CHEM 1361
General Chemistry I
Laboratory Manual
August 2013
10th Edition, 1st Printing
Cameron University
Department of Physical Sciences
Lawton, OK 73505
Dr. Gary S. Buckley
Dr. T. E. Snider
A New Twist in Some of These Experiments
This book is the first of its many editions to include some experiments using modern technology for data
acquisition. Such instrumentation will allow you to acquire measured data more accurately and quickly,
to make judgments related to experimental design on a timely basis, and to do some basic data analysis as
the experiment develops.
This equipment was provided through the use of Cameron University technology fees paid by Cameron
University students each semester.
Acknowledgements
All revenue generated from the sale of this lab manual is used for equipment, scholarships, and research
materials for the Department of Physical Sciences. The materials in this book are original. A debt of
gratitude is owed to Mr. Chris Holtz (Cameron instructor from 1992 -1994) whose earlier efforts were
freely used with his permission. Also Department of Physical Sciences faculty including Dr. Ted Snider
(emeritus), Dr. Jimmy Stanton (emeritus), Dr. Ann Nalley, Dr. Keith Vitense, Dr. Kurtis Koll, Dr. Clint
Bryan, Dr. Danny McGuire, Dr. P. K. Das, Dr. Kyle Moore, and Becky Eden who contributed ideas,
corrections, and even entire experiments to this collection of experiments. Many of these experiments are
widely used in lab books across the country and have been adapted to comply with modern safety
regulations and chemical disposal regulations.
In June of 2010 the Board of Regents of the University of Oklahoma, Cameron University, and Rogers
State University approved the creation of the Frontiers in Chemistry Endowed Lectureship. Funding for
this lectureship is provided through the sales of these lab books. In addition to this funding the Oklahoma
State Regents for Higher Education will match this contribution, multiplying the benefits to students
throughout the years.
A special note of thanks goes to Dr. Ted Snider whose imagination, creativity, and initiative resulted in
the development of this lab book as well as several others throughout the Department of Physical
Sciences. Though Dr. Snider retired in May of 2008 after 40 years of service to Cameron University and
the Department of Physical Sciences, the revenue from this effort will continue to benefit students well
into the future. True to form, Dr. Snider returned the Fall of 2010 to teach and considerably improved
Experiment 15 by suggesting the introduction of a dye to improve visibility.
Thank you to Cameron University Printing Services for the production of this manual.
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Table of Contents
Introductory Notes .................................................................................. 4
Feedback Page......................................................................................... 5
Laboratory Equipment ............................................................................ 7
Departmental Laboratory Policies .......................................................... 8
#1 Laboratory Safety ................................................................................
#2 Measurements, Accuracy, and Precision ............................................
#3 Sugar in Soft Drinks and Fruit Juices ................................................
#4 Using Physical Properties to
Determine the Identity of an Unknown ..............................................
#5 Separation of a Mixture ......................................................................
#6 Determination of an Empirical Chemical Formula ............................
#7 Molecular Modeling............................................................................
#8 Preparation of an Alum .......................................................................
#9 Metathesis Reactions ..........................................................................
#10 Molar Stoichiometry in a Chemical Reaction ..................................
#11 Determination of Acetic Acid in Vinegar .........................................
#12 Comparison of the Energy Content of Fuels by Combustion ...........
#13 Determination of the Enthalpy of Fusion of Ice ...............................
#14 Molar Mass of a Volatile Liquid by the Dumas Method ..................
#15 Regularities in the Properties of the Elements ..................................
Appendix A-1 Basic Error Analysis .................................................. 125
Appendix A-2 Generally Useful Lab Information............................. 126
Appendix A-3 Common Ions. ............................................................ 129
Appendix A-4 Basic Use of the Vernier LabQuest 2 ........................ 133
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Introductory Notes
A student successfully completing the General Chemistry I lab, CHEM 1361, should be able to:
1. Work safely with chemicals and equipment and practice chemical disposal according to designated
protocol.
2. Record data to the proper number of significant figures and properly carry those significant figures
into a final calculated result.
3. Satisfactorily perform routine laboratory techniques, including but not limited to; following both
written and oral instruction; measuring mass; measuring volume accurately in a variety of vessels;
lighting and adjusting burners; filtering; and conducting titrations.
4. Analyze acquired data through given instructions as well as through knowledge gained in the
lecture portion of the class.
5. Use a software program, typically Excel, to make graphs, evaluate trendlines, and apply those
results to experimental data.
6. Evaluate at a basic level both quantitatively and qualitatively the errors associated with
measurement.
One of the best ways of having a successful laboratory experience is to enter the laboratory prepared for
the day’s experiment. To aid in your preparation, each laboratory has a prelaboratory exercise that is
deployed through Blackboard. Your instructor will provide more information on access to Blackboard
and will also provide the due dates for these prelabs. Each prelab is based heavily upon the experiment
you are about to do that week and will provide the means for you to start each week’s experiment
effectively. You will be successful at the prelabs if you read the laboratory, have your lab book with you
when you take the prelab, and pay attention to the Learning Objectives at the start of each lab and the
associated textbook sections.
On occasion a laboratory will appear to not have a direct link to what you are doing in the lecture portion.
This may be either because the laboratory is dealing with a set of skills or topics that we do not cover
explicitly in lecture or perhaps the laboratory and lecture have temporarily lost their synchronization. The
materials and references in the lab book should be sufficient to work you through those times.
Finally, if you would like to provide feedback for future versions of this lab book you will find a form at
the start of the book to provide such feedback. It is probably most helpful if you fill it out each week. At
the end of the semester we will make available an envelope collect these forms. Your name is not
required.
If at any time during the semester I can be of assistance or you find errors in the lab book please feel free
to contact me or let your instructor know.
Gary S. Buckley
e-mail: gbuckley@cameron.edu
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Feedback Page
Your feedback would be helpful as we continually strive to improve this laboratory manual. Please take a
few minutes as you complete and write up experiments to share on this feedback page corrections,
suggestions for improvements in the clarity of instructions, ideas for other experiments, comments on
prelabs – whatever you wish. Keep this log as you go through the semester and at the last lab there will
be an envelope you can place this sheet in anonymously. Use additional sheets if necessary. Thanks for
your help. Dr. Buckley
Experiment
Exp. 1 – Laboratory Safety
Corrections, suggestions, etc.
Exp. 2 – Measurements, Accuracy,
and Precision
Exp. 3 – Using Physical Properties to
Determine the Identity of an
Unknown
Exp. 4 – Sugar in Soft Drinks and
Fruit Juices
Exp. 5 – Separation of a Mixture
Exp. 6 – Determination of an
Empirical Chemical Formula
Exp. 7 – Preparation of a Alum
Exp. 8 – Metathesis Reactions
Feedback Page
-5-
Exp. 9 – Molar Stoichiometry in a
Chemical Reaction
Exp. 10 – Determination of Acetic
Acid in Vinegar
Exp. 11 – Determination of the
Enthalpy of Fusion of Water
Exp. 12 – Comparison of the Energy
Content of Fuels by Combustion
Exp. 13 – Regularities in the
Properties of Elements
Exp. 14 – Molecular Modeling
Exp. 15 – Molar Mass of a Volatile
Liquid by the Dumas Method
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Feedback Page
Laboratory Equipment (not to scale)
General Use Glassware:
←Griffin beakers
– you will have 50
mL, 100 mL, 150
mL, 250 mL, and
400 mL Griffin
beakers in your
lab drawer
←Erlenmeyer
flasks – you
will have three
125 mL
Erlenmeyer
flasks in your
drawer
←Test tubes –
you will have six
(all the same size)
in your drawer.
Size is measured
as width x length
in mm.
Test tubes - measured
by with of mouth X length
Volume Measurement:
←Graduated cylinders
– you will
have 10 mL and 50 mL
graduated
cylinders in your
drawer
←Pipette – will be
distributed as
needed. Various
volumes and types
available.
←Burette – will be
distributed as needed.
We typically use 50mL burettes which
stand about 60 cm tall.
Other Drawer Items:
←Plastic
funnel for
separation of
solids from
liquids
←Glass stirring rod
←Watch
glass
←Metal spatula
←Test tube brush
Equipment in Common Areas:
← Jack Stand
←Bunsen
burner
← Assorted
clamps
← Hot plate or
hot plate/stirrer
← Metal rods
for insertion into
holes on table
tops
Laboratory Equipment
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CHEMISTRY 1361
LABORATORY POLICIES SC Room 217
You are each responsible for the following procedures and schedules. Keep this copy in your laboratory notebook
for future reference. Sign and turn in to your instructor the copy of this form on page 9.
1. Attendance is required for each laboratory session.
Be on time. Attendance records are kept. Tardiness 14. If the instructor is talking, students should be
is not permitted. No results, data sheets, etc., are
listening.
accepted from those absent.
15. Arrive prepared for each laboratory. Read the
2. Goggles are required and must be worn in the
assigned laboratory material before laboratory.
laboratory. You must wear long pants or dresses to
lab. (No shorts, No sandals, No flip-flops. If you
16. Each week safety and experimental technique will
are not properly dressed you will not be allowed
be detailed.
entry into the Lab.)
17. No food or drinks in the laboratory.
3. Each student is responsible for adhering to all safety
rules listed below and all safety instructions given at 18. Dispose of materials in the designated containers.
the start of each lab.
19. No solids should be thrown in the sinks.
4. Generally, each student will have a laboratory
partner. Each pair of experimenters is required to
20. Do not sit on the laboratory tables.
maintain an assigned lab locker/space, return clean,
used glassware to the lab locker assigned, return
21. Never pour excess chemicals back in the storage
hardware to the correct storage drawer, and clean
bottle.
their work area. Make sure that all equipment is
unplugged before you leave the laboratory.
22. Your work area must be clean and all equipment
returned to its proper place before you are dismissed
5. Before a laboratory grade is assigned, you must have
from the laboratory.
a passing grade assigned in the chemistry lecture
course (departmental policy).
23. All spills must be cleaned up immediately.
6. No “horseplay” in the laboratory.
7. Your laboratory instructor will pass out a grading
and attendance policy at the beginning of term. In
addition, s(he) will delineate the order of the labs,
point value, notebook requirements and any other
requirements that will affect your grade.
24. The required text is a lab manual prepared by Dr. T.
E. Snider and Dr. Gary S. Buckley. Monies earned
from the sale of this book are used by the department
to fund scholarships and other academic activities.
8. All pre-lab and post-lab assignments and lab reports
are due as your lab instructor specifies.
9. NO PHOTOCOPIES. Lab exercises are available
in the bookstore.
10. Always show your work when performing
calculations.
11. You alone are responsible for your clothing and
personal property.
12. Store nothing of value in lab lockers.
13. Make sure that you mix the correct chemicals as
instructed in the Laboratory Manual.
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Laboratory Policies
CHEMISTRY 1361
LABORATORY POLICIES SC Room 217
You are each responsible for the following procedures and schedules. Read and initial each of the rules on this
page and submit to your instructor. Please include your name, instructor, and lab time in the lower right hand
corner.
1. Attendance is required for each laboratory
13. Make sure that you mix the correct chemicals as
instructed in the Laboratory Manual.
session. Be on time. Attendance records are
kept. Tardiness is not permitted. No results
data sheets, etc., are accepted from those absent. 14. If the instructor is talking, students should be
listening.
2. Goggles are required and must be worn in the
laboratory. You must wear long pants or dresses to
Lab. (No shorts, No sandals, No flip-flops. If you
are not properly dressed you will not be allowed
entry into the Lab.)
15. Arrive prepared for each laboratory. Read the
assigned laboratory material before laboratory.
16. Each week safety and experimental technique will
be detailed.
3. Each student is responsible for adhering to all safety
rules listed below and all safety instructions given at 17. No food or drinks in the laboratory.
the start of each lab.
18. Dispose of materials in the designated containers.
4. Generally, each student will have a laboratory
19. No solids should be thrown in the sinks.
partner. Each pair of experimenters is required to
maintain an assigned lab locker/space, return clean,
used glassware to the lab locker assigned, return
20. Do not sit on the laboratory tables.
hardware to the correct storage drawer, and clean
their work area. Make sure that all equipment is
21. Never pour excess chemicals back in the storage
unplugged before you leave the laboratory.
bottle.
5. Before a laboratory grade is assigned, you must have 22. Your work area must be clean and all equipment
a passing grade assigned in the chemistry lecture
course (departmental policy).
returned to its proper place before you are dismissed
from the laboratory.
6. No “horseplay” in the laboratory.
23. All spills must be cleaned up immediately.
7. Your laboratory instructor will pass out a grading
24. The required text is a lab manual prepared by Dr. T.
and attendance policy at the beginning of term. In
addition, s(he) will delineate the order of the labs,
point value, notebook requirements and any other
requirements that will affect your grade.
8. All pre-lab and post-lab assignments and lab reports
are due as your lab instructor specifies.
9. NO PHOTOCOPIES. Lab exercises are
available in the bookstore.
10. Always show your work when performing
calculations.
E. Snider and Dr. Gary S. Buckley. Monies earned
from the sale of this book are used by the department
to fund scholarships and other academic activities.
Student Name:
__________________________
Instructor:
______________________________
Lab Day (Circle one): M T
W
R
F
Lab Time: __________________________
11. You alone are responsible for your clothing and
personal property.
12. Store nothing of value in lab lockers.
Laboratory Policies
-9-
This page is intentionally left blank.
Experiment #1: Laboratory Safety and Practice
Objectives: The information in this laboratory is intended to acquaint you with basic safety procedures,
more general regulatory safety responsibilities, and laboratory practice in our labs. By the time you
successfully complete this laboratory, you will be able to:






I.
Identify and use the mandatory and optional safety equipment used in the CHEM 1361 lab
Identify your safety responsibilities toward your partner and others in the laboratory
Extract relevant safety information regarding chemicals to be used from Safety Data Sheets
(SDS), formerly called Material Safety Data Sheets (MSDS)
Apply the “Right-to-Know” law to different situations
Identify the significance of the National Fire Protection Association (NFPA) diamond and its
markings
Recognize the new Globally Harmonized System pictograms
Personal Safety Obligations
Though laboratory safety is not all about you, it certainly begins with you. Three particular aspects
will be considered – personal protective equipment, preparation, and awareness.
a.
Personal Protective Equipment (PPE)
There are four basic routes through which materials can enter your body:




Eye Contact
Skin contact/absorption
Ingestion
Inhalation
Eye contact is easily prevented by using protective goggles, a
requirement in this laboratory at ALL times. The only
acceptable eye protection in this laboratory is splash goggles.
The term “splash” comes from the vents located around the
goggles – a material splashed in the face cannot penetrate into
the eyes. Goggles with simple perforations are not acceptable
since a splash in the face could penetrate that type of
Figure 1.1 Acceptable Safety Goggles
goggles and reach the eyes. Note that contact lenses and
prescription glasses are acceptable in the lab but must be worn in conjunction with splash
goggles.
Even with best practice, it is important to realize that once you leave the laboratory an
important first step is to wash your hands thoroughly. During the course of an experiment
you may come into contact with chemicals that could be irritants to the eyes. Once you
leave the lab and take off your goggles, be aware that you may still carry some of those
chemicals on your hands. If you were to rub your eyes without washing your hands first,
you could quickly undo all of your goggle diligence.
If you do splash something into your eyes or feel as though your eyes have been
contaminated, know the location of the eye wash station (northwest corner of the lab – near
the exit door on the right). Using the eyewash station is simply a matter of pushing the
handle and immersing your open eyes in the stream of water. In the case of a large splash
across the face, it is best to put your face in first with goggles on until the bulk of the spill
Experiment #1: Laboratory Safety and Practice
-11-
is removed. Then remove the goggles and begin the open eye wash. Practice indicates
that, regardless of the cause of the irritation, once started the open eye wash should be
continued for at least 15 minutes.
Skin contact/absorption can occur through many avenues in the lab. The best prevention
comes in the form of adequate coverage of the skin, thus the need to set forth some
guidelines for acceptable clothing:

Closed toe shoes must be worn in the laboratory. This includes tennis shoes, gym
shoes, dress, shoes, etc. Flip flops, sandals, and any shoe exposing a part of the
foot are not acceptable and will result in you being required to leave the lab.
Pants or dresses must be worn in the laboratory. Absolutely no shorts. It is
acceptable to bring (or store in your locker) sweat pants to slip into for the lab.
As much of the upper body as reasonable must be covered. The implication here is
that mid-riff tops or any sort of clothing that exposes the torso is not acceptable.
Tank tops are also not acceptable lab attire.
One way to help protect your clothes in the laboratory is to bring a lab coat. These
are optional. If you do choose to wear a lab coat, it is good practice to either store
it in your locker from week-to-week or be sure to wash it regularly. Remember it
will carry on it anything that you spilled on it during the lab period.



Another important avenue of preventing the absorption of chemicals into your body is to
use protective gloves. The department supplies two types of protective gloves for your use
– latex and nitrile. Some individuals are allergic latex or nitrile. If you are in that
category, please be aware of which gloves are which (they are in labeled boxes). If you are
allergic to BOTH latex AND nitrile, please notify your instructor as soon as possible in the
semester so other gloves can be made available.
There are a few things that are helpful to know about the use of gloves. The chemicals you
will work this semester are not particularly hazardous, but it is helpful to develop good
habits early.

Not all gloves are created equal. Different glove materials have different
compatibilities with a range of chemicals. The gloves are categorized by
degradation, breakthrough rate, and permeation. An example of this sort of
information may be found at
http://www.ansellpro.com/download/Ansell_7thEditionChemicalResistanceGuide.pdf.


-12-
If you are working in the laboratory wearing gloves, do not leave the laboratory
with the gloves on. Since you are trying to protect yourself from contamination by
wearing them, it is not wise to spread any potential contamination outside of the
confines of the laboratory. Remove them and dispose of them before leaving the
room. Do not reuse protective gloves.
There are several acceptable procedures for removing gloves. Remember, they are
liable to be contaminated with something you do not want on your skin. So just
stripping them off with your bare hands will likely lead to the contamination you
have been trying to avoid. One method is as follows:
o Use the thumb and index finger on one of your gloved hands to grasp the
palm of the glove on your other hand.
o Peel the grasped glove off of its hand.
o Once the glove is off, wad it up with your gloved hand and place it in the
palm of your other gloved hand.
Experiment #1: Laboratory Safety and Practice
o Slide a bare finger from your newly ungloved hand under the bottom of the
glove on the still-gloved hand and peel the glove off with the finger – don’t
touch the outside of the glove.
o Pull the glove off in this fashion until you have turned this glove completely
inside out. The other glove will be in the inside of the just-removed glove.
Dispose of the gloves in the trash.
Ingestion is a third avenue through which chemicals can enter your body during the lab. This area is
covered rather directly by the rules that include no smoking, eating, or drinking in the lab. Again, though,
be aware that washing your hands immediately after leaving the lab is an important habit to acquire since
your hands could still contain something that could be ingested.
Inhalation is not a common avenue of entry in the CHEM 1361 lab. During some experiments in which
noxious fumes could be given off, we will work almost exclusively in the fume hoods to prevent the
vapors from entering the room.
II.
Lab Community –wide Safety Considerations
You will be working in a laboratory setting that could have as many as 27 other students working at
the same time in rather close quarters. Though Section I dealt with your personal safety, it is also
important to recognize your role in the safety of others in the laboratory.
Starting with your obligations to your partner, it is essential to recognize that your ability as a pair to
work safely and efficiently in the laboratory is entirely dependent on having both of you come in to
the laboratory prepared. Though getting an experiment “done” in the time allotted is important, it is
not the ultimate goal. This is an instructional laboratory – the goal is for both members of a pair to be
active participants in the laboratory. It is absolutely essential that both members of a pair come in
prepared to conduct the experiment with a full awareness of safety concerns as well as an
understanding of the objectives of the laboratory. You may have heard the expression – “There is no
such thing as a bad question”. In the lab, there is one bad question. When you enter the lab to do an
experiment, do not ask “Which experiment are we doing today?”. You need to be prepared.
In the larger lab community, there are a few things you can do to help ensure a smooth safe lab
experience for everyone. In particular:





Arrive at the lab on time so you can hear the prelab comments. These often involve group
information such as the location of chemicals and equipment, a listing of safety concerns, ideas for
efficiently completing the lab, disposal guidelines, etc. Late arrivals cause a huge disruption to the
lab as the tardy arriver has to constantly ask questions that were covered in the prelab comments.
Be cautious as you move around the lab – look where you are going. There are often stools in the
rows and they can be difficult to see in a full lab.
Pay particular attention to disposal directions. There is typically one fume hood designated as the
disposal hood. Follow the directions in the hood to properly dispose of chemicals. This does fall
under a safety concern since anything that is put down the sink – which will be nothing in this
laboratory – ends up in the environment eventually.
During the safety orientation lecture at the start of the semester pay close attention to the location
and directions for the eyewash station, safety shower, fire blanket, fire extinguisher, and
broom/dust pan for sweeping up broken glass. In the event of an emergency, your role may be to
help someone get to one of these safety devices.
The emergency phone number is posted near the exit door: 581-2911. This is the first number to
call as it brings campus security officers who are trained in first response. They can then direct in
other help as necessary.
Experiment #1: Laboratory Safety and Practice
-13-
III.
Locating Chemical Information
One must be aware of the hazards of particular chemicals in order to use the proper personal
protective equipment, methods, protocol in case of exposure, and disposal methods. This information
is contained in documents called Safety Data Sheets (SDS). These were previously called Material
Safety Data Sheets (MSDS) but the name is changing effective June 1, 2015 due to the
implementation of the Globally Harmonized System of Classification and Labelling of Chemicals
(GHS). The GHS is a UN-devised and endorsed means of standardizing hazard communication
internationally. It replaces several individual national systems and is intended to improve
communications.
There are sixteen sections to the SDS sheets as shown in the table below.
Table 1.1 Safety Data Sheet Components
Short Description
includes product identifier; manufacturer or distributor name,
address, phone number, emergency phone number;
recommended use; restrictions on use
2
Hazard(s) identification
includes all hazards regarding the chemical; required label
elements
3
Composition/information includes information on chemical ingredients; trade secret
on ingredients
claims
4
First-aid measures
includes important symptoms/effects, acute, delayed; required
treatment
5
Fire-fighting measures
lists suitable extinguishing techniques, equipment; chemical
hazards from fire
6
Accidental release
lists emergency procedures; protective equipment; proper
measures
methods of containment and cleanup
7
Handling and storage
lists precautions for safe handling and storage, including
incompatibilities
8
Exposure
lists OSHA’s Permissible Exposure Limits (PELs); Threshold
controls/personal
Limit Values (TLVs); appropriate engineering controls; PPE
protection
9
Physical and chemical
lists the chemicals characteristics
properties
10
Stability and reactivity
lists chemical stability and possibility of hazardous reactions
11
Toxicological
includes routes of exposure; related symptoms, acute and
information
chronic effects, numerical measures of toxicity
12
Ecological information*
13
Disposal consideration*
14
Transport information*
15
Regulatory information*
16
Other information
includes the date of preparation or last revision of the SDS
sheet
Information regarding the GHS was drawn largely from the Occupational Safety and Health
Administration (OSHA) website found at:
https://www.osha.gov/dsg/hazcom/ghs.html#4.8. The United Nations documents detailing all aspects
of the GHS is available as the “Purple Book” at:
http://www.unece.org/trans/danger/publi/ghs/ghs_rev00/00files_e.html
Section #
1
Title
Identification
* Information in Sections 12-15 comes from agencies other than OSHA.
-14-
Experiment #1: Laboratory Safety and Practice
The shaded regions in Table 1.1 suggest areas of the SDS that may be of the most interest to you for
materials used in this CHEM 1361 laboratory. Looking at the shaded areas item-by-item:





Section 2 – It could be helpful to understand the hazards regarding any chemicals you will be
using in the laboratory
Section 4 – First-aid measures could come into play in the laboratory and it is important to know
what the SDS says in this area
Section 8 - Notice that recommended PPE are given in Section 8, though at no point will you
escape at least wearing goggles while in the laboratory regardless of whether or not they are
recommended by the SDS. The limit values are measures of how much exposure an individual
may experience under different circumstances before experiencing ill effects.
Section 10 – Knowledge of the stability of a chemical and products of its hazardous reactions may
be found here and could be of value.
Section 11 – The Toxicological information informs you of the route of exposure (remember there
are four routes), symptoms of exposure, and also provides numerical data regarding the level of
toxicity.
Though this may look a little intimidating, realize that you are working in a controlled environment in the
presence of instructors who are familiar with the chemicals being used. Materials are also selected with
consideration for their safety in the general chemistry laboratory.
Finding the SDS (or currently called MSDS) sheets is not difficult. Each of our laboratories contains a
binder with SDS sheets for the chemicals commonly stored in that laboratory. In addition, the stockroom
(SC 222) contains binders with SDS sheets for ALL of the chemicals in our storeroom. In addition, there
are multiple internet sites which make available SDS sheets. As examples:
http://www.ilpi.com/msds/ provides a variety of sites at which you can find SDS sheets
http://ehs.okstate.edu/links/msds.htm Oklahoma State’s site provides a number of locations
and there are many more. Please be aware that we are not talking about just chemicals used in the
laboratory. As you look, for example at Oklahoma State’s site, you will see sites focusing on inks and
toners. Some sites list janitorial and office supplies. Any time a manufacturer provides a hazardous
chemical they are also required to supply the SDS with it.
The United States (and several other nations) have enacted “Right-to-Know” laws regarding the relaying
of information to affected parties about the materials they are routinely exposed to. A summary from the
OSHA site is given as:
Protection under OSHA's Hazard Communication Standard (HCS) includes all workers
exposed to hazardous chemicals in all industrial sectors. This standard is based on a simple
concept - that employees have both a need and a right to know the hazards and the identities of
the chemicals they are exposed to when working. They also need to know what protective
measures are available to prevent adverse effects from occurring.
Notice this does not apply specifically to students in an educational setting, but it is important for you to
realize that if you want further information about anything you will be working with in the laboratory we
will make every effort to accommodate that request. It is also important to realize that, wherever you
career path takes you, the Right-to-Work concept is in place and you are entitled to information about
substances you will be working with and those working for you have the same right.
Experiment #1: Laboratory Safety and Practice
-15-
IV.
Symbols and Pictograms
There is quite a bit of important information on the SDS sheets and you might imagine there are
circumstances in which someone – a firefighter, first responder, lab manager, etc. – may appreciate a
very quick summary of the pertinent information. You have likely seen information of this nature
presented in the form of pictures and diagrams – consider a tanker truck with a large diamond affixed
to it. In this section we will briefly look at two methods, both a current one and a future one, of
presenting this information pictorially.
The National Fire Protection Agency (NFPA) diamond currently in use is probably one of the more
prevalent public displays of safety information. The diamond and the significance of its sections are
given in Figure 1.2. (I have written the colors in the boxes – normally there would be appropriate
numbers there.) The numbers to be filled in may be found in Section 15 of the SDS sheet for the
specific material.
Red
Yellow
Blue
White
Figure 1.2 NFPA Diamond (from http://www.zeably.com)
-16-
Experiment #1: Laboratory Safety and Practice
Another set of visual symbols (pictograms) associated with the GHS will be starting to s how up fairly
soon. Figure 1.3 gives a summary of those symbols – you do not need to memorize these. Just look them
over to become familiar with the appearance of these emerging pictograms.
Figure 1.3 GHS Pictograms
Experiment #1: Laboratory Safety and Practice
-17-
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-18-
Experiment #1: Laboratory Safety and Practice
Laboratory Safety Orientation
All students participating in chemistry laboratories of the Department of Physical Sciences
of Cameron University must complete Safety orientation before participating in laboratory
activities. The following topics will be discussed with you at the first scheduled laboratory
session of an academic session and some are also covered in writing in Experiment #1.
Those involved in facilitating the safety orientation include the Laboratory
Manager/Supervisor, your assigned Laboratory Instructor and possibly the Departmental
Safety Officer.
I. Laboratory Introduction (policies and procedures)
II. Safety Procedures (reporting injuries, completing incidence reports, first aid
kits/supplies, safety shower, eye wash station, fire blanket, fire extinguisher, fire
hoses, first response concerns)
III. Evacuation Procedures (Each instructor will provide instruction on evacuation
procedures.)
IV. Laboratory Management (chemical usage, broken glassware, chemical disposal
procedures, laboratory/ workplace cleanliness, washing hands, safety goggles,
proper clothing.)
V. Safety Information (general laboratory safety rules, eye protection, safety data
sheets, laboratory preparation)
VI. First Responder (if there is an accident or injury, define the safety evacuation and
first aid procedures). The emergency response number on campus is 581-2911.
I, ______________________________________ (print name), have reviewed the CHEM
1361 laboratory policies and procedures including those procedures related to first aid and
lab safety. Opportunity was given to have questions and concerns answered or explained.
NAME ___________________________________Date ______________
Signature
My instructor’s name and office is _________________ Room SC _____
Sign and submit to your instructor after the safety orientation.
Experiment #1: Laboratory Safety and Practice
-19-
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Experiment #1: Laboratory Safety and Practice
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© Laboratory Safety Institute 2008. Reprinted with permission of the Laboratory Safety Institute (LSI). For a wealth of other
lab safety information, please see the LSI website at http://www.labsafety.org.
Safety Cartoon
-21-
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
Safety Observations
In the preceding cartoon there are numerous safety violations which are numbered. Determine the safety
violation corresponding to each number and write a short description of the violation in the blank
provided to show your understanding of the safety concern. One to three word descriptions are all that is
expected. (According to the LSI answer key, there are actually over 50 errors in the cartoon!)
1. _________________________________
2._________________________________
3. _________________________________
4. _________________________________
5._________________________________
6. _________________________________
7._________________________________
8. _________________________________
9. _________________________________
10. _________________________________
11._________________________________
12._________________________________
13._________________________________
14. _________________________________
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16._________________________________
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21. _________________________________
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23._________________________________
24._________________________________
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30. _________________________________
Safety Cartoon Error List
-22-
Experiment #2: Measurements, Accuracy, and Precision
Learning Objectives
Textbook Reference (Chemistry: The Central
Science, Brown/LeMay/Bursten, 11th Edition)
Measure and record length and volume to the proper number
of significant figures
Select the most appropriate piece of lab glassware to make
volume measurements to desired accuracy
Use mass and density data to calculate volume
Arithmetically manipulate measured quantities reporting
calculated values to the proper number of significant figures
Use the basic statistical concepts of mean and standard
deviation to evaluate precision of measurements
Section 1.5 – Uncertainty in Measurement
Section 1.4, 1.6
Section 1.5 – Uncertainty in Measurement
Appendix A.5 and Appendix of this lab
book – p. 125
I. Basic Concepts of Measurement
Typically two types of numbers occur in science – exact and inexact. Exact numbers include those
quantities that are simply a count of a relatively small number of items – one dozen eggs is 12 eggs,
one ream of paper is 500 sheets, the number of people in a room. Larger counts may be more inexact
– the census, for example.
Inexact numbers, on the other hand, are the result of a measurement using some sort of equipment
such as a ruler, balance, or volumetric glassware. This semester in laboratory you will make a large
number of inexact (not to be read as incorrect) measurements. In this experiment you will acquire a
sense of our measuring devices and techniques as well as an ability to calculate derived quantities
from these measurements.
In dealing with scientific measurement you will routinely encounter two terms – accuracy and
precision. Accuracy refers to the proximity of a measured value to the “true” or “actual” value.
Precision refers to the “scatter” of a series of measurements of the same quantity. For our purposes,
the calculation of the average or mean of a series of measured quantities gives some insight into the
accuracy while the calculation of the standard deviation gives an idea of precision. Note that a
measured (or calculated) quantity may be very accurate, but imprecise, or alternatively could be very
precise but inaccurate.
II. Uncertainty in Measured Quantities
Any measurement made in the laboratory includes an inherent uncertainty due to limitations of the
equipment used to make the measurement and the individual making the measurement. For example,
a meter stick is typically graduated in millimeters. The choice of a meter stick to measure something
on the order of 50 µm (0.000050 m) would not be a practical choice, as the larger markings on the
meter stick of millimeters (0.001 m) would not be adequate to distinguish the sizes of objects on the
order of 50 µm. Similar considerations apply when selecting a volume measuring apparatus – pipette,
burette, volumetric flask, etc.
As a general rule, measurements required in the laboratory should be made and recorded to a number
of significant figures that is one decimal place beyond the graduations on the measuring device – in
other words, an estimate beyond the graduation of the device. For example, when using a meter stick
to measure length, one should record each measurement to the 0.1 mm place if the meter stick is
graduated in millimeters. Similarly, in a graduated cylinder marked to the 0.1 mL, each measurement
should be recorded to the 0.01 place. This is true even if the estimated last digit is a zero.
It should be apparent that such an approach is not appropriate when using a digital device such as an
electronic balance. In such cases, all digits visible in the display window should be recorded.
Experiment #2 – Measurements, Accuracy, and Precision
-23-
III. Measuring Devices in the Laboratory
You will use an assortment of measuring devices during your time in the General Chemistry
laboratory. Some of these are summarized below.
A. Measurement of Mass
Our standard device for measuring mass is the electronic analytical balance (Figure 2.1). The device
is quite easy to read since it has a digital display. There is no requirement
for “estimating” an extra digit as that would be impossible with a digital
instrument.
The left side of the long bar powers on the balance. The right side of the
long bar (→O/T←) allows one to “tare” the balance. Taring the balance
sets the display window to 0.00 no matter what is sitting on the balance.
Though the tare button is useful for determining the mass of a sample
without having to weigh an empty container first, one must be careful to
ensure a tared mass will work in the particular experiment being
conducted.
Figure 2.1 – Analytical Balance
When working with the balance, it is always wise to check before making each measurement to ensure
that the window shows a measurement unit of “g”. With multiple people using the balance, this
occasionally gets changed to “N” (Newtons) which, left unnoticed, will cause a level of distress when
working with your data. To go back and forth between “g” and “N” one simply has to press the
button marked with the two circulating arrows.
B. Measurement of Volume
There are basically two types of volume measuring devices – those
that indicate a volume that has been delivered and those that indicate a
volume contained within a particular type of glassware. Devices that
indicate a volume that has been delivered typically are graduated so
one can read volumes directly, always being sure to estimate one
Figure 2.2 – Assorted Glassware
additional digit beyone the graduation. Devices that are designed to contain and/or deliver a particular
volume of liquid typically have one mark on them to indicate the level at which the device contains
the indicated volume.
Figure 2.2 shows four devices as examples of volumetric glassware you may use this semester. There
will be others, but the principles are the same. From left to right, the devices are a graduated cylinder,
a volumetric flask, a Mohr pipette, and a volumetric pipette. Both the graduated cylinder and the
Mohr pipette may be used to measure and deliver volumes within the range of the device. The
volumetric flask and the volumetric pipette only have value in measuring the precise volume indicated
by the device. For example, a 100-mL volumetric flask has no value unless one is wanting to make
exactly 100-mL of solution. A graduated cylinder and Mohr pipette may be used to measure and/or
deliver a wide range of volumes. NOTICE THAT, WHEN USING A MOHR PIPETTE, YOU DO
NOT DRAIN THE LIQUID ALL THE WAY OUT OF THE PIPETTE – ONLY DELIVER
BETWEEN THE MARKINGS.
To properly read liquid volumes one must recognize the existence and importance of the meniscus, a
curved surface formed due to the varying attractions of molecules between themselves and the
-24-
Experiment #2: Measurements, Accuracy, and Precision
container. Most menisci (plural of meniscus) show a curvature downward though a few, such as
mercury in glass, curve upward. Figure 2.3 shows the menisci of water in
a graduated cylinder and in a volumetric flask. The volume reading in the
graduated cylinder would be recorded as 39.0 mL since all readings will
include one estimated digit, in this case the zero.
On most glassware you will notice some markings that are intended to
Figure 2.3 - Menisci
indicate mode of use as well as accuracy and precision. Two markings of
primary interest are “TD” which stands for “To Deliver” and “TC” which
stands for “To Contain”. A device marked “TD” is calibrated so that the device will deliver the
indicated volume without any special effort. For example, with a “TD” device such as a volumetric
pipette, one does not “blow” out or take any extra measures to remove remaining liquid. The
calibration accounts for the remaining liquid. For a “TC” device, such as a graduated cylinder, the
entire contents must be removed to attain the desired volume. Temperatures are sometimes etched on
the glassware to indicate the temperature at which the calibration is most accurate since glass will
expand and contract some with temperature.
IV. Basic Statistical Treatment of Data
Though statistical treatment of data can become quite complex, we will work with the basic
information necessary to evaluate multiple measurements of the same property. There are two
parameters we will use occasionally to describe repeated measurements that are expected to yield the
same result. The average (or mean) is a common term with which you are familiar and the standard
deviation allows us to evaluate the precision of a set of measurements. Appendix A.5 (page 1110) of
your textbook gives details of calculating these two parameters and an example. The Appendix of this
lab book also describes the process and gives an example.
V. Efficiency Tips
You will notice as you work through the semester that there are a lot of shared pieces of equipment, a
lot of motion within the laboratory, and a lot of line forming at certain stations from week-to-week.
Please note that in many experiments the order in which the experiment is conducted is not important.
This week’s experiment is an example of that. You and your partner could choose to start with Part A
or Part B with no loss of continuity or information. As you work in the lab, be aware of opportunities
for working more efficiently by being aware of underused stations during the lab period.
Procedure
A. Linear Measurements
In this exercise you will use multiple measurements of the length of a piece of string to evaluate the
average and standard deviation of the measurements.
1. For this exercise, consider two pairs of partners to be one unit (four people in the unit). Select
one of the precut lengths of string for your group to use in making these measurements.
2. You will find a plastic four-sided meter stick in the lab. This stick is marked to different levels of
precision on each side – 1 m, 10 dm, 100 cm, and 1000 mm. Each of the four people in the unit
will use this stick to measure the length of the piece of string from Step A.1 using each of the four
sides of the meter stick. Remember that in measurement you will always estimate to one digit
beyond the finest graduation. For example, measurements taken with the meter side will be
estimated and recorded to the 0.1 m place. Each person will record their measurements on their
Laboratory Record form as well as the measurements taken by others in the unit.
Experiment #2 – Measurements, Accuracy, and Precision
-25-
B Volume Measurements
For this exercise you will have a variety of glassware available which will be used to measure 25-mL
of water. You will use a balance and the density of water to determine accurately the actual volume
measured by each device and will thus get a sense of the ability of different pieces of glassware to
accurately measure volume.
1. Record in the second column of the Laboratory Record on page 30 the masses of the 100-mL
beaker, 150-mL beaker, 250-mL beaker, 400-mL beaker, 25-mL volumetric flask and a 50-mL
graduated cylinder.
2. Using the markings on each of the pieces of glassware in Step B.1, measure 25-mL of tap water
to the best of your measuring abilities. Do not measure the water in another vessel and put it into
these containers – use each container to estimate the 25-mL.
3. Reweigh each of the pieces of glassware containing the measured 25-mL of tap water and record
the total masses in the table in the first column of the Laboratory Record on page 30.
4. For the final three pieces of glassware listed on page 30 you will use a volumetric pipette, a Mohr
pipette, and a burette. You will not record the masses of these pieces of glassware. Weigh
your dry 100-mL beaker and record the empty weight in the second column of the Laboratory
Record on page 30. Place approximately 100-mL of tap water in your 150-mL beaker. Using a
pipette bulb, draw 25-mL of tap water into the volumetric pipette past the mark but not all the
way into the pipette bulb. Remove the bulb quickly placing your index finger over the top of
the pipette. Slowly adjust your finger until the water level drops so the bottom of the meniscus is
exactly on the mark on the neck of the pipette. Then hold the pipette over your 100-mL beaker
and let the water drain into it. Weigh the beaker again and record the value in the table.
5. Repeat Step B.4 using the Mohr pipette and a burette to deliver 25-mL. Dry the 100-mL beaker
before draining the pipette into it. In the cases of the Mohr pipette and the burette you must stop
the flow of water when you have delivered the 25-mL – you do not want to drain the entire
pipette or burette.
6. Determine the mass of water in each piece of glassware in the table by subtracting the mass of the
glassware from the mass of the glassware plus water.
7. Use the density of water (0.9982 g/mL at 25 ºC) to convert the mass of water to volume and record
in the table.
8. Record any marking on the glassware such as TC, TD, uncertainties, temperatures, etc. Some may
not have these markings.
-26-
Experiment #2: Measurements, Accuracy, and Precision
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
Lab Time: __________
_________________________________________
Laboratory #2: Measurements, Accuracy, and Precision
Laboratory Record
A. Linear Measurements
1. Fill in the following table with the measurements of each member of your unit.
2. Calculate the average of each column and place it in the appropriate table cell.
3. Fill in the columns designated as “Square of Error” by subtracting each
measurement from the average and squaring that difference.
4. Sum the “Square of Error” columns and record in the indicated row at the bottom.
5. For the standard deviation row, take the sum of each “Square of Error” column,
divide it by the number of readings – 1. For example, if there are four
measurements divide by three.
6. Take the square root of the result in Step 5 to arrive at the standard deviation.
7. The Appendix in this lab book contains a small section on Error Analysis and gives
the mathematical formula for standard deviation for future reference in this course.
1-m markings
m
Square
of Error
10-dm markings
dm
Square
of Error
100-cm markings
cm
Square
of Error
1000-mm
markings
mm
Square
of Error
Person 1
Person 2
Person 3
Person 4
Average
Sum of
“Square of
Error”
Sum of
“Square of
Error”/(n-1)
Standard
Deviation
Experiment #2 – Measurements, Accuracy, and Precision
-27-
Measuring
Device
Mass of
glassware
plus water
(g)
Mass of
empty
glassware
(g)
Mass of
water
(g)
Calculated
volume of
water
(mL)
Markings on
glassware
100-mL beaker
150-mL beaker
250-mL beaker
400-mL beaker
50-mL graduated
cylinder
25-mL
volumetric flask
Table for Recording Data from Volumetric Pipette, Mohr Pipette, and Burette (Procedure Step 4)
Mass of
beaker plus
water (g)
Mass of
empty beaker
(g)
Mass of water
(g)
Calculated
volume of water
(mL)
25-mL
volumetric
pipette
25-mL Mohr
pipette
50-mL burette
-28-
Experiment #2: Measurements, Accuracy, and Precision
Markings on
glassware
Circle
Lab Day: M
Your Name:
T
W
R
F
_________________________________________
Partner’s Name:
_________________________________________
Lab Time: __________
Laboratory #2: Measurements, Accuracy, and Precision
Post Laboratory Exercise
1. State which of the pieces of glassware you have used would be most appropriate for
obtaining the indicated volumes under the following circumstances. More than one
answer may be appropriate in each case.
a. A bowl contains roughly 35 – 40 mL of water.
b. In carrying out an analysis, one needs to deliver as accurately as possible 25.0-mL
of a solution into separate flasks.
c. A procedure calls for mixing about 40 mL of hot water into a beaker containing
cold water.
2. Three measurements were made in an attempt to find the density of a substance resulted
in values of 0.352 g/cc, 0.390 g/cc, and 0.310 g/cc.
a. Find both the average and standard deviation of these measurements. Refer to
the Appendix of this lab book if necessary. Most likely your calculator will
help you conduct these sorts of calculations quickly, but for now show your
work.
b. If the actual density of the substance was 0.351 g/cc, circle the phrase below
that best describes these results.
accurate but not precise
both accurate and precise
precise but not accurate
neither accurate nor precise
Experiment #2: Measurements, Accuracy, and Precision
-29-
3. Give the volume reading of the following buret to the proper number of significant
figures.
Reading to the proper number of significant figures
(include units): _________________________
4. Defining the percent error as:
% error 
-30-
Measured Value  Expected Value
100%
Expected Value
a.
Which of your measuring devices showed the smallest percent error
(remember the expected value in each case was 25 mL)? Show your work.
b.
Which of your measuring devices showed the greatest percent error? Show
your work.
Experiment #2: Measurements, Accuracy, and Precision
Experiment #3: SUGAR IN SOFT DRINKS AND FRUIT JUICES
Learning Objectives
Textbook Reference (Chemistry: Atoms First,
Burdge & Overby, 2012)
Relate density and buoyancy
Section 1.4 – Scientific Measurement
Understand the relationship between sugar concentration and
density
Construct simple hydrometer
Learn principles behind calibration
Use Excel to evaluate calibration data
Use calibrated device to measure unknown
INTRODUCTION
Fat, sugar, and salt are three food components that most Americans consume in excess. For example,
they are present in large amounts in typical fast food meals, fat in potato chips, and sugar in soft drinks.
This experiment focuses on sugar in beverages. Almost all soft drinks that are not milk-based (Coke,
Pepsi, Sprite, Kool-aid, Gatorade, etc.) are essentially solutions of sugar or artificial sweeteners with
small amounts of additives for flavoring and color. Fruit juices are also mostly sugar solutions with small
amounts of other materials. Most of these beverages contain a surprisingly large amount of sugar. In this
experiment you will determine the sugar content of various beverages by measuring the density of each
beverage. When sugar is dissolved in water, the density of the resulting solution increases compared to
pure water. As the concentration of sugar is increased, the density of the solution continues to increase.
By measuring this density increase for various beverages and comparing it with sugar solutions of known
concentrations, it is possible to determine the approximate concentration of sugar in the beverages.
Archimedes’ Principle can be used to evaluate the density of these liquids. According to Archimedes’
Principle, an object immersed wholly or partially in a fluid will experience a buoyant force equal in
magnitude to the weight of the volume of fluid displaced. If the buoyant force on an object is greater than
its weight, the object will float. Notice from the italicized portion of the statement of Archimedes’
principle that the tendency of an object to float depends on its weight – volume relationship, i.e., its
density. If an object is less dense than the fluid it is in, it will float. Thus we can use the ability of the
sugar solutions to float an object to determine the solutions’ densities.
An object will float to different heights in fluids of different densities. The object will sink more deeply
into a less dense fluid than into a more dense fluid since it has to displace a larger volume to match the
weight of displaced fluid. (It is easier for you to float in salt water than fresh water since salt water is
more dense.) This observation is the basis of a hydrometer. A hydrometer is designed to float in fluids
and, based on the height its stem protrudes above the solution, give an indication as to the density of the
fluid. In today’s experiment you will build a simple hydrometer and use it to determine the density of a
variety of common drinks.
In addition to illustrating the above principles, this experiment serves as an introduction the calibration of
instrumentation and its subsequent use in determining physical quantities of interest. It is relatively easy
to get a qualitative idea of the relationships between a large assortment of observations and the properties
of a system of interest. For example, in this experiment, the hydrometer will float higher in the solutions
with a higher percent sugar. However, based on that information we could only arrange the sugar
percents in order. To find the quantitative (numerical) percent, we first record the height of the
hydrometer in solutions of known concentration. From this information, we can then place the
hydrometer in a solution of unknown sugar concentration and determine the %sugar of the unknown by
comparing the hydrometer height to the calibration data.
Experiment #3:
Sugar in Soft Drinks and Fruit Juices
-31-
PROCEDURE
A. Construction of the hydrometer
1. Cut the last segment off of the tip of a plastic pipet.
2. Get some metal pieces that collectively weigh 5.1 grams. Add the metal pieces to the pipet such
that the metal is caught in the bulb of the pipet. Make sure the metal pieces go straight in and all
the way to the bottom. Float your hydrometer in water bulb-end down. It needs to float such that
only the top 1-3 cm projects above the water line. Do not drop it into the water because you might
sink it and have to start over.
3. Roll one of the small paper millimeter scales provided and slide it into the pipet. The scale does
not need to match any particular marking on the pipet and does not need to be in any certain
direction, it needs only to extend from about the bottom of the stem to the top of the stem. Cut off
any excess scale extending beyond the tip of the pipet.
4. Place the pipet, bulb-end down, into a 50-mL graduated cylinder filled to the 50-mL mark with
water. The pipet should float with about two centimeters of the stem sticking out of the water. If it
floats too high, remove the scale and adjust the weight in the pipet until the hydrometer floats with
about two centimeters of stem sticking out of the water.
B. Calibration of the hydrometer
Each hydrometer constructed will behave somewhat differently when placed in the solution. Each
hydrometer needs to be calibrated so that the height it extends above the solution will correspond to a
known % sugar in solution.
1.
2.
3.
Select water or one of the four standard sugar solutions. Place your hydrometer in the solution,
bulb-end down, and be sure it is not touching the walls of the graduated cylinder. If it does touch
the walls, attempt to move it around to avoid contact. If simple moving does not work, take the
hydrometer out of the solution and attempt to level the weights in the bottom.
Once the hydrometer is free-floating in the solution, record the meniscus height on the graduated
scale to the nearest mm. This value will be referred to as the hydrometer reading.
Repeat steps 1 and 2 for the remaining standards and pure water.
C. Measurement of Density of Drinks
Safety Note: Do not drink beverages that have been opened and used in the laboratory. Beverages
have been previously decarbonated by gently boiling and immediately cooling to room temperature.
1.
Repeat the procedure in Part B for each of the drinks arrayed around the laboratory. Record the
data in the table in the Laboratory Record.
D. Data Treatment
You have acquired data for effectively two different parts of the experiment. The data in one part will
be used to calibrate the hydrometer you have constructed. Be aware that your hydrometer and data are
unique – if you lose your calibration data you cannot borrow someone else’s. In this part of the
experiment you will develop a mathematical expression that correlates the hydrometer reading with
the % sugar in the solution.
In the second part of the experiment you will use the result from the first part to calculate the % sugar
in a variety of drinks. Since you will know the relationship between the hydrometer reading and %
sugar from the first part, it will be a straightforward exercise to use the measured height to determine
the % sugar.
1.
-32-
Development of Calibration Relationship
The calibration solutions include water (which is 0% sugar), and 4%, 8%, 12%, and 16% sugar
Experiment #3: Sugar in Soft Drinks and Fruit Juices
solutions. As you review your data from these runs, you will notice that the increase in height of
the hydrometer is about the same for each 4% increase in sugar concentration. A crude way of
developing the calibration relationship would be to take advantage of this observation; determine
the change in height for each change in % sugar, and take an average of the change in % sugar per
mm hydrometer reading. However, this approach does not take full advantage of all of the
accumulated calibration data.
A more sophisticated approach is to recognize that a relationship of this form is equivalent to a
straight line (the familiar y = mx + b form). In this case, the straight line would take the form:
(3.1):
hydrometer reading = m × % sugar + b
(3.1)
where the slope (m) and the intercept (b) come from evaluating the calibration data. A plot of the
hydrometer reading vs. the % sugar (in the form y vs. x, the y-axis is always the vertical axis
representing the dependent variable and the x-axis always the horizontal axis representing the
independent variable; in this case, the hydrometer reading will be vertical and the % sugar will be
horizontal) will yield a straight line of slope m and intercept b. The values of m and b come from
the determination of the best straight line through the calibration data using the linear least-squares
approach. In the linear least-squares approach, the best line is determined by minimizing the sum
of the squares of the different experimental values and a proposed best-fit line. The process will
be carried out in Microsoft Excel. Instructions for using Excel are given in Part D.3 of this
experiment.
2.
Determination of % Sugar in Drinks
Once the values of m and b have been determined it is a simple matter to use Equation (3.2) to
find the % sugar in any solution after determining the height to which the hydrometer floats.
Rearrangement of Equation (3.1) gives:
hydrometer reading - b
(3.2)
m
3. Use of Microsoft Excel 2010 to Determine Best Straight Line
Microsoft Excel 2010 is an example of software called a spreadsheet. Spreadsheets allow one to
handle data sets up to extremely large sizes and to look for relationships, make graphs of the data
in a large number of ways, and to draw statistical information from the data. When questions arise
about particular approaches to using any software, please look at the help menus (usually a small
question mark in the upper right corner of the screen) for help.
% sugar =
Using Excel, you will create a graph of your calibration data from the first part of the experiment
to be used in determining the sugar content in unknowns in the second part of the experiment. The
use of Excel allows an easy method to evaluate the quality of your calibration data and the
numerical values of m (the slope) and b (the intercept) for a straight line.
Data in a spreadsheet are entered in a row and column format with each row labeled with a
number and each column labeled with a letter. The format for referring to cells is to use a
{column letter} {row number sequence}. For example, in Figure 4.1 the cell designated B3 has
the number 5.75 in it. Data are placed into spreadsheet cells by typing the desired entry and then
either hitting the “Enter”, “Tab”, using the keyboard arrows to navigate, or using the mouse to
select cells for the next entry. It is helpful to put a heading above each column of data. Navigate
to cell A1 and type “% Sugar” and then to cell B1 and type “Hydrometer Reading”. Enter your
calibration data in two columns below the headings you have typed.
Experiment #3:
Sugar in Soft Drinks and Fruit Juices
-33-
Figure 3.1 – Microsoft Excel 2010 Spreadsheet with Graph
The next step is to make a graph of the calibration data. A good plot will have a title, axis labels,
and a legend if necessary. First select your two columns of data by clicking with the mouse in the
upper left hand corner (A1 in Figure 4.1) and then dragging the cursor while still holding the
mouse button down to the lower right-hand cell (B5 in Figure 4.1). Click on the Insert tab in the
ribbon at the top of the screen. Choose Scatter and then the upper left-hand choice that appears – it
shows points with no lines. A plot of your data will appear but will not have an appropriate title or
axis labels. The next paragraph shows you how to put on the chart title and axis labels.
Click on the Layout tab in the ribbon. Choose Chart Title. Choose one of the options as to where
you want the chart title. When you do, a box will appear above the plot area that says Chart Title.
You can customize the text by either double-clicking on that box and entering your text or by
entering your text directly in the white area above the column headings. A similar approach is used
to put in the axis titles – just choose axis titles from the same place. You can also delete the legend
(possibly says Series 1) by clicking once on it and hitting the delete key. A legend is helpful if you
have two or more sets of data to plot – we only have one here.
To add the linear least squares best-fit equation, right-click on any data point and select Add
Trendline. Under the Trendline Options, make sure the Linear option is selected. Select the boxes
labeled “Display Equation on Chart” and “Display R-squared Value on Chart” at the very bottom
of the Format Trendline box and close the box. You will see the equation displayed on the chart.
The R-squared is a measure of the agreement of the data to a linear equation. A value of 1.00
means a perfect fit and a value of 0 means there is not much linearity to these data. Note the
equation as presented gives you the m (slope) and b (intercept) values. The way we have made this
plot, the y-axis represents the hydrometer height values and the x-axis is the %sugar.
The last step is to use the slope and intercept determined to find the % sugar in the drinks using
Equation (3.2). Excel is well-designed to calculate numerical data from formulas entered. In our
case, we want to calculate the % sugar in each of several drinks based on the hydrometer height
measurement. If one labels, say cell A10, with the drink name, cell B10 as hydrometer height
corresponding to the drink in A10, and cell C10 as % sugar, a set of formulas can be entered that
will automatically fill in the entries below C10 based on the values entered in B10. From Equation
(3.2), one simply needs to subtract the intercept from the measured height and divide by the slope.
If cell B8 contains the slope and B9 contains the intercept, the % sugar can be calculated in cell
-34-
Experiment #3: Sugar in Soft Drinks and Fruit Juices
C10 from a hydrometer height entered in cell B10 by entering the formula (B10-B9)/B8 into cell
C10. A formula entry in Excel needs to begin with a “+”, “-“, or “=” – otherwise it just thinks it’s
a bunch of letters. A formula can also be copied and pasted and, by default, the pasted formulas
will retain the relative relationships from the first formula. For example, in the formula here the
B10 actually indicates the cell to the immediate left of cell C10. Thus, when copied, the cell
reference will change to indicate the cell to the immediate left of where the formula is entered.
Since we always want to refer to the same cells for the slope and intercept, we will use a “$” to
indicate those are absolute references. Enter the slope and intercept values into cells B8 and B9,
respectively, and enter the drink names into column A starting at A10. Place the hydrometer height
for each drink into its corresponding cell in column B. In cell C10, place the formula:
 (b10  $b$9) / $b$8
(3.3)
Compare this to Equation (3.2). Right-click on cell C10 and select the Copy option. Drag the
mouse cursor down column C until all of the C cells adjacent to your height entries are highlighted
and then right-click and select Paste. You will see the % sugar column filled in with values
calculated from your measured heights. You do not need to plot these – simply list them in the
data table.
Experiment #3:
Sugar in Soft Drinks and Fruit Juices
-35-
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Circle
Lab Day: M
Your Name:
T
W
R
F
_________________________________________
Partner’s Name:
_________________________________________
Lab Time: __________
EXPERIMENT #3: SUGAR IN SOFT DRINKS AND FRUIT JUICES
Laboratory Record
Data from Calibration Samples:
sample
sugar concentration
x axis
Standard 1
0
Standard 2
4%
Standard 3
8%
Standard 4
12 %
Standard 5
16 %
hydrometer reading (mm)
y axis
↑↑↑↑↑↑↑↑DATA TO BE PLOTTED ON CALIBRATION PLOT↑↑↑↑↑↑↑↑
↓↓↓↓MEASURED FROM SOFT DRINKS – CALCULATED FROM CALIBRATION DATA↓↓↓↓
NOT PLOTTED
Name of Drink
Hydrometer reading
(mm)
Calculated % sugar from
spreadsheet
!!!BE SURE TO TURN IN YOUR CALIBRATION GRAPH!!!
Experiment #3: Sugar in Soft Drinks and Fruit Juices
-37-
EXPERIMENT #3: SUGAR IN SOFT DRINKS AND FRUIT JUICES
Post-Lab
1. Let’s assume the manufacturer’s claims that there is no sugar in diet drinks are accurate. (We are
also assuming all sorts of other things but those aren’t important right now.) That would suggest
that all of the diet drinks should read 0% sugar. Based on looking at your diet results compared to
the 0% expected for the diets, how big would you estimate the uncertainty to be in your
measurements? Express your answer in a form like ±1% or ±3% or whatever number you think it
might be.
2.
How does sugar content of fruit juices compare to typical colas and carbonated beverages?
3. Consider a 12-oz. can of soda which is fairly typical. Show your work.
a. If there are approximately 30 mL in a fluid ounce, how many mL of soda are in 12 ounces?
b. Select the highest percentage sugar drink you tested. Which drink is it and what is its
percent sugar?
c. How many grams and how many teaspoons (1 teaspoon of sugar is about 4 g) of sugar are
in 12 ounces of this drink?
4. As sugar is added to water the density of the solution increases so all of the sugar-containing
solutions here had a density greater than 1 g/mL. It is also possible measure the density of
solutions with densities below 1 g/mL using a hydrometer. What modification would you have to
make to your hydrometer to measure the density of solutions significantly less than 1 g/mL?
5. Most people can float in the Great Salt Lake, though they may struggle floating in Lake Lawtonka.
What can be said about the density of the water in the Great Salt Lake compared to Lake
Lawtonka and why do you suppose this is true?
-38-
Experiment #3: Sugar in Soft Drinks and Fruit Juices
EXPERIMENT #4: USING PHYSICAL PROPERTIES TO
DETERMINE THE IDENTITY OF AN UNKNOWN
Learning Objectives
Textbook Reference (Chemistry: Atoms First,
Julia Burdge & Jason Overby, 2012)
Distinguish between physical and chemical properties
Experimentally determine the boiling point of a substance
Section 1.3 – Properties of Matter
Experimentally determine the density of a substance
Experimentally determine the solubility of a substance
Use physical property data to identify an unknown from a list
of possibilities
Section 1.4 – Scientific Measurement
Introduction
Substances typically show two types of properties, referred to as physical properties and chemical
properties. Physical properties are generally considered to be those not dependent on a substance’s
interaction with another substance. Chemical properties, on the other hand, are assigned to a
substance based on its interactions with other substances. Examples of physical properties include
melting point (mp), normal boiling point at atmospheric pressure (bp760mm) or at some other pressure
(e.g., bp0.2mm), density at a certain temperature (d20), the refractive index at a certain wavelength
(usually the “D” emission line of sodium, nD), the specific rotation []D, solubility, specific heat, heat
of fusion, heat of vaporization, hardness, color, spectral behavior, odor, etc. Chemical properties
include ability to burn (or not to burn), ease of oxidation, ability to rust (or not to rust), etc.
In today’s experiment, the identification of an unknown substance will be attempted using three or
four relatively easily measured physical properties: density, boiling point, solubility and refractive
index if the instrumentation is available. These three physical properties will not be unique for every
substance encountered but, if the number of possibilities is kept limited, unambiguous identification is
possible by the use of a combination of these three or four measured properties.
Density
Density provides the relationship between mass and volume of a substance. Defined as the mass a
sample of material divided by the volume it occupies, density has units of mass/volume and is
typically expressed as g/mL, g/cc, kg/L, etc. An alternate method of expressing density is called
specific gravity and gives the ratio of the density of a material divided by the density of water at the
same temperature. The specific gravity is a dimensionless quantity.
Boiling Point
Every liquid consists of both a liquid phase in which the molecules are condensed and also a vapor
phase containing the substance’s molecules in the gas phase. The pressure of the molecules in the gas
phase above the liquid is called the vapor pressure. As a liquid is heated, its vapor pressure increases.
The temperature at which the vapor pressure above the liquid is equal to the external pressure exerted
on the liquid is called its boiling point. If the external pressure is 1 atmosphere, as it will be in the
laboratory, the boiling point is called the normal boiling point.
Solubility
This is probably the easiest concept but one of the hardest observations to make. When two liquids
are placed in contact they either mix (are miscible) or do not mix significantly (are immiscible). Even
liquids considered immiscible really do dissolve in each other to a small extent. For example,
gasoline and water are said to be immiscible, but water is a good antiknock compound so oil
companies add as much water to gasoline as they can to raise the octane rating as cheaply as possible.
If one drop of gasoline is added to a glass of water, it cannot be seen suspended in the water but the
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
-39-
odor and the taste of the water is drastically affected. Generally two substances are considered
miscible if ten percent of one of the liquids dissolves in the other. Solutions are “clear”- not colorless,
but “clear”. “Milkiness”, cloudiness, and opacity are all possible indications that everything is not
dissolved. Obviously, layer formation is also indicative of insolubility.
When substances dissolve in each other, an observable feature is the “concentration gradient” that is
formed. Our eyes are very sensitive to variation in refractive index and we can see these variations.
This is usually expressed by the observer as “wavy lines” but is an observable indication that the two
substances are mixing with each other. If you see this, record it appropriately. Another problem is
losing the added solute in the meniscus of the solvent or in the curvature in the test tube bottom. Only
in careful observation can you circumvent these problems.
Refractive Index
The speed of light is typically quoted as 3 × 108 m/s (or 186,000 miles/s) in a vacuum. As light
travels through other substances, its speed is reduced. The ratio of the speed of light in a vacuum to
that in another substance of interest is called the index of refraction of the substance. In formula form,
speed of light in vacuum
speed of light in substance of interest
where ηD represents the index of refraction of the substance of interest.
D 
Procedure
A. Measurement of Density
1. This first step is only to help you get ready for Part B, the Measurement of Boiling Point. Set up
a hot plate next to a ring stand. Fill a 400-mL beaker about two-thirds full of water and put one
boiling bead (chip, stick, any disturbing material) in the water. Place the beaker on the hot plate
and turn it on low (to approximately the 2 or 3 on the 10 point dial of the hot plate). The boiling
bead will encourage uniform boiling without “bumping”.
2. Get your unknown from the instructor. Record the unknown’s identification code in the blank
provided in the Laboratory Record. Record the unknown number now – it is absolutely
essential that this be recorded in your Laboratory Record.
3. Weigh your 10-mL graduated cylinder. Record the mass to the appropriate number of significant
figures on the data sheet.
4. Add 2.5-3.0 mL of unknown to the graduated cylinder. Weigh the cylinder and the contents and
read the volume exactly. Record both the mass and the exact volume to the correct number of
significant figures as Trial 1 on your Laboratory Record data sheet. Do not empty the cylinder.
5. Add an additional 2.5-3.0 mL of unknown to the cylinder. Weigh the cylinder and the contents
and read the volume exactly. Record the mass and the exact volume to the correct number of
significant figures as Trial 2 on your Laboratory Record data sheet. Do not empty the cylinder.
6. Add an additional 2.5-3.0 mL of unknown to the cylinder. Weigh the cylinder and the contents
and read the volume exactly. Record the mass and the exact volume to the correct number of
significant figures as Trial 3 on your Laboratory Record data sheet. Do not empty the cylinder.
7. Calculate the mass of the unknown by subtracting the mass of the empty cylinder from the mass
of the cylinder and contents for each trial (three calculations) and record the masses. Division of
the mass by the volume will give the density. Record the density for each trial (three
-40-
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
calculations). Average the densities and record the average value and the standard deviation.
B. Measurement of Boiling Point
1. Your hot plate should already be on at a low rate. The water temperature should not exceed 35 oC
at the start of this portion of the experiment. If it does, take out some hot water and add some
cold water or ice.
2. Connect the Vernier stainless steel temperature probe to the LabQuest 2. Following the
instructions in Appendix A-4 set the data acquisition time to 900 seconds.
3. Add the 7-9 mL of unknown in the graduated cylinder to a test tube. If you spill any, wipe it up
and blot the test tube off on the outside. If you need more sample add from your unknown
container. (Caution you cannot have any more sample, so protect it from spills and contamination
and do not waste it.)
4. Clamp the test tube so that the liquid level in it is about
2 cm below the water line of the warm water in the beaker.
5. Clamp the Vernier stainless steel temperature probe into the
test tube so its tip is about 0.5 cm above the liquid level of
your unknown and centered in the test tube. (See Figure
3.1)
6. Turn the hot plate setting to a value of about 5. Start the
data acquisition on the LabQuest 2 by tapping the green
arrow in the lower left-hand side of the screen.
7. When the water starts boiling vigorously turn down the hot
plate setting to a level that just maintains the boiling of the
water.
8. When your data show a plateau for about 1-2 minutes stop
the data acquisition and turn off the hot plate. Do not let it
go for too long as you may boil away all of the unknown.
Figure 3.1 – Boiling Point Setup
9. Determine the mean (average) temperature of the plateau
using the instructions in Appendix A-4. This will be recorded
as your boiling point on the data sheet. Print one copy of the LabQuest 2 screen with the statistics
to turn in as a group. Be sure to include a title that includes the names of the members of your
group before you print it.
C. Determination of Solubility
1. This step is simply to demonstrate the “wavy” lines we are looking for in solubility testing. Place
a small spatula full (about the size of a pea) of sugar in the bottom of a test tube. Add by slowly
pouring down the side of the test tube 3 mL of water. Do not mix or shake the test tube. Place it
in a test tube rack and wait 5 minutes without touching it. After five minutes place a stirring rod
in the water above the sugar (don’t touch the sugar) and then take the rod out. You should be able
to see “wavy lines” streaming up from the sugar as the rod is withdrawn.
2. Get 4 test tubes and place in each 1-3 mL of water. Add 4-8 drops of methanol to one of the test
tubes. As you add the drops place the tip of the eyedropper close to the surface of the water but
not touching. Gently squeeze the dropper until a drop forms on the tip. As the drop forms let the
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
-41-
drop touch the side of the test tube above the surface of the water. Observe. (You must not
“drop” the droplet onto the surface as the “splash” will obscure your observations.) You likely
will see “wavy lines” streaming down into the water as the substance dissolves. You might see
the substance float upon the surface. You might see the substance sink to the bottom. After your
observation upon addition, take the eyedropper, squeeze it to remove air, push it to the bottom of
the test tube and release it, drawing the contents from the bottom. Bring the eyedropper to the top
surface and squeeze it strongly, expelling the contents vigorously into the tube. Do this 2-3 times.
If the test substance layered it might be dissolved after mixing. If the test substance is insoluble,
you will observe droplets and or a “milkiness” in the tube. Any cloudiness is indicative of
insolubility. Record your observations regarding solubility in the blank provided in the
Laboratory Record.
3. Repeat the solubility test (as described above for methanol in water) with 1-propanol,
cyclohexane, and your unknown using water as the solvent and record your observations.
4. Empty the contents of your test tubes into the disposal container. Dry your test tubes well using a
rolled up paper towel. Using a dry test tube, repeat the solubility test (as described above) using
cyclohexane as the solvent. Test the solubility of methanol, 1-propanol, and your unknown, in
the cyclohexane. Record your observations.
Eyepiece
D. Refractive Index
Depending on instrument availability you will be able to determine the
refractive index of your unknown using an Abbe refractometer, similar to the
one pictured in Figure 2.2.
Knob
Switch
1. Place a few drops of liquid on the stage and close the cover.
2. Rotate the light up over the stage.
Stage
3. Look through the ocular and move the switch on the left from the up
position down to its middle position.
Light
4. Look through the eyepiece and you will see a light area and a dark
area separated by a horizontal line (Figure 3.3a). Use the knob on the right
to align the horizontal line with the crosshairs in the window.
Figure 3.2 Abbe Refractometer
Figure 3.3a Switch in middle position
Figure 3.3b Switch in down position
5. Once the horizontal line is aligned with the crosshairs, move the switch on the left to the down
position (Figure 3.3.a). A numerical scale will be displayed. Read the upper number on the
vertical line in the center of the view. Be sure to read all significant figures. This number is the
refractive index and should compare well with one of the numbers in the table on the next page.
-42-
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
E. Unknown Identification
Your unknown is one of the compounds listed in the following table. These values are compiled from
the Aldrich Handbook of Chemicals (a chemical sales catalog) the Chemical Rubber Handbook of
Chemistry and Physics (an annual publication of data commonly called the CRC Handbook) and
MSDS sheets available in a Google™ search of the internet.
Possible Unknowns
Compound
bp
(/oC)
mp
(/oC)
nD
Refractive
index
density
g/mL
solubility
in water
solubility in
cyclohexane
water
100
0
1.3330
1.00
s
i
1-propanol
97
-127
1.3854
.804
s
s
2-propanol
82.4
-89
1.3776
0.785
s
s
cyclohexane
80.7
6.5
1.4262
0.779
i
s
ethyl acetate
77
-84
1.3720
.902
sl. s
s
ethanol (95%)
78
-16
1.3651
.816
s
sl. s
hexane
69
-95
1.3749
0.659
i
s
methanol
64.7
-98
1.3314
0.791
s
sl. s
acetone
56
-94
1.3590
0.791
s
s
*s = soluble; i = insoluble; sl.s = slightly soluble
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
-43-
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
_________________________________________
_________________________________________
Lab Time: __________
EXPERIMENT #3: USING PHYSICAL PROPERTIES TO
DETERMINE THE IDENTITY OF AN UNKNOWN
Laboratory Record
UNKNOWN NUMBER: ________
A Density Determination
Trial 1
Trial 2
Trial 3
Mass of cylinder + contents
________
________
________
Mass of cylinder
________
________
________
Mass of liquid in cylinder
________
________
________
Volume of liquid in cylinder
________
________
________
Density of liquid in cylinder
________
________
Average density of liquid in cylinder
_________________
Standard deviation of three results
_________________
________
B Boiling Point Determination
One member of your group will attach a copy of the plot from the LabQuest 2.
Boiling point (mean temperature of plateau):
____________ °C
C. Solubility Determination
Solubility Determination
methanol
1-propanol
cyclohexane
unknown
water solubility
cyclohexane
solubility
D. Refractive Index (if completed): The refractive index of my unknown is __________________
E. Unknown Identification
Conclusion: The density, boiling point, solubility, and refractive index agree best with the compound
_____________ which is my unknown # _________ .
-44-
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
F
EXPERIMENT #3: USING PHYSICAL PROPERTIES TO
DETERMINE THE IDENTITY OF AN UNKNOWN
Post -Lab
1. Based on your experiences with high altitude directions in cooking, pressure cookers, the
textbook, and anything else you might know, would you expect the boiling points of these
materials to be higher or lower if one conducted this experiment at the top of Pike’s Peak? If you
don’t have actual experience with the high altitude cooking directions, just wander into a store and
read a few boxes. Explain your reasoning.
2. When determining the volume of an unknown, a student read the volume at the top of the
meniscus rather than the bottom. Would this error increase, decrease, or not affect the measured
volume? Explain your reasoning.
3. If the student used the volume read in question 3 to calculate the density of the material, would the
calculated density be high, low, or the same compared to the actual value? Explain your
reasoning. Use the back of the page if necessary.
4. Suppose in determining the density of an unknown the student measured the volume correctly but
forgot to zero the balance before recording the mass. If the balance was reading -0.50 g before the
student placed the object on the balance, how would the mistake of not zeroing the balance affect
the calculated density? Explain your reasoning. Use the back of the page if necessary.
Experiment #4: Using Physical Properties to Determine the Identity of an Unknown
-45-
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EXPERIMENT #5: SEPARATION OF A MIXTURE
Learning Objectives
Textbook Reference (Chemistry: Atoms First,
Julia Burdge and Jason Overby, 2012)
Quantitatively determine the percentage composition of a
mixture
Understand the process of heating a material to a constant
mass
INTRODUCTION
A mixture is a physical combination of two or more pure substances in which each substance retains its
own identity. For example, in a salt (NaCl)/water mixture each component maintains the same chemical
form as it does in the pure state; water is still composed of H2O molecules and salt still exists as Na+ and
Cl- ions in solution. The original components can be recovered by the application of a physical change –
in this case, the evaporation of the water.
Mixtures are either homogeneous or heterogeneous. Homogeneous mixtures are called solutions.
Homogeneous mixtures consist of substances mixed at the molecular-size level to produce a mixture that
has uniform properties and composition throughout. Air, salt water, hot coffee, vodka and bronze are all
examples of homogeneous mixtures. Heterogeneous mixtures consist of distinct regions (or phases)
having different properties and compositions. Common examples of heterogeneous mixtures include
concrete, granite rock, whole milk, and water/oil mixtures.
Most things, whether found in nature or prepared in the laboratory, are impure; that is, they are part of a
mixture. One goal of a chemist is to separate mixtures so that the "good stuff" can be used and the
impurities can be discarded. The selection of a separation method depends on the differences in the
chemical and/or physical properties of its components. Common separation techniques take advantage of
differences in densities (settling or flotation), solubilities and immiscibilities, magnetic properties or
chemical reaction selectivity.
This experiment uses a combination of chemical and physical changes to separate a heterogeneous
mixture of magnesium sulfate heptahydrate (MgSO4 ∙7H2O also known as Epsom salt), silicon dioxide
(SiO2 also known as sand) , and sodium chloride (NaCl also known as table salt).
Magnesium sulfate heptahydrate decomposes at 100 oC according to equation (5.1) .
100°C
MgSO 4 •7H 2 O(s) 
 MgSO 4 (s) + 7H 2 O(g)
(5.1)
When heated, a mixture of MgSO4 ∙7H2O, SiO2 and NaCl will lose mass due only to the water escaping as
vapor from the MgSO4 ∙7H2O. The molar mass of MgSO4 ∙7H2O is 246 g/mol (1 × AWMg + 1 × AWS + 4
× AWO + 7 × FWH2O). Of this 246 g/mol, 126 g/mol (7 × FWH2O) are due to H2O giving a mass
relationship between the mass of water lost and the original mass of MgSO4 ∙7H2O present in the sample
before heating. Equation (5.2) illustrates a method for determining the mass of MgSO4 ∙7H2O originally
present based on the mass of H2O lost
mass MgSO4 •7H2 O in original sample = mass H2 O lost ×
246 g MgSO4 •7H 2 O
126 g H2 O
(5.2)
Only MgSO4, NaCl, and SiO2 remain after heating. Of these, MgSO4 and NaCl are water soluble; SiO2 is
insoluble in water. Addition of water to the remaining mixture will result in the dissolution of the MgSO4
and NaCl with the sand remaining as a solid.
Experiment #5: Separation of a Mixture
-47-
The sand can be separated by filtration or decantation (pouring the liquid layer carefully away from the
solid). The wet SiO2 is then heated, driving off any water adsorbed to the SiO2 particles. Finally, the
mass of the NaCl in the original mixture can be calculated by subtracting the sum of the MgSO4∙7H2O
and SiO2 masses from the total mass of the sample.
SiO2
NaCl
MgSO4
SiO2
NaCl
Add Water
MgSO4·7H2O
Heat
Original Mixture
Schematically, one can consider the experimental procedure to be composed of two steps as illustrated in
Figure 5.1.
SiO2
Figure 5.1 Schematic of Separation Process
-48-
Experiment #5: Separation of a Mixture
You are to complete one trial in determining the percent by mass of each substance in the mixture
and you are to get the data for two other trials from neighbors using the same unknown.
1. Obtain an evaporating dish from your instructor. Carefully inspect the evaporating dish for cracks
– if it is cracked, trade it in for an uncracked one. Record the mass of the clean, dry evaporating
dish on your Laboratory Record data sheet.
2. Use a scoopula, a spatula or a spoon to add between 2.9 and 3.1 grams of the unknown mixture to
your evaporating dish. Record the combined mass of the evaporating dish and sample to all
available significant figures.
3. Place the evaporating dish and sample on a clay triangle/wire
gauze set on an iron ring and ringstand as shown. Gently heat
the dish with the flame from a Bunsen burner for about 3
minutes and then more strongly for 10 minutes. (Gentle
heating means keeping the flame from the burner about 2
inches below the dish; for stronger heating lower the ring so
that the bottom of the evaporating dish is just above the inner
blue flame. Do not heat it so strongly that sand spatters out of
the dish.)
Figure 5.2 – Set-up for
4. Allow the dish to cool so that you can safely carry it to the
Heating Initial Mixture
balance and weigh it. You cannot weigh the dish while it is
hot; this will give erroneous results. Carry the dish on a paper, the triangle, or some other object
but your bare skin should not touch the dish. The grease and oils from your hands will add to the
mass.
5. Reheat the dish/sample for an additional 5 minutes, then cool and reweigh the dish and contents.
The idea is to drive all of the water off, as determined by an unchanging mass. If the mass of the
sample calculated after the second heating is within 1% of the mass of the sample after the first
heating, proceed to Step 6. If not, reheat and reweigh until the 1% criterion is met. This 1%
criterion is explained on page 52.
6. Add 5 mL of distilled water to the solid mixture in the evaporating dish. Carefully swirl the
contents to dissolve the salts and then decant (pour off without losing the solid) the liquid into a
separate beaker. Repeat this washing 5 more times, each time being careful not to pour out any of
the sand. You are extracting the soluble material from the insoluble materials. It is not critical to
get all the solution out each time, just do it carefully so as not to remove any solid.
7. Return the evaporating dish containing the wet sand to the clay triangle/wire gauze and heat until
the sand appears dry. Cool and record the mass of the evaporating dish and dry sand on your data
sheet on the “After 1st Heating” line under the heading “Determination of %SiO2 (sand)”. Repeat
the heating and cooling cycle until you get 1% agreement between sample sizes on successive
heatings (please see the note on the next page about heating to constant mass).
8. Get the data from two other lab pairs with the same unknown and compare results.
9. Clean up your work area, store your equipment, and return the evaporating dish to the instructor.
Experiment #5: Separation of a Mixture
-49-
NOTE ON HEATING TO CONSTANT MASS:
On more than one occasion during the course of the semester you will be asked to heat something to
constant mass. When the instructions say something like “until you get 1% agreement” that means that
the actual sample mass on the previous weighing is within 1% of the actual sample mass on the current
weighing. For example, consider the following set of data for this experiment.
Mass of empty evaporating dish
Mass of dish and sample after heating
After 1st heating
After 2nd heating
After 3rd heating
Mass (g)
33.8792
34.5673
34.5297
34.5243
Compare the masses after the 1st heating and the 2nd heating:
Mass of sample based on 2nd heating: 34.5297 – 33.8792 = 0.6505 g
Mass of sample based on 1st heating: 34.5673 – 33.8792 = 0.6881 g
% difference 
.6881  .6505
100%  5.78%
.6505
Since these two weighings differ by more than the requested 1%, the third heating is done. Checking it
for completion:
Mass of sample based on 3rd heating: 34.5243 – 33.8792 = 0.6451 g
Mass of sample based on 2nd heating: 34.5297 – 33.8792 = 0.6505 g
% difference 
.6505  .6451
100%  0.84%
.6451
By the 1% criterion asked for in the experiment, the 3rd heating is the last that needs to be done. In some
cases the % criterion may be different, but the process is the same. Notice this same approach may be
used whether something is gaining or losing mass.
-50-
Experiment #5: Separation of a Mixture
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
EXPERIMENT #5: SEPARATION OF A MIXTURE
Laboratory Record
Unknown Identifier: ______
Mass of evaporating dish and sample (g)
Mass of empty evaporating dish (g)
Mass of sample (g)
___________
___________
___________
Neighbor 1’s
Trial
____________
____________
____________
Mass of dish and sample after heating
After 1st Heating
After 2nd Heating
After 3rd Heating
Mass of water (H2O) lost (g)
___________
___________
___________
___________
____________
____________
____________
____________
____________
____________
____________
____________
Determination of % MgSO4· 7H2O (Epsom salt)
Mass of MgSO4·7H2O based on Eq (5.2) (g)
___________
Percent MgSO4·7H2O in sample
___________
____________
____________
____________
____________
Average % MgSO4∙7H2O in sample
____________
Your Trial
Determination of %SiO2 (sand)
Mass of dish and SiO2 (g)
After 1st Heating
After 2nd Heating
After 3rd Heating
Mass of SiO2 in sample (g)
Percent SiO2 in sample
Average %SiO2
Determination of %NaCl (sodium chloride)
Mass of NaCl in the original sample (g)
Percent NaCl in sample
Average % NaCl in sample
Summary of results:
Average of three trials:
Standard deviation of three trials:
Neighbor 2’s
Trial
____________
____________
____________
___________
___________
___________
___________
___________
____________
____________
____________
____________
____________
____________
____________
____________
____________
____________
____________
___________
___________
____________
____________
____________
____________
____________
% MgSO4·7H2O
% SiO2
____________
____________
% NaCl
____________
____________
___________
___________
Experiment #5: Separation of a Mixture
-51-
EXPERIMENT #5: SEPARATION OF A MIXTURE
Post Lab
1. The following list contains mistakes that could have been made during the course of conducting the
experiment. Read each statement and circle the appropriate affect the indicated error would have on the
indicated quantity.
In Step 7 of the experiment a student did not dry
the sand completely. What effect would this
Increased
Unchanged
Decreased
have on the % of sand calculated in the mixture
compared to the actual % of sand?
In Step 7 a student did not dry the sand
completely. What effect would this have on the
% NaCl calculated in the mixture compared to
the actual %NaCl?
Increased
Unchanged
Decreased
In Step 1 a student misrecorded the initial
evaporating dish mass by recording 25.45 g
instead of 26.45 g. How would this mistake
affect the apparent sample size compared to the
actual sample size?
Increased
Unchanged
Decreased
Increased
Unchanged
Decreased
The misrecorded evaporating dish mass was still
not caught and the student proceeded to calculate
the mass of sand in the sample. How would this
affect the calculated mass of sand in the sample
compared to the actual mass of sand in the
sample?
Increased
Unchanged
Decreased
In Step 6 a student accidentally dumped some of
the sand out during the washing/extraction step.
How would this affect the mass of NaCl
calculated compared to the actual mass of NaCl
in the sample?
Increased
Unchanged
Decreased
The misrecorded evaporating dish mass was not
caught and the student proceeded to calculate the
mass of MgSO4∙7H2O in the sample. How
would this affect the calculated mass of
MgSO4∙7H2O in the sample compared to the
actual mass MgSO4∙7H2O in the sample?
2.
-52-
What is the percent difference between the mass of your dried sample measured for your second
heating of the sample and the first heating of the sample when removing the water (i.e., the very
first set of heatings)? Show your work.
Experiment #5: Separation of a Mixture
EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULA
Learning Objectives
Textbook Reference (Chemistry: Atoms First, Julia
Burdge and Jason Overby, 2012)
Determine empirical formula from % composition
Section 5.10 –Molar Mass
information
Introduction
A compound is composed of two or more elements chemically combined. The percent composition of a
formula may be determined experimentally either by mass analysis or by synthesis. In mass analysis, a
weighed sample of the compound is separated into its constituent elements and the mass of each element
determined; in synthesis a compound is produced and the mass of the elements that combine to form a
definite mass of the compound is determined.
Results of elemental analysis are commonly expressed as percentages. For example, the results of
analysis might reveal that a particular compound has a composition of 10.4% C, 27.8% S, and 61.7% Cl.
To arrive at the empirical formula for this compound the relative number of atoms in one molecule (or
one mole) of each element must be determined. A simple way to convert from % composition to atoms is
to divide each composition by the atomic mass of the element. This division gives the number of moles
of each element in 100 g of compound. To arrive at simple whole number ratios, our preferred method of
reporting empirical formulas, each of the resulting numbers of moles is divided by the smallest number of
moles. In the example here:
Element % Comp. Atomic Weight %Comp./Atomic
Divide by smallest Empirical
Weight
Formula
C
10.4
12.0107
10.4/12.0107 = 0.867 0.867/0.867=1.0
Determined to
be C1S1Cl2 or
S
27.8
32.065
27.8/32.065 =0.869
0.869/0.867=1.0
simply CSCl2
Cl
61.7
35.453
61.7/35.453=1.74
0.867/1.74=2.0
The empirical ratio is the smallest set of integers giving the relationship between atoms in a compound.
However, the actual molecules of a compound may be comprised of multiple of these empirical units – we
would need other information to figure that out. For example, in the example above, if we also had
determined the molar mass was 230 g/mol, we could determine the molecular formula to be C2S2Cl4;
since each empirical unit has a mass of 115 g/mol there must be two units in one molecule.
On occasion the final ratios may not be whole numbers. As an example, the set of numbers for a
compound after dividing by the smallest number may be 1:1.5:2. In such a case, the entire set of numbers
may be converted to integers by multiplying all of the numbers by an integer to remove the fractional part.
In this case, multiplication by 2 would result in 2:3:4, a set of integers that would be used in the empirical
formula.
In this experiment, magnesium oxide will be produced and the percentages of magnesium and oxygen
determined in the final product. The oxygen source for making magnesium oxide will be the oxygen in
the air. Since air is primarily composed of two elements – nitrogen and oxygen – the two reactions
indicated in Equations (6.1) and (6.2) occur upon heating.
2 Mg(s) + O2 (g)  2MgO(s)
(6.1)
3 Mg(s) + N2 (g)  Mg3 N 2 (s)
(6.2)
It is important in this experiment that ALL of the magnesium be converted to magnesium oxide. Thus
any the Mg3N2 must be removed. This is done by adding water and heating according to Equation (6.3):
Mg3 N2 (s)+ 6 HOH ( )  3 Mg(OH)2 (s) + 2NH3 (g)
(6.3)
The Mg3N2 is converted to Mg(OH)2 which is then converted to MgO through further heating.
Experiment #6: Determination of an Empirical Chemical Formula
-53-
Procedure
1. Inspect a crucible carefully for visible cracks. Though these crucibles can handle extreme
temperatures they are rather delicate. You definitely want your initial crucible to stay with you for
the entire experiment.
2. Support a clean defect-free crucible and cover on a triangle as shown in Figure 6.1.
3. Heat the crucible and cover gently at first, then
strongly for about 5 min. Gentle heating is
accomplished by placing the crucible near the tip of
the outer Bunsen burner flame. The stronger heating is
accomplished by placing the crucible near the tip of
the inner Bunsen Burner flame. Cool the cover and
crucible on a wire gauze.
4. Use steel wool to polish 4-6 inches of a strip of
magnesium to remove oxidation products and use a
KimWipe™ to remove any grease. Crumple the
magnesium (don’t roll it up as we want to have as
much surface area exposed as possible).
5. Place the crucible and lid on a watch glass as well as
the piece of magnesium from Step 4. Carry the watch
glass to the analytical balance room. Record the mass
of the empty crucible and lid on the data sheet and
then the mass of the crucible, lid, and magnesium. Be
sure to record all digits displayed on the analytical
balance.
Figure 6.1 – Experimental Setup
6. Heat the crucible containing the magnesium with the cover ajar (see Figure 6.1). Heat gently for
several minutes until you see ignition. Control the burning rate by covering with the lid. You do
not want to see it smoking as that would indicate material is being lost.
7. Periodically remove the lid and replace it as needed, until you can heat the crucible fully with the
lid ajar. Heat the bottom to a “cherry red” for an additional five minutes.
8. Allow the crucible and lid to cool for approximately 5 minutes. When the crucible is cool add
enough water to wet all the ash in the crucible. (IF THE CRUCIBLE IS NOT COOL AT THIS
POINT WHEN YOU ADD THE WATER THE CRUCIBLE WILL BREAK OR CRACK
AND YOU WILL HAVE TO START OVER.)
9. Again heat the crucible with the burner for about 5 min. Cool the crucible. When the crucible is
cool, weigh the crucible, lid, and magnesium oxide. Record this mass on the Laboratory Record
data sheet using all digits displayed on the analytical balance. This will be your first recorded
mass of the oxide of magnesium (second line on the data sheet).
-54-
Experiment #6: Determination of an Empirical Chemical Formula
10. Heat the crucible again for another 5 minutes. Allow the crucible to cool and again weigh the
crucible, lid, and magnesium oxide. Record this mass on the data page. If the mass of magnesium
oxide formed does not agree within 3% of the mass found after the first heating, heat and find the
mass a third time. The final mass of the magnesium oxide is the mass to be used in the
calculations.
11. Get the data for a second run from another group.
Experiment #6: Determination of an Empirical Chemical Formula
-55-
This page is intentionally left blank.
Circle
Lab Day: M
Your Name:
T
W
R
F
_________________________________________
Partner’s Name:
Lab Time: __________
_________________________________________
EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULA
Laboratory Record
Trial 1
Trial 2 (from a 2nd pair)
Mass of magnesium (g)
____________
____________
Mass of crucible, lid and oxide of magnesium
after first heating (g) (Step 9)
____________
____________
after second heating (g)
(Step 10)
____________
____________
after third heating, if necessary (g)
(Step 10)
____________
____________
Mass of crucible and lid (g)
____________
____________
Mass of magnesium oxide (g)
____________
____________
Mass of oxygen in magnesium oxide (g)
____________
____________
% magnesium in compound
____________
____________
% oxygen in compound
____________
____________
% magnesium/atomic weight Mg
____________
____________
% oxygen/atomic weight O
____________
____________
Empirical formula for compound
____________
____________
Conclusion: Using both sets of data, the
empirical formula for magnesium oxide is:
____________
Experiment 6: Determination of an Empirical Chemical Formula
-57-
EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULA
Post-Lab Record
1. In each of the following laboratory situations an error is described. Circle the correct answer.
The day the experiment was run the air in the lab
was 18% oxygen rather than the usual 21%.
Would the measured mass of oxygen reacted be
larger than, equal to, or smaller than the amount
measured on a day when the oxygen was 21%?
Larger than
Equal to
Smaller than
A student failed to polish the magnesium ribbon
before weighing it. This error would make the
measured starting amount of magnesium larger
than, equal to, or smaller than the actual amount?
Larger than
Equal to
Smaller than
The student accidentally recorded 0.8966 g of
magnesium ribbon instead of the correct 0.8875 g.
Would this error make the calculated mol of
oxygen reacted be larger than, equal to, or smaller
than the actual amount reacted?
Larger than
Equal to
Smaller than
The crucible was heated too strongly after the
addition of water and popping caused some loss
of material. Would this error make the calculated
mol of oxygen reacted be larger than, equal to, or
smaller than the actual amount reacted?
Larger than
Equal to
Smaller than
The crucible was not heated to constant mass after
the addition of water – there was still some water
in the final weighing. Would this error make the
calculated mol of magnesium larger than, equal
to, or smaller than the actual amount reacted?
Larger than
Equal to
Smaller than
2.
-58-
What is the percent difference in your sample mass between the second
heating and the first heating? Show your work.
Experiment #6: Determination of an Empirical Chemical Formula
Experiment #7– Molecular Modeling
Learning Objectives
Draw Lewis structures for simple molecules and
ions
Determine formal charge from Lewis structures
Determine electron-domain geometry and molecular
geometry from Lewis structures
Assess polarity from Lewis structures
Textbook Reference (Chemistry: Atoms First,
Julia Burdge and Jason Overby, 2012)
Section 5.5 – Covalent Bonding and
Molecules
Section 6.3 – Drawing Lewis Structures
Section 6.4 – Lewis Structures and
Formal Charges
Section 7.1 – Molecular Geometry
Section 7.2 – Molecular Geometry and
Polarity
Draw conclusions regarding properties of molecules
from tabulated data
Introduction
(NOTE: YOU WILL BE HELPED IMMENSELY IN THIS LABORATORY IF YOU WORK
OUT THE LEWIS STRUCTURES FOR THE MOLECULES AND IONS IN THE DATA
RECORD PRIOR TO COMING TO THE LABORATORY.)
Bulk properties of materials are very dependent upon the molecular geometry as well as the charge
distributions within the molecule. In this experiment you will use molecular modeling software to study
these properties for a series of molecules and polyatomic ions.
Valence Shell Electron Pair Repulsion Theory (VSEPR) provides a means of approximating molecular
geometry. However, in the basic application of VSEPR, bond angles are initially assigned based on the
standard linear (180º), trigonal planar (120º), tetrahedral (109º), trigonal bipyramidal (90º and 120º), and
octahedral (90º) geometries. These angles will be different depending upon the chemical entities attached
to the central atom, but VSEPR allows only qualitative predictions about the actual bond angles.
Molecular modeling provides a means of more accurately predicting bond angles (and lengths) as well as
a wealth of other information based on the entire chemical species. In this experiment, you will:
A.
B.
C.
D.
E.
Draw Lewis structures for a variety of molecules and ions.
Assign formal charges to each atom based on the Lewis structure.
Assign electron-domain and molecular geometries.
Determine the polarity of each species.
Submit information to and retrieve information from the Scigress molecular modeling
program.
F. Draw some general conclusions from your observations.
The following gives a brief overview of some of these basics in case you have not yet covered these ideas
in class. If you wish further information, refer to your textbook (Brown, LeMay, Bursten, and Murphy) in
sections 8.3 through 8.6 and 9.2 through 9.3.
A. Drawing Lewis Structures (Section 5.5 of Burdge and Overby)
Though computers can do a phenomenal amount of calculational work, the output is only as good as
the input. To successfully complete this experiment you must feed some information culled from
Lewis structures into the software. In particular, the basic types and connections of atoms, types of
bonds (single, double, triple) and formal charges must be supplied. You have covered in class (or will
cover shortly) methods for drawing Lewis structures. The following is not intended to usurp an
approach you may be comfortable with – it just provides a method for you to use if you have not yet
Experiment #7: Molecular Modeling
-59-
seen this in class or are struggling with Lewis structures.
Lewis structures provide a means of understanding the basic bonding structure in a species and can
provide insight into some of its properties. One of the key aspects in drawing Lewis structures is to
recognize the significance of the octet rule which says that atoms tend to have eight valence electrons
around them in a bonded structure. Notice the word “tend” in the octet rule, suggesting there are
exceptions and indeed there are. Hydrogen is a key exception that only accommodates two electrons.
The set of guidelines below is a modification of the instructions on p. 314 of your textbook.
a. Write the symbols for the atoms to show which atoms are attached to which and connect them
with a single bond (consists of two electrons, but typically drawn as a single line). Sometimes
the order of atoms is given in the formula – e.g. HCN. In many polyatomic molecules and
ions, the central atom is written first – e.g. CO2 has carbon in the center, SO42- has S in the
middle, etc.
b. Complete octets of electrons around all atoms. Notice a bond counts as two electrons – they
are shared, but for purposes of counting electrons they both count toward the octet for each
atom.
c. Now count the number of valence electrons the structure should have. The number of valence
electrons for an atom is given by its group number on the periodic table. If the species is a
negative ion, it has extra valence electrons equivalent to its charge. If it is a positive ion, it is
missing the number of electrons equivalent to its charge.
d. If your structure has:
i. the correct number of valence electrons, you are done.
ii. too many valence electrons, delete a nonbonded pair on an atom and slide a nonbonded
pair on an adjacent atom between the two to form a multiple bond. Repeat this process
until your number of valence electrons is correct.
iii. too few electrons, put the necessary “extra” electrons on the central atom. Elements in
Period 3 and below may have more than eight electrons.
A few examples of drawing Lewis structures are given in the table below. Notice how the final structure
of the ion is in brackets and indicates the charge to the upper right of the brackets.
Examples of Drawing Lewis Structures
Step in Process
HBr
SO2
NO3Step a
H ― Br
O―S―O
O―N― O
│
O
Step b
O–S-O
O–N-O
H-Br
O
Step c – number
of valence
electrons
Step d
1 (H) + 7 (Br) =8
6 (S) + 2 × 6 (O) =
18
5 (N) + 3 ×6 (O) + 1
(charge) = 24
Step b and number
of valence electrons
match
Too many electrons
in Step b
Too many electrons in
Step b
O=N-O
O=S-O
H-Br
-60-
Experiment #7: Molecular Modeling
O
B. Assigning Formal Charges Based on the Lewis Structure (Section 6.4 of Burdge and Overby).
The information from the Lewis structure regarding the connection of atoms and types of bonds will
be fed into Scigress. The other important piece of information is the formal charge of each atom.
The formal charge basically compares the number of electrons shared and/or nonbonded around each
atom with its number of valence electrons. If it has more electrons around it in the Lewis structure
than its number of valence electrons, it will have a negative formal charge equal to the difference in
electrons. If it has fewer electrons around it in the Lewis structure than its number of valence
electrons, it will have a positive formal charge equal to the difference between the two numbers.
In counting electrons around each atom, all of the nonbonded electrons count for the atom to which
they are attached. However, in counting bonded electrons, only half of the bonded electrons between
two atoms count for each atom. Consider the table below – the Lewis structures for these examples
are in the previous table.
Examples of Calculating Formal Charge
HBr
SO2
H
Br
O=
=S-O
NO3O=
Looking at atom:
N
-O
1
7
6
6
6
5
6
6
Valence Electrons
0
6
4
2
6
0
4
6
Electrons assigned
due to nonbonded
electrons
½ •2=1 ½ •2=1
½ •4=2
½ •6=3
½ •2=1
½ •8=4
½ •4=2
½ •2=1
Electrons assigned
due to bonded
electrons
0+1=1
6+1=7
2 + 4= 6 2+3= 5
6 + 1=7
0+4= 4
4 + 2= 6 6 + 1=7
Electrons assigned
to atom:
1-1 = 0 7 – 7 = 0 6-6 = 0
6-5 = 1
6–7= -1 5–4 = +1 6 – 6= 0 6–7= -1
Formal charge =
valence – assigned
electrons
The inclusion of the formal charge on each atom is important for Scigress to successfully work with
your structure.
C. Assign Electron-Domain Geometry and Molecular Geometry to Molecules and Ions (Section 7.1
of Burdge and Overby)
The Lewis structure itself tells us nothing immediately about the geometry of the species since it is
drawn on two-dimensional paper and the species exist in three-dimensional space. There are two
principle types of geometry we discuss – the electron-domain geometry and the molecular geometry.
a. Electron-domain geometry
The electron-domain geometry comes directly from the Lewis structure and the application of
the Valence Shell Electron Pair Repulsion model, more commonly referred to as VSEPR.
VSEPR basically suggests the electron domains – nonbonded electrons and bonds (whether
single, double, or triple) – tend to repel each other and as a result will occupy regions of space
around the atom of interest which move the domains as far away from each other as possible.
Application of VSEPR is straightforward. Count the number of electron domains around the
atom of interest – the geometry is predicted to be linear if there are two electron domains,
trigonal planar if three, and tetrahedral if four. The model extends to five, six, and seven
Experiment #7: Molecular Modeling
-61-
electron domains but we will not consider those here. Refer to Table 9.1 on page 345 of your
textbook for a visual representation of these geometries.
b. Molecular Geometry
The molecular geometry focuses on the geometrical shape adopted by the atoms in the
structure. Though the nonbonded electrons are important in determining the geometry of the
species, the molecular geometry basically answers the question “What shape do the atoms
form in this structure?”.
A notation sometimes used to summarize this sort of information uses the letter A to represent
the central atom, B to represent an atom attached to the central atom, and E to represent a
nonbonded pair of electrons. As an example, in this notation SO2 would be written as AB2E.
The table below summarizes the geometries assigned to these sorts of structures. Notice the
electron-domain geometry depends only on the number of electron-domains around the central
atom – not whether they are nonbonded pairs or atoms. The molecular geometry is dependent
on the type of electron-domain.
Designation
AB2
AB3
AB2E
AB4
AB3E
AB2E2
Electron-domain Geometry
Linear
Trigonal Planar
Trigonal Planar
Tetrahedral
Tetrahedral
Tetrahedral
Molecular Geometry
Linear
Trigonal Planar
Bent (or angular)
Tetrahedral
Trigonal Pyramidal
Bent (or angular)
As an example, consider the S atom in the SO2 species from earlier.
O=S-O
There are three electron domains around the S atom – the double bond
going to O, the single bond going to O, and the nonbonded pair of
electrons on the S atom itself. Based on VSEPR, we would expect the electron-domain
geometry to be trigonal planar because of the three electron domains. Drawn more
geometrically accurately the shape of the molecule may look something like:
S
S
O
Electron-domain geometry: three electron-domains
around S so trigonal planar from first structure
O
O
O
Molecular geometry: Bent (or angular) from second
structure where nonbonded electrons on sulfur are
hidden from view.
D. Polarity and Dipole Moments (Sections 7.2 of Burdge and Overby)
Properties of molecules are significantly dependent upon a property of the molecule called its polarity.
The polarity of a molecule depends both upon the polarity of its bonds as well as its geometry.
The polarity is best quantified by considering the dipole moment of a molecule. The dipole moment,
µ, is defined as:
  Qr
where Q is the magnitude of two equal and opposite charges separated by a distance r. To determine
Q, a crude approximation is to find the magnitude and centers of negative and positive charge in a
molecule. The magnitude becomes Q and the distance between the centers becomes r. Notice if the
positive and negative centers are in the same location the dipole moment is zero – the molecule is
nonpolar. Larger values of dipole moment indicate larger polarities of a molecule. Units for the
-62Experiment #7: Molecular Modeling
dipole moment are typically Debye (D) (1 D = 3.34 × 10-34 C∙m where C stands for Coulomb, a unit of
electrical charge).
As mentioned earlier, both the polarity of a molecule’s bonds as well as its geometry contribute to the
polarity of the molecule. A polar bond is one formed between two atoms with differing
electronegativities – the larger the difference the more polar the bond. The geometry of a molecule
may be such that, even though it is composed of polar bonds, they may be arranged in such a way to
effectively counteract each other.
Consider the following two molecules. Notice boron is one of those exceptions to the octet rule – it
handles only six electrons quite well. The nonbonded electrons are omitted.
Cl
Cl
Cl
B
Cl
BCl3 is nonpolar because the Cl atoms pull
electrons equally in opposite directions. The BCl2F
is polar – the F atom is more electronegative and
draws electrons more strongly than the Cl atoms do.
B
Cl
F
E. Submit Information to and Retrieve Information From the Scigress Molecular Modeling
Program
A reasonable question at this point might be: Why do we need the software since the Lewis
structure can provide so much information already? One of the goals of this experiment is to
introduce you to the concept of computer-based molecular modeling. The program you will use is
considerably more powerful than you will see on this first experience. It is helpful to acquire a
basic understanding of providing information to such a program.
The basic flow for setting up these calculations in Scigress is as follows:
Draw the Lewis structure
for your molecule or ion,
being sure to assign
formal charge
Enter the names of the files
of your molecules of
interest into ProjectLeader
and tell it to calculate the
properties.
Draw your molecule
or ion in Workspace
being sure to assign
formal charges
Use the “Beautify” command to
automatically add hydrogens and adjust
bond angles and lengths to standard
values.
Open ProjectLeader and
identify the properties of
interest to be calculated.
Label the bond lengths and angles (if
there are three or more atoms). Save the
molecule under an identifiable name for
later use. Repeat the first four steps for
all of your molecules of interest.
Figure 7.1 Basic procedures for working with molecules in Scigress
Basics of Scigress Workspace
Either double-click the Scigress icon on the desktop or go to Start/Programs/Scigress/Workspace. (In
some older installations, the program may still be called CAChe.)
Figure 14.2 shows a cropped version of the Scigress workspace. There are four primary areas with which
Experiment #7: Molecular Modeling
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to be concerned initially:
1. Menu Bar along the top. This has the usual Windows software functionality as well as options
related to working with a molecule in the workspace.
2. Tool Bar directly below the Menu bar. All of the functions on the Tool bar can also be performed
using the Menu bar.
3. Tool Palette – Vertical set of buttons underneath the C – Carbon entry in Figure 14.2.
a.
The select tool
allows one to select an
atom, bond, or entire
workspace. Selecting an
atom or bond involves
simply clicking on the
desired entity. To select
an entire workspace click
somewhere on the
background.
b.
The molecule
select tool selects an entire
Figure 7.2 Scigress Workspace
molecule by clicking on one
of its atoms or bonds. This is true even if you have multiple molecules in the workspace –
the one you click on will be selected.
c.
The bond select tool selects all single, all double, or all triple bonds depending on
which type is clicked on.
d.
The group select tool allows one to define a group of atoms. This is more helpful
with organic molecules than with the molecules with which we will be working.
e.
The Drawing tool looks like a pen with atoms dripping out of the tip. This must be
selected to place atoms or bonds in the workspace.
f.
The Rotation tool allows for the rotation of the molecule in the workspace.
g.
The Translation tool allows for the translation (movement about the workspace) of
the molecule.
h.
The Magnification tool allows for changing the size of the molecule in the
workspace.
4. Style bar – Directly below the title Chemical Sample1* in Figure 14.2.
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a.
This drop-down allows for the selection of particular elements. When
the drop-down arrow is clicked, there will be a short list of elements and a Periodic Table
entry. If the Periodic Table entry is selected any of the available elements may be selected
from the table. One can choose the element type before placing atoms in the workspace or
can change an element type by selecting the atom and then using this drop-down to change
its type.
b.
This drop-down allows for the selection of the hybridization of
an atom. It is best to set this as unconfigured prior to experimenting with a molecule. If
Experiment #7: Molecular Modeling
this is either set incorrectly or not set to unconfigured, hydrogen atoms may “sprout” into
the structure when it is beautified.
c.
The formal charge on each atom is set through this drop-down. It is
vital to ensure the proper formal charge is entered before Beautifying or Experimenting
with a molecule. If this is not set correctly, hydrogen atoms may “sprout” into the
structure when it is beautified.
d.
This drop-down sets the bond type – single, double, triple, or others. The
bond type may be set with this drop-down before drawing a structure of may be modified
by selecting the bond to be changed and then using this drop-down.
“Experimental” Procedure
There are basically three parts to this experiment.
A. Build the molecules and ions of interest.
B. Set up a ProjectLeader spreadsheet within Scigress with which to conduct calculations.
C. Conduct experiments and retrieve information through ProjectLeader.
A. Build the molecules and ions of interest:
In this experiment you will build and make measurements on twenty-two molecules and/or ions.
1. Build the diatomic molecules: Have the Lewis structure with formal charges handy. Be sure you
have the bonding correct and the formal charges properly assigned
a. Select the atom type in the Workspace to be Hydrogen.
b. Select the drawing tool
c. Left-click once anywhere in the Workspace and a hydrogen atom will appear.
d. Hold the left mouse button down and drag to anywhere else on the workspace and when
you release the button another hydrogen will appear bonded to the first.
e. At this point one hydrogen atom is grayed out and the other highlighted. To highlight the
f.
g.
h.
i.
j.
whole molecule, you may click on the select tool
first and then click anywhere in the
workspace (not on the molecule).
With the select tool still selected, click once on each atom and look at the Style Bar to
ensure its formal charge is correct and its hybridization is set to Unconfigured.
Next label the bond for later use. With the select tool still selected, click once on the bond.
Select Adjust/Atom Distance from the Menu Bar. Then click the check box that says
Define Geometry label. Hit ok and you will see the bond labeled with a distance
Select Beautify/Comprehensive from the Menu Bar. The atoms may adjust distance a
little, but there probably will not be a remarkable change. If other atoms show up, it
indicates that probably your hybridization is not set to Unconfigured, your formal charges
are wrong, and/or your bonding is wrong. Compare your Scigress version to the Lewis
structure, repair if necessary, and try Beautify/Comprehensive again.
At this point, save the molecule under a descriptive name somewhere you can locate it for
later use. For example, save H2 as H2. You may use the desktop, thumb drive, or
wherever you want, remembering that computers in the Sarkey’s lab lose all new
information once they are rebooted.
Repeat steps a – i for the molecules F2, Cl2, HF, HCl, HBr, HI, and CO. As an efficiency
note, after saving your H2, you can use the same workspace as a starting point to draw the
Experiment #7: Molecular Modeling
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others. For example, to go from H2 to F2 do the following:
i. With the H2 workspace still open, do a File/Save As and give the file the name F2.
(This early save prevents you from overwriting H2).
ii. With the Select Tool selected, click anywhere in the workspace to select the entire
molecule. If you only want to change one atom (for example, to make HCl) just
select that atom.
iii. Go to the atom selection drop-down on the Style Bar and select F. Selected atoms
are instantaneously changed to fluorine. Notice the bond label remains – you do
not have to do that again.
iv. Check the bond type as well as each atom (as in Step f above) for the proper formal
charge and a hybridization of Unconfigured.
v. Select Beautify/Comprehensive from the Menu Bar.
vi. Hit the save icon. Now start again with Step j.i. to finish the other diatomic
molecules.
2. Building triatomic and polyatomic molecules.
Building triatomic and polyatomic molecules and ions is no more difficult than building diatomics
– they just have more atoms. There is one addition to the instructions in Steps 1a – 1i. Since three
or more atoms bring in the prospect of bond angles, we will want to label the bond angles
similarly to what we did with bond angles in the diatomics. Any bond angle involves three atoms
– two on the outside and one on the inside. To label a bond angle, click on the Select tool. Leftclick first on one of the outside atoms, hold down the CTRL key and left-click on the middle atom
and, while still holding the CTRL left-click on the other outside atom. Go to Adjust/Bond Angle
on the Menu Bar. Click on the box that says Define Geometry Label and then ok. DON’T
FORGET TO LABEL THE BOND ANGLES IN ANYTHING WITH MORE THAN TWO
ATOMS.
A. Set up a ProjectLeader spreadsheet within Scigress with which to conduct calculations:
ProjectLeader is a segment of Scigress that can be used to efficiently determine several properties
of molecules and ions. Here we will have it calculate the partial charges on atoms, dipole
moments, bond lengths, bond angles, and hybridization for the structures you have already built.
1. Select Start Button►Programs►Scigress►ProjectLeader. (Note: In some older installations
Scigress may be called CAChe – it is still Scigress.) The screen shown in Figure 14.3 pops up.
Figure 7.3 Opening Screen of ProjectLeader
2. In the Chemical Sample column you will enter the molecules you have already built. Double-click
in a grayed-out cell and you will be requested to supply a file name. Fill in the column under
Chemical Sample with the molecules/ions.
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Experiment #7: Molecular Modeling
3. Double-click in the gray area below the column heading B. A box will open to allow the selection
of the information to be placed in column B
(Figure 14.4). Select for Kind: Property of Atom
and click Next. Then choose charge partial from
the drop down list and click Next. For Kind of
Procedure select partial chrg at AM1 geom and
then click OK.
4. Repeat a similar process for Column D to allow
the calculation of the dipole moment. Doubleclick in the gray area below the column heading
D. Select for Kind: Property of Chemical Sample
and click Next. Then choose dipole moment and
click Next. Then choose dipole moment at AM1
geometry and click OK.
5. Conduct a process similar to Step 4 to tabulate
Figure 7.4 Selecting Properties in ProjectLeader
bond lengths. Double-click in the gray-area
under Column E. Select for Kind: Property of bond and click Next. Then choose bond length
and click Next. Then choose extract from sample component and click OK. Two columns will
appear, one labeled bond and the other labeled Bond Length (angstrom).
6. Conduct a process similar to Step 8 to tabulate bond angles for those bonds you have labeled.
Double-click in the gray-area under Column G. Select for Kind: Property of labeled angle and
click Next. Then choose bond angle and click Next. Then choose extract from sample component
and click OK. Two columns will appear, one labeled labeled bond angle and the other labeled
Bond Angle (degree).
7. One last column to include (unless you want to include others for fun). Double-click in the gray
area under Column I. Select for Kind: Property of atom and click Next. Then choose
hybridization and click Next. Then choose extract from sample component and click OK.
8. Drag the mouse across the B through D labels of those three columns while holding down the left
mouse-button. The three columns will become black. From the menu choose Evaluate/Cell or
alternatively right-click on the black area and choose Evaluate Cell from the exposed list.
ProjectLeader will now carry out experiments with your molecules and record the results in those
first three columns.
9. Once those three columns have completed, drag the mouse across any other letters corresponding
to properties you want to calculate (at least through I if you are using only the properties listed
above). Again, from the menu choose Evaluate/Cell or alternatively right-click on the black area
and choose Evaluate Cell from the exposed list.
10. Print out this spreadsheet with your data and hand it in with your report. In the Sarkey’s Lab (SC
203) go to the File menu and select Print. Choose the Science 9500 printer. Then select
Properties, choose Orientation, and set it to Landscape. Regardless of any printer you use, the
Landscape orientation will work best.
Experiment #7: Molecular Modeling
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Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
Experiment #7 – Molecular Modeling
Laboratory Record
Fill in the following information from the experiments you ran using ProjectLeader. Also include the
printout of your ProjectLeader results with this lab report. For designations of polar or nonpolar, assume
a species is nonpolar if its calculated dipole moment is less than 0.10 and polar if it is greater. IT
WOULD BE VERY HELPFUL TO HAVE WORKED OUT THE LEWIS STRUCTURES BEFORE
YOU COME TO LAB – IT WILL MAKE IT GO MUCH MORE SMOOTHLY FOR YOU.
Molecule/Ion
Lewis Structure
Diatomic Molecules
Molecule/Ion
H2
F2
Cl2
HF
HCl
HBr
HI
CO
Lewis Structure
Questions regarding diatomics:
1. From your ProjectLeader output, consider the dipole moments of the diatomic molecules.
a. Separate the diatomic molecules into those that are nonpolar and those that are polar.
Nonpolar: _____________________________
Polar: __________________________
b. What must be true of the atoms in a polar molecule for it to be polar?
2. Do the formal charges assigned from Lewis structures and the partial charges calculated by ProjectLeader agree in
value? Give examples to support your conclusion.
3. Look at the results for the partial charge, bond length, and dipole moments for HF, HCl, HBr, and HI.
a. What generally happens to the dipole moment as one moves down the halogens – i.e., HF → HCl → HBr → HI? Does it
increase, decrease, or stay about the same?
b. Is the trend you see in dipole moment the result of changing partial charge, changing bond length, or both? Explain.
Experiment #7: Molecular Modeling
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Molecule/Ion
Lewis Structure
Molecule/Ion
Lewis Structure
BE SURE TO LABEL YOUR BOND ANGLES ON THESE MOLECULES PER INSTRUCTIONS ON PAGE 114
Triatomic Species
CO2
NO2-
SO2
O3
H2 O
HCN
Questions regarding triatomic species:
1. Both CO2 and SO2 have a central atom surrounded by two oxygen atoms. However, one is polar and
the other nonpolar. Explain what causes this difference.
2. SO2, NO2-, and O3 all have a central atom surrounded by two oxygen atoms – one doubly bonded and one singly bonded. It
would seem a double bond would be shorter than a single bond since it contains more bonding electrons. Is this observation
consistent with the bond lengths calculated in Scigress for these species? If not, explain what might cause the behavior of
the bond lengths observed in Scigress?
3. Consider the geometries of each of the triatomic molecules.
a. What is the electron domain geometry for each of these species?
b. What is the molecular geometry for each species?
4. In a perfectly trigonal planar electron-domain geometry the angles would all be 120°. Consider the molecule SO 2 .
a. What is the O-S-O bond angle calculated by Scigress?
b. The calculated angle differs from 120°. What does the calculated angle tell us about the “amount of space” taken up by a
nonbonded pair of electrons compared to a bonded pair?
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Experiment #7: Molecular Modeling
Molecule/Ion
Lewis Structure
Molecule/Ion
Lewis Structure
BE SURE TO LABEL YOUR BOND ANGLES ON THESE MOLECULES PER INSTRUCTIONS ON PAGE 114
Polyatomic species:
CO32-
SO32-
SO3
PO43-
NH3
CH4
CH3Cl
CH2Cl2
Questions regarding polyatomic species:
1. What is the electron-domain geometry of each of the polyatomic species above?
2. Some of the above species are polar, some nonpolar.
a. Which of the species are polar?
b. Which of the species are nonpolar?
c. Considering the electron-domain geometries and the electron-domains surrounding each central atom, formulate a
general guideline based on these few examples that might be used to predict whether species such as these are polar or
nonpolar.
d. Using the notation from the introduction to this lab, and using your guideline from part c would expect each of the
following
to be polar or nonpolar. The letters B and C represent atoms of two different elements.
AB3E
AB2C2
AB2E2
Experiment #7: Molecular Modeling
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EXPERIMENT #8: PREPARATION OF AN ALUM
Learning Objectives
Learn the technique of gravity filtration
Learn the technique of vacuum filtration
Convert between grams and moles
Convert between molarity and volume and moles
Identify the limiting reactant in a chemical reaction
Calculate theoretical yield in a chemical reaction
Textbook Reference (Chemistry: Atoms First,
Julia Burdge and Jason Overby, 2012)
Section 5.10 – Molar Mass
Section 9.5 – Concentrations of Solutions
Section 8.4 – Limiting Reactants
Section 8.3 – Calculations with Balanced
Chemical Equations
INTRODUCTION
Aluminum is the most abundant metallic element found in the earth's crust. The pure metal is never found
in nature because it reacts over time with other chemicals present in the environment. Most naturally
occurring aluminum compounds are oxides or silicates.
The basic chemical process for obtaining aluminum metal from its ore involves electrolytic reduction of
the aluminum 3+ cation present in the principal ore bauxite, to elemental aluminum. The process is
extremely energy intensive and requires enormous amounts of electricity. Worldwide production of
aluminum metal currently consumes about 5% of the electricity output of the United States. Since the
recycling of aluminum metal only takes about 5% of the energy required to refine the metal from its ore, it
is very cost effective to recycle aluminum scrap metal.
Aluminum metal has many applications since it is very lightweight for a metal (density = 2.7 gram/mL),
is quite resistant to corrosion, and has a high tensile strength. The metal and its alloys are used extensively
in the manufacture of airplanes, automobile components, building materials, and electrical wiring. Other
familiar applications are aluminum foil and beverage cans.
An important property of aluminum is that it reacts with oxygen to form a thin layer of aluminum oxide
which protects the interior metal from further corrosion. Unlike iron or steel, which react with oxygen to
form a layer of rust which tends to flake away easily from the metal thus allowing further corrosion of the
metal surface, the oxide layer on aluminum adheres strongly to the metal thereby preventing continued
corrosion. It is estimated that a discarded aluminum can requires over 100 years in the environment to
corrode away. With available landfill sites rapidly filling up, it is an environmentally sound idea to
recycle all products made of aluminum.
In addition to its many uses as a metal, aluminum compounds also have a wide range of applications.
Some of the most widely used aluminum compounds are known as alums. An alum is a hydrated double
salt combination containing one of the group 1 cations (sodium or potassium) or ammonium ion, a metal
cation with a 3+ oxidation number (aluminum, iron or chromium), two sulfate anions, and n number of
hydrated water molecules. The general formula for an alum can be written as follows:
M1 M3 SO4 2 n H2O
(8.1)
An alum containing aluminum will have an aluminum cation having a 3+ oxidation number substituted
for the M3+ in the above equation. Alums are very useful compounds in many industrial processes. The
pulp and paper industry alone consumes more than a half million tons of the alum produced annually in
Experiment #8: Preparation of an Alum
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the U. S. In order to make writing paper, the open spaces in the cellulose fibers must be filled with
substances that adhere strongly to the fibers to prevent ink from bleeding. This is accomplished by adding
clay and pine rosins to the wet pulp. To fix the rosin strongly to the cellulose fibers, "papermaker's alum"
is added. The positive aluminum ion neutralizes the negative charge on the rosin and allows the rosin to
chemically bind to the fibers. The slurry of treated pulp is poured onto screens in a paper making machine
and the resulting mat is washed, dried, and rolled to produce a smooth sheet of paper like this one. Alum
compounds are also used extensively as flocculating agents and phosphate removal agents in water and
waste treatment plants. Other uses include soaps, greases, fire extinguisher compounds, textiles, drugs,
cosmetics, plastics and in pickling cucumbers- read the label of your pickle jar at home tonight.
The objective of today's lab is to acquaint the student with a method for producing alum from a piece of
scrap aluminum metal. This lab is intended to provide the student with an introduction to several basic
laboratory techniques, including: gravimetric filtration, preparation of chemical solutions and use of
common laboratory glassware. The following section includes a number of chemical reaction equations to
illustrate the chemical reactions that we will be conducting during the lab. If you are not familiar with
chemical reaction equation symbolism, refer to the textbook or ask the instructor to explain what the
symbols and equations represent.
OUTLINE OF THE PREPARATION
Aluminum metal reacts rapidly in a hot aqueous solution of potassium hydroxide, KOH, producing the
water soluble potassium aluminate salt, KAI(OH)4.
2 Al(s) + 2 KOH(aq) + 6 H2 O( )  2 KAl(OH)4 (aq) + 3H2 (g)
(8.2)
When the solution containing the potassium aluminate is reacted with sulfuric acid, H2SO4, two sequential
reactions occur. Initially, the KAl(OH)4 reacts with the acid in an acid/base neutralization reaction to give
a thick precipitate of aluminum hydroxide, Al(OH)3.
2 KAl(OH)4 (aq) + H2SO4 (aq)  2Al(OH)3 (s) + K 2SO4 (aq) + 2 H2O( )
(8.3)
As more acid is added to the solution, the precipitate of Al(OH)3 dissolves according to the following
reaction:
2 Al(OH)3 (s) + 3 H2SO4 (aq)  Al2 (SO4 )3 (aq) + 6 H2O( )
(8.4)
Both of these reactions are very exothermic - meaning they give off heat. If the final saturated solution
containing the aluminum sulfate, Al2(SO4)3, and potassium sulfate, K2SO4, is allowed to cool, crystals of
potassium aluminum sulfate dodecahydrate, KAl(SO4)2 ∙12H2O, slowly precipitate from the solution.
Al2 (SO4 )3 (aq) + K 2SO4 (aq) + 24H2 O( )  2 KAl(SO4 )2 •12H2 O(s)
(8.5)
The product is an alum often used for water purification, sewage treatment and in fire extinguishers. The
overall chemical reaction equation for its synthesis is:
2 Al(s) + 2 KOH(aq) + 4 H 2SO4 (aq) + 22 H2O( )  2 KAl(SO4 )2 •12H2O(s) + 3 H2 (g) (8.6)
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Experiment #8: Preparation of an Alum
EXPERIMENTAL PROCEDURE
1. Use the shears provided by the balance to cut off a piece of aluminum weighing about 1 gram. Cut
the large piece into about ten smaller pieces and place them in a 250 mL beaker. Record the mass
on the Laboratory Record.
2. Prepare a potassium hydroxide (KOH) solution by weighing out 10 grams of KOH pellets and
dissolving them in 40 mL of distilled water in a 250 mL Erlenmeyer flask. (Caution, KOH is
caustic - do not allow it to come in contact with your skin.) Occasionally swirl the mixture to
facilitate the solution and to prevent the KOH from forming a glass. (Caution: fumes, do the
remainder of this step in the hood.) When the crystals have dissolved entirely, carefully add the
KOH solution to the pieces of aluminum in the beaker. Stir the contents with a glass stirring rod.
If the reaction proceeds too slowly warm the contents of the beaker on a hot plate in the hood. Be
prepared (have a hot pad ready) to lift the beaker off the hot plate should frothing occur. If
the solution is heated too strongly the evolution of the hydrogen gas bubbles will produce a
frothing that will boil over. Turn off the hot plate as soon as any froth appears to be accumulating
more than about half an inch above the surface of the solution. Watch the reaction carefully since
the frothing can occur quite suddenly when the solution is being heated.
3. After all of the aluminum metal dissolves, continue heating the solution until the volume of the
liquid decreases to about 20 to 25mL. (The precipitation of the alum crystals in step 5 will occur
more rapidly if the volume of the water is decreased by boiling.)
4. Set up a filter funnel apparatus and obtain a piece of filter paper. Prepare the filter paper by
folding it in half and then folding it again not quite in half. Tear off a corner of the short side
(inside corner). Open the filter paper into a cone and insert it into the filter. Wet the filter paper
with distilled water so that it adheres to the sides of the funnel. Pour the warm solution through the
filter into a 100 mL beaker to remove any pieces of paint or other impurities present in the liquid.
After the filtration process is completed, discard the filter paper containing the paint and scum into
the waste paper trash can.
5. Cool the filtered solution for about five minutes in an ice bath. Then slowly add 30 mL of 6 M
sulfuric acid (H2SO4) while stirring the contents of the beaker. (Caution, do not allow the acid to
come in contact with your skin.) Observe and stir carefully as a precipitate forms.
6. Reheat the mixture until the precipitate dissolves completely. Place the beaker in an ice bath for
about 20 minutes to precipitate the alum crystals. If no crystals begin to form after about ten
minutes, scratch the inside of the beaker with a glass stirring rod to initiate crystallization or seed
the solution with a few crystals from the instructor or a neighboring lab group.
7. After the alum crystals have formed, weigh a piece of filter paper and collect the alum produced
by suction filtration. If some of the alum remains adhered to the bottom and sides of the beaker it
may be necessary to use a stirring rod to scrape the crystals from the beaker into the filter funnel.
Carefully scrape as many of the crystals into the filter as possible. Chill 10 - 15 mL of 95 %
ethanol and use it to rinse out any crystals that remain adhered to the beaker and then pour all the
alcohol and remaining crystals through the filter to rinse the crystals on the filter paper. Allow the
suction to run for several minutes to dry the crystals.
8. Let the crystals air-dry in your lab drawer until next week. You should complete the Laboratory
Record during the intervening week with the exception of the mass and mol of alum produced and
Experiment #8: Preparation of an Alum
-75-
the answer to question #3. When you return after one week, weigh the dried filter paper and the
crystals and complete the report. Deposit your yield in the designated container.
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Experiment #8: Preparation of an Alum
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Circle
Lab Day: M
Your Name:
T
W
R
F
_________________________________________
Partner’s Name:
_________________________________________
Lab Time: __________
EXPERIMENT #8: PREPARATION OF AN ALUM
Laboratory Record and Post-Lab
Reactant
Mass Used
mL used
Molarity (M)
Moles used
(Pay attention
to significant
figures)
Potassium hydroxide
Aluminum
Sulfuric Acid
Plenty – in excess
Water
Product
Mass Produced
Moles Produced
Alum
Hydrogen gas (calculated)
1. Use the data above and Equation (8.6) to determine the limiting reactant. Show your
reasoning.
2. Based on your calculation in question 1, how many grams of alum could theoretically be
produced in your reaction? Show your work.
3. Based on your calculation in question 2 and your actual yield from the data table, what is
your % yield in the reaction? Show your work.
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Experiment #8: Preparation of an Alum
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Experiment #9: Metathesis Reactions
Learning Objectives
Observe evidence of metathesis reactions
Identify spectator ions in solutions
Write net ionic equations for metathesis reactions
Textbook Reference (Chemistry: Atoms First,
Julia Burdge and JasonOverby, 2012)
Section 9.2 – Precipitation Reactions
Section 9.2 – Precipitation Reactions
Section 3.3 – Acids, Bases and
Neutralization Reactions
Section 9.2 – Precipitation Reactions
Section 3.3 – Acids, Bases and
Neutralization Reactions
Introduction
A very common reaction is a double displacement or metathesis reaction. Acid/base reactions are a
specific example of this type of reaction. The reaction is illustrated as follows:
AB + XY  AY + XB
(9.1)
where “AB” represents the formula of a cation “A” and anion “B” and “XY” represents the formula of a
cation “X” and anion “Y”. The reaction does not change even if the relative positive/negative portions are
multivalent (1+, 2+, 3+, 3-, 2-, 1-). The basic approach behind completing and balancing a metathesis
reaction is to first swap the “partner” ions of the cations (above A changes from having B as a “partner” to
having Y, and X changes from having Y as a “partner” to having B), write the correct chemical formulas
for each product based on ionic charges, and then balance the equation.
Evidence for the occurrence of double displacement reactions includes the formation of insoluble
products, the formation of covalent/nonionic substances (such as weak electrolytes and gases), and/or a
temperature change. (Solubility rules are listed in the appendices.) In this experiment you are to
investigate several reactions of this type, correlate them with solubility rules, predict formulas of the
products, write the formulas correctly, write reactions as balanced molecular equations, total ionic
equations, and net ionic equations.
Experimental Procedure
1. Get up to four ceramic 12-well spot plates from the common area. Depending upon class size and
spot plate availability, you may end up with less than four spot plates. This is not a problem – you
will just need to rinse and clean between runs.
2. If your spot plate does not already have a black stripe across the bottom of the wells, draw one
across each well with a permanent marker or grease pencil. The black provides contrast to the
white for purposes of observing light colored precipitates.
3. Each quadrant in the table on the first page of the Laboratory Record represents one spot plate.
Line up the available spot plates to match the arrangement shown in the table as much as possible.
4. Using the available reagents, put 2-3 drops of BaCl2 into each of the wells in the first column of
your spot plates. Repeat adding the chemical in each column heading to the appropriate well.
5. Starting with the first row, add 2-3 drops of 1.0 M NaCl to each cell across the row. As you add
the material, carefully look for precipitates, gases, color changes, or make any other observation
you feel would be helpful. Record this information on the first page of the Laboratory Record.
6. Repeat the additions across each row making and recording carefully your observations each time.
7. If you do not have enough spot plates to accommodate all of the mixtures, you may dispose of the
contents of a spot plate in the appropriate container, rinse it with water, and blot it dry for reuse.
8. As two final experiments, in separate empty wells place 3 drops of 0.3 M Na2CO3 and add 3 drops
of 6 M HCl in one well and 3 drops of 1.5 M Mg(NO3)2 and 3 drops of 3.0 M KOH in the other.
Record your observations at the bottom of your observation table.
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Experiment #9: Metathesis Reactions
9. Dispose of all the contents of your spot plates in the appropriate collector in the fume hood.
Rinse, dry, and return your spot plates to the common area.
Experiment #9: Metathesis Reactions
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Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
Experiment #9: Metathesis Reactions
Laboratory Record
Observations – these may include color change, precipitation, gas bubbles forming, or any other
observation you wish to squeeze in the small boxes below.
0.1 M
BaCl2
0.1 M CaCl2
0.1 M
Cu(NO3)2
0.1 M
Pb(NO3)2
0.3 M
Mg(NO3)2
0.1 M
Sr(NO3)2
1.0 M NaCl
0.3 M Na2CO3
0.1 M
Na2CrO4
0.1 M NaI
0.1 M Na2SO4
0.5 M HCl
0.3 M KOH
0.1 M NaIO4
Observations from mixing 3 drops of 6 M HCl and 3 drops
of 0.3 M Na2CO3:
Observations from mixing 3 drops of 1.5 M
Mg(NO3)2 and 3 drops of 3.0 M KOH:
Experiment #9: Metathesis Reactions
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Experiment #9: Metathesis Reactions
Experiment #9: Metathesis Reactions
Metathesis Reactions Post Lab
For the following reactions, write the formula for the expected products, the observations you made during the experiment, the formula for the product
that led to those observations, the molecular equation, total ionic equation, and net ionic equation. Be sure to write equations that include the physical
state (g, ℓ, s, or aq), the charges on the separated ions, and correct balancing.
Reactants
Formulas of Expected Products
Observations
Product Responsible for Observations
NaCl
Pb(NO3)2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
Na2CO3
BaCl2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
Na2CO3
CaCl2
Molecular Equation:
Total Ionic Equation:
Experiment #9: Metathesis Reactions
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Reactants
Formulas of Expected Products
Observations
Net Ionic Equation:
Na2CrO4
Cu(NO3)2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
Na2CrO4
Pb(NO3)2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
KOH
Pb(NO3)2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation
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Experiment #9: Metathesis Reactions
Product Responsible for Observations
Reactants
Na2SO4
Formulas of Expected Products
Observations
Product Responsible for Observations
BaCl2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
NaI
Sr(NO3)2
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
HCl
Na2CO3
Molecular Equation:
Total Ionic Equation:
Net Ionic Equation:
Experiment #9: Metathesis Reactions
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EXPERIMENT #10: MOLAR STOICHIOMETRY IN A CHEMICAL REACTION
Learning Objectives
Textbook Reference (Chemistry: Atoms First, Sylvia Burdge
and Jason Overby, 2012))
Observe an oxidation-reduction reaction
Relate stoichiometric coefficients to numbers of
moles
Determine number of moles from masses
Section 9.4 – Oxidation-Reduction Reactions
Section 8.3 – Calculations with Balanced Chemical
Equations
Section 5.10 – Molar Mass
INTRODUCTION
The objective of today’s lab is to verify the molar reaction stoichiometry of a chemical reaction. The
reaction being studied is the oxidation-reduction reaction between metallic copper atoms and silver
cations in aqueous solution. When a piece of copper metal is immersed in a solution containing dissolved
silver cations, the following electron transfer reaction occurs spontaneously.
Cu(s) + 2 Ag+ (aq)  Cu 2+ (aq) + 2 Ag(s)
(10.1)
The net result of the reaction is the transfer of two electrons from the copper atom to two silver cations.
In the process the copper dissolves into copper (II) cations while the silver cations precipitate out as
crystals of silver metal. From the balanced equation we see that for every one atom of copper that is
dissolved two silver atoms are formed. This ratio between the numbers of copper and silver atoms
involved in the reaction is constant so we can also say that the dissolution of one mole of metallic copper
must produce two moles of silver metal. We will try to experimentally confirm this molar stoichiometric
relationship by reacting a piece of copper metal with a solution containing silver cations and then
quantitatively determining the mass of copper metal consumed during the reaction and the mass of silver
metal formed. From the masses of the two metals involved in the reaction we can calculate how many
moles of atoms of each were reacted.
number of moles =
mass in g
molar mass
(10.2)
We can use this equation to determine the number of moles of silver precipitated and copper that
dissolved. According to Equation (10.1), the ratio of the moles of silver precipitated to copper metal is
2:1. The experimentally determined mole ratio may vary slightly due to impurities in the metals, but it
should be 2Ag : 1Cu when the numbers are rounded off to the nearest integers.
If the reaction is allowed to proceed to equilibrium, the silver metal will precipitate out of the solution
nearly 100%. In the experimental procedure that is provided below, the reaction is not allowed to occur
long enough to achieve complete displacement of the silver cations from the solution. We can calculate
the extent of the reaction by determining the percent yield of silver that we recover during the time that
the reaction between the silver cations and copper metal has occurred.
% silver recovered =
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mass of silver recovered
× 100%
starting mass of silver
(10.3)
Experiment #10: Molar Stoichiometry in a Chemical Reaction
EXPERIMENTAL PROCEDURE
A block diagram may be helpful for you to understand the processes you will be carrying out in
this experiment. The actual experimental steps are given as numbered steps below.
Fundamentally, you will be stripping electrons from silver atoms to make silver cations and then
taking back electrons from copper to reform the silver atoms.
Start with
silver metal
Ag
Silver
cations
Ag+
Steps 1, 2, 3
Step 4
Copper atoms (Cu)
with silver cations
(Ag+)
Steps 5 and 6
Copper ions (Cu+) and
reclaimed silver atoms
(Note: It is wise to wear gloves during this experiment. Silver ion solution that comes in contact with
your skin will cause a discoloration when exposed to sunlight. The effect is temporary and not
dangerous, but could be somewhat bothersome.)
1. Record accurately the mass of a silver sample of
about 1 g. With a 50 mL graduated cylinder, measure
15 mL of 6 M nitric acid and carefully pour the acid
into the 250 mL beaker. Caution, do not spill. In a
second beaker (400 mL) place approximately 75 mL
of tap water. Heat the water to boiling on a hot plate
in a fume hood. Turn off the hot plate when the water
boils. Place the 250 mL beaker into the 400 mL
beaker with the spouts not aligned (you are making a
double boiler) and add the silver sample to the acid.
Warm the Ag/HNO3 mixture until the Ag dissolves.
The reaction between the silver and nitric acid gives
off noxious fumes of NO2 gas, hence the need to use
the fume hood.
Double Boiler - Silver in inside beaker
Water in outside beaker
2. In this step any remaining nitric acid is neutralized to prevent excess from reacting with the
copper/silver reaction. The neutralization of the nitric acid can be achieved by reacting it with
sodium hydroxide solution. Prepare 75 mL of 1.5 M NaOH solution by weighing out 4.5 grams
of solid sodium hydroxide pellets and dissolving them in 75 mL of distilled water. The dissolution
of the sodium hydroxide occurs more rapidly if the solution is vigorously stirred. Use a graduated
Experiment #10: Molar Stoichiometry in a Chemical Reaction
-89-
plastic eyedropper to add 1 mL portions of sodium hydroxide solution to the solution containing
the dissolved silver. When the NaOH contacts the silver solution, a deep brown precipitate of
silver oxide forms initially but redissolves when the contents are swirled. To initially observe this,
you must hold the eyedropper tip close to the surface of the Ag+ solution touching the inside of the
beaker, release the NaOH such that it runs down the side of the beaker into the Ag+ solution.
Continue adding NaOH until the brown precipitate requires about 30 seconds of swirling to cause
the silver oxide to dissolve. If too much NaOH is added, the precipitate will not redissolve. If this
occurs, add drops of 6 M nitric acid until the silver oxide dissolves.
3. Following the neutralization of most of the nitric acid, the next step is to prepare the solution for
the addition of the copper metal to carry out the oxidation-reduction reaction between the silver
ion and copper. Add distilled water to the 250 mL beaker until it contains about 120 mL of
solution. This will dilute the acid remaining in the beaker so that the reaction between the nitric
acid and copper will not occur to any appreciable extent.
4. Use the wire cutters by the balance to cut off a piece of copper wire about eight inches long.
Polish the wire and wrap it into a loose coil by twisting it around a pencil and stretching it so it
looks like an extended spring. Weigh the copper wire and record its mass on your Laboratory
Record data sheet. Place the copper wire coil into the solution containing the dissolved silver and
allow it to remain undisturbed for about 20 minutes. Observe the contents of the beaker
periodically to witness the reaction between the copper and silver cations. Notice the formation of
the copper cations. The presence of copper cations is indicated by the blue color which is
produced as the copper metal dissolves. As you are waiting for the 20 minutes to expire, place
about 150 mL of water into a 400 mL beaker and start it heating on a hot plate in the fume hood.
This will be used as a boiling water bath in Step 7 to help dry out your silver.
5. Record the mass of a dry evaporating dish on the data sheet. After 20 minutes, carefully lift out
the copper wire which should have the silver crystals adhered to it. Use a wash bottle to rinse the
silver crystals off the wire and into the evaporating dish. Decant the supernatant from the beaker
and wash crystals remaining in the beaker into the evaporating dish. Place the supernatant and all
of the washings in the beaker provided. When all of the silver has been removed from the wire,
use a paper towel to dry the wire off and reweigh it. Record the mass of the copper wire after the
reaction on your data sheet. Place the copper wire in the container in the fume hood labeled
“Copper wires”
6. Decant the liquid from the evaporating dish into a beaker and place the solution into the
designated container. Add about 2 mL of methanol to the silver crystals in the dish. Place the
evaporating dish over the beaker containing the boiling water in the fume hood (set up in Step 4).
Use a scoopula or spatula to occasionally mix the silver up to expose new surface for drying.
When the methanol evaporates completely, dry any excess moisture off of the outside of the
evaporating dish. Weigh the dry evaporating dish with the dry silver crystals. Record this mass as
the “Mass after 1st Heating” on the data sheet. Reheat the evaporating dish, dry the outside of the
dish, and reweigh recording this as the “Mass after 2nd Heating”. Continue to reheat and weigh
until you get the same mass two times in a row. Pour the silver crystals into the beaker in the
fume hood and reweigh the empty evaporating dish.
7. Place the Ag+ /Cu++ solution in the beaker provided in the fume hood. Clean up your work area
and return the evaporating dish. Don’t put it in your drawer.
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Experiment #10: Molar Stoichiometry in a Chemical Reaction
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Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
EXPERIMENT #10: MOLAR STOICHIOMETRY IN A CHEMICAL REACTION
Laboratory Record
Due to time constraints you will only have time to do one trial. If possible, masses should be measured
on the analytical balances in the balance room.
Initial quantity of Ag
Mass (g)
Moles (mol)
_____________________
_____________________
Mass of evaporating dish and dry silver:
After 1st Heating: ______________________
Ag Information
After 2nd Heating: ______________________
After 3rd Heating: ______________________
______________________
Recovered quantity of Ag
_____________________
_____________________
Initial quantity of Cu wire
_____________________
_____________________
Recovered quantity of Cu wire
_____________________
_____________________
Quantity of Cu wire consumed
_____________________
_____________________
Moles Ag recovered / moles Cu consumed (from above): __________________
% Ag recovered
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_____________________
Experiment #10: Molar Stoichiometry in a Chemical Reaction
Cu Information
Mass of empty evaporating dish
EXPERIMENT #10:
MOLAR STOICHIOMETRY IN A CHEMICAL REACTION
Post-Lab
1. In this experiment, the copper is oxidized as it contributes electrons to the silver ions. The silver
ions, on the other hand, are reduced because they gain electrons. This transfer of electrons will
only occur spontaneously in one direction for this pair of elements – silver will not donate
electrons to copper. The activity series, found in the Appendix of this book as well as in textbook
provides a scheme for determining which elements will reduce other elements. In looking at the
activity series you will see that copper is above silver, so the copper reaction in the series remains
as written and the silver reaction is reversed. Based on the activity series, list three other elements
that can reduce silver ions.
a. __________________________
b. __________________________
c. __________________________
2. Explain the cause of the blue color you likely noticed in your final solution.
3. In writing oxidation-reduction reactions, another aspect of balancing equations comes into play
that we did not worry about with metathesis reactions. The charge on each side of a balanced
chemical reaction must be the same. In metathesis reactions this is automatically ensured as long
as the proper reactant and product formulas are written and the equation properly balanced.
Notice in equation 9.1 the “2”s in front of the silver are required to balance the charges – the
number of atoms would be balanced without those coefficients. For the following reactions, first
determine whether or not they occur using the activity series in the Appendix of this lab book (see
Question 1 at the top of this page). If no reaction occurs, just write “No Reaction”. If the reaction
does occur, complete and balance the equation.
a.
Co2+ + Al
b.
Cr + Pb2+
c.
Na+
d.
Mn
→
+ Cu
+
Ca2+
→
→
→
Experiment #10: Molar Stoichiometry in a Chemical Reaction
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Experiment #10: Molar Stoichiometry in a Chemical Reaction
EXPERIMENT #11: Determination of Acetic Acid in Vinegar
Learning Objectives
Work with molarity and volume to determine moles
Learn the terms titration, standard solution, equivalence
point, and end point
Use and read a burette properly
Standardize a sodium hydroxide solution
Determine the % acetic acid in vinegar
Textbook Reference (Chemistry: Atoms First,
Sylvia Burdge and Jason Overby, 2012)
Section 9.5 – Concentrations of Solutions
Section 9.6 – Aqueous Reactions and
Chemical Analysis
Section 9.6 – Aqueous Reactions and
Chemical Analysis
Section 9.6 – Aqueous Reactions and
Chemical Analysis
Introduction
Ordinary table vinegar typically contains between 4.0 and 5.5% acetic acid, HC2H3O2. At higher
strengths, such as 20%, it can be used as a weed killer. In this experiment we will use the method of
titration to determine the percent acetic acid in some commercial grade table vinegars.
A titration is an analytical procedure that takes advantage of the reaction
between two solutions, one with a known concentration or amount of solute
and the other with an unknown concentration or amount of solute.
Measurement of the volume of the solution with known concentration
required to react exactly with the solution of unknown concentration yields
sufficient information to deduce the unknown concentration.
A titration is carried out by placing one of the reactant solutions in a burette
and the other in an Erlenmeyer flask (Figure 11.1). The burette is designed
to deliver volumes with high accuracy. If one knows the molarity of the
solution in the burette, it is a straightforward calculation to use the solution
volume delivered in the burette to find the number of moles delivered (#
mol solute = M × Volume in L). This information can be coupled with the
balanced chemical equation to determine the number of moles of analyte in
the Erlenmeyer flask.
Ideally, in a titration one wishes to find the equivalence point of the reaction
– the point at which stoichiometrically equivalent volumes have been
reacted. In practice, the equivalence point of a titration is usually found by
the use of an indicator which changes color near the equivalence point. If
properly chosen, the indicator usually provides a close approximation to the
equivalence point and the point at which the color changes is called the end
point. Today’s titration will employ phenolphthalein as the indicator – it
will change from colorless to pink as a solution becomes basic.
Figure 11.1 – Titration
Apparatus (Photo Courtesy
of Mr. Herman Curtis)
Today’s experiment involves two parts. In the first part, a sodium hydroxide
solution is standardized (its concentration determined accurately) by reacting the sodium hydroxide
solution from a burette with potassium hydrogen phthalate (KHC8H4O4), also known as KHP, according
to Equation (11.1)
KHC8H4O4 (s) + NaOH (aq)  NaKC8H4O4 (aq) + HOH( )
Experiment #11: Determination of Acetic Acid in Vinegar
(11.1)
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Or, in net ionic form:
HC8 H 4O4- (aq) + OH -  C8 H 4O4 2- (aq) + HOH( )
(11.2)
If one carefully weighs the KHP, the number of moles of KHP reacted may be determined from its mass
and its molar mass (molar mass = 204.2). Since the NaOH and KHP react in a 1:1 mole ratio based on
equation (11.1), the number of moles of KHP reacted is equal to the number of moles of NaOH delivered
from the burette. Since molarity of the NaOH is defined as the number of moles per liter, it may be
determined by dividing the number of moles of KHP measured out by the volume in liters of NaOH
delivered.
The second part of the experiment involves using the standardized NaOH solution to determine the
concentration (and ultimately %) of acetic acid in a sample of vinegar. The chemical equation
representative of the reaction between acetic acid and NaOH is given as Equation (11.3).
HC2 H3O2 (aq) + NaOH (aq)  NaC2H3O2 (aq) + HOH( )
(11.3)
Or, in net ionic form:
HC2 H3O2 (aq) + OH - (aq)  C2 H3O2 - (aq) + HOH( )
(11.4)
If the volume of standardized NaOH required to neutralize the acetic acid in the sample of vinegar is
determined, the number of moles of NaOH may be easily determined from the molarity-volume
relationship. Since, as indicated in Equation (11.3), there is a 1:1 correspondence between the moles of
NaOH and acetic acid reacted, the number of moles of NaOH delivered is equal to the number of moles of
acetic acid present in the vinegar. Thus, the molarity of the acetic acid in the vinegar sample is equal to
the number of moles of NaOH delivered divided by the number of liters of vinegar sample titrated.
Procedure
Standardization
1. Bring three 125-mL Erlenmeyer flasks to the balance room.
2. Tare each flask on the analytical balance and add 0.55-0.65 g of KHP. The mass needs to be in
this range but be sure you record all of the decimal places in your Laboratory Record data sheet.
3. Return to the laboratory and place approximately 25 mL of deionized water into each flask. The
exact amount of water is not important since we only are concerned with the moles of KHP added.
Swirl each flask to accelerate the dissolution of the KHP.
4. Close the stopcock on the burette and fill the burette above the “0” line with NaOH solution. Run
some of the NaOH through the tip of the burette until the liquid level falls below the “0” line. (Do
not spend time trying to make sure the liquid level is right at zero – it is not important that it be at
zero.) Carefully inspect the tip of the burette for air bubbles. If air bubbles exist, remove the
burette from the burette clamp. Holding it over a waste beaker, open the stopcock and with one
motion shake the burette down and up once and it should clear the air bubble. If there is still an air
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Experiment #11: Determination of Acetic Acid in Vinegar
bubble, consult your instructor.
5. Add 3-5 drops of phenolphthalein to each of the Erlenmeyer flasks containing the KHP. The
phenolphthalein is the indicator and will change from clear to pink as one approaches the end
point in the reaction.
6. Record the initial volume reading on your burette to two decimal places. Tips on reading the
burette:
a. All burette readings must be recorded to two decimal places – even if they are “.00”.
b. Be sure to read the bottom of the meniscus. Holding a white piece of paper behind the
burette may make the meniscus more obvious.
c. Burettes read from “0” at the top to “50” at the bottom. Do not do anything exotic with the
reading – read the volume as it is in the burette.
7. Start adding NaOH from your burette into the flask in small portions. Add about 5 mL, swirl, and
observe how fast the pink color dissipates. Use a wash bottle to make sure all of the titrant is in
the bottom of the flask. Add another small portion of NaOH, swirl, and again observe how fast
the pink color dissipates. Continue adding portions of NaOH in smaller and smaller quantities as
the pink color lasts longer and longer. When the pink color stays for 10-15 seconds, add titrant
drop-by-drop. This is accomplished by forming a drop on the end of the burette, touching the drop
to the inside of the Erlenmeyer flask, and rinsing the drop into the bottom of the Erlenmeyer flask
with a wash bottle containing deionized water. With practice one can even deliver less than one
drop.
8. When a faint pink coloration remains for 1 minute the titration is complete.
9. Record the final volume of NaOH in the burette to two decimal places.
10. Refill your burette so the meniscus is between 0 and 1 mL. Repeat Steps 7-10 for the other two
Erlenmeyer flasks containing KHP. When completed, dispose of the contents of the Erlenmeyer
flasks in the usual disposal area.
Determination of % Acetic Acid in Vinegar
1. Pipette 3 mL of the vinegar sample into a 125 mL Erlenmeyer flask using a volumetric or Mohr
pipette. Add about 25 mL of deionized water to each flask.
2. Add 3-5 drops of phenolphthalein indicator to the Erlenmeyer flask.
3. Refill the burette with the standardized NaOH so the meniscus is between 0 and 1 mL. Record the
starting volume to two decimal places.
4. Titrate to the faint pink endpoint, using the techniques from the standardization portion of the
experiment.
5. Repeat the titration of vinegar two more times.
Experiment #11: Determination of Acetic Acid in Vinegar
-97-
Cleaning the burette
It is important to leave the burette clean when completed with your titration. Please follow these
steps.
1. Drain any remaining solution in your burette into the appropriate waste container. DO
NOT PUT IT BACK INTO THE ORIGINAL BOTTLE – ONCE A SOLUTION
COMES OUT OF A BOTTLE IT NEVER GOES BACK IN.
2. Mount the burette in its clamp and close the stopcock.
3. Place about 3-5 mL of deionized water in the burette.
4. Take the burette out of the clamp and rotate it nearly horizontally so the water coats all the
way around the burette.
5. Empty the water from the burette into the appropriate waste container and repeat steps 2-4.
6. When satisfied with its cleanliness, place the burette upside down in it clamp with the
stopcock open.
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Experiment #11: Determination of Acetic Acid in Vinegar
Circle
Lab Day: M
Your Name:
T
W
R
F
_________________________________________
Partner’s Name:
_________________________________________
Lab Time: __________
.
EXPERIMENT #11: Determination of Acetic Acid in Vinegar
Laboratory Record
Standardization
Trial
Mass KHP (g)
Moles KHP
(moles)
Final
burette
reading
(mL)
Initial
burette
reading
(mL)
Volume
NaOH
delivered
(mL)
M NaOH
Trial #1
Trial #2
Trial #3
Average Molarity:
Standard Deviation:
1. Show your work for the calculation of the molarity of NaOH from one of the trials above.
2. How many significant figures should your answers for molarity have? Explain.
Experiment #11: Determination of Acetic Acid in Vinegar
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Experiment #11: Determination of Acetic Acid in Vinegar
Laboratory Record
Molarity of your NaOH solution from the previous page: ______________
Vinegar Unknown Letter or Number: ______________________
Trial
Volume
Vinegar
(mL)
Final
burette
reading
(mL)
Initial
burette
reading (mL)
Volume NaOH
delivered (mL)
Moles NaOH
delivered
(moles)
Molarity
HC2H3O2 in
vinegar
Trial #1
Trial #2
Trial #3
Average Molarity:
Standard Deviation:
1. Calculations: Show your work for the calculation of the molarity of HC2H3O2 in vinegar from one
of the trials above.
2. The molarity indicates the number of moles of solute per liter of solution. Based on your average
molarity of the acetic acid in vinegar solutions, how many grams of acetic acid are contained in
one liter of vinegar solution? Show your work.
3. The percent acetic acid in vinegar is considered as the number of grams of acetic acid per 100 mL
of vinegar. Based on your calculations above, what is the % acetic acid in your vinegar sample?
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Experiment #11: Determination of Acetic Acid in Vinegar
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
_________________________________________
Lab Time: __________
_________________________________________
EXPERIMENT # 11: Determination of Acetic Acid in Vinegar
Post-lab
1. The following errors were made during the standardization of sodium hydroxide. How would the
calculated value of molarity compare to the actual molarity?
The KHP mass was transposed from
0.6605 g to 0.6505 g.
Increased
Decreased
Unchanged
The NaOH ran down the outside of the
burette during filling and some went
into the Erlenmeyer flask.
Increased
Decreased
Unchanged
An air bubble in the tip of the burette
comes out during the titration.
Increased
Decreased
Unchanged
The KHP contains a contaminant that
does not react with the sodium
hydroxide.
Increased
Decreased
Unchanged
The burette leaked sodium hydroxide
during the titration but the leakage did
not fall into the Erlenmeyer flask.
Increased
Decreased
Unchanged
Instead of using 20 mL to dissolve the
KHP a student used 15 mL.
Increased
Decreased
Unchanged
2. A NaOH solution has a concentration of about 0.0.12 M. How many grams of KHP would have to
be used to require approximately 25 mL of the NaOH solution to reach the endpoint? Show your
work.
3. It takes 22.75 mL of 0.0975 M NaOH solution to reach the end point when titrating a 25.00-mL
sample of an H2SO3 solution. What is the molarity of the H2SO3 solution? Show your work.
Experiment #11: Determination of Acetic Acid in Vinegar
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F
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EXPERIMENT #12: COMPARISON OF THE ENERGY
CONTENT OF FUELS BY COMBUSTION
Learning Objectives
Determine enthalpy of combustion through calorimetry
Textbook Reference (Chemistry: Atoms First,
Sylvia Burdge and Jason Overby, 2012)
Section 10.3 Enthalpy
Section 10.4 Calorimetry
Relate enthalpy of combustion to energy content of fuels
Consider differences in energy content of fuels based on
composition
Introduction
Liquid fuels are essential to our way of life. Hydrocarbons/hydrocarbon mixtures are clean burning
energetic fuels. Their liquid state makes them very useful for transportation purposes. In this experiment
you will investigate the energy content of several fuels by using them to heat water. The data generated
will enable you to compare fuels to see which ones are most energetic per unit mass burned and thereby
draw conclusions regarding the energy content of hydrocarbons and hydrocarbon-oxygenated fuels.
Combustion is simply the reaction between a fuel and oxygen. Heat generation is dependent upon the fuel
formula/structure and the quantity burned. Natural gas is a good example. Mainly methane (CH 4), it
burns to form carbon dioxide and water:
CH4 (g) + 2 O2 (g)  CO2 (g) + H2O(g)
(12.1)
Compounds that contain only hydrogen and carbon are called hydrocarbons. Commercial hydrocarbon
fuels include gasoline, kerosene, LNG, propane, and butane. When these are burned with air nitrogen
oxides are formed. The primary air pollutants generated in the automobile are the nitrogen oxides
generated because air (~80% nitrogen) is used instead of oxygen. Some fuels such as ethanol (ethyl
alcohol, C2H5OH) and methyl t-butyl ether (MTBE, called oxygenates) contain oxygen in addition to
carbon and hydrogen. These are proposed to reduce nitrogen oxides by lowering the burning temperature.
In this experiment, your class will measure the amount of heat generated by alcohols such as methanol
(CH3OH), ethanol(C2H5OH), n-propanol (C3H7OH), and n-pentanol (C5H11OH), and compare this energy
with hydrocarbon fuels such as hexane (C6H14), kerosene (app. C10H22), lamp oil (C12H26), and candle
wax (app. C40H82). The experimental procedure is to heat some water in an aluminum can by burning a
measured amount of a fuel sample. The specific heat of water is 4.184 J g-1 °C-1 or 4.184 J g-1 K-1,
indicating it takes 4.184 J to heat one gram of water by 1 °C or 1 K. Therefore, if the mass of water and
the number of degrees the temperature went up are known, the total amount of heat absorbed by the water
can be calculated as follows:
heat generated by burning fuel = heat absorbed by water + heat absorbed by aluminum can (12.2)
heat generated by burning fuel = s water  mwater  ΔTwater + saluminum  maluminum  ΔTaluminum (12.3)
T is a shorthand way of saying "change in temperature". The term swater is the specific heat of water and
the term saluminum is the specific heat of aluminum (0.9000 Jg-1K-1). (You should convince yourself that
combining and cancelling the units on the right side will leave only J.) Theoretically the amount of heat
liberated by the burning fuel should equal the heat absorbed by the water and the can, but in practice some
of the heat will be lost to the surroundings.
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
-103-
PROCEDURE:
Your instructor will tell you which fuel or fuels you are to investigate. You may be asked to do one trial
with each of three different fuels. Alternatively, you may be asked to do several trials with a single fuel to
improve the level of your confidence about an average value for that particular fuel, in which case the
results for the whole class will be assembled for comparison. It would be most helpful in drawing
conclusions if your class, at a minimum, could complete at least three alcohols and three hydrocarbons.
1.
Obtain a dry soda can with the pull tab intact. Wipe off any soot left from previous runs. Don’t
wiggle the pull tab - you do not want to break it off. Take the empty soda can plus your
Laboratory Record data sheet to a balance. Check to be sure that the empty balance is zeroed.
Weigh the can and record the mass to the nearest 0.01 gram.
2.
Using a beaker, add approximately 100 mL of water to the can. Reweigh the can plus water to the
nearest 0.01 gram. By subtraction, calculate the mass of water in the can. The volume is not
critical here - only the mass is critical.
3.
Suspend the can via a glass rod through the pull tab hole. Lay the rod
across an iron ring. This will suspend the can through the ring and you can
raise or lower the can by raising or lowering the ring. See the 11.1.
4.
Obtain a fuel burner, remove the cover, place it under the can, and adjust
the height of the ring so that the bottom of the can is about 2 centimeters
above the top of the wick. Replace the cover on the burner.
5.
Plug the stainless steel temperature probe into the LabQuest 2. Set the data
acquisition time for 480 seconds. Refer to Appendix A-4 for guidance.
6.
Suspend the stainless steel temperature probe above the can via a 3-prong clamp
such that the tip is immersed in the water in the can.
7.
Take the fuel burner plus your Laboratory Record data sheet to a balance. Be sure that the empty
balance reads zero. Weigh the burner and record its mass on the Laboratory Record data sheet,
reporting it to the nearest 0.01 gram.
8.
Start recording data by pressing the green arrow in the lower left-hand side of the LabQuest 2
screen.
9.
Remove the cover from the burner. Place the fuel burner under the soda can and light the burner
with a match. Observe the flame. If necessary, cautiously adjust the height of the can so that the
top of the flame is just below the bottom of the can. ACCURACY WARNING: YOUR
BURNER SHOULD SPEND NO TIME LIT IF IT IS NOT DIRECTLY UNDER THE CAN.
10.
Place a chimney around your burner to reduce the draft from the room air.
11.
Use the stainless steel temperature probe to stir the water occasionally and continue heating the
water until the temperature has increased by about 20 °C; then extinguish the flame.
12.
Continue stirring the water gently until the temperature indicated on the LabQuest 2 begins to
drop. Refer to Figure 12.2 (next page) for guidance on retrieving data from your LabQuest 2.
-104-
Figure 12.1 –
Experimental Setup
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
Figure 12.2 Example of Temperature-Time Data
13.
14.
Use the Statistics option on the LabQuest 2 (See Figure 11.2 and Appendix A-4) to determine the
initial temperature of the water in the can and record it as Tinitial on the data sheet.
15.
Use the Statistics option on the LabQuest 2 (See Figure 11.2 and Appendix A-4) to determine the
final temperature of the water – the maximum temperature it reached – and record it as Tfinal on the
data sheet.
16.
Print one plot similar to Figure 12.2 for your group. Include it with one of the reports of your
group.
17.
Wipe any soot off of the burner. Take the burner and your Laboratory Record data sheet to a
balance. Check that the balance reads zero and then weigh the burner and record the mass to the
nearest 0.01 gram. Calculate the mass of fuel burned by subtracting the final weight of the burner
from the original weight of the burner.
18.
Before doing another measurement, take a few moments to think about problems you might have
encountered. Can you think of any desirable improvements in the procedure? Communicate with
your lab partner about possible improvements.
19.
Now repeat the procedure, either making another measurement with the same fuel, or switching to
a different fuel as directed. Empty the can and add a fresh 100 mL portion of water. Note that you
recorded the mass of the dry empty can before you started. This should not change, except for
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
-105-
possible build-up of soot on the bottom (wipe it off); therefore you need to only measure the mass
of the can with water in it for subsequent trials.
-106-
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
This page is intentionally left blank.
Circle
Lab Day: M
Your Name:
T
W
R
_________________________________________
Partner’s Name:
Lab Time: __________
_________________________________________
EXPERIMENT #12: Enthalpy of Combustion of Various Fuels
Laboratory Record and Post-Lab
Measurement (Calculated)
Fuel
(
1
Fuel
Fuel
)
(
)
(
Specific heat of liquid water (Cs,water): 4.184 Jg-1K-1
Specific heat of Al (Cs,Al): 0.9000 Jg-1K-1
-108-
2
Mass of calorimeter and water (g)
3
Mass of calorimeter (g)
4
Mass of initial water (mwater)(g)
5
Final temperature of water (Tfinal) (°C)
6
Initial temperature of water (Tinitial) (°C)
7
Change in Temperature (T) (°C)
8
Initial mass of burner (g)
9
Final mass of burner (g)
10
Mass of fuel burned (g)
11
Joules gained by aluminum can (J)
(Cs,Al x maluminum x T)
12
Joules gained by water (J)
(Cs,water x mwater x T)
13
Total Joules gained (J)
14
Joules per gram of fuel (J/g)
15
Molar mass of fuel (g/mol)
16
Moles of fuel (mol)
17
kilojoules per mole of fuel (kJ/mol)
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
)
F
Laboratory Record and Post-Lab (continued)
Collect the data from the class and fill in as many of the blanks below as you can.
Fuel
formula
% Oxygen by mass
Heat of combustion in
Joules per gram
methanol
ethanol
n-propanol
n-pentanol
hexane
kerosene
lamp oil
candle wax
Questions:
1. How do the energies available from the alcohols compare to that available from the hydrocarbons on a
per gram basis?
2. Does the percent of oxygen in the fuel appear to make a difference? If so, why do you think this is the
case?
3. How large of a percent error would be introduced if one ignores the heat absorbed by the aluminum
can? Work out the calculation for one of your trials to illustrate the magnitude.
Experiment #12: Comparison of the Energy Content of Fuels by Combustion
-109
EXPERIMENT #13: DETERMINATION OF THE ENTHALPY OF FUSION OF ICE
Learning Objectives
Determine the calorimeter constant in a simple system
Use the conservation of energy principle to determine
missing information
Apply the concepts of specific heat capacity and heat
capacity to a simple calorimeter
Textbook Reference (Chemistry: Atoms First,
Julia Burdge and Jason Overby, 2012)
Section 10.4 - Calorimetry
Section 10.4 - Calorimetry
Section 10.4 - Calorimetry
Introduction
The enthalpy of fusion (also called the heat of fusion) is the heat necessary to change a unit quantity (1 g,
1 mol, etc.) of a solid substance at that substance’s melting (freezing) point to a liquid at the same
temperature. When ice is mixed with water, heat is transferred from the water (and its container) to the
ice resulting in the melting of the ice. For a system composed of ice placed in a known amount of water
in a calorimeter, the equation relating the heat gained by the ice and heat lost by the water and the
calorimeter can be written in the following form:
Heat absorbed by ice  Heat evolved by water  Heat evolved by calorimeter  0
(13.1)
or, in symbols:
qice total  qwater  qcalorimeter  0
(13.2)
where the algebraic sign of q is positive if heat is absorbed and negative if heat is evolved.
The term qice total in Equation (13.2) has two parts to it. One term accounts for the heat absorbed in heating
the ice from 0 °C to the final temperature of the mixture. The second term is the heat absorbed in melting
the ice, let’s call it qfus. Expanding Equation (13.2) to reflect these parts gives:
qice  q fus  qwater  qcalorimeter  0
(13.3)
If three of the four terms in Equation (13.3) are known, the fourth may be determined. In this experiment
the qice, qwater and qcalorimeter will be determined experimentally allowing the calculation of qfus. Division of
this qfus by either the mass of ice used or the number of moles used will yield the enthalpy of fusion of
water.
Determining the heat absorbed or evolved by a substance is not particularly difficult. Equation (13.4)
provides a simple approach to finding the heat flow given the specific heat capacity, mass, and
temperature change.
q  Cs  m  (T f  Ti )
(13.4)
where:
-110-
q = heat absorbed or evolved (negative if heat evolved, positive if absorbed)
Cs = specific heat capacity – amount of heat absorbed or evolved per degree change in
temperature per g of substance. For water, Cs = 4.184 J/g∙°C.
m = mass of the substance
Tf = final temperature
Ti = initial temperature
Experiment #13: Determination of the Enthalpy of Fusion of Ice
Finding qcalorimeter in Equation (13.3) is a little more complicated. One must conduct a calibration
experiment to determine this value, which will be unique for everyone’s calorimeter. The basic
calibration experiment is to mix together known quantities of hot and room temperature water and to
measure the final temperature. The guiding equation is very similar to Equations (13.1) and (13.2)
Heat evolved by hot water + Heat absorbed by room temperature water + Heat absorbed by calorimeter = 0
(13.5)
or, in symbols:
qhot  qRT  qcalorimeter  0
(13.6)
The calorimeter constant, also called its heat capacity, may be found from Equation (13.6) and the
relationship between heat and the calorimeter constant is given as Equation(13.4):
qcalorimeter  Ccalorimeter  (T f  Ti )
(13.7)
Procedure
Since the calorimeters employed in ordinary laboratory work (in our case they will be Styrofoam ® coffee
cups) do not provide complete thermal isolation, the transfer of heat between the calorimeter and the
surroundings is minimized by starting the experiment at a temperature slightly above room temperature.
An initial water temperature of about 5 oC to l0 oC above room temperature should result in a final
temperature of 5 oC to 10 oC below room temperature for the melting of ice.
Experiment #13: Determination of the Enthalpy of Fusion of Ice
-111-
B. Calibration
1. Be sure to record all data in the appropriate blanks of the data sheet. Set up the ring stand as
instructed, get out a hot plate and turn it on to a setting of about “5”. There is no particular
significance to the “5” setting - starting the hot plate at about that setting now saves time later.
2.
Connect the Vernier stainless steel temperature probe to the LabQuest 2 per the instructions in
Appendix A-4. Set the data acquisition time to 600 seconds.
3.
Weigh the empty calorimeter and record the mass on the data sheet.
4.
Place approximately 100 mL of room temperature water in the calorimeter and weigh the
calorimeter and water and record the mass on the data sheet. It is the mass that is critical, not the
volume. Approximately means approximately (using a beaker to measure is adequate).
5.
Place approximately 40 mL of additional water in a 100-mL beaker. Heat the water in the beaker
to boiling. (Turn the hot plate up as needed.)
6.
Once the hot water is boiling, press the green Start arrow in the lower left-hand corner of the
LabQuest 2. The LabQuest 2 will continue to acquire data until this run has been completed
(through Step 10)..
7.
Place the temperature probe into the room temperature water in the calorimeter and record its
temperature for about one minute. Do not stop data collection.
8.
Place the temperature probe into the boiling water and record its temperature for about one minute.
Do not stop data collection.
9.
Insert the probe into the hole in the lid of the calorimeter. Carefully (use a hot pad) pour the
boiling water into the calorimeter and put the lid on the calorimeter with the temperature probe
sticking into the mixed water.
Note: It is important to work quickly. Pour as fast as possible without spilling. Get the lid on.
Initially hold the lid on and swirl the cup a little. Continue to swirl the cup until the temperature
meets the criteria in Step 10..
10. Continue recording temperature data until the temperature plateaus for about 1-2 minutes or until
the temperature begins to drop, whichever comes first. Once the run is over, press the red Stop
button in the lower left-hand corner of the LabQuest 2.
11. Weigh the calorimeter and mixed water and record this information on the data sheet.
12. Use the statistics function to determine the average temperature of the room temperature water in
the plateau region (Figure 13-1 on the next page and See Appendix A-4), the average temperature
of the boiling water, and the maximum temperature of the mixed water. Record these
temperatures in the appropriate places on the data sheet. Print one copy of the LabQuest 2 screen
including the statistics and submit it for the group. Be sure to include the names of the members
of the group as a title on the printout.
13. From this data you will calculate the calorimeter constant, Ccalorimeter.
-112-
Experiment #13: Determination of the Enthalpy of Fusion of Ice
Ti of hot Water
Ti of RT Water
Tf of RT Water
and hot water
Figure 13.1 Appearance of LabQuest2 Data for the Calibration
B. Enthalpy of Fusion Determination:
1.
Place approximately 100 mL of water heated to about 35 ºC in the cup. If you do this part soon
enough after Part A, you can use 100 mL of your mixed water from there because its temperature
is probably in the mid-‘30’s. Determine the mass of the cup and water and record on the data
sheet.
2.
Use the temperature probe to measure the temperature of the water to the tenths place and record it
Ti for the water on the data sheet. Return the temperature probe to the hole in the calorimeter lid.
3.
Place the calorimeter on the balance. Dry about 12-15 g of ice chips. Quickly add the chips to the
cup and put the lid on. A careful recording of the total mass is not important at this stage.
4.
Start recording temperature data by pressing on the green arrow in the lower left-hand corner of
the LabQuest 2.
5.
Stir the contents of the cup as you continue to record the temperature with the LabQuest 2. You
may stop the run when the temperature plateaus for about 1-2 minutes.
6.
Use the statistics function (Appendix A-4) to record the minimum temperature in the plateau
region. This will be your final temperature, Tf, for both the water and the ice. Print one copy of
the LabQuest 2 including the statistics and submit it for the group. Be sure to include the names
of the members of the group aa a title on the printout.
7.
On the data sheet record the mass of the cup and water after melting the ice.
Experiment #13: Determination of the Enthalpy of Fusion of Ice
-113-
Circle
Lab Day: M
Your Name:
T
W
R
_________________________________________
Partner’s Name:
_________________________________________
Lab Time: __________
EXPERIMENT #13: DETERMINATION OF THE
ENTHALPY OF FUSION OF ICE
Laboratory Record
Calibration Data
Heat Evolved by Hot Water:
Heat Absorbed by Room Temperature (RT)
Water
Final mass of calorimeter and
mixed water (g)
Initial mass of calorimeter with
room temperature (RT) water (g)
Initial mass of calorimeter with
room temperature water (g)
Initial mass of empty dry
calorimeter (g)
Mass of hot water (g)
Mass of RT water (g)
Tf hot water (°C)
Tf of RT water (°C)
Ti hot water (°C)
Ti of RT water (°C)
Tf – Ti hot water (°C)
Tf – Ti RT water (°C)
(include algebraic sign)
(include algebraic sign)
qhot (J)
qRT (J)
(=4.184 J/g∙°C × mhot × (Tf – Ti))
(=4.184 J/g∙°C × mRT × (Tf – Ti))
Sum of qhot and qRT (pay attention to algebraic signs) (J):
qcalorimeter (same as previous line with opposite sign) (J)
Ccalorimeter (calorimeter constant, qcalorimeter divided by its temperature change which is
the same as Tf -Ti for the RT water) (J/°C):
-114-
Experiment #13: Determination of the Enthalpy of Fusion of Ice
F
Determination of the Enthalpy of Fusion of Water:
Heat Evolved by Water, qwater:
Heat Absorbed by Ice, qice:
Initial mass of calorimeter, lid, and
water (g)
Final mass of calorimeter, lid,
water, and melted ice (g)
Mass of calorimeter and lid (g)
(from other side)
Initial mass of calorimeter, lid, and
water (g)
Mass of water (g)
Mass of ice (g)
Tf (°C):
Tf (°C):
Ti (°C):
Initial temperature of ice (°C):
Tf – Ti for water (°C):
Tf – Ti (°C):
qwater (J)
qice (J)
(4.184 J/g∙°C × mwater × (Tf –Ti)water)
(4.184 J/g∙°C × mice × (Tf –Ti)ice)
0 °C
Heat Evolved by Calorimeter, qcalorimeter:
Final temperature of mixture (°C):
Initial temperature of calorimeter
(°C):
Ccalorimeter (J/°C) (from other side):
qcalorimeter (J):
(Ccalorimeter ×(Tf – Ti)calorimeter)
Determination of Enthalpy of Fusion of Ice:
qice + qwater + qcalorimeter =
Recall from Equation (13.3) that qice + qfus + qwater + qcalorimeter = 0, so what does the
value of qfus have to be based on your experiment?
The enthalpy of fusion of ice may be expressed as the number of J/g. Based on your
value for qfus and the mass of ice used, what is your experimental enthalpy of fusion of
ice?
What is the enthalpy of fusion of ice on a per mole basis?
The accepted value for the enthalpy of fusion of ice is 335 J/g. Based on this information and your
results, what is the percent error between your experimental enthalpy of fusion and the accepted value?
(Note: For percent error
% error =
experimental value - accepted value
×100% )
accepted value
Show your work:
Experiment #13: Determination of the Enthalpy of Fusion of Ice
-115-
-116-
Experiment #13: Determination of the Enthalpy of Fusion of Ice
EXPERIMENT #13: DETERMINATION OF THE
ENTHALPY OF FUSION OF WATER
Post Lab
1. Following is a list of laboratory errors. Read them and decide how each error would affect the results
of this analysis. Circle the answer in each case.
The balance was reading +0.55-g when a student walked up to it
and she forgot to zero it before weighing the mass of the
calorimeter with the room temperature water in it for calibration
purposes. The student had previously recorded the mass of the
empty dry calorimeter correctly. How would this error affect the
initial mass of water compared to its actual value?
Increased Decreased Unchanged
The above error was not caught and the student proceeded to
calculate the value of the calorimeter constant. How would the
calculated calorimeter constant compare to its actual value? (Hint:
Consider the effect on the calculated qRT and its role in the
relationship qh + qRT + qcalorimeter. Remember the calculated value of
qh does not change.)
Increased Decreased Unchanged
If the error in the first statement was still not caught, how would the
calculated enthalpy of fusion of water compare to the actual value?
(Another hint: The only value in the expression qice + qwater +
qcalorimeter that might be affected is qcalorimeter. )
Increased Decreased Unchanged
A student let the ice sit at room temperature a long time before
adding it to the calorimeter cup and contents. How would the
calculated enthalpy of fusion of water compare to the actual value?
Increased Decreased Unchanged
2. When the water in an ice cube tray freezes it releases heat to the surroundings. Using your value for
the enthalpy of fusion of ice, how much heat is released if 200-mL of room temperature water freezes?
Show your work
Experiment #13: Determination of the Enthalpy of Fusion of Ice
-117-
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-118-
Experiment #13: Determination of the Enthalpy of Fusion of Ice
EXPERIMENT #14: DETERMINING THE MOLAR MASS
OF A VOLATILE LIQUID BY THE DUMAS METHOD
Learning Objectives
Textbook Reference (Chemistry: Atoms First, Julia
Burdge and Jason Overby, 2012)
Work with the ideal gas law
Determine the molecular weight from the gas
law and sufficient measurements
Section 11.5 – The Ideal Gas Equation
Section 11.5 - The Ideal-Gas Equation
INTRODUCTION
One of the earliest methods of molecular mass determination was developed by Dumas’ to determine the
molecular mass of volatile liquids. Dumas’ method takes advantage of the ideal gas law and requires that
it be possible to convert the liquid to a gas at a measurable temperature and pressure without it
decomposing. Consider the ideal gas law:
PV  nRT
(14.1)
where P is the pressure, V the volume, n the number of moles, R the gas constant, and T the temperature
in K. If one conducts an experiment in which any three of the four variables are measured, the fourth may
be calculated. In the case of the Dumas method, P, V, and T are measured. From that information, the
number of moles of gas, n, may be determined. Then a simple determination of the mass, m, of a gas
occupying the volume will lead to the molar mass.
Molar Mass 
mass of gas
number of moles
(14.2)
PROCEDURE
1. Set a hotplate out and put it on a setting of about 5 on its
numerical scale. You do not want it too hot for this
experiment.
Al foil cap with pin hole
2. Get your unknown from your instructor and record the sample
number or code in the blank provided on the Laboratory
Record.
3. Place a boiling bead and a very small quantity of Sudan III or
a similar dye into a clean 250 mL Erlenmeyer flask. Fit the
flask with an aluminum foil cap. Minimize your handling of
the cap to keep from adding mass. Record the mass of the
flask and its contents.
4. Assemble the apparatus shown in Figure 15.1. Use your 800
mL beaker and place about 500 mL of water in the beaker. Be
prepared to remove water if it becomes too high when you
immerse the Erlenmeyer flask into the beaker. If the water
Figure 15.1 – Experimental Setup
level does not cover the Erlenmeyer flask fairly completely,
add more water to the 800 mL beaker until you have covered as much of the Erlenmeyer flask as
possible.
Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method
-119-
5. Place 5 mL of your unknown in the Erlenmeyer flask and
cover with the aluminum cap.
6. Pierce the cap with a pin to allow gas to escape.
7. It is important to have as much of the flask immersed as possible, but the cap needs to remain out
of the water.
8. Turn on the hot plate and boil the water. Watch the unknown in the flask and when it is all gone,
record the temperature of the thermometer, move the thermometer out of the way and raise the
clamp as quickly as possible, removing the flask from the boiling water. Rotate the clamp so that
the flask is not above the steam and turn the hot plate off. Find the barometer in the laboratory
and read and record the atmospheric pressure.
9. Allow the flask to cool to room temperature. You will observe that the flask looks empty when it
is removed from the water bath but that as it cools liquid is condensed inside.
10. Blot the flask and cap dry of any water and reweigh the flask, cap and boiling bead. Record this
new value. The difference between this mass and the dry mass of the Erlenmeyer flask is the mass
of the liquid that occupied the entire volume of the flask at the boiling point of the water.
11. Measure the volume of your Erlenmeyer flask by filling it to the brim with water and then
carefully transferring that water to a large graduated cylinder. Use an eyedropper until the volume
is low enough to pour without spilling. If you spill during the measurement, just empty the
graduated cylinder and refill the flask and start measuring again.
12. If possible, share your results with a group that had the same unknown number as you.
Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method
-120-
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
_________________________________________
_________________________________________
Lab Time: __________
EXPERIMENT # 14: DETERMINING THE MOLAR MASS
OF A VOLATILE LIQUID BY THE DUMAS METHOD
Laboratory Record
Trial 1
Trial 2
(from a second group)
Mass of Erlenmeyer flask, bead, cap, dye,
and condensate
________________ g
________________ g
Mass of empty Erlenmeyer flask, bead,
dye, and cap
________________ g
________________ g
Mass of condensate
________________ g
________________ g
Temperature of hot water
________________ °C
________________ °C
________________
________________
________________ mL
________________ mL
________________
________________
________________
________________
________________
________________
Conversion of volume to units consistent
with your choice of gas constant
________________
________________
Molar mass of unknown
________________
________________
Barometric pressure (include units)
Volume of Erlenmeyer flask
Your choice of gas constant, with units
Conversion of temperature to units
consistent with your choice of gas
constant
Conversion of barometric pressure to
units consistent with your choice of gas
constant
Average molar mass from two trials
________________
Conclusion : The molar mass of unknown # _______ is ________ g/mol.
Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method
-121-
F
EXPERIMENT # 14: DETERMINING THE MOLAR MASS
OF A VOLATILE LIQUID BY THE DUMAS METHOD
Post - Lab
Read the following list of possible mistakes that could be made during this experiment. Circle the
appropriate answer in each case.
A student misrecorded the mass of the final
flask by recording 123.76 g when it actually had
a mass of 120.56 g. How would the calculated
molecular weight of the sample compare to the
actual molecular weight?
Increased
Decreased
Unchanged
A second student did not dry the flask well after
condensing the vapor and proceeded to calculate
the molar mass of the sample How would the
calculated molecular weight of the sample
compare to the actual molecular weight?
Increased
Decreased
Unchanged
There was liquid remaining in the flask when it
was determined it was actually all evaporated.
How would the calculated molecular weight of
the sample compare to the actual molecular
weight?
Increased
Decreased
Unchanged
The pressure was misread as 765 mm Hg
instead of 755 mm Hg. How would the
calculated molecular weight of the sample
compare to the actual molecular weight?
Increased
Decreased
Unchanged
Increased
Decreased
Unchanged
In an interesting development, the liquid
decomposes into other gases when heated
without the experimenter’s knowledge. For
each mole of liquid in the flask three moles of
gaseous products are produced. How would this
situation affect the calculated molecular weight
of the sample compared to the actual molecular
weight?
-122-
Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method
EXPERIMENT #15: Regularities in the Properties of the Elements
Learning Objectives
Become familiar with trends in properties in the
periodic table
Learn the use of Periodic Table Live! as a means
of exploring elements and their properties
Use logic to define searches of large quantities of
data
Textbook Reference (Chemistry: Atoms First, Sylvia Burdge
and Jason Overby, 2012)
Section 4.4 Periodic Trends in Properties of
Elements
J. Chem. Educ. 2008, 85, 22.
http://www.chemeddl.org/collections/ptl/
Introduction
Each element has unique properties that distinguish it from any other element. Although no two elements
are alike, certain similarities and trends among some of the elements are observed. The modern Periodic
Table was developed as a scheme to link elements together.
A general method of classification of elements is the distinction between metals and nonmetals. Since
there are so many metals, the classification of metals and nonmetals is not extremely helpful. As
discoveries of properties of elements were made, similarities between two or three elements could be
observed. The most important compilation for elements was that of atomic weights.
During the development of the Periodic Table, scientists proposed that the elements must be arranged in
repeating groups with similar properties. In 1869, Mendeleev and Mayer arrived at similar classifications
of elements dependent on the periodicity of properties – the appearance of similar properties at regular
intervals.
As the level of understanding of atomic-level properties has increased, it has been observed that the
periodic table, created originally based on bulk properties, is an effective means of predicting and
understanding the importance of microscopic properties and their contributions toward bulk properties.
Thus, we typically look at atomic-level properties such as atomic radius, ionic radius, ionization energy,
electron affinity, and electronegativity and their periodic changes in the chart as a starting point.
A major advantage of the Periodic Table is its usefulness in systematizing chemistry. The Periodic Table
groups elements together to show similarities and differences, and to reveal trends in properties. The
currently adopted (by IUPAC – International Union of Pure and Applied Chemistry) numbering system
places the numbers 1-18 across the groups. However, two other systems are still in use that rely on
Roman numerals and the letters A and B. The back cover of the lab book, as well as the inside front flap
of your textbook, have periodic tables using the system that uses the letter “A” for representative elements
and “B” for transition elements. Some numbering systems use the letter “A” for the first groups
numbered I – VIII followed by “B” for the remaining columns.
Procedure
The study of properties will be conducted here through the use of the web site Periodic Table Live!
(http://www.chemeddl.org/resources/ptl/) produced by the Journal of Chemical Education Online and the
ChemEd Digital Library. This package allows one to efficiently look at a wide range of atomic-level and
bulk properties of elements as well as several other pieces of information such as toxicology, uses, and
several others. Data can be selectively searched and any pairing of numerical properties may be graphed
for comparison purposes.
Experiment #15: Regularities in the Properties of Elements
-123-
The opening window of the software is as shown in Figure 15.1. Descriptions are provided below for
each of the key aspects of the opening window.
Figure 15.1 Opening Window for Periodic Table Live!
The key operations that can be conducted from the opening window include:
Clicking on an elemen:
A click on any element in the periodic table will lead to a wealth of information regarding that element
including a description, physical and atomic properties, images including still images and movies of
reactions, and information regarding the discovery of the element.
About PTL!
This button (upper right) provides information about components needed to use Periodic Table Live! to its
full extent and gives acknowledgements.
Chart/Sort
This option allows one to do searches on elements and their properties and to make plots of one property
vs. another. Further information is below.
Biographies
This button provides biographical information on a large number of people involved in the periodic chart
and its contents.
Glossary
This is what is sounds like – a glossary of terms related to properties.
Preferences
This allows one to set a preference as to the numbering system for groups desired.
-124-
Experiment #15: Regularities in the Properties of Elements
More information on the Chart/Sort button:
1. Using a one-by-one approach would be a laborious method for exploring the periodic properties of
the elements. Rather than using that approach, one can take advantage of the Chart/Sort button at
the upper right-hand side of Figure 15.1. When clicked, the screen shown in Figure 15.2 comes
up.
Figure 15.2 Chart/Sort Window
2. The opening screen of Figure 15.2 has three main sections. The upper-left corner has a periodic
table with selectable elements. Entire categories of elements (Main Group, Transition,
Lanthanide, Actinide, Metal, and Nonmetal) or individual elements may be selected upon which to
search properties. The upper-right hand corner allows for the graphing of any numerical property
as a function of any other numerical property. The lower-left corner has a table which allows the
sorting of values for up to three numerical properties.
3. The first task is to produce a graph of the atomic radius of the elements versus the atomic number.
Click on the X button at the upper-left corner of the graph panel of Figure 15.2. Select the atomic
number (under atomic properties). Click on the Y button at the upper-left corner of the graph
panel of Figure 15.2. Select the Atomic radius (under atomic properties). Figure 15.3 shows the
resulting graph.
Experiment #15: Regularities in the Properties of Elements
-125-
Figure 15.3 Atomic Radius vs Atomic Number
The graph in Figure 15.3 has quite a bit of information in it. Smaller sections of the data may be
looked at by narrowing down the number of points plotted. Suppose one wanted to plot only the
first 36 elements. Click the Options button at the top of the graph panel and choose Search for
elements whose Atomic Number is less than 37. Notice the search box allows for two conditions
logically connected by “AND” or “OR” statements.
4. A similar search box exists for the Table in the lower left-hand side of the panel of Figure 15.3.
This experiment consists of working through the Laboratory Record data sheet and providing
graphs and answers to questions posed about elements on the periodic chart.
-126-
Experiment #15: Regularities in the Properties of Elements
Circle
Lab Day: M
Your Name:
Partner’s Name:
T
W
R
F
_________________________________________
_________________________________________
Lab Time: __________
Experiment #15: Regularities in the Properties of the Elements Laboratory Record
1. Make and print a graph (attach copy) of the atomic radius vs. the atomic number for elements up
through atomic number 36.
2. Make and print a graph (attach copy) of the first ionization energy vs. the atomic number for
elements up through atomic number 36.
3. Make and print a graph (attach copy) of the first ionization energy vs. the atomic radius for
elements up through atomic number 36. Do you see a trend in the change in ionization energy
with atomic radius? If so, explain the reason for this trend.
4. Make and print a graph (attach copy) of the electron affinity vs. the atomic number for elements
through atomic mass number 36.
5. Make and print a graph (attach copy) of the electronegativity vs. the atomic number for elements
up through atomic number 36.
6. Make and print a graph (attach copy) of the ionization energy and electron affinity vs. the
electronegativity. Two Y values may be selected by going to the “Y” button above the graph
panel and selecting “Multiple Fields”. Then select both the ionization energy and the electron
affinity. Is there a relationship evident between the ionization energy and electron affinity and the
electronegativitiy? If so, explain the reason for the trend you see.
7. Using the Search option in the Table panel, what elements were discovered before the year 1700?
8. Which elements are liquids at 25 °C? (Note – this is basically a question about which elements
boil below 25 °C or, those that have a physical state of liquid at 25°C . All temperatures are in K
in the software.)
Experiment #15: Regularities in the Properties of Elements
-127-
9. List the elements that are used in steel.
10. List the elements that are used in rocket fuel.
11. List the elements that are used in both steel and rocket fuel.
12. What is the most expensive element that lists some use.
13. What are the two hardest elements?
14. What elements are used in transistors?
15. What elements are used in fertilizers?
16. What is the highest boiling element?
17. What is the cost in $/100g of the most expensive element used in transistors?
18. What elements are indicated to undergo a vigorous reaction with 15 M HNO3?
19. What elements are found in a range of 0.1 % to 0.5% in the solar system?
20. What elements are found in excess of 7.0% in the Earth’s crust?
-128-
Experiment #15: Regularities in the Properties of Elements
Experiment #15: Determining the Molar Mass of a Volatile Liquid by the Dumas Method
-129-
Appendices
Basic Error Analysis
Several of the experiments in this book require multiple determinations of a value in an attempt to judge both the
accuracy and precision of the technique. This appendix gives a brief overview of the error analysis important for
CHEM 1361 lab. For further information, refer to your lecture textbook, Section 1.5 and Appendix A (p. 1110).
Background
In our experimental attempts to determine a particular physical value, we are really attempting to find the “true” or
“actual” value. Though we may never know the “true” or “actual” value we are able to draw some level of confidence
in our proximity to that value through statistical means. There are two terms that are important in analyzing repeated
attempts to find a “true” or “actual” value.
Accuracy – this is a measure of how closely experimental results approach the “true” or “actual” results
Precision - this is a measure of how closely repeated measurements agree with one another
The approach to accuracy – nearness to the “true” or “actual” value - is typically discussed in terms of confidence, a
topic that will not be covered here. More importantly, we discuss the precision to a first approximation in terms of a
quantity called the standard deviation. The mathematical form of determining the standard deviation is given is
equation (14.3)
N
s
 ( x  x)
i 1
2
i
N 1
(14.3)
where s is the standard deviation, xi represents the ith experimentally determined value; x is the average of all of the
experimental values; and N represents the total number of trials being considered.
As an example, suppose a student measured the molarity of a NaOH solution three times and got the three numbers:
0.1035 M, 0.09826 M, and 0.09956 M. (You don’t need to know what molarity is yet – we are just considering the
numbers.) The average of these three numbers is:
0.1035M  0.09826 M  0.09956 M
 0.1004 M
3
The standard deviation is given as:
s
(0.1035  0.1004) 2  (0.09826  0.1004) 2  (0.09956  0.1004)2
 0.002786M
(3  1)
To report the molarity with the standard deviation, one typically rounds the standard deviation to one significant digit
and maintains the reported average to that decimal place. In this case, one would report 0.100 ± 0.003 M. The
precision is pretty good – about 3% (0.003/0.100 *100). Much more substantial significant tests may be performed but
we will not go beyond the standard deviation.
In terms of accuracy, we would not know how accurate the result is unless we knew the “true” or “actual” value. If the
true value was 0.125 M, we would appear to be precise (about 3%) but not accurate. If, on the other hand, the true value
was 0.101 M, we would have a result that is fairly accurate and fairly precise.
-130-
Appendices
Activity Series of Metals in Aqueous Solution
Metal
Lithium
Potassium
Barium
Calcium
Sodium
Magnesium
Aluminum
Manganese
Zinc
Chromium
Iron
Cobalt
Nickel
Sn
Lead
Hydrogen
Copper
Silver .
Mercury
Platinum
Gold
Oxidation Reaction
(most easily oxidized)
Li
➔
Li+
+
+
K
➔
K
+
2+
Ba
➔
Ba
+
2+
Ca
➔
Ca
+
Na
➔
Na+ +
Mg
➔
Mg2+ +
Al
➔
A13+ +
Mn
➔
Mn2+ +
Zn
➔
Zn2+ +
Cr
➔
Cr3+ +
Fe
➔
Fe2+ +
Co
➔
Co2+ +
Ni
➔
Ni2+ +
Sn
➔
Sn2+ +
Pb
➔
Pb2+ +
H2
➔
2H+ +
Cu
➔
Cu2+ +
Ag
➔
Ag+ +
Hg
➔
Hg2+ +
Pt
➔
Pt2+
+
3+
Au
➔
Au
+
(Most easily reduced)
Appendix A-2
1e1e2e2e1e2e3e2e2e3e2e2e2e2e2e2e2e1e2e2e3e-
Potential
Eo in volts
3.04
2.931
2.912
2.868
2.71
2.377
1.662
1.185
0.762
0.744
0.489
0.280
0.257
0.140
0.126
0.00
-1.17
-0.80
-0.851
-1.118
-1.498
-131-
BASE UNITS IN THE SI SYSTEM, SCIENTIFIC PREFIXES FOR MAGNITUDES, AND
CONVERSION FACTORS RELATING LENGTH, VOLUME, AND MASS UNITS
TABLE I: UNITS OF MEASUREMENT
QUANTITY
Mass
Length
Volume
Time
Temperature
Amount
SI UNIT
kilogram
meter
cubic meter
second
Kelvin
mole
SYMBOL
kg
m
m3
s
K
mol
COMMONLY USED
gram
centimeter
liter
SYMBOL
g
cm
L
degrees Celsius
C
TABLE II: FREQUENTLY USED SCIENTIFIC PREFIXES FOR MAGNITUDES
PREFIX
SYMBOL
giga
mega
kilo
deci
centi
milli
micro
nano
G
M
k
d
c
m
μ
n
NUMBER
1,000,000,000
1,000,000
1,000
1
0.1
0.01
0.001
0.000001
0.000000001
EXPONENTIAL
NOTATION
109
106
103
100
10-1
10-2
10-3
10-6
10-9
TABLE III: CONVERSION FACTORS FOR LENGTH, VOLUME AND MASS UNITS
Length
METRIC
1 km = 103 m
1 cm = 10-2 m
1 mm = 10-3 m
1 nm = 10-9 m
1 A = 10-10 m
ENGLISH
1 ft = 12 in
1 yd = 3 ft
1 mile = 5280 ft
METRIC-ENGLISH
2.54 cm = 1 in
39.37 in = 1 m
1.609 km = 1 mile
Volume
1 mL = 1 cm3
1 L = 1000 mL
1 m3 = 1000 L
1 gal = 4 qt
1 qt = 2 pt
1 pt = 16 fl oz
1 L = 1.057 qt
28.32 L = 1 ft3
Mass
1 kg = 1000 g
1 mg = .001 g
1μ g = 10-6 g
1 m3 = 35.31 ft3
1 lb = 16 oz
1 ton = 2000 lb
-132-
Appendix A-2
453.6 g = 1 lb
1 g = .03527 oz
CONCENTRATIONS OF COMMON CONCENTRATED ACIDS AND BASES
Reagent Solution
Chemical
Formula
Formula
Mass
Density g/mL
(20 oC)
Molarity (M)
(mol/liter)
% wt/wt
Acetic acid
CH3CO2H
60.05
1.05
17.5
100
Ammonium Hydroxide
NH4OH
35.05
0.90
14.5
57
Formic acid
HCO2H
46.03
1.20
23.6
90.5
Hydroiodic acid
HI
127.91
1.70
7.6
57
Hydrobromic acid
HBr
80.92
1.49
8.8
48
Hydrochloric acid
HCl
36.46
1.19
12.1
37
Hydrofluoric acid
HF
20.00
1.18
28.9
49
Nitric acid
HNO3
63.01
1.42
15.9
70
Perchloric acid
HC1O4
100.47
1.67
11.7
70
Phosphoric acid
H3PO4
97.10
1.70
14.8
85
Potassium hydroxide
(soln)
KOH
56.11
1.46
11.7
45
Sodium hydroxide
(soln)
NaOH
40.00
1.54
19.4
50
Sulfuric acid
H2SO4
98.08
1.84
18.0
96
Appendix A-2
-133-
Some Common Ions
Positive
Negative
Li+
Lithium
Cu+
Na+
Sodium
Hg22+
Potassium
Rubidium
Cesium
Hydrogen
Ag+
NH4+
K+
Rb+
1+ Cs+
H+
Copper (I) or
Cuprous
Mercury (I) or
Mercurous
Silver
Ammonium
F-
Fluoride
ClO-
Hypochlorite
Cl-
Chloride
ClO2-
Chlorite
Bromide
Iodide
Nitrate
Nitrite
Acetate
Permanganate
ClO3ClO4IO4OHSCNHCO3-
Cyanide
Hydride
O22-
Chlorate
Perchlorate
Periodate
Hydroxide
Thiocyanate
Hydrogen Carbonate or
Bicarbonate
Peroxide
Oxide
S2-
Sulfide
Sulfate
SO32-
Sulfite
Carbonate
C2O42-
Oxalate
Chromate
Cr2O72- Dichromate
BrI1- NO3NO2C2H3O2MnO4CNH-
Mg2+ Magnesium
Cu2+
Ca2+
Calcium
Hg2+
Sr2+
Strontium
Sn2+
Barium
Iron (II) or
Ferrous
Cadmium
Zn2+
Pb2+
2+ Ba2+
Fe2+
Cd2+
Cr2+
Al3+
Cr3+
3+
Fe3+
Mn2+
Chromium (II) or Co2+
Chromous
Aluminum
As3+
Chromium (III) or Co3+
Chromic
Iron (III) or
Ferric
Copper (II) or
O2Cupric
Mercury (II) or
SO42Mercuric
Tin (II) or
CO32Stannous
Zinc
2- CrO42Lead (II) or
Plumbous
Manganese (II) or
S2O32Manganous
Cobalt (II) or
Cobaltous
Arsenic
N3Cobalt (III) or
PO43Cobaltic
3AsO43BO33-
-134-
Appendix A-3
Thiosulfate
Nitride
Phosphate
P3PO33-
Phosphide
Phosphite
Arsenate
AsO33-
Arsenite
Borate
Appendix A-4. Basic Use of the Vernier LabQuest2
Some of the experiments in this lab book refer to the use of the Vernier LabQuest 2 handheld data acquisition
device and its associated sensors. This Appendix is intended to serve as a primer in the use of that device - its
capabilities extend far beyond what is indicated in this section. If you would like to learn more about it, please
ask your instructor for further information. If you get really carried away, you could look up the LabQuest 2
manual at http://www2.vernier.com/manuals/labquest2_user_manual.pdf.
Appendix A.4 is separated into the following parts:
I.
II.
III.
IV.
V.
VI.
VII.
I.
Powering up and turning off the LabQuest 2
Plugging in and setting up sensors
Modifying Data Collection settings
Starting data collection
Saving data
Analyzing data
Printing data
Powering up and turning off the LabQuest 2
The LabQuest 2 is powered on by holding down the button with the red icon on the top of the device
above the words LabQuest 2. Once the LabQuest 2 is powered up you will see the following screen:
The LabQuest 2 can be powered down by holding down the same button with the red icon for about 5
seconds. It will indicate when it is powering down – let go of the button at that point.
Figure A-4.3 Opening Screen of LabQuest 2
Appendix A-4
-135-
II.
Plugging in and setting up sensors
Plugging in the various sensors is not particularly difficult. They will only fit in one place. If you feel
like you are having to force the connection please check with your instructor.
The sensors we use will be auto-detected so there is not much to do in setting them up. Plug them in and
the LabQuest 2 will recognize them.
III.
Modifying Data Collection Settings
On occasion you may wish to adjust the data collection parameters – how often it takes data and for how
long. Just choose the Sensor menu option and pick Data Collection as shown in Figure A-4.2.
Figure A-4.2 Data Collection option (left) and available parameters after its selection (right). A
keyboard will become visible when you hit any of the numerical entries in the parameter screen (right).
IV.
Starting data collection
Data collection is a matter of hitting the green arrow in the lower left-corner of the screen (See the left
part of Figure A-4.2 above). A graph will appear and you will see your data being collected. The data
collection will stop when the time set in Data Collection above runs out or you hit the red square button
in the lower left that appears while the data are being collected.
Figure A-4.3 The screen during data collection.
-136-
Appendix A-4
V.
Saving data
It is best to save your data before working on it too much. Data in any experiment is not recoverable –
data manipulation is. There are two approaches to saving your data – one will save it in format
compatible with LoggerPro software and the other will save it in a text file that you can easily feed to
Excel. LoggerPro is software made available for data handling on a computer.
Saving in a format compatible with LoggerPro software
Choose the Save option from the File menu from the screen (Figure A-4.4 on the left).
Figure A-4.4 Save option from File Menu (left) and ensuing screen (right)
On the ensuing screen click in the data box at the top right (filled in with untitled in Figure A-4.4 (right).
A keyboard will appear – give your file a name and click Save. If you want to save it on a USB device
or memory card, put the appropriate device in the LabQuest 2 and select it from the icons of available
devices in the upper left before saving.
Saving in a text format suitable for Excel or other software
This process is very similar to the one immediately above. The only difference is that, instead of
selecting the Save option in the File menu, you will select the Export. The ensuing screen is just like the
one on the right. Give it a file name.
A word of caution. If you save a file in the text format and move it to Excel, Excel will not pick it up as
an available file unless you select having it show All Files or *.txt files.
Appendix A-4
-137-
VI.
Data Analysis
The ability to efficiently and effectively analyze data is an important part of what this instrumentation
offers. This section will touch on the basics of data analysis – the focus will be on the types of analysis
we might do in this lab.
As an example of a data set, the data in Figure A-4.5 were gathered by measuring the temperature of
water in a can as it was heated on a gas stove. You will learn more about this in Experiment 12.
Figure A-4.5 A sample data run
-138-
Appendix A-4
Suppose the goal here is to know maximum temperature AFTER the little peak that shows up in the
middle. One can do statistics on the data by just choosing Analyze from the menu, then Statistics, and
then Temperature as shown on the left in Figure A-4.6.
Figure A-4.6 Using the Statistics function to find the maximum temperature
The results are given in the right-hand window of Figure A-4.6. Notice the maximum given (53.7 °C)
was that at the top of the little peak about half-way through the plot – not the one we want. We want the
maximum that occurs after that. That is easy to get.
To determine statistics for a particular region (or any other sort of analysis for that matter) simply touch
the stylus to the window and drag it over the region you are interested in. Once you have it set, you can
fine tune it by using the left and right arrows next to the Time axis label to move in small increments.
Select Analyze/Statistics/Temperature again and you will get the results in the right-hand window.
Once you have chosen the region of interest you will see a screen similar to that in Figure A-4.7. Notice
the maximum in this region is smaller than in the overall set of data (51.6 °C compared to 53.7 °C).
Appendix A-4
-139-
Figure A-4.7 Maximum from Statistics applied to a small region
VII.
Printing Data
Your laboratory will have a wireless printer which can be used to print graphs and other information
directly from the LabQuest 2. When working in pairs it is best that you only print one copy of the graph
and one partner includes it in the report.
To print, go to the Print option under the File menu (the left-hand side of Figure A.4-8). The printer
should already be selected and set up.
Please include a Graph Title that includes at least your name and/or your partner’s name to avoid
confusion at the printer. A keyboard will once again show up when you click on the “Enter graph title
here” entry.
Figure A-4.8 Print Menu selection and ensuing print box
-140-
Appendix A-4