chemistry 113.1 introduction to chemical techniques fall...

CHEMISTRY 113.1INTRODUCTION TO CHEMICAL TECHNIQUES

Fall 2008

SECTION 1. INSTRUCTOR AND COURSE INFORMATION

Instructor: Ms. Luxi LiOffice: Remsen 017 Office hours: Wednesday, 1:00 - 2:00 pmTelephone: (718) 997-4182E-mail: [email protected]: Tuesday, 1:40 – 4:30 PM; Remsen 209

CourseContent: An introductory laboratory course in basic chemistry techniques.

Goals/Objectives: Discovery of basic chemical principles and an introduction to basic chemical

techniques through experimentation. Introduction to data collection, recording,analysis, evaluation and reporting.

Webpage: http://chem.qc.cuny.edu/~introchemlab/CHEM113/home.html

Required Text: Queens College Chemistry 113.1 Laboratory Manual.

SECTION 2. POLICIES/RULES

Attendance: Attendance in laboratory is mandatory. An unexcused absence results in the loss ofall points associated with that laboratory (i.e., 15 pts: prelaboratory quiz, laboratoryproject grade). If you have a university excused absence (such as illness, etc), youmust show the excuse to the laboratory instructor the week after the absence. If youwill miss a laboratory due to religious observance, you must inform theinstructor the week BEFORE the absence, or the absence will not be excused.If the absence is excused, there will be a make-up laboratory (including quiz). Thismake-up laboratory will be administered after the laboratory practical during the 14th

week.

Tardiness: It is your responsibility to attend class and to be punctual. Pre-laboratory quizzeswill be given at the beginning of class and students will be given exactly 30 minutesto complete the quiz. Tardiness will not result in additional time being given for pre-laboratory quizzes.

Laboratories: There are 14 laboratories, including the check-in/introductory laboratory and thefinal examination meeting. The grade will be determined from pre-laboratoryquizzes, laboratory techniques, laboratory reports, and laboratory practical.

Pre-laboratoryquizzes: All pre-laboratory quizzes will be 30 minutes in length and will be given at the

beginning of the laboratory period. Additional time will not be given to studentswho are tardy. These quizzes will consist of three questions, which may containmultiple parts:

Question 1. A post-laboratory question from the previous experiment. Question 2. A safety question about the current experiment.Question 3. A pre-laboratory question from the current experiment.

Laboratoryproject: A complete laboratory project consists of the following documents, which should

be stapled together in the order listed:1. Hard copy of the laboratory report2. The pre-laboratory questions3. The laboratory report sheet4. The post laboratory questions

In addition to this package, an electronic version of the laboratory report should alsobe submitted to Blackboard. The completed project (i.e., Items 1-4) must besubmitted at the beginning of the laboratory period on the date due (see theschedule). An electronic version (i.e., Word or WordPerfect) of the laboratory reportmust be submitted to Blackboard on the date due. Electronic versions oflaboratory reports will be submitted to Turn-it-in software and checked forplagiarism. Turn-it-in checks both internet sources and previously submittedreports. Failure to submit an electronic version of the report will result in a zero onthe laboratory report.

Laboratoryworksheets: Some laboratories do not have projects associated with them. For these laboratories,

the worksheet that accompanies the laboratory must be submitted at the end of thelaboratory in question.

Laboratory reports: A style guide for the laboratory reports is attached to this document and can also be

found on the course website. Unless specified by the laboratory instructor, laboratoryreports are limited in length to 6 pages, excluding Figures and Tables.

Academicdishonesty: While it is natural to discuss experiments among each other, copying and/or

plagiarism will NOT be tolerated on any assignment and will be treated inaccordance with university policy. This policy can be found at:

http://www1.cuny.edu/portal_ur/content/2004/policies/image/policy.pdfElectronic versions of laboratory reports will be submitted to Turn-it-in

software and checked for plagiarism. Turn-it-in checks both internet sourcesand previously submitted reports. We should note here that copying anddownloading graphs from the internet, without express permission of the instructor,constitutes plagiarism (even if referenced). An originality score $ 30% (indicatingthat # 70% of the work is original) from a single reference will result in a gradeof zero on the laboratory report in question.

SECTION 3. GRADING

Course grading: Pre-laboratory quizzes (12 @ 10 pts) 120 pts

Laboratory Projects (12 @ 20 pts) 240 ptsLaboratory practical 60 pts

Total: 420 pts

Grading for Project: 1. Laboratory report* 7 pt

a. Style guide formatting (0.5 pt)b. Introduction and Experiment (1.5 pt)d. Results and Discussion (2 pts)

2. Pre-laboratory questions 4 pt3. Report sheet 3 pt4. Post laboratory questions 4 ptSafety and laboratory technique 2 pt

*Failure to submit an electronic version of the laboratory report will result in the loss of these 7points.

Laboratory Schedule

Lab Pre-Lab (60 min) Lab Work (120 min) Projects due

1 Syllabus and safety.Safety exam.

Check-in. Laboratory techniques.

2 Pre-laboratory quizDensity and graphing.

Finish laboratory techniques.

Density

Laboratory TechniquesProject.

3 Pre-laboratory quizChemical formulas.

Law of definite proportions. Density Project

4 Pre-laboratory quizMoles.

Stoichiometry. Law of definiteproportions project.

5 Pre-laboratory quiz Chemical reactions.

Copper reactions Stoichiometry project.

6 Pre-laboratory quizAcid/base chemistry

Acid/base titrations with anindicator

Copper reaction project.

7 Pre-laboratory quizAcid/base chemistry

Potentiometric analysis. Acid/base project 1.

8 Pre-laboratory quizEnergy in chemistry.

Heat of reactions. Acid/base project 2.

9 Pre-laboratory quizMetal reactivity.

Activity of metals. Heat of reaction project.

10 Pre-laboratory quizElectronic spectroscopy

Emission and Beer’s Law. Activity of metal project.

11 Pre-laboratory quizRedox Chemistry

Redox titration of bleach Beer’s Law Project

12 Chemical kinetics (60 min lecture)

Kinetics of bleaching Redox Titration Project

13 Pre-laboratory quiz Finish Kinetics of bleaching experimentKinetics data analysis using computers (in recitation room)

14 Laboratory check-out and LABORATORY PRACTICAL (60 pts).

Kinetics Project.

Laboratory Drawer Combination:

Chemistry 113.1.

Introduction to Chemical Techniques

Laboratory Manual

Cherice M. Evans, Fred H. Watson andGary L. Findley

Contents

Section 1. General Rules 1

Section 2. Schedule 5

Section 3. Laboratory Safety 7

Section 4. Freshman Chemistry Style Guide 21

Section 5. Instructions for Turn-it-in Assignments 33

Section 6. Useful information 37

Experiment 1. Check-in and Basic Laboratory Techniques 39

Experiment 2. Density 59

Experiment 3. The Law of Definite Proportions 75

Experiment 4. Stoichiometry of a Reaction 93

Experiment 5. Copper reactions 105

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QC Chemistry Laboratory ManualVersion 1.0, 2008

Introduction

SECTION 1General Rules

1.1. Attendance

Attendance in laboratory is mandatory. An unexcused absence results in the loss of allpoints associated with that laboratory (i.e., pre-laboratory quiz, laboratory project grade).If you have a university excused absence (such as illness, etc), you have one week to show theexcuse to the course instructor (not teaching assistant) or the absence will not be excused.If you will miss a laboratory due to religious observance, you must inform thecourse instructor (not teaching assistant) the week BEFORE the absence, orthe absence will not be excused. If you will miss two sequential weeks due toreligious observance, you must inform the instructor (not teaching assistant)that you will be missing two laboratory periods BEFORE the absences, or theabsences will not be excused. If the absence is excused by the course instructor, thecourse instructor will make arrangements with the teaching assistants for the laboratory tobe made-up during a separate laboratory section.

1.2. Laboratories

There are 14 laboratories, including the check-in/introductory laboratory and the check-out/laboratory practical meeting. Grades will be determined from pre-laboratory quizzes,laboratory technique and safety, laboratory reports, laboratory related questions, and thelaboratory practical.

1.3. Tardiness

It is your responsibility to attend class and to be punctual. Pre-laboratory quizzes will begiven at the beginning of class. Tardiness will not result in additional time being allottedfor pre-laboratory quizzes.

1.4. Safety

An open book safety examination, worth 5 points, will be given during the first pre-laboratory period. Students must score a 3/5 on this examination. Failure to do so willresult in the student being barred from the laboratory, until such time as a 3/5 is achieved.(Laboratories that are missed due to failure of the safety examination cannot be made-up.)Moreover, all students will be required to sign the Laboratory Safety Agreement givenat the end of Section 3. Failure to do so will result in removal from the laboratory.

c©2008 QC Chemistry and Biochemistry

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1.5. Pre-laboratory quizzes

All pre-laboratory quizzes will be 30 minutes in length and will be given at thebeginning of the laboratory period. Additional time will not be given to students whoare tardy. These quizzes will have three parts:

(1) Post-laboratory questions and questions from the Report Sheet of the previousexperiment.

(2) Safety questions about the current experiment.(3) Pre-laboratory questions from the current experiment.

1.6. Laboratory projects

A complete laboratory project consists of the following documents, which should be stapledtogether in the order listed:

(1) Hard copy of the laboratory report(2) The pre-laboratory questions(3) The laboratory report sheet(4) The post-laboratory questions

The completed project (i.e., Items 1-4) must be given to the laboratory instructor at thebeginning of the period on the date due (see the schedule in Section 2). In addition tothis package, an electronic version (i.e., Word or WordPerfect) of the laboratory report(excluding pre-laboratory and post-laboratory questions and the laboratory report sheet)must also be uploaded to the Blackboard site for the course on the date due. The electronicreport will check for academic dishonesty using both internet sources and previously sub-mitted reports. Failure to submit an electronic version of the report will resultin a zero on the laboratory report. Instructions for submission of a report throughTurn-it-in are given in Section 5.

1.7. Laboratory reports

A style guide for the laboratory reports is included in Section 4. Unless specified bythe laboratory instructor, laboratory reports are limited in length to 6 pages,excluding Figures and Tables.

1.8. Academic dishonesty

While it is natural to discuss experiments among each other (especially your laboratorypartner), instances of academic dishonesty will NOT be tolerated on any assignment andwill be treated in accordance with university policy. This policy can be found at:

http://www1.cuny.edu/portal ur/content/2004/policies/image/policy.pdfElectronic versions of laboratory reports will be submitted to Turn-it-in soft-ware and checked for academic dishonesty. Turn-it-in checks both internetsources and previously submitted reports. We should note here that copying anddownloading graphs from the internet, without express permission of the instructor, con-stitutes plagiarism (even if referenced).

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1.9. Grading

Course grading:

• Safety examination and Pre-laboratory quizzes: 20% of grade• Laboratory assignments: 60% of grade• Laboratory practical: 20% of grade

Any student missing a laboratory with an university excused absence will be given a make-up laboratory during the fourteenth week.

Grading for a Laboratory Project:

(1) Laboratory report∗: 34% of grade• Style guide formatting (7%)• Written report (27%)

(2) Pre-laboratory questions: 20% of grade(3) Report sheet: 13% of grade(4) Post-laboratory questions: 20% of grade(5) Safety and laboratory technique: 13% of grade

∗Failure to submit an electronic version of the laboratory report through Turn-it-in viaBlackboard will result in the loss of these points.

Grade distribution:A+ 95.5 - 100 % C+ 70.5 - 75.4 %A 87.5 - 95.4 % C 63.5 - 70.4 %B+ 82.5 - 87.4 % D 49.5 - 63.4 %B 75.5 - 82.4 % F < 49.5 %

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Introduction

SECTION 2Schedule

Date Session Pre-laboratory (60 min) Laboratory (120 min) Projects due1 Syllabus and safety Experiment 1: Basic

laboratory techniques2 Safety exam. Significant

figures and error analysis.Experiment 2: Density. Experiment 1

Project.3 Pre-laboratory quiz.

Chemical Formulas.Experiment 3: Law ofdefinite proportions.

Experiment 2Project.

4 Pre-laboratory quiz.Formulas and reactions.

Experiment 4: Stoi-chiometry.


5 Pre-laboratory quiz.Chemical reactions.

Experiment 5: Prepara-tion of a simple salt


6 Aqueous chemistry. Experiment 6: Reac-tions in Solution.


7 Pre-laboratory quiz.Aqueous chemistry.

Experiment 6: Reac-tions in Solution.

8 Pre-laboratory quiz.Aqueous chemistry.

Experiment 6: Reac-tions in Solution.

Experiment6 Project:Knowns.

9 Pre-laboratory quiz.Acid/base chemistry

Experiment 7:Acid/base titrations

Experiment 6Project: Un-knowns.

10 Pre-laboratory quiz.Acid/base chemistry.

Experiment 8: Potentio-metric titrations


11 Pre-laboratory quiz.Heat and heat capacity.

Experiment 9: Heat ofreaction.


12 Pre-laboratory quiz.Electron spectroscopy.

Experiment 10: Emis-sion and Beer’s Law.


13 Pre-laboratory quiz.Metal reactivity andperfect gases

Experiment 11: Activityof metals


14 LABORATORY PRACTICAL. Check-out. Make-up laboratory (examination andlaboratory).


c©2008 QC Chemistry and Biochemistry5

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Introduction

SECTION 3Laboratory Safety

3.1. Self protection

• Wear safety goggles at all times while in the laboratory, even if you have completedyour experiment. Prescription safety glasses can be worn, but students must obtainapproval from the instructor. Contact lenses, if worn in the laboratory, do not takethe place of safety goggles. In fact, some vapors can accumulate under the lens andcause damage to the eyes and, therefore, prescription glasses are a better choicefor laboratory work.

• Bare skin must be minimized while in the laboratory by wearing clothing thatcovers one’s feet, legs and body completely. Hence, closed shoes and full-lengthpants are required, since broken glass and spilled chemicals are all too common onthe floors of chemistry laboratories. Sandals, flip-flops, short skirts, shorts, three-quarter length pants, bare midriffs, bare backs and bare shoulders are not allowed.Long-sleeve shirts are recommended but not required. Long hair should be tiedback. Hats are not allowed. Many synthetic materials are highly flammable andshould not be worn in the laboratory.

• No horseplay, joking or playing is permitted in the lab. Failure to obey this ruleis cause for immediate expulsion from the laboratory.

• No eating or drinking is permitted in the laboratory or in the prelaboratory class-room. The chewing of bubble gum is considered eating in the laboratory.

• No visitors are allowed in the laboratory. Your friends should visit with you beforeor after, but not during the laboratory session.

• Jewelry such as rings, bracelets and watches should be removed, since chemicalscan cause severe irritation when trapped under a piece of jewelry.

• Long hair should be secured. Long necklaces, neckties and/or scarves should beremoved.

• Never taste, smell or touch a chemical or solution unless specifically directed byyour instructor to do so.

• Wear disposable gloves that are appropriate for the chemical you are working with.Take them off and wash hands before you leave the laboratory.


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• Wash your hands after handling chemical and/or reagent bottles. ALWAYS washyour hands with soap and water before leaving the laboratory at the end of theperiod.

• Cell phones are a distraction. Therefore, cell phones must be turned off dur-ing the laboratory period. If a cell phone rings, this phone will be confiscatedby the laboratory instructor for the duration of the laboratory.

3.2. Laboratory accidents

• Know the location of all eye-wash, shower stations, fire extinguishers and fireblankets in the laboratory.

• Report all accidents and injuries to your laboratory instructor immediately.• If an accident occurs, do not panic! Alert your instructor immediately. Then, take

appropriate action regarding the accident (i.e., seek aid for an injury, flush withwater, clean up chemical, etc.)

• Whenever a chemical comes into contact with skin (hands, arms, etc.) flush theaffected area with water for several minutes. Then wash thoroughly with soap andwater. If the area of contact is the eyes or face, use the eyewash fountain to flushthe chemicals. Do not rub the area with your hands before washing the area.

• If the chemical spills over a large part of your body, remove all contaminatedclothing immediately. Modesty is not an issue! Use the safety shower to washarea for at least 15 minutes. Follow any first aid procedures given on the MaterialSafety Data Sheet (MSDS).

• For abrasions, cuts or minor burns, flush the area with water and consult with thelaboratory instructor for further treatment.

• Chemical spills on the bench or the floor should be treated using the followingsteps:(1) Alert your neighbors and the laboratory instructor immediately.(2) Clean up the spill as directed by the laboratory instructor.(3) If the substance is volatile, flammable, or toxic, warn everyone in the labora-

tory of the accident.• Know the individual hazards for all chemicals used during the laboratory experi-

ment. Material Safety Data Sheets (MSDSs) for all chemicals are available at(1) http://msdssearch.com/(2) http://cunyqueens.chemwatchna.com , username: queensmsds , password:

msds

3.3. General laboratory safety

• Do not place any objects (including pens or pencils) that have been placed on thelaboratory bench in your mouth, since these objects may have picked up contam-inants from the laboratory bench.

• Do not work in the laboratory unsupervised.• Do not pipette liquids by mouth. Use a bulb to siphon liquids into a pipette.• Read the experiments and exercises before coming to class. This will familiarize

you with any potential hazards that may exist or evolve during the exercise. Payparticular attention to any information concerning handling or safety of particularchemicals or solutions.

• Do NOT perform unauthorized experiments or deviate from the experimental plan.Report unauthorized experiments to the instructor.

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Fig. 3.1: NFPA label.

• Assemble your laboratory apparatus at least 8 inches from the edge of the labora-tory bench.

• Maintain a clean and orderly laboratory desk and drawer. Keep drawers or cabinetsclosed and aisles free of obstructions. Do not place book bags, athletic equipmentor other items on the floor near the laboratory bench.

• Be aware of the actions of your neighbor as well as yourself. You could be thevictim of a mistake made by a neighbor. Therefore, advise them of impropertechnique or unsafe practices. If necessary report them to the instructor.

3.4. Chemical handling and disposal

• Read the label on a bottle at least twice before using it! Using the wrong chemicalin an experiment will result in erroneous results in your experiment and may leadto a serious accident.

• Avoid removing large excesses of reagent from the bottle. Only dispense from thebottle the amount of reagent that the experiment requires.

• Never return excess chemicals to the reagent bottle.• Do not touch, taste or smell chemicals.• Use the fume hoods to pour noxious or irritating chemicals, and to run chemical

reactions that generate noxious or irritating products.• For additional information about the safe handling of chemicals (including infor-

mation about dealing with laboratory fires), please see the section entitled WorkingSafely with Chemicals.

• Dispose of chemicals using the guidelines given in the section entitled Overview ofHazardous Waste Disposal Procedures for Students.

• Chemicals are often labeled according to NFPA (National Fire Protection Asso-ciation) standard using the label shown in Fig. 3.1. This safety sticker has fourfields, namely blue (health), red (fire), yellow (reactivity) and white (special). Thenumbers in the blocks of the stickers range from zero (0) to four (4) and indicatehazard severity, with zero being the least hazardous.

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3.5. Cleanliness

It is important to keep the laboratory as clean as possible, for safety reasons as well asaesthetic reasons. Each pair of students is responsible for their immediate desk area. Beforeleaving the laboratory, students should make sure that the area of the laboratory bench neartheir assigned drawer is clean and dry, that Bunsen burners and other shared equipmentare put away in the appropriate space, and that the trough to the sink is free of any solidmaterial. Laboratory instructors will check work areas before approving completion ofthe experiment. Pairs of students will be assigned dates for which they are responsiblefor cleaning the reagent shelves, balances and balance tables in the weighing room, andsurfaces under the hoods. This duty will be rotated so that each pair of students will beresponsible for general laboratory cleanliness at least once during the semester.General tips:

• Place broken glassware in the broken glassware box.• Keep drawers and cabinets closed to avoid physical hazards.• Never place materials or chemical bottles on the floor.

3.6. Other information

If you have or suspect you may have any of the following conditions, please inform yourlaboratory instructor before attending your laboratory:

• Pregnancy.• Wear contact lenses.• Wear synthetic finger nails (which are highly flammable).• Chronic breathing problems.• Immune system suppression.• Chronic Anemia.• Treatment with prescription drugs which may affect judgement.

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3.7. Working Safely With Chemicals

Supplied by Queens College Laboratory Safety Officer, Summer 2008.Last modified on August 26, 1998.

Because few laboratory chemicals are without hazards, general precautions for handling alllaboratory chemicals should be adopted.

• It is prudent to minimize all chemical exposures. Precautions should be taken toavoid exposure by the principal routes of entry, that is, contact with skin and eyes,inhalation and ingestion.

I. What Is A Hazardous Chemical?Any chemical that can harm a person or the environment.

Personal hazards of chemicals fall into two major groups: Health Hazards and PhysicalHazards.

(1) Health Hazards (acute or chronic health effects):These chemicals include carcinogens, toxic or highly toxic agents, irritants,

corrosives, sensitizers, and agents which damage the lungs, skin, eyes, or mucousmembranes.

(2) Physical Hazards:Chemicals that are either a combustible liquid, or a compressed gas, explosive,

flammable, an organic peroxide, an oxidizer, pyrophoric, unstable (reactive), orwater-reactive.

II. How Can You Protect Yourself?

(1) Protect Your Skin From Chemical Splashes:Wear long sleeved shirts and a skirt or long pants. Do not wear shorts or a

miniskirt. Long hair and loose clothing or jewelry must be confined when workingin the laboratory. Unrestrained long hair, loose or torn clothing, and jewelrycan dip into chemicals or become ensnared in equipment and moving machinery.Clothing and hair can catch fire. Because synthetic fabrics are flammable and canadhere to the skin, they can increase the severity of a burn. Therefore, cottonis the preferred fabric. It is advisable to wear a laboratory coat. Wear shoesmade of leather; do not wear open-toed shoes, sandals, clogs, or canvas sneakers.

Wear disposable gloves that are appropriate for the chemical you are workingwith. Take them off and wash your hands before you leave the laboratory.

(2) Protect Your Eyes From Chemical Splashes:Wear safety glasses or goggles. They have side shields and a closed top to

protect your eyes from splashes of chemicals. Ordinary prescription glasses do notprovide adequate protection against injury. Wear safety glasses over prescriptionglasses. Contact lenses offer no protection against eye injury and cannot be sub-stituted for safety glasses and goggles. It is best not to wear contact lenses whencarrying out operations where chemical vapors are present or a chemical splashto the eyes or chemical dust is possible, because contact lenses can increase thedegree of harm and can interfere with First Aid and eye-flushing procedures. If

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you must wear contact lenses for medical reasons, then safety glasses with sideshields or tight-fitting safety goggles must be worn over the contact lenses.

(3) Safe Behavior:• Avoid distracting or startling others.• Do not allow practical jokes and horseplay at any time• Use laboratory equipment only for its designated purpose.• Always wash your hands after completing your work.• Do not eat, drink, take medicine, chew gum, smoke, or apply makeup in the

laboratory.

(4) Safety Equipment:Do not pipette by mouth, use pipetting devices. Work inside a laboratory

fume hood when necessary. Check for adequate air flow (a light piece of tissuehanging from the hood sash may help). Keep the sash lowered as far as possible.Keep your hands in and your head out of the hood.

III. What Can You Do In Case Of Emergency?

(1) Report All Accidents and Incidents To Your Supervisor.This may prevent a similar occurrence in the future.

(2) Chemical Spill Emergency:(a) When a chemical has been spilled on the counter top or the floor:

Use the Spill Cleanup Kit to contain and/or neutralize the spill. Collect allcleanup material in a closed container. Label it “Hazardous Chemical Waste”and place in chemical waste tray.

(b) When a chemical has been splashed on a person:Use water from the Safety Shower or the Eyewash Station or cold water fromthe sink for fifteen minutes of continuous flow. Always get medical help assoon as possible.

(3) Medical Emergency:When a person is overcome by fumes do the following:(a) Evacuate the laboratory.(b) Bring the person who is overcome to fresh air.(c) Get medical attention for the person in question.

(4) Fire Emergency:Know your emergency exits!(a) Fire in the Laboratory

• Fight the fire or flee the area? You safety is the MOST importantconsideration for this question.

• Do NOT fight the fire if there is any possibility that you might betrapped by the fire or smoke.

• Do NOT fight the fire if there is considerable heat, smoke or fumes.• Call the Fire Department before you or someone else starts to fight the

fire. You may need backup.• Tell all others to get out. Leave the laboratory.• Close all doors.

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• Do NOT use the elevator.• If you do decide to fight the fire and the fire is a small fire on a bench

top, then smother the fire with a watch glass.• If you do decide to fight the fire and the fire is a larger fire, then grab

an appropriate fire extinguisher and PASS, where PASS stands for:(P) Pull the pin on the fire extinguisher.(A) Aim the extinguisher at the base of the fire.(S) Squeeze the trigger.(S) Sweep the extinguisher from side to side.

(b) A Person on FireDo NOT run. Stop, Drop, and Roll.To use the safety shower, place the person on fire under the shower head, pullthe handle and hold it until the water has extinguished all flames.To use the fire blanket, wrap the person on fire in the blanket and have theperson stop, drop and roll.

(c) Fire Alarm in the BuildingMake sure all bunsen burners are off. Evacuate the laboratory in an orderlyfashion. Close all doors behind you. Do not use the elevator. Meet ata designated meeting place with the rest of your class. Wait until a FireMarshal or Fire Warden says it is safe to return to the building.

Reference:

National Research Council (1995). Prudent Practice in the Laboratory. Washington,D.C. National Academy Press.

Code of Federal Regulations, 29 CFR Part 1910, Subpart Z. U.S. Govt. Printing Office,Washington, DC 20402 (latest edition).

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3.8. Emergency Response Procedure: Queens College – CUNY

Supplied by Queens College Laboratory Safety Officer, Summer 2008.Last modified on May 22, 2008.

FIRE EMERGENCY

In the event of a fire emergency, the following procedure should be followed:(1) Pull the nearest fire alarm. (Do not attempt to put out the fire if you do not

know how to handle a fire extinguisher or if you do not think you can handle theincident.)

(2) Notify others in the immediate area. Close all doors and evacuate to a safe location.(3) Dial 9-911 and provide the operator with the following information:

• The existence of a fire emergency condition.• Specific location of the fire (building, floor, room, etc.).• Your name and location.

(4) Contact Queens College Security Department at 997-5911 or 997-5912 and providethe desk officer with the same information as listed in Item 3.

(5) When fire alarm has sounded, all occupants (faculty, staff and students) shallexit the building and move to a safe location (i.e., gather on the Quadrangle).Follow the directions of the Fire Department personnel, Security personnel and/orassigned Fire Wardens. Use stairways to exit the building. Never use the elevatorsunless directed to do so by Fire Department Personnel. Do not re-enter buildinguntil Fire Department personnel have declared the building safe for occupancy.

MEDICAL EMERGENCY

In the event of a medical emergency, the following procedure should be followed:(1) Dial 9-911 and provide the operator with the following information:

• The existence of a medical emergency condition. Be as specific as possibleabout the incident.

• Specific location of the emergency (building, floor, room, etc.).• Your name and location.

(2) Contact Queens College Security Department at 997-5911 or 997-5912 and providethe desk officer with the same information as in Item 1.

(3) Remain at your location in order to direct emergency personnel unless it mayjeopardize your safety.

HAZARDOUS MATERIALS EMERGENCY RESPONSE

This plan was designed to reduce the potential for overexposure to hazardous chemicals inthe event of a chemical spill. Proper spill containment and cleanup procedures are ordinarilyobtained from material safety data sheets (MSDS’s). However, you should not take it uponyourself to contain or clean up a chemical spill if you:

• Are not familiar with the chemical involved and with the potential hazards asso-ciated with the chemical.

• Do not have the proper personal protective equipment (PPE).• Cannot reasonably be expected to handle the incident.

We should note that you are responsible for familiarizing yourself with thechemicals (including potential hazards and the required PPE) used for a given

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laboratory procedure. If you are unable to address the clean up of a chemical spill basedupon the criteria listed above, then

(1) Evacuate the area.(2) If the spill occurred within a laboratory, notify the instructor or principal investi-

gator.(3) During normal business hours (8:00 am to 4:00 pm), also notify the Environmental

Safety and Health Officer (ESHO) and/or the Laboratory Safety Officer (LSO).During normal business hours, the ESHO and LSO can be reached at:• ESHO: William Graffeo, (718) 997-2881• LSO: Parmanand Panday, (718) 997-4108 or (718) 997-4171• Assistant LSO: Rick Sherrick, (718) 997-4177

(4) During off-hours, notify the Public Safety Desk Officer at (718) 997-5911 or (718)997-5912. The Public Safety Officer will make sure that the instructor (or principalinvestigator), the ESHO and the LSO are informed of the spill.

(5) Let both your instructor (or principal investigator) and Safety Officer (or thePublic Safety Desk Officer) know the following:• The existence of a spill incident.• Specific location of the incident (building, floor, room, etc.).• Type of material spilled, if known.• Any other information that you deem pertinent for the spill clean up.

Once the spill is reported, the ESHO, LSO or Campus Patrol Officer will respond to thescene. The room will be evacuated and the affected area will be cordoned off. This area mayinclude rooms adjacent to the area where the spill has occurred. If an injury has occurred asa result of the spill, the person reporting the spill and/or the Safety Officer will immediatelynotify the New York City Emergency Medical Service (see Medical Emergency above). Ifoccupants are trapped within the affected area, or if the ESHO, LSO or instructor/principalinvestigator cannot be reached, then the Public Safety desk will contact the New York FireDepartment to respond to the scene. For a large chemical spill, an incident report must becompleted and submitted to the ESHO and the LSO. Thus, the Safety Officer will attemptto gather as much information as possible from the individuals present at the site at thetime of the spill. Therefore, do not leave the temporary headquarters until you have talkedto the Safety Officer.

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3.9. Overview of Hazardous Waste Disposal Procedures for Students

Supplied by Queens College Laboratory Safety Officer, Summer 2008.Last modified on May 22, 2008.

Queens College faculty, staff, students, contractors, and other parties that handle or gen-erate hazardous wastes are required to properly handle, store and label hazardous wastesand to comply with applicable federal, state and local regulations. As a student, yourresponsibilities are:

(1) Follow the Queens College Hazardous Waste Program requirements of• Do not dispose of hazardous waste down sink drains.• Do not dispose of hazardous waste in the normal trash.• Do not dispose of hazardous waste by evaporation in fume hood.• Do not dispose of hazardous waste in broken glass container.• Hazardous waste must be collected in a compatible container which is in good

condition.• All containers of hazardous waste MUST be labeled with the word Hazardous

Waste and with other words identifying contents and hazards present.• All hazardous waste containers must be kept tightly capped except when

adding or removing waste.• Do not mix incompatible chemicals together in the same waste container.• Do not store waste containers next to other bottles holding incompatible

chemicals.• Separate incompatible chemicals into separate secondary containment trays.• Store hazardous waste containers at or near point of generation and under

the control of generator.

(2) Review Material Safety Data Sheets prior to working with chemicals.

(3) Use appropriate personal protective equipment when working with chemicals.

(4) Report any accident to laboratory instructor.

(5) Report any emergency to Public Safety (Ext. 7-5911).

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3.10. Laboratory Clean-up Schedule

Session Group Students1

2

3

4

5

6

7

8

9

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Laboratory Safety Agreement

I, the undersigned, have read and understand the safety instructions given in this lab-oratory manual. I agree to abide by these rules. I also understand that failure to obey thesafety rules given above or to follow my instructors advise while in the laboratory can leadto my dismissal from the laboratory for one or more class periods, with a grade of zero forthe missed experiment(s).

As part of the safety lecture during the first laboratory period, an overview of chemicaldisposal and general laboratory safety was discussed. In particular, the items in Section3.9 were reviewed.

All laboratory experiments in this manual have been checked for safety when performedaccording to directions. I, the undersigned, understand that I am responsible for readingall safety precautions required for performing each experiment. Because this is a chemistrylaboratory, I understand that there is the potential for serious accidents if these safetyprecautions are not followed and acknowledge that the fundamental responsibility for safetylies with myself.

Student Name (print): QC Student ID #:

Student Signature: Date:

Instructor Name (print): Instructor Signature:


Introduction

SECTION 4Freshman Chemistry Style Guide

The laboratory report should consist of the following:(1) Abstract – One to four sentences that summarize the experiment(2) Introduction – The introduction should answer, using complete sentences, the

questions given in the pre-laboratory questions.(3) Experimental – The experiment should detail the steps, techniques, and apparatus

used to perform the experiment.(4) Results and Discussion – The result and discussion should give the results obtained

from the experiments and should discuss these results. The details of the calcu-lations are shown on the Report Sheet for the project. Thus, this section shouldnot present the details of the calculations, but should discuss these calculations inwords.

(5) References – Any references including internet sites should be listed.(6) Tables – Any tables referenced in the text.(7) Figures – Any figures referenced in the text.(8) Appendix – The Pre-laboratory question sheet, Report sheet and Post-laboratory

question sheet should be attached to the back of the laboratory report as appen-dices.

4.1. Margins

The margins should be standard paper margins, namelyTop 1” Bottom 0.5”Left 1” Right 1”

The page number should be located 0.5” from the bottom of the page and centered hori-zontally (cf. Figs. 4.1 - 4.7). The text should be 1” from the bottom of the page.

4.2. Line spacing

Double space except for references, tables, table captions, figure captions and title page.See examples for the format of these special cases (cf. Figs. 4.5 - 4.7).

4.3. Font

Times New Roman (or Times Roman or Roman) or Helvetica (or Courier or Arial). 12point type size. Reports written in different fonts will not be accepted.


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4.4. Layout

Section headings should be preceded by a roman numeral and a period, be left justified, useall capital letters and be bold. The body of the text should use full paragraph justification.There should be two lines between each section. Widows (i.e., the last line of a paragraphthat is carried over to a following page) and orphans (i.e., the first line of a paragraphappearing alone at the bottom of a page) will not be tolerated.

4.5. Tables

Tables should be placed at the end of the text after the references in consecutive order.Tables should be numbered sequentially and consecutively from the first table using romannumerals (i.e., I, II, etc.). The table caption should be located at the top of the table. Payattention to significant figures when making tables. No table should be wrapped sothat it exists on two pages. If a table will not fit on a single page, consult your instructorfor how to format long tables. Fig. 4.6 gives an example of a page of tables for a laboratoryreport. Notice that the table is separated from the table caption by a double line and isended by a double line. Also notice that the table headings are separated from the valuesby a single line. You should copy this format when typing tables.

4.6. Figures

Figures should be placed after the tables with one figure/page in consecutive order. Fig-ures should be numbered sequentially and consecutively from the first figure using arabicnumerals. If a figure is composed of multiple parts, each sub-part should be identified witha letter of the alphabet in sequential order from top to bottom. Do not use color on figuresunless required. The figure caption should be placed below the figure.

If data are obtained by scanning a parameter (such as temperature or energy), then thedata should be graphed as a line. If, however, the data are obtained by analyzing differentobjects, then the data should be graphed as markers. If data are presented as markersand a line is drawn through the markers, you should indicate if the line is the result of aregression or is drawn as a help to the reader. Fig. 4.7 gives an example of a line drawnas a regression, while Fig. 4.8 shows an example of a line drawn to help the reader. Notehow the figure captions indicate each type of line.

Figures should not have color backgrounds. The fonts for axis markers and labels shouldmatch the font used in writing the report and should be at least 18 points in size. FiguresWhen a figure is pasted into a word processors (such as Microsoft Word),

4.7. Referencing

References should be placed at the end of the document before the tables and figures. Thereferences should be numbered sequentially and consecutively from the first time of use inthe document. Citing references in the text should use on-line numerals in square brackets(i.e., [1]) which are spaced away from the preceding word or symbol and are placed insidepunctuation. Examples are given below.

• Journal articleA. A. Author, B. B. Author and C. C. Author, “Article title,” Journal Abbr. Vol,start page (year).

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• Non-scientific magazine or newspaperA. A. Author, “Title,” Name of magazine (date published) start page.

• Web siteA. A. Author (if any), Title of document, year. Title of site. http://url.of.site(accessed Month day, year).

• BookA. A. Author, Title of book (Publisher, location, year), start page (if any).

• Edited volumeTitle of book, E. E. Editor, eds. (Publisher, location, year).

• Specific chapter of an edited volumeA. A. Author, “Title of chapter” in Title of book, E. E. Editor, eds. (Publisher,location, year), start page.

4.8. Title page formatting

The title should be simple and concise. The title should be bold, in 14 pt text, with eachword in the title capitalized except for simple words (e.g., a, on, the, an, of, etc). Thetitle is centered on the top of the page. Nonstandard abbreviations and acronyms are notallowed in the title, since these abbreviations cannot be defined in the title. If the title ismore than a single line, then the title is not double spaced. Skip two lines, then list theauthor of the paper. The name of the author should be typed in 12 pt text and centeredon the page. Skip one line, then listed the affiliation of the author. This affiliation shouldbe centered in 12 pt text and in italic. The affiliation is as follows:

Department of Chemistry and Biochemistry, Queens College - CUNYCourse Number, Course Section, Semester

Instructor: Instructor NameSkip two lines, then type the word “Abstract” in 12 pt type, bold, in all capitals, left-justified. Skip a single line and type the abstract following the guidelines given above. Onlythe abstract should be on the title page. All fonts on the title page should be identical tothe fonts in the body of the paper.

An example laboratory report is shown in Figs. 4.1-4.8.

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4.9. Example laboratory report

1

The Density of Various Materials

John X. Smith

Department of Chemistry and Biochemistry, Queens College – CUNY

CHEM 101.1, E5TBA, Fall 2008

Instructor: Ms. Luxi Li

ABSTRACT

The density, defined as mass per unit volume, of a set of regular shaped objects was obtained from

the slope of a graph of mass (measured with an electronic balance) as a function of volume. The

density of 1.16 g/mL indicates that the objects were either polyamide or acrylic. The density of

silver/pink irregular cylinders was determined using Archimede’s principle to be 9.84 g/mL,

indicating that the unknown sample is probably bismuth. Finally, the density of water and

water/ethanol mixtures was obtained. These data were plotted as a function of ethanol concentration

in order to generate a calibration curve. This calibration curve was used to determine that the

ethanol concentration in the unknown sample was 38% by volume.

Fig. 4.1: Example title page for a laboratory report.

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2

I. INTRODUCTION

The identification of a substance is often performed using intensive properties, which are

properties that do not depend on the quantity of the substance. Examples of these properties include

color, odor, melting point, boiling point and density. The density � of a substance is the mass m of

the substance per unit volume V, or � = m / V , (1)

with standard units of kg/m3 (although it is more commonly reported in units of g/cm3 or g/mL).

Density is independent of the quantity of the substance, since both the mass and the volume are

proportional to one another at a fixed temperature. As the temperature changes, the volume of the

substance changes which, in turn, changes the density. In this experiment, we determine the density

of various solid and liquid materials.

II. EXPERIMENTAL

In the first procedure, we obtained a set composed of four regular shaped objects from the

instructor and recorded the code number. We then used a metric ruler to determine the dimensions

of each object with a precision of 0.01 cm. For the cubic object, we calculated the volume V from

these measurements using

V = l × w × h , (2)

where l is the length in cm, w is the width in cm and h is the height in cm. The volume of cylindrical

objects was determined from

V = � × r2 × h , (3)

where r is the radius of the cylinder. To obtain the error in the our volume, we remeasured the

dimensions and recalculated the volume. The mass of the block was determined using an electronic

Fig. 4.2: Example introduction and experimental sections of a laboratory report.

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3

balance. The balance pan was cleaned, a piece of weighing paper was added to the balance, and then

the balance was tared to ensure an accurate measurement of the mass. Each block was individually

placed on the center of the pan (re-zeroing the balance after each measurement) and the mass was

recorded to the nearest 10 mg.

Irregular shaped metal objects (or metal powders) cannot be physically measured in order to

obtain the volume of the sample. However, we were able to determine the volume of the objects by

observing the volume of water displaced by the object. In this procedure, we filled a graduate

cylinder with approximately 5 mL of water. We then obtained a sample of an unknown pink/silver

metal (after recording the code number) and weighed it. After determining the mass, we slowly

added this sample to the water. We tapped the cylinder to dislodge any air bubbles from the metal

sample, and then determined the new water volume. The difference in water volume is equivalent

to the volume of the metal. The density was then calculated using eq. (1) and the identity of the

unknown metal was determined by comparison with data given in [1].

The density of the liquid samples in Part C was evaluated by weighing a 10 mL graduate

cylinder and then adding a volume of the liquid to this cylinder. The mass of the liquid was

determined by the difference in the weight of the graduate cylinder before and after the addition of

the sample. The density was then calculated from eq. (1). The temperature was measured using a

standard alcohol thermometer that was calibrated for the range of 0 – 100 C during Experiment 1.

III. RESULTS AND DISCUSSION

The set obtained from the instructor for Part A consisted of one cubic and three cylinders of red

plastic with code number 2A02. Table I presents the dimensions used in eqs. (2) and (3) to

determine the volume as well as the mass of each objected obtained from the electronic balance.

Fig. 4.3: Continuation of the example experimental section and beginning of the resultsand discussion section of a laboratory report.

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4

Rearranging eq. (1) to give

m = � V . (4)

indicates that a graph of the mass as a function of volume should yield a straight line with an

intercept of zero and with a slope that is equivalent to the density of the objects. Thus, Fig. 1

presents the mass of the red plastic cube and cylinders versus the volume of the objects. A linear

regression of these data gives a density of 1.16 ± 0.01 g/cm3. Using standard densities provided in

[1], we determined that the unknown plastic blocks and cylinders are either acrylic or polyamide.

Distinguishing between acrylic or polyamide would require additional physical or chemical tests.

In Part B, we were given metallic pellets (Code 2B05) that had a slight pinkish tint. The mass

of our sample was 8.451 g, while the volume of displaced water was 0.87 ± 0.01 mL. Substitution

of these data into eq. (1) yields a density of 9.7 ± 0.1 g/mL, or 9.7 ± 0.1 g cm-3 (since 1 mL = 1 cm3).

Comparison of this result to standard densities in [1] indicates that the metallic pellets are bismuth.

Since bismuth has a silver pink color [2], the color of our sample gives additional evidence for our

identification of the unknown sample as bismuth.

Table II gives the volume and mass of ethanol/water mixtures at various concentrations, along

with the density determined using eq. (1). This density is plotted as a function of ethanol

concentration in Fig. 2 to generate a calibration curve for ethanol/water density. A volume of 6.20

mL of the unknown sample 2C03 had a mass of 5.756 g, while a volume of 6.80 mL had a mass of

6.348 g. Thus, the unknown sample 2C03 has a density of 0.931 ± 0.003 g/mL. The calibration

curve in Fig. 2 indicates that a density of 0.931 g/mL corresponds to an ethanol concentration of 38

± 2% by volume, or 76 proof alcohol.

In this experiment, we determined the density of various regular shaped objects by direct

determination of the volume (using a ruler to measure the dimensions) and mass (using an electronic

balance). From our data set, we obtained a density of 1.16 ± 0.01 g/cm3, indicating that the objects

Fig. 4.4: Continuation of the example results and discussion section of a laboratory report.

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were made of either acrylic or polyamide. We also obtained the density of an unknown metal

sample by indirectly measuring the volume (using water displacement) and directly measuring the

mass. Our density of 9.7 ± 0.1 g/mL indicates that the unknown pinkish metal is probably bismuth.

Finally, we generated a calibration curve of density as a function of ethanol concentration for

ethanol/water mixtures. This calibration curve allowed us to determine that the unknown sample

had 38% by volume ethanol.

References

1. “Experiment 2: Density” in Chemistry 113.1. Introduction to Chemical Techniques Laboratory

Manual, C. M. Evans, F. H. Watson and G. L. Findley (Queens College, New York, 2008), p.

36.

2. Bismuth, 2008. Wikipedia. http://en.wikipedia.org/wiki/Bismuth (accessed August 13, 2008).

Fig. 4.5: Continuation of the example results and discussion section and references of alaboratory report.

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Table I. The dimensions, volume V [calculated using eqs. (2) and (3)], and mass m of the four red

plastic objects in package 2A02. The dimensions for the cubic object are given as length × width

× height. The dimensions for the cylindrical objects are given as radius × height.

Object Dimensions (cm) Volume (cm3) Mass (g)

1 2.01 ± 0.01 × 2.00 ± 0.01 × 2.02 ± 0.01 8.12 ± 0.12 9.419

2 1.27 ± 0.02 × 3.01 ± 0.02 15.3 ± 0.6 17.748

3 1.27 ± 0.01 × 6.00 ± 0.01 30.4 ± 0.5 35.26

4 1.27 ± 0.02 × 9.00 ± 0.02 45.6 ± 1.5 52.90

Table II. The volume Vi and mass m

i [where i = A, B for the measurements obtained by myself and

my laboratory partner, respectively] for various ethanol (EtOH) concentrations (% by volume). The

density �

i (i = A, B) for each measurement was determined using eq. (1). The temperature for all

measurements was 23.4 C.

% EtOH VA (mL) m

A (g) V

B (mL) m

B (g) �

A (g/mL) �

B (g/mL)

0 5.60 5.548 7.22 7.182 0.991 0.995

18 6.42 6.193 7.60 7.350 0.965 0.967

36 5.70 5.343 7.80 7.318 0.937 0.938

54 6.42 5.831 7.60 6.848 0.901 0.904

72 6.38 5.447 8.50 7.350 0.865 0.860

90 6.59 5.435 7.40 6.207 0.839 0.832

Fig. 4.6: Example tables for the laboratory report.

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Fig. 1. The mass m (g) of various red plastic objects plotted as a function of the volume V (cm3) of

the object. The solid line represents a linear least square analysis of the experimental data with a

regression equation of m = 1.16 ± 0.01 g cm-3 × V.

60

50

40

30

20

10

0

m (

g)

6050403020100

V (cm3)

Fig. 4.7: Example graph for the laboratory report.

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Fig. 2. The density �

(g/mL) plotted as a function of ethanol concentration (% by volume) for

various ethanol/water mixtures. The solid markers are the data obtained for the stock solutions. The

open marker indicates the density and volume position for unknown sample 2C03. The solid line

is provided as a visual aid.

1.00

0.95

0.90

0.85

0.80

(

g/m

L)

100806040200

% EtOH (by volume)

Fig. 4.8: Example graph for the laboratory report.

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Introduction

SECTION 5Instructions for Turn-it-in Assignments

All laboratory reports must be submitted both in hard copy to the instructor and asan electronic version to Blackboard. The electronic version will be checked for instances ofacademic dishonesty using Turn-it-in software. To submit the electronic version,

(1) Go to http://www.cuny.edu and log-in to the CUNY portal.(2) Go to the Blackboard website and to the Chemistry 113 course site for your section.(3) Select the Lab Reports folder under the Assignment link (cf. Fig. 5.1).(4) Within the Lab Reports folder (cf. Fig. 5.2), select View/Complete under the lab

report that needs to be submitted.(5) Type your first and last name in the appropriate forms (cf. Fig. 5.3).(6) Type the title of the laboratory report in the appropriate form (cf. Fig. 5.3).(7) Click on the Browse button (cf. Fig. 5.3) to load the dialog that will allow

you to select the file to be uploaded. Although the website lists several fileformats that are acceptable, only Word or Wordperfect files are validsubmissions for this course.

(8) Select the file to be uploaded and then click okay. This will return you to theBlackboard site with the path for the file placed in the file form.

(9) Click the Submit button (cf. Fig. 5.3) to submit the paper.(10) Once the paper is successfully uploaded (cf. Fig. 5.4), click the OK button to

return to the Lab Reports folder (cf. Fig. 5.2) on the Blackboard Chemistry 113site.


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Fig. 5.1: The Assignments section of the Blackboard Chemistry 113 site.

Fig. 5.2: The Lab Reports folder within the Assignments section of the Blackboard Chem-istry 113 site.

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Fig. 5.3: The View/Complete page for a Turn-it-in assignment.

Fig. 5.4: The page that appears after successful submission of an electronic laboratoryreport.

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Introduction

SECTION 6Useful information

6.1. Experiment 1. Basic Laboratory Technique

Thermometer calibration line:


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Experiments

EXPERIMENT 1Check-in and Basic Laboratory Techniques

1.1. Safety

During Check-in, you may encounter broken or chipped glassware. This glassware shouldbe handled with care to prevent injury. Broken and/or chipped glassware should be placedin the broken glassware box. Bunsen burners, hot glassware, metal ring stands, and boilingwater can cause painful and serious burns to skin. Hot glassware does not glow and,therefore, looks identical to glassware at room temperature. Thus, be careful when handlinghot glassware.

1.2. Check-in

You will be assigned a laboratory bench and drawer containing your laboratory kit. Youwill be responsible for all items in this kit. Check-in allows you to confirm that your kit hasall required glassware and to replace glassware that is broken or too dirty to clean. This isthe only day when missing or broken items are replaced free of charge. Thus, you shouldreport all items that are missing, scratched, corroded or otherwise unfit for use. You shouldalso use this time to clean all of the glassware in the kit, since this will save time in thefuture. The condition and cleanliness of all items will be spot checked during the semester.At the end of the semester, check-out will be performed to ensure that all items in thelaboratory kit are still in good shape and clean. Illustrations of the items that should bein the laboratory kit are given in Figs. 1.1 and 1.2. General rules about the laboratory kitare:

• This kit is your responsibility. Missing items are replaced free of charge only onthe day of check-in.

• If you drop the course, you must report to room 214 to check-out immediately.• Students completing the course must check out with their regularly scheduled last

lab class. There will be no late check-out.• At Check-out, missing or broken items are billed to the student. Other items must

be clean, dry and flawless to be accepted for check-out.• After the last scheduled day of classes, students who have not checked-

out will be charged a $50 fee in addition to the fee from missing, brokenand dirty items.


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Fig. 1.1: Glassware located in the laboratory kit.

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Fig. 1.2: Other non-glass items located in the laboratory kit.

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These experiments are meant to provide your first experience in the laboratory. If youhave prior laboratory experience, these exercises constitute a good review. Your instructorwill give you the rotation order for the experiment. Submit all results on the Worksheetprovided.

1.3. Experiment 1A. Check-in

Using the Desk Assignment Sheet, check your laboratory kit to ensure that all glasswareand equipment is present and in good working order. Wash and dry all glassware duringthis time. If glassware is too dirty to be cleaned, return this glassware to the stockroom tobe replaced.

1.4. Weighing

You must learn how to use laboratory balances. As always, the limit of readable precisionof the scale should be recorded. When approaching the balance you will need the following:

(1) The substance to be weighed;(2) A container or holder for the sample while on the balance;(3) A sample handling device such as a spatula; and(4) Your report sheet and a pencil or pen to record your measurement.

Balances in this laboratory have a semi-automatic tare (an allowance for mass of the con-tainer or holder). On an electronic balance, ”tare equals zero” is set by depressing the bar,which is also the on/off switch: up for off, down for on, down again to tare. These balanceshave an automatic range selector that will change the readout precision automatically to0.01 g when the gross mass on the pan is over 35 g. The measurement precision for smallmasses is best if very light containers are used, such as the glassine weighing paper for drysolid samples. The precision for samples less than 30 g gross mass is 0.001 g.

1.5. Experiment 1B. Mass

(1) Obtain a bag of pennies from your instructor. Record the code on the bag in theappropriate location on the Experiment 1 Report Sheet.

(2) Without taring, place a sheet of paper on the pan. Read and record its mass inthe table on the Report Sheet. When possible, use the cover to protect the panfrom air drafts to obtain higher precision.

(3) Add the pennies to the pan one at a time, reading the mass after each addi-tion. Enter each in the table given on the Report Sheet in the column enti-tled“Cumulative Mass.” Keep the pennies in order. Calculate the mass of eachpenny by subtracting; record the differences in the column entitled “Mass by Dif-ference.”

(4) Remove all the coins and weigh each individually, taring to zero. Record the massof each coin in the table in the column entitled “Direct Weighing.”

(5) Calculate the mean or average mass m of the pennies, with

m =1n

∑m , (1.1)

where n is the total number of pennies.(6) Calculate the absolute deviation (d = |m−m|) from the mean of each mass (direct

weighing) and then determine the average deviation d, where

d =1n

∑d . (1.2)

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Fig. 1.3: Bunsen burner connected to gas valve.

Consider two hypotheses:

(1) All Lincoln-head pennies are manufactured with equal mass (within 0.001 g), buttheir various histories result in different masses when measured.

(2) New Lincoln-head pennies are lighter than older ones.

Observation is complicated by the various histories of the pennies. Some typical prob-lems in chemistry are illustrated here. Most obviously, the state of corrosion of the penniesrepresents an uncontrolled experimental variable which can be important. Had we usednon-circulated coins, we would have expected better precision. On the other hand, penniesmay not be very uniform even when new. Another question arises: How much experimentaldifference is sufficient and how consistently must it be observed for us to consider two datasets, or groups of data sets as distinctly different? This is an important question for whichstatistical methods provide answers. Which hypothesis do you choose, and why?

1.6. Experiment 1C. Length

While considering precision of reading a balance, select a wooden splint and measure itslength on the inch scale and the centimeter scale. Place your results, along with those ofyour laboratory partner, in the appropriate table on the Report Sheet. Pay attention tothe precision of the ruler during these measurements. What is the deviation and averagedeviation?

Convert the lengths measured in inches to centimeters using the factor 2.54 cm/in.Similarly, convert the lengths measured in centimeters to inches. Pay attention to significantfigures in your results.

1.7. Bunsen Burner

When selecting a burner, check to see that the gas needle valve on the bottom will closecompletely. Also check to see that the barrel of the burner will screw in and out so thatthe air supply to the flame can be controlled.

With the burner gas valve off and the hose connected to the burner and the bench gascock (see Fig. 1.3), turn the bench gas cock fully on and check for leaks around the burnerwith a match. With the air vents closed, open the burner gas valve and light the flame.The flame should be yellow and luminous.

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Stick a test tube into the flame briefly. You should observea deposit of carbon black on the tube. When hydrocarbons suchas methane [CH4 (natural gas)] burn in too little air (oxygen), thereaction is

CH4 (gas) + O2 (gas) → C (sol) + 2 H2O (liq) .With a bit more air, the flame becomes hotter and blue, but carbonmonoxide is formed:

2 CH4 (gas) + 3 O2 (gas) → 2 CO (gas) + 4 H2O (liq) .Now adjust the air supply – you may also have to adjust the gas

with the burner valve – until the flame resembles the schematic tothe right. This is the hottest flame and is characterized by a blueinverted cone shape within the flame that is the so-called reducingflame. A little above the apex of the cone is the hottest area in theflame, reaching temperatures around 1500◦C.

Towards the top of the flame, conditions are oxidizing (high tem-peratures, excess O2). The well-adjusted flame completely convertsmethane and oxygen to carbon dioxide and water:

CH4 (gas) + 2 O2 (gas) → CO2 (gas) + 2 H2O (liq) .To turn off the Bunsen burner, execute the lighting procedure in

reverse. Shut the gas valve at the base of the burner, then close theclose the bench gas cock.

Schematic of abunsen burnerflame when theflame is thehottest.

1.8. Experiment 1D. Understanding Flames

Take a wooden splint and hold it with one end resting on the top of the burner. Notice howthe splint is burned only on the edges of the flame. The flame under the cone is relativelycool (about 350◦C). Place the other end of the splint higher in the flame. Notice that nowthe splint ignites uniformly. Give the splint to your instructor.

1.9. Experiment 1E. Thermometer Calibration

Fill a 100 mL beaker with 50 mL of ice. Then, cover this ice with water and stir. Insertthe thermometer and observe the temperature. When the temperature remains constant,record the temperature as the melting point of ice (or freezing point of water).

Fill a 250 mL beaker with 150 mL of water. Heat the water to boiling over a bunsenburner (cf. Fig. 1.4). When the water is boiling, immerse the thermometer and read thetemperature. (Do not let the thermometer bulb touch the bottom of the beaker during thisprocedure.) Once the temperature is constant, record the temperature as the boiling pointof water.

If the results that you obtained from this procedure give a freezing point and boilingpoint of water that differs from the theoretical temperatures by the same amount, thenthis amount can be used to correct the temperatures recorded using this thermometer.If the difference between observed values and theoretical values is not identical for boththe freezing point and boiling point, then plot the observed values versus the theoreticalvalues. The line between these points will give a calibration line that can be used to correcttemperatures.

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Fig. 1.4: Hot water bath schematic.

1.10. Experiment 1F. Measuring volumes

Fig. 1.5: Reading a graduated cylin-der.

Graduated cylinders are used for measuring vol-umes of a liquid. In your laboratory kit, you havetwo graduated cylinders: a 10 mL graduated cylin-der and a 100 mL graduated cylinder. When aliquid is placed in a graduated cylinder, the liquidlevel in the cylinder will curve with the lowest (orhighest) point being in the middle. This point isthe meniscus and represents the point where oneshould determine the volume. To avoid errors indetermining the volume, the line of sight must beperpendicular to the scale (cf. Fig. 1.5.

Obtain a 250 mL graduated Erlenmeyer flaskand fill with water to the 50 mL mark. Transferthe water, completely and without spilling, to the100 mL graduated cylinder. Record the volume tothe correct precision. Similarly, obtain a 125 mLgraduate Erlenmeyer flask and fill with water to the50 mL mark. Measure the volume and record to the correct precision.

Fill a 50 mL graduated beaker with water to the 40 mL mark. Transfer the water tothe 100 mL graduated cylinder. Record the volume to the correct precision. Similarly, filla 250 mL beaker to the 50 mL mark. Record the volume to the correct precision.

Obtain a sample from the laboratory instructor. Record the volume of this sample tothe correct precision using the 10 mL graduated cylinder.

1.11. Laboratory Project 1 requirements

For this laboratory experiment, the laboratory project consists of the laboratory techniqueand safety score (2 pts), the Experiment 1 Report Sheet (10 pts) and the Experiment 1Post-laboratory questions (3 pts). Pre-laboratory questions and a laboratory report arenot required to complete this laboratory project.

46


47

Desk Assignment Sheet: Chemistry 113.1

Name: Student ID #:

Desk Number: Section: Date:

Combination: Lab Partner:

¤ 1 Beaker, 50 mL ¤ 1 Evaporation Dish-procelain¤ 1 Beaker, 150 mL ¤ 2 Funnels, Short stem¤ 2 Beaker, 250 mL ¤ 1 Iron Ring¤ 1 Beaker, 400 mL ¤ 1 Medicine dropper¤ 1 Beaker, 600 mL ¤ 1 Metric ruler¤ 1 Boiling flask ¤ 1 Nichrome triangle¤ 2 Bottles, wide mouth ¤ 10 Test tubes (sm), 10 × 75 mm¤ 1 Bottles, Plastic wash, 250 mL ¤ 9 Test tubes (md), 15 × 180 mm¤ 2 Clamp, Bunsen ¤ 2 Test tubes (lg), 20 × 250 mm¤ 2 Clamp Fastener ¤ 1 Test tubes rack¤ 3 Crucibles (bottom) ¤ 1 Test tube rack-micro¤ 3 Crucible lids ¤ 1 Test tube brush¤ 1 Crucible tongs ¤ 1 Test tube holder, wood¤ 1 Erlenmeyer flask, 125 mL ¤ 1 Thermometer¤ 1 Erlenmeyer flask, 250 mL ¤ 2 Unknown vials¤ 1 Graduate cylinder, 10 mL ¤ 1 Watch glass, 3”¤ 1 Graduate cylinder, 100 mL ¤ 1 Watch glass, 4”¤ 1 Nickel spatula ¤ 1 Wire gauze

Check In:


Checked in by: Date:

Check Out:


Checked out by: Date:

48


49

Experiment 1. Report Sheet

Name: Section:

Partner: Date:

Results

Fill out all tables and answer all questions below. All values must have appropriatesignificant figures and units. Show these data to your instructor before leaving class.

Experiment 1B. Mass

Table B.I. The cumulative mass of the sample, obtained by adding each penny to the col-lection of pennies on the balance, and the mass of each penny (determined by directlyweighing the individual pennies).

Source Date of penny Cumulative Mass Direct Weighing

Weighing Paper NAN NAN

1st penny

2st penny

3st penny

4th penny

5th penny

Describe each penny used in this study.

1st penny:

2nd penny:

3rd penny:

4th penny:

5th penny:

50

Experiment 1C. Length

Table C.I. The length of a wood splint provided by the instructor measured in inches andin centimeters by myself (Measurement A) and my laboratory partner (Measurement B).Determine the average length for the sample.

Measurement A in cm

Measurement B in cm

Average in cm

Experiment 1D. Understanding Flames

Write your name in the center of the wood splint and give to your instructor.

Experiment 1E. Thermometer Calibration

Table E.I. Measurement of the melting point and boiling point of water by myself (Mea-surement A) and my laboratory partner (Measurement B).

Measurement A Measurement BObserved melting point (Tm) of ice

Thermometer correction using Tm

Observed boiling point (Tb) of water

Barometric pressure (P )

Actual Tb of water at P

Thermometer correction using Tb

51

Experiment 1F. Measuring volumes

Table F.I. The volume of water as determined using the graduations in the glassware (Row1) and the graduated cylinder (Row 2). The error in volume (i.e., |Row 1 - Row 2|) shouldbe given in Row 3.

125 mL Erlen-meyer flask

250 mL Erlen-meyer flask

50 mL Beaker 250 mL Beaker

Volume (mL)in glassware

Volume (mL)using gradu-ated cylinderError in vol-ume of glass-ware

52

Calculations

Fill out all tables and answer all questions below. All values must have appropriatesignificant figures and units. To receive full credit all details for all calculations mustbe shown.

Experiment 1B. Mass

Table B.II. The mass of each penny determined by the difference using the cumulative massin Table B.I and the mass of each penny determined by direct weighing. From these twomeasurements, calculate the average mass of each penny. Then, using the space below,determine the mean [i.e., eq. (1.1)] and the deviation from mean [i.e., eq. (1.2)].

Source Mass byDifference

Mass by DirectWeighing

Average Mass Deviation fromMean

1st penny

2st penny

3st penny

4th penny

5th penny

Mean NAN

53

Experiment 1C. Length

(1) What is the deviation from mean in inches for the values given in Table C.I? incentimeters?

(2) Convert the average length in inches to centimeters using the factor 2.54 cm/in.

(3) Convert the average length in centimeters to inches using the factor 2.54 cm/in.

Experiment 1D. Understanding Flames

(1) When the bunsen burner had a yellow, luminous flame, what was deposited on thetest tube? How is this deposit related to the yellow flame color?

(2) When the flame was set to a two zone flame, in which zone did the wooden splintburn first? Why?

54

Experiment 1E. Thermometer calibration

Plot the average for the two temperatures recorded with your thermometer in ◦C as afunction of the theoretical values on the graph below. Be sure to label all axis of the graphand to use the appropriate number of significant figures when labeling the graph. Use thisgraph and the space provided below to determine the calibration line for your thermometer.This calibration line will be used for all temperatures measured during this semester and,therefore, should be recorded in Section 6.

55

(1) Ethanol has a boiling point of 78.4◦C and a melting point of -114.3◦C. If yourthermometer was placed into boiling ethanol, what temperature would it read? Ifthe ethanol was freezing, what temperature would your thermometer give?

(2) Fahrenheit and Celsius are related by

t(◦F) =(

9◦F5◦C

)t(◦C) + 32◦F ,

while the relationship between Celsius and Kelvin is

T (K) = (t(◦C) + 273.15◦C)(

1K1◦C

).

Using these relationships, convert the boiling point and freezing point of ethanolto Fahrenheit and Kelvin. Be sure to report your answers with the correct unitsand to the correct number of significant figures.

56

Experiment 1F. Measuring volumes

(1) Percent error in this measurement is given by

% Error =|MA −ML|

MA× 100 ,

where MA is the more accurate measurement and ML is the less accurate mea-surement. Determine the percent error in the graduation of the Erlenmeyer flasksand the beakers.

57

Experiment 1. Post-laboratory questions

Name: Section:

Partner: Date:

(1) A student used a graduated cylinder having volume markings every 1 mL to care-fully measure 100 mL of water for an experiment. A fellow student said that byreporting the volume as “100” mL in her lab notebook, she was only entitled toone significant figure. She disagreed. Why did her fellow student say the reportedvolume had only one significant figure? Considering the circumstances, how manysignificant figures are in her measured volume? Justify your answers.

(2) A healthy dog has a temperature ranging from 37.2 to 39.2◦C. Is a dog with atemperature of 103.5◦F within normal range?

(3) Natural gas is mostly methane, a substance that boils at a temperature of 111 K.What is its boiling point in ◦C and in ◦F?

58



Experiments

EXPERIMENT 2Density

2.1. Safety

Certain metals, such as lead, can be toxic when ingested. Wash your hands after handlingall samples. Laboratory ethanol has trace impurities of methanol and toluene. Therefore,laboratory ethanol should not be ingested. Ethanol and ethanol/water mixtures are alsoflammable and, therefore, should be kept away from open flames.

2.2. Introduction

Density ρ is defined as the ratio of the mass m of a sample to its volume V , or

ρ =m

V. (2.1)

Mass and volume are extensive properties of matter – properties that depend on thequantities of substances. Such properties are not of themselves useful in characterizing sub-stances. Intensive properties, on the other hand, are useful in characterizing substances.Intensive properties are often determined by ratioing two extensive properties measured atconstant temperature T and pressure P . Density is an example of this kind of intensiveproperty. When measured under known conditions of T and P , density can be used tocharacterize substances. Of course, two or more substances may have the same density,but for a given substance there is only one density (at constant T and P ). Thus, if youdetermine that a colorless liquid has a density of 1.00 g/mL at 4◦C and 1 atm, this doesnot prove the liquid is water. This fact is simply one piece of evidence that the substancemay be water.

2.3. Experiment 2A. Density of regular-shaped objects

For objects with regular shapes, the volume can be easily determined by measuring thedimensions of the object with a ruler or calipers and then using the appropriate formula.In this experiment, you will determine the density of regular shaped objects.

(1) Obtain a set of objects from your instructor. Record the code on the container inthe appropriate spot of the Report Sheet.

(2) Determine the dimensions of the objects using a ruler. This step should be per-formed by both you and your laboratory partner. Be sure to record the measure-ments to the correct precision. Place this in the appropriate table on the ReportSheet.


59

60

Table 2.1: Densities ρ (in g/mL) of various materials.

Material ρ (g/mL) Material ρ (g/mL)acrylic 1.17 maple 0.71 - 0.83aluminum 2.71 molybdenum 10.22antimony 6.62 polyamide (nylon) 1.15balsa wood 0.10 - 0.20 polypropylene 0.90bismuth 9.80 polytetrafluoroethylene (teflon) 2.2cadmium 8.65 polyurethane 1.23cedar 0.38 polyvinylchloride 1.38copper 8.94 sulfur 1.96iron 7.87 tin 7.29iron wood 1.28 - 1.37 titanium 4.51lead 11.34 tungsten 19.30magnesium 1.74 zinc 7.13

(3) Determine the mass of each object and record this on the Report Sheet.(4) From these data, calculate the volume.(5) Graph the mass as a function of volume.(6) Determine the density of the material from the slope of this graph.(7) Then, identify the material (cf. Table 2.1 for density of some common materials).

2.4. Experiment 2B. Density of irregular-shaped objects

If the object has an irregular shape, the density must be determined using Archimedes’principle, which states [1]

Any body completely or partially submerged in a fluid (gas or liquid)at rest is acted upon by an upward, or buoyant, force the magnitude ofwhich is equal to the weight of the fluid displaced by the body.

Thus, the volume of the displaced fluid is equal to the volume of a fully submerged ob-ject. In this experiment, you will determine the density of irregular shaped objects usingArchimedes’ principle.

(1) You and your laboratory partner should select two samples.(2) Note and record the identity of the sample along with a description of the sample.(3) Weigh each sample on the balance and record the weight.(4) Place approximately 30 mL of water in the 100 mL graduated cylinder. Record the

exact volume. Remember to observe the rules for significant figures in recordingyour data.

(5) Carefully drop the first of your two samples into the graduate cylinder. Recordthe new volume.

(6) Using these data, determine the density of the sample.(7) Carefully remove the water from the graduate cylinder to recover the sample. Dry

the sample and return it to the storage container. Then repeat steps 4-6 for thesecond sample.

(8) Using Table 2.1, determine the material for each sample.

61

2.5. Experiment 2C. Density of liquids

Determining the density of a liquid can be very important. For instance, the density ofaqueous acids (such as HCl) depends on the concentration of the acid dissolved in water.Thus, measuring the density is one way to check the molarity of a concentrated acid. Inthis experiment, you will determine the density of water and ethanol/water mixtures. Theprocedure for determining the density of each sample is as follows:

(1) Collect a small beaker containing ≤ 10 mL of ethanol and measure its temperature.(2) Place a clean dry 10 mL graduate cylinder on the balance and tare it to zero.

Carefully transfer approximately 2 mL of the ethanol into the cylinder, takingcare not to splash the sample on the sides of the cylinder. Record the mass.

(3) Carefully read the volume occupied by the sample at the bottom of the meniscusholding the cylinder at eye level.

(4) Calculate the density of the sample at the current temperature, noting the numberof significant figures.

(5) Add to the ethanol (Sample 1) approximately 1 mL of distilled water and mix usinga clean spatula. Record the mass and the volume of the new sample (Sample 2).

(6) Add approximately 2 mL of distilled water to Sample 2 and mix. Record the massand the volume of the new sample (Sample 3).

(7) add approximately 2 mL of distilled water to Sample 3 and mix. Record the massand volume of the new sample (Sample 4).

(8) Empty the graduated cylinder and dry.(9) Add approximately 2 mL of distilled water (Sample 5) to the 10 mL graduated

cylinder. Record the mass and the volume.Once the volume and mass for all samples has been determined, obtain an unknown

sample of a commercial alcohol from the instructor. Using the same technique above,determine the volume and mass of this unknown. Graph the density as a function ofethanol concentration (% by volume). Using this graph, determine the concentration ofalcohol in the unknown sample.


For this laboratory experiment, the laboratory project consists of the laboratory techniqueand safety score (2 pts), the Laboratory report (5 pts), the Experiment 2 Pre-laboratoryquestions (3 pts), the Experiment 2 Report Sheet (2 pts), and the Experiment 2 Post-laboratory questions (3 pts). Remember that the laboratory report must be sub-mitted as a hard copy to your instructor and as an electronic copy throughBlackboard. The laboratory report is limited to 6 pages, double-spaced, excluding tablesand figures. You should use the questions in the Report Sheet to help guide the writing inthe laboratory report.

References

1. Buoyancy, 2008. Wikipedia. http://en.wikipedia.org/wiki/Buoyancy (accessedAugust 13, 2008).

62

This page was intentional left blank.

63

Appendix I. Experiment 2: Pre-laboratory Questions

Name: Section: Grade

Partner: Date:

(1) List several examples of intrinsic properties.

(2) List several examples of extrinsic properties.

(3) When chemicals are weighed on a balance, how is the pan protected?

(4) How is the density, mass and volume of an object related?

(5) How is the volume of an irregular shaped object determined?

(6) To correctly determine the volume in a graduated cylinder, where should your eyesbe in relation to the rules on the cylinder and the meniscus of the sample?

64


65

Appendix II. Experiment 2: Report Sheet


Partner: Date:

Results

Fill out all tables and answer all questions below. All values must have appropriate signif-icant figures and units. The details for all calculations must be shown for full credit to beobtained.

Experiment 2A. Density of regular-shaped objects

Sample code:

Describe the samples.

Table A.I. The height, area, volume and mass of each object determined by myself.

Object Dimensions MassA

B

C

D

Table A.II. The height, area, volume and mass of each object determined by my labora-tory partner.

Object Dimensions MassA

B

C

D

66

Experiment 2B. Density of irregular-shaped objects

Object A

Your object code:

Describe the object.

Table B.I. The mass m of the object, the initial volume [i.e., Vi (H2O)] of water in thegraduated cylinder, the final volume [i.e., Vf (H2O)] of water in the graduate cylinder, andthe volume V of the object [i.e., V = Vf (H2O) - Vi (H2O)] determined by myself.

m:

Vi (H2O):

Vf (H2O):

V :

Object B

Object code from your laboratory partner:

Describe the object.

Table B.II. The mass m of the object, the initial volume [i.e., Vi (H2O)] of water in thegraduated cylinder, the final volume [i.e., Vf (H2O)] of water in the graduate cylinder, andthe volume V of the object [i.e., V = Vf (H2O) - Vi (H2O)] determined by my laboratorypartner.

m:

Vi (H2O):

Vf (H2O):

V :

67

Experiment 2C. Density of liquids

Table C.I. The volume VEtOH of 95% ethanol, total volume V and mass m for each of theethanol/water mixtures determined by myself.

Sample VEtOH V m1

2

3

4

5

Unknown

Table C.II. The volume VEtOH of 95% ethanol, total volume V and mass m for each of theethanol/water mixtures determined by my laboratory partner.

Sample VEtOH V m1

2

3

4

5

Unknown

68

Calculations

Experiment 2A. Density of regular-shaped objects

Table A.III. The volumes VA and VB, calculated from the dimensions given in Table A.Iand A.II, respectively, the average volume Vavg and the average mass mavg for each objectmeasured. The details for all of the calculations are given below the table.

Object VA VB Vavg mavg

A

B

C

D

69

Experiment 2A (continued)

Figure A.1. Plot the average mass mavg as a function of the average volume Vavg on thegrid provided below. (Remember to label all axes and to use the appropriate significantfigures.)

(1) Perform a linear regression on the data in Table A.III and Fig. A.1. What is thefinal regression equation?

(2) What is the density of the material, as determined from this regression equation?

(3) This density is indicative of the material: .

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Experiment 2B. Density of irregular-shaped objects

Object A

(1) From the data in Table B.I, the density was determined to be:


Object B

(1) From the data in Table B.II, the density was determined to be:


71

Experiment 2C. Density of liquids

Table C.III. The % concentration (by volume) of ethanol CA and CB and the densities ρA

and ρB for myself (A) and my laboratory partner (B). The details of the the calculationsare given below this table.

Sample CA ρA CB ρB

1

2

3

4

5

72

Experiment 2C (continued)

What is the density of the unknown sample for both measurements (yourself and yourlaboratory partner) as determined from the data in Tables C.I and C.II and the averagedensity.

Figure C.1. Graph the average densities ρx for Samples 1 - 5 as functions of the % ethanol(by volume) on the grid provided below. (Remember to label all axes and to use the appro-priate significant figures.)

(1) What is the concentration of the unknown sample, as determined from Fig. C.1?

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Appendix III. Experiment 2: Post-laboratory Questions


Partner: Date:

(1) A miner discovered some yellow nuggets. They weighted 0.0560 kg and had avolume of 2.91 mL at 20◦C. Were the nuggets gold or pyrite (otherwise known asfool’s gold)? Note: The density of gold is 19.3 g/cm3 and that of pyrite is 5.0g/cm3 at 20◦C.

(2) Explain how an alcohol thermometer works.

(3) The density of solid sand (without air spaces) is 2.84 g/mL. The density of gold is19.3 g/mL. If a 1.00 kg sample of sand containing some gold has a density of 3.10g/mL (without air spaces), what is the percentage of gold in the sample?

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(4) Fig. III.1 gives the density ρ of methanol/water mixtures as a function of %methanol (by volume). Using these data, what is the concentration of methanolin a methanol/water mixture with a density of ρ = 0.92 g/mL?

Fig. III.1. The density ρ (g/mL) of methanol/water mixtures plotted as a function of% methanol (by volume) in the mixture. The solid line represents a nonlinear 4th orderpolynomial regression.


Experiments

EXPERIMENT 3The Law of Definite Proportions

3.1. Safety

Copper sulfate pentahydrate and barium chloride dihydrate are harmful if swallowed andcan cause irritation to skin, eyes and respiratory tract. At high concentrations, thesecompound can affect the liver and kidneys. The anhydrous compounds are also harmfulif swallowed. (We should note that 1 gr is the estimated lethal dose for a human.) Thus,gloves should be worn when handling the compounds of this experiment. If any of thesecompounds come in contact with the skin or the eyes, flush with plenty of water for atleast 15 minutes. Bunsen burners, hot glassware, and metal ring stands can cause painfuland serious burns to skin. Hot glassware does not glow and, therefore, looks identical toglassware at room temperature. Thus, be careful when handling hot glassware.

3.2. Introduction

The law of definite proportions states that when two or more elements combine to forma given compound, they do so in fixed proportions by mass. This is the same as sayingthe composition of a compound is fixed. For example, sodium chloride contains 39.3%by mass sodium and 60.7% by mass chlorine. In these experiments, the law of definiteproportions will be used to determine the empirical formulas of hydrated ionic salts. (Anempirical formula expresses the simplest whole number ratio of atoms for each element ina compound.)

The previous two experiments have introduced basic laboratory techniques that will beused throughout the semester. This experiment represents the first laboratory involving ba-sic chemical principles and reactions. However, before these principles can be investigated,an understanding of chemical formulas and nomenclature must first be developed.

3.3. Chemical formulas

A chemical formula represents the composition of a given substance using the basic elemen-tal symbols. If more than one atom of an element is present in a compound, a subscriptis used after the symbol to indicate the number of atoms. Thus, the chemical compoundrepresented by the formula Fe2O3 has two atoms of iron (Fe) and 3 atoms of elementaloxygen (O). However, the mass of a single atom is difficult to measure. (For instance, themass of a single hydrogen cation (or proton) is 1.67 × 10−24 g.) Therefore, the mole isdefined as the number of 12C atoms in exactly 12 grams of 12C. Moreover, the basic unit of


75

76

Fig. 3.1: Flow chart on chemical nomenclature, adapted from [1].

mass for elemental chemistry, namely the atomic mass unit (amu or dalton) is defined as 1amu ≡ 1

12 the mass of an atom of 12C = 1.6605× 10−24 g. Thus,

1 mole = 12 g C atoms × 1 C atom19.926× 10−24 g C atom

= 6.022× 1023 C atoms .

The constant 6.022 × 1023 atoms (or molecules)/mole is known as Avogadro’s number .Since the mole and the atomic mass unit are defined using the same scale, 1 amu = 1g/mole. Thus, the masses given on the periodic table can also be expressed as the numberof grams of the element per mole of element. The molar mass M of a compound (sometimesknown as molecular weight) is obtained by summing the mass of all of the elements in acompound and, therefore, has units of g/mol. Moreover, the definition of a mole whencombined with the law of definite proportions implies that a sample of H2O will have 2moles of atomic hydrogen for every 1 mole of atomic oxygen.

3.4. Chemical nomenclature

Simple chemical compounds are named based on the classification as covalent compounds,ionic compounds, or acids (which are a special class of covalent compounds). Covalentcompounds are formed between non-metals, while ionic compounds are composed of ametal and one or more non-meals. Acids are the class of covalent compounds that resultfrom the interaction of the hydrogen cation H+ with any anion. A flow chart indicatingthe basic rules of nomenclature is given in Fig. 3.1 with a summary of the rules presentedbelow.

3.4.1. Covalent compounds

The numbers in the chemical formula are converted into words using the Greek prefixes of1 → mono 6 → hexa2 → di 7 → hepta3 → tri 8 → octa4 → tetra 9 → nona5 → penta 10 → deca

The elements are named in the same order as in the chemical formula, where elementsare listed in the order of greatest electronegativity. The suffix of the last element becomes

77

-ide. Thus, N2O5 is dinitrogen pentaoxide, while ClO2 is monochlorine dioxide or chlorinedioxide. However, Na2O contains a metal and, therefore, is not disodium oxide. The acids,which are a class of covalent compounds will be discussed later. Organic compounds (i.e.,compounds containing only C, H, O, N, and P) have a different nomenclature which will becovered in more detail in Organic Chemistry. Some simple compounds are more generallyreferred to by common names, instead of the International Union of Pure and AppliedChemistry (IUPAC) name. Table 3.1 gives the names of some simple compounds that youshould know as a student in chemistry. Common names are in italic.

3.4.2. Ionic compounds

Ionic compounds are composed of charged elements or compounds. Positively charged ionsare cations, while negatively charged ions are anions. A neutral ionic compound must havecharge balance (i.e., the net charge on the cations must equal the neat charge of the anions).Thus, the ionic compound formed by the reaction of Fe3+ and Cl− must have 3 atoms ofthe chloride anion (Cl−) for every iron(III) cation (Fe3+) in order for the compound to beneutral. This fact implies a chemical formula of FeCl3. Ionic compounds are named bynaming the cation first and then the anion.

3.4.3. Cations

(1) Cations from periodic table groups I, II and III have charges of +1, +2 and +3,respectively. Since the charge cannot vary for stable compounds with these cations,the unmodified metal name is used for the ion.

(2) Silver, zinc and cadmium only form +1 cations and, therefore, are named usingthe unmodified metal name.

(3) Mercury(I) ion is not stable and, therefore, always forms the cation Hg2+2 in ionic

compounds.(4) All other metal cations are named by adding a roman numeral in parentheses after

the metal name to indicate the charge on the ion (known as the Stock system ofnomenclature). For example, Fe3+ is iron(III), while Pb2+ is lead(II). Some of thesemetal compounds have an older nomenclature, which you will see on some of thereagent bottles in the laboratory. Table 3.2 gives the older system in comparisonto the Stock system for metals that are important in freshman chemistry.

(5) NH+4 , known as the ammonium ion, behaves as a Group I metal ion although it

does not contain a metallic element. Thus, the ammonium ion can act as a cationin an ionic compound.

Table 3.1: Names of simple compounds important in freshman chemistry.

Formula Common nameH2O dihydrogen oxide (water)NH3 azane (ammonia)CH4 methaneCH3CH3 ethaneCH3OH methanol (wood alcohol)CH3CH2OH ethanol (alcohol)PH3 phosphane (phosphine)AsH3 arsine or arsinic trihydride

78

Table 3.2: Metal cation nomenclature: The Stock system in comparison to the oldersystem.

Cation Stock system Older nameCu+ copper(I) ion cuprous ionCu2+ copper(II) ion cupric ionFe2+ iron(II) ion ferrous ionFe3+ iron(III) ion ferric ionHg2+

2 mercury(I) ion mercurous ionHg2+ mercury(II) ion mercuric ionSn2+ tin(II) ion stannous ionSn4+ tin(IV) ion stannic ion

3.4.4. Anions

(1) Monoatomic anions formed from non-metal elements with sufficient extra electronsto have a rare gas configuration have names that end in -ide. Thus O2−, which hasan electron configuration similar to the rare gas Ne, is named oxide. Similarly, P3−(having an electron configuration similar to the rare gas Ar) is named phosphide.

(2) Table 3.3 lists the common polyatomic ions that are important in freshman chem-istry. It is important to known the names and formulas of these anionsand, therefore, this table should be memorized.

(3) Polyatomic anions with names ending in -ite are related to the -ate anions inTable 3.3 but have one less oxygen atom. The standard -ite anions are also givenin Table 3.3.

(4) Anions with names having the form Per-ate are related to the -ate anions in Table3.3, but contain one additional oxygen atom. Thus, ClO−

4 is the perchlorateanion.

(5) Anions with names of the form Hypo-ite are related to the -ite anions in Table3.3, but have one less oxygen atom. Thus, ClO− is the hypochlorite anion.

(6) Anions that are formed from the combination of an anion with a charge > −1and an H+ unit are named hydrogen - ion. Thus, HS− is hydrogen sulfide ion,HPO2−

4 is hydrogen phosphate ion, and H2PO−4 is dihydrogen phosphate ion.

3.4.5. Hydrates

Hydrates are substances formed when water combines chemically in definite proportionswith an ionic salt, thereby giving a constant ratio of water molecules to the ions of the salt.Hydrates are not mixtures, since the water is coordinatively bound to either the cationor anion or both in the salt. In CuSO4 • 5 H2O, for example, the bonding involves fourwater molecules coordinatively bound to the Cu2+ ion in a square planar structure and onemolecule of water bound to the sulfate ion by hydrogen bonds [cf. Fig. 3.2]. The anhydrous(without water) form of a hydrated salt is produced when all the water of hydration is lost.Some examples of hydrates are listed below:

Formula Chemical name Common name(CaSO4)2 • H2O calcium sulfate hemihydrate plaster of parisCaSO4 • 2 H2O calcium sulfate dihydrate gypsumCuSO4 • 5 H2O copper (II) sulfate pentahydrate blue vitriolMgSO4 • 7 H2O magnesium sulfate heptahydrate epsom saltNa2CO3 • 10 H2O sodium carbonate decahydrate washing soda

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Table 3.3: The important polyatomic anions.

Formula Name Formula NameOH− hydroxide CN− cyanide

O2−2 peroxide NH−2 amide

NO−2 nitrite NO−

3 nitrate

SO2−3 sulfite SO2−

4 sulfate

PO3−3 phosphite PO3−

4 phosphate

ClO−2 chlorite ClO−

3 chlorate

CO2−3 carbonate C2H3O−

2 acetate

CrO2−4 chromate Cr2O2−

7 dichromate

MnO−4 permanganate SCN− thiocyanate

Notice that hydrates are named by first naming the ionic salt and then adding x-hydrate,where x− is the appropriate Greek prefix (cf. Section 3.4.1) to indicate the number ofwater molecules associated with salt. The • in the formula indicates a kind of chemicalbond that usually can be easily broken. For example, magnesium sulfate heptahydrate canbe converted to anhydrous magnesium sulfate by heating:

MgSO4 • 7H2O (s) → MgSO4 (s) + 7H2O (g) . (3.1)

3.4.6. Acids

Acids are covalent compounds formed from the hydrogen cation and any of the anions.Although these compounds are covalent, the hydrogen cation can be easily removed whenthe acid is in solution. Thus, acids also have some properties of ionic compounds. Becauseof this uniqueness, acids have a different nomenclature standard with the rules:

(1) Acids formed with -ide anions are named hydro-ic acid. Thus, HCl (whichcontains a chloride anion) is hydrochloric acid.

(2) Acids formed with -ate anions (including per-ates) are named -ic acid. As anexample, H2SO4 is sulfuric acid and HClO4 is perchloric acid.

(3) Acids formed with -ite anions (including hypo-ites) are named -ous acid. Thus,HNO2 is nitrous acid, while HClO is hypochlorous acid.

Fig. 3.2: Schematic picture of the CuSO4 • 5 H2O compound with the Cu2+ ion forminga square planar structure involving four water molecules, while the fifth water molecule ishydrogen bonded to the sulfate counter ion [i.e., SO2−

4 ].

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See Section 3.3.4, Item (6) for the nomenclature of acidic anions (i.e., anions formedfrom the loss of a single H+ from an acid that has more than one hydrogen cation available).

3.5. Experiment 3A. Qualitative properties of hydrates

Efflorescence is the process by which a hydrated salt loses water at room temperature andatmospheric pressures. On the other hand, the property of some salts to collect moisturefrom the air and dissolve in it is called deliquescence. A compound is hygroscopic ifabsorption of water from the atmosphere occurs without dissolution of the compound. Atthe beginning of the laboratory, your instructor placed watch glasses of calcium chloride,sodium carbonate, and lithium chloride on a watch glass in order to expose these compoundsto air. In the appropriate table on the Report Sheet describe each of these compoundsat the beginning of the laboratory and towards the end of the laboratory.

3.6. Experiment 3B. Composition of a hydrate

In this experiment, the laboratory instructor will give you and your laboratory partnertwo hydrated salts chosen from copper sulfate, barium chloride, and sodium sulfate. Thedifference in the mass of the anhydride and the hydrate will then be used to determine themass of water in the hydrate and, therefore, the empirical formula of the hydrate. Theprocedure for this study is as follows:

(1) Weigh a clean, dry, labeled crucible. Record the weight on the Report Sheet.(2) Introduce about 1 - 2 grams of the pulverized hydrated salt. Note the appearance

and color of the solid.(3) Weigh the crucible and contents. Record this weight on the Report Sheet.(4) Setup a wire triangle on the iron ring over a bunsen burner. (Ensuring that the

wire triangle will hold the crucible in an upright position.)(5) Heat the crucible and contents in the hottest part of the flame for 5 - 10 minutes.

(The bottom of the crucible should turn a dull red during heating.) Initially, thehydrate should be heated slowly by waving the burner flame fairly rapidly underthe crucible. If the material begins to boil or crackle, the heating is too intenseand splattering may occur. Within approximately 1 minute, the material shouldbecome drier and stronger heat can be applied. At the end of the 5 - 10 minuteperiod of heating, again heat slowly to allow the crucible to cool slightly beforetransfer.

(6) Using clean crucible tongs, transfer the crucible to a desiccator and allow thecrucible to cool to room temperature.

(7) When cool, weigh the dish and the anhydride and record this weight on the ReportSheet.

(8) Heat the crucible in the flame again for 5 minutes, place in desiccator and allow thecrucible to cool. Once cool, weigh the sample again. Continue the heat/cool/weighcycle until the mass of the sample remains constant. Be sure to record all of yourmeasurements on the Report Sheet.

(9) Place a thermometer in the anhydrous salt and record the temperature.(10) Add water a few drops at a time to convert the anhydride back to the hydrate.

Record the temperature (once the temperature is constant).

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Partner: Date:

(1) What should you do if any of the solid compounds come in contact with your skin?

(2) Heptahydrate implies how many water molecules are complexed to a salt?

(3) Name the following: (a) (NH4)2S, (b) H2CO3, (c) I2, and (d) AlCl3 • 6 H2O.

(4) Write the chemical formula for the following: (a) barium chloride, (b) ferroussulfate, (c) disulfur dichloride, and (d) sodium hydrogen carbonate.

(5) If 3.5 g of CuSO4 • 5 H2O (s) undergoes dehydration, what is the mass of thegaseous water produced?

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Partner: Date:

Results

Fill out all tables and answer all questions below. All values must have appropriatesignificant figures and units.

Experiment 3A. Qualitative properties of hydrates

Table A.I. The information on the reactants obtained during the performance of Experi-ment 3A.

Initial description Final description

calcium chloride

washing soda

lithium chloride

(1) Which crystals deliquesce?

(2) Which crystals effloresce?

(3) Explain why one sample gained water and why one lost water.

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Experiment 3B. Composition of a hydrate

Table B.I. The information, obtain during the performance of Experiment 3B, on the hy-drates given to me.

Hydrate 1 Hydrate 2

Name of salt in hydrate:

Formula of salt in hydrate:

Mass of evaporating dish:

Mass of dish and hydrate:

Mass of hydrate:

Mass of dish after heating*:

Mass of anhydride:

Temperature before water addition:

Temperature after water addition:

Color before heating:

Color after heating:

Color after adding water:

*The mass after heating should only be entered after the crucible and contents havereached constants mass. Use the table below to monitor the mass of the crucible during theheat/cool/weigh cycles for your hydrates. (Please note that all of the cycles listed belowmay not be necessary.)

Mass of dish after heatingHydrate 1 Hydrate 2

Cycle 1

Cycle 2

Cycle 3

Cycle 4

Cycle 5

Cycle 6

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Table B.II. The information, obtain during the performance of Experiment 3B, on thehydrates given to my laboratory partner.

Hydrate 1 Hydrate 2

Name of salt in hydrate:

Formula of salt in hydrate:

Mass of evaporating dish:

Mass of dish and hydrate:

Mass of hydrate:

Mass of dish after heating*:

Mass of anhydride:

Temperature before water addition:

Temperature after water addition:

Color before heating:

Color after heating:

Color after adding water:

*The mass after heating should only be entered after the crucible and contents havereached constants mass.

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Calculations for Experiment 3B

Use the information in Tables B.I and B.II to answer the following questions for the hy-drates studied by both you and your laboratory partner. All values must have appropriatesignificant figures and units.

(1) What evidence of a chemical change did you observe when the hydrated samplewas heated?

(2) What evidence of a chemical change did you observe when water was added to theanhydrous sample?

(3) What is the mass mH2O of the water in the hydrate (determine by subtracting themass man of the anhydride from the mass mhy of the hydrate) for (a) Hydrate 1and (b) Hydrate 2?

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(4) What is the percentage of water in the hydrate for (a) Hydrate 1 and (b) Hydrate2? The percentage of water can be determined by

%H2O =mH2O

mhy

× 100 ,

where mH2O is the mass of water in the hydrate and mhy is the mass of the hydrate.

(5) How many moles of water were in the sample from (a) Hydrate 1 and (b) Hydrate2? The number of moles nH2O of water in the sample can be obtained using

nH2O =mH2O

MH2O

,

where MH2O is the molar mass of water.

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(6) How many moles of anhydride were created in (a) Hydrate 1 and (b) Hydrate 2?The number of moles nan of the anhydride in the sample can be obtained using

nan =man

Man

,

where Man is the molar mass of the anhydride.

(7) How many moles of water are associated with a single mole of anhydride in (a)Hydrate 1 and (b) Hydrate 2? [Determine this by dividing the moles nH2O of waterby the moles nan of anhydride.]

(8) What is the formula for the hydrate [i.e., anhydride • x H2O, where x is the ratioin question (7)] that was thermally decomposed in (a) Hydrate 1 and (b) Hydrate2?

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Partner: Date:

All values must have appropriate significant figures and units.

(1) Rochelle salt is the tetrahydrate of the ionic salt KNaC4H4O6. Write the formulafor Rochelle salt.

(2) Describe (a) the three situations in which Greek prefixes are used and (b) whenRoman numerals are used in naming chemical compounds.

(3) How many moles of atoms are there in 48 g of molecular oxygen?

(4) If 5.051 g of magnesium sulfate heptahydrate is heated to remove all water, whatis the mass of the anhydride formed?

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Experiments

EXPERIMENT 4Stoichiometry of a Reaction

4.1. Safety

Copper sulfate pentahydrate is harmful if swallowed and can cause irritation to skin, eyesand respiratory tract. At high concentrations, this compound can affect the liver andkidneys. If any of this compound comes in contact with the skin or the eyes, flush withplenty of water for at least 15 minutes. Methanol is flammable and should be heated withcare. Bunsen burners, hot glassware, and metal ring stands can cause painful and seriousburns to skin. Hot glassware does not glow and, therefore, looks identical to glassware atroom temperature. Thus, be careful when handling hot glassware.

4.2. Introduction

A very common and useful type of reaction is the displacement reaction, which occurs whena metal displaces another metal in a solution with a single ionic salt. Displacement reactionsbetween metals occur when one metal is more active than another metal and usually involvethe oxidation of one metal and the reduction of the other metal. Oxidation refers to theprocess in which an atom, ion, or molecule loses electrons, while reduction implies that anatom, ion, or molecule gains electrons. A useful mnemonic device to remember this is OILRIG, or Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons). In thisexperiment, we will use stoichiometric principles to determine the oxidation state of ironions (i.e., Fe2+ or Fe3+) formed during the reaction of copper sulfate (i.e, CuSO4) withsolid iron (i.e., Fe0).

4.3. Calculations Involving Concentration

Before we can begin to study solution chemistry, an understanding of concentration mustfirst be developed. Concentration is the ratio of amount of solute to the amount of solventor solution. A solution is a homogeneous mixture of two or more molecules or ions with oneof the molecules (or ions) being the solvent and all others being the solutes. Solutes aredissolved into the solvent. In Experiment 2, you measured the density of ethanol/water so-lutions and plotted these data as a function of percent concentration by volume of ethanol.However, percent concentration is not the most useful measure of concentration when work-ing with chemical reactions. Molar concentration, or molarity M, of a solution is definedas the number of moles of the solute per liter of solution, and is the most widely used mea-sure of concentration. Other measures of concentration [e.g., molality (moles of solute/kg


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of solvent), normality (equivalents of solute/liter of solution) and parts per million (ppm)]will be introduced later. Below, we have shown two examples that illustrate how molarityis calculated and how molarity is used in solution chemistry.

4.3.1. Calculation of molarity

If 8.5 grams of ammonia (having a molar mass of 17.031 g/mole) is dissolved in 500 mL ofwater, the molarity would be determined by first calculating the number of moles of NH3

in solution and by then dividing the number of moles by the total volume of the solution.In other words,

moles NH3 = 8.5 g NH3

(1 mol NH3

17.031 g NH3

)= 0.499 mol NH3

molarity NH3 =0.499 mol NH3

500 mL

(1000 mL

1 L

)= 1.0 mol NH3/L = 1.0 M

4.3.2. Molarity in solution chemistry

If you wanted to make 750 mL (3 significant figures) of a 0.200 M aqueous solution ofsodium chloride, one would calculate the number of grams of sodium chloride (s) by thefollowing steps:

(1) Determine the number of moles of NaCl needed to prepare the solution.(2) Determine the molar mass of NaCl.(3) Multiply the number of moles of NaCl needed by the molar mass of NaCl.

These steps are illustrated below.

moles NaCl = 750 mL solution(

0.200moles NaCl1 L solution

) (1000 mL

1 L

)

= 0.1500 moles NaCl

molar mass NaCl = 23.00 g/mol of Na(

1mole Na1mole NaCl

)+ 35.45 g/mol of Cl

(1mole Cl

1mole NaCl

)

= 58.45 g/molNaCl

grams NaCl = 0.1500 moles NaCl × 58.45 g/mol of NaCl = 8.767 g NaCl= 8.77 g NaCl

4.4. Experiment 4. Reaction of copper sulfate with elemental iron

When elemental iron is oxidized (i.e., losses electrons), it can form two stable cations,namely iron(II) and iron(III). Since elemental iron is more active than a copper(II) cation,adding elemental iron to a solution containing the copper cation will result in a displacementreaction generating iron cations and elemental copper. However, two possible displacementreactions can occur, namely

Fe0(s) + Cu2+

(aq) → Fe2+(aq) + Cu0

(s) , (4.1)

or2 Fe0

(s) + 3 Cu2+(aq) → 2Fe3+

(aq) + 3Cu0(s) . (4.2)

[Notice that both eq. (4.1) and (4.2) are mass and charge balanced and, therefore, representchemical reaction equations (cf. Section 4.5).] In this experiment, we will determine which

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reaction [i.e., eq. (4.1) or (4.2)] is consistent with experiment. Therefore, we will add anexcess of copper sulfate solution to a known amount of elemental iron and will weight theproduct obtained from the chemical reaction. If eq. (4.1) is dominate, the number of molesof copper produced will equal the number of moles of iron reacted. If, however, eq. (4.2)is dominate, the number of moles of copper produced will larger than the number of molesof iron that reacted. The procedure for this study is as follows:

(1) Weigh 1.0 g of iron filings.(2) Transfer the iron filings to a clean, dry, weighed 150 mL beaker.(3) Reweigh the 150 mL beaker containing the iron filings to verify the mass of iron.(4) In Question 1 of the Pre-laboratory questions, you were asked to calculate the

amount of a 1.0 M copper(II) sulfate solution needed for the reaction given in eq.(4.2) to go to completion for a 1.0 g mass of iron. Show this calculation to yourlaboratory instructor for approval.

(5) Add the volume of copper(II) sulfate obtained in (3) plus a 5% excess to a clean,dry Erlenmeyer flask.

(6) Heat the copper sulfate solution to almost boiling.(7) Slowly add the hot copper sulfate solution to the beaker containing the iron filings.

(If the addition is performed quickly, the solution will froth and material will belost during the reaction.)

(8) When the reaction has ceased, allow the copper product to settle. Then carefullydecant the liquid from above the product.

(9) Add approximately 10 mL of distilled water to the product and swirl the beakerto mix. Again, allow the product to settle and decant. This washes the productto remove trace amounts of iron cations. Repeat with a second 10 mL portion ofdistilled water.

(10) Added 5 mL of methanol and swirl. Allow the product to settle and decant.Repeat with a second 5 mL portion of methanol.

(11) Heat the beaker in a hot water bath to remove any remaining methanol. If nec-essary, carefully break up any clumps of copper with the spatula tip. (Check thespatula to ensure that you are not removing any copper from the beaker.)

(12) When the product is thoroughly dry, dry the outside of the beaker and weigh thebeaker to determine the mass of the product.

(13) Transfer the product to a clean, dry, labeled vial for use in Experiment 5.(14) Repeat the entire procedure one more time and added this product to the labeled

vial.

4.5. Chemical reaction equations

A chemical reaction equation (or chemical equation) uses chemical formulas to describe thechemical reaction that occurs. For example, eq. (4.1) states that the reactants Fe0 andCu2+ will react to form the products Fe2+ and Cu0. The right arrow is the symbol defined tomean reacts to yield. The physical state of each substance should be given in any completechemical reaction equation, since the physical state can change the thermodynamics of thereaction. Mass and charge conservation are achieved by placing numbers in front of thechemical formulas. When both the mass and the charge are conserved (i.e., the mass andcharge of the reactants equals the mass and charge of the products), the chemical equationis said to be balanced. You should always check to verify that a chemical equationis balanced before using this equation.

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Partner: Date:

(1) What is the minimum volume needed to react all of 1.000 g of iron with a 1.0M aqueous solution of copper(II) sulfate for the reaction in eq. (4.2). Hint: Thechemical reaction equation for an aqueous solution of copper(II) sulfate is

CuSO4 (aq) → Cu2+(aq) + SO2−

4 (aq) .

(2) In eq. (4.1) which compound is (a) oxidized and (b) reduced?

(3) If copper foil is added to a (colorless) solution of silver nitrate, the solution turnsblue, while the foil turns silvery. (a) What is happening? (b) Which is the moreactive metal: silver or copper?

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(4) The combustion of a thin wire of magnesium metal in an atmosphere of pureoxygen produces the brilliant light of a flashbulb in a camera. The equation forthe reaction is

2 Mg + O2 → 2 MgO .(a) State in words how this equation is read. (b) Give the formula(s) of thereactants. (c) Give the formula(s) of the products. (d) Rewrite the equation toshow that magnesium and magnesium oxide are solids and molecular oxygen is agas.

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Partner: Date:

Results


Table B.I. The data obtain during the performance of Experiment 4 by myself.

Trial 1 Trial 2

Mass of empty 150 mL beaker:

Mass of beaker + iron:

Mass of iron filings:

Volume of 1.0 M copper(II) sulfate used:

Mass of beaker plus dry product:

Mass of product:

Table B.II. The data obtain during the performance of Experiment 4 by my laboratorypartner.

Trial 3 Trial 4

Mass of empty 150 mL beaker:

Mass of beaker + iron:

Mass of iron filings:

Volume of 1.0 M copper(II) sulfate used:

Mass of beaker plus dry product:

Mass of product:

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Calculations


(1) For each trial, calculate the number of moles of iron used. (The calculation foreach trial should be clearly indicated to receive credit.)

(2) For each trial, calculate the number of moles of copper formed.

(3) Write the balanced chemical equation that is indicative of these experimental data?

101

(4) Why was an excess of copper(II) sulfate used?

(5) What would happen if copper metal is added to iron sulfate solution? Why?

(6) Extra credit (2 pts): What is the molarity of the iron solution created duringthis reaction?

(7) Extra credit (2 pts): If the water was evaporated from the aqueous iron solutionafter the reaction, what ionic salt would be left as the residue?

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Partner: Date:


(1) Write the balanced chemical reaction equations for the following: (a) iron reactswith molecular oxygen to yield iron(III) oxide, (b) silver nitrate and calcium chlo-ride react to form calcium nitrate and silver chloride, and (c) ethane reacts withoxygen to form carbon dioxide and water.

(2) The following reaction is used to extract silver impurities from gold:Au (l) + Ag (l) + Cl2 (g) → Au (l) + AgCl (s) .

(a) Write the balanced chemical equation.

(b) How many grams of Cl2 gas would be required to remove the silver impurity in250 grams of 95% (by mass) pure gold? (Assume that silver is the only impurity.)

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Experiments

EXPERIMENT 5Copper reactions

5.1. Safety

Aqueous nitric acid and aqueous sulfuric acid are hazardous. They produce severe burnson the skin and the vapor is a lung irritant. These compounds should be handled in afume hood while wearing safety glasses and gloves. Rinse your hands with water for 5minutes after handling the acid bottles. The gases produced during these reactions aretoxic and must be avoided. Thus, when the experiment states that a procedure should beperformed in the fume hood, DO so. Aqueous sodium hydroxide is also corrosive to theskin and is especially dangerous if splashed into eyes. Again, handle this compound withgloves. Copper sulfate pentahydrate is harmful if swallowed and can cause irritation toskin, eyes and respiratory tract. At high concentrations, this compound can affect the liverand kidneys. If any of this compound comes in contact with the skin or the eyes, flush withplenty of water for at least 15 minutes. Methanol is flammable and should be heated withcare. Bunsen burners, hot glassware, and metal ring stands can cause painful and seriousburns to skin. Hot glassware does not glow and, therefore, looks identical to glassware atroom temperature. Thus, be careful when handling hot glassware.

5.2. Introduction

Chemical reactions are classified into three broad categories: (i) precipitation reactions, (ii)acid/base reactions, (iii) oxidation/reduction reactions and (iv) decomposition reactions.In precipitation reactions (or double displacement reactions) the ions of two soluble saltsreact to form an insoluble neutral compound and a different soluble salt. A good example ofthis type of reaction is the reaction of aqueous sodium chloride with aqueous silver nitrateto yield aqueous sodium nitrate and solid silver chloride, or

NaCl (aq) + AgNO3 (aq) → NaNO3 (aq) + AgCl (s) .

Acid/base reactions involve the reaction of an acid with a base to form a salt and (usually)water. These reactions will be covered in more detail during Experiments 7 and 8. Redoxreactions, such as the reaction investigated in Experiment 4, involve the transfer of electronsfrom one atom to another. Decomposition reactions occur when a compound breaks downto form simpler substances, such as when calcium carbonate decomposes into calcium oxideand carbon dioxide under high temperatures. In this experiment, various chemical reactionsinvolving copper and copper salts will be investigated.


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Table 5.1: Solubility rules for ionic compounds in aqueous solutions.

1. All compounds of the alkali metals (Group IA) are soluble.2. All salts containing ammonium, nitrate, chlorate, perchlo-

rate, and acetate ions are soluble.3. All chlorides, bromides and iodides are soluble except when

combined with silver, lead and mercury(I) cations.4. All sulfates are soluble except those combined with lead,

calcium, strontium, mercury(I), and barium.5. All metal hydroxides and all metal oxides are insoluble ex-

cept those of Group IA and those of calcium, strontium andbarium. When metal oxides dissolve, these compounds reactwith water to form hydroxides.

6. All salts that contain phosphate, carbonate, sulfite, and sul-fide are insoluble except those of Group IA and ammonium.

5.3. Solubility

The reactions that will be investigated below will rely on the precipitation of insolublesalts from aqueous solutions. The formation of a precipitate can be predicted using thesolubility rules, which are important rules that guide inorganic chemical reactions. Thus,the solubility rules in Table 5.3 should be committed to memory.

5.4. Experiment 5A. Preparation of copper oxide

In this study, you will react copper with nitric acid to form copper nitrate in solution, andthen react this product with sodium hydroxide to obtain the solid copper hydroxide. Theprocedure is as follows:

Reaction 1

(1) Place 0.5 g of the copper created in Experiment 4 into a clean, dry, weighed 250mL Beaker.

(2) Take the sample to the fume hood. Slowly add 4.0 mL of concentrated (16 M)nitric acid to the beaker inside the fume hood. Record your observations of thereaction on the Report Sheet.

(3) Once the copper has dissolved, slowly add 75 mL of deionized water. Then, returnto your laboratory bench.

Reaction 2

(1) While stirring the solution created in Reaction 1, add 30 mL of 3.0 M aqueoussodium hydroxide to create a copper precipitate.

Reaction 3

(1) Place the beaker on a wire gauze and ring on a ring stand and heat the solutionfrom Reaction 2 to almost boiling. Do not boil the solution. Stir while heatingto prevent bumping (i.e. large steam bubbles created due to non-uniform heating).

(2) When the reaction is complete, remove the beaker from heat and continue to stirfor a few minutes. Then, allow the product to settle.

(3) Decant the solution. Rinse twice with 10 mL of water and twice with 5 mL ofmethanol.

(4) Dry the product in a hot water bath, then weigh.

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5.5. Experiment 5B. Preparation of copper sulfate pentahydrate

Copper sulfate pentahydrate will be prepared by reacting copper oxide with sulfuric acid.The procedure is as follows:

(1) Place 1.0 g of copper oxide into a clean, dry 50 mL flask. (Since the copper oxidemust be dry, the copper oxide for this experiment will be supplied by the instructor.)

(2) In the fume hood, add about 20 mL of 6 M sulfuric acid to the flask.(3) Return to your laboratory bench and place the flask on a wire gauze and ring for

heating.(4) Heat the sample to almost boiling for a few minutes. Do not boil the solution.(5) If any copper(II) oxide remains, filter using a vacuum filtration apparatus, schemat-

ically shown in Fig. 5.1a. (See Section 5.6 for instructions.)(6) Transfer the filtrate to an evaporating dish and heat over a hot water bath until

the volume has been reduced by one-half. See Fig. 5.1b.(7) Stop heating and allow the solution to slowly cool. If no crystals appear, reheat

the evaporating dish to reduce the solution volume more and then allow the dishto cool again.

(8) Once crystals begin to appear, set the evaporating dish into a shallow ice waterbath to increase crystallization.

(9) Filter the crystals using vacuum filtration.(10) Rinse the crystals with 5 mL of methanol while on the vacuum filtration apparatus.(11) Leave on the vacuum filtration apparatus until the crystals have dried. Then

transfer the crystals to a sheet of weighing paper and weigh the product.

5.6. Vacuum filtration

Vacuum filtration (or suction filtration) uses a vacuum to reduce the pressure on one sideof a piece of filter paper. When a solution is poured on top of this filter paper, the vacuumacts to pull the liquid (i.e., filtrate) through the filter paper quickly, leaving the solid (i.e.,precipitate) behind. To vacuum filter a sample,

(1) Setup a vacuum filtration apparatus as shown in Fig. 5.1a.

Fig. 5.1: Schematics of (a) an vacuum filtration setup and (b) a hot water bath setup.

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(2) Place a piece of filter paper having the correct diameter on the flat area in theBuchner funnel and moisten the filter paper with the solvent in use.

(3) Check the moist filter paper for creases.(4) Make sure that the vacuum line is attached, then turn on the vacuum.(5) Slowly pour the solution to be filtered into the Buchner funnel. Try to keep the

solution in the center of the filter paper, which ensure that all of the precipitateis centered in the middle of the funnel.

(6) Once the solution has been filtered, rinse the container with a small amount of thesolvent to ensure that all of the precipitate was collected.

(7) Use solvents to help dry the precipitate as specified in the experimental procedures.At this point, if clumps are forming a spatula can be used to separate the clumps.

(8) Break the vacuum at the flask or the vacuum line.(9) Turn off the vacuum.



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Partner: Date:

Write the balanced reaction equation for the following reactions. (Remember that a bal-anced reaction equation includes the physical information about the compounds and ischarge and mass balanced.)

(1) reaction of solid copper with aqueous nitric acid and molecular oxygen to producegaseous nitrogen dioxide and copper(II) nitrate.

(2) Write the balanced reaction equation for the reaction of copper(II) nitrate withaqueous sodium hydroxide.

(3) Write the balanced reaction equation for the decomposition of copper(II) hydroxideto copper(II) oxide.

(4) Write the balanced reaction equation for the reaction of copper(II) oxide withaqueous sulfuric acid.

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Partner: Date:

Results


Experiment 5A. Preparation of copper(II) oxide

The mass of elemental copper used in this reaction is: .

Reaction 1: Preparation of copper(II) nitrate

The molarity of nitric acid was: .

The volume of nitric acid added was: .

The volume of distilled water added was: .

What type of reaction is this?

Observations (include color and texture changes):

Reaction 2: Preparation of copper(II) hydroxide

The molarity of sodium hydroxide was: .

The volume of sodium hydroxide was: .



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Reaction 3: Preparation of copper(II) oxide

This reaction took how much time?

The mass of copper(II) oxide created was: .



Experiment 5B. Preparation of copper(II) sulfate pentahydrate

The mass of copper(II) oxide used in this reaction is: .

The molarity of sulfuric acid was .

The volume of sulfuric acid was .

The mass of copper sulfate pentahydrate produced was .



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Calculations


Experiment 5A. Preparation of copper(II) oxide

(1) What is the molarity of the copper(II) nitrate solution after the addition of distilledwater in Reaction 1?

(2) How many moles of hydroxide ion (OH−) are in the 30 mL of 3.0 M aqueoussodium hydroxide solution used in Reaction 2?

(3) The limiting reactant is the reactant that controls the extent of reaction, because itis present in the smallest molar quantity. What is the limiting reactant in Reaction2?

(4) Using the moles of the limiting reactant, determine the mass of copper(II) hydrox-ide created during Reaction 2.

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(5) Using the mass of copper(II) hydroxide calculated above, determine the theoreticalyield of copper(II) oxide in Reaction 3. (The theoretical yield is the mass ofcopper(II) oxide predicted if all reactions proceed at 100% efficiency.)

(6) The percent yield of any reaction is

% yield =experimental masstheoretical mass

× 100 . (5.1)

What is the percent yield of this reaction?

(7) What steps in the reaction introduced error which, in turn, lowered the percentyield?

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Experiment 5B. Preparation of copper sulfate pentahydrate

(1) What is the percent yield for this reaction? Show all necessary calculations below.

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Partner: Date:

Describe whether the error introduced by each of the following problems would result in ahigh or low value for the preparation of copper oxide, or would not affect the results.

(1) Some of the copper nitrate solution is splashed out of the beaker before the additionof sodium hydroxide.

(2) Insufficient sodium hydroxide is added.

(3) The solution bumps during the heating of copper(II) hydroxide to produce cop-per(II) oxide.

(4) Some solid is lost in the decanting process.

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(5) The washings with water and methanol are insufficient to remove all of the solutionresidues from the copper(II) oxide.

(6) The copper(II) oxide crystals are still damp when weighed.

chemistry 113.1 introduction to chemical techniques fall...

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