Bob Jones High School Chemistry I
Updated 3/23/2016
Table of Contents
*** Denotes a lab that must be ordered from Science in Motion
List of Required Labs for Chemistry I ...................................................................................... 3
Chemistry Safety Contract ........................................................................................................ 4
Safety Contract Acknowledgement .......................................................................................... 6
Helpful Aid in Preparing Acid/Base Solutions ......................................................................... 7
Instructions on operating the still.............................................................................................. 7
Cleaning Glassware Basics ....................................................................................................... 8
Laboratory Equipment ............................................................................................................ 11
Laboratory Hazards ................................................................................................................. 13
Safe Laboratory Techniques ................................................................................................ 16
Pre-Lab Writeup...................................................................................................................... 19
Lab Equipment and Safety ...................................................................................................... 20
Conservation of Mass ............................................................................................................. 22
Density Determination ............................................................................................................ 24
Chromatography ..................................................................................................................... 28
Sand and Salt Separation ........................................................................................................ 29
***Energy Content of Foods .................................................................................................. 32
Specific Heat ........................................................................................................................... 35
Calorimetry of Various Foods ................................................................................................ 38
How Sweet It Is ....................................................................................................................... 40
Emission Spectroscopy: Flame Lab with Chloride Salts ........................................................ 46
***Excited Elements .............................................................................................................. 50
Types of Bonding in Solids..................................................................................................... 54
***Evaporation and Intermolecular Attractions ..................................................................... 56
***Moisture Content of Popcorn ............................................................................................ 60
Percent Composition of Bubble Gum ..................................................................................... 63
Percentage Composition of Hydrates...................................................................................... 64
Determining an Empirical Formula ........................................................................................ 67
Chemical Reactions: Single and Double Replacement reactions. ......................................... 70
Types of Chemical Reactions: A Sampler Platter.................................................................. 74
Putting Atoms Together: Synthesis of Zinc Iodide................................................................ 79
Taking Compounds Apart: Decomposing Zinc Iodide .......................................................... 81
Activity Series of Metals ........................................................................................................ 83
Solubility in Double Replacement Reactions ......................................................................... 87
Stoichiometry of Copper (II) Sulfate and Iron ........................................................................ 90
Stoichiometry of a Precipitate ................................................................................................. 93
Limiting Reactant Activity ..................................................................................................... 96
Stoichiometry of HCl and NaHCO3 ........................................................................................ 97
Molar Volume of a Gas........................................................................................................... 99
“Wet” Dry Ice ....................................................................................................................... 102
***Properties of Solutions: Electrolytes and Non-Electrolytes........................................... 103
Acid/Base Indicators ............................................................................................................. 106
39 Drop pH Lab .................................................................................................................... 109
Titration Lab ......................................................................................................................... 112
Qualitative Analysis of the Group I Cations ......................................................................... 115
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Copper into Gold: The Alchemist’s Dream ......................................................................... 120
Making Paint ......................................................................................................................... 121
Guar Gum Slime ................................................................................................................... 122
Ice Cream: Freezing Point Depression of a Solution ............................................................ 124
Periodic Table ....................................................................................................................... 125
Rules of Writing Equations, Solubility Rules, Activity Series of Metals ............................. 126
Demonstration: Air Pressure ................................................................................................. 129
Demonstration: Burning Lycopodium Powder: .................................................................... 130
Demonstration: Burning Magnesium .................................................................................... 131
Demonstration: Density: Coke vs. Diet Coke ...................................................................... 132
Demonstration:Egg In a Bottle ............................................................................................. 134
Demonstration: Elephant Toothpaste .................................................................................... 135
Demonstration: Endothermic Reaction ................................................................................. 136
Demonstration: Gummy Bear Sacrifice: Energy of Oxidation of Carbohydrates ................ 137
Demonstration: Methane Bubbles......................................................................................... 138
Demonstration: Methane Mamba ......................................................................................... 139
Demonstration: Money to Burn ........................................................................................... 143
Demonstration: Thionin — The Two-Faced Solution: Light Energy and Chemical Energy
............................................................................................................................................... 144
Demonstration: Test Tube Thunderstorm ............................................................................. 146
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List of Required Labs for Chemistry I
The following are the labs absolutely required for every chem. I student.
1. Lab Equipment and Safety
2. Sand and salt or other filtration lab
3. Specific heat lab or energy content of foods or calorimetry of various foods
4. One of the Stoichiometry labs
5. Flame Lab and/or Excited Elements Lab
6. Percent composition of hydrates, percent composition of bubble gum and/or percent
composition of Popcorn
7. One of the Reactions labs
8. Determining an empirical formula lab
9. An acid base lab
10. Molar volume of a gas lab.
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Chemistry Safety Contract
PURPOSE
Science is a hands-on laboratory class. You will be doing many laboratory activities which require the use of hazardous chemicals.
Safety in the science classroom is the #1 priority for students, teachers, and parents. To ensure a safe science classroom, a list of rules
has been developed and provided to you in this student safety contract. These rules must be followed at all times. The
acknowledgment sheet must be signed by both you and a parent or guardian before you can participate in the laboratory. Any
questions by the student or parent should be addressed to the teacher before this contract is signed. Lab activities may be videotaped
to encourage safe practices.
GENERAL RULES
1. Conduct yourself in a responsible manner at all times in the laboratory.
2. Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ask the instructor
before proceeding.
3. Never work alone. No student may work in the laboratory without an instructor present.
4. When first entering a science room, do not touch any equipment, chemicals, or other materials in the laboratory area until you are
instructed to do so.
5. Do not eat food, drink beverages, or chew gum in the laboratory. Do not use laboratory glassware as containers for food or
beverages.
6. Perform only those experiments authorized by the instructor. Never do anything in the laboratory that is not called for in the
laboratory procedures or by your instructor. Carefully follow all instructions, both written and oral. Unauthorized experiments are
prohibited.
7. Be prepared for your work in the laboratory. Read all procedures thoroughly before entering the laboratory.
8. Never fool around in the laboratory. Horseplay, practical jokes, and pranks are dangerous and prohibited.
9. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times. Bring only your laboratory
instructions, worksheets, and/or reports to the work area. Other materials (books, purses, backpacks, etc.) should be stored in the
classroom area.
10. Keep aisles clear. The chemical storage area is off limit to all students.
11. Know the locations and operating procedures of all safety equipment including the first aid kit, eyewash station, safety shower, fire
extinguisher, and fire blanket. Know where the fire alarm and the exits are located.
12. Always work in a well-ventilated area. Use the fume hood when working with volatile substances or poisonous vapors. Never
place your head into the fume hood.
13. Be alert and proceed with caution at all times in the laboratory. Notify the instructor immediately of any unsafe conditions you
observe.
14. Dispose of all chemical waste properly. Never mix chemicals in sink drains. Sinks are to be used only for water and those
solutions designated by the instructor. Solid chemicals, metals, matches, filter paper, and all other insoluble materials are to be
disposed of in the proper waste containers, not in the sink. Check the label of all waste containers twice before adding your
chemical waste to the container.
15. Labels and equipment instructions must be read carefully before use. Set up and use the prescribed apparatus as directed in the
laboratory instructions or by your instructor.
16. Keep hands away from face, eyes, mouth and body while using chemicals or preserved specimens. Wash your hands with soap and
water after performing all experiments. Clean all work surfaces and apparatus at the end of the experiment. Return all equipment
clean and in working order to the proper storage area.
17. Experiments must be personally monitored at all times. You will be assigned a laboratory station at which to work. Do not wander
around the room, distract other students, or interfere with the laboratory experiments of others.
18. Know what to do if there is a fire drill during a laboratory period; containers must be closed, gas valves turned off, fume hoods
turned off, and any electrical equipment turned off.
19. If you have a medical condition (e.g., allergies, pregnancy, etc.), check with your physician prior to working in lab.
CLOTHING
20. Any time chemicals, heat, or glassware are used, students will wear laboratory goggles. There will be no exceptions to this rule!
21. Goggles must be worn during all phases of the lab including set-up, cleanup and everything in between. If you have glasses,
goggles must be worn over them. Contact lenses may be worn in the laboratory ONLY with non-directly vented chemical splash
goggles. Certain solvent liquids and vapors may cause the contact to fuse to the eye. It is the student’s responsibility to inform the
teacher if he/she is wearing contacts during the lab.
22. Dress properly during a laboratory activity. Long hair, dangling jewelry, and loose or baggy clothing are a hazard in the
laboratory. Long hair must be tied back and dangling jewelry and loose or baggy clothing must be secured. Shoes must completely
cover the foot. No sandals allowed. It is recommended to bring an old pair of shoes to keep at school in the event that sandals are
inadvertently worn on a lab day.
23. Lab aprons or lab coats have been provided for your use and should be worn during laboratory activities.
ACCIDENTS AND INJURIES
24. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the instructor immediately, no matter how trivial it may
appear.
25. If you or your lab partner are hurt, immediately yell out “Code one, Code one” to get the instructor’s attention.
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26. If a chemical splashes in your eye(s) or on your skin, immediately flush with running water from the eyewash station or safety
shower for at least 20 minutes. Notify the instructor immediately.
HANDLING CHEMICALS
27. All chemicals in the laboratory are to be considered dangerous. Do not touch, taste, or smell any chemicals unless specifically
instructed to do so. The proper technique for smelling chemical fumes will be demonstrated to you.
28. Check the label on chemical bottles twice before removing any of the contents. Take only as much chemical as you need.
29. Never return unused chemicals to their original containers.
30. When transferring reagents from one container to another, hold the containers away from your body.
31. Acids must be handled with extreme care. You will be shown the proper method for diluting strong acids. Always add acid to
water, swirl or stir the solution and be careful of the heat produced, particularly with sulfuric acid. If an acid is spilled on the skin,
first blot with a paper towel, then go to the sink and run water over the affected area. It’s best to remove as much acid as possible
before washing with water.
32. Never dispense flammable liquids anywhere near an open flame or source of heat.
33. Never remove chemicals or other materials from the laboratory area.
HANDLING GLASSWARE AND EQUIPMENT
34. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Place broken or waste
glassware in the designated glass disposal container.
35. Fill wash bottles only with distilled water and use only as intended, e.g., rinsing glassware and equipment, or adding water to a
container.
36. When removing an electrical plug from its socket, grasp the plug, not the electrical cord. Hands must be completely dry before
touching an electrical switch, plug, or outlet.
37. Examine glassware before each use. Never use chipped or cracked glassware. Never use dirty glassware.
38. Report damaged electrical equipment immediately. Look for things such as frayed cords, exposed wires, and loose connections.
Do not use damaged electrical equipment.
39. Do not immerse hot glassware in cold water; it may shatter.
HEATING SUBSTANCES
40. Exercise extreme caution when using a gas burner. Take care that hair, clothing and hands are a safe distance from the flame at all
times. Do not put any substance into the flame unless specifically instructed to do so. Never reach over an exposed flame. Light
gas burners only as instructed by the teacher.
41. Never leave a lit burner unattended. Never leave anything that is being heated or is visibly reacting unattended. Always turn the
burner or hot plate off when not in use.
42. You will be instructed in the proper method of heating and boiling liquids in test tubes. Do not point the open end of a test tube
being heated at yourself or anyone else.
43. Heated metals and glass remain very hot for a long time. They should be set aside to cool and picked up with caution. Use tongs
or heat-protective gloves if necessary.
44. Never look into a container that is being heated.
45. Allow plenty of time for hot apparatus to cool before touching it.
46.Hot and cold glass have the same visual appearance. Determine if an object is hot by bringing the back of your hand close to it
prior to grasping it.
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Student Name (Print clearly)__________________________________________________
Safety Contract Acknowledgement (Please sign, tear out of the lab
your teacher)
QUESTIONS
Do you wear contact lenses?
___ YES ___ NO
Are you color blind?
___ YES ___ NO
Do you have allergies?
___ YES ___ NO
If so, list specific allergies
manual and return to
List any medical conditions
of which the teacher should be aware.
AGREEMENT
I, ___________________________ , (student’s name) have read and agree to follow all of the safety rules set forth
in this contract. I realize that I must obey these rules to ensure my own safety, and that of my fellow students and
instructors. I will cooperate to the fullest extent with my instructor and fellow students to maintain a safe lab
environment. I will also closely follow the oral and written instructions provided by the instructor. I am aware that
any violation of this safety contract that results in unsafe conduct in the laboratory or misbehavior on my part, may
result in being removed from the laboratory, detention, receiving a failing grade, and/or dismissal from the course.
_______________________________________
Student Signature
____________________________
Date
Dear Parent or Guardian:
We feel that you should be informed regarding the school’s effort to create and maintain a safe science classroom/
laboratory environment.
With the cooperation of the instructors, parents, and students, a safety instruction program can eliminate, prevent,
and correct possible hazards.
You should be aware of the safety instructions your son/daughter will receive before engaging in any laboratory
work. Please read the list of safety rules above. No student will be permitted to perform laboratory activities unless
this contract is signed by both the student and parent/guardian and is on file with the teacher. Your signature on this
contract indicates that you have read this Student Safety Contract, are aware of the measures taken to ensure the
safety of your son/daughter in the science laboratory, and will instruct your son/daughter to uphold his/her
agreement to follow these rules and procedures in the laboratory. Please be aware that parents will be notified in the
event of an accident in the lab. Generally, theses accidents are minor in nature but we want you to be aware of any
accidents in the lab.
_______________________________________
Parent/Guardian Signature
_______________________________________
Date
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Helpful Aid in Preparing Acid/Base Solutions
Chemical Name
Ammonium Hydroxide
NH4OH
M.W.: 35.05
14.8 M concentrate
Acetic Acid
HC2H3O2
M.W.: 60.05
83.0
249.0
498.0
64.0
192.0
384.0
67.5
202.5
405.0
56.0
168.0
336.0
14.8 M concentrate
Sulfuric Acid
H2SO4
M.W.: 98.08
345.0
15.8 M concentrate
Phosphoric Acid
H3PO4
M.W.: 97.99
172.5
12.1 M concentrate
Nitric Acid
HNO3
M.W.: 63.02
57.5
17.4 M concentrate
Hydrochloric Acid
HCl
M.W.: 36.46
mL of concentrated acid/base
needed to prepare 1 L of solution
1.0 M
3.0 M
6.0 M
67.5
202.5
405.0
18.0 M concentrate
Instructions on operating the still.
1. Be sure the elements isn’t overly mucked up with scale. If it is heavily scaled, follow the
cleaning directions in the manual (on the wall next to the still) using formic acid solution.
2. Start water running into the still, the outflow should fill a 1 L beaker in one minute. Be sure
the distilled water outlet tube is in the carboy or else it will drip on everything. 
3. After that the still will run on its own, you just have to check it every 20 minutes or so to
make sure the water flow is sufficient. The key is that if the water dripping into the carboy is
hot, you don’t have enough water flowing. If there is too much water flowing the water
won’t boil along the entire length of the element and it won’t be operating at peak efficiency.
4. After you turn off the power, you have to let the water run for 15 min to condense the steam
present and to cool down the system.
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Cleaning Glassware Basics
It's generally easier to clean glassware if you do it right away. When detergent is used, it's
usually one designed for lab glassware, such as Liquinox or Alconox. These detergents
are preferable to any dishwashing detergent you might use on dishes at home.
Much of the time, detergent and tap water are neither required nor desirable. You can
rinse the glassware with the proper solvent, then finish up with a couple of rinses with
distilled water, followed by final rinses with distilled water.
How to Wash Out Common Lab Chemicals
Water Soluble Solutions
(e.g., sodium chloride or sucrose solutions) Rinse 3-4 times with distilled water then put
the glassware away.
Water Insoluble Solutions
(e.g., solutions in hexane or chloroform) Rinse 2-3 times with ethanol or acetone, rinse 34 times with distilled water, then put the glassware away.
In some situations other solvents need to be used for the initial rinse.
Strong Acids
(e.g., concentrated HCl or H2SO4) Under the fume hood, carefully rinse the glassware
with copious volumes of tap water. Rinse 3-4 times with distilled water, then put the
glassware away.
Strong Bases
(e.g., 6M NaOH or concentrated NH4OH) Under the fume hood, carefully rinse the
glassware with copious volumes of tap water. Rinse 3-4 times with distilled water, then
put the glassware away.
Weak Acids
(e.g., acetic acid solutions or dilutions of strong acids such as 0.1M or 1M HCl or H2SO4)
Rinse 3-4 times with distilled water before putting the glassware away.
Weak Bases
(e.g., 0.1M and 1M NaOH and NH4OH) Rinse thoroughly with tap water to remove the
base, then rinse 3-4 times with distilled water before putting the glassware away.
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Washing Special Glassware
Glassware Used for Organic Chemistry
Rinse the glassware with the appropriate solvent. Use distilled water for water-soluble
contents. Use ethanol for ethanol-soluble contents, followed by rinses in distilled water.
Rinse with other solvents as needed, followed by ethanol and finally distilled water. If the
glassware requires scrubbing, scrub with a brush using hot soapy water, rinse thoroughly
with tap water, followed by rinses with distilled water.
Burets
Wash with hot soapy water, rinse thoroughly with tap water, then rinse 3-4 times with
distilled water. Be sure the final rinses sheet off of the glass. Burets need to be thoroughly
clean to be used for quantitative labwork.
Pipets and Volumetric Flasks
In some cases, you may need to soak the glassware overnight in soapy water. Clean
pipets and volumetric flasks using warm soapy water. The glassware may require
scrubbing with a brush. Rinse with tap water followed by 3-4 rinses with distilled water.
Drying or Not Drying Glassware
Not Drying
It is inadvisable to dry glassware with a paper towel or forced air since this can
introduce fibers or impurities that can contaminate the solution. Normally you can
allow glassware to air dry on the shelf. Otherwise, if you are adding water to the
glassware, it is fine to leave it wet (unless it will affect the concentration of the final
solution). If the solvent will be ether, you can rinse the glassware with ethanol or
acetone to remove the water, then rinse with the final solution to remove the alcohol
or acetone.
Rinsing with Reagent
If water will affect the concentration of the final solution, triple rinse the glassware
with the solution.
Drying Glassware
If glassware is to be used immediately after washing and must be dry, rinse it 2-3
times with acetone. This will remove any water and will evaporate quickly. While it's
not a great idea to blow air into glassware to dry it, sometimes you can apply a
vacuum to evaporate the solvent.
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Additional Notes
Remove stoppers and stopcocks when they are not in use. Otherwise they may 'freeze'
in place.
You can degrease ground glass joints by wiping them with a lint-free towel soaked
with ether or acetone. Wear gloves and avoid breathing the fumes.
The distilled water rinse should form a smooth sheet when poured through clean
glassware. If this sheeting action is not seen, more aggressive cleaning methods may
be needed
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Laboratory Equipment
Beakers: Pyrex or Kimax, used as containers or for rough measurements, common sizes are 100 mL, 150 mI., 250 mL,
400 mL, 600 mL, and I L, may be heated.
Burets: glass and Teflon, used to measure volumes for titrations, 50 or 100 mL
Burner: metal, connected to gas supply with rubber tubing to produce a flame to heat substances.
Ceramic square: ceramic material, is placed under hot equipment or glassware.
Clamps: metal or plastic, may be fastened to ring stand to support laboratory apparatus, glassware or other equipment,
types include clamp holder, ring clamp, test tube clamp, 3-prong clamp, and double buret clamp.
Clay triangle: wire frame with porcelain rods; placed on a ring clamp to hold a crucible while heating. Crucible and
cover: porcelain; used to heat small amounts of substances at high temperatures. Crucible tongs: metal; used to pick up
and hold crucibles, crucible covers, and other small objects.
Dropper pipet: glass tube with rubber bulb; used to measure or transfer small volumes of liquids and solutions.
Erlenmeyer flasks: Pyrex, Kimax, used as containers, common sizes are 50 mL, 125 mL, 250 mL; may be heated.
Evaporating dish: porcelain, used to heat and evaporate small volumes of solutions.
Forceps: metal, used to pick up or hold small objects.
Funnel: glass or plastic, used in adding liquids to small-mouth containers or filtering solids from mixtures.
Gas collecting tube: glass tube with graduations, used to collect measured volumes of gas by water displacement.
Glass rod with nichrome wire: used for flame tests.
Graduated cylinder: glass and plastic, used to measure approximate volumes of liquids, common sizes are 10 mL, 50
mL, and 100 mL, do not heat, keep plastic rings at the top of the cylinder.
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Mortar and pestle: glass or porcelain, used to grind large solid chunks into powders.
Ring stand: vertical metal rod fixed to a heavy metal base, used as a support for many pieces of laboratory equipment.
Rubber policeman: glass and rubber, used to scrape solids from containers.
Safety goggles: plastic, chemical splash, with indirect vents, must be worn at all times in the laboratory.
Spatula/scoopula: metal, used to transfer solid chemicals, spatula used for smaller quantities.
Stirring rod: glass, used to stir or to pour liquids.
Test tube: Pyrex or Kimax, used for reactions, common sizes 13mm x 100 mm and 20 mm x 150 mm, may be heated.
Test tube brush: wire handle with bristles, used to clean glassware with small diameters, use correct brush size.
Test tube holder: spring metal used to hold test tubes when heating, squeeze to open, release to close.
Test tube rack: wood or plastic, holds test tubes in vertical position, used during reactions or after cleaning test tubes.
Thermometer: alcohol in glass, used to measure temperatures, temperature range of -10°C to 110°C.
Triple-beam balance: metal, used to measure masses of substances in grams.
Volumetric flask Pyrex or Kimax, used to make solutions with exact amounts, common sizes are 100 mL, 250 mL, 500
mL, and l L, do not heat.
Wash bottle: polyethylene squeeze-bottle, used to dispense distilled water at student workstation.
Watch glass: Pyrex or Kimax, used to cover evaporating dish or beaker or to observe physical properties of solids.
Wire gauze: wire and ceramic, placed on a ring clamp to spread the heat of a burner flame under a container.
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Laboratory Hazards
Students should be aware of possible hazards in the laboratory and take the appropriate safety
precautions. By doing so, the risks of working in the chemistry laboratory will be reduced. This
section addresses laboratory hazards, how to prevent accidents, and what to do if an accident occurs.
Chemical Burns
A chemical burn occurs when the skin or a mucous membrane is damaged by contact with a
substance. Corrosive substances can cause severe burns. An irritant is a chemical that can irritate
the skin and membranes of the eyes, nose, throat, and lungs. Chemicals that are corrosive or
irritating must be treated with special care. Chemical burns can be severe, and permanent damage to
mucous membranes can occur despite the best efforts to rinse chemicals from an affected area. The
best defense against chemical burns is prevention.
Without exception, safety goggles must be worn during all phases of the laboratory period
even during cleanup. Goggles should be put on as soon as you enter the laboratory and remain over
your eyes until you leave the laboratory. Should any chemical splash in your eye, immediately
notify your teacher. Use a continuous flow of running water to flush your eye for 20 minutes. Do
not rub eye. Wear a laboratory coat or apron and lace-up shoes and socks (no sandals) to protect
your clothing, feet, and other areas of your body. If corrosive chemicals come in contact with your
skin, rinse the affected area with water for several minutes. If there is no burning sensation, wash
area with soap and water.
*Estimates for the time required for permanent corneal damage to occur following exposure to 1M
NaOH are in the range of 30 seconds.
An additional burn hazard exists when concentrated acids or bases are mixed with water. The heat
released in mixing these chemicals with water can cause the mixture to boil, spattering corrosive
chemical. The heat can also cause regular glass containers to break, spilling the corrosive chemical.
To avoid these hazards, always add acid or base to water, very slowly and with stirring, and never
the reverse. * As a precaution, Pyrex or Kimax containers, glassware that has been treated to
withstand high temperatures, should always be used.
*Concentrated sulfuric acid causes thermal burns because it reacts with water in the skin releasing
substantial amounts of heat. Nitric acid does not produce thermal burns but denatures the proteins in
the skin destroying tissue. Nitric acid burns are very slow to heal.
Thermal Burns
A thermal burn can occur if you touch hot equipment or get too close to an open flame. You should
be aware that hot and cold glassware look the same. If a gas burner or hot plate is being used, some
of the equipment nearby may be hot. Hold your hand near an item to feel for heat before touching it.
Treat a thermal burn by immediately running cold water on the burned area. Continue to apply the
cold water until the pain is reduced. This usually takes several minutes. In addition to reducing the
pain, cooling the burned area also serves to speed the healing process. Greases and oils should not be
used on burns because they tend to trap heat. Medical assistance should be sought for any serious
burn. Notify your teacher immediately if you are burned.
Cuts from Glass
Many cuts that occur in the laboratory are avoidable by following a few simple rules. You should
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never use broken, cracked, or chipped glassware. If you should break a piece of glassware, do not
pick up the broken pieces with your hands. Use a brush or broom and dustpan to sweep up the shards
of glass. All broken glass should be placed in the box labeled for broken glass. You should never
place broken glass in a regular trashcan.
If you receive a minor cut, briefly allow the cut to bleed by squeezing the cut. Place the injured area
under cold running water, and notify your teacher. Serious cuts and deep puncture wounds require
immediate medical attention. Notify your teacher immediately. Control the bleeding by applying
pressure with the fingertips or by firmly pressing on the wound with a clean towel or gauze.
Cuts frequently occur when thermometers or pieces of glass tubing are inserted into rubber stoppers.
Insert glass tubing only under the supervision of your teacher.
Poisoning
Many of the chemicals used in the experiments in this manual are mildly to moderately toxic. To
prevent poisoning, never eat, chew gum or drink in the laboratory. Do not touch chemicals. Never
taste any chemical in lab. Keep your hands away from your face. Always wash your hands with soap
and water at the end of the lab. In this way you will prevent chemicals that might get on your hands
from reaching your mouth, nose or eyes.
In some cases, the detection of an odor is used to indicate that a chemical reaction has taken place. It
is important to note that many gases are toxic when inhaled. If you must detect an odor, use your
hand to gently fan some of the gas toward your nose. Take a small sniff of the gas instead of a deep
breath. This will minimize the amount of gas sampled.
Fire
A fire may occur if chemicals are mixed improperly or if flammable materials come too close to a
burner flame or hot plate. Use a hot plate as a heat source instead of a burner when flammable
chemicals are being used or produced. When using a hot plate or burner, prevent fires by tying back
long hair and loose fitting clothing.
If hair or clothing should catch fire, DO NOT RUN. Running fans a fire. Stop, drop to the floor,
and roll slowly to smother the flames. Shout for help. If another person is the victim, get a fire
blanket, located at the front of the lab, to smother the flames. If a shower is nearby, help the victim
to use it.
A fire in a container may be extinguished by smothering the flames with the fire blanket, a notebook,
or some other nonflammable object. In case of a fire on a laboratory workbench, turn off all gas jets
and unplug all appliances. Notify the teacher immediately. If a fire extinguisher is needed, the
teacher will call for it. To use a fire extinguisher, pull the ring, point the nozzle at the base of the
fire, and squeeze the handle. Use short bursts from the extinguisher, rather than one continuous
spray. Caution: Never direct the spray of a fire extinguisher into a person's face. If a fire is not
extinguished quickly, leave the laboratory. Crawl to the door if necessary to avoid the smoke. Do
not return to the laboratory until you are told it is safe.
The fire extinguishers available in lab are ABC extinguishers. This designation means that the
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extinguishers may be used on three types of fires. These types of fires are: a. paper and trash, b.
liquids or grease, and c. electrical. ABC fire extinguishers should not be used for flammable solids.
Sand is used to extinguish burning flammable solids.
Fire Warning
The signal for a fire drill or fire is the sound of the fire alarm. If the signal is given while in the
laboratory, students should turn off all gas jets and exit immediately.
Tornado Warning
An announcement over the intercom is the signal for a tornado drill or warning. Go to the area
indicated by your instructor. You should sit on the floor facing the wall and protect your head. You
should remain in this position until the announcement ending the drill or warning is made.
Always/Never Rules
Always
 wear safety goggles
 wear protective clothing
 use proper techniques and procedures
 discard wastes properly
 know the location and use of safety
equipment
 be alert, serious and responsible in lab
Never
 eat or drink in the lab
 clutter your work area
 perform unauthorized experiments
 enter the chemical storage area
 take unnecessary risks
 remove stock chemicals from the supply
area
Report any injury, accident, or chemical spill to the teacher immediately. Know the location of the
eyewash, fire blanket, fire extinguisher, and shower.
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Safe Laboratory Techniques
Pouring Liquids
 Always wear safety goggles when handling chemicals.
 Always read the label on a reagent bottle before using, and then read the label again. Never touch
chemicals with your hands.
 Never return unused chemicals to their original containers. To avoid waste, pour small amounts of
reagents into small beakers and share with students around you.
Follow this procedure when pouring liquids.
1. Remove the lid.
2. Hold bottle with the label in the palm of your hand.
3. When pouring a liquid from a reagent bottle into a beaker or funnel,
the reagent should be poured slowly down a glass stirring rod.
4. When pouring a liquid from a bottle into a test tube or graduated
cylinder, the empty container should be held at eye level. Pour the
liquid slowly until the correct volume is obtained.
5. Place the lid back on the bottle before removing the lid from another
reagent bottle.
Filtering a Mixture
To separate a solid from a liquid, the most common method is gravity
filtration.
1. Fold the filter paper in half and then quarters.
2. Open the folded paper to form a cone with one layer of paper on one
side and three layers on the other side.
3. Put the cone in a funnel. Moisten the filter paper with a small amount
of distilled water and gently press the paper against the sides of the
funnel.
4. Place a beaker beneath the funnel with the tip of the funnel just
touching the inside surface of the beaker about one inch below the
rim.
5. Use a stream of distilled water to wash the solid remaining in the
beaker into the funnel. Wash the solid in the filter with distilled
water to remove all traces of solvent. Dry the solid.
6. Pour the liquid down a glass stirring rod into the funnel. Keep the
liquid below the top edge of the paper at all times to prevent
overflow.
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Using a Gas Burner
Laboratory burners produce various kinds of flames when different
mixtures of gas and air are burned. The Tirrell burner has
adjustable air vents and a gas control valve in the base.
1. Examine a Tirrell burner and identify the parts.
2. Connect the burner to the gas supply with rubber tubing.
3. Close the air vents. Close the gas control valve at the bottom of
the burner and then open both about 1 1/2 full turns.
4. Hold a lighter at the top of the barrel of the burner and turn on
the gas supply at the lab station. With a Tirrell burner, the main
gas supply should be opened fully and the gas flow regulated
by the gas control valve at the base of the burner. The flame
may be yellow or a single blue.
5. Open the air vents slowly, to admit more air into the flame, to
produce a light blue cone-shaped flame. If the flame blows out
after lighting, turn off the gas supply, and relight. Continue to
open the air vents to produce a blue double-cone flame.
6. Adjust the gas supply to produce the desired size of flame. For
a smaller flame, close the air vent slightly and reduce the gas
supply. Practice adjusting the flame.
7. Turn the burner off at the main gas supply valve when finished.
Caution: Tie back long hair and pull back loose clothing when working with a lab burner. Do not reach
across a flame. Do not use a burner around flammables. Never leave a burner flame unattended Know
the location of fire extinguisher, the fire blanket, and safety shower.
Heating a Liquid in a Test Tube
1. Adjust the burner to give a small single blue flame.
2. Fill a test tube no more than one-third full of the liquid to be heated.
3. Hold the test tube with a test tube holder. The test tube holder should
grip the test tube near the mouth of the tube.
4. Place the test tube in a slanting position in the flame, gently heat the
entire length of the test tube and then heat the substance in the test tube a
short distance below the surface of the substance.
5. Gently shake the tube as it is heated until the substance melts, boils, or
reaches the desired temperature.
Caution: Never point the open end of a test tube toward yourself or others. Never
heat the bottom of a test tube held in a vertical position.
Heating a Liquid in a Beaker.
1. Fasten a ring clamp securely to a ring stand so that it is about
three to six centimeters above the top of a gas burner.
2. To set up a hot water bath or boil a liquid in a beaker:
3. Place wire gauze on the ring of a ring clamp.
4. Place a half-filled beaker of liquid on the wire gauze.
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5.
6.
7.
8.
Updated 3/23/2016
Light the burner to produce a hot flame.
Place the burner beneath the wire gauze.
Adjust flame to produce desired temperature.
Caution: Never heat plastic beakers or graduated glassware in
a burner flame. Never let a beaker boil dry; add water to the
beaker as necessary.
Measuring Mass Using an Electronic Balance
To find the mass of an object, follow these general rules.
1. Zero the balance, and place the object on the balance pan. If you are measuring out a
chemical, use a weigh boat. Be sure to zero or tare the balance with the weighing boat on the
electronic balance.
2. If a chemical is spilled on or near the balance, clean it up immediately. If in doubt, check
with the teacher.
3. Never attempt to weigh an object with a mass greater than the maximum capacity of the
balance.
4. Do not attempt to use a balance until the teacher has demonstrated the proper technique.
Measuring Volume
Volume measurements are important in experimental procedures. Accurate laboratory
measurements are made using graduated cylinders, pipets, burets, or volumetric flasks. Although
some beakers have graduation marks, these marks are designed for quick, rough estimates of
volume. Liquid volumes are usually measured in milliliters.
Using a Graduated Cylinder
Place about 50 mL of water in a 100-mL graduated cylinder and
set the cylinder on the laboratory bench. Look at the surface of
the water. The surface curves upwards where the water contacts
the cylinder walls. This curved surface is called a meniscus.
A volume measurement is always read at the bottom of the
meniscus with your eye at the same level as the surface of the
liquid. To make the meniscus more visible place a finger or a
dark piece of paper behind and just below the meniscus while
making the reading.
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Pre-Lab Writeup
To prepare yourself for lab, it is necessary to be familiar with the concepts and procedures of a
particular lab. Therefore, you need to perform the following steps before you go to lab.
On a separate sheet of paper write: (Do NOT number each item on your paper, Label that section
but do not number.)
1. Your name and the date.
2. Title of the Lab
3. The purpose of a lab. Write the purpose in a complete sentence. The purpose can be
found in the discussion portion of each lab.
4. Materials used in the lab.
5. Safety Precautions from each lab.
6. Procedure. Do not copy the procedure in this lab manual. Summarize the steps.
7. Results you expect to see.
8. Write down any questions you have before you start the lab.
Doing this exercise will prepare for the lab, prepare you for a possible lab quiz and cut down on
unnecessary questions for the teacher. In other words, don’t ever ask your teacher a question that
is answered in the lab manual.
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Lab Equipment and Safety
Discussion:
In this activity you will become familiar with some proper lab procedures used in the Chemistry
lab. It is important to choose the correct equipment in lab depending on what you are doing.
You would never use a measuring cup to measure the amount of water in a bathtub. You would
also never use a bathroom scale to figure out much a piece of paper weighs. At the beginning of
chemistry it is important to master certain techniques and skills. No one wants to be the person
in lab who cannot light a burner or who doesn’t know how to use a balance. It’s time to get with
the program and get the 411 on lab equipment.
Procedure
Measuring the volume of a liquid.
1. Read “Using a graduated cylinder” on page Error! Bookmark not defined..
2. What type of graduated cylinder is at your lab station? (10 mL, 50 mL, 100 mL, etc)
3. How much does each gradation represent?
4. What volume of water is your graduated cylinder holding? Write your answer to the tenths.
5. Put 50 mL of water in a beaker using the beaker as a measuring device. Now pour that
amount into a graduated cylinder. How do the two measurements compare?
Measuring mass with an electronic balance.
1. Read “Measuring mass with an electronic balance” on page Error! Bookmark not defined..
2. Place a weigh boat on the balance. Press zero/tare to zero the balance.
3. Add 1.00 g of NaCl to the weigh boat. Discard the salt in the trash. Rinse and dry the weigh
boat for future use.
4. Obtain 3 evaporating dishes, 2 crucibles and 5 pieces of filter paper. Determine and record
the mass for each separate item. Return all the equipment to the drawers and the filter paper
to the teacher desk.
Using a Burner
While commonly called a Bunsen burner, we have Tirrell burners which are an improvement on
the original Bunsen burner.
1. Read page .Error! Bookmark not defined. of the lab manual about using a burner.
2. Read over the fire and thermal burns safety considerations on page.
3. Follow the steps indicated on page to light the burner.
4. Once you have a flame that is burning safely and steadily, you can experiment by completely
closing the air vents. What effect does this have on the flame?
5. Regulate the flow of gas so that the flame extends roughly 8 cm (2.54 cm = 1 inch) above the
burner tube. Now adjust the supply of air until you have a quiet, steady flame with a sharply
defined light blue inner cone. This adjustment gives the highest temperature possible with
your burner. Where is the hottest portion of the flame located? (Consult the poster on the
wall in the lab or classroom)
6. Shut off the gas burner at the gas valve.
7. Draw the burner and label the parts on the data page.
Cleanup
1. Pour all solutions down the drain
2. Clean your lab station.
3. Clean all equipment and leave to dry at your lab station.
4. Wash your hands
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Date_____________
Station #______________
Name(s)________________________________________________________
Data: Lab Equipment
Volume
Type of graduated cylinder
Answer to #3
Volume of water in graduated
cylinder
Comparison of a beaker to a
graduated cylinder
mL
Mass
Evaporating dish 1 mass
g
Evaporating dish 2 mass
g
Evaporating dish 3 mass
g
Crucible 1 mass
g
Crucible 2 mass
g
Filter paper 1 mass
g
Filter paper 2 mass
g
Filter paper 3 mass
g
Filter paper 4 mass
g
Filter paper 5 mass
g
Drawing of burner:
Questions on back of page.
Questions:
1. Why is it important to use a graduated cylinder to measure liquids rather than a beaker?
2. Is it safe to assume that pieces of the same equipment have the same mass? Explain.
3. Why is the nonluminous flame preferred over the yellow luminous flame in the laboratory?
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Conservation of Mass
Discussion:
The law of conservation of mass states that matter is neither created nor destroyed during a
chemical reaction. Therefore, the mass of a system should remain constant during any chemical
process. In this experiment, you will determine whether mass is conserved by examining a
simple chemical reaction and comparing the mass of the system before the reaction with its mass
after the reaction.
MSDS:
Baking soda
Vinegar
Slightly toxic by ingestion. Dust may be irritating to respiratory system.
Substance not considered hazardous. However, not all health aspects of
this substance have been thoroughly investigated. Not for human
consumption.
Procedure
1. Obtain a microplunger and tap it down into a sample of baking soda until the end is packed
with a plug of the powder (4-5 mL) of baking soda should be enough to pack the bulb).
2. Hold the microplunger over a beaker and squeeze the sides of the microplunger to loosen the
plug of baking soda so that it falls into the cup.
3. Use a graduated cylinder to measure 100. mL of vinegar, and pour it into a second beaker.
4. Measure the total mass of the two cups and record.
5. Add the vinegar to the baking soda a little at a time to prevent the reaction from getting out
of control. Allow the vinegar to slowly run down the inside of the cup. Observe and record
your observations about the reaction.
6. When the reaction is complete, measure the mass of the two cups and record. Calculate any
change in mass.
7. Examine the plastic bottle and the hook insert cap. Try to develop a modified procedure that
will test the law of conservation of mass more accurately than the procedure in steps 1-6.
8. Your teacher should approve the procedure you design. Implement the procedure with the
same chemicals and quantities as before but use the bottle and hook insert cap in place of the
two cups. Record your data.
Cleanup
1. Pour all solutions down the drain
2. Clean your lab station.
3. Clean all equipment and leave to dry at your lab station.
4. Wash your hands.
Questions:
1. Answer the questions on page 95 in your textbook.
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Teacher Notes
Conservation of Mass
This lab is designed to be inquiry based. Students are encouraged to create their own procedure.
Much troubleshooting is required since the apparatus(two liter bottle) tends to leak air.
Have modeling clay on hand for students to use to block air leakage from the cap.
The microplunger is made by cutting the bulb of a pipette and inserting a thread through the tip
end of the pipette. See the textbook for a picture.
Students are to redo experiment until they get no change in mass thus proving the law of
conservation of mass.
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Density Determination
Discussion
Density is defined as mass per unit volume. The masses and volumes of different materials may
be measured using direct and indirect methods. Using the measuring skills and techniques
developed in earlier laboratory exercises and in the first part of this experiment, students will
find the masses and volumes of different substances. This data will be used to calculate the
density of three different substances.
Direct Measurement of Volume - Volumes of liquids are measured directly in a graduated
cylinder. Liquid quantities measured in the laboratory are usually expressed in milliliters (mL).
Volumes of regularly shaped geometric solids can be calculated from direct measurements of
their dimensions and the use of the formula, v = l x w x h. Volumes of regularly shaped solids
are expressed in cubic centimeters (cm3).
Indirect Measurement of Volume - The volumes of many solids cannot be measured directly or
easily calculated. These include irregularly shaped objects, such as rocks, and regular solids that
are too small to be measured with any degree of accuracy. Volumes of such solids can be
measured by water displacement. If a solid is immersed in a liquid, the solid will push aside, or
displace a volume of water equal to the volume of the object. Therefore, each milliliter of water
that is displaced by a solid represents one cubic centimeter of solid volume. 1 mL = 1 cm3
Direct Measurement of Mass - The mass of an object is measured directly by placing the object
on the balance pan and reading the mass.
Indirect Measurement of Mass - Many solids and liquids may not be placed directly on a
balance pan. A container must be used. To find the mass of a substance using a container, the
container should be placed on the balance pan. The TARE or Zero button is pressed to rezero the
balance with the container on the pan. Now the liquid or other substance may be placed in the
container, and the mass of the substance read.
Density - In combination with other properties, density can be used to identify substances.
Density is defined as the quantity of matter in a given volume. This relationship, expressed
mass
M
mathematically, is Density 
or D 
. Liquids and solids are reported in g/mL or
volume
V
g/cm3. Gases are represented in g/L or glm3. When the mass and volume of a substance is
known, the density may be calculated.
PROCEDURE:
A. Regular solid - a cube, rectangle, or a cylinder.
1. Measure and record the mass of the object.
2. Measure and record the dimensions of the object.
3. Calculate and record the volume of the object.
4. Calculate and record the density of the object.
Note: Measure dimensions carefully. For geometric solids, any error will be made greater each
time the dimension is multiplied
B. Irregularly shaped solid
1. Measure and record the mass of the object.
2. Add a known volume of water to a graduated cylinder and record the volume.
3. Place the object in the water in the cylinder and record the new volume. The difference in
volumes is the volume of the object. The volume equals the amount of water that was
displaced.
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4. Calculate and record the density of the object.
C. Unknown Liquid - The teacher will assign a number for an unknown liquid.
1. Obtain a clean, dry graduated cylinder from teacher. Record the mass.
2. Take cylinder to teacher to receive liquid. Record the liquid number.
3. Reweigh the cylinder with the liquid in it. Record the mass. The difference in the two
recorded masses is the mass of the liquid.
4. Read and record the volume of the liquid in the cylinder. The amount of liquid in the
cylinder does not have to be a certain value.
5. Return the cylinder to the teacher.
6. Calculate and record the density of the unknown liquid.
Cleanup:
1. 100 mL graduated cylinders get rinsed and left at your station on a paper towel.
2. The irregular object gets dried off and left at the lab station.
3. Your lab area should be free of spilled water, trash, etc
Questions:
1. Which method (A, B, or C) is the method that gives you the most accurate result?
Explain your reasoning.
2. What is a way that you can double check the volume you calculated in the regular
volume?
3. Would the density of the unknown liquid be to high or too low if there was liquid in the
graduated cylinder when you first obtained the graduated cylinder? Explain your
reasoning.
Teacher Notes:
There is a set of density blocks for the regular solids and a small box of random items (screws,
bolts, nuts, etc) for the irregular solids.
For the unknown liquids, organics are used: A list of possibilities and their densities follows:
Acetone
0.790 g/mL
Isopropyl alcohol
0.785 g/mL
t-butyl alcohol
0.780 g/mL
Petroleum ether
0.640 g/mL
Toluene
0.8669 g/mL
Ethyl acetate
0.789 g/mL
Ethanol
0.789 g/mL
Methyl Ethyl Ketone
0.8255 g/mL
Pentane
0.626 g/mL
Hexane
0.655 g/mL
Methylene chloride
1.33 g/mL
Liquids should be kept in the fume hood in small beakers with a watch glass over the beaker.
You need a different 10mL graduated cylinder for each beaker and you can use some of the
liquids more than once. Liquids can be poured back into the stock bottle.
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Date_____________ Station #______________
Name(s)________________________________________________________
Title of Lab: Density Determination
Table 1
Regular solid
Determination of Density by the Direct method
Mass _______
Dimensions_________ x __________ x _________
Volume ______________
Density ______________
Table 2
Determination of Density by the Indirect method
Description of Object ________________________________
Mass of Object
_____________
Initial Volume of Graduated Cylinder
________________
Final Volume of Graduated Cylinder
________________
Volume of Object
________________
Density of Object
________________
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Table 3
Updated 3/23/2016
Determination of Density of an Unknown Liquid
Mass of Dry Graduated Cylinder
____________
Mass of Graduated Cylinder with Liquid
____________
Mass of Liquid
____________
Volume of Liquid
____________
Density of unknown liquid
____________
Liquid’s actual density
Percent Error
____________
Unknown # _______
____________
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Chromatography
Discussion
Most ballpoint pen inks are complex mixtures, containing pigments or dyes that can be separated
by paper chromatography.
Black inks can contain three or more colors; the number of colors depends on the manufacturer.
Each ink formulation has a characteristic pattern that uniquely identifies it.
In this experiment you will develop radial paper chromatograms for four black ballpoint pen
inks, using water as a solvent. You will then repeat this process using isopropanol as the solvent.
You will then measure the distance traveled by each of the individual ink components and the
distance traveled by the solvent front. Finally, you will use these measurements to calculate the
Rf factor for each component.
MSDS:
Isopropanol
Irritant to body tissues. Slightly toxic by ingestion, inhalation, and skin
absorption. The single lethal dose for a human adult is about 250 mL,
although as little as 100 mL can be fatal. Class 1B flammable liquid
.
Procedure
1. Construct an apparatus for paper chromatography as described on page 848 of your book.
You will only make 4 dots. You will use ballpoint pens rather than micropipets to spot your
paper.
2. After 15 minutes or when the water is about 1 cm from the outside edge of the paper, remove
the paper from the Petri dish and allow the chromatogram to dry. Record in the data table the
colors that have separated from each of the four different black inks.
3. Repeat steps 1 and 2 replacing the water with isopropanol.
4. After the chromatogram is dry, use a pencil to mark the point where the solvent front
stopped.
5. With a ruler, measure the distance from the initial ink spot to your mark and record this
distance.
6. Make a small dot with your pencil in the center of each color band.
7. With a ruler, measure the distance form the initial ink spot to each dot separately, and record
each distance.
8. Divide each distance the ink traveled by the distance the solvent traveled. The result is the Rf
for that component. Record Rf values.
Cleanup
1. Throw away the chromatogram.
2. The water and isopropanol solutions can be poured down the drain.
3. Wash your hands.
Questions
1. How do you think they make black ink?
2. If you separated green ink using this technique, what would you expect for results?
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Sand and Salt Separation
Discussion:
Extraction, the separation of substances in a mixture by using a solvent, depends on solubility.
For example, sand can be separated from salt by adding water to the mixture. The salt dissolves
in the water, and the sand settles to the bottom of the container. The sand can be recovered by
decanting the water. The salt can then be recovered by evaporating the water.
Filtration separates substances based on differences in their physical states or in the size of their
particles. For example, a liquid can be separated from a solid by pouring the mixture through a
paper-lined funnel, or if the solid is denser than the liquid, the solid will settle to the bottom of
the container, which will leave the liquid on top. The liquid can then be decanted, which will
leave the solid.
Procedure:
1. Obtain 5.00 g of salt/sand mixture (use the weigh boat and use the procedure learned in the
previous labs). Record the amount of sand/salt in your data table.
2. Place the mixture in a 100 mL beaker and add 10.0 mL of water and stir for 2-3 minutes.
Allow it to settle.
3. Write your names and period on a piece of filter paper using a pencil. NO PENS. Find the
mass of the filter paper and record.
4. Assemble the filtration apparatus as shown on pg 10. Place a 250 mL beaker under the
funnel.
5. Prepare a piece of filter paper as shown on page Error! Bookmark not defined.. Wet
slightly with a wash bottle to make it stay.
6. Grasp the beaker with one hand. With the other hand, pick up a stirring rod and hold it along
the lip of the beaker. Tilt the beaker slightly so that liquid begins to pour out into the funnel
in a slow, steady stream, as shown in Figure 1 pg Error! Bookmark not defined..
7. Wash the remaining solids in the beaker into the filter with a wash bottle.
8. Remove the filter paper, place the wet filter paper on the back counter and allow the contents
to dry.
9. Find the mass of the evaporating dish and watch glass and record.
10. Evaporate liquid filtrate with evaporating dish, watch glass and burner or hot plate. Avoid
spattering.
11. Find the mass of the dry filter paper and contents (it must be dry!!). Record.
12. Find the mass of the evaporating dish, watch glass and contents. Record.
Clean Up
1. Place salt, filter paper and sand in the trash can
2. Wash all glassware and evaporating dish using soap. Rinse thoroughly and allow to air dry.
3. Place all lab equipment back in the appropriate place as directed by your teacher and wipe off
the lab bench where you were working
Questions
1. Why is it necessary to use only pencil on filter paper?
2. Why is it necessary to use a watch glass?
3. If the sand also had iron filings in it, how could you separate the iron from the sand?
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Teacher Notes Sand Salt Lab
1. The sand salt mixture is 70/30. To make this, place 70 grams of sand in a container and
add 30 grams of salt. Make sure to shake the container before use.
2. Remind students to be very careful when heating. It cannot spatter.
3. Sample calculations:
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Date_____________ Station #______________
Name(s)________________________________________________________
Data: Sand and Salt Separation
Mass of salt/sand mixture.
Mass of filter paper
Mass of evaporating dish and watch glass
Mass of dry filter paper and contents (sand)
Mass of evaporating dish, watch glass and
contents (salt)
Calculations:
Mass of the recovered sand
Mass of the recovered salt
Ratio of the sand/salt mixture
Percent yield:
% yield = recovered mass of salt + mass of recovered sand
original mass of salt and sand mixture
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***Energy Content of Foods
Discussion:
All human activity requires “burning” food for energy. In this experiment, you will determine the
energy released (in kJ/g) as various foods, such as cashews, marshmallows, peanuts, and
popcorn, burn. You will look for patterns in the amounts of energy released during burning of the
different foods.
Figure 1
Materials
Vernier LabQuest
Temperature Probe
two food samples
food holder
wooden splint
utility clamp
2 stirring rods
ring stand and 10-cm ring
100-mL graduated cylinder
small can
Procedure
1. Find and record the initial mass of the food sample and food holder. CAUTION: Do not eat
or drink in the laboratory.
2. Determine and record the mass of an empty can. Add 50 mL of cold water to the can.
Determine and record the mass of the can and water.
3. Set up the apparatus as shown in Figure 1. Use a ring and stirring rod to suspend the can
about 2.5 cm (1 inch) above the food sample. Use a test tube clamp to suspend the
Temperature Probe in the water. The probe should not touch the bottom of the can.
Remember: The Temperature Probe must be in the water for at least 30 seconds before you
do Step 5.
4. Record the initial temperature. Remove the food sample from under the can and use a
wooden splint to light it. Quickly place the burning food sample directly under the center of
the can. Allow the water to be heated until the food sample stops burning. CAUTION: Keep
hair and clothing away from open flames.
5. Continue stirring the water until the temperature stops rising. Record this final temperature.
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6. Determine and record the final mass of the food sample and food holder.
7. Repeat the procedure for the second food sample. Use a new 50-mL portion of water.
Cleanup:
When you are done, place burned food, used matches, and partially-burned wooden splints in the
trash provided by the teacher. Pour water in the cans down the drain, make sure your lab area is
clean and dry.
PROCESSING THE DATA
1. Find the mass of water heated for each sample.
2. Find the change in temperature of the water, t, for each sample.
3. Calculate the heat absorbed by the water, q, using the equation
q = Cp•m•t
where q is heat, Cp is the specific heat capacity, m is the mass of water, and t is the change
in temperature. For water, Cp is 4.18 J/g°C. Change your final answer to kJ.
4. Find the mass (in g) of each food sample burned.
5. Use the results of Step 3 and 4 to calculate the energy content (in kJ/g) of each food sample.
6. Record your results and the results of other groups in the Class Results Table.
Questions:
1. Two of the foods in the experiment have a high fat content (peanuts and cashews) and two
have a high carbohydrate content (marshmallows and popcorn). From your results, what
generalization can you make about the relative energy content of fats and carbohydrates?
2. Which food had the highest energy content? The lowest energy content?
3. Food energy is often expressed in a unit called a Calorie. There are 4.18 kJ in 1.00 Calories.
Based on the class average for peanuts, calculate the number of Calories in a 50-g package of
peanuts.
Teacher Notes Energy Content of Foods
1. The food holder is constructed from a paperclip.
2. Sample calculations:
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Date_____________ Station #______________
Name(s)________________________________________________________
Data and calculations Energy Content of Foods
Food type
Initial mass of food and holder
–––––––––––––––––––
g
–––––––––––––––––––
g
Final mass of food and holder
g
g
Mass of food burned
g
g
Mass of can and water
g
g
Mass of empty can
g
g
Final temperature, t2
°C
°C
Initial temperature, t1
°C
°C
Temperature change, t
°C
°C
kJ
kJ
kJ/g
kJ/g
Initial mass of food and holder
–––––––––––––––––––
g
–––––––––––––––––––
g
Final mass of food and holder
g
g
Mass of food burned
g
g
Mass of can and water
g
g
Mass of empty can
g
g
Final temperature, t2
°C
°C
Initial temperature, t1
°C
°C
Temperature change, t
°C
°C
kJ
kJ
kJ/g
kJ/g
Heat, q
Energy content in kJ/g
Food type
Heat, q
Energy content in kJ/g
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Specific Heat
Discussion:
A measure of the efficiency with which a substance can store this heat energy is known as
specific heat capacity, or simply the specific heat, Cp. The greater the material's specific heat,
the more energy must be added to change its temperature.
This lab will determine the specific heat of an unknown metal using the known specific heat of
water.
1. Obtain a solid metal object from your teacher, record its physical characteristics and its
mass.
2. Place the unknown metal into a boiling water bath for 5 minutes to allow it to come to the
same temperature as the boiling water.
3. While the object is in the boiling water, place the Styrofoam calorimeter on the balance,
tare and pour approximately 100. mL of water into the calorimeter. Record the mass of
the water.
4. Record the temperature of the boiling water.
5. Record the temperature of the water in the calorimeter
6. Take the metal out of the boiling water and transfer the metal to the calorimeter.
7. Stir with a stirring rod.
8. Wait for the temperature in the calorimeter to stop changing.
9. Record the final temperature of the system (metal and water)
10. Calculate the heat gained by the water in the cup, the heat lost by the unknown metal and
the specific heat of the solid. Show all work.
11. Identify your metal using the following table
Table 1 Specific Heat Values
Material
Specific Heat (J/g•C)
Aluminum
0.900
Copper
0.385
Iron
0.448
Lead
0.130
Brass
0.385
Magnesium
1.030
Stainless steel
0.500
Tin
0.227
Zinc
3.90
Styrofoam
11.31
Air
10.06
Water
4.19
Ice
20.95
Cleanup: Water gets poured down the sink. The object gets dried off with a paper towel and
returned to the teacher. The calorimeter can sit upside down in the strainers to dry.
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Questions:
1. Which is the better calorimeter: the traditional wooden one, or the styrofoam cup? Justify
your reasoning.
2. Of the materials listed in Table 1, which is the best choice for storing solar heat energy
captured by solar cells? Why?
3. What are the advantages of using styrofoam coffee cups over aluminum ones?
4. In this lab we assumed that there was no heat lost to styrofoam calorimeter. Was this a valid
assumption? Justify your conclusion.
5. Use the fact that cw (the specific heat of water) is large to help explain the role that oceans
play on the world's climate.
6. Why are pots used for cooking often made of copper bottoms with aluminum sides? You
may need to consider the metals' specific heats, densities and price per pound
Teacher notes Specific Heat
1. The metal cylinders are stored in a lab drawer in the Chem I lab
2. It is easier to use two thermometers. One is for the boiling water bath and one for the
calorimeter.
3. Students can practice first by using a known solid and then comparing their results to a
known specific heat of the metal. Students can then calculate their percent error.
4. After practicing give students the unknown. After calculating the specific heat the
students can look at a table of specific heats (pg. 533 textbook) to identify the unknown.
5. Sample calculations:
The formula to use is:
Qlost= Qgained
Q = mC∆T
mmetal X Cmetal X (Tinitialmetal-Tfinalmetal) = mwater X 4.18J/(g K) X (Tfinalwater-Tinitialwater)
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Specific Heat
Object #
Physical Characteristics
Mass of the unknown metal object
Temperature of Boiling Water
Mass of water in the Styrofoam cup
Temperature of water in the Styrofoam cup
Final temperature of water in the
Styrofoam cup after adding the object
Change in temperature of the metal object
(show calculation)
Change in temperature of the water in the
Styrofoam cup (show calculation
Heat gained by the water in the Styrofoam
cup (show calculation)
Heat lost by the unkn)own metal (equal to
heat gained by water
Specific Heat of the unknown metal.
(show calculation)
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Calorimetry of Various Foods
Discussion:
A calorimeter can also be employed to determine how many calories or joules are present in a
given sample of food. Calories are actually kilocalories. A calorie is defined as the amount of
heat necessary to raise one gram of water one degree Celsius. In this lab you will determine the
number of calories in a food sample.
Procedure:
1. Find and record the initial mass of the food sample and food holder. CAUTION: Do not eat
or drink in the laboratory.
2. Determine and record the mass of an empty 125 mL Erlenmeyer flask. Add about 50 mL of
cold water to the flask. Determine and record the mass of the flask and water.
3. Set up the apparatus. Impale your food source on a paperclip. Record the initial temperature
of the water in the flask. Use the lighter or a wood splint to light the food sample. Place the
apparatus over the burning food quickly. Allow the water to be heated until the food sample
stops burning and the temperature stops increasing in the flask. CAUTION: Keep hair and
clothing away from open flames.
4. Repeat the procedure for the same type of food until you have 3 data readings. You can
reuse the water for each trial. Be sure to measure the beginning temperature of the water
before you begin the trial.
5. Calculate the heat gained by the water. This should equal the heat given off by the food.
Calculate the average amount of calories given off
6. Calculate how calorie dense the food sample is (calories/grams of unburned food sample)
7. Record your calorie density for your food on the board in the lab as class results.
Cleanup:
1.
2.
3.
4.
All food remnants get thrown away.
All water gets put down the sink.
The Erlenmeyer flask should be rinsed and the thermometer left at your lab station.
The cans can be left on the back table along with the pieces of Styrofoam.
Questions:
1. How many pieces of the food you tested you could eat in a day if you ate nothing else
(assume a 2000 Calorie (2,000,000 calorie) diet?
2. Based on class results, which food had the most Calories? Which was the least?
3. How could you change the lab apparatus to not allow heat to escape before it warms the
water?
Teacher notes Calorimetry of Various Foods
1. The flask/soup can apparatus is stored in a lab drawer in the Chem I room.
2. unsalted peanuts work well for this experiment.
3. sample calculations:
The formula is:
Qlost = Qgained
Q = mC∆T
Qpeanut = masswater X 1 cal/g C) X (Tfinalwater – Tinitialwater)
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Date_____________
Station #______________
Name(s)________________________________________________________
Data: Calorimetry of Foods
Food type
Trial
1
2
3
Initial mass of food
and holder
g
g
g
Final mass of food
and holder
g
g
g
Mass of food burned
g
g
g
Mass of flask and
water
g
g
g
Mass of empty flask
g
g
g
Final temperature, t2
°C
°C
°C
Initial temperature,
t1
°C
°C
°C
Temperature
change, t
°C
°C
°C
Heat, q
(show calculation)
Cal
Cal
Cal
Energy content
(show calculation)
Cal/g
Cal/g
Cal/g
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How Sweet It Is
Discussion:
Many beverages that we consume on a daily basis are very high in sugar, or sucrose, content.
While many of today’s soft drinks use fructose as a sweetener, the data collected can still be used
to determine the sugar content of a beverage. This lab allows you to determine the sugar content
of your favorite drinks by determining the density of the solution. Since the largest solute by far
is sugar, the density of the drink or soda is comparable to pure sugar solutions.
The purpose is to calculate the density of standard sucrose solutions, construct a calibration
curve from this data, and use the calibration curve to find the sugar content of beverages
Procedure:
1. Acquire a 10 mL graduated cylinder. Place it on an electronic balance and record the mass of the
empty cylinder.
2. Add exactly 10.00 mL of distilled water to the graduated cylinder using a wash bottle. Use an eye
dropper to remove any extra water. Mass the graduated cylinder with water and record the mass on
the date table.
3. Repeat these same steps for all standard sugar solutions available as well as the 3 beverages you have
selected to test. Be sure to rinse the graduated cylinder between samples and to remass the empty
graduated cylinder before each sample.
4. Calculate the density of each solution.
5. Construct a graph of density (y-axis) versus % sugar (x-axis) on the back of the data page. This is
your calibration curve.
6. Use the calibration curve to determine the sugar content of your beverages. On the calibration curve,
find the point on the “y” axis that corresponds to the density of the beverage you tested. Read the
corresponding % sugar content from the “x” axis.
Cleanup:
1. Wash the graduated cylinder with a test tube brush and a drop of soap. Return it to the cabinet.
Questions
1. Sodas today are not sweetened with sucrose but with fructose. Would that have an effect on the
outcomes of this lab? Why or why not.
2. Why is it important to use distilled water instead of tap water in step 2?
3. Would it make any difference in the density if you measured out 100.0 mL of the water and
beverages? Why or why not?
Teacher Notes How Sweet It Is
1. Make sure the solutions are fresh.
2. To make the solutions:
0%
7%
10%
12%
15%
18%
use distilled water
mix 7 grams sugar in 93ml water
mix 10 grams in 90ml water
mix 12 grams in 88ml water
mix 15 grams in 85ml water
mix 18 grams in 82 ml water
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3. The data students gather from their own drinks is not used to create the graph. Emphasize
that students make the graph using only the data from the sugar solutions. After the graph
is made, students use the graph to determine the sugar content of their drink.
4. Sample graph:
Sucrose Calibration Curve
1.2
1.18
1.16
1.14
Density
1.12
1.1
1.08
1.06
1.04
1.02
1
0.98
0
5
10
% Sugar
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Date_____________
Station #______________
Name(s)________________________________________________________
Data How Sweet It Is
Mass of empty graduated cylinder
Liquid Used
Mass
(liquid only)
________________g
Volume
Distilled Water
(0% sucrose)
10.00 mL
7% Sucrose
Solution
10.00 mL
10% Sucrose
Solution
10.00 mL
12% Sucrose
Solution
10.00 mL
15% Sucrose
Solution
10.00 mL
18% Sucrose
Solution
Beverage #1
10.00 mL
Density
(g/mL)
10.00 mL
__________
Beverage #2
10.00 mL
__________
Beverage #3
10.00 mL
__________
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***Half-Life Simulation
Discussion:
Half-life is the time required for one half of a radioactive material to decay or change into
something else. Radioactive atoms have nuclei that are unstable. These nuclei become more
stable by emitting particles or rays. The half-life of an isotope is characteristic of that isotope.
The value can be from fractions of a second to billions of years. Half-life values are constant there is no way to speed up or slow down this natural process.
In this lab, you will simulate the process of radioactive decay and determine the “half-life” for
the process.
Safety:
DO NOT EAT THE CANDIES!!! They have been handled by many students and are not
safe to eat.
Procedure:
1. Count the number of M&Ms in the bag and record on the data sheet as “Trial 0, Radioactive
Atoms.”
2. Return the M&Ms to the bag, shake to mix, and pour the M&Ms onto the platter. Remove
the M&Ms with no letters showing. These represent atoms that have decayed. Count the
numbers of “Radioactive Atoms” and “Decayed Atoms” and record the values in the data
table.
3. Return those M & Ms which have the letters showing to the bag, shake, and pour onto the
platter. Again, count the numbers of “Radioactive Atoms” and “Decayed Atoms” and record
the values in the data table. Record and continue.
4. Continue collecting data until there are no more than 5 “Radioactive Atoms” remaining.
5. Make a graph of Radioactive Atoms (y-axis) versus trial number (x-axis) and draw a smooth
curve through the points.
The value for the half-life is obtained as follows:
1. Select two values on the y-axis (not necessarily data points). One value should be half as
large as the other (60 &30 for example).
2. Draw lines from these points on the y-axis to your curve. This marks the decay of half of a
sample.
3. Next, vertical lines should be drawn from where these lines intersect your lines to the xaxis. The space between these lines on the x-axis is the half-life. This is the amount of time
(trials) required for half of the radioisotope to decay, that is, the half-life.
Questions:
1. If a rock initially contained 10 milligrams of a radioactive parent when it first crystallized,
how much remains after 4 half-lives?
2. Approximately what percentage of parent isotopes remains after 2 half-lives have passed?
3. If a mineral contains1.56% of its original parent isotopes, how many half-lives have passed?
4. Go to the website
http://www.sciencecourseware.org/VirtualDatingDemo/files/3.0_GenericCurve.html
Answer the 4 questions listed. Write each question and the answer.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data: Half Life Simulation
Trial
Radioactive Atoms
0
Decayed Atoms
XXXXXXXXXXX
1
2
3
4
5
6
7
8
9
10
Half-life computed from the graph: __________________________
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Emission Spectroscopy: Flame Lab with Chloride Salts
Discussion:
When solids are heated until they glow, their atoms produce a continuous spectrum. But
substances that are vaporized by heating in a flame can emit light characteristic of the elements
in the substance. As electrons absorb energy they are promoted to higher energy levels. This
energy is released in set amounts (quanta) as the electrons fall back into lower energy levels.
This energy is released in many regions of the electromagnetic spectrum, including the visible
region that you can see. For example, a solution of sodium chloride placed on a platinum wire
and held in a flame emits a bright, yellow light. In this lab, the student will use hydrogen gas as
a flame source to study the light given off by different metals.
Safety precautions
Wear Safety Goggles
Do not ingest any of the salts
Acids are corrosive
Tie back hair around open flames
MSDS:
SrCl2
LiCl
BaCl2
CaCl2
NaCl
KCl
3M HCl
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated.
Moderately toxic by ingestion. Irritant to body tissues. Avoid body
tissue contact.
Highly toxic by ingestion and inhalation. All soluble barium
compounds are poisonous if swallowed and cause nausea, vomiting,
stomach pains and diarrhea.
Slightly toxic by ingestion. Mild irritant to skin, eyes and mucous
membranes. Avoid all body tissue contact
Very slightly toxic by ingestion. Dust may cause minor irritation to
mucous membranes upon inhalation.
Slightly toxic by ingestion. Ingestion of large quantities can cause
weakness, GI and circulatory disturbances. Irritating to body tissues.
Avoid all body tissue content.
Toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
Procedure:
1. Obtain the 0.5 g of the following compounds one at a time from the teacher. Label five weigh
boats. Bring one at a time to the front of the room.
strontium chloride (SrCl2)
sodium chloride (NaCl)
lithium chloride (LiCl)
potassium chloride (KCl)
barium chloride (BaCl2)
Unknown #
calcium chloride (CaCl2)
2. Obtain a clean, dry, petri dish. Place 4 or 5 droppers full of 3M HCl into the dish.
3. Place any one of the compounds in the acid. Dissolve the sample by carefully stirring it with
a clean glass stirring rod.
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4. Light a burner and adjust the flame until it has no color except very pale blue. An occasional
flash of yellow is acceptable.
5. Carefully place a small piece of zinc into the petri dish. Wait until a visible reaction (slight
fizzing or bubbling) begins.
6. Now, being careful not to put any part of the burner into the acid, hold the burner by its hose
connection and place its base just above the dish.
7. Note the color imparted to the flame. The color can best be detected by looking through the
flame in front of a black background. The color may take several seconds to develop and
may only appear in flashes, so be patient. It will probably help to gently shake the dish, but
not so vigorously to spill the acid.
8. Record the observed color in the data table. Be descriptive. You will be observing other
colors that will have subtle differences in shade. For example, don’t just record “yellow”.
Instead, use “yellow-green”, “yellow-orange”, etc.
9. Using tongs, remove the zinc from the dish. Rinse it thoroughly under the tap. Keep this
piece of zinc for the next sample. Then thoroughly rinse the dish and the stirring rod under
the tap.
10. Repeat steps 3 through 10, using the same piece of zinc but a different compound each time.
Be sure to thoroughly rinse the zinc, dish, and stirring rod before using a new compound.
Cleanup
1.Clean and dry your equipment.
2.Dispose of the zinc in the waste container at the front of the room, NOT THE SINK.
Questions:
1. Explain what happens to atoms in the heat of the burner flame that causes the emission of
light. Be sure and include electrons in your explanation.
2. What is it about different atoms that cause them to emit different colors of light?
3. Look at the formulas of each of the compounds you tested. They are all made up of two
elements – one a metal and one a nonmetal. How are all the compounds alike? How are they
different?
4. Judging from your answers to questions 3 and 4 above, what component of these compounds
account for their emitted colors?
5. What color do you think would be imparted to a flame that had sodium nitrate in it? Explain
your answer.
6. What substance do you believe is in the “Unknown” and why?
Teacher Notes Emission Spectroscopy: Flame Lab with Chloride Salts
1. The amount of salt does not matter. It is easy to keep all the salts at the front desk and
distribute them yourself(fill just the tip of the scoopula more or less). Have students
get one salt at a time. Take it back to their desk and perform the experiment before
they come get the second salt.
2. Have a separate scoopula for each salt and keep the lid on the salts since a few are
hygroscopic.(take in water from the air)
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3. Use 3M HCl See the front of this manual for the method of dilution of concentrated
HCl.
4. Rinse the zinc waste with water, dry and place in the used zinc container. We can use
it again.
5. Sample data:
Lithium chloride
Strontium chloride
Barium chloride
Calcium chloride
Sodium chloride
Potassium chloride
pink-red
red
green
orange-red
deep orange
light pink
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Flame Lab with Chloride Salts
Compound
strontium chloride (SrCl2)
Observed Color
lithium chloride (LiCl)
barium chloride (BaCl2)
calcium chloride (CaCl2)
sodium chloride (NaCl)
potassium chloride (KCl)
unknown #
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***Excited Elements
When solids are heated until they glow, their atoms produce a continuous spectrum. But
substances that are vaporized by heating in a flame can emit light characteristic of the elements
in the substance. As electrons absorb energy they are promoted to higher energy levels. This
energy is released in set amounts (quanta) as the electrons fall back into lower energy levels.
This energy is released in many regions of the electromagnetic spectrum, including the visible
region that you can see. For example, a solution of sodium chloride placed on a platinum wire
and held in a flame emits a bright, yellow light. Another method of spectrum analysis involves
the application of high voltage across a gas-filled glass tube. Gases under low pressure and
excited by an electrical discharge give off light in characteristic wavelengths. The emitted light is
passed through a spectroscope, which breaks light into its constituent components for analysis. A
gas viewed through a spectroscope, such as the one shown in Figure 1, forms a series of bright
lines known as a bright-line or emission spectrum. Since each element produces a unique brightline spectrum or pattern, spectroscopy is a valuable branch of science for detecting the presence
of elements. The composition of stars and other objects in outer space is determined using this
technique.
Sodium, for example, gives off bright, yellow light that appears as two adjacent bright lines of
yellow when its gas is viewed through the spectroscope. A gas is identified by comparing the
wavelengths of its emission spectrum to the spectrum produced by a known gas. In this
experiment, you will use a spectroscope to determine the bright-line spectra characteristic of
different elements. The purpose is to observe the characteristic elemental spectra produced by
applying high voltage across a sample of a gas at very low pressure.
Safety Precautions:
 DO NOT TOUCH the spectrum-tube power supply or spectrum tubes when power is
applied. Several thousand volts exist at the power supply and spectrum tubes.
 The spectrum tubes should not be left on for more than 30 seconds at a time. The rule
of thumb is 30 seconds on/ 30 seconds off. The tubes will also get quite hot.
Procedure:
1. Obtain a spectroscope and look through it at an incandescent light bulb. The spectrum should
appear when the slit in the spectroscope is pointed just off center of the glowing filament.
Practice moving the spectroscope until you see a bright, clear image.
2. Darken the room but leave enough background lighting to illuminate the spectroscope scales.
Point the spectroscope away from any exposed window, since daylight will affect the
observed gas spectrum.
3. Helium or hydrogen is a good first choice among the spectral tubes set up around the room.
Adjust the spectroscope until the brightest image is oriented on your scale. Record on the
data sheet the location, width, and color of the brightest lines of the observed spectrum.
Some of the spectrum tubes produce light so dim that you must be very close to them to get
good observations of the spectral lines.
4. Repeat for each of the other spectrum tubes.
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5. Obtain a Labquest system. Connect the Labquest to the fiber optic cable as instructed by
your teacher. Record the emission spectrum using the Labquest system.
6. Use the printer at the front of the room to print the spectrums of your element/compound.
Write your name and the name of the compound/element on the printout and tape it up on
the wall.
7. Your teacher has printed out several spectrums of unknowns. Identify them using the
prinouts of you and your classmates.
Teacher Notes Excited Elements
1. Remind students that the spectrum tubes can only be lit for a short period of time. (30
seconds at a time (30 seconds on, 30 seconds off) helps extend the longevity.)
2. There is a set of spectrometers that was purchased with chem. $$ in the lab in a cabinet.
Please don’t let them out in the general population.
3. The power supplies are in a cabinet in the Chem Lab. There are a total of 6 of them. We
can get more from Science in Motion.
Gas Spectrum tubes we have
Air
Argon
Bromine
Chlorine-2
Helium
Iodine
Krypton
Mercury-3
Neon-2
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Nitrogen-3
Oxygen-2
Water Vapor
Xenon
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Data Excited Elements
Date_____________
Station #______________
Name(s)________________________________________________________
Element
Observations
Draw in lines at the proper locations on the scale and of the correct width and intensity to reflect
what you have observed. On the line at the right of each scale, write the name of the element you
are observing.
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Sample
7
6
5
4
Unknown Sample
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Types of Bonding in Solids
Discussion:
The purpose of this experiment is to relate certain properties of solids to the type of bonding the
solids have. These observable properties depend on the type of bonding that holds the
molecules, atoms, or ions together in each solid. Depending on the type of bonding, solids may
be described as ionic, molecular, metallic, or covalent network solids. The properties to be
studied are relative melting point, solubility in aqueous solution and electrical conductivity.
Safety
Wear lab goggles
Procedure:
1. Place 1g samples of the first substance into an evaporating dish.
2. Touch the electrodes of the conductivity tester to each solid. After each test, rinse with
distilled water and carefully dry the electrodes. Note which substances conducted electricity.
3. Place one evaporating dish on a clay triangle and heat with a burner. As soon as a solid melts
remove the flame.
4. Repeat for each substance. Do not heat any substance for more than 5 minutes. There may
be some substances that will not melt.
5. Note which substances melted and how long it took the substances to melt.
6. Place 5 test tubes in the test tube rack. Place 0.5 g of each solid into its own individual test
tube. Add 5 mL of distilled water to each test tube. Stopper and shake in an attempt to
dissolve the solid.
7. Note which substances dissolved in water.
8. Place the solutions or mixtures into separate 50 or 100 mL beakers and immerse the
electrodes of the conductivity tester. Rinse the electrodes with distilled water before and
after each test. Not which substances conduct electricity.
Cleanup
1. Dispose of solids and solutions as directed by your teacher.
2. Clean all equipment and leave to dry at your lab station.
3. Wash your hands thoroughly after cleaning up.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Type of Bonding in Solids
Substance
Conductivity of
solid
Melting
observations
Dissolution test
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Conductivity of
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***Evaporation and Intermolecular Attractions
Discussion:
In this experiment, Temperature Probes are placed in various liquids. Evaporation occurs when
the probe is removed from the liquid’s container. This evaporation is an endothermic process that
results in a temperature decrease. The magnitude of a temperature decrease is, like viscosity and
boiling temperature, related to the strength of intermolecular forces of attraction. In this
experiment, you will study temperature changes caused by the evaporation of several liquids and
relate the temperature changes to the strength of intermolecular forces of attraction. You will use
the results to predict, and then measure, the temperature change for several other liquids.
You will encounter two types of organic compounds in this experiment—alkanes and alcohols.
The two alkanes are n-pentane, C5H12, and n-hexane, C6H14. In addition to carbon and hydrogen
atoms, alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol,
C2H5OH, are two of the alcohols that we will use in this experiment. You will examine the
molecular structure of alkanes and alcohols for the presence and relative strength of two
intermolecular forces—hydrogen bonding and dispersion forces.
MSDS:
methanol
ethanol
1-propanol
1-butanol
n-pentane
n-hexane
Toxic by ingestion (may cause blindness), inhalation or absorption.
Irritating to body tissues. Avoid body tissue contact. Flammable liquid.
Toxic by ingestion and inhalation. Body tissue irritant. Avoid all body
tissue contact. Denatured with isopropanol and methanol. Not for human
consumption. Flammable liquid.
Severe eye and skin irritant. Slightly toxic by ingestion, inhalation and
skin absorption. Avoid all body contact. Flammable liquid.
Moderately toxic by inhalation or ingestion. Irritant to body tissue.
Absorbed through the skin. Avoid vapors. Flammable liquid.
Irritating to body tissues. Avoid body tissue contact. Vapor is narcotic in
high concentrations.
Irritant to body tissues. Mildly toxic by inhalation. Avoid all body
contact. Flammable liquid.
PRE-LAB EXERCISE
Prior to doing the experiment, complete the Pre-Lab table. The name and formula are given for
each compound. Draw a structural formula for a molecule of each compound. Then determine
the molecular weight of each of the molecules. Dispersion forces exist between any two
molecules, and generally increase as the molecular weight of the molecule increases. Next,
examine each molecule for the presence of hydrogen bonding. Before hydrogen bonding can
occur, a hydrogen atom must be bonded directly to an N, O, or F atom within the molecule. Tell
whether or not each molecule has hydrogen-bonding capability.
PROCEDURE
1. Setup the Vernier Labquest system with two temperature probes, one in channel one and one
in channel two. Open the method Evaporation. Identify which probe is probe 1 and 2 by
holding one in your hand and seeing which temperature increases.
2. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands.
Roll the filter paper around the probe tip in the shape of a cylinder. The paper should be
even with the probe end.
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3. The liquids are color coded in pairs. It doesn’t really matter what order you go in as long as
you stand Probe 1 in one container and Probe 2 in the other container. Make sure the
containers do not tip over.
4. After the probes have been in the liquids for at least 45 seconds, begin data collection by
pressing the ► button. Monitor the temperature for 15 seconds to establish the initial
temperature of each liquid. Then simultaneously remove the probes from the liquids and hold
them so the probe tips extend into the air.
5. When both temperatures have reached minimums and have begun to increase, press the ►
button to end data collection. Click the Statistics button, , then click to display a box for both
probes. Record the maximum (t1) and minimum (t2) values for Temperature 1 and
Temperature 2.
6. For each liquid, subtract the minimum temperature from the maximum temperature to
determine t, the temperature change during evaporation.
7. Roll the rubber band up the probe shaft and dispose of the filter paper as directed by your
teacher. Save the rubberbands; do not throw them away.
8. Repeat Steps 2-7 for the other samples.
Questions
1. Two of the liquids, n-pentane and 1-butanol, had nearly the same molecular weights, but
significantly different t values. Explain the difference in t values of these substances,
based on their intermolecular forces.
2. Which of the alcohols studied has the strongest intermolecular forces of attraction? The
weakest intermolecular forces? Explain using the results of this experiment.
3. Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker
intermolecular forces? Explain using the results of this experiment.
4. Plot a graph of t values of the four alcohols versus their respective molecular weights. Plot
molecular weight on the horizontal axis and t on the vertical axis.
Teacher notes Evaporation and Intermolecular Forces
1. Emphasize students need to hold on to the bottles of liquids while the temperature probe
is in it. Otherwise the bottles fall over.
2. Sample data can be found on the page after next.
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Date_____________ Station #______________
Name(s)________________________________________________________
PRE-LAB
TABLE A: INFORMATION ON DIFFERENT ORGANIC SOLVENTS
AND THEIR HYDROGEN BONDING CAPABILITIES
Substance
Formula
ethanol
C2H5OH
1-propanol
C3H7OH
1-butanol
C4H9OH
n-pentane
C5H12
methanol
CH3OH
n-hexane
C6H14
Structural Formulas
Molecular
Weight
Hydrogen Bond
(Yes or No)
TABLE B: MAXIMUM, MINIMUM AND CHANGE IN TEMP IN DIFFERENT
ORGANIC SOLVENTS
Substance
ethanol
1-propanol
t1
(°C)
t2
(°C)
t (t1–t2)
(°C)
Table C: Predictions of Change in
temperature of different organic solvents
Predicted
t (°C)
1-butanol
n-pentane
methanol
n-hexane
Page 58 of 146
Explanation
Bob Jones High School Chemistry I
Updated 3/23/2016
SAMPLE DATA
Substance
Formula
Structural Formulas
Molecular
Weight
Hydrogen Bond
(Yes or No)
ethanol
C2H5OH
34
Y
1-propanol
C3H7OH
60
Y
1-butanol
C4H9OH
74
Y
n-pentane
C5H12
72
N
methanol
CH3OH
32
Y
n-hexane
C6H14
86
N
TABLE B: MAXIMUM, MINIMUM AND CHANGE IN TEMP IN DIFFERENT
ORGANIC SOLVENTS
(°C)
t2
(°C)
t (t1–t2)
(°C)
ethanol
22.5
13.2
9.3
1-propanol
22.6
15.4
7.2
1-butanol
23.3
21.1
2.2
n-pentane
22.0
5.7
16.3
methanol
24.9
8.9
16
n-hexane
21.6
10.0
11.6
Substance
t1
Table C: Predictions of Change in
temperature of different organic solvents
Predicted
t (°C)
Explanation
Quick evaporation = weakbond and drop in temp. Little evaporation = strong bond
Ethanol had a weak bond compared to propanol’s strong bond formed by more carbons. (Made London
dispersion forces stronger)
Pentane had a weak bond because its weight is so similar to butanol that the lack of a hydrogen bond
makes it much weaker a bond.
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***Moisture Content of Popcorn
Discussion:
Many times when an experiment is done it is difficult to determine what is the correct or “best”
answer. When data is collected, the results of individual trials will give a range of values. In this
experiment, you will collect data for 10 samples and use basic statistics to analyze the data and
arrive at a “best” answer. The purpose of this lab is to determine the moisture content of
popcorn and to calculate the mean and standard deviation for the percent moisture loss.
Equipment/Materials:
popcorn
24 well plate lid
stirring rod
clamp
Bunsen burner
analytical balance
400mL Erlenmeyer flask
Safety:
 Goggles must be worn at all times in the laboratory.
 Care must be taken not to touch the heated flask used in this experiment.
 DO NOT eat the popcorn. Eating food in the laboratory is not permitted.
Procedure:
1. Select 20 kernels of popcorn. Place each kernel in a well of a well plate. Additional kernels
are used in case samples are burned or scorched.
2. Determine the mass of the first 10 kernels and record the mass in your data table.
3. Using a few unweighed kernels, practice popping them. A standard buret clamp makes for a
convenient handle for the flask. Heat the flask gently and keep the popcorn kernel in
constant motion. As soon as the kernel pops, invert the flask and let the popped kernel fall
onto the lab table. If it sticks, a stirring rod may be used to dislodge it.
4. When you are satisfied with your technique, begin popping the weighted samples. Return
each popped kernel to its position in the well plate. Continue until you have 10 popped
kernels that are not burned or scorched.
5. Determine the mass of each kernel after popping. Record this mass in your data table.
6. For the 10 kernels that you have selected, determine the mass lost by each kernel of and the
percent of mass lost. Record these values on the data sheet.
7. Calculate the average or mean percent mass lost. Record this value on the data sheet.
Questions:
1. What was the material that was lost? What evidence did you have for this conclusion?
2. Using the average mass of a popcorn kernel, how many kernels would you expect to find in a
one pound bag of popcorn? One pound = 454grams
3. Do you think there may be a difference in moisture content between brands of popcorn?
How might this affect the popping?
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Date_____________
Station #______________
Name(s)________________________________________________________
Pre Lab Questions Moisture Content of Popcorn
1. The mass of six grapes is measured and the grapes are placed on a shelf to dry. Thirty
days later the mass of each grape is measured to determine how much water mass is lost.
Complete the following table:
Grape
Initial
Final
1
2.4405g
1.1727g
2
2.3104g
1.1588g
3
2.5066g
1.2604g
4
2.5545g
1.2505g
5
2.0747g
1.0199g
6
2.2202g
1.1443g
Mass water
lost
% water
Massfinal
Mass Lost
Calculate mean % water lost.___________
Data and Calculations:
Trial #
Massinitial
1
2
3
4
5
6
7
8
9
10
Mean % Mass Lost __________
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% Mass
Lost
Bob Jones High School Chemistry I
Updated 3/23/2016
Sample Data:
Trial #
1
2
3
4
5
6
7
8
9
10
Mean % Mass Lost
Massinitial
Massfinal
0.1159
0.1325
0.1298
0.1384
0.1696
0.1372
0.1175
0.1821
0.1046
0.1212
0.1193
0.1264
0.1498
0.1231
0.1062
0.1394
11.2569
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Mass Lost
0.0113
0.0113
0.0105
0.012
0.0198
0.0141
0.0113
0.0427
% Mass
Lost
9.7498%
8.5283%
8.0894%
8.6705%
11.6745%
10.2770%
9.6170%
23.4487%
Bob Jones High School Chemistry I
Updated 3/23/2016
Percent Composition of Bubble Gum
Discussion:
Bubble gum is composed of many different ingredients. The five basic ingredients in bubble
gum are gum base, sugar or another sweetener, corn syrup, softener, and flavorings. The most
popular bubble and chewing gum flavors are spearmint, peppermint, cinnamon, and fruit
flavors. The gum base in bubble gum is stretchier and chewier than the base in chewing gum.
Sugar sweetens the gum, corn syrup makes it fresh and flexible, and vegetable oil keeps the
ingredients blended by retaining moisture. In this lab you will calculate the percentage of sugar
in bubble gum and calculate the moles of sugar found in bubble gum.
Procedure:
1. Unwrap your bubble gum carefully and preserve the wrapper.
2. Go to a balance. Place the open wrapper on the balance and zero the balance.
3. Find the mass of the unchewed gum on the wrapper and record the mass.
4. Take your wrapper with you and do not fold or crinkle it.
5. Chew the gum until it’s too disgusting to chew any longer (or at least 10 minutes).
6. Return to the same balance. Place your wrapper on the balance, zero it and determine the
mass of the chewed gum on the wrapper. Record the mass.
Questions:
1. Provide an explanation as to why this procedure may not be the most accurate means of
determining the percent sugar in chewing gum.
2. Create your own original percent composition problem and provide the answer.
Data
Table _____ Determination of Sugar Content of Bubble Gum
Mass of unchewed gum
g
Mass of chewed gum
g
Mass of sugar in gum.
g
Percent sugar in gum(show
calculation)
%
Formula for sucrose
Molar mass of sucrose
g/mol
Moles of sucrose in the
gum.(show calculation)
mol
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Percentage Composition of Hydrates
Discussion:
Hydrates are ionic compounds (salts) that have a definite amount of water as a part of
their structure. This water of hydration is released as vapor when the hydrate is heated.
The remaining solid is known as an anhydrous salt. The general reaction for heating a
hydrate is:
∆
hydrate → anhydrous salt + water vapor
The percent of water in a hydrate can be found experimentally by accurately determining
the mass of the hydrate and the mass of the anhydrous salt. The difference in mass is due
to the water lost by the hydrate. The percent of water in the original hydrate can be
calculated by:
mass _ of _water
%H 2O 
mass _ of _ hydrate
In this experiment, a hydrate will be heated. The change from hydrate to anhydrous salt is
accompanied by a color change for some compounds, while other compounds may
change in particle size or texture. Some changes are subtle, and the student must look
closely to observe the changes. This lab should help students to better understand the
composition of hydrates, simple decomposition reactions, and the Law of Definite
Composition.
Materials: Ring stand, Ring, Clay Triangle, Porcelain Crucible, Crucible Tongs, Analytical
Balance, Bunsen Burner, Hydrated Salt, Scoop
MSDS:
MgSO4
CuSO4
Mg
Avoid inhalation. May irritate eyes and respiratory tract. Avoid
body tissue contact.
Slightly toxic by ingestion. Body tissue irritant. Avoid all body
tissue contact.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Purpose:
The hydrate will either be MgSO4, CuSO4 or BaCl2. Be sure to record the hydrate at your table.
1. Place empty crucible on a clay triangle. An alternate method is to use an evaporating
dish and wire gauze.
2. Heat crucible with the hottest part of the flame for 3 minutes. After heating, do not
touch crucible with hands.
3. Using crucible tongs, remove the crucible from the apparatus.
4. Place on ceramic tile and allow to cool for several minutes. Carry hot crucible over
the ceramic tile when transporting.
5. Find the mass of the crucible. Record mass in the data table. (Never mass an object
when it is hot because heat waves tend to be circular and upward which tends to make
objects appear to have less mass.
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6. While the crucible is cooling, weigh approximately 2.10-2.20g of MgSO4 or 3.003.20 g CuSO4 or 7.50 – 7.70g BaCl2 of the hydrate in a weighing boat. Transfer the
hydrate to the crucible. Find and record the mass of crucible and hydrate.
7. Place crucible with hydrate on the clay triangle. Gently heat the crucible, with a low
single blue flame, by moving the burner back and forth around the bottom of the
crucible. Increase heat gradually. Do not allow the hydrate to pop or spatter.
8. Heat the crucible for 5 more minutes with a hotter single blue flame. If the edges of
the hydrate appear to be turning brown, remove the heat momentarily and resume
heating at a lower temperature.
9. Allow the crucible to cool for 2 minutes. Immediately find the mass of the crucible +
anhydrous salt and record this data. Heat again for 5 minutes, cool, find, and record
the mass of the crucible + anhydrous salt. If the mass is not exactly the same as the
mass after the first heating, heat again for 3 minutes, cool, find the mass and record.
Repeat the heating process until the last two masses are exactly the same.
10. Calculate the moles of anhydrous salt, mass of water, moles of water, and the
empirical formula of the hydrate.
11. Calculate the % water in the hydrated salt and the% error based on the theoretical
percent water in the hydrate from your teacher.
12. % error = experimental % yield - theoretical % yield
theoretical % yield
Cleanup:
1. Empty the cooled crucible into the garbage can.
2. Wipe the cooled crucible out with a dry paper towel. DO NOT WASH WITH WATER OR
SOAP. Leave it at your lab station.
Questions:
1. What could cause you to have a higher percent of water loss than theoretical (i.e. You are
losing 50% water when there is only 36% water in the compound)?
2. What could cause you to have a lower percent of water loss than theoretical (i.e. You are
losing 20% water instead of the expected 36%)?
3. Why is it important that the crucible be cooled before massing?
4. What was the purpose of the second and possibly the third heating during the experiment?
5. Name the following hydrate: Na2CO3·4H2O
6. Write the formula of the following hydrate: calcium sulfate hexahydrate
Teacher comments:
You can use MgSO4, CuSO4 or BaCl2 as the hydrate. Below are the correct formulas, amount
needed for each and percent water
MgSO4 •7H2O
2.10-2.20g
51.172%
CuSO4 •5H2O
3.00-3.20g
36.08%
BaCl2 •2H2O
7.50-7.50g
14.75%
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Percent Composition of Hydrates
Mass of Crucible (empty)
g
Mass of Crucible plus
hydrated salt
Mass of hydrated salt
g
g
Mass of Crucible plus
dehydrated salt 1st time
Mass of Crucible plus
dehydrated salt 2nd time
Mass of Crucible plus
dehydrated salt 3rd time
Mass of Crucible plus
dehydrated salt 4th time
Mass of water driven off
g
g
% water in the hydrated salt
%
Moles of water driven off
moles
Mass of dehydrated salt
g
Moles of dehydrated salt
moles
Ratio of moles of water to
moles of dehyrated salt
Empirical Formula of the
hydrated salt
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Determining an Empirical Formula
DISCUSSION:
In a sample of a compound, regardless of the size of the sample, the number of moles of one
element in the sample divided by the number of moles of another element in the sample will
form a small, whole-number ratio. These small, whole-number ratios can be used to determine
the subscripts in the empirical formula of the compound. For example, suppose that in a 24.0gram sample of a compound, there are 1.5 moles of carbon (18.0 grams of carbon) and 6 moles
of hydrogen (6.0 grams of hydrogen). When these numbers are divided by the smaller number of
moles (1.5 moles of carbon), a small, whole-number ratio of 1:4 is found.
1.5 moles of carbon = 1
6 moles of hydrogen = 4
1.5
1.5
The 1 to 4 ratio means that for every one atom of carbon in the compound there are 4 atoms of
hydrogen. The empirical formula of the compound is CH4, which is methane.
The masses of each of two elements in a compound will be experimentally determined. From this
information, a small, whole-number ratio of moles for the two elements will be calculated, and
the empirical formula of the compound will be determined.
As it heats strongly, magnesium reacts with nitrogen in the air to create nitrides and with oxygen
in the air to make magnesium oxide.. To remove the nitrides, the ash will be mixed with water to
form ammonia, NH3, leaving magnesium oxide.
MSDS
Magnesium
Flammable solid. Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
PROCEDURE:
Magnesium
1. Heat a crucible in the hottest part of a burner flame for 3 minutes without the lid on. Cool 3
minutes.
2. Measure the mass of the empty crucible and record in data table.
3. Obtain approximately 10 cm of magnesium ribbon. Observe and record the physical
characteristics of the magnesium ribbon. Fold the ribbon, place inside crucible and find the
mass of the crucible and magnesium. Record the mass of the crucible and magnesium. Before
heating, make sure the magnesium is resting on the bottom of the crucible.
4. Place crucible on a clay triangle and cover. Heat gently for 2 minutes. Using crucible tongs,
carefully tilt the cover to provide an opening for air to enter the crucible. Heat the partially
covered crucible strongly for 10 minutes.
5. Turn off the burner, remove the cover from the crucible and move the crucible from the clay
triangle to the ceramic square, cover the crucible, and allow the contents to cool for 3 minutes.
When the crucible is cool, remove the cover and examine the contents. If any magnesium has
not reacted, replace the cover at a slight tilt and heat strongly for several more minutes.
6. Turn off the burner.With a wash bottle, slowly add enough distilled water so that it is above
the ash in the crucible. Note any odor that is produced.
7. Replace the cover on the crucible so that it is slightly ajar and heat it gently so as to avoid
spattering. Continue heating until all of the water is evaporated and the ash is very light gray.
8. Measure and record the mass of the crucible and contents (without the lid). Observe and
record the physical characteristics of the newly formed compound.
Clean Up
Wipe the crucible out with a dry paper towel once it is cool. DO NOT wash the crucible or get it
wet. The ash can go in the trash can.
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Questions:
1. Osteoporosis is a disease common in older women who have not had enough calcium in their
diets. Calcium can be added to the diet by tablets that contain either calcium carbonate, calcium
sulfate or calcium phosphate. Determine the chemical formulas for these three calcium
containing compounds. Calculate which will provide the greatest percentage of calcium. Show
your work!
a) calcium carbonate:
b) calcium sulfate:
c) calcium phosphate:
2. What is the simplest formula of a compound containing 19.81g C, 2.2g H, and 77.97g Cl?
Show your work.
Teacher Notes Determining an Empirical Formula
1. When heating the magnesium students may have to lift the lid and blow air into the
crucible.
2. If the flame is too robust, it may deprive the crucible of oxygen. Try having students use
a smaller flame.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
Mass of Crucible
Mass of Crucible and Magnesium
Mass of Magnesium
Characteristics of magnesium
Moles of Magnesium
Moles of Oxygen (Same as # of
moles of magnesium)
Mass of Oxygen
Mass of Magnesium + Mass of
Oxygen = theoretical Mass of
MgO
Mass of Crucible and MgO
(Reacted magnesium)
Physical characteristics of reacted
magnesium
Actual Mass of MgO
Percent Yield (show calculation)
Page 69 of 146
Bob Jones High School Chemistry I
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Chemical Reactions: Single and Double Replacement reactions.
Discussion:
When solutions of electrolytes are mixed, a reaction may occur. The most common type is called
a double replacement. In these reactions the cation of one species reacts with the anion of the
other species to remove both ions from solution by forming a precipitate, water molecule,
molecules of a weak acid or molecules of a gas.
A second type of reaction, called a single replacement, occurs when a metal element displaces a
metal cation or hydrogen from the solution changing the ion into its elemental form of a solid or
a gas while the original element becomes the cation in solution. The purpose of the activity is to
predict the reaction and then observe the experiment to decide if the prediction was correct.
MSDS
PbCl2 solution
Moderately toxic by inhalation and ingestion. Possible carcinogen.
Avoid all body contact. Chronic exposure to inorganic lead via
inhalation or ingestion can result in accumulation in and damage to
the soft tissues and bones.
KI solution
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated.
NaCl solution
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated.
(NH4)2S solution Strong skin and mucous membrane irritant; toxic by skin absorption;
vapor harmful, even low concentrations may cause headache and
general discomfort.
FeSO4 solution
Corrosive to skin, eyes, and mucous membranes. Avoid contact with
body tissues. Solution contains sulfuric acid.
(NH4)2CO3
Substance not considered hazardous. However, not all health aspects
solution
of this substance have been thoroughly investigated.
Mg(NO3)2
Body tissue irritant. Avoid all body tissue contact.
solution
Zn(NO3)2
Slightly toxic by ingestion. Corrosive to body tissues. Avoid all body
solution
tissue contact.
CuCl2 solution
Moderately toxic by ingestion. Avoid contact with body tissues and
mucous membranes.
Mg
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated.
Cu
Irritant to body tissues as dust. Avoid contact with nitric acid, emits
toxic fumes of nitrogen oxides.
Zn
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated. Inhalation of
zinc dust may cause lung irritations. Zinc dust can spontaneously
combust when in contact with moisture.
3M HCl
Toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
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Procedure:
Use an acetate sheet for experiments 1- 6 for each experiment, mix one drop of each
substance. Record your observations in the chart. Use the solubility table to predict the
precipitate
1. Lead (II) chloride and Potassium iodide
2. Sodium chloride and Ammonium sulfide
3. Iron (II) sulfate and Ammonium carbonate
4. Lead (II) chloride and Ammonium sulfide
5. Magnesium nitrate and Potassium iodide
6. Zinc nitrate and Ammonium carbonate
For experiments 7- 12 Use the well plates. Add one small piece of each metal and cover with
the appropriate solution. Record your observations in the chart. Use the activity series to
predict the possibility of a reaction occurring
7. Magnesium and Hydrochloric acid
8. Copper and Hydrochloric acid
9. Zinc and Hydrochloric acid
10. Copper and Zinc nitrate
11. Magnesium and Copper (II) chloride
12. Zinc and Potassium iodide
13.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Chemical Reactions (Version 1)
1. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
2. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
3. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
4. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
5. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
6. Predicted Reaction______________________________________________________
Precipitate_________________________ Observation____________________________
7. Predicted Reaction______________________________________________________
Observation____________________________
8. Predicted Reaction______________________________________________________
Observation____________________________
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Bob Jones High School Chemistry I
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9. Predicted Reaction______________________________________________________
Observation____________________________
10. Predicted Reaction_____________________________________________________
Observation____________________________
11. Predicted Reaction_____________________________________________________
Observation____________________________
12. Predicted Reaction_____________________________________________________
Observation____________________________
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Types of Chemical Reactions: A Sampler Platter
DISCUSSION:
There are many kinds of chemical reactions and several ways to classify them. One useful
method is to classify reactions into five major groups. These are (1) composition or synthesis; (2)
decomposition or analysis; (3) single replacement; (4) double replacement or exchange of ions;
and (5) combustion. Not all chemical reactions can be placed into one of these categories.
In a synthesis reaction, two or more substances (elements or compounds) combine to form a
more complex substance. Equations for synthesis reactions have the general form of A + B 
AB. An example of this reaction is the formation of water from its constituent elements: 2H2(g) +
O2(g)  2H2O(l) .
A decomposition reaction is exactly the opposite of a synthesis reaction. In a decomposition
reaction, a compound breaks down into two or more simpler substances. The general form of the
equation for a decomposition reaction is AB  A + B. The breakdown of water into its
elements is an example: 2H2O(l)  2H2(g) + O2(g).
In a single replacement reaction, one substance in a compound is replaced by another, more
active element. Equations for single replacement reactions have two general· forms. In reactions
in which one metal replaces another metal, the general equation is
A + BY  A Y + B. In reactions in which a nonmetal replaces another nonmetal, the general
form is X + BY  BX + Y. The following equations illustrate each of these types of single
replacement reactions:
Zinc replaces copper (II) ion: Zn(s) + CuSO4 (aq)  ZnSO4 (aq) + Cu(s)
Chlorine replaces bromide ions: Cl2 (g) + 2 KBr(aq)  2 KCI(aq) + Br2 (g)
In a double replacement reaction, the metal ions of two different ionic compounds can be thought
of as trading places. The general equation, AB + CD  AD + CB, represents this type of
reaction. Most single and double replacement reactions take place in aqueous solutions
containing free ions. In a double replacement reaction, one of the products formed must be a
precipitate, an insoluble gas, or a molecular compound.
In a combustion reaction an element or compound is reacted with oxygen, often producing
energy in the form of heat and light. If the reaction is complete combustion, the products are
energy, carbon dioxide and water. If the reaction is incomplete, in addition to energy, carbon
dioxide and water, carbon monoxide and carbon will be produced.
SAFETY PRECAUTIONS:
1. Wear lab coat and safety goggles.
2. Be careful with flame.
3. Dispose of toxic chemicals properly.
4. Tie back long hair and loose
clothing.
5. Wash hands at the end of lab.
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MSDS:
Ethanol
CaCO3
KI solution
Pb(NO3)2 solution
Mg
Zn
6M HCl
Updated 3/23/2016
Toxic by ingestion and inhalation. Body tissue irritant. Avoid all
body tissue contact. Denatured with isopropanol and methanol. Not
for human consumption. Flammable liquid.
Irritant to body tissues. Severe eye and moderate skin irritant.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Moderately toxic by ingestion and inhalation. Possible carcinogen.
Irritating to eyes, skin and mucous membranes. Avoid all body
tissue contact. Chronic exposure to inorganic lead via inhalation or
ingestion can result in accumulation in and damage to the soft
tissues and bones.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Inhalation of zinc dust may cause lung irritations. Zinc dust can
spontaneously combust when in contact with moisture.
Toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
PROCEDURE:
For each of the following reactions, chemical changes should be observed and noted. The reactants
and products should be identified and described. Identify the type of reaction that occurs. List what
evidence is seen to indicate that a chemical reaction has occurred.
1. (Combustion) Place 6 drops of ethanol, C2H5OH, in an evaporating dish. Ignite with a burning
splint. Carefully observe what happens and note any changes in the reactant. Describe the
reaction and the color of the flame.
2. (Decomposition) Place 1 scoop of calcium carbonate, CaCO3 , in a large, clean, dry test tube.
Note the appearance of the salt.
3. Using a test tube holder, heat the CaCO3 strongly in the burner flame for about 3 minutes. Insert,
but do not drop, a burning wood splint in the test tube to test for the presence of carbon dioxide,
CO2 , gas. Carbon dioxide will extinguish the flame. Note any changes in the appearance of the
solid in the test tube.
4. (Double Replacement) Put about 1 mL of 0.1M potassium iodide, KI, solution into a small,
clean test tube. Add 1 drop of 0.1 M lead (II) nitrate solution to the same test tube. Observe what
happens and note any changes in the mixture. CAUTION: Do NOT get this all over your hands.
Wash your hands after this experiment. Dispose of the contents of this test tube in the toxic waste
container.
5. (Synthesis) Place a watch glass near the base of a burner. Examine a piece of magnesium
ribbon. Using crucible tongs, hold the magnesium ribbon in the burner flame until the
magnesium ignites. CAUTION: Do not look directly at the flame. As it burns, the Mg will
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release UV light which can be harmful to your eyes. Hold the burning magnesium away from
you and directly over the watch glass. When the ribbon stops burning put the remains in the
watch glass. Examine the product carefully and note the appearance.
6. (Single Replacment) Stand a small, clean, dry test tube in a test tube rack. Add about 1 mL of
6M hydrochloric acid, HCl (aq), to the tube. CAUTION: Handle acids with care. Acids can
cause painful burns. Carefully drop a small piece of zinc metal, Zn, into the acid in the test tube.
Observe and record what happens.
7. Using a second test tube holder, invert the mouth of a second test tube over the test tube in which
the reaction is taking place. Remove the inverted tube after about 30 seconds and quickly insert a
burning wood splint into the mouth of the tube. A pop indicates the presence of hydrogen gas.
Note the appearance of the substance in the first test tube.
8. Clean Up:
a. Empty the tube with calcium carbonate into the trash can but DO NOT wash.
b. Empty the tube with the lead nitrate in it into the waste beaker on the back bench.
c. Dispose of the magnesium ash and any wooden splints in the garbage can.
d. Empty the tube with zinc in it into the sink away from the drain. Rinse off the piece of zinc,
pick it up with your hand and put it in the zinc waste container on the back bench. Wash
both test tubes with soap and rinse. Leave the test tubes upside down in the test tube rack to
dry.
QUESTIONS:
1.
2.
3.
In this experiment, describe the method that was used to test for the presence of CO2 gas.
Describe the test used to check for the presence of hydrogen gas.
A small quantity of mercury (II) oxide is placed in a clean, dry test tube and heated for about 3
minutes. At the end of this time a glowing splint is placed in the mouth of the test tube. The
splint bursts into flames. Write a balanced equation for the reaction giving the reactants and
products. What type of reaction is this?
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Date_____________ Station #______________
Name(s)________________________________________________________
Data
Table 1
Observations of a Synthesis Reaction
Observations before reaction
Observations during/after
reaction
Proof(s) of Reaction
Balanced Equation with states of matter:
Table 2
Observations of a decomposition reaction
Observations before reaction
Observations during/after
reaction
Proof(s) of Reaction
Balanced Equation with states of matter:
Table 3
Observations of a Single Replacement Reaction
Observations before reaction
Observations during/after
reaction
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Proof(s) of Reaction
Bob Jones High School Chemistry I
Updated 3/23/2016
Balanced Equation with states of matter:
Table 4
Observations of a Double Replacement Reaction
Observations before reaction
Observations during/after
reaction
Proof(s) of Reaction
Balanced Equation with states of matter:
Table 5
Observations of a Combustion Reaction
Observations before reaction
Observations during/after
reaction
Proof(s) of Reaction
Balanced Equation with states of matter:
List all the pieces of evidence you encountered in this lab that a reaction had occurred.
_________________________________________________________________________________
_________________________________________________________________________________
______________________________________________________
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Putting Atoms Together: Synthesis of Zinc Iodide
Discussion:
In a chemical reaction, atoms often combine to form compounds. Chemical properties of the
compounds are quite different from those of the atoms of which they are composed. In this
activity you will take two elements, zinc and iodine, and combine them to form an ionic
compound, zinc iodide. When atoms combine, energy is often involved in the formation of
compounds. We can lower the amount of energy required by using a catalyst. In this activity,
water is the catalyst that initiates the reaction
MSDS:
Zn
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated. Inhalation of zinc dust may
cause lung irritations. Zinc dust can spontaneously combust when in contact
with moisture.
I2
Highly toxic by ingestion and inhalation. Irritating and corrosive to skin.
Avoid all body contact.
ZnI2
Hazardous in case of skin contact (irritant), of eye contact (irritant), of
ingestion, of inhalation.
Procedure:
1. Examine the powdered zinc and iodine crystals. Record their properties.
2. Weigh out about 0.2 g of zinc powder. Place the powder in the watch glass.
3. Weigh out the same amount of iodine crystals. Place the crystals in the watch glass also.
Mix the two by stirring with a stirring rod. Is any evidence of a reaction present?
4. Place the watch glass in the fume hood on a piece of white paper.
5. Carefully add ½ dropper of water to the mixture. Is evidence of a chemical reaction present?
6. When all of the water has been added, stir the solution with a glass stirring rod. Describe the
material.
7. Filter the solution, being careful to transfer all of the product to the filter paper. Catch the
filtrate in a clean beaker.
8. Discard the material on the filter paper and carefully pour the filtrate into a evaporating dish.
9. Place the evaporating dish on a hot plate so the water can evaporate.
10. When the liquid has evaporated, examine the product remaining in the dish. Describe this
product. What is the name of this product? How does it differ from the reactants?
11. Do not discard the product. It will be used in the next lab.
Clean Up:
1. Wash all of your equipment and dry with paper towels.
2. Leave the equipment at your station.
Reactions
A small amount of energy is required to start the reaction between zinc and iodine. Water acts as
a catalyst to permit the reaction to begin at room temperature. When the reaction begins, it
provides enough energy to melt and vaporize the iodine.
water
Zn(s)  I 2 (s) 
 ZnI 2 (s)
Some of the iodine is vaporized by the reaction. This vaporization accounts for the violet
“smoke”.
Questions:
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1. When the Zinc and Iodine are first mixed together in the beaker, how do you know that the
two substances are only forming a mixture rather than combining into a compound?
2. What did you observe occurring as water was added to the Zinc and Iodine mixture? Does
this indicate that this reaction was endothermic or exothermic as the compound, Zinc iodide,
was formed?
3. Zinc iodide is an ionic compound. This means that the elements lose or gain electrons and are
held together by the different electrical charges. When Zinc forms an elemental ion, it loses
two electrons and has a +2 charge. Iodine on the other hand gains one electron and has a -1
charge. Knowing this, what should the chemical formula of Zinc Iodide be? (Hint: a
compound must be electrically neutral)
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Taking Compounds Apart: Decomposing Zinc Iodide
Discussion:
In this activity you will decompose the zinc iodide to produce the original zinc and iodine.
Because energy is required to break this compound into its constituent parts, we will use a
battery. This process is called electrolysis and is often used to separate a compound into its
components.
MSDS:
Zn
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated. Inhalation of zinc dust may
cause lung irritations. Zinc dust can spontaneously combust when in contact
with moisture.
I2
Highly toxic by ingestion and inhalation. Irritating and corrosive to skin.
Avoid all body contact.
ZnI2
Hazardous in case of skin contact (irritant), of eye contact (irritant), of
ingestion, of inhalation.
Procedure:
1. Prepare a sample of zinc iodide according to the directions in Activity 23.
2. Using a dropper or a small pipet, add a small amount of water to the evaporating dish
containing the white precipitate. Carefully stir with a clean glass rod until all of the solid
dissolves.
3. Place two small pieces of copper wire in the solution in the watch glass. The two wires
should be about 1 in. apart. Attach them to the watch glass with a small amount of
masking tape or clay to prevent them from moving.
4. Attach the other ends of the two copper wires to the terminals of a battery, using the wire
leads.
5. Observe. Record your observations.
DO NOT THROW COPPER WIRE AWAY
Reactions:
Two reactions are occurring. One reaction occurs at the cathode. Electrons from the battery
flow into the solution at the cathode. This process produces metallic zinc, which accumulates
on the copper cathode as a zinc coating:
Zn2+(aq) + 2e- → Zn(s)
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The other electrode is the anode. Electrons return from this electrode to the battery. At this
electrode, the iodide ions lose electrons and form iodine.
Iodine crystals will be seen forming on the anode, and they will dissolve in water as well:
2I-(aq) → I2(s) + 2eThe iodine causes a water solution to turn amber.
Questions:
1. What evidence of a chemical reaction is present?
2. How do you know that you· have produced zinc and iodine?
What role did the battery play in this reaction?
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Activity Series of Metals
Discussion
The activity series of metals is a table of metals arranged in the order of their decreasing tendency to
lose electrons and enter into chemical reactions. The table of metals is arranged in the order of
decreasing electropositive character.
Although hydrogen has many physical and chemical properties that are similar to nonmetals, it
frequently functions chemically as a metal, and for this reason, it is included in the activity series of
metals. Its placement indicates that the metals preceding it will displace it from non-oxidizing acids.
The metals that are found uncombined in nature in large amounts are those that are less active than
hydrogen, whereas those metals that are more active than hydrogen are not usually found in the free
state. Two metals that are exceptions are metallic iron and nickel found in meteorites.
The normal test of the chemical activity of an element is its displacing power. If the metal can
displace another metal from a compound, it may be considered to be more chemically active than the
metal it displaces. The relative activity of a metal may be determined by observing: 1 - metal
reactivity with water: cold, warm, or hot; 2 - metal reactivity with acids: hydrogen producing acids
and non-hydrogen producing acids: 3 - metal activity with bases; and 4 - metal reactivity with salt
solutions.
It should be noted that the more finely divided the state of the metals-powdered form rather than
large lumps-the more surface area is exposed, and the greater the activity of the metal. If the reaction
is heated, the reactivity of the elements and compounds tend to increase.
MSDS:
Calcium
Magnesium
Copper
Aluminum
Lead
Zinc
CuCl2 solution
Zn(NO3)2 solution
KI solution
Corrosive solid. Avoid body tissue contact. Violent reaction with
water may evolve explosive hydrogen gas. Flammable solid.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Irritant to body tissues as dust. Avoid contact with nitric acid, emits
toxic fumes of nitrogen oxides.
Silvery odorless solid. Available as sheets, strips, foil, shots, wire,
powder and granules. Avoid inhaling dust.
Lead as a powder or dust is toxic by ingestion or inhalation. Lead
and lead compounds are possible carcinogens. Avoid ingestion and
inhalation. Emits highly toxic fumes of Pb when heated. Chronic
exposure to inorganic lead via inhalation or ingestion can result in
accumulation in and damage to the soft tissues and bones.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Inhalation of zinc dust may cause lung irritations. Zinc dust can
spontaneously combust when in contact with moisture.
Moderately toxic by ingestion. Avoid contact with body tissues and
mucous membranes.
Slightly toxic by ingestion. Corrosive to body tissues. Avoid all
body tissue contact.
Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
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Bob Jones High School Chemistry I
CuSO4 solution
AgNO3 solution
Cu(NO3)2 solution
Pb(NO3)2 solution
MgSO4 solution
6M HCl
Updated 3/23/2016
Mildly toxic by ingestion. Irritant to skin, eyes and mucous
membranes. Avoid contact with body tissues.
Moderately toxic by ingestion. Irritating to body tissues. Avoid all
body tissue contact.
Slightly toxic by ingestion. Irritating to skin, eyes and mucous
membranes. Avoid contact with body tissues.
Moderately toxic by ingestion and inhalation. Possible carcinogen.
Irritating to eyes, skin and mucous membranes. Avoid all body
tissue contact. Chronic exposure to inorganic lead via inhalation or
ingestion can result in accumulation in and damage to the soft
tissues and bones.
Body tissue irritant. Avoid contact with body tissues, especially
eyes.
Toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
Procedure
Pre-lab
For each combination, write the reactants properly and then use the activity series table in your text
or at the end of the lab manual to predict if the reaction is possible. If the reaction is not possible
write NR for the products. If the reaction is possible complete and balance the equation.
In the Lab
For each of the pairs of reactants below place a small piece of the given metal in the well plate. Add
enough of the appropriate solution to just cover the piece of metal. Record your observations below
the equation. After all combinations are placed in the well plate, allow it to sit while you complete
the other two parts of the lab. When you come back and check, use litmus paper, on the water
reactions only, to indicate if a reaction has occurred. Touch a stirring rod into the solution and touch
red litmus paper with the stirring rod. If it turns blue it is an indication of a reaction. Litmus paper
is used because it is sometimes difficult to see if a reaction with water has occurred.
Clean up:
1. After all observations are made, empty the well plates into the sink corner, rinse with water,
remove the metal to a paper towel at your lab station, wash each test tube, rinse with tap water
and do a final rinse with distilled water.
2. Turn the clean well plates upside down and leave for the next class. Rinse the metals and return
to front table.
Questions:
1. Why did you use litmus paper?
2. What was the proof of reaction between metals and acids?
3. What was the proof of reaction between metals and salt solutions?
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Activity Series Lab
REACTIONS WITH WATER (Remember to use litmus paper as evidence of a
reaction)
Calcium and water
Magnesium and water
Copper (II) and water
Aluminum and water
Lead (II) and water
Zinc and water
Get a second well plate and complete the reactions with salts
REACTIONS WITH SALTS
Magnesium and copper (II) chloride
Copper (II) and zinc nitrate
Zinc and potassium iodide
Aluminum and copper (II) sulfate
Copper (II) and silver nitrate
Lead (II) and copper (II) nitrate
Zinc and lead (II) nitrate
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Get a third well plate and complete the reactions with acids
REACTIONS WITH ACIDS
Magnesium and hydrochloric acid
Zinc and hydrochloric acid
Copper (II) and hydrochloric acid
Aluminum and hydrochloric acid
Lead (II) and hydrochloric acid
Zinc and sulfuric acid
Copper (II) and sulfuric acid
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Solubility in Double Replacement Reactions
Discussion:
In general, solubility can be thought of as the tendency of a substance (solute) to dissolve in other
substance (solvent). For qualitative purposes, such terms as "soluble", “insoluble", and
"slightly/soluble" can be used to describe these tendencies. The solubility tables at the end of the lab
manual show the tendencies of several ionic compounds to dissolve in water.
Ionic compounds (salts and bases) dissolve in water by a process known as dissociation. In this
process, the crystal lattice of the solid breaks down and free ions move throughout solution. The total
number of positive charges equals the total number of negative ions in an ionic solution.
In the following double replacement reactions, when two different aqueous (water) solutions are
mixed, one of two things may be observed: a) No reaction will occur. If all the ions remain free,
there is no reaction, and the appearance of the mixture of the ions will remain clear. b) A precipitate
(solid) will form. If two oppositely charged are attracted to one another strongly enough, they may
bond together to form an ionic compound that will not separate into ions in water. This insoluble
solid is a precipitate.
In this experiment, aqueous solutions of several different ionic compounds will be used. Different
combinations of solutions will be mixed and the results observed and noted. In those mixtures in
which reactions occur, precipitates will form. Each reactant and precipitate will be described. The
formula for each precipitate will be written. For each reaction, balanced equations and net ionic
equations will be written, and the results will compared to the results found on the solubility table.
NOTE: Since all salts are soluble to a degree and these concentrations are small there may be some
deviation from the rules.
There are several types of precipitates that may be observed in this lab. Some precipitates will
appear to be swirled. A fine powdered precipitate may look hazy or slightly cloudy. In some
reactions the precipitates will form large clumps. Crystalline precipitates may form when the
reaction is allowed to sit for a short time. The crystals have straight edges and reflect light. Another
type of precipitate is gelatinous, and the precipitate is more gel-like than solid in appearance.
MSDS:
0.1M AgNO3
0.1M BaCl2
0.1M NaOH
0.1M CuSO4
0.1M CoCl2
0.1M K3PO4
Moderately toxic by ingestion. Irritating to body tissues. Avoid all
body tissue contact. Will stain clothing and skin.
Highly toxic by ingestion and inhalation. All soluble barium
compounds are poisonous if swallowed and cause nausea, vomiting,
stomach pains and diarrhea.
Slightly toxic by ingestion and skin absorption. Severely irritating to
body tissues. Causes severe eye burns. Avoid all body tissue contact.
Mildly toxic by ingestion. Irritant to skin, eyes and mucous
membranes. Avoid contact with body tissues.
Toxic by inhalation and injestion. Severe corrosive to all body tissue,
especially skin and eyes. Avoid contact with all body tissues.
Body tissue irritant.
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0.1M Pb(NO3)2
0.1M Fe(NO3)2
0.1M BiCl3
0.1M Na2CO3
0.1M MnCl2
Updated 3/23/2016
Moderately toxic by ingestion and inhalation. Possible carcinogen.
Irritating to eyes, skin and mucous membranes. Avoid all body tissue
contact. Chronic exposure to inorganic lead via inhalation or ingestion
can result in accumulation in and damage to the soft tissues and bones.
Corrosive to body tissues by contact and inhalation. Avoid contact
with skin, eyes and mucous membranes.
Moderately toxic by ingestion and inhalation. Corrosive to body
tissues. Avoid body tissue contact.
Irritating to body tissues.
This material is generally considered nonhazardous, however not all
health aspects of this substance have been fully investigated.
PROCEDURE:
1. Before the lab, write balanced equations with states of matter for either Set A or Set B. Your lab
partner will do the other set. If no precipitate is made, write NR for the products. The equations
can be written on the back of your data sheet.
2. Obtain a set of six solutions (Set A or B). These solutions will be shared with other students
working at the same counter.
3. At each workstation there will be a sheet of white paper with a grid and black ovals on it. There
will also be a clear acetate sheet. Place the acetate sheet on top of the white paper.
4. Use the white sheet of paper as a guide.
5. For the solutions in the boxes across the top of the grid, place 1 drop of each solution in each
black oval under it. Example: AgNO3 has 5 black ovals under it, so there should be 1 drop of
AgNO3 placed in each of those 5 ovals. BaCl2 has only one oval under it. Repeat for all of the
chemicals at the top.
6. For each solution in the boxes on the left side of the grid, place 1 drop of each solution in the
black ovals to the right of the boxes. Example: BaCl2 has 4 ovals 2 beside it, so there should be 1
drop ofBaCl2 placed in each of those 4 ovals. Repeat for all of the chemicals at the left. NOTE:
Only 5 drops of each solution will be used.
7. On the data sheet describe each reactant in terms of color and clarity in the box with the formula
for the reactant.
8. On the data sheet describe each precipitate in terms of color and texture. In the same box under
the description, write the formula for each precipitate. If there is no precipitate, do not describe.
Write N.R. in the box.
Cleanup:
Wipe the solution off with a paper towel and put in the trash can
Questions
1. Which chemical reactions did you predict to produce a precipitate that actually did not?
2. Which chemical reactions did you predict to NOT produce a precipitate that actually did?
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Set A
AgNO3
BaCl2
NaOH
CuSO4
CoCl2
BaCl2
NaOH
CuSO4
CoCl2
K3PO4
Set B
Pb(NO3)2
Fe(NO3)3
NaOH
Fe(NO3)3
NaOH
BiCl3
Na2CO3
MnCl2
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BiCl3
Na2CO3
Bob Jones High School Chemistry I
Updated 3/23/2016
Stoichiometry of Copper (II) Sulfate and Iron
Discussion:
Given a balanced chemical equation and the mass of one of the substances in the equation, the
theoretical masses of all other substances in the equation can be calculated. Calculations in which a
known mass is used to find an unknown mass are called mass-mass calculations. In this experiment,
a single replacement reaction will occur. The iron will replace copper in a compound. A solid will
be separated, by filtration, from the mixture, dried, and massed. As you filter the solution, the liquid
that goes through the filter is called the filtrate. The value of the experimentally measured mass of
product will be compared to the theoretical mass predicted by a mass-mass calculation.
MSDS:
Copper (II) sulfate
Skin and respiratory irritant; moderately toxic by ingestion and
inhalation.
Iron filings
Substance not considered hazardous. However, not all health aspects
of this substance have been thoroughly investigated.
Procedure
1. Weigh out 12.50 g of copper (II) sulfate (CuSO4) into a 100 mL beaker (or use the beaker you
have nearest to 100 mL). While one lab partner does the next step (#2), the other can do step #5.
2. Add 50 mL of distilled water to the beaker. While stirring with a stirring rod, heat the beaker and
contents over a burner, using a ring stand, ring, and wire gauze square. (or use a hot plate)
3. Do not allow the solution to boil.
4. Continue to stir the mixture until the CuSO4 is completely dissolved. Note the color of the
solution. Make observation #1.
5. While your partner is doing step #2, weigh out 2.25 g of iron (Fe) filings. Make Observation #2.
Set the filings aside for later.
6. When the CuSO4 is completely dissolved, turn off the burner.
7. While the CuSO4 solution is still hot, SLOWLY sprinkle in the 2.25 g of iron filings.
8. Remove the beaker from the ring stand using beaker tongs and set it on the ceramic square on the
table.
9. Allow the mixture to cool for about 10 minutes, stirring occasionally. While waiting, answer
Questions 1 and 2.
10. After about 10 minutes, make Observation 3. Then decant off only the liquid portion of the
mixture into the sink, leaving the solid in the bottom as undisturbed as possible.
11. Determine the mass of a piece of filter paper and record it in the data table.
12. Assemble a filtration apparatus using the weighed piece of filter paper.
13. Using a spatula, scrape out as much of the solid as you can into the filter. Then, using a squeeze
bottle with distilled water, wash out the small amount of remaining solid from the beaker into the
filter, being careful never to let the liquid level in the funnel rise above the top edge of the filter
paper.
14. Complete Observation 4.
15. When all of the liquid has passed through the filter, carefully remove the paper from the funnel;
spread it out on a paper towel with your name on it to dry overnight. Then weigh the dried filter
paper with its contents and record in the data table.
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Questions
1. Write a chemical equation for the reaction taking place in the beaker between the Fe(s) and the
CuSO4 (aq).
2. What type of reaction is this?
3. What is the collected solid? Explain.
4. What is in the solution in the beaker after the reaction? Explain.
Teacher notes Stoichiometry
1. The formula weight of CuSO4 •∙ 5H2O = 249.68 g
2. Sample calculations:
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
Observation #1 Color of CuSO4
solution
Observation #2 Fe filings
Observation #3 Description of
contents of beaker after 10 minutes
Observation #4 Contents of filter
paper
Calculations: Show all your work
Mass of filter paper
Mass of filter paper and dry
collected solid
Mass of the dry collected solid.
Determine the limiting reagent and
the reagent in excess.
Calculate the percent yield.
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Stoichiometry of a Precipitate
Discussion:
Given a balanced chemical equation and the mass of one of the substances in the equation, the
theoretical masses of all other substances in the equation can be calculated. Calculations in which a
known mass is used to find an unknown mass are called mass-mass calculations. In this experiment,
a double replacement reaction will occur when two aqueous solutions are mixed. There are two
products for this reaction: an insoluble salt, which precipitates out of solution, and a soluble salt,
which remains in solution. The insoluble solid will be separated, by filtration, from the mixture,
dried, and massed. As you filter the solution, the liquid that goes through the filter is called the
filtrate. The value of the experimentally measured mass of product will be compared to the
theoretical mass predicted by a mass-mass calculation.
Safety Precautions:
MSDS
CaCl2
Slightly toxic by ingestion. Mild irritant to skin, eyes and mucous
membranes. Avoid all body tissue contact.
Na2CO3
Slightly toxic by ingestion. Irritating to body tissues. Avoid all body tissue
contact.
1M HCl
Slightly toxic by inhalation and ingestion. Severe body tissue irritant.
Corrosive to eyes. Avoid all body contact.
Procedure:
1. Using the balance, measure out approximately ___ g of _______. Your teacher will tell you
what numbers and chemicals belong in the blanks.
2. Record the exact mass measured in the data table.
3. Place the salt in a clean 250 mL beaker and add 30 mL of distilled water. Stir thoroughly to be
sure all crystals dissolve. Rinse the stirring rod. Describe solution.
4. Measure out exactly ___g of _________. Record the mass.
5. Place the second salt in a clean 150 mL beaker and add 30 mL of distilled water and stir to
dissolve. Describe solution.
6. Pour the second salt solution into the 250 mL beaker containing the first solution. Record
observations. Describe each product. Rinse the 150 mL beaker twice with distilled water using a
wash bottle. Pour each rinsing into the 250 mL beaker. Wash the 150 mL beaker with soap and
rinse with tap water and then distilled water.
7. Write your initials on a piece of filter paper in pencil, find the mass, and record this mass. Fold
the filter paper and place in the funnel. (See page 11 for instructions on filtering)
8. Place empty 150 mL beaker under funnel, and pour the mixture from the 250 mL beaker into the
filter paper. Pour slowly and do not allow the liquid to rise above the edge of the filter paper.
9. Rinse the beaker with about 5 mL portions of filtrate. Pour the rinse back through the filter until
all precipitate is transferred from the beaker to filter paper.
10. Wash the precipitate by pouring about 10 mL of distilled water through the filter. This should
remove dissolved salts from the filter paper.
11. Calculate theoretical mass of product. Your teacher should initial the calculated mass before you
leave the lab.
12. Carefully remove the filter paper and precipitate from the funnel and place on a lab cart to dry
until the next day.
13. The next day, find the mass of the dry precipitate and filter paper. Record in data table.
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14. Determine the limiting reagent and the excess reagent. Calculate the mass of the remaining
excess reagent. Show your calculation.
15. Calculate the percent yield of product. Show your calculation.
16. Calculate your percent error. Show your calculation.
Cleanup:
1. Wash the all glassware with soap and water
2. Rinse glassware in the washpan under the fume hood with 1M HCl.
3. Rinse with tap water, dry glassware and leave at your station.
Questions:
1. What would happen to the percent yield if there was still water in the precipitate?
2. If there was some precipitate left in the beaker, what effect would it have on the percent
yield?
Teacher Notes:
1. You can either use CaCl2 or SrCl2 and Na2CO3.
2. Have the students weigh out 1.50 – 1.75 g CaCl2 or 2.20-2.45 g SrCl2 and 1.50-1.75 g
Na2CO3. This will give the student approx 1.50 g of ppt.
3. The numbers are very close to being exact ratios so the limiting reactant will depend on
where in the range they picked for the mass of each salt.
4. It is necessary to clean the glassware and funnel in dilute HCl so that it gets rid of the scum
on the glass.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Gravimetric Analysis of a Precipitate
Mass of
Mass of
Balanced equation
for reaction showing
states of matter (aq,
s, g, etc)
Mass of filter paper
Mass of dry
precipitate and filter
paper
Theoretical yield of
product
Show calculations here:
Percent yield
Show calculation here:
Limiting Reagent
Excess Reagent
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Limiting Reactant Activity
Purpose:
The purpose is to develop an understanding of limiting reactant reactions.
Procedure: You will carry out the following "chemical reaction" using the contents of the plastic
bag given you by the instructor:
N + B + 2W → NBW2
1. Examine the contents of your first zip-lock bag and complete the first line of the table. Count
and record the number of nuts, bolts, and washers in the bag.
2. Use the contents of your plastic bag to assemble as many "molecules" of NBW2 as you can.
Record that number in the last column of the table.
3. Decide which "element" was the limiting reactant (N, B, or W?) Circle that number on your
data table.
4. Repeat steps 1-3 using the second plastic bag and its contents. Continue this process until you
have collected 6 sets of data and recorded them on the chart below.
Data Table 1: Modeling a Limiting Reactant Reaction Using Nuts, Bolts and Washers
N + B + 2W →NBW2
Bag #
Nuts
Bolts
Washers
NBW2 Molecules
Questions:
1.
2.
3.
4.
What is a "limiting reactant"?
Was the same element the limiting reactant in each set of data?
Use your first set of data collected to identify the element which was in excess.
When predicting the number of product molecules which reactant do you need to consider - the
one in excess or the limiting reactant?
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Stoichiometry of HCl and NaHCO3
Discussion:
In this experiment, you will measure the mass of the solid reactant, NaHCO3 and the solid product,
NaCl. The experimental determination of these relative masses will enable you to determine their
relative number of moles. As a result of your observations and calculations, you will determine the
mass and mole relationships, the reacting ratios of the solid reactant and product. From your
experimental yield, you will then determine the percent yield of NaCl.
Safety Precautions:
A hot evaporating dish and a cool evaporating dish look exactly the same.
MSDS:
NaHCO3
Slightly toxic by ingestion. Dust may be irritating to respiratory
system.
3M HCl
Toxic by inhalation and ingestion. Severe corrosive to all body
tissues, especially skin and eyes. Avoid all body contact.
Procedure:
1. Place an evaporating dish on top of a watch glass. Measure the mass of the dry evaporating dish
and the dry watch glass. Record this mass in your Data Table.
2. Add 2-3 g of NaHCO3 to the evaporating dish. Measure the mass of the NaHCO3, evaporating
dish, and watch glass. Record this mass in your Data Table.
3. Slowly add about 10 mL of 3M hydrochloric acid to the sodium hydrogen carbonate in the
evaporating dish.
4. Place the evaporating dish on a hot plate. Place the watch glass concave side up on top of the
dish, but tipped slightly so steam can escape.
5. Gently heat the evaporating dish on a low setting until only dry solid remains. Make sure no
water droplets remain on the underside of the watch glass.
6. Turn off the hot plate. Allow the apparatus to cool for at least 15 minutes. Determine the mass of
the cooled apparatus. Record the mass of the dish, residue, and watch glass in your Data Table.
Clean Up
1. Wash the evaporating dish and watch glass and dry thoroughly. Place back at your lab station.
2. Wash your hands
Questions:
1. Why is it important to let everything cool before you weight it?
2. What would happen if you didn’t let it heat until dry? How would it affect your results?
3. After it dried, what two solids would you expect to be in the evaporating dish?
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Date_____________
Station #______________
Name(s)________________________________________________________
Data Mass and Mole Relationships in a Chemical Reaction
Balanced chemical equation for the reaction.
Mass of the dry evaporating
dish and the dry watch glass
Mass of the NaHCO3,
evaporating dish, and watch
glass
Mass of the dish, residue,
and watch glass
Mass of reactant, NaHCO3
Moles of NaHCO3 reacted
Mass of product. NaCl
Moles of NaCl produced
Experimental mole ratio:
NaCl to NaHCO3
Theoretical mole ratio: NaCl
to NaHCO3
Percent yield for the NaCl
produce
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Molar Volume of a Gas
Discussion
The basis of this experiment is the reaction in which a known mass of magnesium will be reacted
with excess hydrochloric acid
Hydrogen gas is the product that is of interest in this experiment. An experimental determination of
the number of moles of hydrogen molecules produced and the volume occupied by these molecules
will be made. The number of moles of hydrogen will be calculated. The volume of hydrogen gas
produced will be measured directly on the scale of a gas measuring tube. Dalton's Law will be used
to correct for the water vapor pressure, and the ideal gas law will be used to correct this volume,
measured under laboratory conditions, to the volume the sample of gas would occupy at STP. The
corrected experimental volume of the hydrogen gas will be compared to the theoretical volume of
hydrogen gas.
MSDS
Mg
6M HCl
Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated.
Toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
PROCEDURE:
1. Obtain a piece of magnesium ribbon, with its mass, from the teacher. Record mass of
magnesium. OR your teacher may have you cut a piece of magnesium ribbon 3.2 cm long.
2. Use some steel wool or sand paper to clean the surface of the magnesium ribbon.
3. Obtain a piece of cotton thread about 15 cm long. Tie one end of the thread around the piece
of magnesium ribbon, leaving about 10 cm of thread free. Bend the piece of magnesium
ribbon so the magnesium will fit easily into the gas measuring tube.
4. Place about 900 mL of room-temperature tap water in a 1 L beaker. Set up a ring stand and
double buret clamp, and place beaker in position.
5. Obtain about 10 mL of 3M or 6M HCl (aq). CAUTION: Handle this acid with care.
Carefully pour the acid into a gas measuring tube.
6. Tilt the gas measuring tube slightly. Using a beaker slowly fill the tube to the top with roomtemperature tap water. Try to avoid mixing the acid and water as much as possible.
7. Lower the piece of magnesium ribbon 4 or 5 cm into the gas measuring tube. Drape the
thread over the edge of the tube and insert a one-hole stopper. Tube must be completely filled
with water.
8. Place finger over the hole in the rubber stopper and invert the gas measuring tube. Lower the
stoppered end of the tube into the beaker of water. Clamp the tube in place so that the
stoppered end is well under water.
9. Describe the reaction.
10. Let the apparatus stand about 5 minutes after the magnesium has completely reacted. Then,
tap the sides of the gas measuring tube to dislodge any bubbles that may have become
attached to the side of the tube.
11. Move the tube up or down, to equalize pressure, until the water level in the tube is the same
as that in the beaker. On the scale of the gas measuring tube, read the volume of the gas in
the tube. Record the volume of gas.
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12. Record the room temperature and barometric pressure.
13. Write a balanced equation for the reaction that occurred.
14. Calculate stoichiometrically how many moles of hydrogen would be made from the mass of
the magnesium. Show your work
15. Calculate the corrected pressure of the gas inside of the tube by subtracting out the partial
pressure of the water. Show your work.
16. Calculate the moles of gas inside the tube. Show your work.
17. Calculate a percent error for moles of hydrogen.
Cleanup:
1. Rinse the gas collecting tube with tap water and then rinse with distilled water from a wash
bottle. Hang upside down in your double buret clamp to dry
2. Wash all beakers and Erlenmeyers with soap and rinse and place on a paper towel at your
station to dry.
Questions:
1. Why was it necessary to clean the surface of the magnesium?
2. For the following possible errors, give the direction they would cause your results to deviate
from the accepted value of 22.4 L/mol at STP. Explain your answers.
a. air bubbles were introduced when tube is turned over in step 7
b. not all the Mg ribbon reacted
c. not all the MgO coating is removed from the Mg ribbon before beginning the
experiment
d. you read the volume when the water level in the tube was higher than outside the tube
in step 9
3. Find the volume of 80 g O2 at the conditions in the lab today.
4. How many liters would 0.25 mol of any gas occupy at pressure in the room but 10 degrees
higher?
Teacher Notes:
3.2 cm of magnesium ribbon is approximately 0.0451 g Mg and would give 45 mL of H2 at 1 atm of
pressure and 22 ºC.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
Balanced Equation:
Table 1
Mass of magnesium
Description of reaction
Gas volume in tube
Room temperature
Atmospheric pressure
Partial pressure of water
Pressure exerted by H2
at room conditions
Experimental moles of
H2
(show calculation)
Expected moles of
hydrogen gas
(show calculation)
Percent Error
(show calculation)
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“Wet” Dry Ice
Discussion
The phase diagram for carbon dioxide shows that CO2 can exist only as a gas at ordinary room
temperature and pressure. To observe the transition of solid CO2 to liquid CO2, you must increase
the pressure until it is at or above the triple point pressure, which is labeled X in the diagram below.
Objectives:
1.
Interpret a phase diagram.
2.
Observe the melting of carbon dioxide due to varying pressure.
3.
Relate observations of carbon dioxide to its phase diagram.
Safety:
Do not touch the dry ice with your hands. Wear goggles at all times.
Procedure:
1. Use forceps or tongs to place 2-3 very small pieces of dry ice on the lab bench, and observe them
until they have completely sublimed.
2. Fill a plastic cup with tap water to a depth of 4-5 cm.
3. Cut the tapered end (tip) off the graduated pipet.
4. Use forceps to carefully slide 8-10 pieces of dry ice down the stem and into the bulb of the pipet.
5. Use a pair of pliers to clamp the opening of the pipet stem securely shut so that no gas can
escape. Use the pliers to hold the tube and to lower the pipet into the cup just until the bulb is
submerged. From the side of the cup, observe the behavior of the dry ice.
6. As soon as the dry ice has begun to melt, quickly loosen the pliers while still holding the bulb in
the water (note: if you wait too long to loosen the pliers. the pipet may burst due to increasing
pressure). Observe the CO2•
7. Tighten the pliers again and observe.
8. Repeat steps 6 and 7 as many times as possible.
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***Properties of Solutions: Electrolytes and Non-Electrolytes
In this experiment, you will discover some properties of strong electrolytes, weak .electrolytes, and
non-electrolytes by observing the behavior of these substances in aqueous solutions. You will
determine these properties using a Conductivity Probe. When the probe is placed in a solution that
contains ions, and thus has the ability to conduct electricity, an electrical circuit is completed across
the electrodes that are located on either side of the hole near the bottom of the probe body (see
Figure 1). This results in a conductivity value that can be read by the computer. The unit of
conductivity used in this experiment is the microsiemens, or S.
The size of the conductivity value depends on the ability of the aqueous solution to conduct
electricity. Strong electrolytes produce large numbers of ions, which results in high conductivity
values. Weak electrolytes result in low conductivity, and non-electrolytes should result in no
conductivity. In this experiment, you will observe several factors that determine whether or not a
solution conducts, and if so, the relative magnitude of the conductivity. Thus, this simple experiment
allows you to learn a great deal about different compounds and their resulting solutions.
In each part of the experiment, you will be observing a different property of electrolytes. Keep in
mind that you will be encountering three types of compounds and aqueous solutions:
Ionic Compounds
These are usually strong electrolytes and can be expected to 100% dissociate in aqueous solution.
Example: NaNO3(S) → Na+(aq) + NO3-(aq)
Molecular Compounds
These are usually non-electrolytes. They do not dissociate to form ions. Resulting solutions do not
conduct electricity.
Example: CH3OH(l) → CH3OH(aq)
Molecular Acids
These are molecules that can partially or wholly dissociate depending on their strength.
Example: Strong electrolyte H2SO4 →H+(aq) + HSO4-(aq) (100% dissociation)
Example: Weak electrolyte HF ↔ H+(aq) + F-(aq) (<100% dissociation)
MATERIALS
Vernier Conductivity Probe
Vernier Labquest system
250-mL beaker
wash bottle with distilled water
ring stand
utility clamp
H2O (tap)
H2O (distilled)
0.05 M NaCl
Safety Precautions:
MSDS
0.05 M NaCl, 0.05 M CaCl2, 0.05 M
0.05 M CaCl2
0.05 M AlCl3
0.05 M HC2H3O2
0.05 M H3PO4
0.05 M H3BO3
0.05 M HCl
0.05 M CH3OH (methanol)
0.05 M C2H6O2 (ethylene glycol)
Substance not considered hazardous. However,
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Bob Jones High School Chemistry I
AlCl3, 0.05 M HC2H3O2, 0.05 M H3BO3
0.05 M H3PO4
0.05 M HCl
0.05 M CH3OH (methanol)
0.05 M C2H6O2 (ethylene glycol)
Updated 3/23/2016
not all health aspects of this substance have
been thoroughly investigated.
Slightly toxic by ingestion and inhalation;
body tissue irritant.
Slightly toxic by inhalation and ingestion.
Severe body tissue irritant. Corrosive to eyes.
Avoid all body contact.
Toxic by ingestion (may cause blindness),
inhalation or absorption.
Mildly toxic by ingestion and inhalation. Skin,
eye, and mucous membrane irritant. Avoid all
body tissue contact.
PROCEDURE
1. Obtain and wear goggles! CAUTION: Handle the solutions in this experiment with care. Do not
allow them to contact your skin. Notify your teacher in the event of an accident.
2. Open the Conductivity lab method on the lab quest. The window will display live conductivity
readings in units of microsiemens (S).
3. The Conductivity Probe is already attached to the labquest. It should be set on the 0-20,000 S
position.
4. Obtain the Group A solution containers. The solutions are: 0.05 M NaCl, 0.05 M CaCl2, and 0.05
M AlCl3.
5.
Measure the conductivity for each of the solutions.
a. Carefully raise each vial and its contents up around the Conductivity Probe until the hole near
the probe end is completely submerged in the solution being tested. Important: Since the two
electrodes are positioned on either side of the hole this part of the probe must be completely
submerged.
b. Briefly swirl the beaker contents. Once the conductivity reading in the Meter window has
stabilized record the value in your data table.
c. Before testing the next solution, clean the electrodes by surrounding them with a 250-mL
beaker and rinse them with distilled water from a wash bottle. Blot the outside of the probe end
dry using a tissue. It is not necessary to dry the inside of the hole near the probe end.
6. Obtain the four Group B solution containers. These include 0.05 M H3PO4, 0.05 M HC2H3O2,
0.05 M H3BO3, and 0.05 M HCl. Repeat the Step 5 procedure.
7. Obtain the five Group C solutions or liquids. These include 0.05 M CH3OH, 0.05 M C2H6O2,
distilled H2O, and tap H2O. Repeat Step 5.
Cleanup:
1. Rinse off the conductivity probe with distilled water, wipe up any spills and thrown any trash
away.
2. Depending on what class you are in, you might be putting the equipment away for the day. Ask
your teacher.
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Date_____________ Station #______________
Name(s)________________________________________________________
Data
Solution
Conductivity (µS)
A – CaCl2
A – AlCl3
A – NaCl
B - HC2H3O2
B – HCl
B - H3PO4
B - H3BO3
C - H2Odistilled
C - H2Otap
C - CH3OH
C - C2H6O2
PROCESSING THE DATA
1. Based on your conductivity values, do the Group A compounds appear to be molecular,
ionic, or molecular acids? Would you expect them to partially dissociate, completely
dissociate, or not dissociate at all?
2. Why do the Group A compounds, each with the same concentration (0.05 M), have such
large differences in conductivity values? Hint: Write an equation for the dissociation of each.
Explain.
3. In Group B, do all four compounds appear to be molecular, ionic, or molecular acids?
Classify each as a strong or weak electrolyte, and arrange them from the strongest to the
weakest, based on conductivity values.
4. Write an equation for the dissociation of each of the compounds in Group B. Use 
 for
 for weak.
strong; 
5. For H3PO4 and H3BO3, does the subscript “3” of hydrogen in these two formulas seem to
result in additional ions in solution as it did in Group A? Explain.
6. In Group C, do all four compounds appear to be molecular, ionic, or molecular acids? Based
on this answer, would you expect them to dissociate?
7. How do you explain the relatively high conductivity of tap water compared to a low or zero
conductivity for distilled water?
8. Did aqueous methanol, CH3OH, have the same conductivity value as aqueous ethylene
glycol, C2H6O2? Explain.
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Acid/Base Indicators
Discussion:
Acids and bases surround you in your everyday life. Fortunately, there are ways to discover whether
something is acidic or basic. You can buy kits for a pool, pH meters for the chemistry lab or make
your own at home. It turns out red cabbage juice makes for a pretty good acid/base indicator. You
will use supporting evidence to predict if common household substances are acids or bases , determine the
pH of the substances and create a calibrated pH scale using cabbage juice as an indicator.
MSDS:
Universal
Contains denatured ethyl alcohol. Moderately toxic by ingestion and
indicator
inhalation. Body tissue irritant. Avoid all body tissue contact. Flammable
liquid.
Ammonia
Liquid and vapor are strongly irritating to skin, eyes, and mucous
membranes. Vapor extremely irritating to eyes. May cause blindness.
Toxic by ingestion or inhalation.
Each Group Will Need:
 litmus paper (both red and blue) rip each strip of litmus paper into thirds
 universal indicator paper and a scale
 pH meter
 distilled water
 5 baby weigh boats; reuse- clean and dry as needed
 Household substances: dishwashing liquid, dishwashing detergent, laundry detergent, laundry
stain remover, fabric softener, bleach, mayonnaise, baking powder, baking soda, white vinegar,
cider vinegar, lemon juice, soft drinks, mineral water, milk
 Fresh red cabbage
 Hot plate
 250 ml beaker
Procedure:
1. Prepare the cabbage water indicator:
Tear a few leaves of cabbage and place them in the 250 mL beaker. The beaker should be about
½ full. Then add distilled water until the cabbage leaves are covered. When the mixture comes to
a boil, turn off and unplug the hotplate.
2. Pick up materials listed above. Get five at a time in the tiny weigh boats. If the item is a solid add
a few drops of distilled water to make it soupy before testing with the indicator papers. Clean and
dry the weigh boats and then get five more substances. Repeat until you have tested all the
substances listed.
3. Predict if each substance is an acid, base, or neutral.
4. Test each substance with the red and blue litmus paper. Record the color change, if any, which
occurs.
5. Test each substance with the universal indicator paper. Record the color that you see, and record
a number from the universal indicator scale.
6. Use the pH meter and record the number indicated.
7. Lastly, add a few drops of the cabbage water indicator to the substance in the weigh boat. Record
the color. Be specific.
Questions:
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1. Create your own pH scale for cabbage water. Use the data from the table and make a colored
scale. See pg 512 of your textbook for an example. (You will have the same numbered scale as
the one on pg. 512, but you will have different colors.)
2. Are cleaning products generally acidic, basic or neutral? What acid, base or neutral
characteristics make it suitable for cleaning?
3. What are acid/base characteristics of foods and beverages?
4. What would you have to do to make your cabbage pH scale a more useful scale?
5. What conflicts of information showed up in your data table? For instance did one indicator say it
was a base and the other type of indicator say you had an acid?
6. What further testing could you do to clear up some of the conflicting data?
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
Material
Prediction
Litmus paper
Universal paper
Page 108 of 146
Acidic, Basic or
Neutral
Bob Jones High School Chemistry I
Updated 3/23/2016
39 Drop pH Lab
Discussion:
The pH of a solution can be determined chemically or instrumentally. The chemical method
uses an indicator, a substance that changes color in some know pH range. A Universal Indicator
will produce an array of colors (red, orange, yellow, green, blue, purple, and colors in between)
depending on the acidity of the solution being tested. The purpose of this experiment is to produce a
standard pH color chart to be used in determining the pH of several household products. You will
also test the conductivity to evaluate whether they are non-electrolytes, strong electrolytes or weak
electrolytes.
Materials/Equipment:
24 well plate
ammonia
bleach
baking soda, muratic acid, sprite, club soda, sugar,
tap water, distilled water, lemon juice, milk, saliva,
egg white, soap, shampoo,
wooden splints, conductivity sensors
Dropper bottles of
solution X
solution Y
Universal Indicator
Safety Considerations:
 Always wear goggles in the lab.
 Carefully wash hands after handling solutions.
MSDS:
Universal Indicator Contains denatured ethyl alcohol. Moderately toxic by ingestion and
inhalation. Body tissue irritant. Avoid all body tissue contact.
Flammable liquid.
0.1 M HCl
Slightly toxic by inhalation and ingestion. Severe body tissue irritant.
Corrosive to eyes. Avoid all body contact.
Procedure:
1.
2.
3.
4.
5.
6.
Place the well plate on a piece of clean white paper.
Add drops of solution X and solution Y to the wells in rows A and B as indicated by the chart
on the next page. Stir solutions with a wooden splint. Be careful to control drops—DON’T
miscount!
Add about 5 drops of each household solution to separate wells in row C & D.
Using a conductivity sensor, test each household product and record the solubility as Very High,
High, Medium, Low, or None. Rinse off the conductivity sensor probe wires with a wash bottle
between each product.
Add one drop of Universal Indicator to each well containing a solution. Stir with a wooden
splint.
Determine the pH values of the household solutions by comparison with the other wells.
Cleanup:
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Rinse the conductivity sensors with a paper towel and dry with a paper towel. Rinse the well plate
with water and then with distilled water. Dry with a paper towel.
Questions:
1. Which household solution was the most acidic?
2. Which household solution was the most alkaline?
3. Was the pH of the cleaners acidic, basic or neutral?
What do you think solution X and Y are?
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
A1
A2
Well #
2.0
pH (±0.1)
empty
drops X
drops Y
B1
Well #
7.0
pH (±0.1)
drops X
drops Y
Well #
pH (±0.1)
A3
3.0
39
0
B2
8.0
20
19
C1
____
A4
4.0
35
4
B3
9.0
17
22
C2
____
A5
5.0
31
8
B4
10.0
14
25
C3
____
A6
6.0
27
12
B5
11.0
11
28
C4
____
24
15
B6
12.0
9
30
C5
____
3
36
C6
____
color
substance
_______
________
________
________
________
________
conductivity
Well #
pH (±0.1)
D1
____
D2
____
D3
____
D4
____
D5
____
D6
____
________
________
________
color
substance
________
________ ________
conductivity
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Titration Lab
Discussion:
A titration is a method of analysis that will allow you to determine the precise endpoint of a reaction
and therefore the precise quantity of reactant in the titration flask. A buret is used to deliver the
second reactant to the flask and an indicator or pH meter is used to detect the endpoint of the
reaction.
MSDS
0.1M NaOH
Moderately toxic by ingestion and skin absorption. Corrosive to body
tissues. Causes severe eye burns. Avoid all body tissue contact.
HCl
Slightly toxic by inhalation and ingestion. Severe body tissue irritant.
Corrosive to eyes. Avoid all body contact.
Phenolphthalein Contains denatured ethyl alcohol. Toxic by ingestion and inhalation.
Irritating to body tissues. Avoid all body tissue contact. Flammable liquid.
1. Read the section on titrations in your book in chapter 15.
2. You will only use one buret for the base. You will use a small graduated cylinder and an eye
dropper to measure out the acid. The acid will be hydrochloric acid and the base will be 1.000 M
NaOH. Your indicator will be phenolphthalein.
3. Get 10.0 mL of 0.10 M HCl and titrate it using base to practice your titrating skills. Your end
point should be very close to 10.00 mL of base.
4. Build a data table. You are running 3 trials for your unknown acid. You must show an initial
and final volume for the base for all 3 trials. Be sure to include the unknown #. Your volume of
acid will always be 10.0 mL.
5. You are to calculate the average mL of base used to titrate 10.0 mL of unknown acid and use this
average to calculate the molarity of the unknown acid using the equation MA x VA = MB x VB
where MA is the molarity of the acid, VA is the volume of the acid MB is the molarity of the base
and VB is the volume of the base. Be sure to show your calculation.
6. Your teacher may have you repeat this titration for more than one unknown acid.
Cleanup:
1. Pour the remainder of the base back into the beaker and pour back into the volumetric flask
under the fume hood.
2. Pour the remainder of the acid down the sink.
3. Pour the contents of the Erlenmeyer down the sink.
4. Wash the burets with a tiny drop of soap and the buret brush. Rinse with tap water twice and
then rinse with distilled water from a wash bottle. Hang upside down in your double buret clamp
with the valve open.
5. Wash all beakers and Erlenmeyers with soap and rinse and put on a paper towel at your station to
dry.
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Teacher Notes Titration
Remind students to be sure the stopcock is closed on the burets before adding the base.
Some students forget to add the phenothalein before beginning the titration.
The concentrations of unknown acids:
#1
0.0500 M
use 8.33 mL 3M HCl to make solution
#2
0.10 M
use 16.67 mL
#3
0.20 M
use 33.30 mL
#4
0.24 M
use 40.00 mL
#5
0.15 M
use 25.00 mL
Add the indicated amount of 3 M HCl to water to 500 mL volumetric flask and dilute to mark.
Each solution is 500. mL.
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Date_____________
Station #______________
Name(s)________________________________________________________
Data
Unknown Acid #
Trial Initial
Final
volume
volume
1
Volume
NaOH used
Molarity of
unknown acid
Calculation
Volume
NaOH used
Molarity of
unknown acid
Calculation
Volume
NaOH used
Molarity of
unknown acid
Calculation
2
3
Unknown Acid #
Trial Initial
Final
volume
volume
1
2
3
Unknown Acid #
Trial Initial
Final
volume
volume
1
2
3
Page 114 of 146
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Qualitative Analysis of the Group I Cations
Discussion
The Group I cations are those species that form chloride precipitates that are insoluble in acid. The
group includes Ag+, Pb2+, and Hg22+ (mercurous ion). These cations are precipitated from any other
cations that might be present in a sample by addition of 6 M hydrochloric acid: Addition of HCl
forms a mixture of AgCl, PbCl2, and Hg2Cl2 solids.
Ag+(aq) + Cl-(aq)  AgCl(s)
Pb2+(aq) + 2Cl-(aq)  PbCl2(s)
Hg+(aq) + 2Cl-(aq)  Hg2Cl2(s)
At this point, the sample is centrifuged and the precipitate of the Group I chlorides isolated
(in a real analysis, the decantate from above the precipitate is saved for further analysis of the cations
of the other groups). After you centrifuge a sample, there is a solid at the bottom called a residue
and a liquid called the supernate. When you pour the supernate off it becomes the decantate because
you have poured it out of the test tube.
Lead ion is then separated from silver and mercury by taking advantage of the fact that PbCl2
is much more soluble in hot water than in cold water (the solubilities of AgCl and Hg2Cl2 do not
vary much with temperature). Distilled water is added to the Group I mixed precipitate, and the
mixture is heated to dissolve PbCl2. The mixture is then centrifuged quickly while still hot, and the
decantate containing lead ion is removed from the remaining silver/mercury precipitate. The
presence of lead ion is then confirmed by addition of dichromate ion, Cr2O72-, which forms a
characteristic yellow precipitate with lead ion, PbCr2O7.
The precipitate containing silver and mercurous ions is then treated with ammonium
hydroxide. Silver ion is complexed by ammonia; the precipitate of AgCl will dissolve and is
removed after centrifugation. A black-gray residue in the centrifuge tube confirms the presence of
mercury.
The decantate containing complexed silver ion is then treated with acid, which reacts with
ammonia, allowing the precipitation of silver chloride. Alternatively, potassium iodide can be added,
which also precipitates the silver (as a creamy yellow-white solid).
In this experiment a known sample containing all three ions, as well as an unknown sample
(containing one or more cations from the specific group), will be analyzed. In real practice, a sample
would not be restricted to the members of one analysis group but, rather, would be a general mixture
of all possible cations.
Safety Precautions
 Protective eyewear approved by your institution must be worn at all times while you are in the
laboratory.
 The centrifuge spins rapidly and can eject the sample tubes with considerable momentum if it is
not correctly balanced. If the centrifuge starts to wobble or creep along the lab bench, immediately
disconnect power and balance the centrifuge.
MSDS
6M HCl
Toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
K2Cr2O7
Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
Page 115 of 146
Bob Jones High School Chemistry I
solution
6M
NH4OH
6M HNO3
AgNO3
HgNO3
Updated 3/23/2016
tissue contact.
Liquid and vapor are strongly irritating to skin, eyes, and mucous
membranes. Vapor extremely irritating to eyes. May cause blindness.
Toxic by ingestion or inhalation. When heated to decomposition, emits
toxic fumes of NH3 and NOx.
Corrosive; will cause severe damage to eyes, skin and mucous membranes.
Moderately toxic by ingestion and inhalation. Strong oxidizer. Avoid
contact with acetic acid and readily oxidized substances.
Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact. Silver compounds photoexpose and will stain skin and
clothing.
Highly toxic by ingestion and inhalation. Severely corrosive to body
tissues. Avoid all body tissue contact.
Procedure
Label all test tubes during the procedure to make certain that samples are not confused or discarded
at the wrong point. All glassware should be cleaned, rinsed and rinsed with distilled water before
starting the lab.
Take notes as you perform the tests on the known so you can evaluate the unknown.
Step
Notes
1.
Transfer approximately 1 mL of the
sample solution to a test tube, and then add
10 drops of 6 M HCl to precipitate the
Group I cations.
2.
Stir the mixture to mix, and then
centrifuge for one minute until the
precipitate is packed firmly in the bottom
of the tube (be sure to balance the
centrifuge).
3.
To ensure that sufficient HCl has
been added to the sample to cause the
precipitation of all the Group I cations, add
1 more drop of HCl to the test tube,
watching the supernatant liquid in the test
tube for the appearance of more precipitate.
4.
If more precipitate forms on the
addition of 1 drop of HCl, recentrifuge and
retest with additional single drops of HCl
until it is certain that all the Group I
chlorides have been precipitated.
Recentrifuge until the precipitate is packed
firmly in the bottom of the centrifuge tube.
Discard the decantate.
5.
Wash the silver/mercury/lead
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precipitate in the centrifuge tube by adding
1 mL of distilled water and 2 drops of 6 M
HCl. Stir with a glass rod and centrifuge.
6.
Remove the wash water and
discard. Wash a second time, centrifuge,
and again discard the wash liquid. The
precipitate now contains the chlorides of
silver, mercury (I), and lead, separated
from all other species.
7.
To the mixed chloride precipitate
add 2-3 mL of distilled hot water.
8.
Heat the mixture in a boiling-water
bath for 3-4 minutes. Stir the precipitate
during the heating period to dissolve the
lead chloride.
9.
While the mixture is still hot,
centrifuge and remove the decantate (which
contains most of the lead ion from the
sample) to a separate test tube.
10.
To the decantate (containing lead
ion),add 1 drop of potassium dichromate
solution (see teacher for solution). Allow
the sample to stand for a few minutes. The
appearance of a yellow precipitate of lead
chromate confirms the presence of lead ion
in the original sample.
11.
If you are doing the know solution
continue on to step 14. Otherwise, to the
remaining Group I precipitate, add 2-3 mL
of hot water and 1-2 drops of 6M HCl. Stir,
and heat in the boiling-water bath for 3-4
minutes. Centrifuge the mixture.
12.
Carefully remove and discard the
supernatant liquid, which may contain lead
ion that was not completely removed
earlier. Add a 2-3-mL portion of hot water
and 1-2 drops of 6M HCl.
13.
Stir the mixture, and heat in the hotPage 117 of 146
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water bath. Centrifuge, remove, and discard
the supernatant liquid. These several
washing steps are necessary to remove all
traces of lead ion from the silver/mercury
precipitate.
14.
Add 2-3 mL of 6M aqueous
ammonium hydroxide to the silver/mercury
precipitate.
15.
Stir the mixture, centrifuge, and
remove the decantate (which contains
dissolved silver ion). The presence of a
gray-black precipitate in the centrifuge tube
at this point confirms the presence of
mercury (I) in the original sample.
16.
Add sufficient 6 M nitric acid to the
decantate to make the solution acidic when
tested with litmus paper. The appearance of
a white precipitate of AgCl confirms the
presence of silver ion in the original
sample.
Cleanup: All precipitates go in the waste beaker in the fume hood. Wash all test tubes with soap
and water, rinse and rinse with distilled water. Leave test tubes upside down in the test tube rack to
dry.
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Unknown #__________
Station#__________________
Name(s)________________________________________________________
Presence of Lead ion
Yes
No
Presence of Silver ion
Yes
No
Presence of Mercury ion
Yes
No
Group I Cation Analysis
Pb2+, Ag+, Hg22+
HCl(aq)
AgCl, PbCl2, Hg2Cl2
cold water wash
hot water wash
Residue 2
Decantate 2
Pb2+
AgCl, Hg2Cl2
NH4OH(aq)
K2Cr2O7
Decantate 3
HgNH2Cl
White and
Hg black
Ag(NH3)2+
HNO3(aq)
AgCl
white
Page 119 of 146
PbCr2O7
yellow
Bob Jones High School Chemistry I
Updated 3/23/2016
Copper into Gold: The Alchemist’s Dream
Discussion:
One of the goals of the ancient alchemists was to convert base metals into gold. Although this goal
was never attained by chemical methods, the alchemists were able to perform many color changes to
make metals resemble gold. In this experiment you will produce some color changes to a copper
token and demonstrate diffusion in the solid state.
In this reaction, zinc dissolves in the hot concentrated sodium hydroxide solution to form sodium
zincate, commonly written as Na2ZnO2 or, as obtained in solid form from concentrated solutions,
NaZn(OH)3. As an ionic equation this can be written:
Zn + 2 OH− → ZnO22- + H2
When the copper token is added to the solution, an electrochemical couple formed by the copperzinc contact causes the zincate ion to migrate to the copper surface where it is decomposed and
reduced to metallic zinc by hydrogen which forms a coating on the token. The resulting token will be
silver in color due to a coating of zinc on its surface.
When the token is heated, the zinc diffuses into the copper to form a layer of the alloy brass, which
results in the gold color. It should be noted that the reduction of the zincate ion to zinc will only take
place if the copper metal is in direct contact with zinc metal. Also, no copper dissolves in the
solution during the reaction.
MSDS:
6M NaOH
Zinc dust
Moderately toxic by ingestion and skin absorption. Corrosive to body
tissues. Causes severe eye burns. Avoid all body tissue contact.
Inhalation of zinc dust may cause lung irritations. Zinc dust can
spontaneously combust when in contact with moisture.
Procedure
1. Obtain a shiny penny from your teacher. It must be nice and clean. (Note: U.S. copper pennies
dated 1982 or earlier work best in this experiment, but any"copper" penny can be used.)
2. Use steel wool to make it shiny and immerse it in a vinegar solution for 30 seconds to clean it.
Remove the penny, rinse it off and dry with a paper towel.
3. The teacher will have placed 10 mL of 6M NaOH and a small scoop (pea sized) of zinc powder
to an evaporating dish. THERE ARE SEVERAL SET UP UNDER THE VENTILATION
HOOD. THESE ARE HOT—BEWARE HOT ITEMS. Do not allow the sodium hydroxide
solution in this experiment to actively boil. Sodium hydroxide is caustic and may splatter causing
severe damage to the skin or eyes. In case of contact, wash it off immediately with cold water
until the skin no longer feels soapy.
4. Drop the penny into the dish and watch it turn “silver” before your very eyes.
5. Remove the silver penny from the dish with the tongs. Rinse it thoroughly with tap water and
dry it with a paper towel.
6. At your station, hold the penny with the tongs and insert it into the hot part of the flame. It
should quickly turn to gold before your very eyes. Do NOT leave it in the flame for too long. It
will cause the penny to react and turn a nasty color.
Questions:
1. Why does it look gold?
2. Why is it necessary to clean the penny before starting the reaction?
Page 120 of 146
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Making Paint
Most paints consist of colored pigments combined with a binder to keep them soft and to allow them
to be spread on paper or canvas. In this activity we will chemically produce several colored
pigments. The paints you make in this activity can be used to paint a picture.
Procedure
1.
Make the pigments:
A. Green Pigment
1. Place 0.3 g of K4Fe(CN)6 in a test tube. Fill the tube half full with warm water. Shake
the tube until the solid dissolves.
2. Add 0.2 g of CoCl2 to the tube. Stopper the tube and shake it thoroughly. Avoid
touching the CoCl2. It is toxic.
3. Filter the solution. Discard the liquid filtrate and save the precipitated pigment on the
filter paper.
B. Brown pigment
1. Dissolve 0.2 g NH4Fe(SO4)2•3H2O in a tube half filled with warm water. Then add 0.2 g
Na2CO3. Shake and filter.
C. Blue pigment
1. Dissolve 0.2 g NH4Fe(SO4)2•3H2O in a tube half filled with warm water. Then add 0.2 g
K4Fe(CN)6. Shake and filter.
D. Orange pigment
1. Dissolve 0.2 g NH4Fe(SO4)2•3H2O in a tube half filled with warm water. Then add 1.0
mL Na2SiO3 solution. Shake and filter.
E. Royal Blue pigment
1. Dissolve 0.2 g CoCl2 in a tube half filled with warm water. Then add 1.0 mL Na2SiO3
solution. Shake and filter.
F. Lavender pigment
1. Dissolve 0.2 g CoCl2 in a tube half filled with warm water. Then add 0.2 g of Na2CO3.
Shake and filter.
Page 121 of 146
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Guar Gum Slime
Discussion:
Slime, a product of the Mattel Toy Corporation, was described by Dr. Maki Papavasiliou, of the
Mattel Materials Laboratory, as a reversible cross-linking gel made from Guar gum, a vegetable gum
used as a protective colloid, stabilizer, and thickening agent for foods, cosmetics, and lotions. The
crosslinking is accomplished by the addition of sodium borate, Na2B4O7•10H2O, commonly called
borax.
Slime is a non-Newtonian fluid that is dilatant . That is, under stress, the material dilates or expands.
Other stress-thickening materials are quicksand, wet sand on the beach, some printer’s inks, starch
solutions, and Silly Putty. Dilatant materials tend to exhibit some unusual properties.
a) Under low stress, such as slowly pulling on the material, it will flow and stretch. If careful, you
can form a thin film.
b) Pull sharply (high stress) and the material breaks.
c) Pour the material from its container then tip the container upward slightly, the gel will self siphon.
d) Put a small amount of the material on a table top and hit it with your hand, there is no splashing or
splattering.
e) Throw a small piece onto a hard surface, it will bounce slightly.
Guar gum is a long-chain polyalcohol with 1,2-diol groupings capable of complexation with the
borate ion, B(OH)4. Guar gum, the main component of Slime, is a vegetable gum derived from the
guar plant, Cyamopsis tetragonolobaus.
In addition to forming complexes with the borate ion, the interaction of long-chain polyalcohols,
such as guar gum, with the borate ion leads to cross linking of different polymer chains, or
sometimes part of the same chain, in such a way that a three-dimensional network of connected
chains is formed. When the concentration of cross-linked chains is high, solvent is immobilized
within the network and a semisolid gel results. Because the borate ion can bond with four alcohol
groups it is particularly effective in creating three-dimensional gel networks from gums such as guar
gum. Other examples of networks and gels are rubber cement, gelatin, fruit jellies, agar, and yogurt.
Page 122 of 146
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Procedure:
1. Pour 10 mL of tap water into a 250 mL beaker. Place a plastic baggie inside the beaker and fold
over the baggie over the top of the beaker.
2. Pour 80 mL of warm distilled water into the baggie. Tap water contains ions that can interfere
with the cross-linking of the slime.
3. Add approximately 2 g of guar gum to the water. Add slowly and stir until completely
dissolved.
4. Add 2 drops of food coloring to the guar gum solution.
5. Obtain 5 mL of sodium borate solution. Add to the guar gum solution. Stir for 1 minute and
then let it sit for 2 minutes.
6. Transfer your slime to a plastic baggie.
7. Enjoy your slime but do not eat the slime and wash your hands after playing with it.
Clean up
1. Wipe up slime which has escaped with a paper towel.
2. DO NOT put any slime into the sink.
3. Remove your slime filled baggie from the beaker, seal it and wipe any water off the outside of
the bag.
4. Rinse out your beaker, dry, and return to the cabinet.
5. If Slime gets on any material, it can be removed by first wetting it with vinegar to break down
the gel, followed by soapy water or foam type upholstery cleaner.
Page 123 of 146
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Ice Cream: Freezing Point Depression of a Solution
Discussion
Colligative properties are properties of solutions that deal with the number of particles dissolved in
solution, not the type of particle in solution. When particles are dissolved in solution, they have the
effect of raising the boiling point and depressing the freezing point. The change in freezing point
can be calculated by the following:
f= molality x -1.86 °C kg/mole (for water solutions only). The molality of solution is calculated
by multiplying the moles of solute by the number of particles it forms in water and dividing by the
kg of solvent. Sucrose (table sugar) is a non-electrolyte and so will only form one particle in
solution. Sodium chloride (table salt) is an electrolyte and will split into two particles, a sodium ion
and chloride ion. In this exercise you will be making a delicious treat and also learning about
colligative properties in a kinesthetic way☺.
Ice Cream Recipe (per person)
½ cup milk
1 tbsp sugar
¼ tsp vanilla
½ bag ice
4-6 tbsp. salt
Two freezer quart sized zip top baggies. (No slider zip bags)
Place the first 3 ingredients in the first bag, seal tightly being sure to get as much air out as possible.
Place the ice and salt in the 2nd bag. Place the bag with the milk mixture into the bag with the salt
and ice and seal the 2nd bag tightly. Toss and knead the bags until the ice cream is frozen and
delicious. Use the electronic thermometers to measure how cold the solution gets. Use all of your
senses to experience the product.
Alternate recipe:
72 mL orange soda
48 mL condensed milk or whipping cream
Bags, salt and ice as described above.
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Periodic Table
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Bob Jones High School Chemistry I
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Rules of Writing Equations, Solubility Rules, Activity Series of Metals
Synthesis (Use Gas Ion Chart)
1. Combination of elements.
2. A metal oxide plus water yields a base.
3. A non-metal oxide plus water yields an acid.
4. A non-metal oxide plus metal oxide yields a salt.
Decomposition (Use Gas Ion Chart)
1. A base when heated decomposes into a metal oxide plus water.
2. An acid when heated decomposes into a non-metal oxide plus water.
3. Metallic carbonates decompose into a metal oxide and carbon dioxide.
4. Metallic chlorates decompose into a metallic chloride and oxygen.
5. Some compounds decompose with electricity or just simply decompose into their basic elements.
Single Replacement Reactions (Use activity series)
1. A metal will replace a less active metal in a compound.
2. Some metals will replace the H in water to produce a metallic hydroxide and hydrogen gas.
3. Some metals will replace the H in acid to produce a salt and hydrogen gas.
4. A halogen (group 17) will replace a less active halogen in a compound.
Double Replacement Reaction States of matter are important. Use the solubility rules.
1. An acid and a base yield a salt and water.
2. A salt and an acid yield a different salt and a different acid.
3. A salt and a salt yield a salt and a salt.
4. Some compounds decompose when made in double replacement reactions
 If carbonic acid is made it decomposes into water and carbon dioxide gas.
 If ammonium hydroxide is made it decomposes into water and ammonia (NH3) gas.
 If sulfurous acid is made it decomposes into water and sulfur dioxide gas.
Combustion
A compound containing at least hydrogen and carbon is mixed with oxygen and produces carbon
dioxide and water.
Gas Ion Chart
Gas
Ion
SO2
SO32SO3
SO42CO2
CO32N2O3
NO2N2O5
NO3P2O3
PO33P2O5
PO43H2O
OHNH3
NH4+
Solubility Rules (Note: If one rule says a compound is soluble, it is soluble regardless of the other
rules)
1.
All ammonium and group 1 compounds are soluble.
2.
All acetates, chlorates, perchlorates and nitrates are soluble.
Page 126 of 146
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3.
4.
5.
6.
7.
Updated 3/23/2016
All bromides, chlorides, and iodides are soluble except lead, mercury 1 and silver.
All hydroxides are insoluble except group 1, barium and strontium. Calcium is slightly
soluble.
The borates, carbonates, chromates, oxides, phosphates, silicates, and sulfites are insoluble
except those of ammonium, potassium and sodium.
All sulfides are insoluble except Group1 and Group 2
All sulfates are soluble except those of Group 2, lead, mercury(I) and silver. Mg, Ca and Sr
are moderately soluble
Most Active
Least Active
Li Rb K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Hg Ag Pt Au
------------------------------------------------------------------------------React w/ oxygen----------------------------------------------------------React w/ acid----------------------------------React w/ steam------------------------------------React w/ cold water--
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The dividing line between soluble and insoluble is 0.1-molar at 25 °C. Any substance that can form
0.1 M or more concentrated is soluble. Any substance that fails to reach 0.1 M is defined to be
insoluble.
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Demonstration: Air Pressure
Concept: Air pressure acts in all directions.
Mateials:
Plastic Cup
Large Notecard
Water
Procedure:
1. Fill plastic cup with water.
2. Place the notecard over the cup.
3. Carefully invert the cup while holding the notecard in place.
4. Remove your hand from the notecard.
The notecard should remain in place and water will remain in the cup.
Page 129 of 146
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Demonstration: Burning Lycopodium Powder:
Simulating a Grain Elevator Explosion
Lycopodium powder is a fine yellow powder derived from the spores of Lycopodium clavatum
(stag's horn club moss, running ground pine). When a lighted match is dropped into a pile of this
powder, it does not burn. However, when the powder is dispersed into a fine mist near a candle
flame, it ignites into a spectacular fireball. This results from an increasing the available surface area
for combustion: when the powder is dispersed into a mist, the particles are surrounded by enough
oxygen to support a combustion reaction.
This demonstration illustrates the basic principle behind a grain elevator explosion. Grain dust, like
lycopodium powder, is not especially flammable, but when grain is dumped into a grain silo or
elevator, some of the finer dust particles can remain suspended in air, surrounded by oxygen. The
mixture can be ignited by a spark or flame, resulting in a devastating explosion. (For more
information, see the following web sites: Kansas Grain Elevator Photographs and a National
Materials Advisory Board report on grain elevator explosions.)
Drop a match into a pile of lycopodium powder in a watch glass. The match is almost immediately
extinguished. Put the same powder through a funnel into a rubber tube, with a glass pipette at the
other end. When the powder is blown through the pipette past a candle flame, the powder bursts into
a fireball.
For Halloween, put a candle in a pumpkin, light it, put a hole in the back of the pumpkin and run a
plastic tube up to it. Put two scoops of lycopodium powder in a funnel connected to the tube and
wiggle the tube to get the powder close to the pumpkin. Take off the funnel, blow fairly hard into
the tube to get the powder to atomize.
!!! Hazards !!!
Lycopodium powder is flammable when dispersed into a fine mist. Usually, with this procedure, the
flame burns out too quickly to light any other combustible materials, but caution should be exercised
nonetheless.
Note: This can be used as a cool Halloween demo when the funnel is placed in the back of a carved
pumpkin. It can also be used to blow the lid off a can to demonstrate pressure increase with
increasing temperature.
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Demonstration: Burning Magnesium
Magnesium is one of the alkaline-earth metals, and is one of the most common elements in the
Earth's crust. In its pure form, it is silvery white, and relatively soft. It burns in air with a brilliant
white light, and for this reason is often used in flares and fireworks.
2Mg(s) + O2(g) ——> 2MgO(s)
The high temperatures reached during the combustion also allow small amounts of magnesium to
react with nitrogen in the air, producing magnesium nitride:
3Mg(s) + N2(g) ——> 2Mg3N2(s)
Magnesium also burns in an environment of carbon dioxide such as in a beaker full of dry ice:
2Mg(s) + CO2(g) ——> 2MgO(s) + C(s)
Magnesium is used in disposable flash bulbs to generate light for photography, but this use has been
largely supplanted by other sources of illumination.
In the demonstration below, a strip of magnesium ribbon is ignited with a Bunsen burner:
!!! Hazards !!!
In addition to being extremely bright, burning magnesium produces some ultraviolet light; avoid
looking directly at it.
The burning magnesium is very hot; do not touch it or let it come in contact with other flammable
materials.
Since magnesium burns in the presence of carbon dioxide, a CO2 fire extinguisher does not put out
the flame from burning magnesium; a dry-chemical fire extinguisher must be used instead.
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Demonstration: Density: Coke vs. Diet Coke
Objectives:
to list similarities of given objects
to list differences of given objects
to brainstorm and find a solution as to why Diet Coke floats and Coke sinks
to define the term density
to see how much sugar we consume drinking one can of soda
Materials:
coke and diet coke
water
fish tank
sugar
nutra sweet
Procedure:
Pass the cans of coke around the room. Have each student take a good look at each can and ask
them to make careful observations about what they see.
Ask the students to name as many similarities as they can about the 2 cans of coke. Make a list on
the board.
Ask the students to list as many differences as they can about the 2 cans. Add to list
Place the regular coke into a small tank of water.
Place the diet coke into the water. (Look surprised and take both out. Have a student come up to
verify that the cans are still sealed and have not been tampered in anyway!)
Place back into water. Ask the students to explain why one is floating.
Possible responses:
they weren't filled right at the plant
the red paint is heavier than the silver paint, or vice versa
one is flat, the carbon dioxide must have leaked out
nutra sweet is lighter than sugar
The "Why":
Show the students what 39 g of sugar looks like ( I found it effective to show the sugar in a small
beaker while holding it next to the can so they can see how much space it would take up in the can)
next to approx *100 mg (on an index card) of Nutra sweet. Explain that ALL that sugar is in the
regular Coke can, and that small amount of Nutra Sweet is in the Diet Coke can. Explain that a
small amount of Nutra Sweet is needed to make the Diet Coke sweet because it is so concentrated.
Most students are surprised to actually SEE how much sugar there is!
Discuss how more "stuff" (matter) is crammed into the same amount of space, or VOLUME, and
that increases the MASS. The relationship of Mass to Volume is Density. The more items (matter)
you place into a defined space, the denser it becomes. For example, New York City is DENSELY
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populated because there are a lot of people in a small area. 20 people in an elevator is DENSER
than 2 people in an elevator.
The Density of water is 1g/cm3. An object will float is the density is less than 1. An object will sink
if its density is greater than 1.
*Note: According to the Coca Cola company : "a soft drink sweetened with aspartame (8 oz., ranges
from approximately 10 to 85 mg)" A can has 12 oz, so I approximated 100 mg for measuring
purposes since my triple beam balance has a 0.1g bar. You can also say that there are 39,000
milligrams of sugar in a can of regular Coke!!!
Extension:
Weigh the Coke and Diet Coke to determine mass of each can. Using water displacement, find the
volume of each can. Use the formula D=M/V and see if you can determine their densities. Is Diet
Coke's density less than 1? Is regular Coke's density greater than 1?
Does this work for all Diet sodas? Try different brands, for example Pepsi, Dr. Pepper, Sprite, etc.
Have the kids form predictions and test them out!
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Demonstration:Egg In a Bottle
A small amount of burning paper is placed into the glass bottle. The egg is placed at the mouth of
the bottle ( "pointy" end down ). The flame goes out and the egg is drawn into the bottle.
Be advised that the incorrect explanation is often given for this Demo. Most people claim that the
fire "uses up the oxygen." Nothing could be further than the truth! What really happens is that the
fire heats the air inside the bottle, which in turn causes it to expand and rise. ( With a good fire the
egg will actually bob up-and-down on the mouth of the bottle as escaping air rushes past it. ) The
fire eventually extinguishes itself, the air in the bottle ( no longer heated ) contracts thus lowering the
air pressure in the bottle ( Remember, some of the air left the bottle so there is less air and air
pressure in the bottle. ) The greater pressure on the outside of the bottle is greater than the pressure
in the bottle so the egg gets pushed in! Plop!
To remove the egg: Invert the bottle to allow the egg to roll back to the bottle neck then blow into
the bottle. Air will go in, around the egg, and into the bottle -- thus raising the pressure within the
bottle. When the pressure in the bottle exceeds room air pressure, the egg will be pushed back out.
The egg can also be drawn into the bottle with no fire at all. If the egg is placed on the bottle and
then the air within the bottle is dramatically cooled -- the air will contract and lower the pressure
within the bottle. This dramatic cooling can be done by placing the bottle into an inch or so of liquid
nitrogen ( LN2 ). After the egg plops into the bottle, additional air from the room will fall into the
bottle, cool, and contract. After a few minutes you can take the bottle away from the LN2 and then
quickly invert the bottle so that the egg blocks the mouth of the bottle. As the air within the bottle
warms up it will expand and push the egg out of the bottle! This helps to show that the fire from
above does not use up the oxygen!
NOTE: Repeated attempts will eventually cause the bottle to break from the temperature induced
stress
Then, of course, this Demo can be done with hot water. Yes, hot water ! Boil some water in a
teakettle or pan. Carefully fill the glass bottle with the hot water. Swirl the water around inside the
bottle then pour the water out. Quickly place the egg over the mouth of the bottle. Plop ... the egg
gets pushed into the jar. The heat from the boiling water causes the air inside the bottle to expand,
forcing some of the air out. As the air inside the bottle begins to cool, it contracts, reducing the air
pressure inside the bottle. The greater air pressure outside the bottle forces the egg into the bottle.
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Demonstration: Elephant Toothpaste
This experiment shows the decomposition of hydrogen peroxide catalyzed by potassium iodide. The
reaction is done in a tall graduated cylinder so that the foamy product shoots out very quickly in a
tall cylindrical shape; hence, the name elephant toothpaste.
Materials:
tall graduated cylinder (at least 500 ml)
food coloring
Dawn detergent
30% hydrogen peroxide (H202)
potassium iodide (KI) (solid or saturated solution)
disposable gloves
Hazards:
1. Wear safety goggles. Also, wear disposable gloves when pouring 30% hydrogen peroxide, as it is
a very strong oxidant.
2. Do not stand over the graduated cylinder because steam and oxygen are produced quickly.
Procedure:
1. Place a garbage bag or other covering on the lab table and possibly on the floor.
2. Put on disposable gloves. Pour 80 ml of 30% hydrogen peroxide into a graduated cylinder.
3. Add about 40 ml of Dawn detergent to the hydrogen peroxide. Swirl to mix
4. Tilt the graduated cylinder and drip red and/or blue food coloring down the sides of the graduated
cylinder to make your toothpaste striped
6. Quickly add a large scoop of KI and stand back. You may also add some saturated KI solution
instead of the solid. Be sure to move your hand away from the top of the graduated cylinder quickly
or the hot foam will get on your hand and arm.
7. You may place a glowing splint in the foam to test for oxygen, but do not drop the splint into the
graduated cylinder. The splint will relight indicating the presence of oxygen.
Discussion:
The rapid catalyzed decomposition of hydrogen peroxide produces O2 gas which forms a foam with
the liquid detergent:
2H2O2 (aq) -> 2H2O + O2 (g)
The I-1ion is a catalyst for the reaction. The brown color of the foam is evidence of iodine in the
reaction. It will stain clothes, skin, and carpet
Disposal:
Leave the gloves on while cleaning up. The foam and solution left in the graduated cylinder may be
rinsed down the drain with excess water.
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Demonstration: Endothermic Reaction
An endothermic process or reaction absorbs energy in the form of heat (endergonic processes or
reactions absorb energy, not necessarily as heat). Examples of endothermic processes include the
melting of ice and the depressurization of a pressurized can. In both processes, heat is absorbed from
the environment. You could record the temperature change using a thermometer or by feeling the
reaction with your hand. The reaction between citric acid and baking soda is a highly safe example
of an endothermic reaction, commonly used as a chemistry demonstration. Do you want a colder
reaction? Solid barium hydroxide reacted with solid ammonium thiocyanate produces barium
thiocyanate, ammonia gas, and liquid water. This reaction gets down to -20°C or -30°C, which is
more than cold enough to freeze water. It's also cold enough to give you frostbite, so be careful! The
reaction proceeds according to the following equation:
Ba(OH)2.8H2O (s) + 2 NH4SCN (s) --> Ba(SCN)2 (s) + 10 H2O (l) + 2 NH3 (g)
Here's what you need to use this reaction as a demonstration:
32g barium hydroxide
17g ammonium thiocyanate (or could use ammonium nitrate or ammonium chloride)
125-ml flask
stirring rod
Perform the Demonstration
Pour the barium hydroxide and ammonium thiocyanate into the flask.
Stir the mixture.
The odor of ammonia should become evident within about 30 seconds. If you hold a piece of
dampened litmus paper over the reaction you can watch a color change showing that the gas
produced by the reaction is basic.
Liquid will be produced, which will freeze into a slush as the reaction proceeds.
If you set the flask on a damp block of wood or piece of cardboard while performing the reaction
you can freeze the bottom of the flask to the wood or paper. You can touch the outside of the flask,
but don't hold it in your hand while performing the reaction.
After the demonstration is completed, the contents of the flask can be washed down the drain with
water. Do not drink the contents of the flask. Avoid skin contact.
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Demonstration: Gummy Bear Sacrifice: Energy of Oxidation of Carbohydrates
Description: This demonstration illustrates the vast amount of energy which is available from the
oxidation of carbohydrates, such as sugar. An excess of oxygen, generated by the decomposition of
sodium chlorate (potassium chlorate can also be used), will react with a gummy bear, releasing a
large amount of energy quickly and dramatically.
Materials Needed





one 25x150 mm test tube
one ring stand with clamp for test tube
5-7 grams sodium chlorate
Bunsen burner
one candy gummy bear (any other candy or a wood splint will work)
Hazards: This reaction produces a large quantity of heat, flame, and smoke (mostly water
vapor). It should be done in a well ventilated room. Sodium chlorate should be used with
caution. It is a strong oxidizing agent, especially when molten. Keep all combustible materials
away from the reaction area. Make sure the test tube used is scrupulously clean and the mouth is
pointed away from the audience.
Procedure: Set up the stand and clamp, and support the test tube in the clamp in a vertical position.
Add 5-7 grams of sodium chlorate to the test tube (about 1 cm in depth). [Note: potassium chlorate
can be substituted here, but the sodium chlorate has a lower melting point and requires less initial
heating.] Gently heat the tube with the burner until the sodium chlorate is completely molten.
Bubbles of oxygen will begin to form. Remove the burner and use crucible tongs to drop in the
gummy bear, then stand back! For added piece of mind, the reaction can be performed behind a
safety shield.
Disposal: Allow the tube to cool, then remove from the clamp. The tube should be soaked in water
for about 15 minutes and then cleaned with a brush to remove the residue. These chemicals may be
washed down the drain.
Discussion: When heated, sodium chlorate decomposes, producing sufficient oxygen to ignite the
sugar in the gummy bear. Since the oxidation of the sugar is very exothermic, sodium chlorate
continues to decompose to oxygen, and the rate of combustion becomes very rapid.
2NaClO3(s)  2NaCl(s) + 3O2(g)
C12H22O11(s) + 3O2(g)  9C(s) + 3CO2(g) + H2O(g) + 5635 kJ
Other carbohydrate materials may be used for this reaction, e.g. an M&M, gum drop or cinnamon
heart. The size of the candy and test tube should be matched so that the candy will easily fit into the
tube.
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Demonstration: Methane Bubbles
Description
Bubbles are blown with ordinary soap solution, but using methane (natural gas) instead of air. As the
bubbles rise in the air, a candle at the end of a meter stick is held up to them. They burst into large,
feathery flames.
Explanation
Methane, or natural gas, is less dense than air, so when bubbles are blown with it, they rise. Methane
(what many of you use to heat your homes) is combustible, that is, when a flame or spark is applied,
it will react rapidly with oxygen to form carbon dioxide and water and produce heat and a flame.
CH4 + 2O2  CO2 + 2H2O
\
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Demonstration: Methane Mamba http://dwb4.unl.edu/chemistry/beckerdemos/BD015.html
Description
Methane gas is bubbled up through a funnel of soapy water and a buoyant column of suds grows
gracefully upward like a large bubbly snake swaying elegantly to the air currents in the room.
Igniting this methane mamba provides for a rather charming effect! See safety precautions.
Chemical Concepts
1. Whether an object floats or sinks in a fluid depends on whether that object's density is less
than or greater than the density of the fluid. This holds true for objects submerged in liquids
as well as gases.
2. For gases under similar conditions, densities are essentially proportional to molecular
masses.
3. Hydrocarbons are generally combustible -- that is, they react exothermally with oxygen. The
products are usually CO2 and H2O.
Safety



Make sure ventilation is adequate.
CAUTION: SAFETY GOGGLES SHOULD BE WORN AT ALL TIMES, AND A FIRE
EXTINGUISHER SHOULD BE ON HAND. REMOVE ALL COMBUSTIBLE
MATERIAL FROM THE VICINITY OF THE DEMONSTRATION.
NOTE: A 10 CM HIGH COLUMN OF METHANE SUDS MAY PRODUCE A FLAME
WELL OVER A METER TALL. LET THE HEIGHT OF YOUR CEILING AND THE
FLAME RETARDANT CAPABILITY OF THE CEILING TILES DICTATE HOW
LARGE A CLUSTER CAN BE SAFELY IGNITED.
Procedure
1. Carefully, and with adequate lubrication, slide the rubber stopper over the glass tube to about
the mid-section. Then insert the stopper securely into the mouth of the funnel.
2. Connect the protruding end of the glass tubing to the rubber hose. If the fit is too loose, wrap
the end of the tube with some electrician's tape to make the hose fit more snugly.
3. Use the ring stand, clamp and large ring support to secure the funnel in a vertical position,
stoppered end down as shown below. Run the hose up over the neck of the ring support (to
avoid leak-back) and then connect it to the gas jet. Set up a FIRE EXTINGUISHER nearby.
4. Pour 300 mL of detergent solution (3% Dawn® by volume) into the funnel. The top of the
glass tubing should be submerged by about 1.5 cm.
5. Begin the snake charming music (if available!). Then, with no open flames nearby, simply
turn on the gas jet full throttle. A column of methane bubbles begins to grow lazily upward,
attaining a height of 2-3 meters (room height permitting) in about 5 minutes. If the column is
allowed to grow for a long enough period of time it begins to arch over, for as the methane
diffuses out of the top bubbles and as air diffuses in, the upper section increases its density
and eventually bends over on its own.
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6. An interesting effect can be achieved by placing a few drops of water on the top of the
growing column. Since the water increases the density, it causes the top to arch over -accentuating the snake-like appearance even more! But as the head of the snake drops down,
water falls from its snout, the density decreases and the snake rears its head back up! Adding
drops to the middle of the column creates a downward curve. With practice, one can fashion
the suds column into a whole assortment of snake-like configurations.
7. At any point, the column of bubbles may be gently scooped up off the funnel and carefully
walked around the room. (Wetting the hands first avoids popping too many bubbles). Large
sweeping motions make for a very graceful display!
8. Perhaps most impressively, these bubble columns may be ignited.
CAUTION: SAFETY GOGGLES SHOULD BE WORN AT ALL TIMES, AND A FIRE
EXTINGUISHER SHOULD BE ON HAND. REMOVE ALL COMBUSTIBLE
MATERIAL FROM THE VICINITY OF THE DEMONSTRATION. NOTE: A 10 CM
HIGH COLUMN MAY PRODUCE A FLAME WELL OVER A METER TALL. LET THE
HEIGHT OF YOUR CEILING AND THE FLAME RETARDANT CAPABILITY OF THE
CEILING TILES DICTATE HOW LARGE A CLUSTER CAN BE SAFELY IGNITED.
The safest manner for presenting this is to turn the gas jet off, scoop the column off the
funnel, set it down on the lab bench at least 2 meters away, and ignite it with a lit candle
attached to the end of a meter stick. The columns may be lit from the bottom (this produces
the fastest burning and therefore the largest flame), from the top (this shows an interesting
fuse-like downward progression of the fire) or from somewhere in between.
9. The diffusion of the methane across the soap film can be demonstrated by holding the flame
a few centimeters above the bubble column; after a few seconds the top of the column
ignites, even though the flame never touches the column! Furthermore, a small cluster of
bubbles may be flattened out on the lab bench and left there for about two minutes. When a
flame is then held directly to the bubbles, nothing happens -- they pop, of course, but they do
not ignite -- even though the cluster appears just as it did when the bubbles were blown. This
indicates not only that the methane has diffused out of the bubbles, but also that air must has
diffused in at a more or less equal rate.
10. If the column is allowed to grow for a long enough period of time to form a full arch, as
described above, and then this arch is ignited at the near (most recently blown) end, the
flames race up the near side and then down the other but invariably stop 10-20 cm before
reaching the tip -- leaving a small cluster of suds dangling momentarily in the air. And if a
flame is held to this slowly descending cluster, again, the bubbles pop, but no combustion
occurs.
Questions
1.
2.
3.
4.
Predict the differences if air were substituted for the gas in this apparatus.
Predict the differences if CO2 were substituted for the gas in this apparatus.
Predict the differences if H2 were substituted for the gas in this apparatus.
Why would a newly-blown cluster of methane suds ignite easily, but a cluster blown five
minute ago not ignite at all?
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TimeTeacher Preparation: 10 minutes
Class Time: 15 minutes
Materials








300 mL of 3% solution of dish detergent (300 mL -- In a bowl, mix approximately 9 mL
detergent with 290 mL of water.) (Joy® or Dawn® work best)
source of natural gas (methane)
a funnel (made from the top half of a 2-L soda bottle)
1 #3 1-holed rubber stopper
glass tubing (6 mm diam) to fit stopper, 9 cm long
rubber hose to fit gas jet nozzle, ca. 1 m long
ring stand and large iron ring support (5 in diam)
ring stand (buret) clamp
candle taped securely to the end of a meter stick
DisposalBe certain the hood is on to enhance the ventilation of the gas.
Lab HintsIf there are too many large bubbles, the column is too buoyant and tends to pinch off before reaching
its full height. If there are too many small bubbles (suds), the column is too dense and simply spills
over the rim of the funnel. The best results are obtained from a combination of large and small
bubbles to produce a column that is just barely less dense than air -- enough to support its own
weight but not cause a substantial upward tug. This might require some "fine tuning", adjusting the
flow rate of the methane, the position of the glass tubing and the depth of the soapy water in the
funnel. Humidity may also play a role, for the top of the column does tend to dry out.
Observations

Methane is about half as dense as air. You can approximate this by comparing their relative
molecular masses: CH4 (molar mass(mw) = 16) compared to air, a mixture of N2 (mw = 28)
and O2 (mw = 32). Because of its lower density, pure methane rises rather rapidly in air. The
soapy water adds significantly to the density, but with just the right proportions, as described
above, you can achieve a soapy water/methane mixture that is just slightly less dense than the
surrounding air, and that ascends very slowly. The adhesive nature of the soapy water, due
mostly to the H-bonding that occurs between the water molecules, helps to hold adjacent
bubbles suds together in a snake-like cluster.
Use the Shrinking Suds (BBExperiment 016) for comparisons.
Answers-
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Q1. Predict the differences if air were substituted for the gas in this apparatus.
A1. The column of bubbles sinks slowly to the desk top because the liquid makes the bubbles
slightly more dense than air. The column would not be not flammable.
Q2. Predict the differences if CO2 were substituted for the gas in this apparatus.
A1. The column of bubbles sinks sharply to the desk top because the molar mass of CO2 is much
greater than the molar mass of air. The column is not flammable; in fact, it could be used to
extinguish a flame, for CO2 does not support combustion.
Q3. Predict the differences if H2 were substituted for the gas in this apparatus.
A3. The column of bubbles rises more rapidly and may break off if the buoyant force (air pushing up
the hydrogen) exceeds the adhesive forces (H-bonds between adjacent water molecules) holding the
bubble cluster together. The column is quite flammable.
Q4. Why does a newly-blown cluster of methane bubbles ignite easily, but a cluster blown five
minute ago not ignite at all?
A4. A newly-blown cluster of methane bubbles contains methane (obviously), which is a highly
combustible hydrocarbon. As time passes, however, that methane diffuses out of the bubble (from
the high concentration inside to the lower concentration outside). Simultaneously, and at apparently
the same rate, air diffuses into the bubbles (this time the high concentration is on the outside, and the
low concentration is on the inside). Thus, if a methane bubble lasts long enough, it no longer
contains an appreciable amount of methane, and is therefore no longer flammable.
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Demonstration: Money to Burn
Application
1. Specific Heat of Water
2. Observation Skill Practice
Materials
1. Dollar bill Metal tongs
2. Solution 50% water & 50% ethanol
3. wood splint
Safety Precautions
1. Ethanol is extremely flammable and a dangerous fire risk, it is also toxic by ingestion.
2. Beware of fumes or residual alcohol igniting! Wear flame resistant gloves, a chemicalresistant apron, chemical splash goggles and always follow laboratory safety rules while
performing demonstrations.
Scientific Concept behind Burning Money
A combustion reaction occurs between alcohol and oxygen, producing heat and light (energy) and
carbon dioxide and water.
C2H5OH + 4 O2 -> 2 CO2 + 3 H2O + energy
When the bill is soaked an alcohol-water solution, the alcohol has a high vapor pressure and is
mainly on the outside of the material (a bill is more like fabric than paper, which is nice, if you've
ever accidentally washed one). When the bill is lit, the alcohol is what actually burns. The
temperature at which the alcohol burns is not high enough to evaporate the water, which has a high
specific heat, so the bill remains wet and isn't able to catch fire on its own. After the alcohol has
burned, the flame goes out, leaving a lightly damp dollar bill.
Preparation
Prepare 50% alcohol- 50% water solution.
Demonstration
1. Dip dollar bill into solution using tongs, remove after a few seconds and light bill with a lite
wood splint.
2. After a few seconds of burning flip the bill with a quick snap of the wrist to put out flame .
Show the intact unburned bill.
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Demonstration: Thionin — The Two-Faced Solution: Light Energy and
Chemical Energy
Discussion
A large beaker contains a bright purple solution. The beaker is placed on an overhead projector that
is half-covered with aluminum foil—half of the purple solution is sitting on the piece of aluminum
foil, the other half is sitting directly on the overhead stage. Switch on the overhead lamp and, in
seconds, the solution on the side of the beaker exposed to light turns colorless, while the unexposed
side remains purple. The result is sharp and stunning—a two-faced solution! The solution appears to
be divided by an invisible line running vertically through the solution. Amazing enough, but switch
off the overhead light and the entire process can be reversed.
Chemical Concepts
• Reversible reactions • Oxidation–reduction • Photochemistry
Materials (for each demonstration)
Aluminum foil
Beaker, glass, 1-L
Iron(II) sulfate, FeSO4•7H2O,
2.0 g Balance
Sulfuric acid solution, H2SO4, 1 M,
100 mL Cylinder, graduated,
10-mL Thionin solution, 0.001 M,
10 mL Stirring rod, glass
Distilled water
500 mL Overhead projector
Safety Precautions
Sulfuric acid solution is corrosive to eyes and skin. Iron (II) sulfate is slightly toxic by ingestion.
Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Please
review current Material Safety Data Sheets for additional safety, handling, and disposal
information.
Preparation
Prepare 0.001 M thionin solution by adding 100 mL distilled water to 0.023 g of thionin. Stir to
dissolve. The thionin solution has a poor shelf life; use within one week.
Procedure
1. Mix together the following chemicals in a 1-L beaker: 10 mL of freshly prepared 0.001 M
thionin solution, 100 mL of 1 M sulfuric acid, and sufficient distilled water to bring the final
volume to about 600 mL. Mix thoroughly.
2. Turn off the room lights and add 2.0 grams of iron(II) sulfate. Stir to dissolve.
3. Place the beaker on the overhead projector stage. Turn on the lamp. Observe that the solution
changes from purple to colorless in a matter of seconds.
4. Now turn off the lamp and allow the purple color to return.
5. On the overhead projector stage, place a piece of aluminum foil several layers thick. The foil
should not cover the entire projector stage.
6. When the solution is purple, place the beaker on the projector stage so that half of the beaker is
sitting on the piece of aluminum foil. (The students should be in direct line with the bisecting
line so they can observe the vertical division.)
7. Turn on the projector lamp, and observe the solution. A distinct vertical division between the
purple side and colorless side should be clearly visible. (The vertical division indicates that the
reaction is initiated by light and not heat.)
8. The reaction can be reversed by turning off the light. The reaction is reversible for several days.
Disposal
The two-faced solution may be rinsed down the drain with excess water.
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Tips
• Prepare the thionin solution fresh. The thionin solution will lose its activity if stored for a long
period of time.
• The color change will fade over time from bright purple to a paler shade of blue or purple. This
color change may still be effective after one week.
• Direct, bright sunlight may also be used as the light source.
Discussion
Thionin is an organic compound that can exist in two forms, an oxidized form which is purple and a
reduced form which is colorless. When a reducing agent such as iron(II) ion (Fe2+) is added to an
acidic thionin solution, the protonated thionin molecule (thio+) accepts two hydrogen atoms and is
reduced to its colorless form—but only in the presence of an intense light source. The reduction is a
photochemical reaction that is catalyzed by light. This demonstration is a dramatic example of the
conversion of light energy to chemical energy. The reaction can also be reversed; when the light
source is removed the purple color due to the oxidized form of thionin returns.
The demonstration also provides a vivid example of a reversible reaction and equilibrium. The
equilibrium is represented by the following chemical equation:
Remember that if one reactant in a balanced chemical equation is oxidized, another reactant must be
reduced. Fe2+ is oxidized to Fe3+ in the forward reaction, while in the reverse reaction Fe3+ is reduced
to Fe2+. Thio+ represents the monoprotonated form of thionin in acidic solution.
Reference
Flinn ChemTopic™ Labs, Volume 15, Equilibrium; Cesa, I., Editor; Flinn Scientific: Batavia IL
(2003).
Materials for Thionin—The Two-Faced Solution are available from Flinn Scientific, Inc.
This activity is also available as a Flinn Chemical Demonstration Kit that contains all the materials
to perform the demonstration seven times.
Catalog No. Description
T0077 Thionin, 2.5 g
S0417 Sulfuric Acid Solution, 3 M, 500 mL
F0016 Ferrous Sulfate, 500 g
A0019 Aluminum Foil, 25-ft. roll
AP8845 Thionin-Chemical Demonstration Kit
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Bob Jones High School Chemistry I
Updated 3/23/2016
Demonstration: Test Tube Thunderstorm
You can react chemicals to produce what looks like a thunderstorm in a test tube. This is a spectacular
chemisty demonstration that is suitable for chemistry class or lab.
Test Tube Thunderstorm Safety
This demonstration involves corrosive acid, flammable alchohol or acetone, and a slight chance of
glassware shattering as a result of the vigorous chemical reaction. The test tube thunderstorm
demonstration should only be performed by qualified individuals, wearing full protective gear and using
proper safety precautions.
Test Tube Thunderstorm Materials
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95% alchohol (any type) or acetone
sulfuric acid
potassium permanganate
glass pipette
test tube
Perform the Test Tube Thunderstorm Demonstration
Wear gloves, a face shield, and protective clothing.
1. Pour some alcohol or acetone into a test tube.
2. Use a glass pipette to introduce a layer of sulfuric acid below the alcohol or acetone. Avoid any
mixing of the two liquids, since the demonstration won't work if too much mixing occurs. Do not
handle the test tube beyond this point.
3. Drop a few crystals of potassium permanganate into the test tube.
4. Turn out the lights. The sulfuric acid and the permanganate react to form manganese heptoxide,
which explodes when it comes into contact with the alcohol or acetone. The reaction looks bit like a
thunderstom in a test tube.
5. When the demonstration is concluded, inactivate the reaction by using metal tongs to place the test
tube into a large container of water. Be very careful! There is a chance the test tube could shatter.
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