis diagram,
 writing resonance forms and diagrams of resonance hybrids,
 using VSEPR theory to predict electronic geometry and molecular geometry, and
 bond polarities, bond lengths, and dipole moments
Objectives
To gain more experience drawing Lewis diagrams of molecules and predicting their geometry, and to learn how
to build and study molecular models with computers.
Theory and Introduction
During the lab period, you will get extensive practice in writing Lewis diagrams, predicting the threedimensional structure of molecules, and constructing simple models of molecules.
Preparing for Lab
The following problems require skills similar to those called for in the report on this experiment. For guidance,
look at the Assignment/Report Form for this experiment.
1. Draw a Lewis diagram for each compound in the table below, showing all unshared electrons, multiple
bonds, and formal charges. Unless otherwise indicated, assume that the first atom in a formula (or in a
group of atoms set off by a bond) is the central atom.
2. Determine whether resonance structures are necessary in representing the structure correctly.
3. For each of your Lewis diagrams, circle the central atom(s).
4. For each central atom, determine and name the electronic geometry and the molecular geometry. The
electronic geometry is sometimes called the "arrangement of electron pairs or groups".
5. Complete the table.
6. The first entry is completed as an example.
Lewis Diagram
Resonance?
EXAMPLE
nitrogen
trifluoride
NF3
no
methylamine
CH3-NH2
.
.
Molecular
Geometry
Number of
Electronic
(name and
Central
Geometry
geometry
Atoms
designation
AXmEn)
1. trigonal
1. tetrahedral
pyrimidal
AX3E
1
2. (only one
central atom)
2.
1.
1.
.
2.
2.
97
tetrachloroiodate
ion
ICl4.
nitrite ion
NO2.
1.
1.
.
.
2.
1.
2.
1.
.
.
2.
2.
11. Carry out the prelaboratory assignment (above).
12. Carry out Part I on the Report Form, drawing all Lewis diagrams, making geometry predictions, and
providing other information requested.
13. Finally, use Spartan to complete Part II of the Report Form.
Here is the completed entry for CH3NH2, which is called methylamine:
Formula Lewis Diagram
Resonance?
Number of
Electronic
Central Atoms Geometry
1. C:
tetrahedral
CH3-NH2
no
2 (C and N)
2. N:
tetrahedral
Molecular Geometry
(name and geometry
designation AXmEn)
1. C: tetrahedral (AX4)
2. N: trigonal
pyramidal (AX3E)
Report
Complete the Report Form for this experiment.
In preparing your report, you will build, save, and turn in additional models. If you can't remember how to carry
out operations requested in the report, refer to this tutorial, which has exposed you to all the tools and
operations you should need.
Procedure
Do Part I, and Part II. Draw Lewis diagrams for each compound listed in the Formula columns of Tables 1
and 2 on the following pages. Be sure to include formal charges and resonance structures where appropriate.
Draw perspective drawings of electronic and molecular geometry, and predict the specific bond angle(s) listed.
It is recommend that you complete as much as possible of Part I in the lab. This will help you to see
geometry better on the computer and in structural drawings throughout your text.
Transfer your work NEATLY to the tables. No credit for illegible work or uninterpretable drawings.
Here are models for your perspective drawings of the various electronic geometries, with central atom
X:
98
These drawings are called wedge/dot diagrams. Solid lines represent bonds that lie in the plane of the paper,
dotted lines imply a bond that lies below the plane of the paper, and dark wedges imply a bond that lies above
the plane of the paper.
Here are wedge/dot diagrams for perspective drawings of the all molecular geometries that are based on each
of the electronic geometries on the previous page. Each model has central atom X and bonded groups A.
As you look over these drawings, notice that names of molecular geometries describe the relationship of
bonded groups A to the central atom X, ignoring the unshared electrons. The unshared electrons influence the
molecular geometry, but are not included in its name. In drawings of molecular geometry, the unshared
electrons are sometimes omitted entirely, as in the examples of water and iodine trichloride on the next page.
But a chemist must “see” those undrawn electrons anyway, because they are often crucial to understanding the
chemical and physical properties of the substance. Notice that ICl3 is t-shaped (instead of trigonal), and H2O is
bent (instead of linear), because of the influence of the unshared electrons.
Examples
Use the entries in this table as examples as you make your entries in Tables 1 and 2.
99
Drawing Lewis Structures
1) Compute the total number of valence electrons, as follows:
For each atom, add its group number of electrons.
For each negative charge, add one electron.
For each positive charge, subtract one electron.
Total = V.
2) Draw single bonds between all connected atoms. H atoms are always terminal. A unique atom is usually
central.
3) On all atoms except H, draw lone pairs of electrons to complete their octets. (Exceptions: In BeX2, Be is
surrounded by only 4 electrons; in BX3, B is surrounded by only 6 electrons.)
4) Count the electrons in the current drawing. Total = E.
5) Compare E and V. Use this comparison to select operation a), b), or c).
a) If E = V, go to step 6.
b) If E > V, find two adjacent atoms that both possess lone pairs. Erase both lone pairs and draw an
additional bond between the atoms. This operation decreases E by 2. Repeat this operation until E = V or
V+1. If E = V + 1, the molecule contains an odd number of electrons, and you should read rule 9 before
continuing. Otherwise, go to step 6.
c) If E < V, add a lone pair to an appropriate atom. This atom will usually be the central atom, and it must
be an atom from the 3rd through 7th rows of the periodic table. This operation increases E by 2. Repeat
this operation until E = V or V+1. If E = V + 1, the molecule contains an odd number of electrons, and you
should read rule 9 before continuing. Otherwise, go to step 6.
6) Compute the formal charge on each atom (if you have not already done so). If any pair of adjacent atoms
carry opposite formal charges, and if the positively charged member of the pair is in the 3rd through 7th rows of
the periodic table, move a lone pair from the atom with negative charge to form an additional bond with the
atom having positive charge. (This action eliminates both charges.)
100
7) If rules 1-6 allow you to draw two or more different diagrams, look again at the formal charges. (As a check on
your work thus far, you should find that all diagrams have the same net charge.) Choose the diagram with the
fewest formal charges as the best Lewis diagram. (For example, a diagram having a formal charge of -1 on one
atom would be chosen over a diagram having 2 +'s and 3 -'s.)
8) If rules 1-6 allow you to draw two or more structures that cannot be distinguished on the basis of formal
charge, these diagrams represent resonance forms.
9) For species with an odd number of electrons, follow rules 1-5 until E = V + 1. Then compute formal charges on
all atoms and remove one electron from an atom with a formal charge of -1. This action gives the atom a formal
charge of zero. If more than one choice is possible, remove the electron from the least electronegative atom.
Return to rule 6.
101
Name________________________________________ Section _________
Experiment 7 Postlab - Covalent Bonding and Molecular Geometry
Part I – Formulas for Analysis Report Sheet
Table 1 – Compounds with one central atom Fill in the missing information
Formula
A. Ammonia, NH3
Lewis Diagram
Resonance
Electronic
Geometry
Predicted Bond
Angles
H-N-H
B. Carbon dioxide,
CO2
O-C-O
C. Carbon
monoxide, CO
O-C
D. Ammonium ion,
NH4+
H-N-H
E. Boron
trichloride, BCl3
Cl-B-Cl
F.
Carbonate, CO32-
O-C-O
O-C=O
G. Phosgene, COCl2
O-C-Cl
H. Nitrate ion, NO3-
O-N-O
102
I.
Tellurium
hexafluoride,
TeF6
F-Te-F (Both types)
J.
Xenon
dichloride, XeCl2
Cl-Xe-Cl
K. Iodine
pentafluoride, IF5
F-I-F (both types)
L. Methylene
chloride, CH2Cl2
Cl-C-Cl
H-C-H
H-C-Cl
M. Phosphate ion,
PO43-
O-P-O
N. Xenon
tetraiodide, XeI4
I-Xe-I
O. Antimony
pentachloride,
SbCl5
Cl-Sb-Cl
P. Hexafluorosilcate
ion, SiF62-
F-Si-F (both
angles)
Q. Nitrogen dioxide,
NO2
O-N-O
103
R. Sulfate ion, SO42-
O-S-O
Table 2 – Compounds with multiple central atoms
Formula
S.
Lewis Diagram
Ethyl alcohol,
CH3CH2OH
Resonance
Electronic
Geometry
Predicted Bond
Angles
H-O-C
Use chem3D to measure the C-O
bond length: ________
T. Formamide,
HCONH2
O-C-N
U. Acetate ion,
CH3CO2-
O-C-O
Use chem3D to measure the C-O
bond length: ________
C=O bond length: ________
V. Propene, CH2CHCH3
C-C-C
W. Benzene, C6H6
C-C-C
104
X. Chloroacetylene,
ClCCH
C-C-H
Part II – Bond Length Exercises
1. For compounds B, C, F, 5, and U, you recorded the length of the C-O bond. From your entries for these
compounds, complete this table:
Formula
B. Carbon
dioxide, CO2
C. Carbon
monoxide,
CO
F.
Carbonate,
CO32-
S.
Ethyl alcohol,
CH3CH2OH
U. Acetate ion,
CH3CO2-
C-O Bond Order
(from your diagram)
C-O Bond Order (if yours
differ from these your
structure is incorrect! –
Try again)
2
Actual C-O bond
length
1.250
3
1.135
1.33
1.412
1
1.540
1.5
1.332
2. Use Excel to prepare a graph of correct C-O bond order (y-axis) versus calculated C-O bond length (x-axis).
Make sure that your graph meets all guidelines for graphs in reports (see lab guide). Attach the graph as the final
page of this report.
3. Describe the observed relationship between bond order and bond length (use the correct orders provided in
the table) . In other words, as bond order increases, does bond length increase, decrease, or stay the same?
4. Explain the observed relationship between bond order and bond length. In other words, how does sharing
different numbers of electrons affect the length of covalent bonds between atoms?
Final Submission
Turn in parts I-II completed and attach the graph to the back of the packet.
105
Experiment 8 - Endothermic and Exothermic Reactions
Objectives
 Learn elementary concepts of calorimetry and thermochemistry
 Practice techniques of careful temperature, mass, and volume measurement
 Determine the sign of ∆H for the reaction of citric acid with sodium hydrogen carbonate
Theory and Introduction
One of the most fundamental measurements you can collect for a reaction is the measurement of its
enthalpy (has the symbol ∆H). Entropy is a measurement of how much energy is released or gained by a reaction
as it progresses. Many reactions will release energy, and they are termed exothermic reactions. The alternative
is a chemical reaction which absorbs energy from the surroundings, and these reactions are called endothermic
reactions. Well not directly proportional to temperature, one can use temperature to determine which type of
reaction is occurring. As we will see in later labs, certain thermodynamic calculations can be conducted that will
allow us to relate temperature to enthalpy. This enthalpy value is an important measurement, because it can
provide information regarding bond strength, and the pathway for which the reaction proceeds through.
In today's experiment we will examine two reactions, and determine the type of thermodynamic
reaction that is occurring; exothermic or endothermic. To do this we will first study the reaction
between citric acid and sodium bicarbonate (a.k.a. baking soda). Reaction will occur is shown below. As
you can see citric acid is a triprotic acid (triprotic means it has three acidic hydrogen), and sodium
bicarbonate is acting as a base in this reaction.
H3C6H5O7(aq) + 3 NaHCO3(s) 
 3 CO2(g) + 3 H2O(l) + Na3C6H5O7(aq)
This type of reaction is often called a neutralization reaction as an acid and base cancel each other out resulting
in the formation of water in the salt.
The second reaction we will explore is the reaction between magnesium metal and hydrochloric acid. In
this reaction electrons will be transferred from magnesium metal to hydrogen resulting in the release of
hydrogen gas. This reaction is shown below.
Mg(s) + 2 HCl(aq) 
 H2(g) + MgCl2(aq)
This type of reaction were electrons are transferred is referred to as a redox reaction. Redox is short for
reduction-oxidation. In a reduction, electrons are received by a molecule; and oxidation electrons are lost by
molecule. Both of these processes must occur simultaneously for a redox reaction to proceed.
Procedure
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
Obtain and wear goggles.
Setup a calorimetry cup on a stir plate and add a magnetic stir bar to the cup.
Measure out 30 mL of citric acid solution into the calorimetry cup.
Place a stainless steel Temperature Probe into the citric acid solution ensuring that it is not hit by the stir bar
as it spins.
Connect the temperature probe to the computer interface.
Prepare the computer for data collection by opening the file “01 Endo Exothermic” from the Chemistry with
Computers folder of Logger Pro.
Weigh out 6.0 g of solid baking soda on a piece of weighing paper.
The Temperature Probe must be in the citric acid solution for at least 30 seconds before this step to ensure
it is at equilibrium with the solution.
Begin data collection by clicking
.
After about 20 seconds have elapsed, add the baking soda to the citric acid solution.
Gently stir the solution with the stirbar to ensure good mixing.
Data will be collected by the computer automatically end will end automatically after 300 seconds.
Dispose of the reaction products down the drain with lots of water.
106
14. To analyze your data:
a. Highlight the graph from the middle of the initial 20 seconds until the end. You can do this by clicking
anywhere in the graph and dragging over the area you wish to select. (Note: you can maximize the graph
to fit your screen by clicking crtl-J).
b. Click the Statistics button, .
c. In the statistics box that appears on the graph, several statistical values are displayed for Temp 1,
including minimum and maximum.
d. Record the maximum and minimum temperatures, noting whether the maximum temperature was
reached after the first 20 seconds or if the temperature of the first 20 seconds was the maximum.
e. Sketch a copy of this graph into your lab notebook. Make sure that you label the graph appropriately
with proper axes.
Part II Hydrochloric Acid Plus Magnesium
15. Make sure your calorimetry cup is clean, and set up for the second party experiment.
16. Measure out 30 mL of HCl solution into the Styrofoam cup.
17. Place the Temperature Probe into the HCl solution. Note: The Temperature Probe must be in the HCl
solution for at least 45 seconds before doing the next step.
CAUTION: Hydrochloric acid is caustic. Avoid spilling it on your skin or clothing. Wear chemical splash
goggles at all times. Notify your teacher in the event of an accident.
18. Obtain approximately 0.10 grams of magnesium powder.
19. Begin data collection by clicking
. After about 20 seconds have elapsed, add the Mg to the HCl
solution.
20. Gently stir the solution with the stir bar to ensure good mixing.
21. Collect data until the computer stops collection automatically.
22. Dispose of the reaction products down the drain with lots of water.
23. To analyze your data as you did in part 1.
24. Sketch the graph in your lab notebook.
Data Table
Part I
Part II
Final temperature, t2
°C
°C
Initial temperature, t1
°C
°C
Temperature change, t
°C
°C
Processing the Data
1. Calculate the temperature change, t, for each reaction by subtracting the initial temperature, t1, from the
final temperature, t2 (t = t2 – t1).
2. Tell which reaction is exothermic. Explain.
3. Which reaction had a negative t value? Is the reaction endothermic or exothermic? Explain.
4. For each reaction, describe three ways you could tell a chemical reaction was taking place.
5. Which reaction took place at a greater rate? Explain your answer.
6. Classify each reaction by its reaction type. Explain your answer.
7. When reactions occur, bonds are broken on the reactants side, and formed on the products side. An
Exothermic reaction releases more energy than it uses, and an Endothermic reaction uses more energy than
it releases. The strength of a bond can be determined by how much energy is released to form it. Given
your temperature changes and the above information, which reaction FORMED the stronger bond?
107
Name________________________________________ Section _________
Experiment 8 Post Lab Sheet – Endothermic and Exothermic Reactions
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Which Reaction was endothermic? ___________________________ T = _________ oC
Write the reaction equation below.
Which Reaction was Exothermic? _____________________________T = _________ oC
Write the reaction equation below.
When reactions occur, bonds are broken on the reactants side, and formed on the products side. An Exothermic
reaction releases more energy than it uses, and an Endothermic reaction uses more energy than it releases.
The strength of a bond can be determined by how much energy is released to form it. The more energy
released, the stronger the bond that is formed. Given your temperature changes and the above information,
which reaction FORMED the stronger bond? Make sure to explain your answer.
108
Experiment 9 - Determination of the Number of Moles in a
Chemical Reaction
Objective
 Measure the enthalpy (temperature) change of a series of reactions.
 Use this temperature data to elucidate the stoichiometry of an oxidation-reduction reaction
Relevant Reading
Kotz et al Chapter 3, Appendix G – Calorimetry
Theory and Introduction
Balanced chemical equations relate the mole ratios of the reactants to that of the products through
coefficients. If you know all the chemicals in the reactants and products you can easily balance an equation and
determine these ratios. However often as chemists we are forced to mix chemicals before knowing the results.
When the products are not known, experiments must be conducted to determine the products formed and their
corresponding mole ratios.
In this experiment, we will use two common chemicals as reactants: The hypochlorite ion OCl- (found in bleach)
and the thiosulfate ion S2O32- (used for developing photos). The reaction will produce an unknown set of products
by the following unbalanced reaction:
a OCl- + b S2O32-  products
It is possible to determine the coefficients a and b without actually knowing the products. Because the law of
conservation of mass and the law of constant proportions tell us the reactants are uniform (i.e. all have the
same formula) and are used at a constant rate (conservation of mass) we can measure how much of a reactant
is used and how much product we make. Using this data and a process called continuous variations you will
prepare a series of mixtures of the two reactants. Each mixture will have the same total volume and the same
total number of moles of reactants. These reactions will be exothermic, and therefore the reaction producing
the most heat will be the reaction completely consuming the reactants. From the volumes of reactants used
you can determine a and b for the reaction, and therefore the mole ratio for the reaction.
Safety:
NaOCl is a strong oxidizer. If you get it on your clothes it will remove the color, it will also cause discomfort and
redness to your skin. Wash your hands without soap if you spill it on yourself.
Procedure
1. Obtain and wear goggles.
2. Connect a Temperature Probe to Channel 1 of the Vernier computer interface. Connect the interface to the
computer with the proper cable.
3. Start the Logger Pro program on your computer. Open the file “09 Mole Ratio” from the Advanced Chemistry
with Vernier folder.
4. Obtain about 200 mL of each of the reactant solutions, NaOCl and Na2S2O3.
5. Measure out precisely 25.0 mL of the 0.50 M NaOCl solution. Pour this solution into a Styrofoam cup.
6. Immerse the tip of the Temperature Probe in the Styrofoam cup of NaOCl solution.
7. Measure out precisely 25.0 mL of the 0.50 M Na2S2O3 solution. Note: Do not mix the two solutions yet.
8. Click
to begin data collection. Let the program gather and graph a few initial temperature readings,
and then add the Na2S2O3 solution. Gently stir the reaction mixture with the Temperature Probe.
109
9. Data collection will stop after 3 minutes. You may click
to end data collection before three minutes
have passed, if the temperature readings are no longer changing.
10. Examine the graph to calculate and record the maximum temperature change. To determine the highest
temperature, click the Statistics button, . The minimum and maximum temperatures are listed in the
statistics box on the graph. It may be necessary to trace the graph to determine the initial temperature, if
the minimum temperature is not suitable.
11. Rinse out and dispose of the reaction mixture as directed.
12. Repeat the necessary steps to continue testing various ratios of the two solutions, keeping the total volume
at 50.0 mL, until you have three measurements on either side of the ratio that produced the greatest
temperature change.
Data Table
Volume OCl– (mL)
Volume S2O32– (mL)
25.0 mL
12.5 mL
37.5 mL
10.0 mL
40.0 mL
20.0 ml
30.0 ml
25.0 mL
37.5 mL
12.5 mL
40.0 mL
10.0 mL
30.0 mL
20.0 mL
Temperature change (°C)
Complete trials as described above until the table is filled out. All solutions may be poured down the drain
with lots of running water.
Data analysis
We can compare volumes directly as the concentrations were equal. If they were not equal, we would
first have to convert to mole amounts using concentration and volume data. From your data, find the solution
mix that provided the greatest temperature change. This mixture is the mole ratio that provides the highest
temperature change, and is therefore the most likely ratio found in the reaction equation. Divide the moles
used by the smallest number of moles to get whole numbers. Use these whole numbers to write the reaction
equation for the reactants (i.e. if you get 1.33, you must multiply both molar values by 3).
Questions
1. The molarities of the reactant solutions are equal in this experiment. How important is it that they are
equal? Why (i.e. explain your answer)?
2. Which solution was the limiting reagent in each trial? Explain your answer.
3. Does the mole ratio that you determined in your experiment match the actual reaction equation’s
coefficients for the two reactants? Explain, especially if your mole ratios do not match the coefficients.
110
Name____________________________________________Section___________
Experiment 9 Post Lab Sheet - Coefficients of a reaction
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Stoichiometric coefficient of OCl-
_________________________
Stoichiometric coefficient of S2O32-
_________________________
Show the calculation for the coefficient determining trial of your experiment
Write a proper chemical equation for your results in the space below.
111
Experiment 10 – Reactions of Copper
Objectives
 Examine the concept of conservation of mass and atom economy.
Relevant Readings
 Kotz et al Chapter 3 and 4,
112
Appendix D – Using the Balances,
113

Appendix E – Handling Chemicals,
114

Appendix F – Filtration,
115

Appendix I – Heating
Theory and Introduction
In this experiment you will do a series of reactions with the element copper. These reactions will involve
the use of some new techniques and some interesting color changes. In addition, they will illustrate some of the
classes of chemical reactions that occur in nature, for example, oxidation-reduction reaction. It is important that
you make careful observations as your carry out these procedures. It is also very important that you do the
questions as they appear in the procedures. This will keep you from falling into the trap of the cookbook
approach that was discussed earlier and help you more fully understand the principles involved in the
experiment.
In the first reaction you will dissolve elemental copper and make a solution of copper nitrate. This
solution will then be treated with a base that causes a precipitation. The resulting precipitate will then be
heated to produce an oxide of copper. The oxide will then be treated with sulfuric acid to produce a blue
colored solution of copper (II) sulfate.
Finally, the copper ion will be reduced by zinc to produce copper metal that you will collect and dry. In
some steps of this experiment you are given amounts of chemicals to use so that you have an idea of the scale
at which you should work. In other steps you are left to make your own decisions. Keep in mind the things you
learned about chemical reactions in the last experiment. They will be of use to you here.
Procedure
You will work in pairs on this experiment.
1. Making a Solution of Copper (II) Nitrate: Accurately weigh out 0.500 - 0.530 g sample of copper wire and
record the measurement in your notebook. Using your pencil or pen, coil the wire into a flat spiral and place
it into a clean 250 mL Erlenmeyer flask. Measure 4 mL of concentrated nitric acid found in the FUMEHOOD
into a graduated cylinder. Keep the flask and graduated cylinder in the hood and pour the nitric acid into
the flask. Make observations regarding any changes that occur during the reaction. Be careful not to
breathe any of the fumes! Make notes of your observations in your notebook. Swirl the flask to be sure all
the copper dissolves. If all the copper fails to dissolve within 10 minutes, then add an additional mL of nitric
acid. Once all the copper has dissolved then add about 100 mL of deionized water and take the flask to your
laboratory bench. Be sure that no more fumes are being created before you remove the flask from the
fume hood.
This reaction is an oxidation-reduction reaction that is somewhat complicated.
8 HNO3 (aq) + 3 Cu(s) + O2 (g) → 3 Cu(NO3)2(aq) + 4 H2O(l) + 2 NO2(g)
Identify the elements oxidized (lose electrons and become positively charged) and reduced (gain electrons and
become negatively charged) in this reaction and indicate how many moles of electrons are transferred between
the elements being oxidized and reduced. You may need to look up the rules on oxidation states in your
textbook before coming to lab.
2. Synthesis of solid copper (II) hydroxide: The second reaction that you should perform is the synthesis of
solid copper (II) hydroxide from the copper (II) nitrate you made in the last step. This may be
accomplished by adding 6 M sodium hydroxide to the copper solution. Add a minimum of sodium
hydroxide, but also be sure that you have added enough to complete the reaction. This volume will be
around 10 - 20 mL.
116
How do you determine that you have added enough sodium hydroxide?
Write the balanced chemical equation for the reaction in this step.
3. Formation of the Oxide: To change the hydroxide to the oxide, slowly and carefully heat the solution
containing the precipitate over on a Bunsen burner. Be sure to stir the solution continuously during the
heating with a glass rod (Do not use the brown rubber policemen end!). This will avoid the possibility of
"bumping" which can occur when a large steam bubble forms within a solution due to a local region of
overheating. Bumping will cause a loss of product in addition to possible injury! Watch the solution
carefully and record observations. When the reaction is complete, remove the burner and continue to
stir for a couple of minutes to avoid bumping of the cooling solution. Then stop stirring and allow the
product to settle. Once the product has settled, decant (pour off the liquid leaving the solid behind) the
solution and wash the liquid down the drain with lots of water. Be careful not to lose any solid. Next
wash the solid with hot deionized water.
After the solid has settled, decant off the wash and rinse it down the drain. Save the solid for the next step.
Write the balanced chemical equation for the reaction in this step. Assume the products are copper (II)
oxide and water.
4. Formation of Copper (II) Sulfate: Copper (II) oxide will react with sulfuric acid to produce copper (II)
sulfate. Perform this reaction using 3 M sulfuric acid. As with all reactions, use enough sulfuric acid to
get the job done, but do not generate excess waste. In preparation for the formation of the solid copper,
quantitatively transfer your solution from your 250 mL Erlenmeyer flask to a 250 mL beaker. The term
quantitatively transfer indicates that after you have poured the solution out of your 250mL Erlenmeyer
flask you should wash any solid left in the flask with a small amount of deionized water (less than 5 mL),
adding the rinse to the solution in the 250 mL beaker with your solution. You should repeat this wash
and rinse 2 more times. Now move back to the fume hood.
Write the balanced chemical equation for this reaction. Assume the products are copper (II) sulfate and water.
5. Formation of Copper Metal: Working in the fume hood, quickly add 1.0 g of zinc metal and stir. The
reaction will be complete when the solution is colorless. Record your observations. When the reaction
is complete, decant the solution into a clean beaker. Examine the solid, if you note any solid zinc then
add 10 mL of 6 M HCl and stir the solution until the zinc is no longer present. When gas is no longer
being produced, decant this solution into your beaker containing the previous decantant. This solution
can be poured down the drain with copious amounts of water.
Write the balanced chemical equation for the reaction in this step. Assume the reactants are copper(II) sulfate
and zinc, and that the products are copper metal and zinc sulfate.
Answer this question before proceeding. Was the reaction between zinc and hydrochloric acid endothermic or
exothermic? Use your textbook if you do not know what these words mean. What was the gas produced in the
reaction?
117
Recovery of Copper Metal: You will now transfer the copper from the beaker into a small dish. First weigh your
dry dish. Using your rubber policeman, transfer the entire solid to the dish. Use approximately 5 mL of deionized
water from your wash bottle to aid this transfer (See
118
Appendix F – Filtration for more information on solids transfer). Stir the solid in the water to wash away water
soluble impurities. Decant the water and wash two more times with 5 mL aliquots of deionized water. Carefully
and slowly heat the dish on a hot plate. Watch this carefully to keep the solid from bumping. Continue the
heating and stir the precipitate vigorously with a policeman, be careful not to melt the rubber policeman. Once
the solid is completely dry, allow the dish to cool, and weigh it. The difference between the mass with the solid
and empty dish is the weight of copper you recovered.
Clean-Up. Clean all glassware that was used before leaving the laboratory using soap and water. Hang the
glassware carefully on the drying rack over the sink on the side of the lab. Clean the dish by soaking it in 6 M
nitric acid solution in the hood. Do not use concentrated nitric acid for this step.
Calculate the percent yield you obtained. This can be found by dividing the actual yield you obtained by the
theoretical yield. The theoretical yield is the ideal yield you would have if there were no loss or contamination
during the experiment. The actual yield is the yield you actually obtained in the experiment. In this reaction
conversion to moles is not necessary since you started with copper metal and finished with copper metal.
NOTE: YOU SHOULD MAKE IT A HABIT TO DETERMINE A PERCENT YIELD FOR ALL REACTIONS YOU COMPLETE
IN CHEMISTRY! THEY ARE MANDATORY IN THIS COURSE, AND POINTS WILL BE REMOVED FROM POST LABS IF
THEY ARE MISSING.
Give reasons as to why you might expect your percent yield to be low. Give reasons as to why you might expect
the percent yield to be high. Explain why your percent yield is high (>100%) or low (<100%)?
Data Table
Substance
Copper wire
Copper Precipitate
Step
1
Mass (g)
Moles
6
Data Analysis
1.
In step 1 of the experimental procedure, copper metal is added to concentrated nitric acid. The reaction
between copper metal and concentrated nitric acid is an oxidation-reduction reaction that is somewhat
complicated.
8 HNO3 (aq) + 3 Cu(s) + O2 (g) → 3 Cu(NO3)2(aq) + 4 H2O(l) + 2 NO2(g)
Identify the element or elements that undergo OXIDATION. Identify the element or elements that
undergo REDUCTION. Considering which elements are being oxidized and reduced, determine how
many moles of electrons are transferred between the elements being oxidized and reduced.
2. In step 2 of the experimental procedure, you added 6 M sodium hydroxide to the copper (II) nitrate
solution to form copper (II) hydroxide. You added a minimum amount of sodium hydroxide, but enough
to complete the reaction. How do you determine that enough sodium hydroxide has been added?
3. In step 2 of the experimental procedure, you added 6 M sodium hydroxide to the copper (II) nitrate
solution to form solid copper (II) hydroxide. What is the BALANCED chemical equation for the reaction in
this step?
119
4.
5.
6.
7.
8.
9.
10.
In step 3 of the experimental procedure, the copper (II) hydroxide was heated and decomposed to
copper (II) oxide and water. What is the BALANCED chemical equation for this reaction?
In step 4 of the experimental procedure, 3 M sulfuric acid is added to the copper (II) oxide to produce
copper (II) sulfate and water. What is the BALANCED chemical equation for this reaction?
In step 5 of the experimental procedure, 1.0 g of zinc is added to the copper (II) sulfate solution to
produce copper metal and zinc sulfate. What is the BALANCED chemical equation for this reaction?
In step 5 of the experimental procedure, 1.0 g of zinc is added to the copper (II) sulfate solution to
produce copper metal and zinc sulfate. What are the oxidizing and reducing agents?
In a 2nd reaction in step 5 of the experimental procedure, 6 M HCl is added to copper metal to remove
any excess zinc metal. Is the reaction between the HCl and Zn exothermic or endothermic and what gas
is released in this reaction? Hint: An exothermic reaction releases heat and an endothermic reaction
absorbs heat.
Calculate the percent yield you obtained. Report your percentage to 3 significant digits, e.g. 89.3%.
a. What factors would cause the percent yield to be LOW?
b. What factors would cause the percent yield to be HIGH?
Summarize the main results of this week's experiment in your conclusion. Explain why your percent yield
is greater than 100% or lower than 100%.
120
Name__________________________________
Section___________
Experiment 10 - Post Lab Sheet – Reactions of Copper
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
1.
In step 1 of the experimental procedure, copper metal is added to concentrated nitric acid.
What substance was:
__________ Oxidized
__________ Reduced
2. In step 2 of the experimental procedure, you added 6 M sodium hydroxide to the copper (II) nitrate
solution to form copper (II) hydroxide. What process was used to determine that the reaction is
complete in step 2? Use a complete sentence to describe what you looked for.
3. In step 2 of the experimental procedure, you added 6 M sodium hydroxide to the copper (II) nitrate
solution to form solid copper (II) hydroxide. Write the BALANCED chemical equation for the reaction in
step 2?
4. In step 3 of the experimental procedure, the copper (II) hydroxide was heated and decomposed to
copper (II) oxide and water. Write the BALANCED chemical equation for the reaction in step 3?
5. In step 4 of the experimental procedure, 3 M sulfuric acid is added to the copper (II) oxide to produce
copper(II) sulfate and water. Write the BALANCED chemical equation for the reaction in step 4?
6. In step 5 of the experimental procedure, 1.0 g of zinc is added to the copper (II) sulfate solution to
produce copper metal and zinc sulfate. Write the BALANCED chemical equation for the reaction in step
5?
7.
Calculate your percent yield and show the calculations below.
121
Experiment 11 - Aqueous Reaction Stoichiometry: Analysis of
Vinegar by acid-base titration
Objectives
 Determine the molarity of a commercially available vinegar solution.
 Determine a mass/volume percent.
 Gain experience with accurately conducting an acid-base titration.
 Gain experience with high precision measurement devices such as burets and pipets.
 Determine the equivalence point of a weak acid-strong base titration.
 Calculate the molar concentrations of an acid solution.
Relevant Readings
 Kotz et al Chapter 3-4,
122

Appendix B – Using a Buret,
123

Appendix C – Conducting a Titration
Theory and Introduction
Use of vinegar throughout history is well documented. This week acid has served as a medicinal agent
for thousands of years. One of the first instances of its documented use was in year 400 BC when Hippocrates,
the father of modern medicine, described the use of vinegar as an antiseptic in treating his patients. The word
vinegar is derived from the French word for sour wine, a description where wanting that has been exposed to
oxygen ferments further by bacteria producing a sour taste. Today, vinegar is described as a solution composed
acetic acid, water, and occasionally other substances (mainly stabilizers and preservatives or flavorings). In order
to be sold legally, store-bought vinegar must contain at least 4 g of acetic acid per 100 mL of solution or 4%
(m/v).
Acetic acid ranks amongst the most commercially important chemicals in the world. It is prepared
commercially utilizing bioreactors in a two-step process. In the first step, sugars extracted from fruits such as
apples and grapes are converted by yeast into ethyl alcohol and carbon dioxide as shown below.
C6H12O6  2 C2H5OH + 2 CO2
Once this ethyl alcohol is produced the second step converts it to acetic acid utilizing aerobic bacterial
metabolism.
C2H5OH + O2  HC2H3O2 + H2O
In this lab will utilize the technique of titration to determine the acetic acid content of a commercially
bought vinegar solution. A titration is a process used to determine the volume of a solution of known
concentration that is needed to react with a given amount of another substance whose concentration is
unknown. We conduct this reaction well meeting two conditions:
1. The titration is conducted in a controlled manner to determine the point where the unknown reactant is
completely consumed. This point is called the equivalence point or endpoint.
2. We know the reaction that is occurring, and therefore the stoichiometric ratio between the two reactants.
This allows us to determine a molar equivalency at the endpoint between the known and the unknown
components.
The reaction we are conducting today is known as an acid-base reaction. Generically, these types of
reactions follow a pattern:
Acid + base  salt + water
In this experiment, our acid will be acetic acid, and are standard base will be sodium hydroxide. These two
components will react to form sodium acetate salt and water.
HC2H3O2 (l) + NaOH (aq)  NaC2H3O2 (aq) + H2O (l)
As you can see from this balanced equation for each acetic acid that reacts, one sodium hydroxide is also used.
This one-to-one relationship is the stoichiometric ratio between the reactants. Because of this relationship, if we
determine the volume of sodium hydroxide used at the equivalence point and knowing the concentration of the
sodium hydroxide we will be able to determine the moles of sodium hydroxide used. This molar amount is equal
to the amount of acetic acid in the original solution thereby yielding moles of acetic acid in our sample. This
calculation is shown below.
mol HC2 H 3O2  mL NaOH x
1L
0.35mol NaOH 1 mol HC2 H 3O2
x
x
1000mL
1L
1 mol NaOH
As you can see from the math above, it is fundamental to the calculations that we know the exact location of the
equivalence point. To determine this location we will use two methods; a pH indicator and we will monitor pH
with the pH probe. Using an acid-base indicator to detect the end point is an older less accurate method.
However, it is still useful to include an indicator to ensure electronics (i.e. the pH probe ) are functioning
properly. A pH indicator is a colored substance with two or more differing color phases that change depending
124
on the value pH solution. Indicators themselves are weak acids or bases that will react with the acid or base
added as the experiment progresses. The drawback to using them to identify an end point is mainly that the
color change does not always occur exactly at the pH of our endpoint. This adds a small error to our equivalence
point measurement and ultimately our concentration. In this experiment we will utilize indicator
phenophthalein. This indicator is colorless under acid conditions, but changes to a bright pink color in base.
A much more accurate method of determining the location of the equivalence point for an acid-base
titration is to monitor pH. By monitoring the pH we will be able to locate the most rapid change in pH, and the
equivalence point in the middle of this region. Because this region of rapid change is so important, it is this
region we must accurately measure pH changes and therefore use the most care. Once our graph is complete
we will locate the equivalence point using a calculus-based technique (don't worry the computer does this for
you) called the second derivative. The second derivative will cross the x-axis at the exact location of our
equivalence point.
Once we have determined the equivalence point accurately, and use this information to determine the
number of moles of acid in the sample, we can then calculate concentrations for our unknown vinegar sample.
In general chemistry labs, molarity is the concentration unit most commonly used. It is defined as the number of
moles of substance per liter of solution.
M
mol HC2 H 3O2
L sample
For this experiment molarity (M) of the acetic acid sample is of greatest interest. So we will define molarity as
moles of acetic acid per liter of sample.
M
mol HC2 H 3O2
L sample
In many medical situations and food science labs, concentration is often measured by percentage (i.e. m/v%).
This measurement is convenient as percentage refers to per 100, or in the case of solutions grams chemical per
100 mL. For example before grams of acetic acid per 100 mL commercial requirement would represent 4%(m/v).
The vinegar sample I have purchased for use in this lab indicates that it is 5.1% vinegar. You will test a sample of
this vinegar to determine whether or not the company is being truthful.
In order to determine this percentage, we must modify our concentration calculation slightly. Percent
concentration is in grams per milliliter and then we will multiply by 100 to determine the percent. For this we
must first calculate the grams of acetic acid in the sample using the calculation below:
g HC2 H 3O2  mol HC2 H 3O2 x
60.01g HC2 H 3O2
1 mol HC2 H 3O2
Once we have these grams it is simply a matter of dividing by the milliliters of sample used and multiplying by
100.
% HC2 H 3O2 
g HC2 H 3O2
x100%
mL sample
Prelab Questions
To be completed prior to the start of lab as part of your prelab writeup.
1. A 10.00 mL sample of acetic acid solution requires 25.35 mL of a 0.2255 M sodium hydroxide solution to
completely neutralize it. Calculate the molarity of the acetic acid solution, and the percent concentration.
The reaction that is occurring is the one listed in the introduction above.
Procedure
125
1. Obtain and wear goggles.
2. Obtain 10.0 mL of a vinegar solution of unknown concentration. The volume does not have to be 10.0 mL,
you only must know exactly how much you start with. Do not waste your time attempting to obtain exactly
10.0 mL. In fact it would be better if you did all three titrations using different starting volumes; anything
between 9 and 10 mL will work just fine. Remember to allow the math the hard work!
3. Add about 50 mL of distilled water to a 250 mL beaker.
4. Add the vinegar solution to the beaker. CAUTION: Handle the vinegar solution with care as it is an acid. It
can cause painful burns if it comes in contact with the skin.
5. Place the beaker on a magnetic stirrer and add a stirring bar.
6. Add 2 to 3 drops of the indicator.
7. Connect a pH Sensor to Channel 1 of a Vernier computer interface. Connect the interface to the computer
using the proper cable.
8.
Figure 9: Typical Buret Setup
9. Set up a ring stand, buret clamp, and 50.0 mL buret to conduct the titration (see Figure 9: Typical Buret
SetupFigure 9).
10. Rinse and fill the buret with 0.100 M NaOH solution. In this case it is extremely important that you begin the
titration with the correct level set to 0.00 mL. To do this you must first prepare the buret for filling.
a. Wash the buret at first with soap and tap water. There are long handled brushes hanging in lab that you
can use to clean the inside of the buret.
b. Rinse the buret several times with distilled water. Add 5 to 10 mL of distilled water and allow some of
this to run through the stopcock. With a few milliliters of distilled water remaining, pour it slowly from
the end of the buret while rotating it in your fingers. This will ensure distilled water covers every internal
surface and washes out the impurities.
c. The last step requires that we condition the buret. Conditioning a buret involves rinsing and with the
chemical that will be used inside it. At about 5 mL of sodium hydroxide to the buret, and repeat pouring
it through the end while rotating as you did before.
d. Finally, filled the buret with sodium hydroxide ensuring that it is exactly at 0.00 mL when you're done.
i. You should note several things before filling it. First of all it will take more than 50 mL to fill the
buret. At the bottom below the scale are several milliliters that are not included in the volume
indicated on the side of the glassware.
ii. You will be able to reach zero quickest by filling the buret above the zero mark, and draining into a
waste beaker until it is exactly at 0.00. Doing this will ensure that you only need read the buret and
enter the number directly into the computer. If you are not at zero mark when you begin you must
subtract that volume from every number you enter in the computer.
126
CAUTION: Sodium hydroxide solution is caustic. Avoid spilling it on your skin or clothing.
9. Use a utility clamp to suspend the pH Sensor on the ring stand, as shown in Figure 9.
a. Position the pH Sensor so that its tip is immersed in the vinegar solution but is not struck by the stirring
bar.
b. Gently stir the beaker of acid solution.
10. Run the Logger Pro program on your computer. Open the file “07a Acid-Base” from the Advanced Chemistry
with Vernier folder.
11. You are now ready to begin a titration. The first titration you will conduct will only be to determine the
approximate location of the equivalence point.
a. Open the valve on the stopcock and allow sodium hydroxide to flow into the beaker while stirring.
b. When the solution goes pink, stop the flow immediately.
c. Read the volume from the side of the buret. This is the approximate location of the equivalence point.
12. Dispose of the contents from the beaker down the drain and set up for another titration as you did before.
This time however we will use the pH probe to determine the precise location of the equivalence point.
a. Fill the buret as before ensuring that the volume is 0.00 mL.
b. Before adding NaOH titrant, click
.
c. Once the displayed pH reading has stabilized, click
.
d. In the edit box, type “0” (for 0.00 mL added) and press the ENTER key to store the first data pair for this
experiment.
e. While we are not in equivalence point region, we need not be as accurate or accumulate as many data
points as possible. To speed this process up, take your relative equivalence point and divide by three.
Add that amount of NaOH. When the pH stabilizes, again click
. In the edit box, type the current
buret reading as accurately as possible. Press ENTER. You have now saved the second data pair for the
experiment.
A note on reading burets…. The buret is one of the most precise measuring devices that
you will have used all semester. It is incredibly important that you use it properly or your
readings and equivalents point will possess high amounts of error. The burets you are
using read to two decimal places. Make sure that you read properly at every
measurement. Failure to do so could lead to as much as 1% error at each measurement.
f. Repeat the addition as you did in the previous step until you are within 5 mL of your approximate
equivalence point.
g. Once you are within 5 mL of the approximate equivalence point you should begin taking much smaller
readings in the 0.10 mL or less range. When the graph begins increasing nearly vertically, it is extremely
important that you have as many points as possible. Continue to do this until the pH increase begins to
flatten out.
h. Add three data points once the graph flattens out using the one third volume you calculated to start the
titration off.
13. When you have finished collecting data, click
.
14. Dispose of the reaction mixture as down the drain with lots of water.
15. Rinse the pH Sensor with distilled water in preparation for the second titration.
127
16. Follow the steps below to find the equivalence point, which is the largest increase in pH upon the addition of
a very small amount of NaOH solution. A good method of determining the precise equivalence point of the
titration is to take the second derivative of the pH-volume data, a plot of 2pH/vol2.
a. Right click on the graph, and open the graph properties. Check the interpolate box.
b. On the second tab of this window is a check box for d2. Check this box and a graph of the second
derivative should appear. You may need to adjust the scale to see it properly.
c. Find the point where this new graph crosses the x-axis using the cursor (put the cursor on the point it
crosses and look for the x coordinate in the lower left corner).
16. Sketch a copy of the graph in your lab notebook.
17. Repeat the necessary steps to test the acetic acid solution 2 more times (total of 3 trials).
Data Table
Trial
Equivalence point
(mL)
1
2
3
Data Analysis
1. Calculate the molar amounts of NaOH used in the reaction with the vinegar (HC2H3O2) solution.
2. Calculate the molar concentration (molarity) of the HC2H3O2 solution.
3. Determine the molarity of the vinegar solution.
4. Determine the %vinegar for the solution.
5. Compare the actual molar concentrations of your acetic acid solutions with your calculated molarities. Were the
calculated molarities of your acid solutions within a reasonable range (about 5%) of the actual values? If not,
suggest reasons for the inaccuracy. You will get this data from your instructor when you complete the lab.
128
Name__________________________________
Section___________
Experiment 11 Post Lab Sheet – Aqueous Titration of Vinegar
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Average % Vinegar
__________________
Average Vinegar Molarity
___________________
Show the calculations for one trial in your experiment from volume of NaOH to % and M
129
Experiment 12 - The Potentiometric Titration of Hydrogen Peroxide
Objectives
 Conduct the potentiometric titration of the reaction between commercially available hydrogen peroxide and
potassium permanganate.
 Measure the potential change of the reaction.
 Determine the concentration of the hydrogen peroxide solution.
Relevant Reading
Kotz et al Chapter 4,
130
Appendix B – Using a Buret,
131
Appendix C – Conducting a Titration
Theory and Introduction
Under aqueous conditions, chemical systems are often out of balance and seeking equilibrium. One of
the key methods by which this equilibrium is reached is by the transfer of electrons from one chemical species
to another. This process is referred to as a redox reaction or reduction-oxidation reaction. These redox
reactions are very common in both chemical and biological sciences. For example, the processes of
photosynthesis and respiration are such reactions. In photosynthesis, shown below, the oxidation state of
carbon (in CO2) begins at +4 and is reduced to zero (in glucose).
Respiration, essentially the reverse of the photosynthetic process, involves a similar exchange of electrons.
In studying redox reactions, the most common method of tracking their progress is ORP (oxidationreduction potential). It is a measure of the potential for electrons to travel from one species to another
molecule with higher affinity for them. It is worth mentioning that the name oxidation does not imply the need
for oxygen, only that oxygen has a great affinity for electrons.
A titration, as you recall, is a convenient method of learning more about a solution by reacting it with a
second solution of known molar concentration. There are a number of ways to measure the progress of a
titration. In this experiment, you will use an ORP (Oxidation-Reduction Potential) Sensor to measure the
electrical potential of the reaction being studied in a titration. Your data will look like an acid-base titration
curve, however pH will not be measured. The volume of KMnO4 titrant used at the equivalence point will be
used to determine the concentration of the H2O2 solution. Your sample of H2O2 will come from a bottle of
ordinary, over-the-counter hydrogen peroxide purchased at a grocery or a drug store. According to the bottle, it
has a 3% (m/v) concentration; you will confirm this.
The reaction being studied, the reduction of manganese is shown below.
5 H2O2 (aq) + 2MnO4– (aq) + 6 H+ (aq)  5 O2 (g) + 2 Mn2+ (aq) + 8 H2O (l)
As you can see, the stoichiometry is more complex than the reaction previously studied. However, the
mathematical approach is the same as before. From your titration data, you will determine the equivalence
point of the reaction. Using this volume and the balanced equation above, you can then determine the amount
of peroxide in the store bought solution.
mol H 2O2  mL MnO4  x
0.020 mol MnO4  5 mol H 2O2
1L
x
x
1000mL
1L
2 mol MnO4 
With this information, we can determine the molarity (M) of the peroxide by simply dividing by the liters of
solution (L).
M
mol H 2O2
L sample
As with the vinegar lab, we need to measure the concentration in terms of mass-volume percentage. For this
we need to convert moles to grams again.
g H 2O2  mol H 2O2 x
34.01g H 2O2
1 mol H 2O2
Once we have these grams it is simply a matter of dividing by the milliliters of sample used and multiplying by
100.
% H 2O2 
g H 2O2
x100%
mL sample
132
Prelab Questions
1. Assign oxidation numbers to each element in this unbalanced equation. (remember peroxide is an
exception)
H2O2 (aq) + KMnO4 (aq) + H2SO4 (aq) → O2 (g) + MnSO4 (aq) + K2SO4 (aq) + H2O (l)
2. The apparatus for your titration is shown in the illustration.
a. Identify each piece of equipment.
b. Based on the information in the Experimental Procedures, which solution(s) will you put in the glassware
labeled i?
c. Based on the information in the Experimental Procedures, which solution(s) will you put in the glassware
labeled ii?
Procedure
1. Obtain and wear goggles.
2. Prepare an acidified and diluted hydrogen peroxide, H2O2, solution for the titration.
a. Measure out precisely 10.0 mL of a 3% hydrogen peroxide solution
b. Add 90.0 mL of distilled water. Mix the dilute H2O2 thoroughly (keep this step in mind as we have just
diluted the peroxide 1:10…so it is now 0.3%).
3. Measure out precisely 10.0 mL of the dilute H2O2 solution.
4. Add about 25 mL of distilled water and about 10 mL of 4.5 M sulfuric acid, H2SO4, solution. CAUTION: H2SO4
is a strong acid, and should be handled with care.
5. Transfer the solution to a 250 mL beaker.
6. Place the beaker of H2O2 solution on a magnetic stirrer and add a stirring bar.
7. Connect the ORP Sensor to LabQuest and open the “08a Potentiometric Titration” file in the Advanced
Chemistry with Computers Folder.
8. Set up a ring stand, a buret (butterfly) clamp, and a buret to conduct the titration (see Figure 9).
133
9. Clean, rinse and fill the 50 mL buret with 0.020 M MnO4– solution. CAUTION: Handle theKMnO4 solution with
care; it has been mixed with H2SO4, which can cause painful burns if it comes in contact with the skin.
10. Use a utility clamp to suspend the ORP Sensor on the ring stand, as shown in Figure 9. Position the ORP
Sensor so that its tip is immersed in the H2O2 solution but does not interfere with the movement of the
magnetic stirring bar.
11. Gently stir the beaker of solution.
12. The objective of your first trial is to determine the region of the titration curve near the equivalence point,
and not to precisely determine the equivalence point. At the equivalence point, you will see a faint pink
color of unreacted MnO4– solution.
a. Start data collection.
b. Before you have added any of the MnO4–solution, select Keep and enter 0 as the buret volume in mL.
Select OK to store the first data pair.
c. Add 1 mL of the MnO4– titrant. Stir the solution gently at all times.
d. When the potential stabilizes, select Keep and enter the current buret reading. Make this reading as
precise as possible5.
e. Select OK to store the second data pair.
f. Add MnO4– solution in 1 mL increments and enter the buret reading after each increment.
g. Continue adding MnO4– solution until the potential value remains constant (levels off for 3 data points).
h. Stop data collection to view a graph of potential vs. volume.
i. Examine the titration curve and estimate the volume of MnO4– solution used to reach the
equivalence point of the titration.
ii. Record this value in your data table for Trial 1.
13. Dispose of the reaction mixture as down the drain.
14. Rinse the ORP Sensor with distilled water in preparation for the second titration.
15. Repeat the necessary steps to conduct a second titration with a new sample of H2O2 solution. You may draw
your H2O2 sample from the remaining 90 mL of H2O2 that you diluted in Step 2a.
16. When you conduct the second titration, carefully add the MnO4– solution drop by drop in the +/-5 mL region
near the estimated equivalence point from trial 1, so that you can precisely identify the equivalence point of
the reaction. You should be adding less than 0.05 mL amounts in this range.
17. Examine your titration data to identify the region where the potential made the greatest increase. The
equivalence point is in this region.
18. Identify the equivalence point as precisely as possible and record the volume of MnO4– solution used to
reach the equivalence point of the titration for Trial 2.
a. A good method of determining the precise equivalence point of the titration is to take the second
derivative of the pH-volume data, a plot of 2pH/vol2.
b. Right click on the graph, and open the graph properties. Check the interpolate box.
c. On the second tab of this window is a check box for d2. Check this box and a graph of the second
derivative should appear. You may need to adjust the scale to see it properly.
d. Find the point where this new graph crosses the x-axis using the cursor (put the cursor on the point it
crosses and look for the x coordinate in the lower left corner).
19. Conduct a third trial as you did with the second trial.
20. Sketch the titration curve in your notebook.
Data Table
Trial 1
5
Trial 2
Trial 3
Remember how many decimal places a buret reads to!
134
Volume of H2O2 solution
Volume of MnO4– solution used
at equivalence point (mL)
Data Analysis Questions
1. Calculate the moles of MnO4– used to reach the equivalence point of the reaction for each trial.
2. Use your answer to question 1, along with the balanced redox equation in the introduction, to calculate the
moles of H2O2 in the sample of solution for each trial.
3. Calculate the molar concentration of the H2O2 solution for each trial.
4. Calculate the %(m/v) concentration of the H2O2 solution for each trial.
5. The hydrogen peroxide solution that you tested is a commercial product with a concentration, as described
on the label of the container, as 3%. As stated in the introductory remarks, a 3% H2O2 solution converts to a
molarity of 0.88 M. Compare your experimentally determined molarity of H2O2 to the label description.
a. Calculate the average molarity and average %.
b. Calculate the %error between your molarity and the %(m/v)
135
Name__________________________________
Section___________
Experiment 12 Post Lab Sheet - Potentiometric Titration of Hydrogen Peroxide
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Average % peroxide
____________________
Average peroxide molarity
____________________
%Error for %m/v
____________________
Show the calculations for one trial of your experiment
136
Experiment 13 - The Synthesis and Analysis of Aspirin
Objectives
 Conduct a synthesis reaction.
 Gain experience with vacuum filtration.
 Conduct a colorimetric analysis of your aspirin sample.
Relevant Reading
 Kotz et al Chapter 4,
137

Appendix A – Common Glassware,
138

Appendix D – Using the Balances,
139

Appendix F – Filtration,
140

Appendix I – Heating
Introduction and Theory
Chemistry is a diverse field which can be divided into many sub disciplines. One of the most common
methods used to divide chemistry is by the type of chemistry being studied. By type of study we refer to the source
of the chemicals for the technique. Using this method we see divisions of chemistry such as analytical, inorganic,
physical and organic chemistry. It is this last division that we devote our study to today. The primary focus of organic
chemistry is the study of molecules based on a carbon structural frame. Organic chemistry itself can be further
divided into chemistry focused on the properties of molecules, and the synthetic chemists, those chemists who
study the manufacture of chemicals from smaller starting molecules. Synthetic chemistry is not limited to organic
chemists, and can be found in several other disciplines. This lab today focuses on the organic synthesis of
acetylsalicylic acid, more commonly known as aspirin.
Organic chemistry itself covers many different types of molecules, again all based on a simple carbon
hydrogen structure. These structures are modified by adding small groups of atoms such as nitrogen, oxygen, sulfur,
and phosphorous to create a molecule whose reactivity changes. These additional side groups change the properties
of a molecule, and are therefore often referred to as functional groups. Because functional groups often change a
molecule in a predictable way, functional groups are often used to classify compounds in organic chemistry. The
reagents used in this lab contain a few of these functional groups, in particular some containing oxygen atoms. The
two functional groups of main interest are the carboxylic acid and the ester. These are shown in the figure below.
So far in this lab we have focused on measurements and chemical analysis labs. Today will synthesize a
chemical compound that may perhaps be common to your everyday life. Acetylsalicylic acid, more commonly
known as aspirin will be synthesized from salicylic acid and a molecule called acetic anhydride.
In this sequence we will utilize phosphoric acid as a catalyst to initiate the reaction.
Once you have synthesized the product, we will purify it and perform an analysis to determine if we have indeed
created acetylsalicylic acid by testing it with a device known as a colorimeter.
Prior to 1800s, willow bark was known to facilitate pain relief when brewed into a tea or sucked upon to
relieve pains such as headaches and toothaches. It is believed that Western society learned of this technique
141
from the Native Americans of both North and South America who extracted salicylic acid for this purpose. As its
use expanded, a drawback to salicylic acid was discovered. Many people found the taste incredibly bitter, and
often irritating to the lining of the stomach often leading to ulcers. 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. To
combat this problem the German pharmaceutical company in 1899 known as Bayer pharmaceuticals recognized
the problem as an acid buildup in the stomach and adopted Felix Hoffmann's synthesis. They named this
chemical aspirin, and have been selling it ever since.
Prelab Questions – To be completed as part of your prelab write-up
1. What is a catalyst?
2. What does it mean to quantitatively transfer a solution? Describe the process of doing it in your response.
3. Determine the molar mass of the following chemicals:
a. salicylic acid
b. acetic anhydride
c. acetylsalicylic acid
4. 6.84 grams of salicylic acid and 5.29 mL of acetic anhydride (density: 1.080 g/mL) are mixed together. 3.79
grams of acetylsalicylic acid are obtained in the reaction.
a. What is the theoretical yield of the reaction?
b. What is the percent yield of the reaction?
Procedure
Part I: Synthesize Aspirin
1.
2.
3.
4.
5.
a.
b.
6.
7.
8.
9.
10.
11.
12.
Obtain and wear goggles.
Measure out 2.0 grams of salicylic acid into a 50 mL Erlenmeyer flask.
Add 5.0 mL of acetic anhydride and 5 drops of 85% phosphoric acid.
Swirl the mixture. If necessary, use a sparingly small amount of distilled water to rinse down any bits of solid
that may be on the inner walls of the flask. CAUTION: Handle the phosphoric acid and acetic anhydride
with care. Both substances can cause painful burns if they come in contact with the skin.
Heat the mixture on a hot plate, at 75°C, for 15 minutes, or when the mixture ceases releasing vapors.
Stir the mixture occasionally during heating.
After about 10 minutes, add 2 mL of distilled water to the flask.
Set up a Büchner funnel and filter flask so that you are ready to filter the reaction mixture after it has
cooled.
When you are confident that the reaction has reached completion (no vapors appearing), carefully remove
the flask from the hot plate.
Add 20 mL of distilled water.
Allow the mixture to cool to near room temperature.
Transfer the flask to an ice bath for about five minutes. As the mixture cools, crystals of aspirin should form
in the flask.
Transfer the contents of the cooled flask to a Büchner funnel assembly. Filter the mixture with vacuum
suction.
When most of the liquid has been drawn through the funnel, turn off the suction and wash the crystals with
5 mL of cold, distilled water.
a. After about 15 seconds, turn the suction back on.
b. Wash the crystals with cold, distilled water twice more in this manner.
Data Table
Synthesis of Aspirin
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Mass of salicylic acid used (g)
Trial 1
Volume of acetic anhydride used (mL)
Mass of acetic anhydride used (vol. × 1.08
g/mL)
Mass of aspirin synthesized (g)
Data Analysis
1. What is the theoretical yield of aspirin in your synthesis? The mole ratio is 1:1 between salicylic
acid and acetic anhydride in this reaction.
2. Based on the results of the absorbance testing with the Colorimeter, what is the percent purity of
your sample of aspirin?
3. Using the reaction for the synthesis today, write a molecular equation for what has occurred.
143
Name__________________________________
Section___________
Experiment 13 Post Lab Sheet – Synthesis of Aspirin
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
Theoretical Yield
____________________
Actual Yield
____________________
% Yield
____________________
% Purity
____________________
Show your Theoretical and % Yield calculations below.
144
Experiment 14 - Additivity of Heats of Reaction: Hess’s Law
Objectives
•
•
•
•
•
Determine the heats of solution for three reactions.
Combine equations for two reactions to obtain the equation for a third reaction.
Use a calorimeter to measure the temperature change in several reactions.
Calculate the heat of reaction, ΔH, for the three reactions.
Use the results to confirm Hess’s law.
Relevant Reading

145

Appendix D – Using the Balances,
146

Appendix E – Handling Chemicals, Appendix G – Calorimetry,
147

Appendix I – Heating
Theory and Introduction
The study of energy transformations in chemistry is known as thermodynamics. The study of
thermodynamic properties of chemical reactions is extremely important. From this information it is possible to
deduce the reaction mechanism (the physical path from reactants to products) the reaction follows. When
studying a chemical reaction we must clearly define the conditions we are observing. The chemicals under study
that change are called the system. Anything in the immediate presence but not including the system, is referred
to as the surroundings. Once this is defined, we can study the transfer of heat between the two. Heat we define
as energy flowing between the reaction in the surroundings to establish a thermal equilibrium. The flow of heat
will always occur from a region with higher temperature to one with a lower temperature until temperature is
the same in both regions. At this point the flow of heat stops.
The quantity of heat transferred is denoted using a q. If this quantity is positive then the reaction is
endothermic, and absorbing heat. If instead the quantity of heat is negative the result is an exothermic reaction
where heat is leaving the reaction system. The flow of heat must follow the first law of thermodynamics, no
energy may be created or destroyed during a physical process. For this to be true all heat transferred from a
reaction must be absorbed by the surroundings or vice versa. Mathematically we can define this as follows:
qsystem =-qsurroundings
When studying a chemical system, we often define the heat transfer using an energy term called
enthalpy. Enthalpy is a property inherent to a substance that allows us to relate heat absorbed or released by a
reaction in terms of moles of chemicals reacted. Its value can be related to q, and is unique to a reaction at a
given temperature and pressure. We do note the change in enthalpy as H. For heat and enthalpy to be directly
related a reaction must occur at a constant pressure. Once this is occurred we can write the change in enthalpy
as directly equal to heat as follows:
H = qp
Here the subscript p refers to the constant pressure requirement.
When the surroundings began to heat up, the speed and/or magnitude of the temperature change is
dependent upon the substance being heated. Each substance has a unique property of how it absorbs and
distributes heat. This property is known as the specific heat capacity (s). The specific heat capacity is the
quantity of heat required to raise the temperature of 1 g of the substance by 1°C at a constant pressure. As long
as pressure is constant we can relate the heat required to change the temperature using the specific heat
capacity. Combining specific heat, the mass of the substance and the temperature change the amount of heat
required is mathematically defined below.
q=msT
Where m is the mass of the substance in grams, S is the specific heat capacity of the substance in J/goC and T is
the change in temperature (a state function defined as Tf-Ti).
The type of chemistry you are performing today is known as calorimetry, and the device we will use to
measure the heat change is called a calorimeter. In a constant pressure scenario such as this lab is a simple
apparatus consisting of an insulated container (in our case a Styrofoam cup with a phone lid) a temperature
probe and a magnetic stir bar. This kind of device is often referred to as a coffee cup calorimeter. To determine
the heat of reaction we will mix to chemicals in the presence of water, and measure the change of temperature
of the water. Here the reaction is our system and the water is our surroundings. By determining how much heat
the water absorbs or releases we will determine how much energy is absorbed or released by the reaction. From
that, we will equate the heat released directly to enthalpy change H.

In order for this to be true, our solutions must have essentially the properties of water ( e.g. density and
specific heat capacity). This means a milliliter of our solution is assumed to weigh 1 g. As we are using relatively
low concentration solutions, this assumption holds. We will then be able to conduct our reaction in water, the
148
heat will transfer from the reaction to the water changing the temperature. Substituting into the equations
above, we see the following.
qsystem =-qsurroundings
qsystem=-mH2OsH2OTH2O
H = qsystem
Once we have this entropy, you should recognize that the units are only in joules. As stated earlier entropy is
measured in kilojoules per mole. In order for this value to represent the enthalpy of reaction, we must divide Q
by the number of moles of the reactant to obtain the proper value and units.
We will conduct this process on three reactions using constant pressure calorimetry. What is unique
about the situation is that two of the reactions can be manipulated to add up to the third one. According to
Hess's law the sum of the enthalpy of reactions for the two that add up to the third should be the same as the
enthalpy of reaction for the third reaction. Today you will confirm whether this law holds at least within the
realm of equipment we possess.
Experiment today we utilize three reactions:
(1) Solid sodium hydroxide dissolves in water to form an aqueous solution of ions.
NaOH(s) 
 Na+(aq) + OH–(aq) H1 = ?
(2) Solid sodium hydroxide reacts with aqueous hydrochloric acid to form water and an aqueous solution of
sodium chloride.
NaOH(s) + H+(aq) ) + Cl–(aq) 
 H2O(l) + Na+(aq) + Cl–(aq) H2 = ?
(3) Solutions of aqueous sodium hydroxide and hydrochloric acid react to form water and aqueous sodium
chloride.
Na+(aq) + OH–(aq) + H+(aq) ) + Cl–(aq) 
 H2O(l) + Na+(aq) + Cl–(aq) H3 = ?
Prelab Questions – To be completed as part of your prelab write-up
1. Verify that reactions one and three do in fact add up to reaction two. Show all your work
2. Using the information in the introduction of this experiment, calculate the heat (in J) of the reaction if 50.0
mL of HCl is added to 50.0 mL of NaOH in a coffee-cup calorimeter. The initial temperature for both
o
solutions is 23.2 C. At the end of the reaction after the data is graphed, the final temperature is determined
o
to be 38.5 C.
3. Using the enthalpy changes given below, calculate the enthalpy change for the following equation:
N2 (g) + 2 O2 (g)
using the following equations:
2 NO (g) → N2 (g) + O2 (g)
NO2 (g) → NO (g) + 1/2 O2 (g)
→
2 NO2 (g)
∆H = -180.0 kJ
∆H = 112.0 kJ
Procedure
Reaction 1
1. Obtain and wear goggles.
2. Connect the probe to the computer interface.
3. Prepare the computer for data collection by opening the file “18 Hess’s Law” from the Chemistry with
Computers folder.
4. Setup a calorimetry cup as shown in the figure above.
5. Measure out 100.0 mL of water into the Styrofoam cup.
6. Add a magnetic stirbar and stir the solution.
149
7. Lower the Temperature Probe into the solution.
8. Use a utility clamp to suspend a Temperature Probe from a ring stand as shown in Figure 1.
9. Weigh out about 2 grams of solid sodium hydroxide, NaOH, and record the mass to the nearest 0.01 g. Since
sodium hydroxide readily picks up moisture from the air, it is necessary to weigh it and proceed to the next
step without delay. CAUTION: Handle the NaOH and resulting solution with care.
10. Click on
to begin data collection and obtain the initial temperature, t1. It may take several seconds
for the Temperature Probe to equilibrate at the temperature of the solution.
a. After three or four readings at the same temperature have been obtained.
b. Add the solid NaOH to the Styrofoam cup.
c. Stir continuously for the remainder of the 200 seconds or until the temperature maximizes.
d. As soon as the temperature has begun to drop after reaching a maximum, you may terminate the trial
by clicking
.
11. Examine the initial readings in the table to determine the initial temperature, t1.
12. To determine the final temperature, t2, click the Statistics button,
The maximum temperature is listed in
the statistics box on the graph. Record t1 and t2 in your data table.
13. Rinse and dry (not with paper towels!) the Temperature Probe, Styrofoam cup, and stir bar.
14. Dispose of the solutions down the drain.
Reaction 2
15. Repeat Steps 3-14 using 100.0 mL of 0.50 M hydrochloric acid, HCl, instead of water. CAUTION: Handle the
HCl solution and NaOH solid with care.
Reaction 3
16. Repeat Steps 3-8, initially measuring out 50.0 mL of 1.0 M HCl (instead of water) into the Styrofoam
calorimeter. In Step 5, instead of solid NaOH, measure 50.0 mL of 1.0 M NaOH solution into a graduated
cylinder. After t1 has been determined for the 1.0 M HCl, add the 1.0 M NaOH solution to the Styrofoam cup.
CAUTION: Handle the HCl and NaOH solutions with care.
Processing Data
1. Determine the mass of 100 mL of solution for each reaction (assume the density of each solution is 1.00
g/mL).
2. Determine the temperature change, t, for each reaction.
3. Calculate the heat released by each reaction, q, by using the formula:
q = Cp•m•t (Cp = 4.18 J/g°C)
4. Convert joules to kJ in your final answer.
5. Find H (Remember, H = –q).
6. Calculate moles of NaOH used in each reaction.
a. In Reactions 1 and 2, this can be found from the mass of the NaOH.
b. In Reaction 3, it can be found using the molarity, M, of the NaOH and its volume, in L. Remember to
include the fact that it is being diluted by the addition of HCl.
7. Use the results of the Step 5 and Step 6 calculations to determine H/mol NaOH in each of the three
reactions.
8. To verify the results of the experiment, combine the heat of reaction (H/mol) for Reaction 1 and Reaction 3
and solve for H for these two reactions combined.
9. This sum should be similar to the heat of reaction (H/mol) for Reaction 2. Using the value you determined
in Reaction 2 as the accepted value and the sum of Reactions 1 and 3 as the experimental value find the
percent error for the experiment.
150
Data and Calculations
Reaction 1
Reaction 2
Reaction 3
1. Mass of solid
NaOH
g
g
2. Mass (total)
of solution
g
g
g
3. Final
temperature, t2
°C
°C
°C
4. Initial
temperature, t1
°C
°C
°C
5. Change in
temperature, t
°C
°C
°C
kJ
kJ
kJ
kJ
kJ
kJ
mol
mol
mol
kJ/mol
kJ/mol
kJ/mol
(no solid NaOH mass)
6. Heat, q
7. H
8. Moles of NaOH
9. H/mol
10. Experimental value kJ/mol
11. Accepted value
12. Percent error
kJ/mol
%
151
Name_____________________________________________Section___________
Experiment 14 Post Lab Sheet - Hess’ Law
This sheet is due before the start of lab one week after the lab is completed. Attach your carbons and
supporting material to the BACK of this sheet. Everything must be completed on this page in detail for
full credit.
H for Rxn 1 ____________________
H for Rxn 2 ____________________
H for Rxn 3 ____________________
Show the Chemical Equation manipulations to add reactions 1 and 3 to get reaction 2, and the math
calculations for comparing reactions 1 and 3 reactions to the 2nd.
152
Appendix A – Common Glassware
Erlenmeyer Flasks and Beakers
Erlenmeyer flasks and beakers are used for mixing, transporting, and reacting, but not for accurate
measurements. The volumes stamped on the sides are approximate and accurate to within about 5%.
Graduated Cylinders
Graduated cylinders are useful for measuring liquid volumes to within about 1% accuracy. They are for
153
general purpose use, but not for quantitative analysis. If greater accuracy is needed, use a pipet or
volumetric flask.
154
Appendix B – Using a Buret
A buret is used to deliver solution in precisely-measured, variable volumes. Burets are used primarily
for titration, to deliver one reactant until the precise end point of the reaction is reached.
With practice, the position of the meniscus of a liquid in the 50 mL burets used in the Chemistry
labs can be estimated to within 0.01 mL. Figure 10 shows the use of a card with a dark mark on it to
sharpen the image of the meniscus. You will find by experiment that if the top of the strip is positioned
slightly below the level of the liquid in the buret, the bottom of the meniscus will be very easy to see.
Figure 10: Reading a buret meniscus with a reading card.
You should always use the following procedure when changing the solution in a buret. First, empty the
buret out the top and half-fill it with deionized water. Open the stopcock and drain about 5 mL out of the tip.
Over the sink, empty the buret out the top by inverting it swiftly, and then repeat the water washing, this time
also opening the stopcock when the buret is inverted to allow most of the water to drain back out of the tip.
Wait about 30 seconds for drainage and then close the stopcock. While it is still upside down, blot/wipe off the
top of the buret with a laboratory tissue. Then turn it upright, and using a clean, dry beaker for the transfer, add
enough of the new solution to bring the liquid level up to about the 48 mL mark. Next, drain part of the liquid
out of the tip into a waste receiver, close the stopcock, and wipe off the tip with a laboratory tissue. Then, at the
sink, cradle the top of the buret between the thumb and index finger of one hand.
While holding it by the tip with your other hand, turn the buret horizontal. While twirling the buret by the tip,
slowly empty it through the top, being careful to wet the entire interior wall with the new solution. Repeat this
operation two more times. Finally, fill the buret above the zero mark and drain the excess out the tip until the
meniscus is within the calibrated portion of the buret. Be sure that no air bubbles are trapped in the tip. Do not
attempt to bring the meniscus to 0.00. This method is both time consuming and unwise, since the 0.00 line may
not be in precisely the right place.
Using a Buret
1. Clean the buret with soap and the long handled bottle brushes found in the lab. Rinse thoroughly with tap
water several times, including rinsing several milliliters of water through the stopcock and tip. Use a rolling
motion when pouring to ensure all the soap and dirt is rinsed out of the buret.
2. Rinse the buret several times with deionized or distilled water several times using a rolling motion to pour
the water out of the buret. This ensures that all the tap water is rinsed from the buret.
3. Before titrating, condition the buret with the titrant solution and check that the buret is flowing freely. To
condition a piece of glassware, rinse it so that all surfaces are coated with the solution to be used, and then
155
drain with the usual rolling motion. Conditioning two or three times will insure that the concentration of
titrant is not changed by a stray drop of water.
4. To fill a buret, close the stopcock at the bottom and use a funnel. You may need to lift up on the funnel
slightly, to allow the solution to flow in freely.
5. Check the tip of the buret for an air bubble. To remove an air bubble, flick the side of the buret tip while
solution is flowing. If an air bubble is present during a titration, volume readings may be in error.
6. Rinse the tip of the buret with water from a wash bottle and dry it carefully. After a minute, check for
solution on the tip to see if your buret is leaking. The tip should be clean and dry before you take an initial
volume reading.
156
7. For the purposes of this lab, you should fill the buret over the 0.00 mL mark and carefully drain it until it is
exactly 0.00 mL. This will make entering computer data much easier and faster. In practice however, this is
generally a bad technique for technically specific labs (e.g. where you are measuring patient’s blood
chemistry). The calibrated portion of the buret is only BETWEEN the 0 and 50 mL marks, not outside of
them. In this lab where we are stressing technique over precision, the small error created by this is
inconsequential.
8. When your buret is conditioned and filled, with no air bubbles or leaks, take an initial volume reading. A
buret reading card with a black rectangle can help you to take a more accurate reading. Read the bottom of
the meniscus. Be sure your eye is at the level of meniscus, not above or below. You can ensure precision by
using the circular marks on the buret near the meniscus. Using the closest ring to the meniscus, move your
eye up or down until the ring becomes a line. You are now at the same level as the meniscus and the
bottom will line up with the measurement mark. Reading from an angle, rather than straight on, results in a
parallax error.
9. Deliver solution to the titration flask by turning the stopcock. The solution should be delivered quickly until a
couple of mL from the endpoint.
157
158
Appendix C – Conducting a Titration
Titration
A titration is a method of analysis that will allow you to determine the precise
endpoint of a reaction and therefore the precise quantity of reactant in the titration
flask. A buret is used to deliver the second reactant to the flask and an indicator or
pH Meter is used to detect the endpoint of the reaction.
Doing a Titration
Begin by preparing your buret, as described on
the
Appendix B – Using a Buret page. Your buret
should be conditioned and filled with
titrant solution. You should check for air
bubbles and leaks, before proceding with
the titration.
Take an initial volume reading and record it in your notebook. Before beginning a
titration, you should always calculate the expected endpoint volume.
159
Prepare the solution to be analyzed by placing it in a clean Erlenmeyer flask or
beaker. If your sample is a solid, make sure it is completely dissoloved. Put a
magnetic stirrer in the flask and add indicator.
Use the buret to deliver a stream of titrant to within a couple of mL of your
expected endpoint. You will see the indicator change color when the titrant hits the
solution in the flask, but the color change disappears upon stirring.
160
Approach the endpoint more slowly and watch the color of your flask carefully. Use
a wash bottle to rinse the sides of the flask and the tip of the buret; be sure all
titrant is mixed in the flask.
As you approach the endpoint, you may
need to add a partial drop of titrant. You
can do this with a rapid spin of a teflon
stopcock or by partially opening the
stopcock and rinsing the partial drop into
the flask with a wash bottle.
Make sure you know what the endpoint should look
like. For phenolphthalein, the endpoint is the first
permanent pale pink. The pale pink fades in 10 to 20
minutes due to the addition of CO2 from the
atmosphere.
As you are generally collecting endpoints not by
indicator, but by pH or conductivity you will
continue adding titrant until your graph levels out
flat again. The indicator however serves as visual
proof that you are making progress and allows you
to recognize problems with the data collection
deviecs used in lab.
161
If the flask looks like this, you have
gonepast the endpoint.
When you have reached the endpoint,
read the final volume in the buret and
record it in your notebook.
Use the computer to determine how
much titrant was added. For more on this
practice see the Using LoggerPro
appendix.
For more information on burets, see
Appendix B – Using a Buret.
Titrating with a pH meter
162
Titration with a pH meter follows the
same procedure as a titration with an
indicator, except that the endpoint is
detected by a rapid change in pH, rather
than the color change of an indicator.
Arrange the sample, stirrer, buret, and
pH meter electrode so that you can read
the pH and operate the buret with ease.
The computer will detect the pH
accurately and allow you to find the
endpoint. This is best conducted through
the use of the second derivative method.
See Logger Pro use in the titration labs
for more on how to do this.
163
Appendix D – Using the Balances
A balance is used to measure the mass of an object. Each laboratory room contains two electronic
balances that are very easy to use. To use the balance, turn it on by pushing the tare bar down. The electronic
readout should then be lit. Open one of the sliding doors and be sure the balance pan and surrounding area is
clean. You can clean it with a balance brush or KimWipe. Next shut the doors and press the tare bar to set the
balance at zero. Now simply place the object to be weighed on the balance and measure the mass to 0.001
grams.
Always use weighing paper when weighing solids to protect the balance. To do this simply place the weighing
paper on the balance pan and be sure it is not touching the side. Press the tare bar on the right side and the
balance will then read 0.000 g. Now add the desired mass of solid and record the mass. Always clean the balance
carefully after use. At the end of the period, turn off the balance by pressing the on/off button. Always use the
balance with extreme care as it is very expensive.
Top-loading Balance
Use a top loading balance to weigh solid material when a precision of 0.001 g is adequate. For more
accurate mass measurements or small amounts, use a precision analytical balance.
1. Check if the balance is turned on. If not, press the on/off button and wait until the display reads 0.000 g. In
practice it is best to let a balance sit for 10-15 minutes to warm-up for the best accuracy.
2. Place a container or large, creased weighing paper on the balance pan. Push tare button to zero the balance.
164
3. Carefully add your substance to the container or paper and record the mass.
4. Use the brush provided to clean any spills. Discard any disposable tare containers, weighing paper, or
Kimwipes in the nearest wastebasket.
Precision Analytical Balance
A precision analytical balance measures masses to within 0.0001 g. Use these balances when you need
this high degree of precision.
1. Turn the balance on by pressing the control bar. The display lights up for several seconds, then
resets to 0.0000.
165
2. Place creased, small weighing paper on the balance pan.
3. Close the sliding glass doors. Wait for the green dot on the left to go out. This is the stability
indicator light, indicating that the weight is stable.
4. Press the control bar to cancel out the weight of the container or paper. The display will again
read 0.0000.
166
5. Carefully add the substance to be weighed up to the desired mass. Do not attempt to reach a
particular mass exactly.
6. Before recording the mass, close the glass doors and wait until the stability detector lamp goes
out. Record mass of solid.
7. There are several things you shouldn’t do using this balance in particular:
a. Don't pick up tare containers with bare hands since your fingerprints add mass. Use
Kimwipes or tongs to prevent this.
b. Don't lean on the bench while weighing.
c. Do record the mass of your container, if you will need it later.
d. Do check the level indicator bubble before weighing. The two rear balance feet serve as
leveling screws.
8. Use the brush provided to clean spills in the weighing chamber. Discard any disposable tare
containers, weighing paper, or Kimwipes in the nearest wastebasket.
167
Appendix E – Handling Chemicals
Handling Solids
Use a clean spatula to transfer solid from bottles. Never use a contaminated spatula. Also, never
return unused solid to the reagent bottle. Simply discard it or pass it on to someone who needs it. To
avoid waste, never remove more solid from a bottle than is necessary. See Error! Reference source not
ound. for more information on weighing and transferring solids.
1. If you are creating a solution, you should rinse the wax paper you use to weigh objects with a
distilled water bottle to ensure all solid is transferred into the container.
2. If the solid is not to be used in a solution it is best to weigh the solid directly in the container it
will be used in if possible. This minimizes loss of solid that might have stuck to the wax paper.
Figure 11: Using Wax Paper to Transfer a Solid
Handling Liquids
When transferring liquids from a reagent bottle, always remove the cap/stopper and hold it in your
hand so that the chemical side faces out. Never place the cap/stopper on the bench or contamination could
result. Pour the liquid slowly and carefully to avoid spillage. You may find the use of a glass rod helpful in
pouring from larger containers. It should be placed over the spout and the liquid allowed to drain down the rod
into your container.
Figure 12 below indicates the proper technique for pouring a fluid from a beaker to another
vessel.
168
Figure 12: Proper technique for transferring a liquid.
Capping a Flask
During many experiments you will have to cap a flask to protect the contents from contamination.
Figure 13 illustrates the proper method using Parafilm. Note: When the flask contains an
organic solvent you should never use a rubber stopper for storage as the solvent will deteriorate the
rubber. Additionally, cork stoppers are usually appropriate for dry chemical storage or those chemicals
that are non-volitile.
Figure 13: Covering a flask for temporary storage using parafilm
169
Appendix F – Filtration
You will often need to separate a liquid from a solid. At times you will simply decant, that is, you
will carefully pour out the liquid, leaving the solid behind. At other times you will need to filter the
solution. To do this you will use filter paper and a funnel.
There are two primary methods of filtration utilized in an undergraduate chemistry lab: Gravity
filtration and vacuum filtration. Gravity filtration is primarily used when it is desired that the liquid
saved; whereas vacuum filtration is primarily used to remove an insoluble solid from an unwanted
liquid. The distinction may sound simple, but the ramification on chemical yield can be very large.
Under vacuum many liquids are more volatile and will evaporate, lowering the yield or increasing the
concentration of dissolved solutes.
Gravity Filtration
Paper will stick to smooth glass or plastic funnels reducing the porosity of the paper. This makes
it harder for any substances to pass through the filter. In order to speed this process up, filter discs
should be fluted before use.
170
Figure 14: Fluting filter paper for gravity filtration.
You will then set the paper in the funnel using your wash bottle. To do this simply place the
paper into the funnel and add a small amount of water to the bottom of the filter. Slowly add water to
the sides with a circular motion to avoid air bubbles between the paper and the funnel. Once the paper
has set, transfer the solution to be filtered. If the solid has settled, decant the liquid through the filter
first in order to save time. Never overwhelm the filter; don't add the solution too quickly and never
come to within one centimeter of the top of the paper. Transfer the solid using a wash bottle and rubber
policeman, and then wash the solid as directed by the experimental procedure.
Vacuum Filtration
Vacuum filtration is a technique for separating a solid product from a solvent or liquid reaction
mixture. The mixture of solid and liquid is poured through a filter paper in a Buchner funnel. The solid is
trapped by the filter and the liquid is drawn through the funnel into the flask below, by a vacuum.
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1. To prepare for a vacuum filtration, gather together a filter flask, Buchner funnel, tubing, filter
paper, clean solvent, disposable dropper, and your sample.
2. In the general chemistry lab, you will use the faucet-based vacuum outlets on the sinks or the
Brinkmann Aspirators. You will find tubing in the chemical equipment drawers. Ideally, the
tubing should be connected from the vacuum nozzle to a solvent trap and from the trap to the
side arm of your filter flask. This prevents your solvent from being sucked into the pump.
3. Turn on the vacuum by switching on the aspirator or by turning on the water in a sink with a
white aspirator attachment. Check the vacuum by feeling for suction at the end of your tubing.
The vacuum should be strong enough to hold the tubing to your finger without falling off.
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4. Connect the tubing to the side arm of your filter flask and check the suction at the top of the
flask (left). Place the black rubber ring adapter in the top of the flask and then the Buchner
funnel. Check again for good suction by placing your gloved hand across the top of the funnel
(right). If you do not feel strong suction, there is a poor connection and a leak somewhere in
your system. DO NOT TRY THIS WHEN YOUR SOLID IN THE FUNNEL! You are likely to spread the
solid all over the bench and yourself when you try to break the vacuum.
5. Prepare to filter your sample by placing a filter paper in the Buchner funnel and wetting it with
clean solvent. You should see the paper being sucked down against the holes in the funnel and
the solvent should quickly pass through into the filter flask.
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6. To filter your sample, slowly pour into the center of the filter paper.
7. Use more clean solvent to rinse your beaker, so that the entire solid is collected.
8. Rinse the solid on the filter paper with more clean solvent. Continue to draw air through the
solid, to evaporate any remaining solvent in your sample.
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9. When you are finished, break the vacuum at the connection between the flask and the trap.
Then turn off the vacuum. Switching off the vacuum before breaking the seal often sucks water
back into the trap or flask where your product is.
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Appendix G – Calorimetry
Calorimetry
Calorimetry is used to determine the heat released or absorbed in a chemical reaction. The
calorimeters shown here can determine the heat of a solution reaction at constant (atmospheric)
pressure. The calorimeter is a styrofoam cup fitted with a plastic top in which there is a hole for a
thermometer. (It's crude, but very effective!) Key techniques for obtaining accurate results are starting
with a dry calorimeter, measuring solution volumes precisely, and determining change in temperature
accurately.
Using a Calorimeter
Solutions volumes should be carefully measured with a graduated cylinder. Add solution completely, to
a dry calorimeter. Don't forget to add the stir bar each time!
Set up the calorimeter with the stainless steel temperature probe supported from a stand so that the
bulb does not touch the bottom of the cup. Clamp the calorimeter so that it rests on the stirrer. Be
careful not to turn on the heat or you will melt the styrofoam.
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The change in temperature is determined by measuring the initial temperature, T1, of the reactants,
and the maximum temperature, T2, of the contents of the calorimeter during the exothermic reaction.
Use a magnifying glass to measure temperature values precisely.
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Appendix H – Using a Centrifuge
Centrifugation
A centrifuge separates a heterogeneous mixture of solid and liquid by spinning it. After a successful
centrifugation, the solid precipitate settles to the bottom of the test tube and the solution, called the
cetrifugate, is clear.
Using a Centrifuge
Place test tube in centrifuge holder.
Balance with another test tube filled to the
same level in the opposite holder.
Close cover and turn knob. Centrifugation takes
a minute or more. Note that you must turn off
the centrifuge with the switch and wait for it to
stop spinning, to effectively separate the
precipitate and solution.
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Appendix I – Heating
Heating
You will use both a hot plate and a Bunsen burner to heat solids and solutions. Always be
careful to avoid burns and never heat a material too quickly or explosive "bumping" can occur.
Hot plates
When using a hot plate always begin at the setting indicated in the manual. However, this
setting may vary depending on the hot plate so you will have to experiment.
Bunsen Burners
In using a Bunsen burner, always use a tight blue flame as shown in Figure 15. Control the heat
transfer by adjusting the distance from the burner to the object. Note that the distances suggested in
the manual are measured from the hottest part of the flame to the object.
Figure 15: Parts and proper use of a Bunsen burner.
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