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Table of ContentsChemistry Safety Contract .............................................................................................................. 3
Safe Laboratory Techniques ...................................................................................................... 10 Lab Notebook Guidelines ............................................................................................................. 13 Lab Equipment and Safety ............................................................................................................ 15 The Great Gas Plot ........................................................................................................................ 16
How Sweet It Is! ........................................................................................................................... 19 Half-Life Simulation ..................................................................................................................... 22 Excited Elements .......................................................................................................................... 23 Molecule Construction .................................................................................................................. 25 Evaporation and Intermolecular Attractions ................................................................................. 26
Percent Composition of Hydrates ................................................................................................. 26 Determining an Empirical Formula .............................................................................................. 26 Types of Chemical Reactions: A Sampler Platter........................................................................ 26
Solubility in Double Replacement Reactions ............................................................................... 26 Whats for dinner? Leftovers. ....................................................................................................... 26 Stoichiometry of a Precipitate ....................................................................................................... 26 LabQuest-Determination of Wavelength of Maximum Absorbance ............................................ 26
LabQuest Spectrophotometric Analysis of Aspirin ...................................................................... 26 Thin-Layer Separation of Lipstick ................................................................................................ 26
Determination of Melting Points .................................................................................................. 26 Molar Volume of a Gas................................................................................................................. 26 Airbags .......................................................................................................................................... 26
“Wet” Dry Ice ............................................................................................................................... 26 Properties of Solutions: Electrolytes and Non-Electrolytes......................................................... 26
39 Drop pH Lab ............................................................................................................................ 26 Titration Lab ................................................................................................................................. 26
Qualitative Analysis of the Group I Cations ................................................................................. 26 Copper into Gold: The Alchemist’s Dream ................................................................................. 26
Freezing Point Depression of a Solution (Ice Cream) .................................................................. 26 Periodic Table ............................................................................................................................... 26 Rules of Writing Equations, Solubility Rules, Activity Series of Metals ..................................... 26
Polyatomic Ions ............................................................................................................................ 26 Apples ........................................................................................................................................... 26 Avocado ........................................................................................................................................ 26
Tomato .......................................................................................................................................... 26 Jalapeno......................................................................................................................................... 26 Banana........................................................................................................................................... 26 Firehouse ....................................................................................................................................... 26
KFC ............................................................................................................................................... 26 Marcos........................................................................................................................................... 26 Outback ......................................................................................................................................... 26
Pistachios ...................................................................................................................................... 26 Basil .............................................................................................................................................. 26 Oregano ......................................................................................................................................... 26 Steak .............................................................................................................................................. 26 Brisket ........................................................................................................................................... 26 Ribs ............................................................................................................................................... 26 Chicken ......................................................................................................................................... 26
Sausage ......................................................................................................................................... 26 Pork Butt ....................................................................................................................................... 26
BACON......................................................................................................................................... 26 Tofu ............................................................................................................................................... 26 Nutmeg .......................................................................................................................................... 26 Salami ........................................................................................................................................... 26 Reuben .......................................................................................................................................... 26
Quinoa ........................................................................................................................................... 26 Turkey ........................................................................................................................................... 26 Stuffing ......................................................................................................................................... 26 Cranberry Sauce ............................................................................................................................ 26 Green Bean Casserole ................................................................................................................... 26
Mashed Potatoes ........................................................................................................................... 26 Pumpkin Pie .................................................................................................................................. 26
Reeses Cup .................................................................................................................................... 26 Nerds ............................................................................................................................................. 26 Milk Duds ..................................................................................................................................... 26 Candy Corn ................................................................................................................................... 26
Milky Way .................................................................................................................................... 26 Vanilla Tootsie Roll ...................................................................................................................... 26
Dots ............................................................................................................................................... 26 Cream of tartar .............................................................................................................................. 26
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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.Students are never
permitted in the science storage rooms or preparation areas unless given specific permission by their instructor.
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.
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25. If you or your lab partner are hurt, immediately yell out “Code one, Code one” to get the instructor’s attention.
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 manual and return to your
teacher)
QUESTIONS List any medical conditions
Do you wear contact lenses? of which the teacher should be aware.
___ YES ___ NO
Are you color blind?
___ YES ___ NO
Do you have allergies?
___ YES ___ NO
If so, list specific allergies
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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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 shoes that cover the entire foot
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.
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Cuts from Glass
Many cuts that occur in the laboratory are avoidable by following a few simple rules. You should
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.
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.
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The fire extinguishers available in lab are ABC extinguishers. This designation means that the
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 and shoes
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 tubing.
3. Close the air vents. Close the gas control valve at the bottom of
the burner and then open both about 1 ½ 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 triple-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.
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Heating a Liquid in a Beaker.
1. Place wire gauze on a hot plate.
2. Place a half-filled beaker of liquid on the wire gauze.
3. Turn on the hot plate and adjust the knob to 7.
4. Caution: Never heat plastic beakers or graduated glassware on a
hotplate. Never let a beaker boil dry; add water to the beaker as
necessary. Never adjust the hot plate to 10.
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. Never put chemicals directly on the pan. 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. There are brushes
available to clean off solids. 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.
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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Lab Notebook Guidelines
A lab notebook should be used to explain lab procedures, record all lab data, and show how
calculations are made. You may also use the notebook to discuss the results of an experiment and to
explain the theories involved.
A record of lab work is an important document which will show the quality of the lab work
that you have done. You may need to show your notebook and your lab reports to the Chemistry
Department at a college or university in order to obtain credit for the lab part of an AP Chemistry
class. As you record information in your notebook, keep in mind that someone who is unfamiliar
with your work may be using this notebook to evaluate your lab experience in chemistry. When you
explain your work, list your data, calculate values and answer questions, be sure that the meaning
will be obvious to anyone who reads your notebook.
Guidelines for the notebook:
1. Write your name and class on the front cover.
2. In black or blue ink, number all the right hand pages on the lower right corner if they are not
already numbered.
3. Save the first 2 pages for a Table of Contents. This should be kept current as you proceed. Each
time you write up a lab, place the title and page numbers where the lab report begins in the Table
of Contents.
4. Write in blue or black ink. Use only the right hand pages. The left hand pages can be used to do
calculations or other scratch work.
5. If you make a mistake, DO NOT ERASE OR SCRIBBLE. Just draw ONE LINE through your
error, and continue. It is expected that some errors will occur. A lab notebook is a working
document, not a perfect, error-free, polished product. Errors should be corrected by drawing one
line through the mistake, and then proceeding with the new data.
6. Do not use the first person or include personal comments.
Lab Reports (Lab reports will be worth 25 points)
Include the following information in your lab reports. Label each section
1. Title – The title should be descriptive. Experiment 5 is not a descriptive title.
2. Date – This is the date you performed the experiment.
3. Also record your lab partner and your lab station.
4. Purpose – A brief statement of what you are attempting to do in a complete sentence.
5. Procedure – A description of the method you are using which provides a summary of the
procedure. Do not include lengthy, detailed directions. A person who understands chemistry
should be able to read this section and know what you are doing.
6. Data-Record all your data directly into your lab notebook on the right-hand pages. Each person
must have all the data in their lab notebook and have it stamped by the teacher before you leave.
Organize your data in a neat, orderly form. Label all data very clearly. Use correct sig figs and
always include proper units. Underline, use capital letters or use any device you choose to help
organize this section well. Space things out – don’t try to cram everything on one page. Use
tables as much as possible. A data table must have a label and a title. e.g. – Table 1: Density
Values for Sugar Solutions.
7. Calculations and Graphs- You should show how calculations are carried out. Give the equation
used and show how your values are substituted into it. Give the calculated values. If graphs are
included, make the graphs an appropriate size. Label all axes and give each graph a title. If
experiments are not quantitative, this section may be omitted.
8. Conclusions – Make a simple statement concerning what you conclude from the experiment.
This is not a place to give your opinion of the lab and whether or not it was “fun”. It is not
your job to review the lab like you would if you saw a movie.
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9. Experimental error – If there is a known value for something you are doing in lab, calculate
the experimental error.
Accuracy
Accuracy is a measure of how close an experimental value is to a value which is
accepted as correct. The measure of the accuracy of an experimental value is reported as
Percent Error.
Accepted
AcceptedalExperimenterror
%
10. Error Analysis – What are some specific sources of error, and how do they influence the
data? Do they make the values obtained larger or smaller than they should be? Which
measurement was the least precise? Instrumental error and human error exist in all
experiments, and should not be mentioned as a source of error unless they cause a significant
fault. Significant digits and mistakes in calculations are NOT a valid source of error. In
writing this section it is sometimes helpful to ask yourself what you would do differently if
you were to repeat the experiment and wanted to obtain better precision.
11. Questions – Answer any questions included in the lab directions. Answer in such a way that
the meaning of the question is obvious from your answer.
Reporting Lab Data
Graphing Data
1. All graphs should have a descriptive title (“Graph” is not a title) and a label. e.g. – Graph A:
Density of Solutions with Varying Sugar Concentrations.
2. Both the vertical and horizontal axes should both have labels and units clearly marked. Use a
ruler to draw the axes.
3. The scales chosen should reflect the precision of the measurements. For example, if
temperature is known to be ±0.1ºC, you should be able to plot the value this closely. Don’t
have each block of the graph equal to 10ºC.
4. There should be a table in which the data values are listed. Don’t put data in a graph unless
you have first listed it in a table.
5. The controlled or independent variable is placed on the horizontal axis. The dependent
variable is graphed on the vertical axis.
6. There should be an obvious small point on the graph for each experimental value. It is not
necessary to include the coordinates of each point since they will be in the data table.
7. A smooth line should be drawn that lies as close as possible to most of the points. Do NOT
draw a line connecting one point to the next as in a dot-to-dot drawing. If the line is a
straight line, use a ruler to draw it.
8. If a computer program is used to draw the graph, the rules still apply.
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Lab Equipment and Safety
No writeup in lab notebook.
Discussion:
In this activity you will become familiar with some proper lab procedures used in the Chemistry lab.
Procedure (Write the answer to each question on a sheet of paper, not your lab notebook)
Measuring the volume of a liquid.
1. Read “Using a graduated cylinder”. See your TOC.
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. Add some water to your graduated cylinder. 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” in this lab manual. See your TOC.
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 the lab manual about using a burner.
2. Read over the fire and thermal burns safety considerations in this lab manual.
3. Follow the steps in this lab manual 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.
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. 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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The Great Gas Plot
Write up in your lab notebook.
Using Balloons and Graphs to Analyze Relationships
Data collected in scientific investigations is often graphed and analyzed to establish meaningful
relationships between variables such that explanations can be constructed and predictions made. In
all science, understanding the effect of material quantities on various outcomes is essential, such as
seen in the following scenarios.
A chemist or chemical engineer optimizes a process to synthesize durable substances from less
expensive, readily available materials.
A biologist investigates the effects of nutrient availability on the life cycle or growth potential of an
organism.
• A geologist studies how pH levels in rainfall affects the weathering of exposed rock, which can be
linked back to certain automobile emissions.
• An environmentalist evaluates the impact of leakage from a previously capped oil well on the
biodiversity and health of plants in the area.
• A physicist explores mechanisms for storing solar or geothermal energy through processes
involving phase changes or other reversible reactions.
In this experiment, you will be exploring the specific relationships between quantities of reactants
and products involved in the chemical reaction between vinegar and baking soda. A successful
outcome in this experiment is dependent on replicable data collection and your ability to construct
and analyze a graph. Clear thought processes and well-written responses contribute to your success
in this task.
In each trial of this experiment, you will use balloons to capture the gas generated by the reaction. A
graph of the volume of gas produced versus the amount of baking soda reacted will be plotted and
analyzed to make inferences and draw conclusions about conditions that affect the relationship
between these two quantities.
PROCEDURE
1. Create a table for the data you collect in this lab using the table below as an aid. Label eight
balloons with a quantity of baking soda, NaHCO3, in 0.5-g increments, starting with “0.5 g” and
continuing until the last balloon is labeled with “4.0 g”.
2. Using an electronic balance, measure out the required quantity of baking soda corresponding to
each balloon.
3. Carefully transfer the baking soda into the appropriately labeled balloon using a funnel. Repeat
this process until you have placed the appropriate quantity of baking soda in each of the labeled
balloons.
4. Add water to the large container until it is mostly full. Carefully place the container filled with
water inside the overflow pan. Slowly fill the remainder of the container with water. Take care not
to spill any water into the overflow pan.
5. Obtain a sample of vinegar and load it into the syringe. To do so, completely compress the
plunger on the syringe, place the tip of the syringe below the surface of the vinegar, and then pull
up on the plunger until the syringe contains more than 30 mL of vinegar. Carefully push the
plunger in until the syringe contains exactly 30 mL of vinegar.
6. Select one of the labeled balloons and stretch the neck of the balloon over the lower barrel of the
syringe. Before releasing the balloon, squeeze as much air as possible out of the balloon. Make
sure the balloon is secured tightly around the barrel of the syringe.
7. Push the plunger down and deliver all 30 mL of vinegar into the balloon. Take care to not push
any of the air out of the syringe—only depress the plunger enough to deliver all of the vinegar
into the balloon.
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8. Working with a partner, carefully twist the balloon, restricting its neck as close to the syringe as
possible. Snap the balloon clip closed around the neck, trapping the gas in the balloon. For added
insurance against a leak, tie a knot in the neck of the balloon.
9. Shake the balloon and allow the vinegar and baking soda to react. Continue shaking until the
balloon returns to room temperature.
10. Place the balloon in the container filled with water and carefully press down with a flat object
until the balloon is completely submerged and the flat object is sitting firmly against the top of the
container.
11. Carefully remove the balloon and the container from the overflow pan. Measure the amount of
water that overflowed into the pan and record this value in your data table.
12. Pour the water from the overflow pan back into the container, and then place the container back
inside the overflow pan. The water should completely refill the container unless some was spilled.
13. Repeat Step 4 through Step 12 until you have measured the volume of water displaced by all
eight balloons.
Table 1. Measuring Displacement
Mass of Baking Soda Added
(g)
Volume of Vinegar Added
(mL)
Volume of Water Displaced
(mL)
0.50 30
1.00 30
1.50 30
2.00 30
2.50 30
3.00 30
3.50 30
4.00 30
14. On your lab table, arrange the balloons in a logical sequence that establishes a correlation
between the volume of gas produced and the amount of baking soda used. Sketch, label, and
create a caption for the relationship depicted by your arrangement in the space provided2. In
Table 1, label the “Volume of Water Displaced” column as “Volume of CO2 Produced,” as well.
State your reasoning for why those two measurements are the same.
15. Use a graphing program to plot a graph of the volume of carbon dioxide gas produced versus the
mass of baking soda.
a. Provide axes labels and include units along with the title.
b. For any and all regions of the graph that appear to be linear, perform a linear fit on the data
using the program’s curve-fitting function. Be sure to display and record the equation(s) for the
line of best fit in y = mx + b format.
16. Compare your graph to the arrangement you sketched in Question 1. Explain how both
representations are useful to the experimenter. Be sure to offer at least one way that each
representation provides information the other does not.
Questions:
1. On your graph, you should notice an obvious change in the relationship that exists between the
amount of baking soda added and the volume of gas produced. Cite where this occurs and
explain how the relationship changed even though the amount of baking soda was increasing and
the amount of vinegar remained constant. Justify your explanation.
2. The reaction between vinegar and baking soda caused a change in the temperature of the balloon.
Based on your observations, justify whether the reaction is exothermic or endothermic.
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3. A student measured the volume of the balloon by displacement without allowing the balloon to
return to room temperature. Describe the effect this would have on the volume of the balloon.
Justify your answer.
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How Sweet It Is!
Most have heard the saying, “An apple a day keeps the doctor away.” Can the same be said for apple
juice? Certain beverages have a really bad name with regard to health, particularly the health of
children, and that name is surprisingly not “calories” but “carbohydrates.” Some schools, districts,
and even cities are taking a stand against the accessibility of many sweet treats on campuses. As a
Pre-Lab Exercise for this investigation, arm yourself with a little background information. Which
drinks are getting cut in schools (perhaps yours), which get to stay, and why? Consider what the
American Beverage Association has to say and whether it lines up with other health organizations
such as the Centers for Disease Control and Prevention (CDC). In this investigation, you will
conduct a comparative study to experimentally determine the “carb” concentration of store-bought
beverages. Using your data, you will develop a position about what products should be accessible to
students at Bob Jones. To take quantitative measurements, you will create a tool called a hydrometer
using a straw and some simple materials. Before you can conduct your study, you will first need to
make the hydrometer and determine how its behavior in liquids provides you with useful
information.
PART I: CONSTRUCT THE HYDROMETER
1. Obtain a straw, ruler, and marker. Starting from one end, mark the straw at 0.5 cm increments
along three quarters of the straw’s length.
2. Fold down approximately 1 cm of straw at the end that is not marked. Slide a small paper clip
over the fold until half the length of the clip is on the straw. To secure the clip in place and
ensure the straw’s end stays folded, wrap a piece of transparent tape tightly around the bottom of
the straw. The end of the straw that is now sealed will be referred to as the “bottom” of the straw.
The open end of the straw represents the “0” mark.
3. Mark every five lines moving down the straw such that you can read the numbers when the straw
is vertical with the open end facing up.
4. Using materials provided by your teacher, attach an appropriate amount of ballast to the bottom
of your straw. Your ballast should be compact enough to fit inside a 100-mL graduated cylinder
and heavy enough to make the hydrometer float upright in the cylinder when filled to the 100-
mL mark. In distilled water, the bottom of the hydrometer with ballast should be approximately 1
in. from the bottom of the cylinder but not more than 2 in.
5. To ensure that the hydrometer is floating freely, give the straw a spin. Take a reading of the
distilled water with the hydrometer to the nearest tenth of an interval. Record your value
PART II: CALIBRATE THE HYDROMETER
6. What relationship exists between the hydrometer readings and the percent mass/mass (%m/m) of
sugar in solution? Design an experiment that will allow you to answer this question. In your
investigation, you may use up to 100 g of sugar. Record this procedure in your lab notebook in
the procedure section.
7. Perform your procedure and record all your data in an appropriate data table in your lab
notebook.
8. Using excel or some other graphing tool graph your data from part II. Identify the dependent and
independent variable. Either copy the graph into your lab notebook or print out and tape into the
lab notebook. Calculate and record the slope intercept formula for the line of best fit.
PART III: USE THE HYDROMETER
9. Consider the assortment of juices, sodas, and beverages provided by your teacher. The nutrition
labels have been removed but you have been provided with all names and brands.
10. Use your hydrometer to collect data on each teacher provided sample and record that data. Use
your graph to calculate the percent sugar in each sample.
Questions:
1. Would a straw with a larger diameter require more or less ballast? Justify your answer.
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2. If poured slowly enough, you can float a layer of ethanol in water. Predict how the hydrometer
would behave in a sample of pure ethanol compared to pure water, and explain your reasoning.
3. What is the significance of the y-intercept on your graph?
4. Identify the physical property of the solution that is responsible for the trend in hydrometer
readings as concentration increases.
5. Obtain the nutrition facts from your teacher for the beverages tested, does the data from the
nutrition facts confirm your experimental results? Why or why not?
6. All the soft drinks tested were allowed to go “flat.” If freshly opened soda were used in this
experiment, would the calculated concentration be greater than, less than, or the same as a “flat”
sample? Justify your answer.
7. Compose a position about what products should be accessible to students at Bob Jones. Justify
your position using evidence from your research and your investigation.
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Calorimetry of Various Foods
Write up in your lab notebook.
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. Food is rated in Calories (notice the capital C) A Calorie is really a
kilocalorie so you need to know this general fact: 1.000 Calorie = 1000. calories = 4.180 kiloJoules.
Procedure:
1. You will need to find the mass of the food before and after burning, the flask empty and with
water and some other data. The procedure is up to you. After you have written your procedure
bring it to the teacher and get it approved before moving on with the experiment.
2. Add cold water to the flask (amount is up to you).
3. 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.
4. Determine the mass of the unburned food and calculate the mass of food burned.
5. Calculate the average amount of calories and joules given off by the food and absorbed by the
water.
6. Calculate how calorie dense the food sample is (calories/grams of burned food sample)
7. Post your calorie density for your food by following your teacher’s directions on edmodo.
Cleanup:
1. All food remnants get thrown away.
2. All water gets put down the sink.
3. The Erlenmeyer flask should be rinsed and the thermometer left at your lab station.
4. 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?
Page 22 of 127
Half-Life Simulation
Data, Calculations and Questions only
Discussion:
Every element on the periodic table has one or more radioactive isotopes. Recall that isotopes are
atoms of the same element that differ in the number of neutrons their nucleus contains. Depending
on the ratio of neutrons to protons, these isotopes may be stable or radioactive. Charts have been
produced that identify the band of stability (Figure 1). Isotopes falling within the band are stable
whereas those outside the band are radioactive. Notice that as the elements get heavier, the neutron
to proton ratio drifts greater than 1:1. For those isotopes whose neutron to proton ratio lands them
outside the band of stability, the nucleus will undergo radioactive decay until a stable atom is
formed. The amount of time it takes for a sample to decay is specific to the type of atom that is
decaying. The amount of time it takes for one half of a radioactive sample to decay is called its half-
life. Half-lives can range from fractions of a second to millions of years.
.
In this activity, you will model radioactive decay with beads. The analysis of data and the
determination of half-life will be done by graphical means.
PROCEDURE
1.Count the beads in your cup. Record this number in the data table in your lab notebook. This is the
value for 0.0 seconds.
2.Put the beads in the cup and then pour them out onto the paper plate. Remove the beads that landed
in the starred section. These beads will be considered “decayed.” Replace the decayed beads with
the same number of the other color of beads you were given so that the total number of particles
remains constant. Count the remaining candies and record this number in your data table. Each
trial is to be counted as 10 seconds.
3.Continue this procedure until you have 10 trials or until you have fewer than 5 beads left,
whichever comes first. Record each trial in your data table.
4.Construct a graph of time in the x axis and number of beads remaining on the y axis. Draw a
smooth curve line of best fit. Include in your graph title the numbers of sections on your paper
plate
5.Using your graph, calculate a half life.
6.Compare your results with those of other groups. What effect did having different numbers of
segments in the paper plate have on half life?
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Excited Elements
No writeup required
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. 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 bright-line 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.
Purpose:
The purpose is to observe the characteristic elemental spectra produced by applying high voltage
across a sample of a gas at very low pressure.
Materials/Equipment:
High voltage power supplies Spectral tubes
Vernier LabQuest Spectrophotometer
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 because of the danger of
UV ray exposure. The rule of thumb is 30 seconds on/ 30 seconds off. The tubes will also get quite
hot.
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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.
5. Obtain a Labquest system. Connect the Labquest to the fiber optic cable as instructed by
your teacher. Record the emission spectrum of your assigned tube using the Labquest
system.
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Molecule Construction
For the following compounds you will compile the information below over the next
couple of days. You will initially write all your work on scrap paper and it will be
copied into your lab notebook later.
A. formula
B. lewis structure
C. electronegativity difference of EACH type of bond
D. non polar, polar or ionic quality of EACH type of bond
E. molecular geometry (# 15 and 19 will have multiple geometries)
F. Hybridization of EACH atom
G. Count of pi and sigma bonds in each molecule
1. H2
2. Cl2
3. O2
4. N2
5. HCl
6. BrCl
7. water
8. carbon dioxide
9. H2S
10. BF3
11. NH3
12. CH4 (methane)
13. C2H4
14. CCl4
15. CH3COOH (vinegar)
16. C2H2 (acetylene)
17. CH3CH2OH (ethanol)
18. H2CO (formaldehyde)
19. H3CCOCH3 (acetone)
20. HCN (cyanide gas)
21. SO3.
22. Nitrate (see end of lab manual)
23. Carbonate
24. CO
25. ozone (O3)
Day 1 On scrap paper because it will be a hot mess for some of them, draw the
lewis structure for all 25. Calculate item C and D above for each molecule
Day 2 Use kits and phet to model each molecule to determine its geometry(ies) and
F and G from the list above.
Once you have all the information compiled, record neatly in your lab notebook.
Madison City Schools PreAP Chemistry Updated 7/29/2015
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Evaporation and Intermolecular Attractions
Write up in your lab notebook.
Discussion:
In this part 1 of 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 London dispersion forces.
Part 2: The interaction of particles governs the physical world. The competing forces of attraction
and repulsion between subatomic particles are responsible for properties of the atom that include
size, ionization energy, and electronegativity.
Between two atoms, those same forces influence the way valence electrons associate with each atom.
Electrons can be transferred or shared to create an alliance, or bond, between two atoms. When
electrons are transferred, the resulting ions are so strongly attracted to each other that they arrange
themselves into a crystalline structure that maximizes those attractions. These interactions can all be
understood qualitatively by Coulomb’s law: The force of the attraction, or repulsion, is related to the
magnitude of the charges, q1 and q2, and the distance between the particles (Figure 1).
The same principle that holds cations and anions together also gives us a way to understand the
physical properties of substances. Metallurgy, cooking, the development of plastics and polymers,
petroleum engineering, and even candle making are examples of where differences in the physical
properties of the materials used and created have application.
In physical processes like phase changes, collections of particles interact without any change in
chemical composition. The nature of these interactions depends on a set of factors that influence the
degree of coulombic attractions between particles. Understanding these factors allows us to make
informed predictions about the relative properties of pure substances, including the melting point. SDS: methanol Toxic by ingestion (may cause blindness), inhalation or absorption.
Irritating to body tissues. Avoid body tissue contact. Flammable liquid. ethanol Toxic by ingestion and inhalation. Body tissue irritant. Avoid all body
tissue contact. Denatured with isopropanol and methanol. Not for human
Madison City Schools PreAP Chemistry Updated 7/29/2015
Page 27 of 127
consumption. Flammable liquid. 1-propanol Severe eye and skin irritant. Slightly toxic by ingestion, inhalation and
skin absorption. Avoid all body contact. Flammable liquid. 1-butanol Moderately toxic by inhalation or ingestion. Irritant to body tissue.
Absorbed through the skin. Avoid vapors. Flammable liquid. n-pentane Irritating to body tissues. Avoid body tissue contact. Vapor is narcotic in
high concentrations. Flammable liquid. n-hexane Irritant to body tissues. Mildly toxic by inhalation. Avoid all body
contact. Flammable liquid. Paraffin Substance is not considered hazardous. Not for topical use. Combustible
solid. Wax does not always smoke before it ignites. If melting wax use a thermometer.
Sucrose Substance is not considered hazardous. Not for human consumption or use.
Dextrose Substance is not considered hazardous. Not for human consumption or use.
Salt Very slightly toxic by ingestion. Dust may cause minor irritation to mucous membranes upon inhalation. Not for human use.
PRE-LAB EXERCISE
Part 1:Prior to doing the experiment, complete the Pre-Lab table in your lab manual. The name and formula are given for each compound. Look up the structural formula for a molecule of each compound on the internet. Then determine the molecular weight (molar mass) 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.
Part 2: Answer the following 5 questions on the left hand page of the lab notebook where you wrote your Purpose, procedure and data table.
1. Within an atom, list the forces that exist between charged particles and categorize them as attractive or repulsive forces.
2. Explain why a strong attraction occurs between particles as a result of the transfer of an electron from the less electronegative atom to the more electronegative atom.
3. Using Coulomb’s law, explain why energy is required to move two positively charged particles closer together.
4. Define the phrase “physical properties” in your own words, and give two examples.
5. In your own words, describe what is happening on a particulate level during the melting of a solid.
PROCEDURE
Part 1: 1. Wrap 2 electronic thermometers 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.
Madison City Schools PreAP Chemistry Updated 7/29/2015
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2. 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.
3. After the probes have been in the liquids for at least 45 seconds, begin data collection by reading the thermometer or probe. 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.
4. When both temperatures have reached minimums and have begun to increase, end data collection. Record the maximum (t1) and minimum (t2) values for Temperature 1 and Temperature 2.
5. For each liquid, subtract the minimum temperature from the maximum temperature to determine t, the temperature change during evaporation.
6. 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.
7. Repeat Steps 2-7 for the other samples. 8. 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. Part 2: 1. In the data section of your lab notebook, write a hypothesis predicting the order in which the
salt, paraffin, sucrose, and glucose will melt.
2. Take a piece of aluminum foil and shape it into a dish as demonstrated by Mr. Elegante. Use the
sharpie to divide the lid into four quadrants as shown in Figure 2.
1. Use scoopulas to transfer small samples of each of the four compounds to the outermost groove
of the can lid. You only need a very small sample, just enough to be able to see the compound.
2. Carefully place the can lid onto an unplugged hot plate. Plug in the hot plate and turn the heat up
to a medium setting.
3. Watch carefully as the compounds melt. Record your observations. You will need to record the
relative order of melting. As soon as three of the compounds melt, turn off the hot plate and
unplug it.
4. In a table titled Properties of Four Substances , write the names and formulas of the four
substances in part 2 in order of increasing strength of intermolecular forces. The columns of
the table should be: Name ,Formula , Bond type , Polarity , Molar Mass, and Type of
Intermolecular forces
5. Identify the bond type as either covalent or ionic. Finally, complete the rest of the table for
covalent molecules only.
Cleanup: Save all the rubber bands, throw all the pieces of filter paper away, turn off the thermometers and return the bottles to the front desk. Allow the can lid to cool and dispose of it in the trash
Questions
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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. An intermolecular force is defined as the force of attraction between neighboring molecules.
Considering this definition and what you know about melting points, what is the general
relationship that exists between the strength of intermolecular forces and melting point?
5. Coulomb’s law describes why oppositely charged particles are attracted to each other. Using the concept of dipoles, explain why the attractive forces between neighboring molecules are not as strong as the attractive forces between ions.
6. From what you know about attractive forces between particles, explain the relative order of melting points for the four substances investigated in part 2.
7. In determining the strength of the intermolecular forces, it is important to consider the net forces present. Use this argument to defend why water, despite having hydrogen bonding, has a much lower melting point than paraffin.
8. Hydrogen sulfide is a gas at room temperature whereas water is a liquid, yet hydrogen sulfide has more electrons than water. Explain this anomaly.
9. Consider the halogens at room temperature and 1 atmosphere of pressure. Explain why fluorine and chlorine are gases at room temperature whereas bromine is a liquid and iodine is a solid.
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PRE-LAB Table A: INFORMATION ON DIFFERENT ORGANIC SOLVENTS
AND THEIR HYDROGEN BONDING CAPABILITIES
Substance Formula Structural Formulas Molecular
Weight
Hydrogen Bond
(Yes or No)
ethanol C2H5OH
1-propanol C3H7OH
1-butanol C4H9OH
n-pentane C5H12
methanol CH3OH
n-hexane C6H14
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Percent Composition of Hydrates
Data, Calculations and Questions only
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:
hydrateofmass
waterofmassOH
__
__% 2
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.
The name of a hydrate follows a set pattern: the name of the ionic compound followed by
a numerical prefix and the suffix "-hydrate." For example, CuSO4 · 5H2O is "copper(II)
sulfate pentahydrate." When the chemical formula for a hydrated ionic compound is
written, the formula for the ionic compound is separated from the waters of hydration by
a centered "dot". The dot means “is associated with” and does NOT mean multiply. The
notation of hydrous compound · nH2O, where n is the number of water molecules per
formula unit of the salt, is commonly used to show that a salt is hydrated. The n is usually
a low integer, though it is possible for fractional values to exist. In a monohydrate n is
one; in a hexahydrate n is 6, etc. (Typical prefixes are mono-, di-, tri-, tetra-, penta-,
hexa-, hepta-, octa-, nona-, and deca-
SDS:
MgSO4 Avoid inhalation. May irritate eyes and respiratory tract. Avoid body tissue
contact.
CuSO4 Slightly toxic by ingestion. Body tissue irritant. Avoid all body tissue contact.
BaCl2 Highly toxic by ingestion and inhalation. All soluble barium compounds are
poisonous if swallowed and cause nausea, vomiting, stomach pains and
diarrhea.
Procedure:
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.
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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.
6. While the crucible is cooling, weigh approximately 2.10-2.20g of MgSO4 or 3.00-
3.20 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.
Cleanup:
1. Empty the cooled crucible into the garbage can.
2. Wipe the cooled crucible out with a dry paper towel.
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
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Determining an Empirical Formula
Data, Calculations and Questions only 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 24.0-
gram 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. SDS
Magnesium Flammable solid. Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
PROCEDURE:
1. Heat a crucible in the hottest part of a burner flame for 3 minutes without the lid on. Turn off
the burner and cool 3 minutes.
2. Measure the mass of the empty crucible and record in data table.
3. Obtain approximately 10 cm of magnesium ribbon. Clean the magnesium ribbon with steel
wool. Observe and record the physical characteristics of the magnesium ribbon. Coil the
ribbon so it will fit in the bottom of the crucible, 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. Measure and record the mass of the crucible and contents (without the lid). Observe and
record the physical characteristics of the newly formed compound.
7. Calculate the empirical formula of the compound. Show your work.
Clean Up
Wipe the crucible out with a dry paper towel once it is cool. The ash can go in the trash can.
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
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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.
3. What was the purpose of using steel wool?
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Types of Chemical Reactions: A Sampler Platter
Write up in your lab notebook.
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.
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SDS:
Ethanol Toxic by ingestion and inhalation. Body tissue irritant. Avoid all
body tissue contact. Denatured with isopropanol and methanol. Not
for human consumption. Flammable liquid.
CaCO3 Irritant to body tissues. Severe eye and moderate skin irritant.
KI solution Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
Pb(NO3)2 solution 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.
Mg Substance not considered hazardous. However, not all health
aspects of this substance have been thoroughly investigated.
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.
6M HCl 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.
Write a balanced equation including states of matter for each reaction.
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.
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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 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 inverted tube.
A pop indicates the presence of hydrogen gas. Note the appearance of the substance
in the first test tube.
Clean Up:
1. Empty the tube with calcium carbonate into the trash can but DO NOT wash.
2. Empty the tube with the lead nitrate in it into the waste beaker on the back bench.
3. Dispose of the magnesium ash and any wooden splints in the garbage can.
4. 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. In this experiment, describe the method that was used to test for the presence of
CO2 gas.
2. Describe the test used to check for the presence of hydrogen gas.
3. 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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Solubility in Double Replacement Reactions
Data, Calculations and Questions only
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. It is possible to write a different kind of chemical reaction called a net
ionic equation. It shows only the ions actually changing in a reaction and leaves out
those that do not change. The ones that do not change are called “spectator ions”. To
write a net ionic equation you separate anything that is (aq) and ionic or a strong acid into
separate ions. Anything that remains unchanged in the reaction can be marked out. For
example:
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) is an equation for a precipitate forming.
Separate the (aq) that are ionic and you get:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) AgCl(s) + Na+(aq) +NO3
-(aq)
Cancel out the nitrate ions and sodium ions because they do not change at all and you’re
left with:
Ag+(aq)+ Cl-(aq) AgCl(s)
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 ions 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.
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SDS:
0.1M AgNO3 Moderately toxic by ingestion. Irritating to body tissues. Avoid all
body tissue contact. Will stain clothing and skin.
0.1M BaCl2 Highly toxic by ingestion and inhalation. All soluble barium
compounds are poisonous if swallowed and cause nausea, vomiting,
stomach pains and diarrhea.
0.1M NaOH Slightly toxic by ingestion and skin absorption. Severely irritating to
body tissues. Causes severe eye burns. Avoid all body tissue contact.
0.1M CuSO4 Mildly toxic by ingestion. Irritant to skin, eyes and mucous
membranes. Avoid contact with body tissues.
0.1M CoCl2 Toxic by inhalation and injestion. Severe corrosive to all body tissue,
especially skin and eyes. Avoid contact with all body tissues.
0.1M K3PO4 Body tissue irritant.
0.1M Pb(NO3)2 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.
0.1M Fe(NO3)3 Corrosive to body tissues by contact and inhalation. Avoid contact
with skin, eyes and mucous membranes.
0.1M BiCl3 Moderately toxic by ingestion and inhalation. Corrosive to body
tissues. Avoid body tissue contact.
0.1M Na2CO3 Irritating to body tissues.
0.1M MnCl2 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. Write a net ionic equation for each equation that produces a precipitate.
2. At each workstation there will be a sheet of white paper with a grid and black ovals
on it in a protective sleeve. Use the white sheet of paper as a guide.
3. 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.
4. 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 of BaCl2 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.
5. On the data sheet describe each precipitate in terms of amount, 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:
1. Wipe the sheet with a paper towel, rinse it off and dry with another paper towel.
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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 BiCl3 Na2CO3
Fe(NO3)3
NaOH
BiCl3
Na2CO3
MnCl2
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Whats for dinner? Leftovers.
Write up in your lab notebook.
Discussion:
In this activity, you will design an experiment to determine which reactant is the limiting
reactant, aluminum metal or copper(II) chloride solution and which is leftover. You will
calculate the experimental molarity of the copper(II) chloride solution and percent error
for this value given the theoretical value obtained from your teacher. You will then
calculate the percent yield for the solid copper produced.
SDS:
Copper (II) chloride Harmful if swallowed. Do not eat, drink or smoke when using this
product. Causes skin irritation.
Aluminum Flammable solid. Keep away from heat, sparks, open flames, and
hot surfaces..
Procedure
1. Watch the podcast “Aluminum Leftovers” of possible steps to include in your
experiment. Be sure to write down pertinent information. With a lab partner, design a
procedure that will allow you calculate the limiting reactant. Write steps for
conducting your experiment in the procedure section of your lab writeup. Obtain
teacher approval before beginning.
2. Write the balanced chemical equation for this reaction including state symbols for
each reactant and product.
3. Was the reaction endothermic or exothermic? Cite evidence from the experiment to
support your claim. Write the energy term on the appropriate side of the equation
written in step 2.
4. Refer to the balanced equation written in step 2 and the observations made in the
experiment to complete the pictorial representation in Figure 1 for the “final” beaker.
Copy this figure into your data section of your lab notebook.
5. When all of the liquid has passed through the filter, carefully remove the paper from
the funnel; spread it out on a hotplate set on medium to dry. When you think it is dry,
mass the filter paper with its contents, put it back on the hotplate for one more
minutes, remass the filter paper and its contents to ensure all the water has
evaporated. Be sure to record all data.
6. Make a statement in the data section indicating the limiting reactant? What evidence
do you have to support this claim?
7. Calculate the experimental molarity of the copper(II) chloride solution your teacher
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provided. Molarity is the number of moles of copper (II) chloride divided by the
Liters of solution you used.
8. Obtain the theoretical molarity of the copper(II) chloride solution from your teacher
and calculate the percent error.
9. Calculate the theoretical mass of copper.
10. Calculate the percent yield for the copper produced in your experiment.
Cleanup:
Wash all dishes and equipment with soapy water and rinse well. Refill the wash
bottle if needed for the next class. Leave the equipment at your lab station on a paper
towel for the next class.
Questions
1. What type of results would you expect to see if you had placed solid copper into a
solution of aluminum chloride? Explain your reasoning.
2. A student forgets to dry the piece of aluminum before measuring the final mass.
Explain how this error will affect the calculated molarity of the copper(II) chloride
solution. Clearly state whether the calculated molarity will be too large, too small, or
unchanged by this error and include mathematical justification for your answer.
3. The copper obtained in an experiment was allowed to dry for several days. Before
weighing the sample, the student noticed that some of the copper flakes appeared a bit
green. Will this color change increase, decrease, or not change the percent yield of
copper produced in the experiment? Clearly justify your claim with appropriate
evidence.
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Stoichiometry of a Precipitate
Write up in your lab notebook.
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:
SDS
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.
1 M 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 12 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
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before you leave the lab.
12. Record your observations of the precipitate. 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.
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 and leave at your station.
Questions:
1. What would happen to the percent yield if there was still water in the precipitate?
Justify your answer.
2. If there was some precipitate left in the beaker, what effect would it have on the
percent yield? Justify your answer.
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LabQuest-Determination of Wavelength of Maximum Absorbance
Write up in your lab notebook.
Discussion:
Colored solutions are colored because they absorb certain wavelengths of light while
allowing other wavelengths of light to pass through. As observers, we see the
wavelengths of light that are not absorbed. Each solution will have one or more
wavelength at which it absorbs strongly. The wavelength of light at which a colored
solution absorbs the most is called its “wavelength of maximum absorbance.” Light at
this wavelength can be used to analyze the solution in spectrophotometric analysis.
Getting Started: Assemble the LabQuest/Spectrometer system. Make sure to use the
AC power adapter to power the LabQuest unit. The spectrometer connects to the
LabQuest unit with a USB cable. Once assembled turn on the LabQuest unit and allow it
to go thru its boot process. Once finished you should see a meter screen with an orange
background with “USB:Abs” listed as the unit of measurement.
Calibration: Use the stylus to tap “Sensors”, “Calibrate”, and “USB Spectrometer.”
The “Calibrate Spectrometer” window will open. The lamp in the spectrometer will
warm for 60 seconds. Once prompted place a cuvette filled half full with tap water into
the sample holder on the spectrometer. Your teacher will describe the proper alignment
for the cuvette. Use the stylus to tap the “Finish Calibration” button. When you get the
“Calibration Completed” message, tap OK to return to the meter window.
Procedure:
1. Use a disposable pipet to fill a plastic cuvette approximately half full with each of
the colored ionic solutions. Record the color and concentration of each solution on
the data sheet. Make sure to wipe any excess liquid off the exterior of the cuvette
with a paper towel.
2. Insert the first cuvette into the sample holder on the spectrometer. Press the
“Collect/Stop” button with your finger. It looks like a “PLAY” button on a DVD
player. Once the spectrum is acquired, press the “Collect/Stop” button again. Be
Patient! It takes a moment for the process to stop. Do not press the button twice. The
graph will autozero and the wavelength of maximum absorbance will be displayed.
Record the absorbance value, the wavelength, and the color of the wavelength of
max absorbance on the data sheet.
3. Remove the cuvette containing the first sample from the spectrometer and insert the
cuvette containing the second sample. Press the “Collect/Stop” button. You will be
asked whether you wish to store or discard the previous sample data. Use the stylus
to select “Discard.” Once the spectrum is acquired, press the “Collect/Stop” button
again. The graph will autozero and the wavelength of maximum absorbance will be
displayed. Record the absorbance value, the wavelength, and the color of the
wavelength of max absorbance on the data sheet.
4. Repeat step 3 for the remaining samples.
5. Predict the wavelength of max absorbance of the yellow pH=7.00 buffer solution.
Measure the wavelength to check your prediction.
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6. Predict the wavelength of max absorbance of the reddish pH=4.00 buffer solution.
Measure the wavelength to check your prediction.
Cleanup:
Rinse out all cuvettes and leave on them a paper towel at your station to dry. Your
teacher will give you instructions on using the labquest.
Questions:
1. If the wavelength of max absorbance appears in the white area to the right of the
color spectrum on the graph window what does this tell us about the type of
electromagnetic radiation best absorbed by the sample?
2. Which solution showed the greatest analytical sensitivity? Hint: Divide the
absorbance by the concentration.
3. Why does greenish NiCl2 have a similar wavelength of max absorbance to that of
yellowish FeCl3? Hint: What two primary colors combine to make green?
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LabQuest Spectrophotometric Analysis of Aspirin
Write up in your lab notebook.
Discussion: Aspirin is the common name for acetylsalicylic acid. Acetylsalicylic acid has the
molecular formula C9H8O4. Aspirin reacts with sodium hydroxide according to the
following equation:
C9H8O4 + 3NaOH (aq) Na2C7H4O3 (aq) + NaC2H3O2 (aq) +
2H2O Aspirin sodium hydroxide sodium salicylate sodium acetate
water
If an acidified solution of iron (III) nitrate is added the following reaction takes place:
Na2C7H4O3 (aq) + Fe(NO3)3 (aq) FeC7H4O3+ (aq) + 2Na+ (aq)
+ 3NO3- (aq)
sodium salicylate iron (III) nitrate salicylatroiron (III) complex
The salicylatroiron (III) complex is a purple colored substance that has a max absorbance
at or near 536nm (in the blue-green region of the visible light spectrum). In this
experiment we will take an off the shelf pain reliever and measure its aspirin content
using a spectrometer.
Getting Started: Assemble the LabQuest/Spectrometer system. Make sure to use the
AC power adapter to power the LabQuest unit. The spectrometer connects to the
LabQuest unit with a USB cable. Once assembled turn on the LabQuest unit and allow it
to go through its boot process. Once finished you should see a meter screen with an
orange background with “USB:Abs” listed as the unit of measurement.
Spectrometer Calibration: Use the stylus to tap “Sensors”, “Calibrate”, and “USB
Spectrometer.” The “Calibrate Spectrometer” window will open. The lamp in the
spectrometer will warm for 60 seconds. Once prompted place a cuvette filled half full
with tap water into the sample holder on the spectrometer. Your teacher will describe the
proper alignment for the cuvette. Use the stylus to tap the “Finish Calibration” button.
When you get the “Calibration Completed” message, tap OK to return to the meter
window.
Preparation of a Calibration Curve:
1. Use a disposable pipet to fill a plastic cuvette approximately half full with each of the
A-E calibration standard solutions. Make sure to wipe any excess liquid off the
exterior of the cuvette with a paper towel or kimwipe.
2. Insert the “A” cuvette into the sample holder on the spectrometer. Press the
“Collect/Stop” button with your finger. It looks like a “PLAY” button on a DVD
player. Once the spectrum is acquired, press the “Collect/Stop” button again. Be
Patient! It takes a moment for the process to stop. Do not press the button twice. The
graph will autozero and the absorbance and the wavelength of maximum absorbance
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will be displayed. Record the max absorbance wavelength and the absorbance at the
max absorbance wavelength on the data sheet. Use the left/right arrows to move the
data circle to 536nm and record the absorbance at this wavelength.
3. Remove the cuvette containing the “A” standard from the spectrometer and insert the
cuvette containing the “B” standard. Press the “Collect/Stop” button. You will be
asked whether you wish to store or discard the previous sample data. Use the stylus
to select “Discard.” Once the spectrum is acquired, press the “Collect/Stop” button
again. The graph will autozero and the wavelength of maximum absorbance will be
displayed. Record the max absorbance wavelength and the absorbance at the max
absorbance wavelength on the data sheet. Use the left/right arrows to move the data
circle to 536nm and record the absorbance at this wavelength.
4. Repeat step 3 for the remaining calibration standards.
5. Make a graph of absorbance versus concentration. Your teacher may have you
prepare two graphs, one using the absorbance at the wavelength of max absorbance
and another using the absorbance at 536nm.
Sample Preparation: Select one of the three pre-digested aspirin samples and pipet
2.0mL of the solution to a clean 50mL volumetric flask and dilute to the mark with the
acidified iron (III) solution. Mix thoroughly.
Sample Analysis:
1. Half fill a cuvette with the aspirin sample (in acidified iron (III) solution) prepared
earlier. Use a paper towel or kimwipe to wipe any excess liquid from the outside of
the cuvette. Insert the cuvette into the spectrometer and read and record the data
exactly as you did with the A-E calibration standard solutions.
2. Use the graph to find the concentration of aspirin in the diluted sample. Record this
value on the data sheet.
3. Calculate the mg of aspirin in the tablet and compare this value to the value given on
the manufacturer’s label.
Brand Name ___________________ Aspirin Content, mg (from label)
_______________mg
Table 1: Data for Calibration Curve
Solution Concentration
(mg/mL)
A 0.080
B 0.064
C 0.048
D 0.032
E 0.016
Calculation of mg aspirin in the commercial product:
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________mg/mL X 50mL / 2.0mL X 250mL =
____________mg
(from graph) (dilution factor) (sample volume)
Calculate % difference between experimental value and value from manufacturer’s
label (% error)
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Thin-Layer Separation of Lipstick
Write up in your lab notebook.
Discussion:
In thin layer chromatography (TLC), the stationary phase is a thin layer of adsorbent
particles attached to a solid plate. The ones we use have silica gel attached to plastic
sheets. When handling the sheets it is best not to touch the silica gel surface. A small
amount of sample is applied (spotted) near the bottom of the plate and the plate is placed
in the mobile phase (solvent). This solvent is drawn up by capillary action. Separation
occurs as each component, being different in chemical and physical composition,
interacts with the stationary and mobile phases to a different degree creating the
individual bands on the plate. The retention factor. Rf value, is used to characterize and
compare components of various samples. This value is calculated in the following way:
Rf value = distance from origin to component spot
distance from origin to solvent front.
In this activity, you will determine how many pigments are contained in a lipstick sample
as well as whether or not different lipstick samples contain any of the same pigments.
You will also look at the resulting chromatogram under UV light in order to identify any
fluorescent components.
SDS
Solvent A
65 mL toluene: Highly flammable liquid and vapor. Keep away from heat, sparks,
open flames, and hot surfaces. Harmful if swallowed. May be fatal if swallowed and
enters airways. Causes skin and eye irritation. May cause drowsiness or dizziness.
Avoid breathing mist, vapors or spray
30 mL methanol Highly flammable liquid and vapor. Keep away from
heat, sparks, open flames, and hot surfaces. Toxic if swallowed or in contact with skin
+H311). Do not eat, drink or smoke when using this product (P270).
Hazard class: Skin and serious eye damage, corrosion or irritation (Category 2, 2A).
Causes skin and serious
eye irritation (H315+H319).
Hazard class: Specific target organ toxicity, single exposure (Category 1). Causes
damage to organs (H370). Do
not breathe mist, vapors or spray
5 mL glacial acetic
Solvent B
92 mL butanol
8 mL 2 M HCl
Procedure:
1. Touching the TLC strip as little as possible and only at the edge, obtain two (2) TLC
strips (4 x 10 cm each) and 2 or 3 lipstick samples that you would like to test.
2. Use a pencil to lightly draw a line approximately cm from the bottom edge of the TLC
strip. Mark the line with an “x” for each lipstick sample you will test and identify the
different samples in some way at the top of the TLC strip.
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3. Use a toothpick to make a small dot in the center of each “x” with the proper lipstick
sample. A small sample size (about cm in diameter and not too thick) will result in
better separation of pigments.
4. Repeat steps 2 and 3 using the same lipstick samples so that you have two identical
TLC strips to be tested.
5. Put a very small amount of solvent A into one 100 mL beaker and a very small amount
of Solvent B in another 100 mL beaker. The solvent level must be below the point of
application!!! Mark these beakers in some way so that you know which solvent is
which.
6. Handling the TLC strips by the edges only, place one strip in each solvent and cover
each with a 250 mL beaker. Be sure that the sample spots are at the bottom of the TLC
strip when it is place in the beaker, but also be sure that the sample spots are not below
the solvent level in the beaker.
7. Check the progress of the solvent as it wets the TLC strip and moves toward the top.
Do not let the solvent reach the top!
8, When the solvent is close to the top of the TLC strip, remove the chromatogram and
place on a paper towel to dry.
9. Immediately mark the solvent front and the center of each pigment spot with a pencil
line.
10. Make a dot with your pencil in the center of each spot. If the spot is irreuular or
bandshaped, use your judgement as to where the approximate center might be if it
were a round spot.
11. Measure the distance from the point of origin to the solvent front and record in the
Data Table. Be sure that collected data is recorded in the appropriate Data Table based
on the solvent used (A or B) as well as the lipstick tested.
12. Calculate the R value for each spot by dividing the spot distance (origin to spot) by
the solvent distance (origin to solvent front). Record properly in your Data Table.
13. Identify any pigments that appear to be the same in the different lipstick samples.
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Determination of Melting Points
Write up in your lab notebook.
Discussion: The melting point of a compound is the temperature at which it changes
from a solid to a liquid. This is a physical property often used to identify compounds or to
check the purity of the compound. It is difficult, though, to find a melting point. Usually,
chemists can only obtain a melting range of 2-3°C accuracy. This is usually sufficient for
most uses of the melting point. The purpose of this experiment is to become familiar
with equipment used to find the melting points of organic compounds and to identify
unknown compounds based on their melting point.
Procedure:
1. Obtain a capillary tube and one of the labeled organic compounds. Each person in
the group should do a different sample.
2. Push the open end of the capillary tube into the solid until the solid has moved about
0.5 cm into the capillary tube. Invert the tube and tap the closed end on the table to
pack the solid into the closed end of the tube.
3. Place the capillary tube into the Mel-Temp chamber as instructed by your teacher.
Three capillary tubes will fit into this chamber, so up to three members of the group
can run their samples at the same time.
4. Turn on the Mel-Temp and ensure the thermometer is inserted properly. Start with a
setting of one to two. Observation of the sample chamber and the thermometer
should be made frequently in order not to miss the melting point of the sample. Heat
slowly, so that accurate results are achieved. Use small incremental increases in the
setting when needed.
5. Watch the temperature on the thermometer and solid inside the capillary tube. As the
melting range of the solid approaches, you will see the solid begin to sweat. This is
the low end of the melting range. The temperature at which the solid has completely
turned to liquid is the high end of the melting range. Record these values.
6. Turn off the Mel-Temp to allow it to cool before trying your unknowns. It should
take about 15 minutes for the temperature to cool down.
7. Once you are comfortable with the procedure for determining the melting point of a
solid, prepare a sample of your unknown in the same way that you prepared the
known solids and find its melting range.
8. From the melting range you found, identify your unknown solid using the data from
the knowns you tested initially.
Questions:
1. Define the "melting point" of a substance.
2. What is the purpose of determining melting points?
3. Why is this method not used for finding the melting points of inorganic compounds?
4. What would you expect the melting point to be if you mixed two substances and took
the melting point of the mixture?
5. How could the rate of heating influence the melting point?
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Molar Volume of a Gas
Data, Calculations and Questions only
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.
SDS
Mg Substance not considered hazardous. However, not all health aspects of this
substance have been thoroughly investigated.
6M HCl 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 room-temperature 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.
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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.
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?
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Airbags
Discussion;
Airbags were mandated as a piece of safety equipment for all vehicles in 1998. All cars
and trucks are required to have airbags on both driver and passenger sides. Many vehicles
today include side passenger airbags. To date, the National Highway Traffic Safety
Administration, NHSA, calculates that using a seat belt and having an airbag reduces the
risk of death by 61 percent.
Automobile airbags are simply gas-inflated cushions. The gas that fills the bag is the
result of a series of rapid chemical reactions triggered by an impact. In this activity, you
will simulate the process of generating the gas using a different but similar chemical
reaction inside a resealable plastic bag.
This experiment connects stoichiometry and gas laws. Your mission is to generate
enough gas to just fill a small, resealable plastic bag using baking soda and 6.0 M
hydrochloric acid, HCl. The ideal “airbag” will be filled to plumpness yet not
overinflated or underinflated. The bag may contain unreacted chemicals and other
products of the reaction.
You will be asked to describe the method you develop to solve the problem. You must
complete this assignment (including the report) during the assigned period. You may not
share information between groups.
You and your partner will be graded on the reproducibility of the experimental design,
data collection skills, and the accuracy of your results. Your written responses must be as
clear and as concise as possible in communicating your thought processes. You must
follow proper safety procedures. PROCEDURE
1. Write a balanced chemical equation for the reaction that will produce the gas in your
simulated airbag before beginning your procedure.
2. With your partner, design a plan for inflating your airbag and record your plan on your
student answer page. Have your teacher initial your work before beginning.
3. Carry out your plan, and record your observations.
4. Revise your design. Repeat, and show your teacher the filled bag.
QUESTIONS
1. Using evidence from your experiment, go back to your balanced chemical equation
and include energy and state symbols.
2. What was the limiting reagent in your experiment? Justify your answer with data.
3. Sketch the system of the inflated plastic bag after the reaction has finished and label all
of the species present. Be sure to include the state symbol along with charges for any
ions.
4. The reaction in this activity is only a simulation for that in a real automobile airbag.
Write the series of balanced chemical equations for the reaction that actually takes
place in an automobile airbag. You may need to do some research, and many online
resources are available.
5. The volume for the driver airbag has remained an average of 56 L over the years.
Consider only the first step in the reaction written in Question 4 and calculate the mass
of the reactant needed to produce this volume of gas. Assume standard temperature
and pressure.
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“Wet” Dry Ice
Do Not Write up in your lab notebook.
Discussion How can ice be “dry”? Dry ice is the common name for compressed carbon dioxide, CO2.
The term “dry” refers to the fact that no liquid is left behind when a sample of dry ice is left
out at room temperature. The solid carbon dioxide undergoes a phase change directly from a
solid to a gas; it sublimes.
Refer to the phase diagrams of water (Figure 1) and carbon dioxide (Figure 2). Both graphs
contain a point at a specific temperature and pressure where all three phases exist in
equilibrium. This coordinate on the graph is referred to as the triple point of the substance.
You will also notice that both graphs indicate a critical point. At temperatures above the
critical temperature, it is not possible to condense the gas into liquid state even with an
increase in pressure. 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.
Safety Precautions:
Do not touch the dry ice with your hands. Dry ice is very cold(-78 °C) and can cause
tissue damage to your skin. Wear goggles at all times.
Prelab instructions: Answer these questions on a separate sheet of paper.
1. Draw Lewis structures for water and for carbon dioxide.
2. Refer to Figure 1 and Figure 2 in the discussion.
a. Compare the scales of the y-axes for the two phase diagrams. Give the
approximate range of pressures for each diagram, including units.
b. Consider samples at standard pressure, 1 atm, and room temperature, 20°C.
Determine the state of matter in which each sample of water and carbon dioxide
would exist.
c. State a set of conditions under which both water and carbon dioxide would be
solid.
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d. Compare the solid-liquid equilibrium lines for water and carbon dioxide. What
does the slope of this line indicate about the density of the solid?
3. Find the triple point pressure of CO2 online and fill in the ? in figure 2 above.
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.
Part 2:
1. Construct a micro-pressure gauge by cutting the stem from a thin-stemmed pipette.
2. Use the hot glue gun to place a small amount of glue to seal the end. Tie a piece of
thread around this end, as shown in Figure 5.
3. Beginning with the sealed end, mark the stem of the pipette in 1.0 cm or otherwise
equal increments.
4. Prepare another large pipette as you did in Part 1 by cutting off the tip.
5. Check the length of the micro-pressure gauge by placing it inside the large pipette.
Cut the pressure gauge so that its length just comes to the top of the bulb.
6. Place a drop of colored water into the pressure gauge by squeezing the center of the
stem with pliers and placing the open end into a cup of colored water. Gently release
the grip to allow the water to be drawn into the pressure gauge. Based on the position
of your colored water droplet, count the number of increments occupied by the
trapped air sample. Record this value in Table 2 on your student answer page.
7. Obtain a small sample of dry ice and fill the bulb about one quarter full. Slide the
pressure gauge into the large pipette, fold the tip, and firmly grip both with the pliers. Be
sure to place the setup into the water as shown in Figure 6 and observe.
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8. As the pressure climbs inside the pipette, you will notice the colored water plug rising in
the micro-pressure gauge. Note the position of the colored water droplet when the dry ice
begins to liquefy.
9. Gently release the grip and repeat until there is no more dry ice in the bulb.
10. Bring your data from your pipette back into class.
Clean up:
1. Clean up any water, dry the plastic cup and leave it at your station. Throw away all
pipettes you used but save your pressure gauge. Do NOT throw away the plastic cups
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Properties of Solutions: Electrolytes and Non-Electrolytes
Write up in your lab notebook.
Discussion:
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. 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 ionize depending on their strength.
Example: Strong electrolyte H2SO4 →H+(aq) + HSO4-(aq) (100% dissociation)
Example: Weak electrolyte HF ↔ H+(aq) + F-(aq) (<100% dissociation)
SDS
0.05 M NaCl, 0.05 M CaCl2, 0.05 M
AlCl3, 0.05 M HC2H3O2, 0.05 M H3BO3
Substance not considered hazardous. However,
not all health aspects of this substance have
been thoroughly investigated.
0.05 M H3PO4 Slightly toxic by ingestion and inhalation; body
tissue irritant.
0.05 M HCl Slightly toxic by inhalation and ingestion.
Severe body tissue irritant. Corrosive to eyes.
Avoid all body contact.
0.05 M CH3OH (methanol) Toxic by ingestion (may cause blindness),
inhalation or absorption.
0.05 M C2H6O2 (ethylene glycol) Mildly toxic by ingestion and inhalation. Skin,
eye, and mucous membrane irritant. Avoid all
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body tissue contact.
Procedure
1. Open the Conductivity lab method on the lab quest. The window will display live
conductivity readings in units of microsiemens (S).
2. The Conductivity Probe is already attached to the labquest. It should be set on the 0-
20,000 S position.
3. Obtain the Group A solution containers. The solutions are: 0.05 M NaCl, 0.05 M
CaCl2, and 0.05 M AlCl3.
4. 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.
5. 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.
6. 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 4.
Cleanup:
1. Rinse off the conductivity probe with distilled water, wipe up any spills and thrown
any trash away.
Questions:
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 strong; for weak.
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?
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8. Did aqueous methanol, CH3OH, have the same conductivity value as aqueous
ethylene glycol, C2H6O2? Explain.
39 Drop pH Lab
Data, Calculations and Questions only
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 Dropper bottles of
ammonia solution X
bleach solution Y
baking soda, muratic acid, sprite, club soda,
sugar, tap water, distilled water, lemon juice,
milk, saliva, egg white, soap, shampoo,
Universal Indicator
wooden splints, conductivity sensors
SDS:
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. Place the well plate on a piece of clean white paper.
2. 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!
3. Add about 5 drops of each household solution to separate wells in row C & D.
4. Using a conductivity sensor, test each household product and record the conductivity
as Very High, High, Medium, Low, or None. Rinse off the conductivity sensor
probe wires with a wash bottle between each product.
5. Add one drop of Universal Indicator to each well containing a solution. Stir with a
wooden splint.
6. Determine the pH values of the household solutions by comparison with the other
wells.
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Well #
pH (±0.1)
drops X
drops Y
A1
empty
A2
2.0
39
0
A3
3.0
35
4
A4
4.0
31
8
A5
5.0
27
12
A6
6.0
24
15
Well #
pH (±0.1)
drops X
drops Y
B1
7.0
20
19
B2
8.0
17
22
B3
9.0
14
25
B4
10.0
11
28
B5
11.0
9
30
B6
12.0
3
36
Well #
pH (±0.1)
color
substance
conductivity
C1
____
_______
C2
____
________
C3
____
________
C4
____
________
C5
____
________
C6
____
________
Well #
pH (±0.1)
color
substance
conductivity
D1
____
________
D2
____
________
D3
____
________
D4
____
________
D5
____
________
D6
____
________
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Titration Lab
Data, Calculations and Questions only
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.
Titrations are calculated using dimensional analysis: For example, 13.20 mL of 0.100 M
NaOH is used to titrate 11.19 mL of Sulfuric acid. The molarity of the Sulfuric acid is
unknown.
First, write a balanced equation:
2NaOH + H2SO4 Na2SO4 + 2H2O 0.01320 mL NaOH x 0.100 moles NaOH x 1 mole H2SO4 x ? mol__
1 L 2 moles NaOH 0.01119 L H2SO4
This give an answer in moles of H2SO4 per L H2SO4 which is molarity.
SDS
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. The acid will be hydrochloric acid and the base will be 0.100 M NaOH. Your
indicator will be phenolphthalein.
3. Get two 150 mL beakers, label one base with a sharpie, label the other acid. Go to
the fume hood and put 150 mL of base in the base beaker and 50 mL of acid in the
acid beaker. Pour the acid into one buret and the base into the other. Drain both
burets a little to move the liquid into the tip of the buret.
4. Get a 125 mL Erlenmeyer flask. You will use this to do the titration. Rinse it out
with water twice and distilled water once.
5. Using the buret get 10.0 mL of 0.10 M HCl and titrate it using base to practice your
titrating skills. If you add too much base, back titrate using a little more acid until the
solution goes colorless. Your end point should be very close to 10.00 mL of base.
6. 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 around 10.0 mL. Get a 50 mL
beaker, rinse it with distilled water and bring to the teacher for your unknown sample.
7. 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. Be sure to show
your calculation.
8. Your teacher may have you repeat this titration for more than one unknown acid.
Cleanup:
1. Pour the remainder of the base down the sink. Pour the remainder of the acid down
the sink. Pour the contents of the Erlenmeyer down the sink.
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2. Pour 50 mL of 0.1M HCl into the now empty base buret to neutralize the base. Rinse
both burets 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.
3. Wash all beakers and the Erlenmeyer flask with soap and rinse and put on a paper
towel at your station to dry.
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Qualitative Analysis of the Group I Cations
Do Not Write up in your lab notebook.
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 precipi-
tated from any other cations that might be present in a sample by addition of 6 M hydro-
chloric 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. 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.
SDS
6M HCl Toxic by inhalation and ingestion. Severe corrosive to all body tissues,
especially skin and eyes. Avoid all body contact.
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K2Cr2O7
solution
Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact.
6M NH4OH 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.
6M HNO3 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.
AgNO3 Moderately toxic by ingestion. Irritating to body tissues. Avoid all body
tissue contact. Silver compounds photoexpose and will stain skin and
clothing.
HgNO3 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
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.
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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 known 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 hot-
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
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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
Pb2+, Ag+, Hg22+
HCl(aq)
cold water wash
AgCl, Hg2Cl2
HgNH2Cl Ag(NH3)2+
White and
Hg black HNO3(aq)
AgCl
white
Pb2+
Decantate 2 Residue 2
K2Cr2O7
PbCr2O7
yellow
Group I Cation Analysis
AgCl, PbCl2, Hg2Cl2
hot water wash
NH4OH(aq)
Decantate 3
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Copper into Gold: The Alchemist’s Dream
Do Not Write up in your lab notebook.
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− → ZnO2
2- + H2
When the copper token is added to the solution, an electrochemical couple formed by the
copper-zinc 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.
SDS:
6M NaOH Moderately toxic by ingestion and skin absorption. Corrosive to body
tissues. Causes severe eye burns. Avoid all body tissue contact.
Zinc dust 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?
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2. Why is it necessary to clean the penny before starting the reaction?
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Freezing Point Depression of a Solution (Ice Cream)
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 (123 mL) cup milk
½ (123 mL) cup cream (you can skip the cream but double the milk)
2 TBS (30 mL) sugar
¼ tsp (3.75 mL) vanilla
½ bag ice
½ cup (123 mL) salt
Two freezer sized zip top baggies- one quart and one gallon.. (No slider zip bags) (two
bags per person
Spoons and/or cups to eat it.
Place the first 4 ingredients in the quart bag bag, seal tightly being sure to get as much air
out as possible. Place the ice and salt in the gallon bag. Place the bag with the milk
mixture into the bag with the salt and ice and seal the gallon 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
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Periodic Table
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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 SO32-
SO3 SO42-
CO2 CO32-
N2O3 NO2-
N2O5 NO3-
P2O3 PO33-
P2O5 PO43-
H2O OH-
NH3 NH4+
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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.
3. All bromides, chlorides, and iodides are soluble except lead, mercury 1 and silver.
4. All hydroxides are insoluble except group 1, barium and strontium.
5. The borates, carbonates, chromates, oxides, phosphates, silicates, and sulfites are
insoluble except those of ammonium, and group 1.
6. All sulfides are insoluble except Group1, Group 2 and ammonium
7. All sulfates are soluble except those of Group 2, lead, mercury(I) and silver. 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.
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
Least active
F
Cl
Br
I
Most active
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Polyatomic Ions
Ammonium NH4+
Acetate C2H3O2- or CH3COO-
Arsenate AsO43-
Borate BO33-
Bromate BrO3-
Carbonate CO32-
Chlorate ClO3-
Chromate CrO42-
Cyanide CN-
Dichromate Cr2O72-
Dihydrogen phosphate H2PO4-
Hydrogen carbonate (bicarbonate) HCO3-
Hydrogen phosphate HPO42-
Hydrogen sulfate (bisulfate) HSO4-
Hydroxide OH-
Iodate IO3-
Nitrate NO3-
Oxalate C2O4-2
Permanganate MnO4-
Peroxide O22-
Phosphate PO43-
Silicate SiO32-
Sulfate SO42-
Thiocyanate SCN-
Thiosulfate S2O32-
Rules to Derive other PAIs
-ate to –ite one less oxygen, same charge
-ite to hypo-ite one less oxygen same charge
-ate to per-ate one more oxygen same charge
For example
ClO4- perchlorate
ClO3- chlorate
ClO2- chlorite
ClO- hypochlorite
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Apples
UNITS
Parsec = 1.91738 X 1013 miles
furlong=1/8 mile-exactly
Rod = Pearch -exactly
Rod = 16.5 links -exactly
Football field = 100 yards -exactly
Soccer field = 100 meters -exactly
Rod=5 ½ yards-exactly
Fathom = 6 feet -exactly
Yard = 3 feet -exactly
foot = 12 inches -exactly
inch = 2.54 centimeters
Walking pace (avg) = 22.00 inches
Story on a building = 3.333 meters
Light year = 9.467 X 1015 meters
Barn = 10-24 meters -exactly
League = 3 miles -exactly
1 mile = 5280 feet -exactly
1 kilometer (km) = 1000 meters -exactly
1 cubit = 25.00 inches
1 cubit = 2.08333 ft
1 cubit2 = 4.34027777 ft2
1 cubit3 = 9.042245 ft3
TIME UNITS millennium = 1,000 years -exactly
century = 100 years -exactly
decade = 10 years -exactly
year = 365.25 days
day = 24 hours -exactly
hour = 60 minutes -exactly
minute = 60 seconds -exactly
blink of an eye = 0.1 second
fortnight = 14 days -exactly
score = 20 years -exactly
WEIGHTS & METRIC MASSES
Pound = 16 ounces -exactly
Ton = 2000 pounds -exactly
Tonne = 1000 kilograms (metric ton) -
exactly
Gram = 1000 milligrams -exactly
Kilogram = 1000 grams -exactly
Kilogram = 2.20462280 pounds
1 pound = 454 grams
Poundal = 14.09808 grams
Dram = 1.771 845 195 3125 grams.
Grain = 65 mg
1 carat = 0.200 g for diamonds
VOLUME MEASUREMENTS
1 liter = 1000 milliliters -exactly
2 liters = 64.75 ounces
1 gallon = 128 ounces –exactly
1 gallon = 3.76 L
1 milliliter = 1 cm3 -exactly
1 milliliter = 20 drops
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Avocado
Do the following on your own sheet of paper. Do not write on this sheet of paper.
Write down the problem/question and then write down your answer. If it is a math
problem, circle your final answer.
1. For each of the following, how many sig figs are present:
a. 1.0098 b. 16000 c. 0.007054
d. 54000.010 e. 1000 f. 0.0680250
2. Express of the following problems using the correct number of sig figs.
a. 6.54 + 7.329 b. 7.98 X 6.5423 c. 1.45-0.078
d. 6.02 X 10.000 e. 5.23 / 2.4 f. 7.89 X 107 x 6.81 X 109
3. For each of the following, solve the math problem given and write the answer in the
correct units.
a. 2.34 cm+ 6.58cm +1.23cm +32.4cm +1.05cm
b. 18.357 ml – 1.34 ml
c. 16.57cm x 1.4567cm x 2.0030 cm
d. 6547.008 g2 / 19.7 g
e. 12.45 cm + 1 cm + 19.713 cm + 67.54 cm
f. 1.23 m x 6.54637 m x 54.321 m x 1.6 s
g. 1.23 x 1012 g + 6.21 x 1011 g + 7.4 x 1010g
h. 1.3 x 107 cm x 9.325 x 102 cm
i. 1.44 x 108 joul2 / 1.270 x 104joul
4. Write each of the following in scientific notation:
a. 657894 b. 45.32 c. 87901
d. 85 e. 2 f. 1000000
g. 0.00345 h. 0.064 i. 10.2
5. Convert 1.78 x 109 meters /second to km/hour.
6. Convert 87 cm to km
7. Convert 0.007812 dm to m.
8. If you rode on a beam of light for two days, how far in km would you travel? Speed
of light is 3.00 X 108 m/s.
9. If one kilogram = 2.2 lbs, convert your weight to milligrams.
10. Convert 16.7 inches to feet
11. Convert 25 yards to feet (there are 3 feet in a yard)
12. Convert 90 centuries to months
13. Convert 84 miles to kilometers (there are 1.609 kilometers in a mile)
14. Convert 4.75 centimeters to inches (there are 2.54 centimeters in an inch)
15. Convert 48,987 minutes to days
16. Convert 27 months to fortnights (there are 14 days in a fortnight and ~30 days in a
month)
17. Convert 0.09 miles to inches (there are 36 inches in a yard and 1760 yards in a mile)
18. Convert 4.66 centimeters to miles (2.54 centimeters in an inch, 36 inches in a yard,
1760 yards in a mile)
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Tomato
Reasons for not using Dimensional Analysis
1. Let's say you're super-intelligent and enjoy solving relatively simple problems in the
most complex manner.
2. Let's say you're tired of always getting the correct answers.
3. Let's say you're an arty type and you can't be confined by the structure of DA. You
like messy solutions scribbled all over the page in every which direction. It's not that
you want to make a mistake. But you really don't care that much about the answer.
You just like the abstract design created by the free-wheeling solution... and the
freedom from being confined by structure.
4. Let's say that you have no interest in going to the prom or making the soccer team,
and you don't mind being unpopular, unattractive, ignorant, insecure, uninformed,
and unpleasant.
Otherwise, You Need Dimensional Analysis!
Personal Testimony:
I was at home, sick with the flu when Mr. Elegante taught my class about Dimensional
Analysis. Despite opportunities given to me to make up the assignments that I had
missed, I chose to not do them. I thought that my mathematical abilities were already
sufficient. How wrong I was! It’s been five years since I took that class--Now I spend
my afternoons panhandling at traffic lights, hoping for passersby to give me spare
change. If I ‘m lucky enough to scam a buck after a day’s work, I’m still not sure if my
hourly rate makes cents.
--Mario
All answers must be in the correct unit with correct significant figures. Show all of your
work and circle your final answer. Do all work in your notebook.
1. A box is 19.3 inches by 0.86 ft x 1.289 dm. Calculate the volume of the box in
cm3.
2. What is the volume, in cm3, of a tank that can hold 18 754 Kg of methanol whose
density is 0.788g/cm3.
3. CaCl2 is used as a de-icer on roads in the winter. It has a density of 2.50 g/ml.
What is the mass of 15.0 L this substance?
4. Convert the speed of light in meters per second to miles per hour. Given: 1 mile
= 1.609 km. speed of light is 3.00 x 108 m/s
5. On the strange planet of Bookoo they have a strange monetary system. The ugle
is a type of currency used on one of the islands. A person wants to buy a zugu for
a special someone. Zugus are found only on 1 remote island. The person must
travel from island to island exchanging money to buy the zugu. You must figure
out the exchange rate given the following information: 3 ugles = 7 kapus, 17
kapus = 5 linpeps, 9 fincors = 23 kapus, 15 rinpers = 8 bloors and 2 fincors = 11
rinpers. Find the final price in ugles if the price on the island is 358 rinpers.
6. If you travel 169.246 km in 15.2 hours, what is your speed in km/hour?
7. Calculate the mass of a person who weighs 96.55 lbs in kg. (See notes for kg to
lbs conversion)
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8. Nonsense words taken from the poem Jabberwocky (from Lewis Carroll’s
Through the Looking Glass)
There are 20 tumtum trees in the tulgey wood.
In each tulgey wood is one frumious Bandersnatch.
There are 5 slithy toves in 2 borogoves.
There are 2 mome raths per Jabberwock.
There are 2 Jubjub birds in 200 tumtum trees.
There are 200 mome raths in each borogove.
There are 5 Jubjub birds per slithy tove.
The question is: If there are 5 frumious Bandersnatches, how many Jabberwocks are
there?
9. Because you never learned dimensional analysis, you have been working at a fast
food restaurant for the past 35 years wrapping hamburgers. Each hour you wrap
184 hamburgers. You work 8 hours per day. You work 5 days a week. You get
paid every 2 weeks with a salary of $840.34. How many hamburgers will you
have to wrap to make your first one million dollars? If you can solve the
problem, you will have learned dimensional analysis and you can get a better job.
But, since you won't be working there any longer, your solution will be wrong. If
you can't solve the problem, you can continue working which means the problem
is solvable, but you can't solve it. We have decided to overlook this impasse and
allow you to solve the problem as if you had continued to wrap hamburgers.
Jalapeno Metric Conversion
1. 40 ml = _______ L 2. 5000 L = _______ kl 3. 8 g = _______ kg
4. 12000 L = _______ kl 5. 50 mg = _______ g 6. 6000 m = _______ km
7. 200 kg = _______ g 8. 10000 g = _______ kg 9. 500 ml = _______ L
10. 1 L = _______ ml 11. 4000 L = _______ kl 12. 400 cm = _______ Mm
13. 20 ml = _______ kl 14. 7000 ml = _______ L 15. 7 cm = _______ mm
16. 9000 L = _______ ml 17. 6 m = _______ m 18. 1000 cm = _______ m
19. 11 km = _______ m 20. 80 mg = _______ kg 21. 3 m = _______ mm
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Banana
1. A diamond is a pure sample of the element carbon. The tabulated density of diamond
is 3.51 g/cm3. Jewelers use a unit called a carat to describe the mass of diamond. What is
the volume of the stone in a 2.00 carat diamond engagement ring?
2. A farmer owns a rectangular field that fronts along a road. State highway engineers
surveyed the frontage and found it measures 138.3 meters in length. The farmer, who
wants to build a fence around the field, paces off the field’s width and estimates it to be
52 meters. How many meters of fence will the farmer have to build? What mass of
fertilizer will the farmer need to fertilize the field with 0.0050 kg of fertilizer per square
meter of field?
3. The diameter of metal wire is given by its wire gauge number. For example, 16-gauge
wire has a diameter of 0.0508 inch. Calculate the length in meters of a 5.00 pound spool
of 16-gauge copper wire. The density of copper is 8.92 g/cm3.
4. The solder used by plumbers to fasten copper pipes consists of 67% lead and 33% tin.
If you have a 1.0 pound (454 g) block of solder, how many grams of lead do you have?
Grams of tin?
5. You are fortunate to own a ring made of white gold, which actually consists of 60.%
gold and 40.% platinum. If your ring has a mass of 6.89 g, how many grams of platinum
do you have? How many grams of gold?
6. At 25 oC the density of water is 0.997 g/cm3, whereas the density of ice is 0.917
g/cm3. (a) If a soft drink can (volume = 250. mL) is filled completely with pure water
and then frozen at –10 oC, what volume will the solid occupy? (b) Could the ice be
contained within the can?
7. Amethyst is a colored form of the mineral quartz in which the color purple comes from
traces of the element Manganese. (a) To determine the density of amethyst, you take a
stone having a mass of 15.25 g and place it in a 100. mL graduate cylinder containing
45.0 mL of water. On adding the stone, the water surface rises to the 50.8 mL mark.
What is the density of amethyst? (b) If you have a piece of amethyst that is 1.8cm x 0.95
cm, how thick does it have to be to have a mass of 6.50 g?
8. Battery plates in lead storage batteries (the type used in automobiles) are made from a
mixture of two chemical elements: lead (94.0%) and antimony (6.0%). If you have a
battery plate with a mass of 25.0 g, how many grams of lead and how many grams of
antimony are present?
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Firehouse
Use the combined gas law to solve the following problems:
1) If I initially have a gas at a pressure of 12 atm, a volume of 23 liters, and a temperature of 200. K, and
then I raise the pressure to 14 atm and increase the temperature to 300. K, what is the new volume of the
gas?
2) A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and a temperature of 299 K. If I raise the
temperature to 350 K and lower the pressure to 153.2 kPa , what is the new volume of the gas in cm3?
3) A gas that has a volume of 28 liters, a temperature of 45 0C, and an unknown pressure has its volume
increased to 34 liters and its temperature decreased to 35 0C. If I measure the pressure after the change
to be 2.0 atm, what was the original pressure of the gas?
4) A gas has a temperature of 14 0C, and a volume of 4.5 liters. If the temperature is raised to 29 0C and
the pressure is not changed, what is the new volume of the gas?
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5) If I have 17 liters of gas at a temperature of 67 0C and a pressure of 88.89 atm, what will be the pressure
of the gas if I raise the temperature to 94 0C and decrease the volume by 5 liters?
6) I have an unknown volume of gas at a pressure of 7.15 psi and a temperature of 325 K. If I raise the
pressure to 1.2 atm, decrease the temperature to 320. K, and measure the final volume to be 48 liters,
what was the initial volume of the gas?
7) If I have 21 liters of gas held at a pressure of 78 atm and a temperature of 637 0C, what will be the
volume of the gas if I increase the pressure to 45 atm and decrease the temperature to 751 K?
8) If I have 2.9 L of gas at a pressure of 5 atm and a temperature of 50 0C, what will be the temperature of
the gas if I decrease the volume of the gas to 2.4 L and decrease the pressure to 2280. torr?
9) I have an unknown volume of gas held at a temperature of 115 K in a container with a pressure of 607.8
kPa. If increasing the temperature to 225 K and increasing the pressure to 30. atm causes the volume of the gas
to be 29 liters, how many liters of gas did I start with?
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KFC
Use your knowledge of the ideal and combined gas laws to solve the following problems. Hint: Figuring out
which equation you need to use is the hard part!
1) If four moles of a gas at a pressure of 5.4 atmospheres have a volume of 120 liters, what is the
temperature?
2) If I initially have a gas with a pressure of 84 kPa and a temperature of 350 C and I heat it an additional
230 degrees, what will the new pressure be? Assume the volume of the container is constant.
3) My car has an internal volume of 2600 liters. If the sun heats my car from a temperature of 200 C to a
temperature of 550 C, what will the pressure inside my car be? Assume the pressure was initially 760
mm Hg.
4) How many moles of gas are in my car in problem #3?
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5) A toy balloon filled with air has an internal pressure of 1.25 atm and a volume of 2.50 L. If I take the
balloon to the bottom of the ocean where the pressure is 95 atmospheres, what will the new volume of
the balloon be? How many moles of gas does the balloon hold? (Assume T = 285 K)
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Marcos
Ideal Gas Law Problems
Use the ideal gas law to solve the following problems: Do all work on this sheet. No work, no credit.
1) If I have 4.02 moles of a gas at a pressure of 576 kPa and a volume of 12.4 liters, what is the
temperature?
2) If I have an unknown quantity of gas at a pressure of 984 torr, a volume of 33.1 liters, and a temperature
of 87 0C, how many moles of gas do I have?
3) If I contain 3.00 moles of gas in a container with a volume of 60. liters and at a temperature of 400. K,
what is the pressure inside the container?
4) If I have 7.7 moles of gas at a pressure of 0.090 atm and at a temperature of 56 0C, what is the volume of
the container that the gas is in?
5) If I have 89.23 g of iodine gas at a temperature of 67 0C, and a volume of 88.89 liters, what is the
pressure of the gas?
6) If I have an unknown quantity of gas at a pressure of 0.523 atm, a volume of 25 liters, and a temperature
of 300. K, how many moles of gas do I have?
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7) If I have 21.79 g of carbon monoxide gas held at a pressure of 78 atm and a temperature of 900. K, what
is the volume of the gas?
8) If I have 1.9 moles of gas held at a pressure of 594.6 kPa and in a container with a volume of 500.
milliliters, what is the temperature of the gas?
9) If I have 2.46 moles of gas held at a temperature of 97 0C and in a container with a volume of 4500 cm3,
what is the pressure of the gas?
10) If I have an unknown quantity of gas held at a temperature of 1195 K in a container with a volume of
25.1 liters and a pressure of 560. psi, how many moles of gas do I have?
11) What is the density of carbon tetrachloride gas if the pressure is 11.74 psi and the temperature is 32 ºF?
12) If I have 72 liters (15.79 g) of gas held at a pressure of 3.4 atm and a temperature of 225 K, what is the
molar mass of the gas?
13) What pressure would a mixture of 3.20g of O2, 6.40 g of CH4, and 6.40 g of SO2 exert if the gases were
placed in a 40.0 L container at 127 °C?
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Outback Directions: For each of the following problems, show all work including units, circle your final answer and
maintain appropriate significant figures.
1. A tin can with a volume of 20.0 ml contains a gas at 1.00 atm and 27.0 °C. If 800.00 calories of heat is
applied to raise the temperature to 67.0 °C and the volume is unchanged, what is the new pressure inside the
can?
2. A sample of gas occupies 10.0 L at 240. °C under a pressure of 80.0 kPa. At what temperature will the gas
occupy 20.0 L if the pressure is increased to 107 kPa?
3. What is the volume of a gas balloon filled with 4.00 moles of helium when the atmospheric pressure is 748
mm Hg and the temperature is 30.0°C?
4. A sample of nitrogen occupies 300. ml at STP. Under what pressure would this sample occupy 150. ml if
the temperature were increased to 546 °C?
5. A sample of helium occupies 250. ml at 100 °C under a pressure of 1.00 atm. If the pressure were increased
to 1900. mm Hg, to what temperature would the sample have to be heated to occupy a volume of 300. ml?
6. How many moles of nitrogen are contained in 328 ml of the gas under a pressure of 3040. mm Hg at 527 °C.
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7. Calculate the volume exerted by 80.0 g of CH4 at 127 °C under a pressure of 1520 mm Hg.
8. What pressure would a mixture of 3.20g of O2, 6.40 g of CH4, and 6.40 g of SO2 exert if the gases were
placed in a 40.0 L container at 127 °C? Hint. Find the number of moles of each gas and then add them up to
find total number of moles. Then use the ideal gas law.
9. Fluorine will effuse at a rate of 7.2 m/s at a certain temperature. What speed will sulfur dioxide gas effuse at
if the temperature is the same?
10. If 12.2 grams of aluminum chlorate decomposes when heated, how many Liters of oxygen will be produced
at a temperature of -34 °C and a pressure of 19.2 psi?
11. If 18.3 g of tin are mixed with steam at a temperature of 134° C and a pressure of 18 atm, how many liters
of hydrogen will be produced?
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Pistachios Answer each of the following questions on a separate sheet of paper. You must show work in order to
receive credit. Bring it up to me if you have questions or if you are done.
1. Calculate the mass of solute needed to make 750. mL of 2.50 M sodium chloride solution.
2. What mass of silver nitrate must be used to make 200. mL of 0.100 M solution.
3. Describe in detail how you would make 1.25 L of 0.125 M potassium permanganate solution.
4. How much concentrated acid and how much water must be mixed together to make 1.75 L of 6.00 M
sulfuric acid? Sulfuric Acid concentrate is 18.00 M.
5. What is the final molarity of the solution if you add 17.00g of strontium oxalate and bring to volume of
1.82 L with water, then you take out 15 mL of the initial solution and add it to 77 mL of water?
6. Calculate the molarity of a solution made with 30.0 g of oxalic acid with 500.0 mL of solution.
7. Calculate the molality of a solution made with 24.2 g of copper II nitrate trihydrate (copper II nitrate
with 3 molecules of water attached) in 1300. g of water.
8. If I have a 1.00 M solution of potassium dichromate and I need to make 750. mL of 0.23 M solution,
how do I go about that?
9. If I combine 40.0 mL of 2.00 M sulfuric acid with 40.0 mL of 3.50 M potassium hydroxide, what is the
amount of precipitate that will be formed?
10. If I add 19.2 g of calcium chloride to 2.32 kg of water, what temperature will it freeze at?
11. Calculate the number of grams of sodium chloride needed to make 4.00 kg of water boil at 105 °C.
12. If you were to add 1.93 moles of a non-electrolyte to 1.7 kg of naphthalene, what would the new boiling
point be?
Basil Do all work in your notebook. No work, no credit. You will do all of the odds. If you need more practice
you may do the evens if you want.
1. How many grams of magnesium phosphate, Mg3(PO4)2 are found in 2.8 moles of it?
2. In a lab, you needed to calculate the molar mass of calcium sulfate CaSO42H2O. For 125 grams of it,
how many moles are present?
3. Calculate how many atoms of carbon are found in 7 formula units of iron (III) cyanide, Fe(CN)3.
4. What is the mass of 5 molecules of carbon dioxide?
5. For 567 grams of strontium sulfate, SrSO4, how many formula units are present?
6. For 6.7 x 1012 atoms of xenon, how many moles are present?
7. For 8.9 x 1015 formula units of calcium oxide, CaO, how many grams are present?
8. In 56 moles of sulfur, how many atoms are present?
9. How many grams of magnesium are present in 36.1 moles of it?
10. How many atoms of potassium are present in 11.2 moles of it?
11. For 7 x 1033 atoms of aluminum, how many grams are present?
12. In 15 moles of fluorine, how many grams are present?
13. In 15 moles of sulfur, how many atoms are present?
14. In 89.0 grams of nitrogen, how many atoms are present?
15. For 2.02 x 1026 formula units of magnesium chloride, MgCl2 how many grams are present?
16. For 1.00 x 1012 formula units of ammonium hydroxide, NH4OH, how many moles are present?
17. For 235.1 grams of sodium hydroxide, NaOH, how many formula units are present?
18. For 13 grams of calcium bromide, CaBr2, how many moles are present?
19. What is the mass of 4 molecules of carbon dioxide?
20. What is the mass of 3999 formula units of beryllium bromideBeBr2?
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21. Aspartame is an artificial sweetener that is 160 times sweeter than sucrose (table sugar) when dissolved
in water. It is marketed as Nutra-Sweet. The molecular formula for aspartame is C14H18N2O5.
a. Calculate the molar mass of aspartame.
b. How many moles of molecules are present in 10.0 g of aspartame?
c. Calculate the mass in grams of 1.56 mol aspartame.
d. How many molecules area in 5.0 mg aspartame?
e. How many atoms of nitrogen are in 1.2 g aspartame?
f. What is the mass in grams of 1.0 x 109 molecules of aspartame?
g. What is the mass in grams of one molecule of aspartame?
22. Dimethylnitrosamine, (CH3)2N2O is a carcinogenic (cancer causing) substance that may be formed in
foods, beverages or gastric juices from the reaction of nitrite ion (used as a food preservative) with other
substances.
a. What is the molar mass of dimethylnitrosamine?
b. How many molecules are present in 250 mg dimethylnitrosamine?
c. What is the mass of 0.050 mole dimethylnitrosamine?
d. How many atoms of hydrogen are in 1.0 mole dimethylnitrosamine?
e. What is the mass of 1.0 x106 molecules of dimethylnitrosamine?
f. What is the mass in grams of one molecule of dimethylnitrosamine?
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Oregano
DON’T WRITE ON ME. ALL ANSWERS MUST APPEAR IN YOUR NOTEBOOK. ALL MATH
PROBLEMS INVOLVING FREQ, WAVELENGTH OR ENERGY MUST INCLUDE THE WORK, ALL OF
THE WORK AND LAST OF ALL, THE WORK.
1 meter = 1 x 109 nanometers (nm)
Fill in the blanks in the following table:
"floating point" scientific notation
4.21x10-3
12,900
0.0000834
659,000,000,000,000
3.00x108
1. Convert 522 nm to m.
2. Convert 4.44x10-7 m to nm.
3. What is the wavelength of a photon with frequency 5.7x1014 Hz? What color is it?
4. What is the frequency of a photon with wavelength 6.2x10-7 m? What color is it?
5. What is the energy of a photon with frequency 7.3x1014 Hz? What color is it?
6. What is the energy of a photon with wavelength 600 nm? What color is it?
7. What is the frequency of a photon with wavelength 565 nm? What color is it?
8. What is the frequency of a photon with energy of 6.22x10-19 J? What color is it?
9. What is the wavelength of a photon with energy 5.5x10-19 J? What color is it?
Table12 Colors, & Information
color wavelength () frequency ()
[nm] [x1014 Hz]
UV <400 >7.5
violet 400 to 420 7.5 to 7.1
blue 420 to 490 7.1 to 6.1
green 490 to 580 6.1 to 5.2
yellow 580 to 590 5.2 to 5.1
orange 590 to 650 5.1 to 4.6
red 650 to 700 4.6 to 4.3
IR >700 <4.3
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Steak
Write the formulas for the following covalent compounds on your paper:
1. antimony tribromide
2. hexaboron monosilicide
3. chlorine dioxide
4. fluorine moniodide
5. iodine pentafluoride
6. dinitrogen trioxide
7. phosphorus triiodide
Write the names for the following covalent compounds:
8. P4S5
9. O2
10. SeF6
11. Si2Br6
12. SCl4
13. CH4
14. B2Si
15. NF3
Brisket
Name the following chemical compounds on your paper:
1) NaBr
2) Ca(C2H3O2)2
3) P2O5
4) Ti(SO4)2
5) FePO4
6) K3N
7) SO2
8) CuOH
9) Zn(NO2)2
10) V2S3
11) HCl
Write the formulas for the following chemical compounds:
12) silicon dioxide
13) nickel (III) sulfide
14) manganese (II) phosphate
15) silver acetate
16) diboron tetrabromide
17) magnesium sulfate
18) potassium carbonate
19) ammonium hydrogen phosphate
20) tin (IV) selenide
21) carbon tetrachloride
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Ribs Do the following in your notebook. On your paper. The sheet is yours to keep.
Ionic Compounds. Given the name, write the formula. Given the formula, write the name.
1. sodium hydroxide
2. mercury (I) sulfate
3. calcium hypochlorite
4. lead (II) phosphate
5. aluminum chlorate
6. ammonium sulfide
7. copper (I) oxalate
8. antimony (V) oxide
9. manganese (III) sulfite
10. silver oxide
11. zinc nitrite
12. chromium (III) silicate
13. ammonium dichromate
14. iron (III) selenide
15. UCl5
16. Sn3(PO3)2
17. WH6
18. Co(ClO4)2
19. KCN
20. (NH4)2SO4
21. Sc3AsO4
Molecular Compounds. Given the name, write the formula. Given the formula, write the name.
22. N2F
23. S2O7
24. PCl5
25. SiO2
26. CCl4
27. sulfur heptoxide
28. tetraselenium diiodide
29. dinitrogen monoxide
30. pentaphosphorus nonabromide
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Chicken
Given the formula, write the name. Given the name, write the formula. All answers must be on your paper.
You must also record whether it is an acid, ionic or molecular.
1. NaHSO4
2. KClO4
3. H3As
4. calcium hydroxide
5. barium cyanide
6. KH2PO4
7. P4O10
8. Ba(NO2)2
9. magnesium bromide
10. ammonium dichromate
11. cadmium arsenate
12. silver chloride
13. ZnC2O4
14. carbon disulfide
15. PCl3
16. carbonic acid
17. sulfur hexafluoride
18. H2SO3
19. HCN
20. Copper I peroxide
21. chromium II bicarbonate
22. ammonium acetate
23. H2CrO4
24. C2O4
25. gold I nitride
26. disulphur heptaiodide
27. hexarsenic nonafluoride
28. oxygen
29. zinc nitrate
30. ammonium sulfide
31. perchloric acid
32. calcium thiosulfate
33. Ni(NO3)2
34. Pb(SO4)2
35. aluminum thiocyanate
36. manganese IV dihydrogen phosphate
Sausage
Do all work in your notebook. Do not write on this sheet. Follow all sig fig rules.
Calculate the percentage composition of the following compounds:
1. a. Fe2O3 b. Ag2O c. HgO d. sodium sulfide
e. sodium sulfate
2. Calculate the percentage of nitrogen in each of the following compounds:
a. NH4NO3 b. ammonium sulfate c. nitrous acid
3. Calculate the percent composition of iron III dichromate.
4. Calculate the percent composition of potassium oxalate.
5. Calculate the percent composition of aluminum chlorite.
6. Calculate the percent composition of Lead (IV) chloride
7. Calculate the percent composition of disulfur heptabromide.
8. Calculate the percent composition of calcium nitrate.
9. If you have 27.92 g of iron III dichromate, how many grams of iron are in the compound? (See percent
you calculated in #3)
10. If you have 1920. g of sodium sulfide, how many grams of sulfur are there in the compound? (See
percent calculated in #1d)
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Pork Butt
Naming Acids Worksheet. All answers on your paper.
1. Nitric acid
2. Hydrocyanic acid
3. Chloric acid
4. Acetic acid
5. Hydrobromic acid
6. Sulfurous acid
7. Chlorous acid
8. Boric acid
9. Hydrochloric acid
10. Phosphoric acid
11. Nitrous acid
12. Hydrofluoric acid
13. Perchloric acid
14. Hydroiodic acid
15. Phosphorous acid
16. Carbonic acid
17. Sulfuric acid
18. Formic acid (the formate ion is COOH-)
19. Thiocyanic acid
Name each of the following acids:
1. HClO4
2. HCOOH
3. H3PO4
4. HCl
5. H3BO3
6. H2SO4
7. HNO2
8. HI
9. HCH3COO
10. HF
11. H3PO3
12. HCN
13. HClO3
14. H2CO3
15. H2SO3
16. HClO2
17. HNO3
18. HBr
19. H2S2O3
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BACON
1. calcium iodide
2. magnesium nitrate
3. iron(II) acetate
4. lead(IV) oxide
5. calcium hydride
6. copper(II) sulfide
7. ammonia
8. hydrosulfuric acid
9. magnesium nitrate
10. iron(II) thiocyanate
11. calcium selenide
12. magnesium sulfide
13. sodium hydride
14. lead(IV) oxide
15. potassium chromate
16. cadmium bromide
17. copper(II) sulfide
18. aluminum acetate
19. ammonium hydrogen phosphate
20. antimony(III) phosphide
21. antimony(V) sulfite
22. vanadium(V) carbonate
23. nickel(II) sulfate
24. arsenic(II) sulfite
25. tin(II) fluoride
26. lead(IV) sulfide
27. hydrobromic acid
28. lead(II) chlorate
29. manganese(II) oxalate
30. carbon tetrachloride
31. manganese(III) dichromate
32. ammonium sulfate
33. mercury(II) hydrogensulfate
34. mercury(I) nitrite
35. cobalt(II) hydroxide
36. tin(IV) sulfide
37. lithium perchlorate
38. manganese(III) permanganate
39. tungsten(VI) phosphate
40. sulfur dioxide
41. copper(II) sulfate
42. gold(I) phosphide
43. zinc nitrite
44. rubidium hydride
45. mercury(I) phosphate
46. barium nitride
47. mercury(II) phosphate
48. cesium cyanide
49. tetraphosphorus decoxide
50. silver cyanide
51. strontium peroxide
52. carbon disulfide
53. cesium dihydrogen phosphate
54. rubidium chromate
55. tungsten(VI) sulfite
56. ammonium phosphate
57. disulfur dichloride
58. sodium hydrogen phosphate
59. platinum(IV) oxide
60. lithium fluoride
61. tin(II) hydrogen carbonate
62. chromium(VI) phosphide
63. Lead(IV) dichromate
64. chromium(III) acetate
65. gold arsenide
66. nickel(II) carbonate
67. sodium thiocyanate
68. aluminum oxalate
69. ammonium perbromate
70. tin(IV) hydrogen sulfite
71. KCN
72. titanium(IV) sulfite
73. sulfur trioxide
74. lithium permanganate
75. SF6
76. dinitrogen pentoxide
77. NO
78. KSCN
79. CO
80. HF
81. calcium hypochlorite
82. magnesium hydroxide
83. ammonium hypoiodite
84. bismuth(III) bromide
85. HCl(aq)
86. HCN
87. CO2
88. CO2
89. NO2
90. KOCN
91. sulfite ion
92. (NH4)2C2O4
93. ZnS
94. CuCl2
95. MgCl2
96. PCl5
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97. Ca(ClO3)2
98. LiNO2
99. CaSO4
100. KH2PO4
101. AgNO3
102. CuCN
103. H2S
104. KHCO3
105. CaO
106. NaHSO4
107. H2CO3
108. Li2HPO4
109. Mg3(PO4)2
110. H3PO3
111. KCl
112. MgSO4
113. K2O
114. Ca(IO2)2
115. Al(NO2)3
116. SiO2
117. MgO
118. CuCl
119. SnI
120. KClO2
121. AsCl5
122. CaSO3
123. CuSO3
124. NaBr
125. HF
126. P2O3
127. FeSO4
128. HClO
129. SnCl4
130. NO2
131. AsCl3
132. NaH
133. KCN
134. ZnS
135. NH4OH
136. Pb(NO3)2
137. Fe(ClO4)3
138. H2Se
139. HNO2
140. H3PO4
141. CS2
142. CaH2
143. bromic acid
144. persulfuric acid
145. ammonia
146. carbon tetrachloride
147. CO2
148. hydrosulfuric acid
149. dinitrogen pentoxide
150. H2Se
151. P2O3
152. CS2
153. HF
154. NO2
155. H2S
156. PCl5
157. HCl
158. sulfur dioxide
159. CO
160. CO2
161. NO
162. HF
163. SF6
164. NO2
165. sulfur trioxide
166. dinitrogen pentoxide
167. tetraphosphorus decoxide
168. carbon disulfide
169. disulfur dichloride
170. HI
171. boron trifluoride
172. diphosphorus trioxide
173. carbon tetrachloride
174. P4O10
175. iodine monochloride
176. hydrobromic acid
177. sulfur trioxide
178. hydrofluoric acid
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Tofu
Write formulas for the following
1. sodium nitrite
2. sodium carbonate
3. sodium sulfate
4. potassium hydroxide
5. potassium nitrate
6. potassium nitrite
7. potassium phosphate
8. cadmium hydroxide
9. cadmium carbonate
10. cadmium sulfate
11. cadmium phosphate
12. aluminum bromide
13. aluminum nitrate
14. aluminum sulfide
15. magnesium nitrate
16. magnesium oxalate
17. magnesium hypochlorite
18. barium bromide
19. barium nitrate
20. barium carbonate
21. iron II arsenate
22. iron II chloride
23. iron II hydroxide
24. iron II carbonate
25. iron II sulfite
26. Copper II permanganate
27. iron II phosphate
28. iron III fluoride
29. strontium chloride
30. strontium hydroxide
31. strontium nitride
32. strontium sulfide
33. strontium sulfite
34. iron III sulfate
35. iron III arsenate
36. mercury II bromide
37. lead IV carbonate
38. tin IV hydrogen phosphate
39. chlorine heptoxide
40. chlorine trifluoride
41. sulfur hexafluoride
42. dichlorine monoxide
43. chlorine dioxide
44. carbon tetrasulfide
45. diphosporus heptafluoride
46. sulfur nonafluoride
47. hypochlorous acid
48. nitric acid
49. nitrous acid
50. chloric acid
51. hydrochloric acid
52. chlorous acid
53. sulfuric acid
54. sulfurous acid
55. hydrosulfuric acid
56. hydrofluoric acid
57. oxalic acid
Write names for the following
:
58. NaNO3
59. Na2SO4
60. Na3PO4
61. KNO2
62. K2CO3
63. K2SO4
64. CdBr2
65. Cd(NO3)2
66. CdSO3
67. CdS
68. Al4C3
69. Al(OH)3
70. Mg(NO2)2
71. MgSO3
72. Mg3(PO4)2
73. Ba(NO2)2
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74. BaSO3
75. BaCO3
76. Al(HCO3)3
77. Al2(SO4)3
78. AlAsO4
79. FeBr2
80. Co2(SO3)3
81. CoS
82. Co(C2O4)2
83. Ba(CH3COO)2
84. SrBr2
85. Sr(H2PO4)2
86. Sn(ClO4)2
87. Sr(CN)2
88. Fe(NO3)2
89. NaOH
90. K2S
91. NO
92. N2O3
93. N3O7
94. NCl7
95. P4O10
96. F2
97. PCl5
98. N2O4
99. HClO4
100. H2S
101. HBr
102. H3P
103. H2C2O4
104. HClO
105. HClO3
106. HClO2
107. H2CO3
108. HCH3COO
109. HNO2
110. HNO3
111. H2SO3
112. H2SO4
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Nutmeg
Recall that the symbol [H+] means “H+ ion concentration in moles/liter (M).”
Recall that pH is defined in the following fashion: pH = -log[H+] it is the negative log of the
concentration of H+ ions.
If you want to find [H+] from pH you have to take the inverse log: [H+] = 10(-pH)
Remember to change sign before you take the inverse log.
Inverse logging is done with the 10x key on your calculator. (2nd function, LOG key)
1. Using universal indicator we determine the pH of a solution to be 6.5. What is the [H+] ?
2. Using pH paper we determine the pH of a solution to be 13.2. What is the [H+] ?
3. We measure the concentration of H+ ions to be 1.00 x 10-5 M. What is the pH of this solution?
4. Very strong acids such as concentrated H2SO4 can cause severe injuries, or put large holes in
clothing. This acid has a [ H+] equal to 36.0 M. What is the pH of concentrated sulfuric acid?
5. What is the pH if [H+] is 4.3 x 10-6 M ?
6. What is the pH if [H+] is 3.87 x 10-14 M ?
7. What is the pH if [H+] is 9.1 x 10-3 M ?
8. What is the pH if [H+] is 2.22 x 10-8 M ?
9. What is the pH if [H+] is 1.5 x 10-9 M ?
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10. What is the pH if [H+] is 5.0 x 10-11 M ?
11. What is the pOH if [H+] is 2.5 x 10-5 M ?
12. What is the pH if [H+] is 1.0 x 10-10 M ?
13. What is the pH if [H+] is 6.7 x 10+1 M ?
14. What is the pOH if [H+] is 8.0 x 10-13 M ?
15. What is the [OH-] if the pH is 8.7 ?
16. What is the [H+] if the pH is 3.2 ?
17. What is the [H+] if the pOH is 4.5 ?
18. What is the [H+] if the pH is -1.5 ?
19. What is the [H+] if the pOH is 6.9 ?
20. What is the [H+] if the pH is 12.8 ?
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Salami
Procedure:
Open the internet browser and enter the address: http://phet.colorado.edu
Click on “Play with Sims” and select “Chemistry” from the menu on the left.
Open the “States of Matter” Simulation and select “Run Now”
Investigation:
1. Predict what the molecules of a solid, liquid and gas look like. Illustrate your prediction with a drawing.
Solid Liquid Gas
2. Complete the table below by exploring the “Solid, Liquid, Gas” tab in the simulation. Test your
predictions and record your observations by recording the temperature and illustrations of each
substance in the three states of matter.
Substances Observations
Solid Liquid Gas
Neon
Temperature:
Illustration:
Temperature:
Illustration:
Temperature:
Illustration:
Argon
Temperature:
Illustration:
Temperature:
Illustration:
Temperature:
Illustration:
Oxygen
Temperature:
Illustration:
Temperature:
Illustration:
Temperature:
Illustration:
Learning Goal: Students will be able to demonstrate their knowledge of the states of matter through illustrations
and descriptions. These illustrations and descriptions should include:
How the molecules in a solid, liquid and gas compare to each other.
How temperature relates to the kinetic energy of molecules.
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Water
Temperature:
Illustration:
Temperature:
Illustration:
Temperature:
Illustration:
3. Sketch a graph of Kinetic Energy vs. Temperature. Use this graph to describe the relationship between
the two concepts.
4. Write a summary paragraph, which includes drawings, to demonstrate you have mastered the learning
goal. Be sure to incorporate both concepts of the learning goal:
How the molecules in a solid, liquid and gas compare to each other.
How temperature relates to the kinetic energy of molecules.
5. Explain this phase diagram by relating what you know about temperature, states of matter and pressure.
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Reuben
Today you are going to use the “Reactants, Products and Leftovers” simulation to explore how many products you can make given the initial amounts of two reactants. Go to the link on edmodo for the simulation. Using your laptop fill in this document and upload to edmodo. Each person should upload the file into their edmodo account
CONCEPT AREA 1: MAKING SANDWICHES 1. If you have 6 pieces of bread and 4 slices of cheese, predict how many cheese sandwiches of type A you can make. Then predict how many of type B you can make.
A: B: How did you figure this out? Now check your predictions using the “Sandwich Shop” tab. Do the results make sense? Revise your answers or reasoning as needed. In case A, the bread could be called the “limiting reactant.” How would you define a “limiting reactant”? What is the limiting reactant for case B and why? What is leftover when all the sandwiches are made?
CONCEPT AREA 2: MAKING AMMONIA 2. Consider the chemical equation: 1 N2 + 3 H2 → 2 NH3 For the 3 scenarios below, predict which one will produce the most ammonia, and predict which ones will have leftovers.
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A: B: C: Explain your reasoning: Now check your predictions using the “Real Reaction” tab. Do the results make sense? Revise your answers or reasoning as needed. How did the “Real Reaction” tab relate to the “Sandwich Shop” tab? 3. Play at least one “Game!” at each level with your partner (estimated time = 5 minutes per game). Record your score for each level in the table below.
Level Type one person’s
name here
Type the other
person’s name here
1 2 3
How did you solve the problems? Write your strategy in the space below. Did your strategy change as you played the game? If so, write how it changed.
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Quinoa
http://phet.colorado.edu/new/simulations/sims.php?sim=Gas_Properties
If the direct link does not work, use Google, and use the search terms “Phet Gas properties”.
Part 1: Play
Purpose: Play! See how everything works before we try to find any relationships.
Procedure:
1. Pump the handle once for the heavy particles. Watch the temperature and pressure gauge. See how long it takes for
the values to stabilize. Record the values in Table 1 on the next page.
2. Using the little person pushing on the wall, decrease the volume of the box in about half. Record the results in Table 1.
If it explodes reset it and this time don’t pump as much gas.
3. Using the little person pushing on the wall, make the volume as small as possible and watch for at least 60 seconds.
Thoroughly describe what happens over the course of the full minute. Be sure to address the following in your answer:
a. What happens to the container and the little man pushing on the container?
b. What happens to temperature and pressure?
4. Hit the “Reset” button. Give the pump a push with the heavy species. Wait for the values to stabilize and record the
results in table 1 on the next page.
Are these values the exact same as the first time you did it? __________ Why or why not?
5. There is a box in the lower right corner entitled “Gas in Pump”. Select the “light species”. You will notice that the
pump turns red. Give the pump a press. Wait for the values to stabilize and record the results in table 1 on the next
page.
6. In the box entitled “Heat Control”, grab the arrow and move it to add. What happens?
7. Using the “Heat Control”, grab the arrow and move it to “remove.” What happens?
8. Feel free to play with the simulator. Try the “Pause” and “Step” buttons at the bottom. Try sliding the top of the
container. Take about 3-5 minutes to play with the different options.
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Table 1: Playing with the simulation program
After pumping the
“heavy” handle
once (1st time)
After cutting the
volume in about
half
Reset: Constant
volume 1 pump of
heavy species
Added 1
pump of light
species
Temperature (K)
Pressure (atm)
Number of Heavy
Species
Number of Light
Species
Part 2: Comparing the effect of particle size.
Purpose: Your goal is to see if there is a difference between using heavy or light particles.
Procedure:
1. Hit the Reset button.
2. Under the “Measurement Tools” section, be sure to have the “Energy Histograms” checked.
3. Under the “Advanced Options” tab, be sure that the temperature is check with a value of 300 K.
4. Set the “Constant Parameter” to be volume.
5. Type in the appropriate number of particles in the boxes on the far right of the screen.
6. Fill in the data table below.
7. When finished with one trial, repeat from step 1 (be sure to reset, just to be safe).
8. For the last rows of the table, chose a combination of heavy and light particles to a total of 500.
Table 6: Measurement of Pressure, Temperature, and moles.
Trial 1 Trial 2 Trial 3 Average
100 Heavy
Particles
Pressure (atm)
Temperature (K)
100 Light
Particles
Pressure (atm)
Temperature (K)
50 of each
Type
Pressure (atm)
Temperature (K)
___ heavy
___ light
Pressure (atm)
Temperature (K)
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Questions:
1. Is there a significant difference between the averages of the 100 light particles and the 100 heavy?
2. How does the particle size and mass change the pressure or temperature?
3. Describe how the lighter particles move compared to the heavy particles.
4. Considering that all the particles in the container are at the same temperature, why do the particles move at
different speeds?
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Turkey
Mass to Mass Problems
In the following problems, calculate how much of the indicated product is made. Show all your work.
1) NaOH + HCl ---> NaCl + H2O
If you start with 10.0 grams of sodium hydroxide, how many grams of sodium chloride will be
produced?
2) C2H4 + O2 ---> CO2 + H2O
If you start with 25 grams of ethylene (C2H4), how many grams of carbon dioxide will be produced?
3) Li + CaCl2 ---> LiCl + Ca
If you start with 5.5 grams of lithium chloride, how many grams of calcium chloride will be produced?
4) HCl + NaCH3COO ---> NaCl + CH3COOH
If you start with 194.67 grams of sodium acetate, how many grams of acetic acid will be produced?
Answer the following problems:
5) CaBr2 + KCl ---> CaCl2 + KBr
How many grams of potassium chloride would you need to make 723 grams of potassium bromide?
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6) C4H8 + O2 ---> CO2 + H2O
How many grams of 2-butene (C4H8) would you need to make 50.2 grams of carbon dioxide?
7) FeCl3 + H2O ---> FeCl3 + H2O
How many grams of water would you need to completely react with 250.5 grams of iron (III) chloride?
8) H2CO3 ---> H2O + CO2
If 15 grams of water was made in this reaction, how many grams of carbon dioxide were made?
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Stuffing
Stoichiometry Calculations For the following questions, balance the equation and calculate grams of the underlined product. Show all your
work.
1) Zn + H2SO4 --> ZnSO4 + H2 (You begin with 10. grams of zinc)
2) CaO + H2O --> Ca(OH)2 (You begin with 55.8 grams CaO)
3) ZnS + O2 --> ZnO + SO2 (45 grams of ZnO is produced)
4) Fe(NO3)3 + NH4OH --> Fe(OH)3 + NH4NO3 (17.8 g. NH4NO3 produced)
5) Fe + O2 --> Fe2O3 (You start with 500.0 grams of iron)
6) NaCl + MgBr2 --> NaBr + MgCl2 (You start with 34 grams of MgBr2)
7) PbCO3 + HCl --> PbCl2 + H2CO3 (You start with 21 grams of HCl)
For the following questions, write a balanced equation and calculate grams called for in the problem. Show all
your work.
8) How many g of magnesium sulfide is formed when 35.0 grams of magnesium reacts with an excess of
sulfur?
9) How many grams of beryllium hydroxide is needed to react completely with 100 grams of hydrochloric
acid?
10) How many g of iron (II) nitrate is formed when 25 grams of iron is dissolved in pure nitric acid?
11) How many grams of lead (II) sulfide is precipitated in the reaction between 21.0 grams of lead (II)
bromide and excess hydrogen sulfide?
12) How many g of lithium chloride is produced by the reaction of 230 grams of lithium metal and excess
chlorine gas?
13) How many L of carbon dioxide is formed when 23 grams of C6H8 is burned in an excess of oxygen gas
at STP? ( 1.00 mole of any gas at STP = 22.4 L)
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Cranberry Sauce
WRITE ON ME but only work and answers on your own paper will be graded.
Assume all volumes at STP. You must answer all questions on your own paper. All work must
be shown and all units must be labeled.
1. How many liters of hydrogen are produced if 4.00g of zinc react with hydrochloric
acid?
2. In the reaction between aluminum and oxygen, how many grams of aluminum are
required to react with 50.10 L of oxygen?
3. In the Haber process, hydrogen gas and nitrogen gas are combined to produce
ammonia. If 60.0 L of ammonia are produced, how many liters of hydrogen and
nitrogen are necessary?
4. If 15.0 L of ethane gas, C2H6, is burned completely to form carbon dioxide and water,
calculate the following:
a. Liters of oxygen necessary for combustion.
b. Liters of water vapor produced.
5. Magnesium is placed in a crucible and allowed to react with oxygen in the
atmosphere. If 46.20 g of magnesium is used, how many liters of oxygen will be used
in the reaction?
6. Silver nitrate and barium chloride solutions are combined in a test tube. How many
grams of the precipitate will be made if 8.43 g of silver nitrate is used?
7. Calcium carbonate is heated. If 99.24g of calcium carbonate is used, how many moles
of gas will be produced?
8. Silver and hydrochloric acid are combined in a beaker. If 7.00 moles of hydrochloric
acid are used, how many grams of silver will be dissolved by the acid?
9. You take water and apply an electric current to it to separate it. If 36.02g of water are
used, how many Liters of gas will be produced?
10. Mr. Elegante can eat 37 pieces of bacon an hour and he does so for 3 hours, then
proceeds to eat slices of salami at a rate of 89 an hour for 3 additional hours. He now
thinks better of himself and proceeds to ingest oat bran at a rate of 370 g per 10
minute period for a period of 18.3 minutes. Assuming that bacon and salami can be
found on the periodic table and calculated as such and also assuming that oat bran will
not negate the effects of ingesting large amounts of fat, how much time did you waste
reading this question?
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Green Bean Casserole
Do your work on this sheet. Show all work and label all units. Remember your significant figures.
1. Molten iron and carbon monoxide are produced in a blast furnace by the reactions of iron III oxide and
carbon. If 2.50 x 104 g of iron III oxide are used, how many moles of iron can be produced?
2. How many grams of copper II oxide can be reduced to copper metal with 10.0 liters of H2 at STP?
3. In a reaction between 19.5 g potassium and excess chloric acid at STP, 3.2 L of gas is produced. What is
the percent yield?
4. If 15.0 liters of ethane gas (C2H6) at STP are burned completely to form carbon dioxide and water, calculate
the liters of oxygen required to burn the ethane.
5. What reagent is limiting if 3.00 liters of chlorine gas at STP react with a solution containing 25.0 grams of
NaBr?
6. If 20.0 grams of KOH react with 15.0 grams of (NH4)2SO4, calculate the following:
a. The moles of K2SO4 produced.
b. The grams of NH4OH produced.
7. How much iron must be combined with silver nitrate to produce 18.29 g Ag assuming the reactions runs at
7.23 % yield and iron takes on a +1 charge?
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8. A detergent, sodium dodecanesulfonate, C12H25OSO3Na, is made by the reaction of H2SO4 and NaOH on
dodecanol, C12H25OH. Water is also a product. If a manufacturer needs to produce 3.00 x 107 grams of
detergent daily, how many grams of dodecanol are needed and the reaction runs at 71.7% yield?
9. Calcium carbide (CaC2) reacts with water to produce calcium hydroxide and acetylene (C2H2). What
volume of the gas acetylene could be produced if you react 50.0 grams of CaC2 and 50.0 grams of water?
10. In the reaction at STP between hydrogen and nitrogen to produce ammonia, NH3, 50.0 liters of hydrogen
react with 15.0 liters of nitrogen. Calculate the grams of ammonia produced.
11. Methyl orange is a colored indicator, which is red in the presence of acids and yellow in the presence of
bases. What will be the color of the solution which results from mixing two separate solutions containing
10.0 grams of H2SO4 and 7.00 grams NaOH after a small amount of methyl orange is added?
11. Magnesium acetate can be prepared by a reaction involving 15.0 grams of iron II acetate with either 10.0
grams of MgCrO4 or 15.0 grams of MgSO4. Which reaction will give the most magnesium acetate? How many
grams of magnesium acetate will that be?
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Mashed Potatoes
Limiting Reagents and Percentage Yield Worksheet 1. a) 80.0 grams of diiodine pentoxide, reacts with 28.0 grams of carbon monoxide to form iodine and carbon
dioxide. Determine the mass of iodine which could be produced.
b) If, in the above situation, only 0.160 moles, of iodine, I2 was produced.
i) what mass of iodine was produced?
ii) what percentage yield of iodine was produced.
2. Zinc and sulphur react to form zinc sulphide.
If 25.0 g of zinc and 30.0 g of sulphur are mixed,
a) Which chemical is the limiting reactant?
b) How many grams of ZnS will be formed?
c) How many grams of the excess reactant will remain after the reaction is over?
3. Which element is in excess when 3.00 grams of Mg is ignited in 2.20 grams of pure oxygen? What mass of
excess remains? What mass of MgO is formed?
4.How many grams of Al2S3 are formed when 5.00 grams of Al is heated with 10.0 grams S?
5. When Molybdenum VI oxide and Zn are heated together they react to form molybdenum III oxide and zinc
oxide
What mass of ZnO is formed when 20.0 grams of MoO3 is reacted with 10.0 grams of Zn?
6. Silver nitrate reacts with iron III chloride to give silver chloride and iron III nitrate. In a particular
experiment, it was planned to mix a solution containing 25.0 g of silver nitrate with another solution
containing 45.0 grams of iron III chloride.
a) Write the chemical equation for the reaction.
b) Which reactant is the limiting reactant?
c) What is the maximum number of moles of AgCl that could be obtained from this mixture?
d) What is the maximum number of grams of AgCl that could be obtained?
e) How many grams of the reactant in excess will remain after the reaction is over?
7. Solid calcium carbonate, CaCO3, is able to remove sulfur dioxide from waste gases by the balanced reaction:
CaCO3 + SO2 + other reactants ------> CaSO3 + other products
In a particular experiment, 255 g of CaCO3 was exposed to 135 g of SO2 in the presence of an excess amount of
the other chemicals required for the reaction.
a) What is the theoretical yield of CaSO3?
b) If only 198 g of CaSO3 was isolated from the products, what was the percentage yield of CaSO3 in this
experiment?
8. A research supervisor told a chemist to make 100. g of chlorobenzene from the reaction of benzene with
chlorine and to expect a yield no higher that 65%. What is the minimum quantity of benzene that can give
100. g of chlorobenzene if the yield is 65%? The equation for the reaction is:
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C6H6 + Cl2 -----------> C6H5Cl + HCl
benzene chlorobenzene
9. Certain salts of benzoic acid have been used as food additives for decades. The potassium salt of benzoic
acid, potassium benzoate, can be made by the action of potassium permanganate on toluene.
C7H8 + 2 KMnO4 -------> KC7H5O2 + 2 MnO2 + KOH + H2O
toluene potassium
benzoate
If the yield of potassium benzoate cannot realistically be expected to be more than 68%, what is the minimum
number of grams of toluene needed to achieve this yield while producing 10.0 g of KC7H5O2?
10. Aluminum dissolves in an aqueous solution of NaOH according to the following reaction:
2 NaOH + 2 Al + 2 H2O -----> 2 NaAlO2 + 3 H2
If 84.1 g of NaOH and 51.0 g of Al react:
i) Which is the limiting reagent?
ii) How much of the excess reagent remains?
iii) What mass of hydrogen is produced?
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Pumpkin Pie
Super Fun Stoichiometry Problems!!!!!
1. The energy used to power one of the Apollo lunar missions was supplied by the following overall reation:
N2H4 + (CH3)2N2H2 + N2O4 N2 + CO2 + H2O
For the phase of the mission when the lunar module ascended from the surface of the moon, a total of 1200.0 kg
N2H4 was available to react with 1000.0 kg (CH3)2N2H2 and 4500.0 kg N2O4.
a. For this portion of the flight, which of the allocated components was used up first?
b. How much water in kilograms was put into the lunar atmosphere.
2. The Ostwald process for producing nitric acid from ammonia consists of the following steps:
NH3 + O2 NO + H2O
NO + O2 NO2
NO2 + H2O HNO3 + NO
If the yield in each step is 94%, how many grams of nitric acid can be produced from 5.00 kg of ammonia?
3. Magnesium is obtained from seawater when calcium hydroxide is added to it to precipitate magnesium
hydroxide. The precipitate is filtered and reacted with hydrochloric acid to produce magnesium chloride. The
magnesium chloride is then electrolyzed to produce Mg and Cl2. If you are working at the magnesium making
factory and wanted to produce 500.0 g of Mg, how many grams of magnesium hydroxide would have to be
precipitated from the seawater?
4. Phosphate baking powder is a mixture of starch, sodium bicarbonate, and calcium dihydrogen phosphate.
When mixed with water, phosphate baking powder releases carbon dioxide gas, causing a dough to rise.
NaHCO3 + Ca(H2PO4)2 Na2HPO4 + CaHPO4 + CO2 + H2O
If 0.750 L of carbon dioxide is needed for a cake and each kg of baking powder contains 168 g of sodium
bicarbonate, how many grams of baking powder must be used to generate this amount of carbon dioxide?
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Reeses Cup
Writing Complete Equations Practice
For each of the following problems, write complete chemical equations to describe the chemical process taking
place. Write them in the space to the right of the equations below
1) When lithium hydroxide pellets are added to a solution of sulfuric acid, lithium sulfate solution and
water are formed.
2) If a copper coil is placed into a solution of silver nitrate, silver crystals form on the surface of the
copper. Additionally, highly soluble (meaning it dissolves in water) copper (I) nitrate is generated.
3) When crystalline C6H12O6 is burned in gaseous oxygen, carbon dioxide gas and water vapor are formed.
Balance the equations below:
1) ____ N2 + ____ H2 ____ NH3
2) ____ KClO3 ____ KCl + ____ O2
3) ____ NaCl + ____ F2 ____ NaF + ____ Cl2
4) ____ H2 + ____ O2 ____ H2O
5) ____ Pb(OH)2 + ____ HCl ____ H2O + ____ PbCl2
6) ____ AlBr3 + ____ K2SO4 ____ KBr + ____ Al2(SO4)3
7) ____ CH4 + ____ O2 ____ CO2 + ____ H2O
8) ____ C3H8 + ____ O2 ____ CO2 + ____ H2O
9) ____ C8H18 + ____ O2 ____ CO2 + ____ H2O
10) ____ FeCl3 + ____ NaOH ____ Fe(OH)3 + ____NaCl
11) ____ P + ____O2 ____P2O5
12) ____ Na + ____ H2O ____ NaOH + ____H2
13) ____ Ag2O ____ Ag + ____O2
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14) ____ S8 + ____O2 ____ SO3
15) ____ CO2 + ____ H2O ____ C6H12O6 + ____O2
16) ____ K + ____ MgBr2 ____ KBr + ____ Mg
17) ____ HCl + ____ CaCO3 ____ CaCl2 + ____H2O + ____ CO2
18) ____ HNO3 + ____ NaHCO3 ____ NaNO3 + ____ H2O + ____ CO2
19) ____ H2O + ____ O2 ____ H2O2
20) ____ NaBr + ____ CaF2 ____ NaF + ____ CaBr2
21) ____ H2SO4 + ____ NaNO2 ____ HNO2 + ____ Na2SO4
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Nerds Synthesis and Decomposition Worksheet
Directions: For each reaction, identify it as Synthesis or Decomposition. Also, write down the
rule number that applies to each reaction. Predict the products. Be sure to balance the equation
and show states of matter if you can. Do all answers in your notebook.
1. Solid sodium chlorate is heated.
2. Solid potassium oxide is mixed with water.
3. Dinitrogen trioxide gas is combined with water
4. Solid lithium carbonate is heated.
5. Water undergoes electrolysis.
6. Carbon dioxide gas and barium oxide are reacted.
7. Aluminum carbonate is heated.
8. Sulfur trioxide gas and water are allowed to react
9. Phosphoric acid is heated.
10. Cadmium hydroxide solution is heated.
Solid rubidium oxide is combined with water.
11. Diphophorus trioxide gas is combined with silver oxide solution.
Milk Duds
In your notebook please
1. Sulfur dioxide gas is bubbled through water
2. Lithium oxide is added to water.
3. Aluminum chloride is electrolyzed
4. A small piece of sodium is added to a container of iodine vapor.
5. Carbonic acid is heated.
6. Lead is added to a solution of magnesium acetate
7. Copper is added to a solution of hydrochloric acid
8. Iodine crystals are added to a solution of sodium chloride
9. Calcium metal is added to a solution of nitrous acid
10. A solution of iron III chloride is poured over a piece of platinum wire.
11. Zinc pellets are added to hydrobromic acid
12. Zinc acetate solution and cesium hydroxide solution are combined.
13. Ammonium sulfide solution is reacted with hydrochloric acid
14. Solid calcium carbonate is reacted with sulfuric acid
15. Carbon dioxide gas is reacted with solid potassium oxide
16. Cobalt II hydroxide is added to hydroiodic acid.
17. A piece of chromium is immersed in a solution of tin IV thiosulfate.
18. You pick up a straw and blow carbon dioxide into a glass of water.
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Candy Corn
In your notebook please
1. iron + hydrochloric acid iron III
2. silver nitrate(aq) + potassium chloride(aq)
3. sodium oxide + sulfur dioxide
4. copper + hydrofluoric acid copper II
5. magnesium chlorate
6. calcium + water
7. propane (C3H8) + O2
8. dinitrogen pentoxide + water
9. sodium sulfite(aq) + hydrochloric acid (aq)
10. barium + phosphoric acid
11. chlorine + iron III fluoride
12. water
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Milky Way
Synthesis and Decomposition Worksheet In your notebook please
For each equation, predict the products, balance the equation and give the rule (#1, #2, etc) used to work each
problem.
1. SO2 + H2O
2. CO2 + H2O
3. P2O5 + H2O
4. CaO + H2O
5. SO2 + MgO
6. Sr + P
7. H2CO3 (when heated)
8. Ba(OH)2 (when heated)
9. NaCl (with electricity)
10. KClO3 (when heated)
11. CaCO3 (when heated)
12. Ag2O (when heated)
Single Replacement
1. Fe + Cu(NO3)2
2. Na + H2O
3. Zn + HCl
4. Br2 + KI
5. Li + H2O
6. Fe + KCl
7. Pb + H2O
8. NaCl + Br2
Double Replacement (Use solubility rules)
1. KOH(aq) + HCl(aq)
2. 2KNO3(aq) + H2SO4(aq)
3. AgNO3(aq) + NaCl (aq)
4. HClO3(aq) + NH4OH(aq)
Combustion. For each one of the following, combine the given with oxygen to produce carbon dioxide and
water.
1. CH4
2. C6H12O6
3. C3H8
4. C2H5OH
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Vanilla Tootsie Roll
Chemistry equation practice Directions: Do all work on this sheet, (preferably in pencil). For each of the following word
equations, write the formulas, predict the products, and balance the equation. If no reaction
occurs, write NR in place of products. Write down whether it is a synthesis (S), decomposition
(D), single replacement (SR), double replacement (DR) or combustion (C).
1. Magnesium + oxygen
2. Aluminum + hydrochloric acid
3. Sodium oxide (s) + sulfur dioxide (g)
4. Phosphoric acid (heated)
5. Sodium chlorate
6. Zinc chloride (aq) + ammonium sulfide (aq) \
7. Strontium carbonate (s) (heated)
8. Mercury II sulfate (aq) + ammonium nitrate (aq)
9. Iron + copper II sulfate (aq)
10. Zinc + sulfuric acid
11. Dinitrogen pentoxide (g) + water
12. Chlorine + magnesium iodide (aq)
13. Potassium + water
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14. Iron + hydrochloric acid
15. Cobalt III hydroxide (aq) + nitric acid
16. Bromine + sodium iodide (aq)
17. Sodium hydroxide (aq) + phosphoric acid
18. Ammonium sulfate (aq) + calcium hydroxide (aq)
19. Silver nitrate (aq) + potassium chloride (aq)
20. Magnesium hydroxide (aq) + phosphoric acid
21. Iron II sulfide (s) + hydrochloric acid
22. Ammonium sulfide (aq) + iron II nitrate (aq)
23. Sulfuric acid + potassium hydroxide (aq)
24. Aluminum sulfate (aq) + calcium phosphate (s)
25. C12H22O11 burns
26. Silver acetate (aq) + potassium chromate (aq)
27. Ammonium phosphate (aq) + barium hydroxide (aq)
28. Calcium oxide(s) + diphosphorus pentoxide (s)
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Dots
1. sodium + iodine →
2. magnesium + oxygen →
3. hydrogen + oxygen →
4. nickel (II) chlorate →
5. barium carbonate →
6. zinc hydroxide →
7. aluminum + sulfuric acid →
8. potassium iodide + chlorine →
9. iron + copper (II) nitrate → iron (II)
10. nickel + zinc nitrate → nickel (II)
11. sodium bromide + iodine →
12. potassium + water →
13. silver nitrate + zinc chloride →
14. copper (II) hydroxide + acetic acid →
15. iron (II) sulfate + ammonium sulfide →
16. C3H8 burns
17. C12H22O11 burns
18. Radium bromide + electricity
19. C2H5OH burns
20. C6H12 burns
21. C4H6 burns
22. Magnesium + iodine
23. Copper (II) chloride solution + hydrosulfuric acid (aq)
24. Sodium hydroxide solution + perchloric acid (aq)
25. Zinc carbonate is heated
26. Sodium + magnesium chloride solution
27. Potassium + chlorine
28. Aluminum carbonate is heated
29. Aluminum plus nitrogen
30. Sulfurous acid is heated
31. Iron III hydroxide solution is heated
32. Osmium (IV) oxide + diphosphorus trioxide
33. sodium oxalate solution + aluminum nitrate (oxalates follow the rules of carbonates)
34. Strontium + lithium carbonate
35. Cadmium chloride + fluorine
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Cream of tartar Use your own paper to show all work. 1. What is the concentration of hydroiodic acid if 59.69 mL of it is neutralized by
40.39 mL of a .2968 M lithium hydroxide solution?
2. It requires 65.95 mL of potassium hydroxide solution to neutralize 89.29 mL of
0.2118-M nitric acid. What is the concentration of the potassium hydroxide solution?
3. What is the molarity of a copper(I) hydroxide solution if 50.50 mL of the solution
is titrated to the endpoint with 51.99 mL of 0.3574-M carbonic acid?
4. What volume of a 0.8351-M iron(II) hydroxide solution is needed to neutralize 98.35 mL of 0.5417-M phosphoric acid?
5. Complete the table
pH pOH [H3O
+] [OH
-]
2.05
5.05
9.90 X 10-9
7.27 X 10-4
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