Syllabus - Queens College Academic Senate - CUNY

CHEMISTRY 113.1 

INTRODUCTION TO CHEMICAL TECHNIQUES 

Fall 2008 

SECTION 1. INSTRUCTOR AND COURSE INFORMATION 

Instructor: Ms. Luxi Li 

Office: Remsen 017 

Office hours: Wednesday, 1:00 - 2:00 pm 

Telephone: (718) 997-4182 

E-mail: luxi.li@qc.cuny.edu 

Laboratory: Tuesday, 1:40 – 4:30 PM; Remsen 209 

Course 

Content: 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 of 

all points associated with that laboratory (i.e., 15 pts: prelaboratory quiz, laboratory 

project grade). If you have a university excused absence (such as illness, etc), you 

must show the excuse to the laboratory instructor the week after the absence. If you 

will miss a laboratory due to religious observance, you must inform the 

instructor 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). This 

make-up laboratory will be administered after the laboratory practical during the 14 th 

week. 

Tardiness: It is your responsibility to attend class and to be punctual. Pre-laboratory quizzes 

will be given at the beginning of class and students will be given exactly 30 minutes 

to complete the quiz. Tardiness will not result in additional time being given for prelaboratory 

quizzes.

Laboratories: There are 14 laboratories, including the check-in/introductory laboratory and the 

final examination meeting. The grade will be determined from pre-laboratory 

quizzes, laboratory techniques, laboratory reports, and laboratory practical. 

Pre-laboratory 

quizzes: 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 students 

who are tardy. These quizzes will consist of three questions, which may contain 

multiple 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. 

Laboratory 

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

be stapled together 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 

In addition to this package, an electronic version of the laboratory report should also 

be submitted to Blackboard. The completed project (i.e., Items 1-4) must be 

submitted at the beginning of the laboratory period on the date due (see the 

schedule). An electronic version (i.e., Word or WordPerfect) of the laboratory report 

must be submitted to Blackboard on the date due. Electronic versions of 

laboratory reports will be submitted to Turn-it-in software and checked for 

plagiarism. Turn-it-in checks both internet sources and previously submitted 

reports. Failure to submit an electronic version of the report will result in a zero on 

the laboratory report. 

Laboratory 

worksheets: 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 the 

laboratory 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, laboratory 

reports are limited in length to 6 pages, excluding Figures and Tables. 

Academic 

dishonesty: 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 in 

accordance with university policy. This policy can be found at: 

http://www1.cuny.edu/portal_ur/content/2004/policies/image/policy.pdf 

Electronic versions of laboratory reports will be submitted to Turn-it-in

SECTION 3. GRADING 

software and checked for plagiarism. Turn-it-in checks both internet sources 

and previously submitted reports. We should note here that copying and 

downloading graphs from the internet, without express permission of the instructor, 

constitutes plagiarism (even if referenced). An originality score $ 30% (indicating 

that # 70% of the work is original) from a single reference will result in a grade 

of zero on the laboratory report in question. 

Course 

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

Laboratory Projects (12 @ 20 pts) 240 pts 

Laboratory practical 60 pts 

Total: 420 pts 

Grading for 

Project: 1. Laboratory report * 

a. Style guide formatting (0.5 pt) 

b. Introduction and Experiment (1.5 pt) 

d. Results and Discussion (2 pts) 

7 pt 

2. Pre-laboratory questions 4 pt 

3. Report sheet 3 pt 

4. Post laboratory questions 4 pt 

Safety and laboratory technique 2 pt 

* 

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

points.

Laboratory Schedule 

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

1 Syllabus and safety. 

Safety exam. 

2 Pre-laboratory quiz 

Density and graphing. 


Chemical formulas. 


Moles. 


Chemical reactions. 


Acid/base chemistry 




Energy in chemistry. 


Metal reactivity. 


Electronic spectroscopy 


Redox Chemistry 

12 Chemical kinetics 

(60 min lecture) 

Check-in. 

Laboratory techniques. 

Finish laboratory techniques. 

Density 

Laboratory Techniques 

Project. 

Law of definite proportions. Density Project 

Stoichiometry. Law of definite 

proportions project. 

Copper reactions Stoichiometry project. 

Acid/base titrations with an 

indicator 


Finish Kinetics of bleaching experiment 

Kinetics data analysis using computers (in recitation room) 

14 Laboratory check-out and 

LABORATORY PRACTICAL (60 pts). 

Copper reaction project. 

Potentiometric analysis. Acid/base project 1. 

Heat of reactions. Acid/base project 2. 

Activity of metals. Heat of reaction project. 

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

Redox titration of bleach Beer’s Law Project 

Kinetics of bleaching Redox Titration Project 

Kinetics Project.

Laboratory Drawer Combination: 

Chemistry 113.1. 

Introduction to Chemical Techniques 

Laboratory Manual 

Cherice M. Evans, Fred H. Watson and 

Gary 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 

ii

QC Chemistry Laboratory Manual 

Version 1.0, 2008 

1.1. Attendance 

Introduction 

SECTION 1 

General Rules 

Attendance in laboratory is mandatory. An unexcused absence results in the loss of all 

points 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 the 

excuse 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 the 

course instructor (not teaching assistant) the week BEFORE the absence, or 

the absence will not be excused. If you will miss two sequential weeks due to 

religious observance, you must inform the instructor (not teaching assistant) 

that you will be missing two laboratory periods BEFORE the absences, or the 

absences will not be excused. If the absence is excused by the course instructor, the 

course instructor will make arrangements with the teaching assistants for the laboratory to 

be made-up during a separate laboratory section. 

1.2. Laboratories 

There are 14 laboratories, including the check-in/introductory laboratory and the checkout/laboratory 

practical meeting. Grades will be determined from pre-laboratory quizzes, 

laboratory technique and safety, laboratory reports, laboratory related questions, and the 

laboratory practical. 

1.3. Tardiness 

It is your responsibility to attend class and to be punctual. Pre-laboratory quizzes will be 

given at the beginning of class. Tardiness will not result in additional time being allotted 

for pre-laboratory quizzes. 

1.4. Safety 

An open book safety examination, worth 5 points, will be given during the first prelaboratory 

period. Students must score a 3/5 on this examination. Failure to do so will 

result 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 given 

at the end of Section 3. Failure to do so will result in removal from the laboratory. 

1 

c○2008 QC Chemistry and Biochemistry

2 

1.5. Pre-laboratory quizzes 

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 students who 

are tardy. These quizzes will have three parts: 

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

experiment. 

(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 stapled 

together 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 the 

beginning of the period on the date due (see the schedule in Section 2). In addition to 

this 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 electronic 

report will check for academic dishonesty using both internet sources and previously submitted 

reports. Failure to submit an electronic version of the report will result 

in a zero on the laboratory report. Instructions for submission of a report through 

Turn-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 by 

the 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 laboratory 

partner), instances of academic dishonesty will NOT be tolerated on any assignment and 

will be treated in accordance with university policy. This policy can be found at: 

http://www1.cuny.edu/portal ur/content/2004/policies/image/policy.pdf 

Electronic versions of laboratory reports will be submitted to Turn-it-in software 

and checked for academic dishonesty. Turn-it-in checks both internet 

sources and previously submitted reports. We should note here that copying and 

downloading graphs from the internet, without express permission of the instructor, constitutes 

plagiarism (even if referenced).

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 makeup 

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 via 

Blackboard 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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Version 1.0, 2008 

Introduction 

SECTION 2 

Schedule 

Date Session Pre-laboratory (60 min) Laboratory (120 min) Projects due 

1 Syllabus and safety Experiment 1: Basic 

2 Safety exam. Significant 

laboratory techniques 

Experiment 2: Density. Experiment 1 

figures and error analysis. 

Project. 

3 Pre-laboratory quiz. Experiment 3: Law of Experiment 2 

Chemical Formulas. definite proportions. Project. 

4 Pre-laboratory quiz. Experiment 4: Stoi- Experiment 3 

Formulas and reactions. chiometry. 

Project. 

5 Pre-laboratory quiz. Experiment 5: Prepara- Experiment 4 

Chemical reactions. tion of a simple salt 

6 Aqueous chemistry. Experiment 6: Reactions 

in Solution. 

7 Pre-laboratory quiz. Experiment 6: Reac- 

Aqueous chemistry. tions in Solution. 

8 Pre-laboratory quiz. Experiment 6: Reac- 

Aqueous chemistry. tions in Solution. 

9 Pre-laboratory quiz. 


Experiment 7: 

Acid/base titrations 

10 Pre-laboratory quiz. Experiment 8: Potentio- 

Acid/base chemistry. metric titrations 

11 Pre-laboratory quiz. Experiment 9: Heat of 

Heat and heat capacity. reaction. 

12 Pre-laboratory quiz. Experiment 10: Emis- 

Electron spectroscopy. sion and Beer’s Law. 

13 Pre-laboratory quiz. Experiment 11: Activity 

Metal reactivity and of metals 

14 

perfect gases 

LABORATORY PRACTICAL. Checkout. 

Make-up laboratory (examination and 

laboratory). 

5 

Project. 

Experiment 5 

Project. 

Experiment 

6 Project: 

Knowns. 

Experiment 6 

Project: 

knowns.Un- 

Experiment 

Project. 

7 

Experiment 

Project. 

8 

Experiment 9 

Project. 

Experiment 10 

Project. 

Experiment 11 

Project. 


6 



Version 1.0, 2008 

3.1. Self protection 

Introduction 

SECTION 3 

Laboratory Safety 

• Wear safety goggles at all times while in the laboratory, even if you have completed 

your experiment. Prescription safety glasses can be worn, but students must obtain 

approval from the instructor. Contact lenses, if worn in the laboratory, do not take 

the place of safety goggles. In fact, some vapors can accumulate under the lens and 

cause damage to the eyes and, therefore, prescription glasses are a better choice 

for laboratory work. 

• Bare skin must be minimized while in the laboratory by wearing clothing that 

covers one’s feet, legs and body completely. Hence, closed shoes and full-length 

pants are required, since broken glass and spilled chemicals are all too common on 

the floors of chemistry laboratories. Sandals, flip-flops, short skirts, shorts, threequarter 

length pants, bare midriffs, bare backs and bare shoulders are not allowed. 

Long-sleeve shirts are recommended but not required. Long hair should be tied 

back. Hats are not allowed. Many synthetic materials are highly flammable and 

should not be worn in the laboratory. 

• No horseplay, joking or playing is permitted in the lab. Failure to obey this rule 

is cause for immediate expulsion from the laboratory. 

• No eating or drinking is permitted in the laboratory or in the prelaboratory classroom. 

The chewing of bubble gum is considered eating in the laboratory. 

• No visitors are allowed in the laboratory. Your friends should visit with you before 

or after, but not during the laboratory session. 

• Jewelry such as rings, bracelets and watches should be removed, since chemicals 

can cause severe irritation when trapped under a piece of jewelry. 

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

removed. 

• Never taste, smell or touch a chemical or solution unless specifically directed by 

your 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. 

7 


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• Wash your hands after handling chemical and/or reagent bottles. ALWAYS wash 

your hands with soap and water before leaving the laboratory at the end of the 

period. 

• Cell phones are a distraction. Therefore, cell phones must be turned off during 

the laboratory period. If a cell phone rings, this phone will be confiscated 

by 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 fire 

blankets 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 with 

water, clean up chemical, etc.) 

• Whenever a chemical comes into contact with skin (hands, arms, etc.) flush the 

affected area with water for several minutes. Then wash thoroughly with soap and 

water. If the area of contact is the eyes or face, use the eyewash fountain to flush 

the 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 contaminated 

clothing immediately. Modesty is not an issue! Use the safety shower to wash 

area for at least 15 minutes. Follow any first aid procedures given on the Material 

Safety Data Sheet (MSDS). 

• For abrasions, cuts or minor burns, flush the area with water and consult with the 

laboratory instructor for further treatment. 

• Chemical spills on the bench or the floor should be treated using the following 

steps: 

(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 laboratory 

of the accident. 

• Know the individual hazards for all chemicals used during the laboratory experiment. 

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 the 

laboratory bench in your mouth, since these objects may have picked up contaminants 

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. Pay 

particular attention to any information concerning handling or safety of particular 

chemicals or solutions. 

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

Report unauthorized experiments to the instructor.

Fig. 3.1: NFPA label. 

• Assemble your laboratory apparatus at least 8 inches from the edge of the laboratory 

bench. 

• Maintain a clean and orderly laboratory desk and drawer. Keep drawers or cabinets 

closed and aisles free of obstructions. Do not place book bags, athletic equipment 

or other items on the floor near the laboratory bench. 

• Be aware of the actions of your neighbor as well as yourself. You could be the 

victim of a mistake made by a neighbor. Therefore, advise them of improper 

technique 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 chemical 

in an experiment will result in erroneous results in your experiment and may lead 

to a serious accident. 

• Avoid removing large excesses of reagent from the bottle. Only dispense from the 

bottle 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 information 

about dealing with laboratory fires), please see the section entitled Working 

Safely with Chemicals. 

• Dispose of chemicals using the guidelines given in the section entitled Overview of 

Hazardous Waste Disposal Procedures for Students. 

• Chemicals are often labeled according to NFPA (National Fire Protection Association) 

standard using the label shown in Fig. 3.1. This safety sticker has four 

fields, namely blue (health), red (fire), yellow (reactivity) and white (special). The 

numbers in the blocks of the stickers range from zero (0) to four (4) and indicate 

hazard severity, with zero being the least hazardous. 

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10 

3.5. Cleanliness 

It is important to keep the laboratory as clean as possible, for safety reasons as well as 

aesthetic reasons. Each pair of students is responsible for their immediate desk area. Before 

leaving the laboratory, students should make sure that the area of the laboratory bench near 

their assigned drawer is clean and dry, that Bunsen burners and other shared equipment 

are put away in the appropriate space, and that the trough to the sink is free of any solid 

material. Laboratory instructors will check work areas before approving completion of 

the experiment. Pairs of students will be assigned dates for which they are responsible 

for cleaning the reagent shelves, balances and balance tables in the weighing room, and 

surfaces under the hoods. This duty will be rotated so that each pair of students will be 

responsible 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 your 

laboratory 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.

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 all 

laboratory chemicals should be adopted. 

• It is prudent to minimize all chemical exposures. Precautions should be taken to 

avoid 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 Physical 

Hazards. 

(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 mucous 

membranes. 

(2) Physical Hazards: 

Chemicals that are either a combustible liquid, or a compressed gas, explosive, 

flammable, an organic peroxide, an oxidizer, pyrophoric, unstable (reactive), or 

water-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 working 

in the laboratory. Unrestrained long hair, loose or torn clothing, and jewelry 

can dip into chemicals or become ensnared in equipment and moving machinery. 

Clothing and hair can catch fire. Because synthetic fabrics are flammable and can 

adhere to the skin, they can increase the severity of a burn. Therefore, cotton 

is the preferred fabric. It is advisable to wear a laboratory coat. Wear shoes 

made of leather; do not wear open-toed shoes, sandals, clogs, or canvas sneakers. 

Wear disposable gloves that are appropriate for the chemical you are working 

with. 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 not 

provide adequate protection against injury. Wear safety glasses over prescription 

glasses. Contact lenses offer no protection against eye injury and cannot be substituted 

for safety glasses and goggles. It is best not to wear contact lenses when 

carrying out operations where chemical vapors are present or a chemical splash 

to the eyes or chemical dust is possible, because contact lenses can increase the 

degree of harm and can interfere with First Aid and eye-flushing procedures. If 

11

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you must wear contact lenses for medical reasons, then safety glasses with side 

shields 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 tissue 

hanging 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 all 

cleanup 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 from 

the sink for fifteen minutes of continuous flow. Always get medical help as 

soon 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 important 

consideration for this question. 

• Do NOT fight the fire if there is any possibility that you might be 

trapped 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.

Reference: 

• 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 Fire 

Do NOT run. Stop, Drop, and Roll. 

To use the safety shower, place the person on fire under the shower head, pull 

the 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 the 

person stop, drop and roll. 

(c) Fire Alarm in the Building 

Make sure all bunsen burners are off. Evacuate the laboratory in an orderly 

fashion. Close all doors behind you. Do not use the elevator. Meet at 

a designated meeting place with the rest of your class. Wait until a Fire 

Marshal or Fire Warden says it is safe to return to the building. 

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 


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 the 

incident.) 

(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 provide 

the desk officer with the same information as listed in Item 3. 

(5) When fire alarm has sounded, all occupants (faculty, staff and students) shall 

exit 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/or 

assigned Fire Wardens. Use stairways to exit the building. Never use the elevators 

unless directed to do so by Fire Department Personnel. Do not re-enter building 

until 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 possible 

about 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 provide 

the desk officer with the same information as in Item 1. 

(3) Remain at your location in order to direct emergency personnel unless it may 

jeopardize your safety. 

HAZARDOUS MATERIALS EMERGENCY RESPONSE 

This plan was designed to reduce the potential for overexposure to hazardous chemicals in 

the event of a chemical spill. Proper spill containment and cleanup procedures are ordinarily 

obtained from material safety data sheets (MSDS’s). However, you should not take it upon 

yourself to contain or clean up a chemical spill if you: 

• Are not familiar with the chemical involved and with the potential hazards associated 

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 the 

chemicals (including potential hazards and the required PPE) used for a given

laboratory procedure. If you are unable to address the clean up of a chemical spill based 

upon the criteria listed above, then 

(1) Evacuate the area. 

(2) If the spill occurred within a laboratory, notify the instructor or principal investigator. 

(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 principal 

investigator), the ESHO and the LSO are informed of the spill. 

(5) Let both your instructor (or principal investigator) and Safety Officer (or the 

Public 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 the 

scene. The room will be evacuated and the affected area will be cordoned off. This area may 

include rooms adjacent to the area where the spill has occurred. If an injury has occurred as 

a result of the spill, the person reporting the spill and/or the Safety Officer will immediately 

notify the New York City Emergency Medical Service (see Medical Emergency above). If 

occupants are trapped within the affected area, or if the ESHO, LSO or instructor/principal 

investigator cannot be reached, then the Public Safety desk will contact the New York Fire 

Department to respond to the scene. For a large chemical spill, an incident report must be 

completed and submitted to the ESHO and the LSO. Thus, the Safety Officer will attempt 

to gather as much information as possible from the individuals present at the site at the 

time of the spill. Therefore, do not leave the temporary headquarters until you have talked 

to the Safety Officer. 

15

16 

3.9. Overview of Hazardous Waste Disposal Procedures for Students 


Last modified on May 22, 2008. 

Queens College faculty, staff, students, contractors, and other parties that handle or generate 

hazardous wastes are required to properly handle, store and label hazardous wastes 

and to comply with applicable federal, state and local regulations. As a student, your 

responsibilities 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).

3.10. Laboratory Clean-up Schedule 

Session Group Students 

1 

2 

3 

4 

5 

6 

7 

8 

9 

10 

11 

12 

13 

17

18 


Laboratory Safety Agreement 

I, the undersigned, have read and understand the safety instructions given in this laboratory 

manual. I agree to abide by these rules. I also understand that failure to obey the 

safety rules given above or to follow my instructors advise while in the laboratory can lead 

to my dismissal from the laboratory for one or more class periods, with a grade of zero for 

the missed experiment(s). 

As part of the safety lecture during the first laboratory period, an overview of chemical 

disposal and general laboratory safety was discussed. In particular, the items in Section 

3.9 were reviewed. 

All laboratory experiments in this manual have been checked for safety when performed 

according to directions. I, the undersigned, understand that I am responsible for reading 

all safety precautions required for performing each experiment. Because this is a chemistry 

laboratory, I understand that there is the potential for serious accidents if these safety 

precautions are not followed and acknowledge that the fundamental responsibility for safety 

lies with myself. 

Student Name (print): QC Student ID #: 

Student Signature: Date: 

Instructor Name (print): Instructor Signature:



Version 1.0, 2008 

Introduction 

SECTION 4 

Freshman 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 calculations 

are shown on the Report Sheet for the project. Thus, this section should 

not present the details of the calculations, but should discuss these calculations in 

words. 

(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 appendices. 

4.1. Margins 

The margins should be standard paper margins, namely 

Top 1” Bottom 0.5” 

Left 1” Right 1” 

The page number should be located 0.5” from the bottom of the page and centered horizontally 

(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). 12 

point type size. Reports written in different fonts will not be accepted. 

21 


22 

4.4. Layout 

Section headings should be preceded by a roman numeral and a period, be left justified, use 

all 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 paragraph 

that is carried over to a following page) and orphans (i.e., the first line of a paragraph 

appearing 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 roman 

numerals (i.e., I, II, etc.). The table caption should be located at the top of the table. Pay 

attention to significant figures when making tables. No table should be wrapped so 

that it exists on two pages. If a table will not fit on a single page, consult your instructor 

for how to format long tables. Fig. 4.6 gives an example of a page of tables for a laboratory 

report. Notice that the table is separated from the table caption by a double line and is 

ended by a double line. Also notice that the table headings are separated from the values 

by 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. Figures 

should be numbered sequentially and consecutively from the first figure using arabic 

numerals. If a figure is composed of multiple parts, each sub-part should be identified with 

a letter of the alphabet in sequential order from top to bottom. Do not use color on figures 

unless required. The figure caption should be placed below the figure. 

If data are obtained by scanning a parameter (such as temperature or energy), then the 

data should be graphed as a line. If, however, the data are obtained by analyzing different 

objects, then the data should be graphed as markers. If data are presented as markers 

and a line is drawn through the markers, you should indicate if the line is the result of a 

regression or is drawn as a help to the reader. Fig. 4.7 gives an example of a line drawn 

as a regression, while Fig. 4.8 shows an example of a line drawn to help the reader. Note 

how the figure captions indicate each type of line. 

Figures should not have color backgrounds. The fonts for axis markers and labels should 

match the font used in writing the report and should be at least 18 points in size. Figures 

When 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. The 

references should be numbered sequentially and consecutively from the first time of use in 

the 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 inside 

punctuation. Examples are given below. 

• Journal article 

A. A. Author, B. B. Author and C. C. Author, “Article title,” Journal Abbr. Vol, 

start page (year).

• Non-scientific magazine or newspaper 

A. A. Author, “Title,” Name of magazine (date published) start page. 

• Web site 

A. A. Author (if any), Title of document, year. Title of site. http://url.of.site 

(accessed Month day, year). 

• Book 

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

• Edited volume 

Title of book, E. E. Editor, eds. (Publisher, location, year). 

• Specific chapter of an edited volume 

A. 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 each 

word in the title capitalized except for simple words (e.g., a, on, the, an, of, etc). The 

title is centered on the top of the page. Nonstandard abbreviations and acronyms are not 

allowed in the title, since these abbreviations cannot be defined in the title. If the title is 

more than a single line, then the title is not double spaced. Skip two lines, then list the 

author of the paper. The name of the author should be typed in 12 pt text and centered 

on the page. Skip one line, then listed the affiliation of the author. This affiliation should 

be centered in 12 pt text and in italic. The affiliation is as follows: 

Department of Chemistry and Biochemistry, Queens College - CUNY 

Course Number, Course Section, Semester 

Instructor: Instructor Name 

Skip two lines, then type the word “Abstract” in 12 pt type, bold, in all capitals, leftjustified. 

Skip a single line and type the abstract following the guidelines given above. Only 

the abstract should be on the title page. All fonts on the title page should be identical to 

the fonts in the body of the paper. 

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

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24 

4.9. Example laboratory report 

ABSTRACT 

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 

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. 

1 

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

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/m 3 (although it is more commonly reported in units of g/cm 3 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 = × r ¡ 2 × 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 

2 

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

25

26 

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 – 100C 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 results 

and discussion section of a laboratory report. 

3

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/cm 3 . 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 cm 3 ). 

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/cm 3 , indicating that the objects 

4 

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

27

28 

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 a 

laboratory report. 

5

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 (cm 3 ) 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 V i 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 23.4C. 

% EtOH V A (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 

6 

Fig. 4.6: Example tables for the laboratory report. 

29

30 

m (g) 

60 

50 

40 

30 

20 

10 

0 

0 

10 

20 

Fig. 1. The mass m (g) of various red plastic objects plotted as a function of the volume V (cm 3 ) 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. 

7 

30 

V (cm 3 ) 

40 

50 

60 

Fig. 4.7: Example graph for the laboratory report.

(g/mL) 

1.00 

0.95 

0.90 

0.85 

0.80 

0 

20 

40 

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. 

8 

60 

% EtOH (by volume) 

80 

100 

Fig. 4.8: Example graph for the laboratory report. 

31

32 



Version 1.0, 2008 

Introduction 

SECTION 5 

Instructions for Turn-it-in Assignments 

All laboratory reports must be submitted both in hard copy to the instructor and as 

an electronic version to Blackboard. The electronic version will be checked for instances of 

academic 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 file 

formats that are acceptable, only Word or Wordperfect files are valid 

submissions for this course. 

(8) Select the file to be uploaded and then click okay. This will return you to the 

Blackboard 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 113 

site. 

33 


34 

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 Chemistry 

113 site.

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 laboratory 

report. 

35

36 



Version 1.0, 2008 

6.1. Experiment 1. Basic Laboratory Technique 

Thermometer calibration line: 

37 

Introduction 

SECTION 6 

Useful information 


38 



Version 1.0, 2008 

1.1. Safety 

Experiments 

EXPERIMENT 1 

Check-in and Basic Laboratory Techniques 

During Check-in, you may encounter broken or chipped glassware. This glassware should 

be handled with care to prevent injury. Broken and/or chipped glassware should be placed 

in the broken glassware box. Bunsen burners, hot glassware, metal ring stands, and boiling 

water 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 handling 

hot glassware. 

1.2. Check-in 

You will be assigned a laboratory bench and drawer containing your laboratory kit. You 

will be responsible for all items in this kit. Check-in allows you to confirm that your kit has 

all required glassware and to replace glassware that is broken or too dirty to clean. This is 

the only day when missing or broken items are replaced free of charge. Thus, you should 

report all items that are missing, scratched, corroded or otherwise unfit for use. You should 

also use this time to clean all of the glassware in the kit, since this will save time in the 

future. 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 the 

laboratory kit are still in good shape and clean. Illustrations of the items that should be 

in the laboratory kit are given in Figs. 1.1 and 1.2. General rules about the laboratory kit 

are: 

• This kit is your responsibility. Missing items are replaced free of charge only on 

the 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 checkedout 

will be charged a $50 fee in addition to the fee from missing, broken 

and dirty items. 

39 


40 

Fig. 1.1: Glassware located in the laboratory kit.

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

41

42 

These experiments are meant to provide your first experience in the laboratory. If you 

have prior laboratory experience, these exercises constitute a good review. Your instructor 

will give you the rotation order for the experiment. Submit all results on the Worksheet 

provided. 

1.3. Experiment 1A. Check-in 

Using the Desk Assignment Sheet, check your laboratory kit to ensure that all glassware 

and equipment is present and in good working order. Wash and dry all glassware during 

this time. If glassware is too dirty to be cleaned, return this glassware to the stockroom to 

be replaced. 

1.4. Weighing 

You must learn how to use laboratory balances. As always, the limit of readable precision 

of 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 container 

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 balances 

have an automatic range selector that will change the readout precision automatically to 

0.01 g when the gross mass on the pan is over 35 g. The measurement precision for small 

masses is best if very light containers are used, such as the glassine weighing paper for dry 

solid 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 the 

appropriate location on the Experiment 1 Report Sheet. 

(2) Without taring, place a sheet of paper on the pan. Read and record its mass in 

the table on the Report Sheet. When possible, use the cover to protect the pan 

from air drafts to obtain higher precision. 

(3) Add the pennies to the pan one at a time, reading the mass after each addition. 

Enter each in the table given on the Report Sheet in the column entitled“Cumulative 

Mass.” Keep the pennies in order. Calculate the mass of each 

penny by subtracting; record the differences in the column entitled “Mass by Difference.” 

(4) Remove all the coins and weigh each individually, taring to zero. Record the mass 

of each coin in the table in the column entitled “Direct Weighing.” 

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

m = 1 

m , (1.1) 

n 

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 = 1 

n 

d . (1.2)

Consider two hypotheses: 

Fig. 1.3: Bunsen burner connected to gas valve. 

(1) All Lincoln-head pennies are manufactured with equal mass (within 0.001 g), but 

their 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 problems 

in chemistry are illustrated here. Most obviously, the state of corrosion of the pennies 

represents an uncontrolled experimental variable which can be important. Had we used 

non-circulated coins, we would have expected better precision. On the other hand, pennies 

may not be very uniform even when new. Another question arises: How much experimental 

difference is sufficient and how consistently must it be observed for us to consider two data 

sets, or groups of data sets as distinctly different? This is an important question for which 

statistical 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 its 

length on the inch scale and the centimeter scale. Place your results, along with those of 

your laboratory partner, in the appropriate table on the Report Sheet. Pay attention to 

the precision of the ruler during these measurements. What is the deviation and average 

deviation? 

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 significant 

figures in your results. 

1.7. Bunsen Burner 

When selecting a burner, check to see that the gas needle valve on the bottom will close 

completely. Also check to see that the barrel of the burner will screw in and out so that 

the air supply to the flame can be controlled. 

With the burner gas valve off and the hose connected to the burner and the bench gas 

cock (see Fig. 1.3), turn the bench gas cock fully on and check for leaks around the burner 

with a match. With the air vents closed, open the burner gas valve and light the flame. 

The flame should be yellow and luminous. 

43

44 

Stick a test tube into the flame briefly. You should observe 

a deposit of carbon black on the tube. When hydrocarbons such 

as methane [CH4 (natural gas)] burn in too little air (oxygen), the 

reaction is 

CH4 (gas) + O2 (gas) → C (sol) + 2 H2O (liq) . 

With a bit more air, the flame becomes hotter and blue, but carbon 

monoxide 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 to 

the right. This is the hottest flame and is characterized by a blue 

inverted cone shape within the flame that is the so-called reducing 

flame. A little above the apex of the cone is the hottest area in the 

flame, reaching temperatures around 1500 ◦ C. 

Towards the top of the flame, conditions are oxidizing (high temperatures, 

excess O2). The well-adjusted flame completely converts 

methane 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 the 

close the bench gas cock. 

1.8. Experiment 1D. Understanding Flames 

Schematic of a 

bunsen burner 

flame when the 

flame is the 

hottest. 

Take a wooden splint and hold it with one end resting on the top of the burner. Notice how 

the splint is burned only on the edges of the flame. The flame under the cone is relatively 

cool (about 350 ◦ C). Place the other end of the splint higher in the flame. Notice that now 

the 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. Insert 

the 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 bunsen 

burner (cf. Fig. 1.4). When the water is boiling, immerse the thermometer and read the 

temperature. (Do not let the thermometer bulb touch the bottom of the beaker during this 

procedure.) Once the temperature is constant, record the temperature as the boiling point 

of water. 

If the results that you obtained from this procedure give a freezing point and boiling 

point of water that differs from the theoretical temperatures by the same amount, then 

this 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 both 

the freezing point and boiling point, then plot the observed values versus the theoretical 

values. The line between these points will give a calibration line that can be used to correct 

temperatures.

Fig. 1.4: Hot water bath schematic. 

1.10. Experiment 1F. Measuring volumes 

Graduated cylinders are used for measuring volumes 

of a liquid. In your laboratory kit, you have 

two graduated cylinders: a 10 mL graduated cylinder 

and a 100 mL graduated cylinder. When a 

liquid is placed in a graduated cylinder, the liquid 

level in the cylinder will curve with the lowest (or 

highest) point being in the middle. This point is 

the meniscus and represents the point where one 

should determine the volume. To avoid errors in 

determining the volume, the line of sight must be 

perpendicular to the scale (cf. Fig. 1.5. 

Obtain a 250 mL graduated Erlenmeyer flask 

and fill with water to the 50 mL mark. Transfer 

the water, completely and without spilling, to the 

100 mL graduated cylinder. Record the volume to 

the correct precision. Similarly, obtain a 125 mL Fig. 1.5: Reading a graduated cylin- 

graduate Erlenmeyer flask and fill with water to the der. 

50 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 to 

the 100 mL graduated cylinder. Record the volume to the correct precision. Similarly, fill 

a 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 to 

the 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 technique 

and safety score (2 pts), the Experiment 1 Report Sheet (10 pts) and the Experiment 1 

Post-laboratory questions (3 pts). Pre-laboratory questions and a laboratory report are 

not required to complete this laboratory project. 

45

46 


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: 

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48 


Experiment 1. Report Sheet 

Name: Section: 

Partner: Date: 

Results 

Fill out all tables and answer all questions below. All values must have appropriate 

significant 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 collection 

of pennies on the balance, and the mass of each penny (determined by directly 

weighing 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: 

49

50 

Experiment 1C. Length 

Table C.I. The length of a wood splint provided by the instructor measured in inches and 

in 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 (Measurement 

A) and my laboratory partner (Measurement B). 

Observed 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 

Measurement A Measurement B

Experiment 1F. Measuring volumes 

Table F.I. The volume of water as determined using the graduations in the glassware (Row 

1) and the graduated cylinder (Row 2). The error in volume (i.e., |Row 1 - Row 2|) should 

be given in Row 3. 

Volume (mL) 

in glassware 

Volume (mL) 

using graduated 

cylinder 

Error in volume 

of glassware 

125 mL Erlenmeyer 

flask 

250 mL Erlenmeyer 

flask 

50 mL Beaker 250 mL Beaker 

51

52 

Calculations 


significant figures and units. To receive full credit all details for all calculations must 

be shown. 

Experiment 1B. Mass 

Table B.II. The mass of each penny determined by the difference using the cumulative mass 

in Table B.I and the mass of each penny determined by direct weighing. From these two 

measurements, 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 

1st penny 

2st penny 

3st penny 

4th penny 

5th penny 

Mass by 

Difference 

Mass by Direct 

Weighing 

Average Mass 

Deviation from 

Mean 

Mean NAN

Experiment 1C. Length 

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

centimeters? 

(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 the 

test 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 splint 

burn first? Why? 

53

54 

Experiment 1E. Thermometer calibration 

Plot the average for the two temperatures recorded with your thermometer in ◦ C as a 

function of the theoretical values on the graph below. Be sure to label all axis of the graph 

and to use the appropriate number of significant figures when labeling the graph. Use this 

graph 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.

(1) Ethanol has a boiling point of 78.4 ◦ C and a melting point of -114.3 ◦ C. If your 

thermometer was placed into boiling ethanol, what temperature would it read? If 

the ethanol was freezing, what temperature would your thermometer give? 

(2) Fahrenheit and Celsius are related by 

t( ◦ 

9◦F F) = 

5◦ 

t( 

C 

◦ C) + 32 ◦ F , 

while the relationship between Celsius and Kelvin is 

T (K) = (t( ◦ C) + 273.15 ◦ 

1K 

C) 

1◦ 

. 

C 

Using these relationships, convert the boiling point and freezing point of ethanol 

to Fahrenheit and Kelvin. Be sure to report your answers with the correct units 

and to the correct number of significant figures. 

55

56 

Experiment 1F. Measuring volumes 

(1) Percent error in this measurement is given by 

% Error = |MA − ML| 

× 100 , 

MA 

where MA is the more accurate measurement and ML is the less accurate measurement. 

Determine the percent error in the graduation of the Erlenmeyer flasks 

and the beakers.

Experiment 1. Post-laboratory questions 

Name: Section: 


(1) A student used a graduated cylinder having volume markings every 1 mL to carefully 

measure 100 mL of water for an experiment. A fellow student said that by 

reporting the volume as “100” mL in her lab notebook, she was only entitled to 

one significant figure. She disagreed. Why did her fellow student say the reported 

volume had only one significant figure? Considering the circumstances, how many 

significant 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 a 

temperature 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? 

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Version 1.0, 2008 

2.1. Safety 

Experiments 

EXPERIMENT 2 

Density 

Certain metals, such as lead, can be toxic when ingested. Wash your hands after handling 

all samples. Laboratory ethanol has trace impurities of methanol and toluene. Therefore, 

laboratory ethanol should not be ingested. Ethanol and ethanol/water mixtures are also 

flammable 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 

. (2.1) 

V 

Mass and volume are extensive properties of matter – properties that depend on the 

quantities of substances. Such properties are not of themselves useful in characterizing substances. 

Intensive properties, on the other hand, are useful in characterizing substances. 

Intensive properties are often determined by ratioing two extensive properties measured at 

constant temperature T and pressure P . Density is an example of this kind of intensive 

property. When measured under known conditions of T and P , density can be used to 

characterize 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 you 

determine that a colorless liquid has a density of 1.00 g/mL at 4◦C and 1 atm, this does 

not prove the liquid is water. This fact is simply one piece of evidence that the substance 

may be water. 

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

For objects with regular shapes, the volume can be easily determined by measuring the 

dimensions 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 in 

the appropriate spot of the Report Sheet. 

(2) Determine the dimensions of the objects using a ruler. This step should be performed 

by both you and your laboratory partner. Be sure to record the measurements 

to the correct precision. Place this in the appropriate table on the Report 

Sheet. 

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 - 0.83 

aluminum 2.71 molybdenum 10.22 

antimony 6.62 polyamide (nylon) 1.15 

balsa wood 0.10 - 0.20 polypropylene 0.90 

bismuth 9.80 polytetrafluoroethylene (teflon) 2.2 

cadmium 8.65 polyurethane 1.23 

cedar 0.38 polyvinylchloride 1.38 

copper 8.94 sulfur 1.96 

iron 7.87 tin 7.29 

iron wood 1.28 - 1.37 titanium 4.51 

lead 11.34 tungsten 19.30 

magnesium 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 of 

which 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 object. 

In this experiment, you will determine the density of irregular shaped objects using 

Archimedes’ 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 recording 

your data. 

(5) Carefully drop the first of your two samples into the graduate cylinder. Record 

the 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 the 

second sample. 

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

2.5. Experiment 2C. Density of liquids 

Determining the density of a liquid can be very important. For instance, the density of 

aqueous 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. In 

this experiment, you will determine the density of water and ethanol/water mixtures. The 

procedure 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, taking 

care 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 meniscus 

holding the cylinder at eye level. 

(4) Calculate the density of the sample at the current temperature, noting the number 

of significant figures. 

(5) Add to the ethanol (Sample 1) approximately 1 mL of distilled water and mix using 

a 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 mass 

and the volume of the new sample (Sample 3). 

(7) add approximately 2 mL of distilled water to Sample 3 and mix. Record the mass 

and 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 of 

ethanol concentration (% by volume). Using this graph, determine the concentration of 

alcohol in the unknown sample. 



and safety score (2 pts), the Laboratory report (5 pts), the Experiment 2 Pre-laboratory 

questions (3 pts), the Experiment 2 Report Sheet (2 pts), and the Experiment 2 Postlaboratory 

questions (3 pts). Remember that the laboratory report must be submitted 

as a hard copy to your instructor and as an electronic copy through 

Blackboard. The laboratory report is limited to 6 pages, double-spaced, excluding tables 

and figures. You should use the questions in the Report Sheet to help guide the writing in 


References 

1. Buoyancy, 2008. Wikipedia. http://en.wikipedia.org/wiki/Buoyancy (accessed 

August 13, 2008). 

61

62 

This page was intentional left blank.

Appendix I. Experiment 2: Pre-laboratory Questions 

Name: Section: Grade 


(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 eyes 

be in relation to the rules on the cylinder and the meniscus of the sample? 

63

64 


Results 

Appendix II. Experiment 2: Report Sheet 



Fill out all tables and answer all questions below. All values must have appropriate significant 

figures and units. The details for all calculations must be shown for full credit to be 

obtained. 

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 Mass 

A 

B 

C 

D 

Table A.II. The height, area, volume and mass of each object determined by my laboratory 

partner. 

Object Dimensions Mass 

A 

B 

C 

D 

65

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 the 

graduated cylinder, the final volume [i.e., Vf (H2O)] of water in the graduate cylinder, and 

the volume V of the object [i.e., V = Vf (H2O) - Vi (H2O)] determined by myself. 

Object B 

m: 

Vi (H2O): 

Vf (H2O): 

V : 

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 the 

graduated cylinder, the final volume [i.e., Vf (H2O)] of water in the graduate cylinder, and 

the volume V of the object [i.e., V = Vf (H2O) - Vi (H2O)] determined by my laboratory 

partner. 

m: 

Vi (H2O): 

Vf (H2O): 

V :

Experiment 2C. Density of liquids 

Table C.I. The volume VEtOH of 95% ethanol, total volume V and mass m for each of the 

ethanol/water mixtures determined by myself. 

Sample VEtOH V m 

1 

2 

3 

4 

5 

Unknown 

Table C.II. The volume VEtOH of 95% ethanol, total volume V and mass m for each of the 

ethanol/water mixtures determined by my laboratory partner. 

Sample VEtOH V m 

1 

2 

3 

4 

5 

Unknown 

67

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.I 

and A.II, respectively, the average volume Vavg and the average mass mavg for each object 

measured. The details for all of the calculations are given below the table. 

Object VA VB Vavg mavg 

A 

B 

C 

D

Experiment 2A (continued) 

Figure A.1. Plot the average mass mavg as a function of the average volume Vavg on the 

grid provided below. (Remember to label all axes and to use the appropriate significant 

figures.) 

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

final regression equation? 

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

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

69

70 

Experiment 2B. Density of irregular-shaped objects 

Object A 

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

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

Object B 

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

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

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 calculations 

are given below this table. 

Sample CA ρA CB ρB 

1 

2 

3 

4 

5 

71

72 

Experiment 2C (continued) 

What is the density of the unknown sample for both measurements (yourself and your 

laboratory partner) as determined from the data in Tables C.I and C.II and the average 

density. 

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 appropriate 

significant figures.) 

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

Appendix III. Experiment 2: Post-laboratory Questions 



(1) A miner discovered some yellow nuggets. They weighted 0.0560 kg and had a 

volume of 2.91 mL at 20 ◦ C. Were the nuggets gold or pyrite (otherwise known as 

fool’s gold)? Note: The density of gold is 19.3 g/cm 3 and that of pyrite is 5.0 

g/cm 3 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 is 

19.3 g/mL. If a 1.00 kg sample of sand containing some gold has a density of 3.10 

g/mL (without air spaces), what is the percentage of gold in the sample? 

73

74 

(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 methanol 

in 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 order 

polynomial regression.


Version 1.0, 2008 

3.1. Safety 

Experiments 

EXPERIMENT 3 

The Law of Definite Proportions 

Copper sulfate pentahydrate and barium chloride dihydrate are harmful if swallowed and 

can cause irritation to skin, eyes and respiratory tract. At high concentrations, these 

compound can affect the liver and kidneys. The anhydrous compounds are also harmful 

if 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 these 

compounds come in contact with the skin or the eyes, flush with plenty of water for at 

least 15 minutes. Bunsen burners, hot glassware, and metal ring stands 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 handling hot glassware. 


The law of definite proportions states that when two or more elements combine to form 

a given compound, they do so in fixed proportions by mass. This is the same as saying 

the 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 definite 

proportions will be used to determine the empirical formulas of hydrated ionic salts. (An 

empirical formula expresses the simplest whole number ratio of atoms for each element in 

a compound.) 

The previous two experiments have introduced basic laboratory techniques that will be 

used throughout the semester. This experiment represents the first laboratory involving basic 

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 elemental 

symbols. If more than one atom of an element is present in a compound, a subscript 

is used after the symbol to indicate the number of atoms. Thus, the chemical compound 

represented by the formula Fe2O3 has two atoms of iron (Fe) and 3 atoms of elemental 

oxygen (O). However, the mass of a single atom is difficult to measure. (For instance, the 

mass of a single hydrogen cation (or proton) is 1.67 × 10 −24 g.) Therefore, the mole is 

defined as the number of 12 C atoms in exactly 12 grams of 12 C. 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 1 

amu ≡ 1 

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

1 mole = 12 g C atoms × 

1 C atom 

19.926 × 10 −24 g C atom = 6.022 × 1023 C atoms . 

The constant 6.022 × 10 23 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 = 1 

g/mole. Thus, the masses given on the periodic table can also be expressed as the number 

of grams of the element per mole of element. The molar mass M of a compound (sometimes 

known as molecular weight) is obtained by summing the mass of all of the elements in a 

compound and, therefore, has units of g/mol. Moreover, the definition of a mole when 

combined with the law of definite proportions implies that a sample of H2O will have 2 

moles 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). Covalent 

compounds are formed between non-metals, while ionic compounds are composed of a 

metal and one or more non-meals. Acids are the class of covalent compounds that result 

from the interaction of the hydrogen cation H + with any anion. A flow chart indicating 

the basic rules of nomenclature is given in Fig. 3.1 with a summary of the rules presented 

below. 

3.4.1. Covalent compounds 

The numbers in the chemical formula are converted into words using the Greek prefixes of 

1 → mono 6 → hexa 

2 → di 7 → hepta 

3 → tri 8 → octa 

4 → tetra 9 → nona 

5 → penta 10 → deca 

The elements are named in the same order as in the chemical formula, where elements 

are listed in the order of greatest electronegativity. The suffix of the last element becomes

-ide. Thus, N2O5 is dinitrogen pentaoxide, while ClO2 is monochlorine dioxide or chlorine 

dioxide. 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 be 

covered in more detail in Organic Chemistry. Some simple compounds are more generally 

referred to by common names, instead of the International Union of Pure and Applied 

Chemistry (IUPAC) name. Table 3.1 gives the names of some simple compounds that you 

should 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 ions 

are cations, while negatively charged ions are anions. A neutral ionic compound must have 

charge 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 Fe 3+ and Cl − must have 3 atoms of 

the chloride anion (Cl − ) for every iron(III) cation (Fe 3+ ) in order for the compound to be 

neutral. This fact implies a chemical formula of FeCl3. Ionic compounds are named by 

naming 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 using 

the unmodified metal name. 

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

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 of 

nomenclature). For example, Fe 3+ is iron(III), while Pb 2+ is lead(II). Some of these 

metal compounds have an older nomenclature, which you will see on some of the 

reagent bottles in the laboratory. Table 3.2 gives the older system in comparison 

to the Stock system for metals that are important in freshman chemistry. 

, 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 cation 

in an ionic compound. 

(5) NH + 4 

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

Formula Common name 

H2O dihydrogen oxide (water) 

NH3 

azane (ammonia) 

CH4 

methane 

CH3CH3 

ethane 

CH3OH methanol (wood alcohol) 

CH3CH2OH ethanol (alcohol) 

PH3 

phosphane (phosphine) 

AsH3 

arsine or arsinic trihydride 

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

system. 

3.4.4. Anions 

Cation Stock system Older name 

Cu + copper(I) ion cuprous ion 

Cu 2+ copper(II) ion cupric ion 

Fe 2+ iron(II) ion ferrous ion 

Fe 3+ iron(III) ion ferric ion 

Hg 2+ 

2 mercury(I) ion mercurous ion 

Hg 2+ mercury(II) ion mercuric ion 

Sn 2+ tin(II) ion stannous ion 

Sn 4+ tin(IV) ion stannic ion 

(1) Monoatomic anions formed from non-metal elements with sufficient extra electrons 

to have a rare gas configuration have names that end in -ide. Thus O2− , which has 

an 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 chemistry. 

It is important to known the names and formulas of these anions 

and, therefore, this table should be memorized. 

(3) Polyatomic anions with names ending in -ite are related to the -ate anions in 

Table 3.3 but have one less oxygen atom. The standard -ite anions are also given 

in Table 3.3. 

(4) Anions with names having the form Per-ate are related to the -ate anions in Table 

3.3, but contain one additional oxygen atom. Thus, ClO − 4 

is the perchlorate 

anion. 

(5) Anions with names of the form Hypo-ite are related to the -ite anions in Table 

3.3, 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 > −1 

and an H + unit are named hydrogen - ion. Thus, HS− is hydrogen sulfide ion, 

is dihydrogen phosphate ion. 

HPO 2− 

4 is hydrogen phosphate ion, and H2PO − 4 

3.4.5. Hydrates 

Hydrates are substances formed when water combines chemically in definite proportions 

with 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 cation 

or anion or both in the salt. In CuSO4 • 5 H2O, for example, the bonding involves four 

water molecules coordinatively bound to the Cu 2+ ion in a square planar structure and one 

molecule 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 paris 

CaSO4 • 2 H2O calcium sulfate dihydrate gypsum 

CuSO4 • 5 H2O copper (II) sulfate pentahydrate blue vitriol 

MgSO4 • 7 H2O magnesium sulfate heptahydrate epsom salt 

Na2CO3 • 10 H2O sodium carbonate decahydrate washing soda

Table 3.3: The important polyatomic anions. 

Formula Name Formula Name 

OH − hydroxide CN − cyanide 

O 2− 

2 peroxide NH − 2 amide 

NO − 2 nitrite NO − 3 nitrate 

SO 2− 

3 sulfite SO 2− 

4 

PO 3− 

3 phosphite PO 3− 

4 

sulfate 

phosphate 

ClO − 2 chlorite ClO − 3 chlorate 

CO 2− 

3 carbonate C2H3O − 2 acetate 

CrO 2− 

4 chromate Cr2O 2− 

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 of 

water molecules associated with salt. The • in the formula indicates a kind of chemical 

bond that usually can be easily broken. For example, magnesium sulfate heptahydrate can 

be converted to anhydrous magnesium sulfate by heating: 

3.4.6. Acids 

MgSO 4 • 7H2O (s) → MgSO 4 (s) + 7H2O (g) . (3.1) 

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 when 

the acid is in solution. Thus, acids also have some properties of ionic compounds. Because 

of this uniqueness, acids have a different nomenclature standard with the rules: 

(1) Acids formed with -ide anions are named hydro-ic acid. Thus, HCl (which 

contains a chloride anion) is hydrochloric acid. 

(2) Acids formed with -ate anions (including per-ates) are named -ic acid. As an 

example, 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 Cu 2+ ion forming 

a square planar structure involving four water molecules, while the fifth water molecule is 

hydrogen bonded to the sulfate counter ion [i.e., SO 2− 

4 ]. 

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80 

See Section 3.3.4, Item (6) for the nomenclature of acidic anions (i.e., anions formed 

from 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 and 

atmospheric pressures. On the other hand, the property of some salts to collect moisture 

from the air and dissolve in it is called deliquescence. A compound is hygroscopic if 

absorption of water from the atmosphere occurs without dissolution of the compound. At 

the 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 compounds 

to air. In the appropriate table on the Report Sheet describe each of these compounds 

at 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 partner 

two hydrated salts chosen from copper sulfate, barium chloride, and sodium sulfate. The 

difference in the mass of the anhydride and the hydrate will then be used to determine the 

mass of water in the hydrate and, therefore, the empirical formula of the hydrate. The 

procedure 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, the 

hydrate should be heated slowly by waving the burner flame fairly rapidly under 

the crucible. If the material begins to boil or crackle, the heating is too intense 

and splattering may occur. Within approximately 1 minute, the material should 

become drier and stronger heat can be applied. At the end of the 5 - 10 minute 

period of heating, again heat slowly to allow the crucible to cool slightly before 

transfer. 

(6) Using clean crucible tongs, transfer the crucible to a desiccator and allow the 

crucible to cool to room temperature. 

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

Sheet. 

(8) Heat the crucible in the flame again for 5 minutes, place in desiccator and allow the 

crucible to cool. Once cool, weigh the sample again. Continue the heat/cool/weigh 

cycle until the mass of the sample remains constant. Be sure to record all of your 

measurements 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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(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) ferrous 

sulfate, (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 the 

gaseous water produced? 

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Results 





significant figures and units. 

Experiment 3A. Qualitative properties of hydrates 

Table A.I. The information on the reactants obtained during the performance of Experiment 

3A. 

calcium chloride 

washing soda 

lithium chloride 

(1) Which crystals deliquesce? 

(2) Which crystals effloresce? 

Initial description Final description 

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

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86 

Experiment 3B. Composition of a hydrate 

Table B.I. The information, obtain during the performance of Experiment 3B, on the hydrates 

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 have 

reached constants mass. Use the table below to monitor the mass of the crucible during the 

heat/cool/weigh cycles for your hydrates. (Please note that all of the cycles listed below 

may not be necessary.) 

Mass of dish after heating 


Cycle 1 

Cycle 2 

Cycle 3 

Cycle 4 

Cycle 5 

Cycle 6

Table B.II. The information, obtain during the performance of Experiment 3B, on the 

hydrates given to my laboratory partner. 


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 have 

reached 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 hydrates 

studied by both you and your laboratory partner. All values must have appropriate 


(1) What evidence of a chemical change did you observe when the hydrated sample 

was heated? 

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

anhydrous sample? 

(3) What is the mass m H2 O of the water in the hydrate (determine by subtracting the 

mass m an of the anhydride from the mass m hy of the hydrate) for (a) Hydrate 1 

and (b) Hydrate 2?

(4) What is the percentage of water in the hydrate for (a) Hydrate 1 and (b) Hydrate 

2? The percentage of water can be determined by 

%H2O = mH2O × 100 , 

mhy 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) Hydrate 

2? The number of moles n H2 O of water in the sample can be obtained using 

n H2 O = m H 2 O 

M H2 O 

where M H2 O is the molar mass of water. 

, 

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90 

(6) How many moles of anhydride were created in (a) Hydrate 1 and (b) Hydrate 2? 

The number of moles n an of the anhydride in the sample can be obtained using 

n an = m an 

M an 

where M an 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 n H2 O of water 

by the moles n an of anhydride.] 

(8) What is the formula for the hydrate [i.e., anhydride • x H2O, where x is the ratio 

in question (7)] that was thermally decomposed in (a) Hydrate 1 and (b) Hydrate 

2? 

,




All values must have appropriate significant figures and units. 

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

for Rochelle salt. 

(2) Describe (a) the three situations in which Greek prefixes are used and (b) when 

Roman 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, what 

is the mass of the anhydride formed? 

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Version 1.0, 2008 

4.1. Safety 

Experiments 

EXPERIMENT 4 

Stoichiometry of a Reaction 

Copper sulfate pentahydrate is harmful if swallowed and can cause irritation to skin, eyes 

and respiratory tract. At high concentrations, this compound can affect the liver and 

kidneys. If any of this compound comes in contact with the skin or the eyes, flush with 

plenty of water for at least 15 minutes. Methanol is flammable and should be heated with 

care. Bunsen burners, hot glassware, and metal ring stands 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 handling hot glassware. 


A very common and useful type of reaction is the displacement reaction, which occurs when 

a metal displaces another metal in a solution with a single ionic salt. Displacement reactions 

between metals occur when one metal is more active than another metal and usually involve 

the oxidation of one metal and the reduction of the other metal. Oxidation refers to the 

process in which an atom, ion, or molecule loses electrons, while reduction implies that an 

atom, ion, or molecule gains electrons. A useful mnemonic device to remember this is OIL 

RIG, or Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons). In this 

experiment, we will use stoichiometric principles to determine the oxidation state of iron 

ions (i.e., Fe 2+ or Fe 3+ ) formed during the reaction of copper sulfate (i.e, CuSO4) with 

solid iron (i.e., Fe 0 ). 

4.3. Calculations Involving Concentration 

Before we can begin to study solution chemistry, an understanding of concentration must 

first be developed. Concentration is the ratio of amount of solute to the amount of solvent 

or solution. A solution is a homogeneous mixture of two or more molecules or ions with one 

of the molecules (or ions) being the solvent and all others being the solutes. Solutes are 

dissolved into the solvent. In Experiment 2, you measured the density of ethanol/water solutions 

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 working 

with chemical reactions. Molar concentration, or molarity M, of a solution is defined 

as the number of moles of the solute per liter of solution, and is the most widely used measure 

of concentration. Other measures of concentration [e.g., molality (moles of solute/kg 

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94 

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 molarity 

is 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 of 

water, 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, 

 

1 mol NH3 

moles NH3 = 8.5 g NH3 = 0.499 mol NH3 

17.031 g NH3 molarity NH 3 = 

0.499 mol NH3 

500 mL 

4.3.2. Molarity in solution chemistry 

 

1000 mL 

= 1.0 mol NH3/L = 1.0 M 

1 L 

If you wanted to make 750 mL (3 significant figures) of a 0.200 M aqueous solution of 

sodium chloride, one would calculate the number of grams of sodium chloride (s) by the 

following 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 

= 0.1500 moles NaCl 

molar mass NaCl = 23.00 g/mol of Na 

= 58.45 g/molNaCl 

0.200moles NaCl 

1 L solution 

 

1mole Na 

1mole NaCl 

 

1000 mL 

1 L 

+ 35.45 g/mol of Cl 

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 

 

1mole Cl 

1mole NaCl 

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 displacement 

reaction generating iron cations and elemental copper. However, two possible displacement 

reactions can occur, namely 

Fe 0 (s) + Cu 2+ (aq) → Fe 2+ (aq) + Cu 0 (s) , (4.1) 

or 

2 Fe 0 (s) + 3 Cu 2+ (aq) → 2 Fe 3+ (aq) + 3 Cu 0 (s) . (4.2) 

[Notice that both eq. (4.1) and (4.2) are mass and charge balanced and, therefore, represent 

chemical reaction equations (cf. Section 4.5).] In this experiment, we will determine which

eaction [i.e., eq. (4.1) or (4.2)] is consistent with experiment. Therefore, we will add an 

excess of copper sulfate solution to a known amount of elemental iron and will weight the 

product obtained from the chemical reaction. If eq. (4.1) is dominate, the number of moles 

of 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 moles 

of 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 your 

laboratory 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 be 

lost during the reaction.) 

(8) When the reaction has ceased, allow the copper product to settle. Then carefully 

decant the liquid from above the product. 

(9) Add approximately 10 mL of distilled water to the product and swirl the beaker 

to mix. Again, allow the product to settle and decant. This washes the product 

to remove trace amounts of iron cations. Repeat with a second 10 mL portion of 

distilled 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 necessary, 

carefully break up any clumps of copper with the spatula tip. (Check the 

spatula 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 the 

beaker 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 the 

chemical reaction that occurs. For example, eq. (4.1) states that the reactants Fe 0 and 

Cu 2+ will react to form the products Fe 2+ and Cu 0 . The right arrow is the symbol defined to 

mean reacts to yield. The physical state of each substance should be given in any complete 

chemical reaction equation, since the physical state can change the thermodynamics of the 

reaction. Mass and charge conservation are achieved by placing numbers in front of the 

chemical formulas. When both the mass and the charge are conserved (i.e., the mass and 

charge of the reactants equals the mass and charge of the products), the chemical equation 

is said to be balanced. You should always check to verify that a chemical equation 

is balanced before using this equation. 

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the laboratory report.




(1) What is the minimum volume needed to react all of 1.000 g of iron with a 1.0 

M aqueous solution of copper(II) sulfate for the reaction in eq. (4.2). Hint: The 

chemical reaction equation for an aqueous solution of copper(II) sulfate is 

CuSO4 (aq) → Cu 2+ (aq) + SO 2− 

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 turns 

blue, while the foil turns silvery. (a) What is happening? (b) Which is the more 

active metal: silver or copper? 

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98 

(4) The combustion of a thin wire of magnesium metal in an atmosphere of pure 

oxygen produces the brilliant light of a flashbulb in a camera. The equation for 

the reaction is 

2 Mg + O2 → 2 MgO . 

(a) State in words how this equation is read. (b) Give the formula(s) of the 

reactants. (c) Give the formula(s) of the products. (d) Rewrite the equation to 

show that magnesium and magnesium oxide are solids and molecular oxygen is a 

gas.

Results 






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

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: 

Trial 1 Trial 2 

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

partner. 

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: 

Trial 3 Trial 4 

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Calculations 


(1) For each trial, calculate the number of moles of iron used. (The calculation for 

each 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?

(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 during 

this reaction? 

(7) Extra credit (2 pts): If the water was evaporated from the aqueous iron solution 

after the reaction, what ionic salt would be left as the residue? 

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(1) Write the balanced chemical reaction equations for the following: (a) iron reacts 

with molecular oxygen to yield iron(III) oxide, (b) silver nitrate and calcium chloride 

react to form calcium nitrate and silver chloride, and (c) ethane reacts with 

oxygen 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. 

103 

(b) How many grams of Cl2 gas would be required to remove the silver impurity in 

250 grams of 95% (by mass) pure gold? (Assume that silver is the only impurity.)

104 



Version 1.0, 2008 

5.1. Safety 

Experiments 

EXPERIMENT 5 

Copper reactions 

Aqueous nitric acid and aqueous sulfuric acid are hazardous. They produce severe burns 

on the skin and the vapor is a lung irritant. These compounds should be handled in a 

fume hood while wearing safety glasses and gloves. Rinse your hands with water for 5 

minutes after handling the acid bottles. The gases produced during these reactions are 

toxic and must be avoided. Thus, when the experiment states that a procedure should be 

performed in the fume hood, DO so. Aqueous sodium hydroxide is also corrosive to the 

skin and is especially dangerous if splashed into eyes. Again, handle this compound with 

gloves. Copper sulfate pentahydrate is harmful if swallowed and can cause irritation to 

skin, eyes and respiratory tract. At high concentrations, this compound can affect the liver 

and kidneys. If any of this compound comes in contact with the skin or the eyes, flush with 

plenty of water for at least 15 minutes. Methanol is flammable and should be heated with 

care. Bunsen burners, hot glassware, and metal ring stands 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 handling hot glassware. 


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 salts 

react to form an insoluble neutral compound and a different soluble salt. A good example of 

this type of reaction is the reaction of aqueous sodium chloride with aqueous silver nitrate 

to 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. Redox 

reactions, such as the reaction investigated in Experiment 4, involve the transfer of electrons 

from one atom to another. Decomposition reactions occur when a compound breaks down 

to form simpler substances, such as when calcium carbonate decomposes into calcium oxide 

and carbon dioxide under high temperatures. In this experiment, various chemical reactions 

involving copper and copper salts will be investigated. 

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5.3. Solubility 

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, perchlorate, 

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 except 

those of Group IA and those of calcium, strontium and 

barium. When metal oxides dissolve, these compounds react 

with water to form hydroxides. 

6. All salts that contain phosphate, carbonate, sulfite, and sulfide 

are insoluble except those of Group IA and ammonium. 

The reactions that will be investigated below will rely on the precipitation of insoluble 

salts from aqueous solutions. The formation of a precipitate can be predicted using the 

solubility 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, and 

then react this product with sodium hydroxide to obtain the solid copper hydroxide. The 

procedure is as follows: 

Reaction 1 

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

mL 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 the 

reaction on the Report Sheet. 

(3) Once the copper has dissolved, slowly add 75 mL of deionized water. Then, return 

to your laboratory bench. 

Reaction 2 

(1) While stirring the solution created in Reaction 1, add 30 mL of 3.0 M aqueous 

sodium 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 solution 

from Reaction 2 to almost boiling. Do not boil the solution. Stir while heating 

to 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 stir 

for 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 of 

methanol. 

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

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 oxide 

must 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, schematically 

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 dish 

to cool again. 

(8) Once crystals begin to appear, set the evaporating dish into a shallow ice water 

bath 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 side 

of a piece of filter paper. When a solution is poured on top of this filter paper, the vacuum 

acts 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 the 

Buchner 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 precipitate 

is centered in the middle of the funnel. 

(6) Once the solution has been filtered, rinse the container with a small amount of the 

solvent 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. 









the laboratory report.




Write the balanced reaction equation for the following reactions. (Remember that a balanced 

reaction equation includes the physical information about the compounds and is 

charge and mass balanced.) 

(1) reaction of solid copper with aqueous nitric acid and molecular oxygen to produce 

gaseous nitrogen dioxide and copper(II) nitrate. 

(2) Write the balanced reaction equation for the reaction of copper(II) nitrate with 

aqueous sodium hydroxide. 

(3) Write the balanced reaction equation for the decomposition of copper(II) hydroxide 

to copper(II) oxide. 

(4) Write the balanced reaction equation for the reaction of copper(II) oxide with 

aqueous sulfuric acid. 

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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 . 


Observations (include color and texture changes):

Calculations 


Experiment 5A. Preparation of copper(II) oxide 

(1) What is the molarity of the copper(II) nitrate solution after the addition of distilled 

water in Reaction 1? 

(2) How many moles of hydroxide ion (OH − ) are in the 30 mL of 3.0 M aqueous 

sodium hydroxide solution used in Reaction 2? 

(3) The limiting reactant is the reactant that controls the extent of reaction, because it 

is present in the smallest molar quantity. What is the limiting reactant in Reaction 

2? 

(4) Using the moles of the limiting reactant, determine the mass of copper(II) hydroxide 

created during Reaction 2. 

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(5) Using the mass of copper(II) hydroxide calculated above, determine the theoretical 

yield of copper(II) oxide in Reaction 3. (The theoretical yield is the mass of 

copper(II) oxide predicted if all reactions proceed at 100% efficiency.) 

(6) The percent yield of any reaction is 

experimental mass 

% yield = 

theoretical mass 

What is the percent yield of this reaction? 

× 100 . (5.1) 

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

yield?

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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Describe whether the error introduced by each of the following problems would result in a 

high 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 addition 

of sodium hydroxide. 

(2) Insufficient sodium hydroxide is added. 

(3) The solution bumps during the heating of copper(II) hydroxide to produce copper(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 solution 

residues from the copper(II) oxide. 

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

Syllabus - Queens College Academic Senate - CUNY

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