Madison City Schools PreAP Chemistry Updated 7/29/2015 Page 1 of 127 Table of Contents Chemistry 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
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Madison City Schools PreAP Chemistry Updated 7/29/2015
Page 1 of 127
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
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,
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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
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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
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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.
Madison City Schools PreAP Chemistry Updated 7/29/2015
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.