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1CHEM 1361General Chemistry I Laboratory Manual

August 201310th Edition, 1st Printing

Cameron UniversityDepartment of Physical Sciences

Lawton, OK 73505Dr. 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.

2

Table of ContentsIntroductory 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

3

Appendix A-4 Basic Use of the Vernier LabQuest 2........................133

4

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. Buckleye-mail: [email protected]

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mailto:[email protected]

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 Corrections, suggestions, etc.Exp. 1 – Laboratory Safety

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

Exp. 9 – Molar Stoichiometry in a Chemical Reaction

Feedback Page 7

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

8 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.

Volume Measurement:

Other Drawer Items:

Equipment in Common Areas:

Laboratory Equipment 9

CHEMISTRY 1361LABORATORY 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 is not permitted. No results, data sheets, etc., are accepted from those absent.

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.)

3. Each student is responsible for adhering to all safety rules listed below and all safety instructions given at the start of each lab.

4. Generally, each student will have a laboratory partner. Each pair of experimenters is required to maintain an assigned lab locker/space, return clean, used glassware to the lab locker assigned, return hardware to the correct storage drawer, and clean their work area. Make sure that all equipment is unplugged before you leave the laboratory.

5. Before a laboratory grade is assigned, you must have a passing grade assigned in the chemistry lecture course (departmental policy).

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.

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.

14. If the instructor is talking, students should be listening.

15. Arrive prepared for each laboratory. Read the assigned laboratory material before laboratory.

16. Each week safety and experimental technique will be detailed.

17. No food or drinks in the laboratory.

18. Dispose of materials in the designated containers.

19. No solids should be thrown in the sinks.

20. Do not sit on the laboratory tables.

21. Never pour excess chemicals back in the storage bottle.

22. Your work area must be clean and all equipment returned to its proper place before you are dismissed from the laboratory.

23. All spills must be cleaned up immediately.

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.

10 Laboratory Policies

1CHEMISTRY 1361LABORATORY 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 session. Be on time. Attendance records are kept. Tardiness is not permitted. No results data sheets, etc., are accepted from those absent.

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.)

3. Each student is responsible for adhering to all safety rules listed below and all safety instructions given at the start of each lab.

4. Generally, each student will have a laboratory partner. Each pair of experimenters is required to maintain an assigned lab locker/space, return clean, used glassware to the lab locker assigned, return hardware to the correct storage drawer, and clean their work area. Make sure that all equipment is unplugged before you leave the laboratory.

5. Before a laboratory grade is assigned, you must have a passing grade assigned in the chemistry lecture course (departmental policy).

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.

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.

14. If the instructor is talking, students should be listening.

15. Arrive prepared for each laboratory. Read the assigned laboratory material before laboratory.

16. Each week safety and experimental technique will be detailed.

17. No food or drinks in the laboratory.

18. Dispose of materials in the designated containers.

19. No solids should be thrown in the sinks.

20. Do not sit on the laboratory tables.

21. Never pour excess chemicals back in the storage bottle.

22. Your work area must be clean and all equipment returned to its proper place before you are dismissed from the laboratory.

23. All spills must be cleaned up immediately.

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.

Student Name: __________________________

Instructor: ______________________________

Lab Day (Circle one): M T W R F

Lab Time: __________________________

Laboratory Policies 11

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:

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

I. 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 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 -13-

Figure 1.1 Acceptable Safety Goggles

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.

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.

-14- Experiment #1: Laboratory Safety and Practice

http://www.ansellpro.com/download/Ansell_7thEditionChemicalResistanceGuide.pdf

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 ConsiderationsYou 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.


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 ComponentsSection # Title Short Description

1 Identification 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 on ingredients

includes information on chemical ingredients; trade secret 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 measures

lists emergency procedures; protective equipment; proper methods of containment and cleanup

7 Handling and storage lists precautions for safe handling and storage, including incompatibilities

8 Exposure controls/personal protection

lists OSHA’s Permissible Exposure Limits (PELs); Threshold Limit Values (TLVs); appropriate engineering controls; PPE

9 Physical and chemical properties

lists the chemicals characteristics

10 Stability and reactivity lists chemical stability and possibility of hazardous reactions11 Toxicological

informationincludes routes of exposure; related symptoms, acute and 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

sheetInformation 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

* Information in Sections 12-15 comes from agencies other than OSHA.


http://www.unece.org/trans/danger/publi/ghs/ghs_rev00/00files_e.html

https://www.osha.gov/dsg/hazcom/ghs.html%234.8

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 sheetshttp://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:

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.


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.

http://ehs.okstate.edu/links/msds.htm

http://www.ilpi.com/msds/

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.


Blue

Red

Yellow

White

Figure 1.2 NFPA Diamond (from http://www.zeably.com)

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

1Laboratory 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 _____1

Sign and submit to your instructor after the safety orientation.


3.4.

5. 7.

8.

10.11.

12.

14.13.

15.

17. 19.20.

21.27.

30.

22.

1. 2.

6.

16.

18.

23.24.

25.

26.

28.

29.

9.

© 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 -23-

Your Name: _________________________________________

CircleLab Day: M T W R F

Partner’s Name: _________________________________________Lab Time: __________

1Safety ObservationsIn 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. _________________________________

15._________________________________ 16._________________________________

17._________________________________ 18._________________________________

19._________________________________ 20._________________________________

21. _________________________________ 22._________________________________

23._________________________________ 24._________________________________

25. _________________________________ 26. _________________________________

Safety Cartoon Error List -24-

27. _________________________________ 28. _________________________________

29. _________________________________ 30. _________________________________

Cartoon Error List -25-

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

Section 1.5 – Uncertainty in Measurement

Select the most appropriate piece of lab glassware to make volume measurements to desired accuracyUse mass and density data to calculate volume Section 1.4, 1.6Arithmetically manipulate measured quantities reporting calculated values to the proper number of significant figures

Section 1.5 – Uncertainty in Measurement

Use the basic statistical concepts of mean and standard deviation to evaluate precision of measurements

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 thosequantities 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.

-26- Experiment #2: Measurements, Accuracy, and Precision

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 -27-

Figure 2.2 – Assorted Glassware

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.

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.


Figure 2.2 – Analytical Balance

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 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 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 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.


Figure 2.3 - Menisci

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.


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.


Laboratory #2: Measurements, Accuracy, and PrecisionLaboratory 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 10-dm markings 100-cm markings 1000-mm markings

m Square of Error

dm Square of Error

cm Square of Error

mm Square of Error

Person 1Person 2Person 3Person 4AverageSum of “Square of Error”Sum of “Square of Error”/(n-1)Standard Deviation


Your Name: _________________________________________



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)

Markings on glassware

25-mL volumetric

pipette25-mL Mohr

pipette

50-mL burette


Laboratory #2: Measurements, Accuracy, and PrecisionPost 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 precise but not accurateboth accurate and precise neither accurate nor precise


Your Name: _________________________________________



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:

a. Which of your measuring devices showed the smallest percent error (remember the expected value in each case was 25 mL)? Show your work.

Experiment #2: Measurements, Accuracy, and Precision -35-

b. Which of your measuring devices showed the greatest percent error? Show your work.


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 MeasurementUnderstand the relationship between sugar concentration and densityConstruct simple hydrometerLearn principles behind calibrationUse Excel to evaluate calibration dataUse calibrated device to measure unknown INTRODUCTIONFat, 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 -37-

PROCEDUREA. 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 hydrometerEach 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. 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.

2. 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.

3. Repeat steps 1 and 2 for the remaining standards and pure water.

C. Measurement of Density of DrinksSafety 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 TreatmentYou 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. Development of Calibration Relationship The calibration solutions include water (which is 0% sugar), and 4%, 8%, 12%, and 16% sugar

-38- 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: 32:133Equation Chapter 3 Section 3

232\* MERGEFORMAT (.)

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 33 to find the % sugar in any solution after determining the height to which the hydrometer floats. Rearrangement of Equation 32 gives:

333\* MERGEFORMAT (.)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.

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.


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.


Figure 3.1 – Microsoft Excel 2010 Spreadsheet with Graph

The last step is to use the slope and intercept determined to find the % sugar in the drinks using Equation 33. 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 33, 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 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:


Compare this to Equation 33. 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 JUICESLaboratory Record

Data from Calibration Samples:

sample sugar concentrationx axis

hydrometer reading (mm)y axis

Standard 1 0

Standard 2 4 %

Standard 3 8 %

Standard 4 12 %

Standard 5 16 %

↑↑↑↑↑↑↑↑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


Your Name: _________________________________________



!!!BE SURE TO TURN IN YOUR CALIBRATION GRAPH!!!

EXPERIMENT #3: SUGAR IN SOFT DRINKS AND FRUIT JUICESPost-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?


1EXPERIMENT #4: USING PHYSICAL PROPERTIES TODETERMINE THE IDENTITY OF AN UNKNOWN

Learning Objectives Textbook Reference (Chemistry: Atoms First, Julia Burdge & Jason Overby, 2012)

Distinguish between physical and chemical properties Section 1.3 – Properties of MatterExperimentally determine the boiling point of a substance

Experimentally determine the density of a substance Section 1.4 – Scientific MeasurementExperimentally determine the solubility of a substanceUse physical property data to identify an unknown from a list of possibilities

IntroductionSubstances 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. 1Examples 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.

DensityDensity 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 PointEvery 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.

SolubilityThis 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 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 IndexThe 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,

where ηD represents the index of refraction of the substance of interest.

ProcedureA. 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 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.

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

Figure 3.1 – Boiling Point Setup

Eyepiece

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 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.

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.

1. Place a few drops of liquid on the stage and close the cover.

2. Rotate the light up over the stage.

3. Look through the ocular and move the switch on the left from the up position down to its middle position.

4. Look through the eyepiece and you will see a light area and a darkarea separated by a horizontal line (Figure 3.3a). Use the knob on the rightto align the

horizontal line with the crosshairs in the window.


Switch

Figure 3.2 Abbe Refractometer

Light

Stage

Figure 3.3a Switch in middle position Figure 3.3b Switch in down position

Knob

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.

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

Your Name: _________________________________________



EXPERIMENT #3: USING PHYSICAL PROPERTIES TODETERMINE THE IDENTITY OF AN UNKNOWN

Laboratory RecordUNKNOWN 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 DeterminationOne 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 IdentificationConclusion: The density, boiling point, solubility, and refractive index agree best with the compound _____________ which is my unknown # _________ .


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.

1EXPERIMENT #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 mixtureUnderstand the process of heating a material to a constant mass

INTRODUCTIONA 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 56 .55Equation Section 5


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 57 illustrates a method for determining the mass of MgSO4 ∙7H2O originally present based on the mass of H2O lost


-54- Experiment #5: Separation of a Mixture

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.

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.

Schematically, one can consider the experimental procedure to be composed of two steps as illustrated in Figure 5.1.

Figure 5.1 Schematic of Separation Process

Experiment #5: Separation of a Mixture -55-

Orig

inal

Mix

ture MgSO4·7H2O

SiO2

NaCl

Heat MgSO4

SiO2

NaCl Add

Wat

er SiO2

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.)

4. Allow the dish to cool so that you can safely carry it to the 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.


Figure 5.2 – Set-up for Heating Initial Mixture

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 (g)Mass of empty evaporating dish 33.8792Mass of dish and sample after heating

After 1st heating 34.5673After 2nd heating 34.5297After 3rd heating 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 gMass of sample based on 1st heating: 34.5673 – 33.8792 = 0.6881 g

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 gMass of sample based on 2nd heating: 34.5297 – 33.8792 = 0.6505 g

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.


Your Name: _________________________________________



EXPERIMENT #5: SEPARATION OF A MIXTURE

Laboratory Record

Unknown Identifier: ______

Your Trial Neighbor 1’s Trial Neighbor 2’s Trial

Mass of evaporating dish and sample (g) ___________ ____________ ____________

Mass of empty evaporating dish (g) ___________ ____________ ____________

Mass of sample (g) ___________ ____________ ____________

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 57 (g) ___________ ____________ ____________

Percent MgSO4·7H2O in sample ___________ ____________ ____________

Average % MgSO4∙7H2O in sample ____________

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: % MgSO4·7H2O % SiO2 % NaCl

Average of three trials: ___________ ____________ ____________

Standard deviation of three trials: ___________ ____________ ____________


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 have on the % of sand calculated in the mixture compared to the actual % of sand?

Increased Unchanged Decreased

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?


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?


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?


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?


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?


2. 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 -59-

EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULA


Determine empirical formula from % composition information

Section 5.10 –Molar Mass

IntroductionA 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 Weight Divide by smallest Empirical

FormulaC 10.4 12.0107 10.4/12.0107 = 0.867 0.867/0.867=1.0 Determined to be

C1S1Cl2 or simply CSCl2

S 27.8 32.065 27.8/32.065 =0.869 0.869/0.867=1.0Cl 61.7 35.453 61.7/35.453=1.74 0.867/1.74=2.0The 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.86Equation Section 6In 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 69 and 610 occur upon heating.


10610\* MERGEFORMAT (.)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 611:


The Mg3N2 is converted to Mg(OH)2 which is then converted to MgO through further heating.

-60- Experiment #6: Determination of an Empirical Chemical Formula

Procedure1. 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.

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).

Experiment #6: Determination of an Empirical Chemical Formula -61-

Figure 6.1 – Experimental Setup

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.

-62- Experiment #6: Determination of an Empirical Chemical Formula

Your Name: _________________________________________



EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULALaboratory 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: ____________

-64- Experiment #14: Molecular Modeling

EXPERIMENT #6: DETERMINATION OF AN EMPIRICAL CHEMICAL FORMULAPost-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?


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?


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?


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?


2. What is the percent difference in your sample mass between the second heating and the first heating? Show your work.

Experiment #7: Molecular Modeling -65-

Experiment #7– Molecular ModelingLearning Objectives Textbook Reference (Chemistry: Atoms First,

Julia Burdge and Jason Overby, 2012)Draw Lewis structures for simple molecules and ions Section 5.5 – Covalent Bonding and Molecules

Section 6.3 – Drawing Lewis StructuresDetermine formal charge from Lewis structures Section 6.4 – Lewis Structures and Formal

ChargesDetermine electron-domain geometry and molecular geometry from Lewis structures

Section 7.1 – Molecular Geometry

Assess polarity from Lewis structures Section 7.2 – Molecular Geometry and PolarityDraw 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. Draw Lewis structures for a variety of molecules and ions.B. Assign formal charges to each atom based on the Lewis structure.C. Assign electron-domain and molecular geometries.D. Determine the polarity of each species.E. 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 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


H-Br O – S - O O – N - O

O

H-BrO = S - O O = N - O

O

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, SO4

2- 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 StructuresStep in Process HBr SO2 NO3

-

Step a H ― Br O―S―O O―N― O│O

Step b

Step c – number of valence electrons

1 (H) + 7 (Br) =8 6 (S) + 2 × 6 (O) = 18

5 (N) + 3 ×6 (O) + 1 (charge) = 24

Step d Step b and number of valence electrons

match

Too many electrons in Step b

Too many electrons in Step b

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 ChargeHBr SO2 NO3

-

Looking at atom: H Br O= =S- -O N O= -OValence Electrons 1 7 6 6 6 5 6 6Electrons assigned due to nonbonded electrons

0 6 4 2 6 0 4 6

Electrons assigned due to bonded 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 to atom:

0+1=1 6+1=7 2 + 4= 6 2+3= 5 6 + 1=7 0+4= 4 4 + 2= 6 6 + 1=7

Formal charge = valence – assigned electrons

1-1 = 0 7 – 7 = 0 6-6 = 0 6-5 = 1 6–7= -1 5–4 = +1 6 – 6= 0 6–7= -1

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 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


O

S

O O

S Electron-domain geometry: three electron-domains around S so trigonal planar from first structure

Molecular geometry: Bent (or angular) from second structure where nonbonded electrons on sulfur are hidden from view.

O

O = S - O

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 Electron-domain Geometry Molecular GeometryAB2 Linear LinearAB3 Trigonal Planar Trigonal Planar

AB2E Trigonal Planar Bent (or angular)AB4 Tetrahedral Tetrahedral

AB3E Tetrahedral Trigonal PyramidalAB2E2 Tetrahedral Bent (or angular)

As an example, consider the S atom in the SO2 species from earlier. 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:

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:

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 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


Cl Cl

B

F

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.

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.

E. Submit Information to and Retrieve Information From the Scigress Molecular Modeling ProgramA 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:

Basics of Scigress WorkspaceEither 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 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.


Figure 7.1 Basic procedures for working with molecules in Scigress

Draw the Lewis structure for your molecule or ion, being sure to assign formal charge

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.

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.

Open ProjectLeader and identify the properties of interest to be calculated.

Enter the names of the files of your molecules of interest into ProjectLeader and tell it to calculate the properties.

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 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.

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 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


Figure 7.2 Scigress Workspace

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” ProcedureThere 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

whole molecule, you may click on the select tool first and then click anywhere in the workspace (not on the molecule).

f. 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.

g. 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

h. 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.

i. 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.

j. 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 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. Left-click 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.

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.


Figure 7.3 Opening Screen of ProjectLeader

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. Double-click 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 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.


Figure 7.4 Selecting Properties in ProjectLeader

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Your Name: _________________________________________



Experiment #7 – Molecular ModelingLaboratory 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 Molecule/Ion Lewis StructureDiatomic Molecules

H2 F2

Cl2 HF

HCl HBr

HI CO

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.


Molecule/Ion Lewis Structure Molecule/Ion Lewis StructureBE SURE TO LABEL YOUR BOND ANGLES ON THESE MOLECULES PER INSTRUCTIONS ON PAGE 114

Triatomic Species

CO2 NO2-

SO2 O3

H2O 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 SO2 .

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?

BE SURE TO LABEL YOUR BOND ANGLES ON THESE MOLECULES PER INSTRUCTIONS ON PAGE 114


Molecule/Ion Lewis Structure Molecule/Ion Lewis Structure

Polyatomic species:

CO32- SO3

2-

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 #6: Determination of an Empirical Chemical Formula -79-

EXPERIMENT #8: PREPARATION OF AN ALUM


Learn the technique of gravity filtrationLearn the technique of vacuum filtrationConvert between grams and moles Section 5.10 – Molar MassConvert between molarity and volume and moles Section 9.5 – Concentrations of SolutionsIdentify the limiting reactant in a chemical reaction Section 8.4 – Limiting ReactantsCalculate theoretical yield in a chemical reaction 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: 128Equation Section 8


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 the U. S. In order to make writing paper, the open spaces in the cellulose fibers must be filled with -80- Experiment #8: Preparation of an Alum

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.


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.


As more acid is added to the solution, the precipitate of Al(OH)3 dissolves according to the following reaction:


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.


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:

Experiment #8: Preparation of an Alum -81-


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

-82- Experiment #8: Preparation of an Alum

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 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.


Your Name: _________________________________________



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

Water Plenty – in excess

Product Mass Produced Moles Produced

Alum

Hydrogen gas (calculated)

1. Use the data above and Equation 818 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.


Experiment #9: Metathesis Reactions

Learning Objectives Textbook Reference (Chemistry: Atoms First, Julia Burdge and JasonOverby, 2012)

Observe evidence of metathesis reactions Section 9.2 – Precipitation ReactionsIdentify spectator ions in solutions Section 9.2 – Precipitation Reactions

Section 3.3 – Acids, Bases and Neutralization Reactions

Write net ionic equations for metathesis reactions Section 9.2 – Precipitation ReactionsSection 3.3 – Acids, Bases and Neutralization Reactions

IntroductionA 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:199Equation Section 9


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 Procedure1. 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.

Experiment #9: Metathesis Reactions -87-

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.

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.

-88- Experiment #9: Metathesis Reactions

Your Name: _________________________________________



Experiment #9: Metathesis ReactionsLaboratory 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)20.3 M

Mg(NO3)20.1 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 ReactionsMetathesis 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:


Net Ionic Equation:

Na2CO3 CaCl2

Molecular Equation:


Net Ionic Equation:



Na2CrO4 Cu(NO3)2

Molecular Equation:


Net Ionic Equation:

Na2CrO4 Pb(NO3)2

Molecular Equation:


Net Ionic Equation:

KOH Pb(NO3)2

Molecular Equation:


Net Ionic Equation

Na2SO4 BaCl2

Molecular Equation:




Net Ionic Equation:

NaI Sr(NO3)2

Molecular Equation:


Net Ionic Equation:

HCl Na2CO3

Molecular Equation:


Net Ionic Equation:


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 Section 9.4 – Oxidation-Reduction ReactionsRelate stoichiometric coefficients to numbers of moles

Section 8.3 – Calculations with Balanced Chemical Equations

Determine number of moles from masses Section 5.10 – Molar Mass

INTRODUCTIONThe 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.21104Equation Chapter (Next) Section 10


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.


We can use this equation to determine the number of moles of silver precipitated and copper that dissolved. According to Equation 1022, 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.


-94- Experiment #10: Molar Stoichiometry in a Chemical Reaction

Start with silver metalAg

Silver cations Ag+ Copper atoms (Cu) with silver cations (Ag+)

Copper ions (Cu+) and reclaimed silver atoms

Steps 1, 2, 3 Step 4

Steps 5 and 6

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.

(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.

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 plastic eyedropper to add 1 mL portions of sodium hydroxide solution to the solution containing

Experiment #10: Molar Stoichiometry in a Chemical Reaction -95-

Double Boiler - Silver in inside beakerWater in outside beaker

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.


Your Name: _________________________________________



EXPERIMENT #10: MOLAR STOICHIOMETRY IN A CHEMICAL REACTIONLaboratory 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.

Mass (g) Moles (mol)

Initial quantity of Ag _____________________

_____________________

Ag Inform

ationMass of evaporating dish and dry silver:

After 1st Heating: ______________________

After 2nd Heating: ______________________

After 3rd Heating: ______________________

Mass of empty evaporating dish ______________________

Recovered quantity of Ag _____________________ _____________________


Initial quantity of Cu wire _____________________

_____________________

Cu Inform

ation

Recovered quantity of Cu wire _____________________ _____________________

Quantity of Cu wire consumed _____________________ _____________________

Moles Ag recovered / moles Cu consumed (from above): __________________

% Ag recovered _____________________

Experiment #10: Molar Stoichiometry in a Chemical Reaction -99-

EXPERIMENT #10: MOLAR STOICHIOMETRY IN A CHEMICAL REACTIONPost-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+ + Cu →

d. Mn + Ca2+ →


EXPERIMENT #11: Determination of Acetic Acid in Vinegar

Learning Objectives Textbook Reference (Chemistry: Atoms First, Sylvia Burdge and Jason Overby, 2012)

Work with molarity and volume to determine moles Section 9.5 – Concentrations of SolutionsLearn the terms titration, standard solution, equivalence point, and end point

Section 9.6 – Aqueous Reactions and Chemical Analysis

Use and read a burette properlyStandardize a sodium hydroxide solution Section 9.6 – Aqueous Reactions and

Chemical AnalysisDetermine the % acetic acid in vinegar Section 9.6 – Aqueous Reactions and

Chemical Analysis

IntroductionOrdinary 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.

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 11262511Equation Section 11


Experiment #11: Determination of Acetic Acid in Vinegar -101-

Figure 11.1 – Titration Apparatus (Photo Courtesy of Mr. Herman Curtis)

Or, in net ionic form:

271127\* MERGEFORMAT (.) 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 1126, 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 1128.


Or, in net ionic form:


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 1128, 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

-102- Experiment #11: Determination of Acetic Acid in Vinegar

motion shake the burette down and up once and it should clear the air bubble. If there is still an air 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 Vinegar1. 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.


Cleaning the buretteIt 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.


Your Name: _________________________________________



. EXPERIMENT #11: Determination of Acetic Acid in Vinegar

Laboratory Record

Standardization

Trial Mass KHP (g)

Moles KHP (moles)

Finalburette 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 VinegarLaboratory Record

Molarity of your NaOH solution from the previous page: ______________

Vinegar Unknown Letter or Number: ______________________

TrialVolume Vinegar

(mL)

Finalburette 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?


Your Name: _________________________________________



EXPERIMENT # 11: Determination of Acetic Acid in VinegarPost-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.


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.


1EXPERIMENT #12: COMPARISON OF THE ENERGY CONTENT OF FUELS BY COMBUSTION


Determine enthalpy of combustion through calorimetry Section 10.3 EnthalpySection 10.4 Calorimetry

Relate enthalpy of combustion to energy content of fuelsConsider differences in energy content of fuels based on composition

IntroductionLiquid 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 3015Equation Chapter (Next) Section 1draw 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 (CH4), it burns to form carbon dioxide and water:3112Equation Section 12


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 1calculated as follows:


341234\* MERGEFORMAT (.) 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

Experiment #12: Comparison of the Energy Content of Fuels by Combustion -109-

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.

-110- Experiment #12: Comparison of the Energy Content of Fuels by Combustion

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 1stir 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.

Figure 12.2 Example of Temperature-Time Data13.

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 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.


Your Name: _________________________________________



EXPERIMENT #12: Enthalpy of Combustion of Various FuelsLaboratory Record and Post-Lab

Measurement (Calculated) Fuel( )

Fuel( )

Fuel( )

1 Specific heat of liquid water (Cs,water): 4.184 Jg-1K-1

Specific heat of Al (Cs,Al): 0.9000 Jg-1K-1

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 -115

1Laboratory 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 #13: DETERMINATION OF THE ENTHALPY OF FUSION OF ICE


Determine the calorimeter constant in a simple system Section 10.4 - CalorimetryUse the conservation of energy principle to determine missing information

Section 10.4 - Calorimetry

Apply the concepts of specific heat capacity and heat capacity to a simple calorimeter

Section 10.4 - Calorimetry

IntroductionThe 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: 3513Equation Section 13


or, in symbols:


where the algebraic sign of q is positive if heat is absorbed and negative if heat is evolved.

The term qice total in Equation 1337 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 1337 to reflect these parts gives:


If three of the four terms in Equation 1338 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 1339 provides a simple approach to finding the heat flow given the specific heat capacity, mass, and temperature change.


where: 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 substanceTf = final temperature

Experiment #13: Determination of the Enthalpy of Fusion of Ice -117-

Ti = initial temperature

Finding qcalorimeter in Equation 1338 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 1336 and 1337


or, in symbols:


The calorimeter constant, also called its heat capacity, may be found from Equation 1341 and the relationship between heat and the calorimeter constant is given as Equation1339:


1Procedure 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.

-118- Experiment #13: Determination of the Enthalpy of Fusion of Ice

B. Calibration1. 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.


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.


Figure 13.1 Appearance of LabQuest2 Data for the Calibration

Your Name: _________________________________________



1EXPERIMENT #13: DETERMINATION OF THE

ENTHALPY OF FUSION OF ICELaboratory 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)(include algebraic sign)

Tf – Ti RT water (°C)(include algebraic sign)

qhot (J) (=4.184 J/g∙°C × mhot × (Tf – Ti))

qRT (J) (=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):


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): 0 °C

Tf – Ti for water (°C): Tf – Ti (°C):

qwater (J)(4.184 J/g∙°C × mwater × (Tf –Ti)water)

qice (J)(4.184 J/g∙°C × mice × (Tf –Ti)ice)

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 1338 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 )Show your work:


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?


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.)


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. )


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?


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 #14: DETERMINING THE MOLAR MASS OF A VOLATILE LIQUID BY THE DUMAS METHOD


Work with the ideal gas law Section 11.5 – The Ideal Gas EquationDetermine the molecular weight from the gas law and sufficient measurements

Section 11.5 - The Ideal-Gas Equation

INTRODUCTIONOne 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:4314Equation Section 14

441444\* MERGEFORMAT (.) 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.

451445\* MERGEFORMAT (.)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.

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 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.

5. Place 5 mL of your unknown in the Erlenmeyer flask and cover with the aluminum cap.

Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method -127-


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.


Your Name: _________________________________________



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

Barometric pressure (include units) ________________ ________________

Volume of Erlenmeyer flask ________________ mL ________________ mL

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 ________________ ________________

Conversion of volume to units consistent with your choice of gas constant ________________ ________________

Molar mass of unknown ________________ ________________

Average molar mass from two trials ________________


Conclusion : The molar mass of unknown # _______ is ________ g/mol.

-130- Experiment #14: Determining the Molar Mass of a Volatile Liquid by the Dumas Method

EXPERIMENT # 14: DETERMINING THE MOLAR MASSOF A VOLATILE LIQUID BY THE DUMAS METHOD

Post - LabRead 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?


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?


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?


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?


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?



EXPERIMENT #15: Regularities in the Properties of the Elements


Become familiar with trends in properties in the periodic table

Section 4.4 Periodic Trends in Properties of Elements

Learn the use of Periodic Table Live! as a means of exploring elements and their properties

J. Chem. Educ. 2008, 85, 22.http://www.chemeddl.org/collections/ptl/

Use logic to define searches of large quantities of data

IntroductionEach 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.

ProcedureThe 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.

-132- Experiment #13: Regularities in the Properties of Elements

http://www.chemeddl.org/resources/ptl/

http://www.chemeddl.org/collections/ptl/

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.

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/SortThis option allows one to do searches on elements and their properties and to make plots of one property vs. another. Further information is below.

BiographiesThis button provides biographical information on a large number of people involved in the periodic chart and its contents.

GlossaryThis is what is sounds like – a glossary of terms related to properties.

PreferencesThis allows one to set a preference as to the numbering system for groups desired.

Experiment #15: Regularities in the Properties of Elements -133-

Figure 15.1 Opening Window for Periodic Table Live!

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.

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.


Figure 15.2 Chart/Sort Window

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.


Figure 15.3 Atomic Radius vs Atomic Number

Your Name: _________________________________________



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.)

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?


-138- Determining the Molar Mass of a Volatile Liquid by the Dumas Method

Appendices

Basic Error AnalysisSeveral 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).

BackgroundIn 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 1446


where s is the standard deviation, xi represents the ith experimentally determined value; 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:

The standard deviation is given as:

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.

Appendix A-1 -139-

1Activity Series of Metals in Aqueous SolutionMetal Oxidation Reaction Potential

(most easily oxidized) Eo in voltsLithium Li ➔ Li+ + 1e- 3.04Potassium K ➔ K+ + 1e- 2.931Barium Ba ➔ Ba2+ + 2e- 2.912Calcium Ca ➔ Ca2+ + 2e- 2.868Sodium Na ➔ Na+ + 1e- 2.71Magnesium Mg ➔ Mg2+ + 2e- 2.377Aluminum Al ➔ A13+ + 3e- 1.662Manganese Mn ➔ Mn2+ + 2e- 1.185Zinc Zn ➔ Zn2+ + 2e- 0.762Chromium Cr ➔ Cr3+ + 3e- 0.744Iron Fe ➔ Fe2+ + 2e- 0.489Cobalt Co ➔ Co2+ + 2e- 0.280Nickel Ni ➔ Ni2+ + 2e- 0.257Sn Sn ➔ Sn2+ + 2e- 0.140Lead Pb ➔ Pb2+ + 2e- 0.126Hydrogen H2 ➔ 2H+ + 2e - 0.00Copper Cu ➔ Cu2+ + 2e- -1.17Silver . Ag ➔ Ag+ + 1e- -0.80Mercury Hg ➔ Hg2+ + 2e- -0.851Platinum Pt ➔ Pt2+ + 2e- -1.118Gold Au ➔ Au3+ + 3e- -1.498

(Most easily reduced)

-140- Appendix A-2

1BASE UNITS IN THE SI SYSTEM, SCIENTIFIC PREFIXES FOR MAGNITUDES, AND CONVERSION FACTORS RELATING LENGTH, VOLUME, AND MASS UNITS

TABLE I: UNITS OF MEASUREMENT

QUANTITY SI UNIT SYMBOL COMMONLY USED SYMBOLMass kilogram kg gram gLength meter m centimeter cmVolume cubic meter m3 liter LTime second s Temperature Kelvin K degrees Celsius CAmount mole mol

1TABLE II: FREQUENTLY USED SCIENTIFIC PREFIXES FOR MAGNITUDES

EXPONENTIALPREFIX SYMBOL NUMBER NOTATION

giga G 1,000,000,000 109

mega M 1,000,000 106

kilo k 1,000 103

- - 1 100

deci d 0.1 10-1

centi c 0.01 10-2

milli m 0.001 10-3

micro μ 0.000001 10-6

nano n 0.000000001 10-9

1TABLE III: CONVERSION FACTORS FOR LENGTH, VOLUME AND MASS UNITS

Length METRIC ENGLISH METRIC-ENGLISH1 km = 103 m 1 ft = 12 in 2.54 cm = 1 in 1 cm = 10-2 m 1 yd = 3 ft 39.37 in = 1 m 1 mm = 10-3 m 1 mile = 5280 ft 1.609 km = 1 mile 1 nm = 10-9 m 1 A = 10-10 m

Volume 1 mL = 1 cm3 1 gal = 4 qt 1 L = 1.057 qt 1 L = 1000 mL 1 qt = 2 pt 28.32 L = 1 ft3 1 m3 = 1000 L 1 pt = 16 fl oz

Mass 1 kg = 1000 g 1 m3 = 35.31 ft3 1 mg = .001 g 1 lb = 16 oz 453.6 g = 1 lb

1μ g = 10-6 g 1 ton = 2000 lb 1 g = .03527 oz

Appendix A-2 -141-

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

-142- Appendix A-2

Some Common IonsPositive Negative

1+

Li+ Lithium Cu+ Copper (I) orCuprous

1-

F- Fluoride ClO- Hypochlorite

Na+ Sodium Hg22+ Mercury (I) or

MercurousCl- Chloride ClO2

- Chlorite

K+ Potassium Ag+ Silver Br- Bromide ClO3- Chlorate

Rb+ Rubidium NH4+ Ammonium I- Iodide ClO4

- PerchlorateCs+ Cesium NO3

- Nitrate IO4- Periodate

H+ Hydrogen NO2- Nitrite OH- Hydroxide

C2H3O2- Acetate SCN- Thiocyanate

MnO4- Permanganate HCO3

- Hydrogen Carbonate orBicarbonate

CN- Cyanide O22- Peroxide

H- Hydride

2+

Mg2+ Magnesium Cu2+ Copper (II) orCupric

2-

O2- Oxide S2- Sulfide

Ca2+ Calcium Hg2+ Mercury (II) orMercuric

SO42- Sulfate SO3

2- Sulfite

Sr2+ Strontium Sn2+ Tin (II) or Stannous

CO32- Carbonate C2O4

2- Oxalate

Ba2+ Barium Zn2+ Zinc CrO42- Chromate Cr2O7

2- DichromateFe2+ Iron (II) or

FerrousPb2+ Lead (II) or

PlumbousCd2+ Cadmium Mn2+ Manganese (II)

or Manganous

S2O32- Thiosulfate

Cr2+ Chromium (II) orChromous

Co2+ Cobalt (II) orCobaltous

3+ Al3+ Aluminum As3+ Arsenic 3- N3- Nitride P3- PhosphideCr3+ Chromium (III)

orChromic

Co3+ Cobalt (III) or Cobaltic

PO43- Phosphate PO3

3- Phosphite

Fe3+ Iron (III) orFerric

AsO43- Arsenate AsO3

3- Arsenite

Appendix A-3 -143-

BO33- Borate

-144- Appendix A-3

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. Powering up and turning off the LabQuest 2II. Plugging in and setting up sensorsIII. Modifying Data Collection settingsIV. Starting data collectionV. Saving dataVI. Analyzing dataVII. Printing data

I. 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.

Appendix A-4 -145-

Figure A-4.3 Opening Screen of LabQuest 2

http://www2.vernier.com/manuals/labquest2_user_manual.pdf

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.

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.

-146- Appendix A-4

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).

Figure A-4.3 The screen during data collection.

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 softwareChoose the Save option from the File menu from the screen (Figure A-4.4 on the left).

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 softwareThis 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 -147-

Figure A-4.4 Save option from File Menu (left) and ensuing screen (right)

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.

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.

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.

-148- Appendix A-4

Figure A-4.5 A sample data run

Figure A-4.6 Using the Statistics function to find the maximum temperature

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).

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.

Appendix A-4 -149-

Figure A-4.7 Maximum from Statistics applied to a small region

Figure A-4.8 Print Menu selection and ensuing print box

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