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Madison City Schools PreAP Chemistry Updated 7/29/2015

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Table of ContentsChemistry Safety Contract .............................................................................................................. 3

Safe Laboratory Techniques ...................................................................................................... 10 Lab Notebook Guidelines ............................................................................................................. 13 Lab Equipment and Safety ............................................................................................................ 15 The Great Gas Plot ........................................................................................................................ 16

How Sweet It Is! ........................................................................................................................... 19 Half-Life Simulation ..................................................................................................................... 22 Excited Elements .......................................................................................................................... 23 Molecule Construction .................................................................................................................. 25 Evaporation and Intermolecular Attractions ................................................................................. 26

Percent Composition of Hydrates ................................................................................................. 26 Determining an Empirical Formula .............................................................................................. 26 Types of Chemical Reactions: A Sampler Platter........................................................................ 26

Solubility in Double Replacement Reactions ............................................................................... 26 Whats for dinner? Leftovers. ....................................................................................................... 26 Stoichiometry of a Precipitate ....................................................................................................... 26 LabQuest-Determination of Wavelength of Maximum Absorbance ............................................ 26

LabQuest Spectrophotometric Analysis of Aspirin ...................................................................... 26 Thin-Layer Separation of Lipstick ................................................................................................ 26

Determination of Melting Points .................................................................................................. 26 Molar Volume of a Gas................................................................................................................. 26 Airbags .......................................................................................................................................... 26

“Wet” Dry Ice ............................................................................................................................... 26 Properties of Solutions: Electrolytes and Non-Electrolytes......................................................... 26

39 Drop pH Lab ............................................................................................................................ 26 Titration Lab ................................................................................................................................. 26

Qualitative Analysis of the Group I Cations ................................................................................. 26 Copper into Gold: The Alchemist’s Dream ................................................................................. 26

Freezing Point Depression of a Solution (Ice Cream) .................................................................. 26 Periodic Table ............................................................................................................................... 26 Rules of Writing Equations, Solubility Rules, Activity Series of Metals ..................................... 26

Polyatomic Ions ............................................................................................................................ 26 Apples ........................................................................................................................................... 26 Avocado ........................................................................................................................................ 26

Tomato .......................................................................................................................................... 26 Jalapeno......................................................................................................................................... 26 Banana........................................................................................................................................... 26 Firehouse ....................................................................................................................................... 26

KFC ............................................................................................................................................... 26 Marcos........................................................................................................................................... 26 Outback ......................................................................................................................................... 26

Pistachios ...................................................................................................................................... 26 Basil .............................................................................................................................................. 26 Oregano ......................................................................................................................................... 26 Steak .............................................................................................................................................. 26 Brisket ........................................................................................................................................... 26 Ribs ............................................................................................................................................... 26 Chicken ......................................................................................................................................... 26

Sausage ......................................................................................................................................... 26 Pork Butt ....................................................................................................................................... 26

BACON......................................................................................................................................... 26 Tofu ............................................................................................................................................... 26 Nutmeg .......................................................................................................................................... 26 Salami ........................................................................................................................................... 26 Reuben .......................................................................................................................................... 26

Quinoa ........................................................................................................................................... 26 Turkey ........................................................................................................................................... 26 Stuffing ......................................................................................................................................... 26 Cranberry Sauce ............................................................................................................................ 26 Green Bean Casserole ................................................................................................................... 26

Mashed Potatoes ........................................................................................................................... 26 Pumpkin Pie .................................................................................................................................. 26

Reeses Cup .................................................................................................................................... 26 Nerds ............................................................................................................................................. 26 Milk Duds ..................................................................................................................................... 26 Candy Corn ................................................................................................................................... 26

Milky Way .................................................................................................................................... 26 Vanilla Tootsie Roll ...................................................................................................................... 26

Dots ............................................................................................................................................... 26 Cream of tartar .............................................................................................................................. 26


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Chemistry Safety Contract PURPOSE Science is a hands-on laboratory class. You will be doing many laboratory activities which require the use of hazardous

chemicals. Safety in the science classroom is the #1 priority for students, teachers, and parents. To ensure a safe science

classroom, a list of rules has been developed and provided to you in this student safety contract. These rules must be followed at

all times. The acknowledgment sheet must be signed by both you and a parent or guardian before you can participate in the

laboratory. Any questions by the student or parent should be addressed to the teacher before this contract is signed. Lab activities

may be videotaped to encourage safe practices.

GENERAL RULES

1. Conduct yourself in a responsible manner at all times in the laboratory.

2. Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ask the

instructor before proceeding.

3. Never work alone. No student may work in the laboratory without an instructor present.

4. When first entering a science room, do not touch any equipment, chemicals, or other materials in the laboratory area until you

are instructed to do so.

5. Do not eat food, drink beverages, or chew gum in the laboratory. Do not use laboratory glassware as containers for food or

beverages.

6. Perform only those experiments authorized by the instructor. Never do anything in the laboratory that is not called for in the

laboratory procedures or by your instructor. Carefully follow all instructions, both written and oral. Unauthorized

experiments are prohibited.

7. Be prepared for your work in the laboratory. Read all procedures thoroughly before entering the laboratory.

8. Never fool around in the laboratory. Horseplay, practical jokes, and pranks are dangerous and prohibited.

9. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times. Bring only your laboratory

instructions, worksheets, and/or reports to the work area. Other materials (books, purses, backpacks, etc.) should be stored in

the classroom area.

10. Keep aisles clear. The chemical storage area is off limit to all students.

11. Know the locations and operating procedures of all safety equipment including the first aid kit, eyewash station, safety

shower, fire extinguisher, and fire blanket. Know where the fire alarm and the exits are located.

12. Always work in a well-ventilated area. Use the fume hood when working with volatile substances or poisonous vapors.

Never place your head into the fume hood.

13. Be alert and proceed with caution at all times in the laboratory. Notify the instructor immediately of any unsafe conditions

you observe.

14. Dispose of all chemical waste properly. Never mix chemicals in sink drains. Sinks are to be used only for water and those

solutions designated by the instructor. Solid chemicals, metals, matches, filter paper, and all other insoluble materials are to be

disposed of in the proper waste containers, not in the sink. Check the label of all waste containers twice before adding your

chemical waste to the container.

15. Labels and equipment instructions must be read carefully before use. Set up and use the prescribed apparatus as directed in

the laboratory instructions or by your instructor.

16. Keep hands away from face, eyes, mouth and body while using chemicals or preserved specimens. Wash your hands with

soap and water after performing all experiments. Clean all work surfaces and apparatus at the end of the experiment. Return

all equipment clean and in working order to the proper storage area.

17. Experiments must be personally monitored at all times. You will be assigned a laboratory station at which to work. Do not

wander around the room, distract other students, or interfere with the laboratory experiments of others.Students are never

permitted in the science storage rooms or preparation areas unless given specific permission by their instructor.

18. Know what to do if there is a fire drill during a laboratory period; containers must be closed, gas valves turned off, fume

hoods turned off, and any electrical equipment turned off.

19. If you have a medical condition (e.g., allergies, pregnancy, etc.), check with your physician prior to working in lab.

CLOTHING

20. Any time chemicals, heat, or glassware are used, students will wear laboratory goggles. There will be no exceptions to this

rule!

21. Goggles must be worn during all phases of the lab including set-up, cleanup and everything in between. If you have glasses,

goggles must be worn over them. Contact lenses may be worn in the laboratory ONLY with non-directly vented chemical

splash goggles. Certain solvent liquids and vapors may cause the contact to fuse to the eye. It is the student’s responsibility

to inform the teacher if he/she is wearing contacts during the lab.

22. Dress properly during a laboratory activity. Long hair, dangling jewelry, and loose or baggy clothing are a hazard in the

laboratory. Long hair must be tied back and dangling jewelry and loose or baggy clothing must be secured. Shoes must

completely cover the foot. No sandals allowed. It is recommended to bring an old pair of shoes to keep at school in the event

that sandals are inadvertently worn on a lab day.

23. Lab aprons or lab coats have been provided for your use and should be worn during laboratory activities.

ACCIDENTS AND INJURIES

24. Report any accident (spill, breakage, etc.) or injury (cut, burn, etc.) to the instructor immediately, no matter how trivial it may

appear.


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25. If you or your lab partner are hurt, immediately yell out “Code one, Code one” to get the instructor’s attention.

26. If a chemical splashes in your eye(s) or on your skin, immediately flush with running water from the eyewash station or

safety shower for at least 20 minutes. Notify the instructor immediately.

HANDLING CHEMICALS

27. All chemicals in the laboratory are to be considered dangerous. Do not touch, taste, or smell any chemicals unless

specifically instructed to do so. The proper technique for smelling chemical fumes will be demonstrated to you.

28. Check the label on chemical bottles twice before removing any of the contents. Take only as much chemical as you need.

29. Never return unused chemicals to their original containers.

30. When transferring reagents from one container to another, hold the containers away from your body.

31. Acids must be handled with extreme care. You will be shown the proper method for diluting strong acids. Always add acid to

water, swirl or stir the solution and be careful of the heat produced, particularly with sulfuric acid. If an acid is spilled on the

skin, first blot with a paper towel, then go to the sink and run water over the affected area. It’s best to remove as much acid as

possible before washing with water.

32. Never dispense flammable liquids anywhere near an open flame or source of heat.

33. Never remove chemicals or other materials from the laboratory area.

HANDLING GLASSWARE AND EQUIPMENT

34. Never handle broken glass with your bare hands. Use a brush and dustpan to clean up broken glass. Place broken or waste

glassware in the designated glass disposal container.

35. Fill wash bottles only with distilled water and use only as intended, e.g., rinsing glassware and equipment, or adding water to

a container.

36. When removing an electrical plug from its socket, grasp the plug, not the electrical cord. Hands must be completely dry

before touching an electrical switch, plug, or outlet.

37. Examine glassware before each use. Never use chipped or cracked glassware. Never use dirty glassware.

38. Report damaged electrical equipment immediately. Look for things such as frayed cords, exposed wires, and loose

connections. Do not use damaged electrical equipment.

39. Do not immerse hot glassware in cold water; it may shatter.

HEATING SUBSTANCES

40. Exercise extreme caution when using a gas burner. Take care that hair, clothing and hands are a safe distance from the flame

at all times. Do not put any substance into the flame unless specifically instructed to do so. Never reach over an exposed

flame. Light gas burners only as instructed by the teacher.

41. Never leave a lit burner unattended. Never leave anything that is being heated or is visibly reacting unattended. Always turn

the burner or hot plate off when not in use.

42. You will be instructed in the proper method of heating and boiling liquids in test tubes. Do not point the open end of a test

tube being heated at yourself or anyone else.

43. Heated metals and glass remain very hot for a long time. They should be set aside to cool and picked up with caution. Use

tongs or heat-protective gloves if necessary.

44. Never look into a container that is being heated.

45. Allow plenty of time for hot apparatus to cool before touching it.

46.Hot and cold glass have the same visual appearance. Determine if an object is hot by bringing the back of your hand close to it

prior to grasping it.


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Student Name (Print clearly)__________________________________________________

Safety Contract Acknowledgement (Please sign, tear out of the lab manual and return to your

teacher)

QUESTIONS List any medical conditions

Do you wear contact lenses? of which the teacher should be aware.

___ YES ___ NO

Are you color blind?

___ YES ___ NO

Do you have allergies?

___ YES ___ NO

If so, list specific allergies

AGREEMENT

I, ___________________________ , (student’s name) have read and agree to follow all of the safety rules set

forth in this contract. I realize that I must obey these rules to ensure my own safety, and that of my fellow

students and instructors. I will cooperate to the fullest extent with my instructor and fellow students to maintain

a safe lab environment. I will also closely follow the oral and written instructions provided by the instructor. I

am aware that any violation of this safety contract that results in unsafe conduct in the laboratory or

misbehavior on my part, may result in being removed from the laboratory, detention, receiving a failing grade,

and/or dismissal from the course.

_______________________________________ ____________________________

Student Signature Date

Dear Parent or Guardian:

We feel that you should be informed regarding the school’s effort to create and maintain a safe science

classroom/ laboratory environment. With the cooperation of the instructors, parents, and students, a safety

instruction program can eliminate, prevent, and correct possible hazards. You should be aware of the safety

instructions your son/daughter will receive before engaging in any laboratory work. Please read the list of safety

rules above. No student will be permitted to perform laboratory activities unless this contract is signed by both

the student and parent/guardian and is on file with the teacher. Your signature on this contract indicates that

you have read this Student Safety Contract, are aware of the measures taken to ensure the safety of your

son/daughter in the science laboratory, and will instruct your son/daughter to uphold his/her agreement to

follow these rules and procedures in the laboratory. Please be aware that parents will be notified in the event of

an accident in the lab. Generally, theses accidents are minor in nature but we want you to be aware of any

accidents in the lab.

_______________________________________ _______________________________________

Parent/Guardian Signature Date


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Laboratory Hazards

Students should be aware of possible hazards in the laboratory and take the appropriate safety

precautions. By doing so, the risks of working in the chemistry laboratory will be reduced. This

section addresses laboratory hazards, how to prevent accidents, and what to do if an accident

occurs.

Chemical Burns

A chemical burn occurs when the skin or a mucous membrane is damaged by contact with a

substance. Corrosive substances can cause severe burns. An irritant is a chemical that can

irritate the skin and membranes of the eyes, nose, throat, and lungs. Chemicals that are corrosive

or irritating must be treated with special care. Chemical burns can be severe, and permanent

damage to mucous membranes can occur despite the best efforts to rinse chemicals from an

affected area. The best defense against chemical burns is prevention.

Without exception, safety goggles must be worn during all phases of the laboratory period,

even during cleanup. Goggles should be put on as soon as you enter the laboratory and remain

over your eyes until you leave the laboratory. Should any chemical splash in your eye,

immediately notify your teacher. Use a continuous flow of running water to flush your eye for

20 minutes. Do not rub eye. Wear a laboratory coat or apron and shoes that cover the entire foot

and socks (no sandals) to protect your clothing, feet, and other areas of your body. If corrosive

chemicals come in contact with your skin, rinse the affected area with water for several minutes.

If there is no burning sensation, wash area with soap and water.

*Estimates for the time required for permanent corneal damage to occur following exposure to

1M NaOH are in the range of 30 seconds.

An additional burn hazard exists when concentrated acids or bases are mixed with water. The

heat released in mixing these chemicals with water can cause the mixture to boil, spattering

corrosive chemical. The heat can also cause regular glass containers to break, spilling the

corrosive chemical. To avoid these hazards, always add acid or base to water, very slowly and

with stirring, and never the reverse. * As a precaution, Pyrex or Kimax containers, glassware that

has been treated to withstand high temperatures, should always be used.

*Concentrated sulfuric acid causes thermal burns because it reacts with water in the skin

releasing substantial amounts of heat. Nitric acid does not produce thermal burns but denatures

the proteins in the skin destroying tissue. Nitric acid burns are very slow to heal.

Thermal Burns

A thermal burn can occur if you touch hot equipment or get too close to an open flame. You

should be aware that hot and cold glassware look the same. If a gas burner or hot plate is being

used, some of the equipment nearby may be hot. Hold your hand near an item to feel for heat

before touching it. Treat a thermal burn by immediately running cold water on the burned area.

Continue to apply the cold water until the pain is reduced. This usually takes several minutes. In

addition to reducing the pain, cooling the burned area also serves to speed the healing process.

Greases and oils should not be used on burns because they tend to trap heat. Medical assistance

should be sought for any serious burn. Notify your teacher immediately if you are burned.


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Cuts from Glass

Many cuts that occur in the laboratory are avoidable by following a few simple rules. You should

never use broken, cracked, or chipped glassware. If you should break a piece of glassware, do

not pick up the broken pieces with your hands. Use a brush or broom and dustpan to sweep up

the shards of glass. All broken glass should be placed in the box labeled for broken glass. You

should never place broken glass in a regular trashcan.

If you receive a minor cut, briefly allow the cut to bleed by squeezing the cut. Place the injured

area under cold running water, and notify your teacher. Serious cuts and deep puncture wounds

require immediate medical attention. Notify your teacher immediately. Control the bleeding by

applying pressure with the fingertips or by firmly pressing on the wound with a clean towel or

gauze.

Poisoning

Many of the chemicals used in the experiments in this manual are mildly to moderately toxic. To

prevent poisoning, never eat, chew gum or drink in the laboratory. Do not touch chemicals.

Never taste any chemical in lab. Keep your hands away from your face. Always wash your hands

with soap and water at the end of the lab. In this way you will prevent chemicals that might get

on your hands from reaching your mouth, nose or eyes.

In some cases, the detection of an odor is used to indicate that a chemical reaction has taken

place. It is important to note that many gases are toxic when inhaled. If you must detect an

odor, use your hand to gently fan some of the gas toward your nose. Take a small sniff of the

gas instead of a deep breath. This will minimize the amount of gas sampled.

Fire

A fire may occur if chemicals are mixed improperly or if flammable materials come too close to

a burner flame or hot plate. Use a hot plate as a heat source instead of a burner when flammable

chemicals are being used or produced. When using a hot plate or burner, prevent fires by tying

back long hair and loose fitting clothing.

If hair or clothing should catch fire, DO NOT RUN. Running fans a fire. Stop, drop to the floor,

and roll slowly to smother the flames. Shout for help. If another person is the victim, get a fire

blanket, located at the front of the lab, to smother the flames. If a shower is nearby, help the

victim to use it.

A fire in a container may be extinguished by smothering the flames with the fire blanket, a

notebook, or some other nonflammable object. In case of a fire on a laboratory workbench, turn

off all gas jets and unplug all appliances. Notify the teacher immediately. If a fire extinguisher

is needed, the teacher will call for it. To use a fire extinguisher, pull the ring, point the nozzle at

the base of the fire, and squeeze the handle. Use short bursts from the extinguisher, rather than

one continuous spray. Caution: Never direct the spray of a fire extinguisher into a person's

face. If a fire is not extinguished quickly, leave the laboratory. Crawl to the door if necessary to

avoid the smoke. Do not return to the laboratory until you are told it is safe.


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The fire extinguishers available in lab are ABC extinguishers. This designation means that the

extinguishers may be used on three types of fires. These types of fires are: a. paper and trash, b.

liquids or grease, and c. electrical. ABC fire extinguishers should not be used for flammable

solids. Sand is used to extinguish burning flammable solids.

Fire Warning

The signal for a fire drill or fire is the sound of the fire alarm. If the signal is given while in the

laboratory, students should turn off all gas jets and exit immediately.

Tornado Warning

An announcement over the intercom is the signal for a tornado drill or warning. Go to the area

indicated by your instructor. You should sit on the floor facing the wall and protect your head.

You should remain in this position until the announcement ending the drill or warning is made.

Always/Never Rules

Always

wear safety goggles

wear protective clothing and shoes

use proper techniques and procedures

discard wastes properly

know the location and use of safety

equipment

be alert, serious and responsible in lab

Never

eat or drink in the lab

clutter your work area

perform unauthorized experiments

enter the chemical storage area

take unnecessary risks

remove stock chemicals from the supply

area

Report any injury, accident, or chemical spill to the teacher immediately. Know the location of

the eyewash, fire blanket, fire extinguisher, and shower.

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Safe Laboratory Techniques

Pouring Liquids

Always wear safety goggles when handling chemicals.

Always read the label on a reagent bottle before using, and then read the label again. Never touch

chemicals with your hands.

Never return unused chemicals to their original containers. To avoid waste, pour small amounts of reagents

into small beakers and share with students around you.

Follow this procedure when pouring liquids.

1. Remove the lid.

2. Hold bottle with the label in the palm of your hand.

3. When pouring a liquid from a reagent bottle into a beaker or

funnel, the reagent should be poured slowly down a glass stirring

rod.

4. When pouring a liquid from a bottle into a test tube or graduated

cylinder, the empty container should be held at eye level. Pour

the liquid slowly until the correct volume is obtained.

5. Place the lid back on the bottle before removing the lid from

another reagent bottle.

Filtering a Mixture

To separate a solid from a liquid, the most common method is

gravity filtration.

1. Fold the filter paper in half and then quarters.

2. Open the folded paper to form a cone with one layer of paper on

one side and three layers on the other side.

3. Put the cone in a funnel. Moisten the filter paper with a small

amount of distilled water and gently press the paper against the

sides of the funnel.

4. Place a beaker beneath the funnel with the tip of the funnel just

touching the inside surface of the beaker about one inch below

the rim.

5. Use a stream of distilled water to wash the solid remaining in the

beaker into the funnel. Wash the solid in the filter with distilled

water to remove all traces of solvent. Dry the solid.

6. Pour the liquid down a glass stirring rod into the funnel. Keep

the liquid below the top edge of the paper at all times to prevent

overflow.

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Using a Gas Burner

Laboratory burners produce various kinds of flames when different

mixtures of gas and air are burned. The Tirrell burner has adjustable

air vents and a gas control valve in the base.

1. Examine a Tirrell burner and identify the parts.

2. Connect the burner to the gas supply with tubing.

3. Close the air vents. Close the gas control valve at the bottom of

the burner and then open both about 1 ½ full turns.

4. Hold a lighter at the top of the barrel of the burner and turn on

the gas supply at the lab station. With a Tirrell burner, the main

gas supply should be opened fully and the gas flow regulated by

the gas control valve at the base of the burner. The flame may be

yellow or a single blue.

5. Open the air vents slowly, to admit more air into the flame, to

produce a light blue cone-shaped flame. If the flame blows out

after lighting, turn off the gas supply, and relight. Continue to

open the air vents to produce a blue triple-cone flame.

6. Adjust the gas supply to produce the desired size of flame. For a

smaller flame, close the air vent slightly and reduce the gas

supply. Practice adjusting the flame.

7. Turn the burner off at the main gas supply valve when finished.

Caution: Tie back long hair and pull back loose clothing when working with a lab burner. Do not reach

across a flame. Do not use a burner around flammables. Never leave a burner flame unattended. Know the

location of fire extinguisher, the fire blanket, and safety shower.

Heating a Liquid in a Test Tube 1. Adjust the burner to give a small single blue flame.

2. Fill a test tube no more than one-third full of the liquid to be heated.

3. Hold the test tube with a test tube holder. The test tube holder should grip the

test tube near the mouth of the tube.

4. Place the test tube in a slanting position in the flame, gently heat the entire

length of the test tube and then heat the substance in the test tube a short

distance below the surface of the substance.

5. Gently shake the tube as it is heated until the substance melts, boils, or

reaches the desired temperature. Caution: Never point the open end of a test tube toward yourself or others.

Never heat the bottom of a test tube held in a vertical position.

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Heating a Liquid in a Beaker.

1. Place wire gauze on a hot plate.

2. Place a half-filled beaker of liquid on the wire gauze.

3. Turn on the hot plate and adjust the knob to 7.

4. Caution: Never heat plastic beakers or graduated glassware on a

hotplate. Never let a beaker boil dry; add water to the beaker as

necessary. Never adjust the hot plate to 10.

Measuring Mass Using an Electronic Balance

To find the mass of an object, follow these general rules.

1. Zero the balance, and place the object on the balance pan. If you are measuring out a chemical,

use a weigh boat. Never put chemicals directly on the pan. Be sure to zero or tare the balance

with the weighing boat on the electronic balance.

2. If a chemical is spilled on or near the balance, clean it up immediately. There are brushes

available to clean off solids. If in doubt, check with the teacher.

3. Never attempt to weigh an object with a mass greater than the maximum capacity of the balance.

Measuring Volume

Volume measurements are important in experimental procedures. Accurate laboratory measurements

are made using graduated cylinders, pipets, burets, or volumetric flasks. Although some beakers

have graduation marks, these marks are designed for quick, rough estimates of volume. Liquid

volumes are usually measured in milliliters.

Using a Graduated Cylinder

Place about 50 mL of water in a 100-mL graduated cylinder

and set the cylinder on the laboratory bench. Look at the

surface of the water. The surface curves upwards where the

water contacts the cylinder walls. This curved surface is

called a meniscus.

A volume measurement is always read at the bottom of the

meniscus with your eye at the same level as the surface of the

liquid. To make the meniscus more visible place a finger or a

dark piece of paper behind and just below the meniscus while

making the reading.

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Lab Notebook Guidelines

A lab notebook should be used to explain lab procedures, record all lab data, and show how

calculations are made. You may also use the notebook to discuss the results of an experiment and to

explain the theories involved.

A record of lab work is an important document which will show the quality of the lab work

that you have done. You may need to show your notebook and your lab reports to the Chemistry

Department at a college or university in order to obtain credit for the lab part of an AP Chemistry

class. As you record information in your notebook, keep in mind that someone who is unfamiliar

with your work may be using this notebook to evaluate your lab experience in chemistry. When you

explain your work, list your data, calculate values and answer questions, be sure that the meaning

will be obvious to anyone who reads your notebook.

Guidelines for the notebook:

1. Write your name and class on the front cover.

2. In black or blue ink, number all the right hand pages on the lower right corner if they are not

already numbered.

3. Save the first 2 pages for a Table of Contents. This should be kept current as you proceed. Each

time you write up a lab, place the title and page numbers where the lab report begins in the Table

of Contents.

4. Write in blue or black ink. Use only the right hand pages. The left hand pages can be used to do

calculations or other scratch work.

5. If you make a mistake, DO NOT ERASE OR SCRIBBLE. Just draw ONE LINE through your

error, and continue. It is expected that some errors will occur. A lab notebook is a working

document, not a perfect, error-free, polished product. Errors should be corrected by drawing one

line through the mistake, and then proceeding with the new data.

6. Do not use the first person or include personal comments.

Lab Reports (Lab reports will be worth 25 points)

Include the following information in your lab reports. Label each section

1. Title – The title should be descriptive. Experiment 5 is not a descriptive title.

2. Date – This is the date you performed the experiment.

3. Also record your lab partner and your lab station.

4. Purpose – A brief statement of what you are attempting to do in a complete sentence.

5. Procedure – A description of the method you are using which provides a summary of the

procedure. Do not include lengthy, detailed directions. A person who understands chemistry

should be able to read this section and know what you are doing.

6. Data-Record all your data directly into your lab notebook on the right-hand pages. Each person

must have all the data in their lab notebook and have it stamped by the teacher before you leave.

Organize your data in a neat, orderly form. Label all data very clearly. Use correct sig figs and

always include proper units. Underline, use capital letters or use any device you choose to help

organize this section well. Space things out – don’t try to cram everything on one page. Use

tables as much as possible. A data table must have a label and a title. e.g. – Table 1: Density

Values for Sugar Solutions.

7. Calculations and Graphs- You should show how calculations are carried out. Give the equation

used and show how your values are substituted into it. Give the calculated values. If graphs are

included, make the graphs an appropriate size. Label all axes and give each graph a title. If

experiments are not quantitative, this section may be omitted.

8. Conclusions – Make a simple statement concerning what you conclude from the experiment.

This is not a place to give your opinion of the lab and whether or not it was “fun”. It is not

your job to review the lab like you would if you saw a movie.

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9. Experimental error – If there is a known value for something you are doing in lab, calculate

the experimental error.

Accuracy

Accuracy is a measure of how close an experimental value is to a value which is

accepted as correct. The measure of the accuracy of an experimental value is reported as

Percent Error.

Accepted

AcceptedalExperimenterror

%

10. Error Analysis – What are some specific sources of error, and how do they influence the

data? Do they make the values obtained larger or smaller than they should be? Which

measurement was the least precise? Instrumental error and human error exist in all

experiments, and should not be mentioned as a source of error unless they cause a significant

fault. Significant digits and mistakes in calculations are NOT a valid source of error. In

writing this section it is sometimes helpful to ask yourself what you would do differently if

you were to repeat the experiment and wanted to obtain better precision.

11. Questions – Answer any questions included in the lab directions. Answer in such a way that

the meaning of the question is obvious from your answer.

Reporting Lab Data

Graphing Data

1. All graphs should have a descriptive title (“Graph” is not a title) and a label. e.g. – Graph A:

Density of Solutions with Varying Sugar Concentrations.

2. Both the vertical and horizontal axes should both have labels and units clearly marked. Use a

ruler to draw the axes.

3. The scales chosen should reflect the precision of the measurements. For example, if

temperature is known to be ±0.1ºC, you should be able to plot the value this closely. Don’t

have each block of the graph equal to 10ºC.

4. There should be a table in which the data values are listed. Don’t put data in a graph unless

you have first listed it in a table.

5. The controlled or independent variable is placed on the horizontal axis. The dependent

variable is graphed on the vertical axis.

6. There should be an obvious small point on the graph for each experimental value. It is not

necessary to include the coordinates of each point since they will be in the data table.

7. A smooth line should be drawn that lies as close as possible to most of the points. Do NOT

draw a line connecting one point to the next as in a dot-to-dot drawing. If the line is a

straight line, use a ruler to draw it.

8. If a computer program is used to draw the graph, the rules still apply.

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Lab Equipment and Safety

No writeup in lab notebook.

Discussion:

In this activity you will become familiar with some proper lab procedures used in the Chemistry lab.

Procedure (Write the answer to each question on a sheet of paper, not your lab notebook)

Measuring the volume of a liquid.

1. Read “Using a graduated cylinder”. See your TOC.

2. What type of graduated cylinder is at your lab station? (10 mL, 50 mL, 100 mL, etc)

3. How much does each gradation represent?

4. Add some water to your graduated cylinder. What volume of water is your graduated cylinder

holding? Write your answer to the tenths.

5. Put 50 mL of water in a beaker using the beaker as a measuring device. Now pour that amount

into a graduated cylinder. How do the two measurements compare?

Measuring mass with an electronic balance.

1. Read “Measuring mass with an electronic balance” in this lab manual. See your TOC.

2. Place a weigh boat on the balance. Press zero/tare to zero the balance.

3. Add 1.00 g of NaCl to the weigh boat. Discard the salt in the trash. Rinse and dry the weigh

boat for future use.

4. Obtain 3 evaporating dishes, 2 crucibles and 5 pieces of filter paper. Determine and record the

mass for each separate item. Return all the equipment to the drawers and the filter paper to the

teacher desk.

Using a Burner

While commonly called a Bunsen burner, we have Tirrell burners which are an improvement on the

original Bunsen burner.

1. Read the lab manual about using a burner.

2. Read over the fire and thermal burns safety considerations in this lab manual.

3. Follow the steps in this lab manual to light the burner.

4. Once you have a flame that is burning safely and steadily, you can experiment by completely

closing the air vents. What effect does this have on the flame?

5. Regulate the flow of gas so that the flame extends roughly 8 cm (2.54 cm = 1 inch) above the

burner tube. Now adjust the supply of air until you have a quiet, steady flame with a sharply

defined light blue inner cone. This adjustment gives the highest temperature possible with your

burner. Where is the hottest portion of the flame located? (Consult the poster on the wall in the

lab or classroom)

6. Shut off the gas burner at the gas valve.

7. Draw the burner and label the parts.

Cleanup

1. Pour all solutions down the drain

2. Clean your lab station.

3. Clean all equipment and leave to dry at your lab station.

4. Wash your hands

Questions:

1. Why is it important to use a graduated cylinder to measure liquids rather than a beaker?

2. Is it safe to assume that pieces of the same equipment have the same mass? Explain.

3. Why is the nonluminous flame preferred over the yellow luminous flame in the laboratory?

Page 16 of 127

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


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.

Page 25 of 127

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.


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Evaporation and Intermolecular Attractions


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


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.


Page 28 of 127

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


Page 29 of 127

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.


Page 30 of 127

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


Page 31 of 127

Percent Composition of Hydrates


Discussion:

Hydrates are ionic compounds (salts) that have a definite amount of water as a part of

their structure. This water of hydration is released as vapor when the hydrate is heated.

The remaining solid is known as an anhydrous salt. The general reaction for heating a

hydrate is:

∆

hydrate → anhydrous salt + water vapor

The percent of water in a hydrate can be found experimentally by accurately determining

the mass of the hydrate and the mass of the anhydrous salt. The difference in mass is due

to the water lost by the hydrate. The percent of water in the original hydrate can be

calculated by:

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


Page 33 of 127

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


Page 34 of 127

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?


Page 35 of 127

Types of Chemical Reactions: A Sampler Platter


Discussion:

There are many kinds of chemical reactions and several ways to classify them. One useful

method is to classify reactions into five major groups. These are (1) composition or synthesis; (2)

decomposition or analysis; (3) single replacement; (4) double replacement or exchange of ions;

and (5) combustion. Not all chemical reactions can be placed into one of these categories.

In a synthesis reaction, two or more substances (elements or compounds) combine to form a

more complex substance. Equations for synthesis reactions have the general form of

A + B AB. An example of this reaction is the formation of water from its constituent

elements: 2H2(g) + O2(g) 2H2O(l) .

A decomposition reaction is exactly the opposite of a synthesis reaction. In a decomposition

reaction, a compound breaks down into two or more simpler substances. The general form of the

equation for a decomposition reaction is AB A + B. The breakdown of water into its

elements is an example: 2H2O(l) 2H2(g) + O2(g).

In a single replacement reaction, one substance in a compound is replaced by another, more

active element. Equations for single replacement reactions have two general· forms. In reactions

in which one metal replaces another metal, the general equation is

A + BY A Y + B. In reactions in which a nonmetal replaces another nonmetal, the general

form is X + BY BX + Y. The following equations illustrate each of these types of single

replacement reactions:

Zinc replaces copper (II) ion: Zn(s) + CuSO4 (aq) ZnSO4 (aq) + Cu(s)

Chlorine replaces bromide ions: Cl2 (g) + 2 KBr(aq) 2 KCI(aq) + Br2 (g)

In a double replacement reaction, the metal ions of two different ionic compounds can be thought

of as trading places. The general equation, AB + CD AD + CB, represents this type of

reaction. Most single and double replacement reactions take place in aqueous solutions

containing free ions. In a double replacement reaction, one of the products formed must be a

precipitate, an insoluble gas, or a molecular compound.

In a combustion reaction an element or compound is reacted with oxygen, often producing

energy in the form of heat and light. If the reaction is complete combustion, the products are

energy, carbon dioxide and water. If the reaction is incomplete, in addition to energy, carbon

dioxide and water, carbon monoxide and carbon will be produced.


Page 36 of 127

SDS:

Ethanol Toxic by ingestion and inhalation. Body tissue irritant. Avoid all

body tissue contact. Denatured with isopropanol and methanol. Not

for human consumption. Flammable liquid.

CaCO3 Irritant to body tissues. Severe eye and moderate skin irritant.

KI solution Substance not considered hazardous. However, not all health


Pb(NO3)2 solution Moderately toxic by ingestion and inhalation. Possible carcinogen.

Irritating to eyes, skin and mucous membranes. Avoid all body

tissue contact. Chronic exposure to inorganic lead via inhalation or

ingestion can result in accumulation in and damage to the soft

tissues and bones.

Mg Substance not considered hazardous. However, not all health


Zn Substance not considered hazardous. However, not all health


Inhalation of zinc dust may cause lung irritations. Zinc dust can

spontaneously combust when in contact with moisture.

6M HCl Toxic by inhalation and ingestion. Severe corrosive to all body

tissues, especially skin and eyes. Avoid all body contact.

PROCEDURE:

For each of the following reactions, chemical changes should be observed and noted. The

reactants and products should be identified and described. Identify the type of reaction

that occurs. List what evidence is seen to indicate that a chemical reaction has occurred.

Write a balanced equation including states of matter for each reaction.

1. (Combustion) Place 6 drops of ethanol, C2H5OH, in an evaporating dish. Ignite with a

burning splint. Carefully observe what happens and note any changes in the reactant.

Describe the reaction and the color of the flame.

2. (Decomposition) Place 1 scoop of calcium carbonate, CaCO3 , in a large, clean, dry

test tube. Note the appearance of the salt.

3. Using a test tube holder, heat the CaCO3 strongly in the burner flame for about 3

minutes. Insert, but do not drop, a burning wood splint in the test tube to test for the

presence of carbon dioxide, CO2 , gas. Carbon dioxide will extinguish the flame.

Note any changes in the appearance of the solid in the test tube.

4. (Double Replacement) Put about 1 mL of 0.1M potassium iodide, KI, solution into a

small, clean test tube. Add 1 drop of 0.1 M lead (II) nitrate solution to the same test

tube. Observe what happens and note any changes in the mixture. CAUTION: Do

NOT get this all over your hands. Wash your hands after this experiment. Dispose of

the contents of this test tube in the toxic waste container.


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5. (Synthesis) Place a watch glass near the base of a burner. Examine a piece of

magnesium ribbon. Using crucible tongs, hold the magnesium ribbon in the burner

flame until the magnesium ignites. CAUTION: Do not look directly at the flame. As it

burns, the Mg will release UV light which can be harmful to your eyes. Hold the

burning magnesium away from you and directly over the watch glass. When the

ribbon stops burning put the remains in the watch glass. Examine the product

carefully and note the appearance.

6. (Single Replacment) Stand a small, clean, dry test tube in a test tube rack. Add about

1 mL of 6M hydrochloric acid, HCl (aq), to the tube. CAUTION: Handle acids with

care. Acids can cause painful burns. Carefully drop a small piece of zinc metal, Zn,

into the acid in the test tube. Observe and record what happens.

7. Using a second test tube holder, invert the mouth of a second test tube over the test

tube in which the reaction is taking place. Remove the inverted tube after about 30

seconds and quickly insert a burning wood splint into the mouth of the inverted tube.

A pop indicates the presence of hydrogen gas. Note the appearance of the substance

in the first test tube.

Clean Up:

1. Empty the tube with calcium carbonate into the trash can but DO NOT wash.

2. Empty the tube with the lead nitrate in it into the waste beaker on the back bench.

3. Dispose of the magnesium ash and any wooden splints in the garbage can.

4. Empty the tube with zinc in it into the sink away from the drain. Rinse off the piece

of zinc, pick it up with your hand and put it in the zinc waste container on the back

bench. Wash both test tubes with soap and rinse. Leave the test tubes upside down in

the test tube rack to dry.

QUESTIONS:

1. In this experiment, describe the method that was used to test for the presence of

CO2 gas.

2. Describe the test used to check for the presence of hydrogen gas.

3. A small quantity of mercury (II) oxide is placed in a clean, dry test tube and heated

for about 3 minutes. At the end of this time a glowing splint is placed in the mouth

of the test tube. The splint bursts into flames. Write a balanced equation for the

reaction giving the reactants and products. What type of reaction is this?


Page 38 of 127

Solubility in Double Replacement Reactions


Discussion:

In general, solubility can be thought of as the tendency of a substance (solute) to dissolve

in other substance (solvent). For qualitative purposes, such terms as "soluble",

“insoluble", and "slightly/soluble" can be used to describe these tendencies. The

solubility tables at the end of the lab manual show the tendencies of several ionic

compounds to dissolve in water.

Ionic compounds (salts and bases) dissolve in water by a process known as dissociation.

In this process, the crystal lattice of the solid breaks down and free ions move throughout

solution. The total number of positive charges equals the total number of negative ions in

an ionic solution. It is possible to write a different kind of chemical reaction called a net

ionic equation. It shows only the ions actually changing in a reaction and leaves out

those that do not change. The ones that do not change are called “spectator ions”. To

write a net ionic equation you separate anything that is (aq) and ionic or a strong acid into

separate ions. Anything that remains unchanged in the reaction can be marked out. For

example:

AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) is an equation for a precipitate forming.

Separate the (aq) that are ionic and you get:

Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) AgCl(s) + Na+(aq) +NO3

-(aq)

Cancel out the nitrate ions and sodium ions because they do not change at all and you’re

left with:

Ag+(aq)+ Cl-(aq) AgCl(s)

In the following double replacement reactions, when two different aqueous (water)

solutions are mixed, one of two things may be observed: a) No reaction will occur. If all

the ions remain free, there is no reaction, and the appearance of the mixture of the ions

will remain clear. b) A precipitate (solid) will form. If two oppositely charged ions are

attracted to one another strongly enough, they may bond together to form an ionic

compound that will not separate into ions in water. This insoluble solid is a precipitate.

In this experiment, aqueous solutions of several different ionic compounds will be used.

Different combinations of solutions will be mixed and the results observed and noted. In

those mixtures in which reactions occur, precipitates will form. Each reactant and

precipitate will be described. The formula for each precipitate will be written. For each

reaction, balanced equations and net ionic equations will be written, and the results will

compared to the results found on the solubility table. NOTE: Since all salts are soluble to

a degree and these concentrations are small there may be some deviation from the rules.

There are several types of precipitates that may be observed in this lab. Some precipitates

will appear to be swirled. A fine powdered precipitate may look hazy or slightly cloudy.

In some reactions the precipitates will form large clumps. Crystalline precipitates may

form when the reaction is allowed to sit for a short time. The crystals have straight edges

and reflect light. Another type of precipitate is gelatinous, and the precipitate is more gel-

like than solid in appearance.


Page 39 of 127

SDS:

0.1M AgNO3 Moderately toxic by ingestion. Irritating to body tissues. Avoid all

body tissue contact. Will stain clothing and skin.

0.1M BaCl2 Highly toxic by ingestion and inhalation. All soluble barium

compounds are poisonous if swallowed and cause nausea, vomiting,

stomach pains and diarrhea.

0.1M NaOH Slightly toxic by ingestion and skin absorption. Severely irritating to

body tissues. Causes severe eye burns. Avoid all body tissue contact.

0.1M CuSO4 Mildly toxic by ingestion. Irritant to skin, eyes and mucous

membranes. Avoid contact with body tissues.

0.1M CoCl2 Toxic by inhalation and injestion. Severe corrosive to all body tissue,

especially skin and eyes. Avoid contact with all body tissues.

0.1M K3PO4 Body tissue irritant.

0.1M Pb(NO3)2 Moderately toxic by ingestion and inhalation. Possible carcinogen.

Irritating to eyes, skin and mucous membranes. Avoid all body tissue

contact. Chronic exposure to inorganic lead via inhalation or ingestion

can result in accumulation in and damage to the soft tissues and bones.

0.1M Fe(NO3)3 Corrosive to body tissues by contact and inhalation. Avoid contact

with skin, eyes and mucous membranes.

0.1M BiCl3 Moderately toxic by ingestion and inhalation. Corrosive to body

tissues. Avoid body tissue contact.

0.1M Na2CO3 Irritating to body tissues.

0.1M MnCl2 This material is generally considered nonhazardous, however not all

health aspects of this substance have been fully investigated.

PROCEDURE:

1. Before the lab, write balanced equations with states of matter for either Set A or Set

B. Your lab partner will do the other set. If no precipitate is made, write NR for the

products. Write a net ionic equation for each equation that produces a precipitate.

2. At each workstation there will be a sheet of white paper with a grid and black ovals

on it in a protective sleeve. Use the white sheet of paper as a guide.

3. For the solutions in the boxes across the top of the grid, place 1 drop of each solution

in each black oval under it. Example: AgNO3 has 5 black ovals under it, so there

should be 1 drop of AgNO3 placed in each of those 5 ovals. BaCl2 has only one oval

under it. Repeat for all of the chemicals at the top.

4. For each solution in the boxes on the left side of the grid, place 1 drop of each

solution in the black ovals to the right of the boxes. Example: BaCl2 has 4 ovals 2

beside it, so there should be 1 drop of BaCl2 placed in each of those 4 ovals. Repeat

for all of the chemicals at the left. NOTE: Only 5 drops of each solution will be used.

5. On the data sheet describe each precipitate in terms of amount, color and texture. In

the same box under the description, write the formula for each precipitate. If there is

no precipitate, do not describe. Write N.R. in the box.

Cleanup:

1. Wipe the sheet with a paper towel, rinse it off and dry with another paper towel.


Page 40 of 127

Set A

AgNO3 BaCl2 NaOH CuSO4 CoCl2

BaCl2

NaOH

CuSO4

CoCl2

K3PO4

Set B

Pb(NO3)2 Fe(NO3)3 NaOH BiCl3 Na2CO3

Fe(NO3)3

NaOH

BiCl3

Na2CO3

MnCl2


Page 41 of 127

Whats for dinner? Leftovers.


Discussion:

In this activity, you will design an experiment to determine which reactant is the limiting

reactant, aluminum metal or copper(II) chloride solution and which is leftover. You will

calculate the experimental molarity of the copper(II) chloride solution and percent error

for this value given the theoretical value obtained from your teacher. You will then

calculate the percent yield for the solid copper produced.

SDS:

Copper (II) chloride Harmful if swallowed. Do not eat, drink or smoke when using this

product. Causes skin irritation.

Aluminum Flammable solid. Keep away from heat, sparks, open flames, and

hot surfaces..

Procedure

1. Watch the podcast “Aluminum Leftovers” of possible steps to include in your

experiment. Be sure to write down pertinent information. With a lab partner, design a

procedure that will allow you calculate the limiting reactant. Write steps for

conducting your experiment in the procedure section of your lab writeup. Obtain

teacher approval before beginning.

2. Write the balanced chemical equation for this reaction including state symbols for

each reactant and product.

3. Was the reaction endothermic or exothermic? Cite evidence from the experiment to

support your claim. Write the energy term on the appropriate side of the equation

written in step 2.

4. Refer to the balanced equation written in step 2 and the observations made in the

experiment to complete the pictorial representation in Figure 1 for the “final” beaker.

Copy this figure into your data section of your lab notebook.

5. When all of the liquid has passed through the filter, carefully remove the paper from

the funnel; spread it out on a hotplate set on medium to dry. When you think it is dry,

mass the filter paper with its contents, put it back on the hotplate for one more

minutes, remass the filter paper and its contents to ensure all the water has

evaporated. Be sure to record all data.

6. Make a statement in the data section indicating the limiting reactant? What evidence

do you have to support this claim?

7. Calculate the experimental molarity of the copper(II) chloride solution your teacher


Page 42 of 127

provided. Molarity is the number of moles of copper (II) chloride divided by the

Liters of solution you used.

8. Obtain the theoretical molarity of the copper(II) chloride solution from your teacher

and calculate the percent error.

9. Calculate the theoretical mass of copper.

10. Calculate the percent yield for the copper produced in your experiment.

Cleanup:

Wash all dishes and equipment with soapy water and rinse well. Refill the wash

bottle if needed for the next class. Leave the equipment at your lab station on a paper

towel for the next class.

Questions

1. What type of results would you expect to see if you had placed solid copper into a

solution of aluminum chloride? Explain your reasoning.

2. A student forgets to dry the piece of aluminum before measuring the final mass.

Explain how this error will affect the calculated molarity of the copper(II) chloride

solution. Clearly state whether the calculated molarity will be too large, too small, or

unchanged by this error and include mathematical justification for your answer.

3. The copper obtained in an experiment was allowed to dry for several days. Before

weighing the sample, the student noticed that some of the copper flakes appeared a bit

green. Will this color change increase, decrease, or not change the percent yield of

copper produced in the experiment? Clearly justify your claim with appropriate

evidence.


Page 43 of 127

Stoichiometry of a Precipitate


Discussion:

Given a balanced chemical equation and the mass of one of the substances in the

equation, the theoretical masses of all other substances in the equation can be calculated.

Calculations in which a known mass is used to find an unknown mass are called mass-

mass calculations. In this experiment, a double replacement reaction will occur when two

aqueous solutions are mixed. There are two products for this reaction: an insoluble salt,

which precipitates out of solution, and a soluble salt, which remains in solution. The

insoluble solid will be separated, by filtration, from the mixture, dried, and massed. As

you filter the solution, the liquid that goes through the filter is called the filtrate. The

value of the experimentally measured mass of product will be compared to the theoretical

mass predicted by a mass-mass calculation.

Safety Precautions:

SDS

CaCl2 Slightly toxic by ingestion. Mild irritant to skin, eyes and mucous membranes.

Avoid all body tissue contact.

Na2CO3 Slightly toxic by ingestion. Irritating to body tissues. Avoid all body tissue

contact.

1 M HCl Slightly toxic by inhalation and ingestion. Severe body tissue irritant.

Corrosive to eyes. Avoid all body contact.

PROCEDURE:

1. Using the balance, measure out approximately ___ g of _______. Your teacher will

tell you what numbers and chemicals belong in the blanks.

2. Record the exact mass measured in the data table.

3. Place the salt in a clean 250 mL beaker and add 30 mL of distilled water. Stir

thoroughly to be sure all crystals dissolve. Rinse the stirring rod. Describe solution.

4. Measure out exactly ___g of _________. Record the mass.

5. Place the second salt in a clean 150 mL beaker and add 30 mL of distilled water and

stir to dissolve. Describe solution.

6. Pour the second salt solution into the 250 mL beaker containing the first solution.

Record observations. Describe each product. Rinse the 150 mL beaker twice with

distilled water using a wash bottle. Pour each rinsing into the 250 mL beaker. Wash

the 150 mL beaker with soap and rinse with tap water and then distilled water.

7. Write your initials on a piece of filter paper in pencil, find the mass, and record this

mass. Fold the filter paper and place in the funnel. (See page 12 for instructions on

filtering)

8. Place empty 150 mL beaker under funnel, and pour the mixture from the 250 mL

beaker into the filter paper. Pour slowly and do not allow the liquid to rise above the

edge of the filter paper.

9. Rinse the beaker with about 5 mL portions of filtrate. Pour the rinse back through the

filter until all precipitate is transferred from the beaker to filter paper.

10. Wash the precipitate by pouring about 10 mL of distilled water through the filter.

This should remove dissolved salts from the filter paper.

11. Calculate theoretical mass of product. Your teacher should initial the calculated mass


Page 44 of 127

before you leave the lab.

12. Record your observations of the precipitate. Carefully remove the filter paper and

precipitate from the funnel and place on a lab cart to dry until the next day.

13. The next day, find the mass of the dry precipitate and filter paper. Record in data

table.

14. Determine the limiting reagent and the excess reagent. Calculate the mass of the

remaining excess reagent. Show your calculation.

15. Calculate the percent yield of product. Show your calculation.

16. Calculate your percent error. Show your calculation.

Cleanup:

1. Wash the all glassware with soap and water

2. Rinse glassware in the washpan under the fume hood with 1M HCl.

3. Rinse with tap water and leave at your station.

Questions:

1. What would happen to the percent yield if there was still water in the precipitate?

Justify your answer.

2. If there was some precipitate left in the beaker, what effect would it have on the

percent yield? Justify your answer.


Page 45 of 127

LabQuest-Determination of Wavelength of Maximum Absorbance


Discussion:

Colored solutions are colored because they absorb certain wavelengths of light while

allowing other wavelengths of light to pass through. As observers, we see the

wavelengths of light that are not absorbed. Each solution will have one or more

wavelength at which it absorbs strongly. The wavelength of light at which a colored

solution absorbs the most is called its “wavelength of maximum absorbance.” Light at

this wavelength can be used to analyze the solution in spectrophotometric analysis.

Getting Started: Assemble the LabQuest/Spectrometer system. Make sure to use the

AC power adapter to power the LabQuest unit. The spectrometer connects to the

LabQuest unit with a USB cable. Once assembled turn on the LabQuest unit and allow it

to go thru its boot process. Once finished you should see a meter screen with an orange

background with “USB:Abs” listed as the unit of measurement.

Calibration: Use the stylus to tap “Sensors”, “Calibrate”, and “USB Spectrometer.”

The “Calibrate Spectrometer” window will open. The lamp in the spectrometer will

warm for 60 seconds. Once prompted place a cuvette filled half full with tap water into

the sample holder on the spectrometer. Your teacher will describe the proper alignment

for the cuvette. Use the stylus to tap the “Finish Calibration” button. When you get the

“Calibration Completed” message, tap OK to return to the meter window.

Procedure:

1. Use a disposable pipet to fill a plastic cuvette approximately half full with each of

the colored ionic solutions. Record the color and concentration of each solution on

the data sheet. Make sure to wipe any excess liquid off the exterior of the cuvette

with a paper towel.

2. Insert the first cuvette into the sample holder on the spectrometer. Press the

“Collect/Stop” button with your finger. It looks like a “PLAY” button on a DVD

player. Once the spectrum is acquired, press the “Collect/Stop” button again. Be

Patient! It takes a moment for the process to stop. Do not press the button twice. The

graph will autozero and the wavelength of maximum absorbance will be displayed.

Record the absorbance value, the wavelength, and the color of the wavelength of

max absorbance on the data sheet.

3. Remove the cuvette containing the first sample from the spectrometer and insert the

cuvette containing the second sample. Press the “Collect/Stop” button. You will be

asked whether you wish to store or discard the previous sample data. Use the stylus

to select “Discard.” Once the spectrum is acquired, press the “Collect/Stop” button

again. The graph will autozero and the wavelength of maximum absorbance will be

displayed. Record the absorbance value, the wavelength, and the color of the

wavelength of max absorbance on the data sheet.

4. Repeat step 3 for the remaining samples.

5. Predict the wavelength of max absorbance of the yellow pH=7.00 buffer solution.

Measure the wavelength to check your prediction.


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6. Predict the wavelength of max absorbance of the reddish pH=4.00 buffer solution.

Measure the wavelength to check your prediction.

Cleanup:

Rinse out all cuvettes and leave on them a paper towel at your station to dry. Your

teacher will give you instructions on using the labquest.

Questions:

1. If the wavelength of max absorbance appears in the white area to the right of the

color spectrum on the graph window what does this tell us about the type of

electromagnetic radiation best absorbed by the sample?

2. Which solution showed the greatest analytical sensitivity? Hint: Divide the

absorbance by the concentration.

3. Why does greenish NiCl2 have a similar wavelength of max absorbance to that of

yellowish FeCl3? Hint: What two primary colors combine to make green?


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LabQuest Spectrophotometric Analysis of Aspirin


Discussion: Aspirin is the common name for acetylsalicylic acid. Acetylsalicylic acid has the

molecular formula C9H8O4. Aspirin reacts with sodium hydroxide according to the

following equation:

C9H8O4 + 3NaOH (aq) Na2C7H4O3 (aq) + NaC2H3O2 (aq) +

2H2O Aspirin sodium hydroxide sodium salicylate sodium acetate

water

If an acidified solution of iron (III) nitrate is added the following reaction takes place:

Na2C7H4O3 (aq) + Fe(NO3)3 (aq) FeC7H4O3+ (aq) + 2Na+ (aq)

+ 3NO3- (aq)

sodium salicylate iron (III) nitrate salicylatroiron (III) complex

The salicylatroiron (III) complex is a purple colored substance that has a max absorbance

at or near 536nm (in the blue-green region of the visible light spectrum). In this

experiment we will take an off the shelf pain reliever and measure its aspirin content

using a spectrometer.

Getting Started: Assemble the LabQuest/Spectrometer system. Make sure to use the

AC power adapter to power the LabQuest unit. The spectrometer connects to the

LabQuest unit with a USB cable. Once assembled turn on the LabQuest unit and allow it

to go through its boot process. Once finished you should see a meter screen with an

orange background with “USB:Abs” listed as the unit of measurement.

Spectrometer Calibration: Use the stylus to tap “Sensors”, “Calibrate”, and “USB

Spectrometer.” The “Calibrate Spectrometer” window will open. The lamp in the

spectrometer will warm for 60 seconds. Once prompted place a cuvette filled half full

with tap water into the sample holder on the spectrometer. Your teacher will describe the

proper alignment for the cuvette. Use the stylus to tap the “Finish Calibration” button.

When you get the “Calibration Completed” message, tap OK to return to the meter

window.

Preparation of a Calibration Curve:

1. Use a disposable pipet to fill a plastic cuvette approximately half full with each of the

A-E calibration standard solutions. Make sure to wipe any excess liquid off the

exterior of the cuvette with a paper towel or kimwipe.

2. Insert the “A” cuvette into the sample holder on the spectrometer. Press the

“Collect/Stop” button with your finger. It looks like a “PLAY” button on a DVD

player. Once the spectrum is acquired, press the “Collect/Stop” button again. Be

Patient! It takes a moment for the process to stop. Do not press the button twice. The

graph will autozero and the absorbance and the wavelength of maximum absorbance


Page 48 of 127

will be displayed. Record the max absorbance wavelength and the absorbance at the

max absorbance wavelength on the data sheet. Use the left/right arrows to move the

data circle to 536nm and record the absorbance at this wavelength.

3. Remove the cuvette containing the “A” standard from the spectrometer and insert the

cuvette containing the “B” standard. Press the “Collect/Stop” button. You will be

asked whether you wish to store or discard the previous sample data. Use the stylus

to select “Discard.” Once the spectrum is acquired, press the “Collect/Stop” button

again. The graph will autozero and the wavelength of maximum absorbance will be

displayed. Record the max absorbance wavelength and the absorbance at the max

absorbance wavelength on the data sheet. Use the left/right arrows to move the data

circle to 536nm and record the absorbance at this wavelength.

4. Repeat step 3 for the remaining calibration standards.

5. Make a graph of absorbance versus concentration. Your teacher may have you

prepare two graphs, one using the absorbance at the wavelength of max absorbance

and another using the absorbance at 536nm.

Sample Preparation: Select one of the three pre-digested aspirin samples and pipet

2.0mL of the solution to a clean 50mL volumetric flask and dilute to the mark with the

acidified iron (III) solution. Mix thoroughly.

Sample Analysis:

1. Half fill a cuvette with the aspirin sample (in acidified iron (III) solution) prepared

earlier. Use a paper towel or kimwipe to wipe any excess liquid from the outside of

the cuvette. Insert the cuvette into the spectrometer and read and record the data

exactly as you did with the A-E calibration standard solutions.

2. Use the graph to find the concentration of aspirin in the diluted sample. Record this

value on the data sheet.

3. Calculate the mg of aspirin in the tablet and compare this value to the value given on

the manufacturer’s label.

Brand Name ___________________ Aspirin Content, mg (from label)

_______________mg

Table 1: Data for Calibration Curve

Solution Concentration

(mg/mL)

A 0.080

B 0.064

C 0.048

D 0.032

E 0.016

Calculation of mg aspirin in the commercial product:


Page 49 of 127

________mg/mL X 50mL / 2.0mL X 250mL =

____________mg

(from graph) (dilution factor) (sample volume)

Calculate % difference between experimental value and value from manufacturer’s

label (% error)


Page 50 of 127

Thin-Layer Separation of Lipstick


Discussion:

In thin layer chromatography (TLC), the stationary phase is a thin layer of adsorbent

particles attached to a solid plate. The ones we use have silica gel attached to plastic

sheets. When handling the sheets it is best not to touch the silica gel surface. A small

amount of sample is applied (spotted) near the bottom of the plate and the plate is placed

in the mobile phase (solvent). This solvent is drawn up by capillary action. Separation

occurs as each component, being different in chemical and physical composition,

interacts with the stationary and mobile phases to a different degree creating the

individual bands on the plate. The retention factor. Rf value, is used to characterize and

compare components of various samples. This value is calculated in the following way:

Rf value = distance from origin to component spot

distance from origin to solvent front.

In this activity, you will determine how many pigments are contained in a lipstick sample

as well as whether or not different lipstick samples contain any of the same pigments.

You will also look at the resulting chromatogram under UV light in order to identify any

fluorescent components.

SDS

Solvent A

65 mL toluene: Highly flammable liquid and vapor. Keep away from heat, sparks,

open flames, and hot surfaces. Harmful if swallowed. May be fatal if swallowed and

enters airways. Causes skin and eye irritation. May cause drowsiness or dizziness.

Avoid breathing mist, vapors or spray

30 mL methanol Highly flammable liquid and vapor. Keep away from

heat, sparks, open flames, and hot surfaces. Toxic if swallowed or in contact with skin

+H311). Do not eat, drink or smoke when using this product (P270).

Hazard class: Skin and serious eye damage, corrosion or irritation (Category 2, 2A).

Causes skin and serious

eye irritation (H315+H319).

Hazard class: Specific target organ toxicity, single exposure (Category 1). Causes

damage to organs (H370). Do

not breathe mist, vapors or spray

5 mL glacial acetic

Solvent B

92 mL butanol

8 mL 2 M HCl

Procedure:

1. Touching the TLC strip as little as possible and only at the edge, obtain two (2) TLC

strips (4 x 10 cm each) and 2 or 3 lipstick samples that you would like to test.

2. Use a pencil to lightly draw a line approximately cm from the bottom edge of the TLC

strip. Mark the line with an “x” for each lipstick sample you will test and identify the

different samples in some way at the top of the TLC strip.


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3. Use a toothpick to make a small dot in the center of each “x” with the proper lipstick

sample. A small sample size (about cm in diameter and not too thick) will result in

better separation of pigments.

4. Repeat steps 2 and 3 using the same lipstick samples so that you have two identical

TLC strips to be tested.

5. Put a very small amount of solvent A into one 100 mL beaker and a very small amount

of Solvent B in another 100 mL beaker. The solvent level must be below the point of

application!!! Mark these beakers in some way so that you know which solvent is

which.

6. Handling the TLC strips by the edges only, place one strip in each solvent and cover

each with a 250 mL beaker. Be sure that the sample spots are at the bottom of the TLC

strip when it is place in the beaker, but also be sure that the sample spots are not below

the solvent level in the beaker.

7. Check the progress of the solvent as it wets the TLC strip and moves toward the top.

Do not let the solvent reach the top!

8, When the solvent is close to the top of the TLC strip, remove the chromatogram and

place on a paper towel to dry.

9. Immediately mark the solvent front and the center of each pigment spot with a pencil

line.

10. Make a dot with your pencil in the center of each spot. If the spot is irreuular or

bandshaped, use your judgement as to where the approximate center might be if it

were a round spot.

11. Measure the distance from the point of origin to the solvent front and record in the

Data Table. Be sure that collected data is recorded in the appropriate Data Table based

on the solvent used (A or B) as well as the lipstick tested.

12. Calculate the R value for each spot by dividing the spot distance (origin to spot) by

the solvent distance (origin to solvent front). Record properly in your Data Table.

13. Identify any pigments that appear to be the same in the different lipstick samples.


Page 52 of 127

Determination of Melting Points


Discussion: The melting point of a compound is the temperature at which it changes

from a solid to a liquid. This is a physical property often used to identify compounds or to

check the purity of the compound. It is difficult, though, to find a melting point. Usually,

chemists can only obtain a melting range of 2-3°C accuracy. This is usually sufficient for

most uses of the melting point. The purpose of this experiment is to become familiar

with equipment used to find the melting points of organic compounds and to identify

unknown compounds based on their melting point.

Procedure:

1. Obtain a capillary tube and one of the labeled organic compounds. Each person in

the group should do a different sample.

2. Push the open end of the capillary tube into the solid until the solid has moved about

0.5 cm into the capillary tube. Invert the tube and tap the closed end on the table to

pack the solid into the closed end of the tube.

3. Place the capillary tube into the Mel-Temp chamber as instructed by your teacher.

Three capillary tubes will fit into this chamber, so up to three members of the group

can run their samples at the same time.

4. Turn on the Mel-Temp and ensure the thermometer is inserted properly. Start with a

setting of one to two. Observation of the sample chamber and the thermometer

should be made frequently in order not to miss the melting point of the sample. Heat

slowly, so that accurate results are achieved. Use small incremental increases in the

setting when needed.

5. Watch the temperature on the thermometer and solid inside the capillary tube. As the

melting range of the solid approaches, you will see the solid begin to sweat. This is

the low end of the melting range. The temperature at which the solid has completely

turned to liquid is the high end of the melting range. Record these values.

6. Turn off the Mel-Temp to allow it to cool before trying your unknowns. It should

take about 15 minutes for the temperature to cool down.

7. Once you are comfortable with the procedure for determining the melting point of a

solid, prepare a sample of your unknown in the same way that you prepared the

known solids and find its melting range.

8. From the melting range you found, identify your unknown solid using the data from

the knowns you tested initially.

Questions:

1. Define the "melting point" of a substance.

2. What is the purpose of determining melting points?

3. Why is this method not used for finding the melting points of inorganic compounds?

4. What would you expect the melting point to be if you mixed two substances and took

the melting point of the mixture?

5. How could the rate of heating influence the melting point?


Page 53 of 127

Molar Volume of a Gas


Discussion

The basis of this experiment is the reaction in which a known mass of magnesium will be

reacted with excess hydrochloric acid

Hydrogen gas is the product that is of interest in this experiment. An experimental

determination of the number of moles of hydrogen molecules produced and the volume

occupied by these molecules will be made. The number of moles of hydrogen will be

calculated. The volume of hydrogen gas produced will be measured directly on the scale

of a gas measuring tube. Dalton's Law will be used to correct for the water vapor

pressure, and the ideal gas law will be used to correct this volume, measured under

laboratory conditions, to the volume the sample of gas would occupy at STP. The

corrected experimental volume of the hydrogen gas will be compared to the theoretical

volume of hydrogen gas.

SDS

Mg Substance not considered hazardous. However, not all health aspects of this

substance have been thoroughly investigated.

6M HCl Toxic by inhalation and ingestion. Severe corrosive to all body tissues,

especially skin and eyes. Avoid all body contact.

PROCEDURE:

1. Obtain a piece of magnesium ribbon, with its mass, from the teacher. Record

mass of magnesium. OR your teacher may have you cut a piece of magnesium

ribbon 3.2 cm long.

2. Use some steel wool or sand paper to clean the surface of the magnesium ribbon.

3. Obtain a piece of cotton thread about 15 cm long. Tie one end of the thread

around the piece of magnesium ribbon, leaving about 10 cm of thread free. Bend

the piece of magnesium ribbon so the magnesium will fit easily into the gas

measuring tube.

4. Place about 900 mL of room-temperature tap water in a 1 L beaker. Set up a ring

stand and double buret clamp, and place beaker in position.

5. Obtain about 10 mL of 3M or 6M HCl (aq). CAUTION: Handle this acid with

care. Carefully pour the acid into a gas measuring tube.

6. Tilt the gas measuring tube slightly. Using a beaker slowly fill the tube to the top

with room-temperature tap water. Try to avoid mixing the acid and water as much

as possible.

7. Lower the piece of magnesium ribbon 4 or 5 cm into the gas measuring tube.

Drape the thread over the edge of the tube and insert a one-hole stopper. Tube

must be completely filled with water.

8. Place finger over the hole in the rubber stopper and invert the gas measuring tube.

Lower the stoppered end of the tube into the beaker of water. Clamp the tube in

place so that the stoppered end is well under water.

9. Describe the reaction.


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10. Let the apparatus stand about 5 minutes after the magnesium has completely

reacted. Then, tap the sides of the gas measuring tube to dislodge any bubbles that

may have become attached to the side of the tube.

11. Move the tube up or down, to equalize pressure, until the water level in the tube is

the same as that in the beaker. On the scale of the gas measuring tube, read the

volume of the gas in the tube. Record the volume of gas.

12. Record the room temperature and barometric pressure.

13. Write a balanced equation for the reaction that occurred.

14. Calculate stoichiometrically how many moles of hydrogen would be made from

the mass of the magnesium. Show your work

15. Calculate the corrected pressure of the gas inside of the tube by subtracting out

the partial pressure of the water. Show your work.

16. Calculate the moles of gas inside the tube. Show your work.

17. Calculate a percent error for moles of hydrogen.

Cleanup:

1. Rinse the gas collecting tube with tap water and then rinse with distilled water

from a wash bottle. Hang upside down in your double buret clamp to dry

2. Wash all beakers and Erlenmeyers with soap and rinse and place on a paper towel

at your station to dry.

Questions:

1. Why was it necessary to clean the surface of the magnesium?

2. For the following possible errors, give the direction they would cause your results

to deviate from the accepted value of 22.4 L/mol at STP. Explain your answers.

a. air bubbles were introduced when tube is turned over in step 7

b. not all the Mg ribbon reacted

c. not all the MgO coating is removed from the Mg ribbon before beginning

the experiment

d. you read the volume when the water level in the tube was higher than

outside the tube in step 9

3. Find the volume of 80 g O2 at the conditions in the lab today.

4. How many liters would 0.25 mol of any gas occupy at pressure in the room but 10

degrees higher?


Page 55 of 127

Airbags

Discussion;

Airbags were mandated as a piece of safety equipment for all vehicles in 1998. All cars

and trucks are required to have airbags on both driver and passenger sides. Many vehicles

today include side passenger airbags. To date, the National Highway Traffic Safety

Administration, NHSA, calculates that using a seat belt and having an airbag reduces the

risk of death by 61 percent.

Automobile airbags are simply gas-inflated cushions. The gas that fills the bag is the

result of a series of rapid chemical reactions triggered by an impact. In this activity, you

will simulate the process of generating the gas using a different but similar chemical

reaction inside a resealable plastic bag.

This experiment connects stoichiometry and gas laws. Your mission is to generate

enough gas to just fill a small, resealable plastic bag using baking soda and 6.0 M

hydrochloric acid, HCl. The ideal “airbag” will be filled to plumpness yet not

overinflated or underinflated. The bag may contain unreacted chemicals and other

products of the reaction.

You will be asked to describe the method you develop to solve the problem. You must

complete this assignment (including the report) during the assigned period. You may not

share information between groups.

You and your partner will be graded on the reproducibility of the experimental design,

data collection skills, and the accuracy of your results. Your written responses must be as

clear and as concise as possible in communicating your thought processes. You must

follow proper safety procedures. PROCEDURE

1. Write a balanced chemical equation for the reaction that will produce the gas in your

simulated airbag before beginning your procedure.

2. With your partner, design a plan for inflating your airbag and record your plan on your

student answer page. Have your teacher initial your work before beginning.

3. Carry out your plan, and record your observations.

4. Revise your design. Repeat, and show your teacher the filled bag.

QUESTIONS

1. Using evidence from your experiment, go back to your balanced chemical equation

and include energy and state symbols.

2. What was the limiting reagent in your experiment? Justify your answer with data.

3. Sketch the system of the inflated plastic bag after the reaction has finished and label all

of the species present. Be sure to include the state symbol along with charges for any

ions.

4. The reaction in this activity is only a simulation for that in a real automobile airbag.

Write the series of balanced chemical equations for the reaction that actually takes

place in an automobile airbag. You may need to do some research, and many online

resources are available.

5. The volume for the driver airbag has remained an average of 56 L over the years.

Consider only the first step in the reaction written in Question 4 and calculate the mass

of the reactant needed to produce this volume of gas. Assume standard temperature

and pressure.


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“Wet” Dry Ice

Do Not Write up in your lab notebook.

Discussion How can ice be “dry”? Dry ice is the common name for compressed carbon dioxide, CO2.

The term “dry” refers to the fact that no liquid is left behind when a sample of dry ice is left

out at room temperature. The solid carbon dioxide undergoes a phase change directly from a

solid to a gas; it sublimes.

Refer to the phase diagrams of water (Figure 1) and carbon dioxide (Figure 2). Both graphs

contain a point at a specific temperature and pressure where all three phases exist in

equilibrium. This coordinate on the graph is referred to as the triple point of the substance.

You will also notice that both graphs indicate a critical point. At temperatures above the

critical temperature, it is not possible to condense the gas into liquid state even with an

increase in pressure. The phase diagram for carbon dioxide shows that CO2 can exist only

as a gas at ordinary room temperature and pressure. To observe the transition of solid

CO2 to liquid CO2, you must increase the pressure until it is at or above the triple point

pressure.

Safety Precautions:

Do not touch the dry ice with your hands. Dry ice is very cold(-78 °C) and can cause

tissue damage to your skin. Wear goggles at all times.

Prelab instructions: Answer these questions on a separate sheet of paper.

1. Draw Lewis structures for water and for carbon dioxide.

2. Refer to Figure 1 and Figure 2 in the discussion.

a. Compare the scales of the y-axes for the two phase diagrams. Give the

approximate range of pressures for each diagram, including units.

b. Consider samples at standard pressure, 1 atm, and room temperature, 20°C.

Determine the state of matter in which each sample of water and carbon dioxide

would exist.

c. State a set of conditions under which both water and carbon dioxide would be

solid.


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d. Compare the solid-liquid equilibrium lines for water and carbon dioxide. What

does the slope of this line indicate about the density of the solid?

3. Find the triple point pressure of CO2 online and fill in the ? in figure 2 above.

Procedure:

1. Use forceps or tongs to place 2-3 very small pieces of dry ice on the lab bench, and

observe them until they have completely sublimed.

2. Fill a plastic cup with tap water to a depth of 4-5 cm.

3. Cut the tapered end (tip) off the graduated pipet.

4. Use forceps to carefully slide 8-10 pieces of dry ice down the stem and into the bulb

of the pipet.

5. Use a pair of pliers to clamp the opening of the pipet stem securely shut so that no gas

can escape. Use the pliers to hold the tube and to lower the pipet into the cup just

until the bulb is submerged. From the side of the cup, observe the behavior of the dry

ice.

6. As soon as the dry ice has begun to melt, quickly loosen the pliers while still holding

the bulb in the water (note: if you wait too long to loosen the pliers. the pipet may

burst due to increasing pressure). Observe the CO2

7. Tighten the pliers again and observe.

8. Repeat steps 6 and 7 as many times as possible.

Part 2:

1. Construct a micro-pressure gauge by cutting the stem from a thin-stemmed pipette.

2. Use the hot glue gun to place a small amount of glue to seal the end. Tie a piece of

thread around this end, as shown in Figure 5.

3. Beginning with the sealed end, mark the stem of the pipette in 1.0 cm or otherwise

equal increments.

4. Prepare another large pipette as you did in Part 1 by cutting off the tip.

5. Check the length of the micro-pressure gauge by placing it inside the large pipette.

Cut the pressure gauge so that its length just comes to the top of the bulb.

6. Place a drop of colored water into the pressure gauge by squeezing the center of the

stem with pliers and placing the open end into a cup of colored water. Gently release

the grip to allow the water to be drawn into the pressure gauge. Based on the position

of your colored water droplet, count the number of increments occupied by the

trapped air sample. Record this value in Table 2 on your student answer page.

7. Obtain a small sample of dry ice and fill the bulb about one quarter full. Slide the

pressure gauge into the large pipette, fold the tip, and firmly grip both with the pliers. Be

sure to place the setup into the water as shown in Figure 6 and observe.


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8. As the pressure climbs inside the pipette, you will notice the colored water plug rising in

the micro-pressure gauge. Note the position of the colored water droplet when the dry ice

begins to liquefy.

9. Gently release the grip and repeat until there is no more dry ice in the bulb.

10. Bring your data from your pipette back into class.

Clean up:

1. Clean up any water, dry the plastic cup and leave it at your station. Throw away all

pipettes you used but save your pressure gauge. Do NOT throw away the plastic cups


Page 59 of 127

Properties of Solutions: Electrolytes and Non-Electrolytes


Discussion:

In this experiment, you will discover some properties of strong electrolytes, weak

electrolytes, and non-electrolytes by observing the behavior of these substances in

aqueous solutions. You will determine these properties using a Conductivity Probe.

When the probe is placed in a solution that contains ions, and thus has the ability to

conduct electricity, an electrical circuit is completed across the electrodes that are located

on either side of the hole near the bottom of the probe body. This results in a

conductivity value that can be read by the computer. The unit of conductivity used in this

experiment is the microsiemens, or S.

The size of the conductivity value depends on the ability of the aqueous solution

to conduct electricity. Strong electrolytes produce large numbers of ions, which results in

high conductivity values. Weak electrolytes result in low conductivity, and non-

electrolytes should result in no conductivity. In this experiment, you will observe several

factors that determine whether or not a solution conducts, and if so, the relative

magnitude of the conductivity. Thus, this simple experiment allows you to learn a great

deal about different compounds and their resulting solutions.

In each part of the experiment, you will be observing a different property of electrolytes.

Keep in mind that you will be encountering three types of compounds and aqueous

solutions:

Ionic Compounds

These are usually strong electrolytes and can be expected to 100% dissociate in aqueous

solution.

Example: NaNO3(s) → Na+(aq) + NO3-(aq)

Molecular Compounds

These are usually non-electrolytes. They do not dissociate to form ions. Resulting

solutions do not conduct electricity.

Example: CH3OH(l) → CH3OH(aq)

Molecular Acids

These are molecules that can partially or wholly ionize depending on their strength.

Example: Strong electrolyte H2SO4 →H+(aq) + HSO4-(aq) (100% dissociation)

Example: Weak electrolyte HF ↔ H+(aq) + F-(aq) (<100% dissociation)

SDS

0.05 M NaCl, 0.05 M CaCl2, 0.05 M

AlCl3, 0.05 M HC2H3O2, 0.05 M H3BO3

Substance not considered hazardous. However,

not all health aspects of this substance have

been thoroughly investigated.

0.05 M H3PO4 Slightly toxic by ingestion and inhalation; body

tissue irritant.

0.05 M HCl Slightly toxic by inhalation and ingestion.

Severe body tissue irritant. Corrosive to eyes.

Avoid all body contact.

0.05 M CH3OH (methanol) Toxic by ingestion (may cause blindness),

inhalation or absorption.

0.05 M C2H6O2 (ethylene glycol) Mildly toxic by ingestion and inhalation. Skin,

eye, and mucous membrane irritant. Avoid all


Page 60 of 127

body tissue contact.

Procedure

1. Open the Conductivity lab method on the lab quest. The window will display live

conductivity readings in units of microsiemens (S).

2. The Conductivity Probe is already attached to the labquest. It should be set on the 0-

20,000 S position.

3. Obtain the Group A solution containers. The solutions are: 0.05 M NaCl, 0.05 M

CaCl2, and 0.05 M AlCl3.

4. Measure the conductivity for each of the solutions.

a. Carefully raise each vial and its contents up around the Conductivity Probe until

the hole near the probe end is completely submerged in the solution being tested.

Important: Since the two electrodes are positioned on either side of the hole this

part of the probe must be completely submerged.

b. Briefly swirl the beaker contents. Once the conductivity reading in the Meter

window has stabilized record the value in your data table.

c. Before testing the next solution, clean the electrodes by surrounding them with a

250-mL beaker and rinse them with distilled water from a wash bottle. Blot the

outside of the probe end dry using a tissue. It is not necessary to dry the inside of

the hole near the probe end.

5. Obtain the four Group B solution containers. These include 0.05 M H3PO4, 0.05 M

HC2H3O2, 0.05 M H3BO3, and 0.05 M HCl. Repeat the Step 5 procedure.

6. Obtain the five Group C solutions or liquids. These include 0.05 M CH3OH, 0.05 M

C2H6O2, distilled H2O, and tap H2O. Repeat Step 4.

Cleanup:

1. Rinse off the conductivity probe with distilled water, wipe up any spills and thrown

any trash away.

Questions:

1. Based on your conductivity values, do the Group A compounds appear to be

molecular, ionic, or molecular acids? Would you expect them to partially dissociate,

completely dissociate, or not dissociate at all?

2. Why do the Group A compounds, each with the same concentration (0.05 M), have

such large differences in conductivity values? Hint: Write an equation for the

dissociation of each. Explain.

3. In Group B, do all four compounds appear to be molecular, ionic, or molecular acids?

Classify each as a strong or weak electrolyte, and arrange them from the strongest to

the weakest, based on conductivity values.

4. Write an equation for the dissociation of each of the compounds in Group B. Use for strong; for weak.

5. For H3PO4 and H3BO3, does the subscript “3” of hydrogen in these two formulas

seem to result in additional ions in solution as it did in Group A? Explain.

6. In Group C, do all four compounds appear to be molecular, ionic, or molecular acids?

Based on this answer, would you expect them to dissociate?

7. How do you explain the relatively high conductivity of tap water compared to a low

or zero conductivity for distilled water?


Page 61 of 127

8. Did aqueous methanol, CH3OH, have the same conductivity value as aqueous

ethylene glycol, C2H6O2? Explain.

39 Drop pH Lab


Discussion:

The pH of a solution can be determined chemically or instrumentally. The

chemical method uses an indicator, a substance that changes color in some know pH

range. A Universal Indicator will produce an array of colors (red, orange, yellow, green,

blue, purple, and colors in between) depending on the acidity of the solution being tested.

The purpose of this experiment is to produce a standard pH color chart to be used in

determining the pH of several household products. You will also test the conductivity to

evaluate whether they are non-electrolytes, strong electrolytes or weak electrolytes.

Materials/Equipment:

24 well plate Dropper bottles of

ammonia solution X

bleach solution Y

baking soda, muratic acid, sprite, club soda,

sugar, tap water, distilled water, lemon juice,

milk, saliva, egg white, soap, shampoo,

Universal Indicator

wooden splints, conductivity sensors

SDS:

Universal

Indicator

Contains denatured ethyl alcohol. Moderately toxic by ingestion and

inhalation. Body tissue irritant. Avoid all body tissue contact.

Flammable liquid.

0.1 M HCl Slightly toxic by inhalation and ingestion. Severe body tissue irritant.


Procedure:

1. Place the well plate on a piece of clean white paper.

2. Add drops of solution X and solution Y to the wells in rows A and B as indicated by

the chart on the next page. Stir solutions with a wooden splint. Be careful to control

drops—DON’T miscount!

3. Add about 5 drops of each household solution to separate wells in row C & D.

4. Using a conductivity sensor, test each household product and record the conductivity

as Very High, High, Medium, Low, or None. Rinse off the conductivity sensor

probe wires with a wash bottle between each product.

5. Add one drop of Universal Indicator to each well containing a solution. Stir with a

wooden splint.

6. Determine the pH values of the household solutions by comparison with the other

wells.


Page 62 of 127

Well #

pH (±0.1)

drops X

drops Y

A1

empty

A2

2.0

39

0

A3

3.0

35

4

A4

4.0

31

8

A5

5.0

27

12

A6

6.0

24

15

Well #

pH (±0.1)

drops X

drops Y

B1

7.0

20

19

B2

8.0

17

22

B3

9.0

14

25

B4

10.0

11

28

B5

11.0

9

30

B6

12.0

3

36

Well #

pH (±0.1)

color

substance

conductivity

C1

____

_______

C2

____

________

C3

____

________

C4

____

________

C5

____

________

C6

____

________

Well #

pH (±0.1)

color

substance

conductivity

D1

____

________

D2

____

________

D3

____

________

D4

____

________

D5

____

________

D6

____

________


Page 63 of 127

Titration Lab


Discussion:

A titration is a method of analysis that will allow you to determine the precise endpoint of

a reaction and therefore the precise quantity of reactant in the titration flask. A buret is

used to deliver the second reactant to the flask and an indicator or pH meter is used to

detect the endpoint of the reaction.

Titrations are calculated using dimensional analysis: For example, 13.20 mL of 0.100 M

NaOH is used to titrate 11.19 mL of Sulfuric acid. The molarity of the Sulfuric acid is

unknown.

First, write a balanced equation:

2NaOH + H2SO4 Na2SO4 + 2H2O 0.01320 mL NaOH x 0.100 moles NaOH x 1 mole H2SO4 x ? mol__

1 L 2 moles NaOH 0.01119 L H2SO4

This give an answer in moles of H2SO4 per L H2SO4 which is molarity.

SDS

0.1M NaOH Moderately toxic by ingestion and skin absorption. Corrosive to body

tissues. Causes severe eye burns. Avoid all body tissue contact.

HCl Slightly toxic by inhalation and ingestion. Severe body tissue irritant.


Phenolphthalein Contains denatured ethyl alcohol. Toxic by ingestion and inhalation.

Irritating to body tissues. Avoid all body tissue contact. Flammable liquid.

1. Read the section on titrations in your book in chapter 15.

2. The acid will be hydrochloric acid and the base will be 0.100 M NaOH. Your

indicator will be phenolphthalein.

3. Get two 150 mL beakers, label one base with a sharpie, label the other acid. Go to

the fume hood and put 150 mL of base in the base beaker and 50 mL of acid in the

acid beaker. Pour the acid into one buret and the base into the other. Drain both

burets a little to move the liquid into the tip of the buret.

4. Get a 125 mL Erlenmeyer flask. You will use this to do the titration. Rinse it out

with water twice and distilled water once.

5. Using the buret get 10.0 mL of 0.10 M HCl and titrate it using base to practice your

titrating skills. If you add too much base, back titrate using a little more acid until the

solution goes colorless. Your end point should be very close to 10.00 mL of base.

6. Build a data table. You are running 3 trials for your unknown acid. You must show

an initial and final volume for the base for all 3 trials. Be sure to include the

unknown #. Your volume of acid will always be around 10.0 mL. Get a 50 mL

beaker, rinse it with distilled water and bring to the teacher for your unknown sample.

7. You are to calculate the average mL of base used to titrate 10.0 mL of unknown acid

and use this average to calculate the molarity of the unknown acid. Be sure to show

your calculation.

8. Your teacher may have you repeat this titration for more than one unknown acid.

Cleanup:

1. Pour the remainder of the base down the sink. Pour the remainder of the acid down

the sink. Pour the contents of the Erlenmeyer down the sink.


Page 64 of 127

2. Pour 50 mL of 0.1M HCl into the now empty base buret to neutralize the base. Rinse

both burets with tap water twice and then rinse with distilled water from a wash

bottle. Hang upside down in your double buret clamp with the valve open.

3. Wash all beakers and the Erlenmeyer flask with soap and rinse and put on a paper

towel at your station to dry.


Page 65 of 127

Qualitative Analysis of the Group I Cations


Discussion

The Group I cations are those species that form chloride precipitates that are insoluble in

acid. The group includes Ag+, Pb2+, and Hg22+ (mercurous ion). These cations are precipi-

tated from any other cations that might be present in a sample by addition of 6 M hydro-

chloric acid: Addition of HCl forms a mixture of AgCl, PbCl2, and Hg2Cl2 solids.

Ag+(aq) + Cl-(aq) AgCl(s)

Pb2+(aq) + 2Cl-(aq) PbCl2(s)

Hg+(aq) + 2Cl-(aq) Hg2Cl2(s)

At this point, the sample is centrifuged and the precipitate of the Group I

chlorides isolated. After you centrifuge a sample, there is a solid at the bottom called a

residue and a liquid called the supernate. When you pour the supernate off it becomes

the decantate because you have poured it out of the test tube.

Lead ion is then separated from silver and mercury by taking advantage of the fact

that PbCl2 is much more soluble in hot water than in cold water (the solubilities of AgCl

and Hg2Cl2 do not vary much with temperature). Distilled water is added to the Group I

mixed precipitate, and the mixture is heated to dissolve PbCl2. The mixture is then

centrifuged quickly while still hot, and the decantate containing lead ion is removed from

the remaining silver/mercury precipitate. The presence of lead ion is then confirmed by

addition of dichromate ion, Cr2O72-, which forms a characteristic yellow precipitate with

lead ion, PbCr2O7.

The precipitate containing silver and mercurous ions is then treated with

ammonium hydroxide. Silver ion is complexed by ammonia; the precipitate of AgCl will

dissolve and is removed after centrifugation. A black-gray residue in the centrifuge tube

confirms the presence of mercury.

The decantate containing complexed silver ion is then treated with acid, which

reacts with ammonia, allowing the precipitation of silver chloride. Alternatively,

potassium iodide can be added, which also precipitates the silver (as a creamy yellow-

white solid).

In this experiment a known sample containing all three ions, as well as an

unknown sample (containing one or more cations from the specific group), will be

analyzed. In real practice, a sample would not be restricted to the members of one

analysis group but, rather, would be a general mixture of all possible cations.

Safety Precautions

Protective eyewear approved by your institution must be worn at all times while you

are in the laboratory.

The centrifuge spins rapidly and can eject the sample tubes with considerable

momentum if it is not correctly balanced. If the centrifuge starts to wobble or creep

along the lab bench, immediately disconnect power and balance the centrifuge.

SDS

6M HCl Toxic by inhalation and ingestion. Severe corrosive to all body tissues,

especially skin and eyes. Avoid all body contact.


Page 66 of 127

K2Cr2O7

solution

Moderately toxic by ingestion. Irritating to body tissues. Avoid all body

tissue contact.

6M NH4OH Liquid and vapor are strongly irritating to skin, eyes, and mucous

membranes. Vapor extremely irritating to eyes. May cause blindness.

Toxic by ingestion or inhalation. When heated to decomposition, emits

toxic fumes of NH3 and NOx.

6M HNO3 Corrosive; will cause severe damage to eyes, skin and mucous membranes.

Moderately toxic by ingestion and inhalation. Strong oxidizer. Avoid

contact with acetic acid and readily oxidized substances.

AgNO3 Moderately toxic by ingestion. Irritating to body tissues. Avoid all body

tissue contact. Silver compounds photoexpose and will stain skin and

clothing.

HgNO3 Highly toxic by ingestion and inhalation. Severely corrosive to body

tissues. Avoid all body tissue contact.

Procedure

Label all test tubes during the procedure to make certain that samples are not confused or

discarded at the wrong point. All glassware should be cleaned, rinsed and rinsed with

distilled water before starting the lab.

Take notes as you perform the tests on the known so you can evaluate the unknown.

Step Notes

1. Transfer approximately 1 mL of the

sample solution to a test tube, and then add

10 drops of 6 M HCl to precipitate the

Group I cations.

2. Stir the mixture to mix, and then

centrifuge for one minute until the

precipitate is packed firmly in the bottom of

the tube (be sure to balance the centrifuge).

3. To ensure that sufficient HCl has

been added to the sample to cause the

precipitation of all the Group I cations, add

1 more drop of HCl to the test tube,

watching the supernatant liquid in the test

tube for the appearance of more precipitate.

4. If more precipitate forms on the

addition of 1 drop of HCl, recentrifuge and

retest with additional single drops of HCl

until it is certain that all the Group I

chlorides have been precipitated.

Recentrifuge until the precipitate is packed

firmly in the bottom of the centrifuge tube.

Discard the decantate.

5. Wash the silver/mercury/lead

precipitate in the centrifuge tube by adding 1

mL of distilled water and 2 drops of 6 M

HCl. Stir with a glass rod and centrifuge.


Page 67 of 127

6. Remove the wash water and discard.

Wash a second time, centrifuge, and again

discard the wash liquid. The precipitate now

contains the chlorides of silver, mercury (I),

and lead, separated from all other species.

7. To the mixed chloride precipitate

add 2-3 mL of distilled hot water.

8. Heat the mixture in a boiling-water

bath for 3-4 minutes. Stir the precipitate

during the heating period to dissolve the

lead chloride.

9. While the mixture is still hot,

centrifuge and remove the decantate (which

contains most of the lead ion from the

sample) to a separate test tube.

10. To the decantate (containing lead

ion),add 1 drop of potassium dichromate

solution (see teacher for solution). Allow the

sample to stand for a few minutes. The

appearance of a yellow precipitate of lead

chromate confirms the presence of lead ion

in the original sample.

11. If you are doing the known solution

continue on to step 14. Otherwise, to the

remaining Group I precipitate, add 2-3 mL

of hot water and 1-2 drops of 6M HCl. Stir,

and heat in the boiling-water bath for 3-4

minutes. Centrifuge the mixture.

12. Carefully remove and discard the

supernatant liquid, which may contain lead

ion that was not completely removed earlier.

Add a 2-3-mL portion of hot water and 1-2

drops of 6M HCl.

13. Stir the mixture, and heat in the hot-

water bath. Centrifuge, remove, and discard

the supernatant liquid. These several

washing steps are necessary to remove all

traces of lead ion from the silver/mercury

precipitate.

14. Add 2-3 mL of 6M aqueous

ammonium hydroxide to the silver/mercury

precipitate.

15. Stir the mixture, centrifuge, and

remove the decantate (which contains

dissolved silver ion). The presence of a

gray-black precipitate in the centrifuge tube


Page 68 of 127

at this point confirms the presence of

mercury (I) in the original sample.

16. Add sufficient 6 M nitric acid to the

decantate to make the solution acidic when

tested with litmus paper. The appearance of

a white precipitate of AgCl confirms the

presence of silver ion in the original sample.

Cleanup: All precipitates go in the waste beaker in the fume hood. Wash all test

tubes with soap and water, rinse and rinse with distilled water. Leave test tubes

upside down in the test tube rack to dry.


Page 69 of 127

Unknown #__________ Station#__________________

Name(s)________________________________________________________

Presence of Lead ion Yes No

Presence of Silver ion Yes No

Presence of Mercury ion Yes No

Pb2+, Ag+, Hg22+

HCl(aq)

cold water wash

AgCl, Hg2Cl2

HgNH2Cl Ag(NH3)2+

White and

Hg black HNO3(aq)

AgCl

white

Pb2+

Decantate 2 Residue 2

K2Cr2O7

PbCr2O7

yellow

Group I Cation Analysis

AgCl, PbCl2, Hg2Cl2

hot water wash

NH4OH(aq)

Decantate 3


Page 70 of 127

Copper into Gold: The Alchemist’s Dream


Discussion:

One of the goals of the ancient alchemists was to convert base metals into gold. Although

this goal was never attained by chemical methods, the alchemists were able to perform

many color changes to make metals resemble gold. In this experiment you will produce

some color changes to a copper token and demonstrate diffusion in the solid state.

In this reaction, zinc dissolves in the hot concentrated sodium hydroxide solution to form

sodium zincate, commonly written as Na2ZnO2 or, as obtained in solid form from

concentrated solutions, NaZn(OH)3. As an ionic equation this can be written:

Zn + 2 OH− → 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?


Page 71 of 127

2. Why is it necessary to clean the penny before starting the reaction?


Page 72 of 127

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


Page 73 of 127

Periodic Table


Page 74 of 127

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+


Page 75 of 127

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


Page 76 of 127

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


Page 77 of 127

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


Page 78 of 127

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)


Page 79 of 127

Tomato

Reasons for not using Dimensional Analysis

1. Let's say you're super-intelligent and enjoy solving relatively simple problems in the

most complex manner.

2. Let's say you're tired of always getting the correct answers.

3. Let's say you're an arty type and you can't be confined by the structure of DA. You

like messy solutions scribbled all over the page in every which direction. It's not that

you want to make a mistake. But you really don't care that much about the answer.

You just like the abstract design created by the free-wheeling solution... and the

freedom from being confined by structure.

4. Let's say that you have no interest in going to the prom or making the soccer team,

and you don't mind being unpopular, unattractive, ignorant, insecure, uninformed,

and unpleasant.

Otherwise, You Need Dimensional Analysis!

Personal Testimony:

I was at home, sick with the flu when Mr. Elegante taught my class about Dimensional

Analysis. Despite opportunities given to me to make up the assignments that I had

missed, I chose to not do them. I thought that my mathematical abilities were already

sufficient. How wrong I was! It’s been five years since I took that class--Now I spend

my afternoons panhandling at traffic lights, hoping for passersby to give me spare

change. If I ‘m lucky enough to scam a buck after a day’s work, I’m still not sure if my

hourly rate makes cents.

--Mario

All answers must be in the correct unit with correct significant figures. Show all of your

work and circle your final answer. Do all work in your notebook.

1. A box is 19.3 inches by 0.86 ft x 1.289 dm. Calculate the volume of the box in

cm3.

2. What is the volume, in cm3, of a tank that can hold 18 754 Kg of methanol whose

density is 0.788g/cm3.

3. CaCl2 is used as a de-icer on roads in the winter. It has a density of 2.50 g/ml.

What is the mass of 15.0 L this substance?

4. Convert the speed of light in meters per second to miles per hour. Given: 1 mile

= 1.609 km. speed of light is 3.00 x 108 m/s

5. On the strange planet of Bookoo they have a strange monetary system. The ugle

is a type of currency used on one of the islands. A person wants to buy a zugu for

a special someone. Zugus are found only on 1 remote island. The person must

travel from island to island exchanging money to buy the zugu. You must figure

out the exchange rate given the following information: 3 ugles = 7 kapus, 17

kapus = 5 linpeps, 9 fincors = 23 kapus, 15 rinpers = 8 bloors and 2 fincors = 11

rinpers. Find the final price in ugles if the price on the island is 358 rinpers.

6. If you travel 169.246 km in 15.2 hours, what is your speed in km/hour?

7. Calculate the mass of a person who weighs 96.55 lbs in kg. (See notes for kg to

lbs conversion)


Page 80 of 127

8. Nonsense words taken from the poem Jabberwocky (from Lewis Carroll’s

Through the Looking Glass)

There are 20 tumtum trees in the tulgey wood.

In each tulgey wood is one frumious Bandersnatch.

There are 5 slithy toves in 2 borogoves.

There are 2 mome raths per Jabberwock.

There are 2 Jubjub birds in 200 tumtum trees.

There are 200 mome raths in each borogove.

There are 5 Jubjub birds per slithy tove.

The question is: If there are 5 frumious Bandersnatches, how many Jabberwocks are

there?

9. Because you never learned dimensional analysis, you have been working at a fast

food restaurant for the past 35 years wrapping hamburgers. Each hour you wrap

184 hamburgers. You work 8 hours per day. You work 5 days a week. You get

paid every 2 weeks with a salary of $840.34. How many hamburgers will you

have to wrap to make your first one million dollars? If you can solve the

problem, you will have learned dimensional analysis and you can get a better job.

But, since you won't be working there any longer, your solution will be wrong. If

you can't solve the problem, you can continue working which means the problem

is solvable, but you can't solve it. We have decided to overlook this impasse and

allow you to solve the problem as if you had continued to wrap hamburgers.

Jalapeno Metric Conversion

1. 40 ml = _______ L 2. 5000 L = _______ kl 3. 8 g = _______ kg

4. 12000 L = _______ kl 5. 50 mg = _______ g 6. 6000 m = _______ km

7. 200 kg = _______ g 8. 10000 g = _______ kg 9. 500 ml = _______ L

10. 1 L = _______ ml 11. 4000 L = _______ kl 12. 400 cm = _______ Mm

13. 20 ml = _______ kl 14. 7000 ml = _______ L 15. 7 cm = _______ mm

16. 9000 L = _______ ml 17. 6 m = _______ m 18. 1000 cm = _______ m

19. 11 km = _______ m 20. 80 mg = _______ kg 21. 3 m = _______ mm


Page 81 of 127

Banana

1. A diamond is a pure sample of the element carbon. The tabulated density of diamond

is 3.51 g/cm3. Jewelers use a unit called a carat to describe the mass of diamond. What is

the volume of the stone in a 2.00 carat diamond engagement ring?

2. A farmer owns a rectangular field that fronts along a road. State highway engineers

surveyed the frontage and found it measures 138.3 meters in length. The farmer, who

wants to build a fence around the field, paces off the field’s width and estimates it to be

52 meters. How many meters of fence will the farmer have to build? What mass of

fertilizer will the farmer need to fertilize the field with 0.0050 kg of fertilizer per square

meter of field?

3. The diameter of metal wire is given by its wire gauge number. For example, 16-gauge

wire has a diameter of 0.0508 inch. Calculate the length in meters of a 5.00 pound spool

of 16-gauge copper wire. The density of copper is 8.92 g/cm3.

4. The solder used by plumbers to fasten copper pipes consists of 67% lead and 33% tin.

If you have a 1.0 pound (454 g) block of solder, how many grams of lead do you have?

Grams of tin?

5. You are fortunate to own a ring made of white gold, which actually consists of 60.%

gold and 40.% platinum. If your ring has a mass of 6.89 g, how many grams of platinum

do you have? How many grams of gold?

6. At 25 oC the density of water is 0.997 g/cm3, whereas the density of ice is 0.917

g/cm3. (a) If a soft drink can (volume = 250. mL) is filled completely with pure water

and then frozen at –10 oC, what volume will the solid occupy? (b) Could the ice be

contained within the can?

7. Amethyst is a colored form of the mineral quartz in which the color purple comes from

traces of the element Manganese. (a) To determine the density of amethyst, you take a

stone having a mass of 15.25 g and place it in a 100. mL graduate cylinder containing

45.0 mL of water. On adding the stone, the water surface rises to the 50.8 mL mark.

What is the density of amethyst? (b) If you have a piece of amethyst that is 1.8cm x 0.95

cm, how thick does it have to be to have a mass of 6.50 g?

8. Battery plates in lead storage batteries (the type used in automobiles) are made from a

mixture of two chemical elements: lead (94.0%) and antimony (6.0%). If you have a

battery plate with a mass of 25.0 g, how many grams of lead and how many grams of

antimony are present?


Page 82 of 127

Firehouse

Use the combined gas law to solve the following problems:

1) If I initially have a gas at a pressure of 12 atm, a volume of 23 liters, and a temperature of 200. K, and

then I raise the pressure to 14 atm and increase the temperature to 300. K, what is the new volume of the

gas?

2) A gas takes up a volume of 17 liters, has a pressure of 2.3 atm, and a temperature of 299 K. If I raise the

temperature to 350 K and lower the pressure to 153.2 kPa , what is the new volume of the gas in cm3?

3) A gas that has a volume of 28 liters, a temperature of 45 0C, and an unknown pressure has its volume

increased to 34 liters and its temperature decreased to 35 0C. If I measure the pressure after the change

to be 2.0 atm, what was the original pressure of the gas?

4) A gas has a temperature of 14 0C, and a volume of 4.5 liters. If the temperature is raised to 29 0C and

the pressure is not changed, what is the new volume of the gas?


Page 83 of 127

5) If I have 17 liters of gas at a temperature of 67 0C and a pressure of 88.89 atm, what will be the pressure

of the gas if I raise the temperature to 94 0C and decrease the volume by 5 liters?

6) I have an unknown volume of gas at a pressure of 7.15 psi and a temperature of 325 K. If I raise the

pressure to 1.2 atm, decrease the temperature to 320. K, and measure the final volume to be 48 liters,

what was the initial volume of the gas?

7) If I have 21 liters of gas held at a pressure of 78 atm and a temperature of 637 0C, what will be the

volume of the gas if I increase the pressure to 45 atm and decrease the temperature to 751 K?

8) If I have 2.9 L of gas at a pressure of 5 atm and a temperature of 50 0C, what will be the temperature of

the gas if I decrease the volume of the gas to 2.4 L and decrease the pressure to 2280. torr?

9) I have an unknown volume of gas held at a temperature of 115 K in a container with a pressure of 607.8

kPa. If increasing the temperature to 225 K and increasing the pressure to 30. atm causes the volume of the gas

to be 29 liters, how many liters of gas did I start with?


Page 84 of 127

KFC

Use your knowledge of the ideal and combined gas laws to solve the following problems. Hint: Figuring out

which equation you need to use is the hard part!

1) If four moles of a gas at a pressure of 5.4 atmospheres have a volume of 120 liters, what is the

temperature?

2) If I initially have a gas with a pressure of 84 kPa and a temperature of 350 C and I heat it an additional

230 degrees, what will the new pressure be? Assume the volume of the container is constant.

3) My car has an internal volume of 2600 liters. If the sun heats my car from a temperature of 200 C to a

temperature of 550 C, what will the pressure inside my car be? Assume the pressure was initially 760

mm Hg.

4) How many moles of gas are in my car in problem #3?


Page 85 of 127

5) A toy balloon filled with air has an internal pressure of 1.25 atm and a volume of 2.50 L. If I take the

balloon to the bottom of the ocean where the pressure is 95 atmospheres, what will the new volume of

the balloon be? How many moles of gas does the balloon hold? (Assume T = 285 K)


Page 86 of 127

Marcos

Ideal Gas Law Problems

Use the ideal gas law to solve the following problems: Do all work on this sheet. No work, no credit.

1) If I have 4.02 moles of a gas at a pressure of 576 kPa and a volume of 12.4 liters, what is the

temperature?

2) If I have an unknown quantity of gas at a pressure of 984 torr, a volume of 33.1 liters, and a temperature

of 87 0C, how many moles of gas do I have?

3) If I contain 3.00 moles of gas in a container with a volume of 60. liters and at a temperature of 400. K,

what is the pressure inside the container?

4) If I have 7.7 moles of gas at a pressure of 0.090 atm and at a temperature of 56 0C, what is the volume of

the container that the gas is in?

5) If I have 89.23 g of iodine gas at a temperature of 67 0C, and a volume of 88.89 liters, what is the

pressure of the gas?

6) If I have an unknown quantity of gas at a pressure of 0.523 atm, a volume of 25 liters, and a temperature

of 300. K, how many moles of gas do I have?


Page 87 of 127

7) If I have 21.79 g of carbon monoxide gas held at a pressure of 78 atm and a temperature of 900. K, what

is the volume of the gas?

8) If I have 1.9 moles of gas held at a pressure of 594.6 kPa and in a container with a volume of 500.

milliliters, what is the temperature of the gas?

9) If I have 2.46 moles of gas held at a temperature of 97 0C and in a container with a volume of 4500 cm3,

what is the pressure of the gas?

10) If I have an unknown quantity of gas held at a temperature of 1195 K in a container with a volume of

25.1 liters and a pressure of 560. psi, how many moles of gas do I have?

11) What is the density of carbon tetrachloride gas if the pressure is 11.74 psi and the temperature is 32 ºF?

12) If I have 72 liters (15.79 g) of gas held at a pressure of 3.4 atm and a temperature of 225 K, what is the

molar mass of the gas?

13) What pressure would a mixture of 3.20g of O2, 6.40 g of CH4, and 6.40 g of SO2 exert if the gases were

placed in a 40.0 L container at 127 °C?


Page 88 of 127

Outback Directions: For each of the following problems, show all work including units, circle your final answer and

maintain appropriate significant figures.

1. A tin can with a volume of 20.0 ml contains a gas at 1.00 atm and 27.0 °C. If 800.00 calories of heat is

applied to raise the temperature to 67.0 °C and the volume is unchanged, what is the new pressure inside the

can?

2. A sample of gas occupies 10.0 L at 240. °C under a pressure of 80.0 kPa. At what temperature will the gas

occupy 20.0 L if the pressure is increased to 107 kPa?

3. What is the volume of a gas balloon filled with 4.00 moles of helium when the atmospheric pressure is 748

mm Hg and the temperature is 30.0°C?

4. A sample of nitrogen occupies 300. ml at STP. Under what pressure would this sample occupy 150. ml if

the temperature were increased to 546 °C?

5. A sample of helium occupies 250. ml at 100 °C under a pressure of 1.00 atm. If the pressure were increased

to 1900. mm Hg, to what temperature would the sample have to be heated to occupy a volume of 300. ml?

6. How many moles of nitrogen are contained in 328 ml of the gas under a pressure of 3040. mm Hg at 527 °C.


Page 89 of 127

7. Calculate the volume exerted by 80.0 g of CH4 at 127 °C under a pressure of 1520 mm Hg.

8. What pressure would a mixture of 3.20g of O2, 6.40 g of CH4, and 6.40 g of SO2 exert if the gases were

placed in a 40.0 L container at 127 °C? Hint. Find the number of moles of each gas and then add them up to

find total number of moles. Then use the ideal gas law.

9. Fluorine will effuse at a rate of 7.2 m/s at a certain temperature. What speed will sulfur dioxide gas effuse at

if the temperature is the same?

10. If 12.2 grams of aluminum chlorate decomposes when heated, how many Liters of oxygen will be produced

at a temperature of -34 °C and a pressure of 19.2 psi?

11. If 18.3 g of tin are mixed with steam at a temperature of 134° C and a pressure of 18 atm, how many liters

of hydrogen will be produced?


Page 90 of 127

Pistachios Answer each of the following questions on a separate sheet of paper. You must show work in order to

receive credit. Bring it up to me if you have questions or if you are done.

1. Calculate the mass of solute needed to make 750. mL of 2.50 M sodium chloride solution.

2. What mass of silver nitrate must be used to make 200. mL of 0.100 M solution.

3. Describe in detail how you would make 1.25 L of 0.125 M potassium permanganate solution.

4. How much concentrated acid and how much water must be mixed together to make 1.75 L of 6.00 M

sulfuric acid? Sulfuric Acid concentrate is 18.00 M.

5. What is the final molarity of the solution if you add 17.00g of strontium oxalate and bring to volume of

1.82 L with water, then you take out 15 mL of the initial solution and add it to 77 mL of water?

6. Calculate the molarity of a solution made with 30.0 g of oxalic acid with 500.0 mL of solution.

7. Calculate the molality of a solution made with 24.2 g of copper II nitrate trihydrate (copper II nitrate

with 3 molecules of water attached) in 1300. g of water.

8. If I have a 1.00 M solution of potassium dichromate and I need to make 750. mL of 0.23 M solution,

how do I go about that?

9. If I combine 40.0 mL of 2.00 M sulfuric acid with 40.0 mL of 3.50 M potassium hydroxide, what is the

amount of precipitate that will be formed?

10. If I add 19.2 g of calcium chloride to 2.32 kg of water, what temperature will it freeze at?

11. Calculate the number of grams of sodium chloride needed to make 4.00 kg of water boil at 105 °C.

12. If you were to add 1.93 moles of a non-electrolyte to 1.7 kg of naphthalene, what would the new boiling

point be?

Basil Do all work in your notebook. No work, no credit. You will do all of the odds. If you need more practice

you may do the evens if you want.

1. How many grams of magnesium phosphate, Mg3(PO4)2 are found in 2.8 moles of it?

2. In a lab, you needed to calculate the molar mass of calcium sulfate CaSO42H2O. For 125 grams of it,

how many moles are present?

3. Calculate how many atoms of carbon are found in 7 formula units of iron (III) cyanide, Fe(CN)3.

4. What is the mass of 5 molecules of carbon dioxide?

5. For 567 grams of strontium sulfate, SrSO4, how many formula units are present?

6. For 6.7 x 1012 atoms of xenon, how many moles are present?

7. For 8.9 x 1015 formula units of calcium oxide, CaO, how many grams are present?

8. In 56 moles of sulfur, how many atoms are present?

9. How many grams of magnesium are present in 36.1 moles of it?

10. How many atoms of potassium are present in 11.2 moles of it?

11. For 7 x 1033 atoms of aluminum, how many grams are present?

12. In 15 moles of fluorine, how many grams are present?

13. In 15 moles of sulfur, how many atoms are present?

14. In 89.0 grams of nitrogen, how many atoms are present?

15. For 2.02 x 1026 formula units of magnesium chloride, MgCl2 how many grams are present?

16. For 1.00 x 1012 formula units of ammonium hydroxide, NH4OH, how many moles are present?

17. For 235.1 grams of sodium hydroxide, NaOH, how many formula units are present?

18. For 13 grams of calcium bromide, CaBr2, how many moles are present?

19. What is the mass of 4 molecules of carbon dioxide?

20. What is the mass of 3999 formula units of beryllium bromideBeBr2?


Page 91 of 127

21. Aspartame is an artificial sweetener that is 160 times sweeter than sucrose (table sugar) when dissolved

in water. It is marketed as Nutra-Sweet. The molecular formula for aspartame is C14H18N2O5.

a. Calculate the molar mass of aspartame.

b. How many moles of molecules are present in 10.0 g of aspartame?

c. Calculate the mass in grams of 1.56 mol aspartame.

d. How many molecules area in 5.0 mg aspartame?

e. How many atoms of nitrogen are in 1.2 g aspartame?

f. What is the mass in grams of 1.0 x 109 molecules of aspartame?

g. What is the mass in grams of one molecule of aspartame?

22. Dimethylnitrosamine, (CH3)2N2O is a carcinogenic (cancer causing) substance that may be formed in

foods, beverages or gastric juices from the reaction of nitrite ion (used as a food preservative) with other

substances.

a. What is the molar mass of dimethylnitrosamine?

b. How many molecules are present in 250 mg dimethylnitrosamine?

c. What is the mass of 0.050 mole dimethylnitrosamine?

d. How many atoms of hydrogen are in 1.0 mole dimethylnitrosamine?

e. What is the mass of 1.0 x106 molecules of dimethylnitrosamine?

f. What is the mass in grams of one molecule of dimethylnitrosamine?


Page 92 of 127

Oregano

DON’T WRITE ON ME. ALL ANSWERS MUST APPEAR IN YOUR NOTEBOOK. ALL MATH

PROBLEMS INVOLVING FREQ, WAVELENGTH OR ENERGY MUST INCLUDE THE WORK, ALL OF

THE WORK AND LAST OF ALL, THE WORK.

1 meter = 1 x 109 nanometers (nm)

Fill in the blanks in the following table:

"floating point" scientific notation

4.21x10-3

12,900

0.0000834

659,000,000,000,000

3.00x108

1. Convert 522 nm to m.

2. Convert 4.44x10-7 m to nm.

3. What is the wavelength of a photon with frequency 5.7x1014 Hz? What color is it?

4. What is the frequency of a photon with wavelength 6.2x10-7 m? What color is it?

5. What is the energy of a photon with frequency 7.3x1014 Hz? What color is it?

6. What is the energy of a photon with wavelength 600 nm? What color is it?

7. What is the frequency of a photon with wavelength 565 nm? What color is it?

8. What is the frequency of a photon with energy of 6.22x10-19 J? What color is it?

9. What is the wavelength of a photon with energy 5.5x10-19 J? What color is it?

Table12 Colors, & Information

color wavelength () frequency ()

[nm] [x1014 Hz]

UV <400 >7.5

violet 400 to 420 7.5 to 7.1

blue 420 to 490 7.1 to 6.1

green 490 to 580 6.1 to 5.2

yellow 580 to 590 5.2 to 5.1

orange 590 to 650 5.1 to 4.6

red 650 to 700 4.6 to 4.3

IR >700 <4.3


Page 93 of 127

Steak

Write the formulas for the following covalent compounds on your paper:

1. antimony tribromide

2. hexaboron monosilicide

3. chlorine dioxide

4. fluorine moniodide

5. iodine pentafluoride

6. dinitrogen trioxide

7. phosphorus triiodide

Write the names for the following covalent compounds:

8. P4S5

9. O2

10. SeF6

11. Si2Br6

12. SCl4

13. CH4

14. B2Si

15. NF3

Brisket

Name the following chemical compounds on your paper:

1) NaBr

2) Ca(C2H3O2)2

3) P2O5

4) Ti(SO4)2

5) FePO4

6) K3N

7) SO2

8) CuOH

9) Zn(NO2)2

10) V2S3

11) HCl

Write the formulas for the following chemical compounds:

12) silicon dioxide

13) nickel (III) sulfide

14) manganese (II) phosphate

15) silver acetate

16) diboron tetrabromide

17) magnesium sulfate

18) potassium carbonate

19) ammonium hydrogen phosphate

20) tin (IV) selenide

21) carbon tetrachloride


Page 94 of 127

Ribs Do the following in your notebook. On your paper. The sheet is yours to keep.

Ionic Compounds. Given the name, write the formula. Given the formula, write the name.

1. sodium hydroxide

2. mercury (I) sulfate

3. calcium hypochlorite

4. lead (II) phosphate

5. aluminum chlorate

6. ammonium sulfide

7. copper (I) oxalate

8. antimony (V) oxide

9. manganese (III) sulfite

10. silver oxide

11. zinc nitrite

12. chromium (III) silicate

13. ammonium dichromate

14. iron (III) selenide

15. UCl5

16. Sn3(PO3)2

17. WH6

18. Co(ClO4)2

19. KCN

20. (NH4)2SO4

21. Sc3AsO4

Molecular Compounds. Given the name, write the formula. Given the formula, write the name.

22. N2F

23. S2O7

24. PCl5

25. SiO2

26. CCl4

27. sulfur heptoxide

28. tetraselenium diiodide

29. dinitrogen monoxide

30. pentaphosphorus nonabromide


Page 95 of 127

Chicken

Given the formula, write the name. Given the name, write the formula. All answers must be on your paper.

You must also record whether it is an acid, ionic or molecular.

1. NaHSO4

2. KClO4

3. H3As

4. calcium hydroxide

5. barium cyanide

6. KH2PO4

7. P4O10

8. Ba(NO2)2

9. magnesium bromide

10. ammonium dichromate

11. cadmium arsenate

12. silver chloride

13. ZnC2O4

14. carbon disulfide

15. PCl3

16. carbonic acid

17. sulfur hexafluoride

18. H2SO3

19. HCN

20. Copper I peroxide

21. chromium II bicarbonate

22. ammonium acetate

23. H2CrO4

24. C2O4

25. gold I nitride

26. disulphur heptaiodide

27. hexarsenic nonafluoride

28. oxygen

29. zinc nitrate

30. ammonium sulfide

31. perchloric acid

32. calcium thiosulfate

33. Ni(NO3)2

34. Pb(SO4)2

35. aluminum thiocyanate

36. manganese IV dihydrogen phosphate

Sausage

Do all work in your notebook. Do not write on this sheet. Follow all sig fig rules.

Calculate the percentage composition of the following compounds:

1. a. Fe2O3 b. Ag2O c. HgO d. sodium sulfide

e. sodium sulfate

2. Calculate the percentage of nitrogen in each of the following compounds:

a. NH4NO3 b. ammonium sulfate c. nitrous acid

3. Calculate the percent composition of iron III dichromate.

4. Calculate the percent composition of potassium oxalate.

5. Calculate the percent composition of aluminum chlorite.

6. Calculate the percent composition of Lead (IV) chloride

7. Calculate the percent composition of disulfur heptabromide.

8. Calculate the percent composition of calcium nitrate.

9. If you have 27.92 g of iron III dichromate, how many grams of iron are in the compound? (See percent

you calculated in #3)

10. If you have 1920. g of sodium sulfide, how many grams of sulfur are there in the compound? (See

percent calculated in #1d)


Page 96 of 127

Pork Butt

Naming Acids Worksheet. All answers on your paper.

1. Nitric acid

2. Hydrocyanic acid

3. Chloric acid

4. Acetic acid

5. Hydrobromic acid

6. Sulfurous acid

7. Chlorous acid

8. Boric acid

9. Hydrochloric acid

10. Phosphoric acid

11. Nitrous acid

12. Hydrofluoric acid

13. Perchloric acid

14. Hydroiodic acid

15. Phosphorous acid

16. Carbonic acid

17. Sulfuric acid

18. Formic acid (the formate ion is COOH-)

19. Thiocyanic acid

Name each of the following acids:

1. HClO4

2. HCOOH

3. H3PO4

4. HCl

5. H3BO3

6. H2SO4

7. HNO2

8. HI

9. HCH3COO

10. HF

11. H3PO3

12. HCN

13. HClO3

14. H2CO3

15. H2SO3

16. HClO2

17. HNO3

18. HBr

19. H2S2O3


Page 97 of 127

BACON

1. calcium iodide

2. magnesium nitrate

3. iron(II) acetate

4. lead(IV) oxide

5. calcium hydride

6. copper(II) sulfide

7. ammonia

8. hydrosulfuric acid


10. iron(II) thiocyanate

11. calcium selenide

12. magnesium sulfide

13. sodium hydride

14. lead(IV) oxide

15. potassium chromate

16. cadmium bromide

17. copper(II) sulfide

18. aluminum acetate

19. ammonium hydrogen phosphate

20. antimony(III) phosphide

21. antimony(V) sulfite

22. vanadium(V) carbonate

23. nickel(II) sulfate

24. arsenic(II) sulfite

25. tin(II) fluoride

26. lead(IV) sulfide

27. hydrobromic acid

28. lead(II) chlorate

29. manganese(II) oxalate

30. carbon tetrachloride

31. manganese(III) dichromate

32. ammonium sulfate

33. mercury(II) hydrogensulfate

34. mercury(I) nitrite

35. cobalt(II) hydroxide

36. tin(IV) sulfide

37. lithium perchlorate

38. manganese(III) permanganate

39. tungsten(VI) phosphate

40. sulfur dioxide

41. copper(II) sulfate

42. gold(I) phosphide

43. zinc nitrite

44. rubidium hydride

45. mercury(I) phosphate

46. barium nitride

47. mercury(II) phosphate

48. cesium cyanide

49. tetraphosphorus decoxide

50. silver cyanide

51. strontium peroxide


53. cesium dihydrogen phosphate

54. rubidium chromate

55. tungsten(VI) sulfite

56. ammonium phosphate

57. disulfur dichloride

58. sodium hydrogen phosphate

59. platinum(IV) oxide

60. lithium fluoride

61. tin(II) hydrogen carbonate

62. chromium(VI) phosphide

63. Lead(IV) dichromate

64. chromium(III) acetate

65. gold arsenide

66. nickel(II) carbonate

67. sodium thiocyanate

68. aluminum oxalate

69. ammonium perbromate

70. tin(IV) hydrogen sulfite

71. KCN

72. titanium(IV) sulfite

73. sulfur trioxide

74. lithium permanganate

75. SF6

76. dinitrogen pentoxide

77. NO

78. KSCN

79. CO

80. HF

81. calcium hypochlorite

82. magnesium hydroxide

83. ammonium hypoiodite

84. bismuth(III) bromide

85. HCl(aq)

86. HCN

87. CO2

88. CO2

89. NO2

90. KOCN

91. sulfite ion

92. (NH4)2C2O4

93. ZnS

94. CuCl2

95. MgCl2

96. PCl5


Page 98 of 127

97. Ca(ClO3)2

98. LiNO2

99. CaSO4

100. KH2PO4

101. AgNO3

102. CuCN

103. H2S

104. KHCO3

105. CaO

106. NaHSO4

107. H2CO3

108. Li2HPO4

109. Mg3(PO4)2

110. H3PO3

111. KCl

112. MgSO4

113. K2O

114. Ca(IO2)2

115. Al(NO2)3

116. SiO2

117. MgO

118. CuCl

119. SnI

120. KClO2

121. AsCl5

122. CaSO3

123. CuSO3

124. NaBr

125. HF

126. P2O3

127. FeSO4

128. HClO

129. SnCl4

130. NO2

131. AsCl3

132. NaH

133. KCN

134. ZnS

135. NH4OH

136. Pb(NO3)2

137. Fe(ClO4)3

138. H2Se

139. HNO2

140. H3PO4

141. CS2

142. CaH2

143. bromic acid

144. persulfuric acid

145. ammonia


147. CO2



150. H2Se

151. P2O3

152. CS2

153. HF

154. NO2

155. H2S

156. PCl5

157. HCl

158. sulfur dioxide

159. CO

160. CO2

161. NO

162. HF

163. SF6

164. NO2

165. sulfur trioxide


167. tetraphosphorus decoxide


169. disulfur dichloride

170. HI

171. boron trifluoride

172. diphosphorus trioxide


174. P4O10

175. iodine monochloride

176. hydrobromic acid

177. sulfur trioxide

178. hydrofluoric acid


Page 99 of 127

Tofu

Write formulas for the following

1. sodium nitrite

2. sodium carbonate

3. sodium sulfate

4. potassium hydroxide

5. potassium nitrate

6. potassium nitrite

7. potassium phosphate

8. cadmium hydroxide

9. cadmium carbonate

10. cadmium sulfate

11. cadmium phosphate

12. aluminum bromide

13. aluminum nitrate

14. aluminum sulfide


16. magnesium oxalate

17. magnesium hypochlorite

18. barium bromide

19. barium nitrate

20. barium carbonate

21. iron II arsenate

22. iron II chloride

23. iron II hydroxide

24. iron II carbonate

25. iron II sulfite

26. Copper II permanganate

27. iron II phosphate

28. iron III fluoride

29. strontium chloride

30. strontium hydroxide

31. strontium nitride

32. strontium sulfide

33. strontium sulfite

34. iron III sulfate

35. iron III arsenate

36. mercury II bromide

37. lead IV carbonate

38. tin IV hydrogen phosphate

39. chlorine heptoxide

40. chlorine trifluoride

41. sulfur hexafluoride

42. dichlorine monoxide

43. chlorine dioxide

44. carbon tetrasulfide

45. diphosporus heptafluoride

46. sulfur nonafluoride

47. hypochlorous acid

48. nitric acid

49. nitrous acid

50. chloric acid

51. hydrochloric acid

52. chlorous acid

53. sulfuric acid

54. sulfurous acid


56. hydrofluoric acid

57. oxalic acid

Write names for the following

:

58. NaNO3

59. Na2SO4

60. Na3PO4

61. KNO2

62. K2CO3

63. K2SO4

64. CdBr2

65. Cd(NO3)2

66. CdSO3

67. CdS

68. Al4C3

69. Al(OH)3

70. Mg(NO2)2

71. MgSO3

72. Mg3(PO4)2

73. Ba(NO2)2


Page 100 of 127

74. BaSO3

75. BaCO3

76. Al(HCO3)3

77. Al2(SO4)3

78. AlAsO4

79. FeBr2

80. Co2(SO3)3

81. CoS

82. Co(C2O4)2

83. Ba(CH3COO)2

84. SrBr2

85. Sr(H2PO4)2

86. Sn(ClO4)2

87. Sr(CN)2

88. Fe(NO3)2

89. NaOH

90. K2S

91. NO

92. N2O3

93. N3O7

94. NCl7

95. P4O10

96. F2

97. PCl5

98. N2O4

99. HClO4

100. H2S

101. HBr

102. H3P

103. H2C2O4

104. HClO

105. HClO3

106. HClO2

107. H2CO3

108. HCH3COO

109. HNO2

110. HNO3

111. H2SO3

112. H2SO4


Page 101 of 127

Nutmeg

Recall that the symbol [H+] means “H+ ion concentration in moles/liter (M).”

Recall that pH is defined in the following fashion: pH = -log[H+] it is the negative log of the

concentration of H+ ions.

If you want to find [H+] from pH you have to take the inverse log: [H+] = 10(-pH)

Remember to change sign before you take the inverse log.

Inverse logging is done with the 10x key on your calculator. (2nd function, LOG key)

1. Using universal indicator we determine the pH of a solution to be 6.5. What is the [H+] ?

2. Using pH paper we determine the pH of a solution to be 13.2. What is the [H+] ?

3. We measure the concentration of H+ ions to be 1.00 x 10-5 M. What is the pH of this solution?

4. Very strong acids such as concentrated H2SO4 can cause severe injuries, or put large holes in

clothing. This acid has a [ H+] equal to 36.0 M. What is the pH of concentrated sulfuric acid?

5. What is the pH if [H+] is 4.3 x 10-6 M ?






Page 102 of 127


11. What is the pOH if [H+] is 2.5 x 10-5 M ?


13. What is the pH if [H+] is 6.7 x 10+1 M ?

14. What is the pOH if [H+] is 8.0 x 10-13 M ?

15. What is the [OH-] if the pH is 8.7 ?

16. What is the [H+] if the pH is 3.2 ?

17. What is the [H+] if the pOH is 4.5 ?

18. What is the [H+] if the pH is -1.5 ?

19. What is the [H+] if the pOH is 6.9 ?

20. What is the [H+] if the pH is 12.8 ?


Page 103 of 127

Salami

Procedure:

Open the internet browser and enter the address: http://phet.colorado.edu

Click on “Play with Sims” and select “Chemistry” from the menu on the left.

Open the “States of Matter” Simulation and select “Run Now”

Investigation:

1. Predict what the molecules of a solid, liquid and gas look like. Illustrate your prediction with a drawing.

Solid Liquid Gas

2. Complete the table below by exploring the “Solid, Liquid, Gas” tab in the simulation. Test your

predictions and record your observations by recording the temperature and illustrations of each

substance in the three states of matter.

Substances Observations

Solid Liquid Gas

Neon

Temperature:

Illustration:

Temperature:

Illustration:

Temperature:

Illustration:

Argon

Temperature:

Illustration:

Temperature:

Illustration:

Temperature:

Illustration:

Oxygen

Temperature:

Illustration:

Temperature:

Illustration:

Temperature:

Illustration:

Learning Goal: Students will be able to demonstrate their knowledge of the states of matter through illustrations

and descriptions. These illustrations and descriptions should include:

How the molecules in a solid, liquid and gas compare to each other.

How temperature relates to the kinetic energy of molecules.

http://phet.colorado.edu/


Page 104 of 127

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.


Page 105 of 127

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.


Page 106 of 127

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.


Page 107 of 127

Quinoa

http://phet.colorado.edu/new/simulations/sims.php?sim=Gas_Properties

If the direct link does not work, use Google, and use the search terms “Phet Gas properties”.

Part 1: Play

Purpose: Play! See how everything works before we try to find any relationships.

Procedure:

1. Pump the handle once for the heavy particles. Watch the temperature and pressure gauge. See how long it takes for

the values to stabilize. Record the values in Table 1 on the next page.

2. Using the little person pushing on the wall, decrease the volume of the box in about half. Record the results in Table 1.

If it explodes reset it and this time don’t pump as much gas.

3. Using the little person pushing on the wall, make the volume as small as possible and watch for at least 60 seconds.

Thoroughly describe what happens over the course of the full minute. Be sure to address the following in your answer:

a. What happens to the container and the little man pushing on the container?

b. What happens to temperature and pressure?

4. Hit the “Reset” button. Give the pump a push with the heavy species. Wait for the values to stabilize and record the

results in table 1 on the next page.

Are these values the exact same as the first time you did it? __________ Why or why not?

5. There is a box in the lower right corner entitled “Gas in Pump”. Select the “light species”. You will notice that the

pump turns red. Give the pump a press. Wait for the values to stabilize and record the results in table 1 on the next

page.

6. In the box entitled “Heat Control”, grab the arrow and move it to add. What happens?

7. Using the “Heat Control”, grab the arrow and move it to “remove.” What happens?

8. Feel free to play with the simulator. Try the “Pause” and “Step” buttons at the bottom. Try sliding the top of the

container. Take about 3-5 minutes to play with the different options.

http://phet.colorado.edu/new/simulations/sims.php?sim=Gas_Properties


Page 108 of 127

Table 1: Playing with the simulation program

After pumping the

“heavy” handle

once (1st time)

After cutting the

volume in about

half

Reset: Constant

volume 1 pump of

heavy species

Added 1

pump of light

species

Temperature (K)

Pressure (atm)

Number of Heavy

Species

Number of Light

Species

Part 2: Comparing the effect of particle size.

Purpose: Your goal is to see if there is a difference between using heavy or light particles.

Procedure:

1. Hit the Reset button.

2. Under the “Measurement Tools” section, be sure to have the “Energy Histograms” checked.

3. Under the “Advanced Options” tab, be sure that the temperature is check with a value of 300 K.

4. Set the “Constant Parameter” to be volume.

5. Type in the appropriate number of particles in the boxes on the far right of the screen.

6. Fill in the data table below.

7. When finished with one trial, repeat from step 1 (be sure to reset, just to be safe).

8. For the last rows of the table, chose a combination of heavy and light particles to a total of 500.

Table 6: Measurement of Pressure, Temperature, and moles.

Trial 1 Trial 2 Trial 3 Average

100 Heavy

Particles

Pressure (atm)

Temperature (K)

100 Light

Particles

Pressure (atm)

Temperature (K)

50 of each

Type

Pressure (atm)

Temperature (K)

___ heavy

___ light

Pressure (atm)

Temperature (K)


Page 109 of 127

Questions:

1. Is there a significant difference between the averages of the 100 light particles and the 100 heavy?

2. How does the particle size and mass change the pressure or temperature?

3. Describe how the lighter particles move compared to the heavy particles.

4. Considering that all the particles in the container are at the same temperature, why do the particles move at

different speeds?


Page 110 of 127

Turkey

Mass to Mass Problems

In the following problems, calculate how much of the indicated product is made. Show all your work.

1) NaOH + HCl ---> NaCl + H2O

If you start with 10.0 grams of sodium hydroxide, how many grams of sodium chloride will be

produced?

2) C2H4 + O2 ---> CO2 + H2O

If you start with 25 grams of ethylene (C2H4), how many grams of carbon dioxide will be produced?

3) Li + CaCl2 ---> LiCl + Ca

If you start with 5.5 grams of lithium chloride, how many grams of calcium chloride will be produced?

4) HCl + NaCH3COO ---> NaCl + CH3COOH

If you start with 194.67 grams of sodium acetate, how many grams of acetic acid will be produced?

Answer the following problems:

5) CaBr2 + KCl ---> CaCl2 + KBr

How many grams of potassium chloride would you need to make 723 grams of potassium bromide?


Page 111 of 127

6) C4H8 + O2 ---> CO2 + H2O

How many grams of 2-butene (C4H8) would you need to make 50.2 grams of carbon dioxide?

7) FeCl3 + H2O ---> FeCl3 + H2O

How many grams of water would you need to completely react with 250.5 grams of iron (III) chloride?

8) H2CO3 ---> H2O + CO2

If 15 grams of water was made in this reaction, how many grams of carbon dioxide were made?


Page 112 of 127

Stuffing

Stoichiometry Calculations For the following questions, balance the equation and calculate grams of the underlined product. Show all your

work.

1) Zn + H2SO4 --> ZnSO4 + H2 (You begin with 10. grams of zinc)

2) CaO + H2O --> Ca(OH)2 (You begin with 55.8 grams CaO)

3) ZnS + O2 --> ZnO + SO2 (45 grams of ZnO is produced)

4) Fe(NO3)3 + NH4OH --> Fe(OH)3 + NH4NO3 (17.8 g. NH4NO3 produced)

5) Fe + O2 --> Fe2O3 (You start with 500.0 grams of iron)

6) NaCl + MgBr2 --> NaBr + MgCl2 (You start with 34 grams of MgBr2)

7) PbCO3 + HCl --> PbCl2 + H2CO3 (You start with 21 grams of HCl)

For the following questions, write a balanced equation and calculate grams called for in the problem. Show all

your work.

8) How many g of magnesium sulfide is formed when 35.0 grams of magnesium reacts with an excess of

sulfur?

9) How many grams of beryllium hydroxide is needed to react completely with 100 grams of hydrochloric

acid?

10) How many g of iron (II) nitrate is formed when 25 grams of iron is dissolved in pure nitric acid?

11) How many grams of lead (II) sulfide is precipitated in the reaction between 21.0 grams of lead (II)

bromide and excess hydrogen sulfide?

12) How many g of lithium chloride is produced by the reaction of 230 grams of lithium metal and excess

chlorine gas?

13) How many L of carbon dioxide is formed when 23 grams of C6H8 is burned in an excess of oxygen gas

at STP? ( 1.00 mole of any gas at STP = 22.4 L)


Page 113 of 127

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?


Page 114 of 127

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?


Page 115 of 127

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?


Page 116 of 127

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:


Page 117 of 127

C6H6 + Cl2 -----------> C6H5Cl + HCl

benzene chlorobenzene

9. Certain salts of benzoic acid have been used as food additives for decades. The potassium salt of benzoic

acid, potassium benzoate, can be made by the action of potassium permanganate on toluene.

C7H8 + 2 KMnO4 -------> KC7H5O2 + 2 MnO2 + KOH + H2O

toluene potassium

benzoate

If the yield of potassium benzoate cannot realistically be expected to be more than 68%, what is the minimum

number of grams of toluene needed to achieve this yield while producing 10.0 g of KC7H5O2?

10. Aluminum dissolves in an aqueous solution of NaOH according to the following reaction:

2 NaOH + 2 Al + 2 H2O -----> 2 NaAlO2 + 3 H2

If 84.1 g of NaOH and 51.0 g of Al react:

i) Which is the limiting reagent?

ii) How much of the excess reagent remains?

iii) What mass of hydrogen is produced?


Page 118 of 127

Pumpkin Pie

Super Fun Stoichiometry Problems!!!!!

1. The energy used to power one of the Apollo lunar missions was supplied by the following overall reation:

N2H4 + (CH3)2N2H2 + N2O4 N2 + CO2 + H2O

For the phase of the mission when the lunar module ascended from the surface of the moon, a total of 1200.0 kg

N2H4 was available to react with 1000.0 kg (CH3)2N2H2 and 4500.0 kg N2O4.

a. For this portion of the flight, which of the allocated components was used up first?

b. How much water in kilograms was put into the lunar atmosphere.

2. The Ostwald process for producing nitric acid from ammonia consists of the following steps:

NH3 + O2 NO + H2O

NO + O2 NO2

NO2 + H2O HNO3 + NO

If the yield in each step is 94%, how many grams of nitric acid can be produced from 5.00 kg of ammonia?

3. Magnesium is obtained from seawater when calcium hydroxide is added to it to precipitate magnesium

hydroxide. The precipitate is filtered and reacted with hydrochloric acid to produce magnesium chloride. The

magnesium chloride is then electrolyzed to produce Mg and Cl2. If you are working at the magnesium making

factory and wanted to produce 500.0 g of Mg, how many grams of magnesium hydroxide would have to be

precipitated from the seawater?

4. Phosphate baking powder is a mixture of starch, sodium bicarbonate, and calcium dihydrogen phosphate.

When mixed with water, phosphate baking powder releases carbon dioxide gas, causing a dough to rise.

NaHCO3 + Ca(H2PO4)2 Na2HPO4 + CaHPO4 + CO2 + H2O

If 0.750 L of carbon dioxide is needed for a cake and each kg of baking powder contains 168 g of sodium

bicarbonate, how many grams of baking powder must be used to generate this amount of carbon dioxide?


Page 119 of 127

Reeses Cup

Writing Complete Equations Practice

For each of the following problems, write complete chemical equations to describe the chemical process taking

place. Write them in the space to the right of the equations below

1) When lithium hydroxide pellets are added to a solution of sulfuric acid, lithium sulfate solution and

water are formed.

2) If a copper coil is placed into a solution of silver nitrate, silver crystals form on the surface of the

copper. Additionally, highly soluble (meaning it dissolves in water) copper (I) nitrate is generated.

3) When crystalline C6H12O6 is burned in gaseous oxygen, carbon dioxide gas and water vapor are formed.

Balance the equations below:

1) ____ N2 + ____ H2 ____ NH3

2) ____ KClO3 ____ KCl + ____ O2

3) ____ NaCl + ____ F2 ____ NaF + ____ Cl2

4) ____ H2 + ____ O2 ____ H2O

5) ____ Pb(OH)2 + ____ HCl ____ H2O + ____ PbCl2

6) ____ AlBr3 + ____ K2SO4 ____ KBr + ____ Al2(SO4)3

7) ____ CH4 + ____ O2 ____ CO2 + ____ H2O

8) ____ C3H8 + ____ O2 ____ CO2 + ____ H2O

9) ____ C8H18 + ____ O2 ____ CO2 + ____ H2O

10) ____ FeCl3 + ____ NaOH ____ Fe(OH)3 + ____NaCl

11) ____ P + ____O2 ____P2O5

12) ____ Na + ____ H2O ____ NaOH + ____H2

13) ____ Ag2O ____ Ag + ____O2


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14) ____ S8 + ____O2 ____ SO3

15) ____ CO2 + ____ H2O ____ C6H12O6 + ____O2

16) ____ K + ____ MgBr2 ____ KBr + ____ Mg

17) ____ HCl + ____ CaCO3 ____ CaCl2 + ____H2O + ____ CO2

18) ____ HNO3 + ____ NaHCO3 ____ NaNO3 + ____ H2O + ____ CO2

19) ____ H2O + ____ O2 ____ H2O2

20) ____ NaBr + ____ CaF2 ____ NaF + ____ CaBr2

21) ____ H2SO4 + ____ NaNO2 ____ HNO2 + ____ Na2SO4


Page 121 of 127

Nerds Synthesis and Decomposition Worksheet

Directions: For each reaction, identify it as Synthesis or Decomposition. Also, write down the

rule number that applies to each reaction. Predict the products. Be sure to balance the equation

and show states of matter if you can. Do all answers in your notebook.

1. Solid sodium chlorate is heated.

2. Solid potassium oxide is mixed with water.

3. Dinitrogen trioxide gas is combined with water

4. Solid lithium carbonate is heated.

5. Water undergoes electrolysis.

6. Carbon dioxide gas and barium oxide are reacted.

7. Aluminum carbonate is heated.

8. Sulfur trioxide gas and water are allowed to react

9. Phosphoric acid is heated.

10. Cadmium hydroxide solution is heated.

Solid rubidium oxide is combined with water.

11. Diphophorus trioxide gas is combined with silver oxide solution.

Milk Duds

In your notebook please

1. Sulfur dioxide gas is bubbled through water

2. Lithium oxide is added to water.

3. Aluminum chloride is electrolyzed

4. A small piece of sodium is added to a container of iodine vapor.

5. Carbonic acid is heated.

6. Lead is added to a solution of magnesium acetate

7. Copper is added to a solution of hydrochloric acid

8. Iodine crystals are added to a solution of sodium chloride

9. Calcium metal is added to a solution of nitrous acid

10. A solution of iron III chloride is poured over a piece of platinum wire.

11. Zinc pellets are added to hydrobromic acid

12. Zinc acetate solution and cesium hydroxide solution are combined.

13. Ammonium sulfide solution is reacted with hydrochloric acid

14. Solid calcium carbonate is reacted with sulfuric acid

15. Carbon dioxide gas is reacted with solid potassium oxide

16. Cobalt II hydroxide is added to hydroiodic acid.

17. A piece of chromium is immersed in a solution of tin IV thiosulfate.

18. You pick up a straw and blow carbon dioxide into a glass of water.


Page 122 of 127

Candy Corn

In your notebook please

1. iron + hydrochloric acid iron III

2. silver nitrate(aq) + potassium chloride(aq)

3. sodium oxide + sulfur dioxide

4. copper + hydrofluoric acid copper II

5. magnesium chlorate

6. calcium + water

7. propane (C3H8) + O2

8. dinitrogen pentoxide + water

9. sodium sulfite(aq) + hydrochloric acid (aq)

10. barium + phosphoric acid

11. chlorine + iron III fluoride

12. water


Page 123 of 127

Milky Way

Synthesis and Decomposition Worksheet In your notebook please

For each equation, predict the products, balance the equation and give the rule (#1, #2, etc) used to work each

problem.

1. SO2 + H2O

2. CO2 + H2O

3. P2O5 + H2O

4. CaO + H2O

5. SO2 + MgO

6. Sr + P

7. H2CO3 (when heated)

8. Ba(OH)2 (when heated)

9. NaCl (with electricity)

10. KClO3 (when heated)

11. CaCO3 (when heated)

12. Ag2O (when heated)

Single Replacement

1. Fe + Cu(NO3)2

2. Na + H2O

3. Zn + HCl

4. Br2 + KI

5. Li + H2O

6. Fe + KCl

7. Pb + H2O

8. NaCl + Br2

Double Replacement (Use solubility rules)

1. KOH(aq) + HCl(aq)

2. 2KNO3(aq) + H2SO4(aq)

3. AgNO3(aq) + NaCl (aq)

4. HClO3(aq) + NH4OH(aq)

Combustion. For each one of the following, combine the given with oxygen to produce carbon dioxide and

water.

1. CH4

2. C6H12O6

3. C3H8

4. C2H5OH


Page 124 of 127

Vanilla Tootsie Roll

Chemistry equation practice Directions: Do all work on this sheet, (preferably in pencil). For each of the following word

equations, write the formulas, predict the products, and balance the equation. If no reaction

occurs, write NR in place of products. Write down whether it is a synthesis (S), decomposition

(D), single replacement (SR), double replacement (DR) or combustion (C).

1. Magnesium + oxygen

2. Aluminum + hydrochloric acid

3. Sodium oxide (s) + sulfur dioxide (g)

4. Phosphoric acid (heated)

5. Sodium chlorate

6. Zinc chloride (aq) + ammonium sulfide (aq) \

7. Strontium carbonate (s) (heated)

8. Mercury II sulfate (aq) + ammonium nitrate (aq)

9. Iron + copper II sulfate (aq)

10. Zinc + sulfuric acid

11. Dinitrogen pentoxide (g) + water

12. Chlorine + magnesium iodide (aq)

13. Potassium + water


Page 125 of 127

14. Iron + hydrochloric acid

15. Cobalt III hydroxide (aq) + nitric acid

16. Bromine + sodium iodide (aq)

17. Sodium hydroxide (aq) + phosphoric acid

18. Ammonium sulfate (aq) + calcium hydroxide (aq)

19. Silver nitrate (aq) + potassium chloride (aq)

20. Magnesium hydroxide (aq) + phosphoric acid

21. Iron II sulfide (s) + hydrochloric acid

22. Ammonium sulfide (aq) + iron II nitrate (aq)

23. Sulfuric acid + potassium hydroxide (aq)

24. Aluminum sulfate (aq) + calcium phosphate (s)

25. C12H22O11 burns

26. Silver acetate (aq) + potassium chromate (aq)

27. Ammonium phosphate (aq) + barium hydroxide (aq)

28. Calcium oxide(s) + diphosphorus pentoxide (s)


Page 126 of 127

Dots

1. sodium + iodine →

2. magnesium + oxygen →

3. hydrogen + oxygen →

4. nickel (II) chlorate →

5. barium carbonate →

6. zinc hydroxide →

7. aluminum + sulfuric acid →

8. potassium iodide + chlorine →

9. iron + copper (II) nitrate → iron (II)

10. nickel + zinc nitrate → nickel (II)

11. sodium bromide + iodine →

12. potassium + water →

13. silver nitrate + zinc chloride →

14. copper (II) hydroxide + acetic acid →

15. iron (II) sulfate + ammonium sulfide →

16. C3H8 burns

17. C12H22O11 burns

18. Radium bromide + electricity

19. C2H5OH burns

20. C6H12 burns

21. C4H6 burns

22. Magnesium + iodine

23. Copper (II) chloride solution + hydrosulfuric acid (aq)

24. Sodium hydroxide solution + perchloric acid (aq)

25. Zinc carbonate is heated

26. Sodium + magnesium chloride solution

27. Potassium + chlorine

28. Aluminum carbonate is heated

29. Aluminum plus nitrogen

30. Sulfurous acid is heated

31. Iron III hydroxide solution is heated

32. Osmium (IV) oxide + diphosphorus trioxide

33. sodium oxalate solution + aluminum nitrate (oxalates follow the rules of carbonates)

34. Strontium + lithium carbonate

35. Cadmium chloride + fluorine


Page 127 of 127

Cream of tartar Use your own paper to show all work. 1. What is the concentration of hydroiodic acid if 59.69 mL of it is neutralized by

40.39 mL of a .2968 M lithium hydroxide solution?

2. It requires 65.95 mL of potassium hydroxide solution to neutralize 89.29 mL of

0.2118-M nitric acid. What is the concentration of the potassium hydroxide solution?

3. What is the molarity of a copper(I) hydroxide solution if 50.50 mL of the solution

is titrated to the endpoint with 51.99 mL of 0.3574-M carbonic acid?

4. What volume of a 0.8351-M iron(II) hydroxide solution is needed to neutralize 98.35 mL of 0.5417-M phosphoric acid?

5. Complete the table

pH pOH [H3O

+] [OH

-]

2.05

5.05

9.90 X 10-9

7.27 X 10-4

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