chemistry 100 lab - san diego miramar...

Chemistry 100 Lab Fundamentals of Chemistry

Fall 2019

Dr. Fred O. Garces

Dr. Leigh Plesniak

Miramar Chem. Laboratory Policy and Procedure 08/19 -FG

Welcome to Miramar College's Chemistry 100 laboratory. There are a number of policies that should be followed when you are in S5-209.

Please be sure to read everything carefully and to practice these policies so your safety is not compromised.

1. All students are required to have, and wear, safety goggles when conducting experiments. There are NO exceptions to this rule. Goggles

are available for purchase at the bookstore. If you are not wearing your goggles, do not proceed with the experiments. You instructor will not

allow you to be present in the lab classroom. You should bring safety goggles by the time you begin the first experiment. Some instructors will

penalize your lab safety quiz up to 50% off for not bringing or using your goggles during experiments.

2. Each group who share a locker must purchase a Master™ Combination lock with the following serial numbers V99XXX, V629XXX, or

10976xxx (the last three numbers are different from lock to lock). The Miramar Bookstore sells these locks and it is a departmental policy that all

students must exchange the department's brass lock with these Master™ combination locks by the second week of the term. If you do not have

a lock, the instructor will ask you to go to the bookstore to purchase these locks before being allowed to perform an experiment. All other locks

not conforming to the Master™ Combination lock just described, will be cut off. These Master™ locks can be open by the lab technicians with a

master key. Lab techs must have access to student lockers at all times in case of emergency.

3. Students are not allowed into the classroom earlier than ten minutes prior to the start of class. Furthermore, students are not allowed in the

lab room without instructor supervision.

4. The department has a very strong HAZMAT policy. Absolutely no chemicals—not even things like sodium chloride and sucrose are

allowed down the sinks. Contamination of the drainage system with toxic chemicals cannot be compromised. The department has a blanket

policy that all chemicals must be placed in the proper waste container. Violation of this rule will result in a grade adjustment on the student’s lab

reports.

5. The MSDS library contains hazardous information for all chemicals used in this course and can be found in the cabinet above the printer.

Also available in these cabinet are the Merck and Handbook of Chemistry and Physics. The reference texts are for student use.

6. Students are not allowed to use the phone in the lab or the prep room. All cell phones should be turned off when entering the laboratory. If it

is an emergency, the instructor will place the phone call. There is an emergency box at the front of the room to contact the district's dispatch

office.

7. Students are not allowed in the prep room (instrument and chemical storage), the lab technician’s office, or behind the instructor’s table at

any time.

8. The lab room is equipped with hot plates, balances and Bunsen burners and other community equipment. Your instructor will inform you of

where these are stored. Students must return all equipment back to its proper location before leaving the classroom. Under no circumstances

are community equipment to be kept in the student's lockers.

9. A first-aid kit in located at the rear of the class, near the entrance to the prep room. The hand broom and dustpan are found just to the right

of the instructor’s table. The emergency shower is on the north west corner of the room. Contact a lab tech or your instructor in the event of any

chemical spill. It is an absolute policy to report all accidents to the instructor.

10. Students are not allowed to keep other supplies/equipment that has been set out for each experiment such as: buret, pH meters,

thermometer, rulers in their lockers. Violating this rule will result in poor lab technique grade.

11. It is the responsibility of all students to clean the laboratory before leaving. The lab techs work very hard to keep clean and uncluttered. All

students should clean messy lab benches and sink areas, push in their chairs before leaving the classroom; in general, pick up after themselves.

Regularly leaving the lab without consideration of this policy will result in a poor lab technique grade.

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Table of Contents

Check In and Safety .................................................................................................................................. 4

Lab Drawer Equipment ..................................................................................................................................... 4

Safety Statement............................................................................................................................................... 5

Safety Quiz ........................................................................................................................................................ 7

Exercise 1A: Math Basics ......................................................................................................................... 9

Exercise 1B: Dimensional Analysis ........................................................................................................ 11

Experiment 2: A Penny for Your Thought; Scientific Method Introduction ............................................. 13

Objective ......................................................................................................................................................... 13

Materials .......................................................................................................................................................... 13

Introduction ..................................................................................................................................................... 13

Procedure – Part I ........................................................................................................................................... 14

Experiment 2: Data Sheets ............................................................................................................................ 15

Procedure – Part II .......................................................................................................................................... 17

Data Sheets - Part II........................................................................................................................................ 19

Experiment 3: Measurements, the Metric System & Density ................................................................. 23

Objective ......................................................................................................................................................... 23

Materials .......................................................................................................................................................... 23

Introduction ..................................................................................................................................................... 23

Procedure ........................................................................................................................................................ 24


Experiment: Energy of Food ................................................................................................................... 31

Objective ......................................................................................................................................................... 31

Materials .......................................................................................................................................................... 31

Introduction ..................................................................................................................................................... 31

Procedure – ..................................................................................................................................................... 32


Exercise 6A: Naming Compounds.......................................................................................................... 39

Exercise 6B: Lewis Structures & VSEPR .............................................................................................. 41

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Experiment: Separation of a Ternary Mixture......................................................................................... 45

Objective ......................................................................................................................................................... 45

Materials .......................................................................................................................................................... 45

Introduction ..................................................................................................................................................... 45

Procedure – ..................................................................................................................................................... 46


Exercise 8A: Balancing Chemical Equations ....................................................................................... 53

Exercise 8B: Stoichiometry Exercise ..................................................................................................... 55

Experiment: Chemical Reactions & Chemical Equations ...................................................................... 59

Objective ......................................................................................................................................................... 59

Materials .......................................................................................................................................................... 59

Introduction ..................................................................................................................................................... 59

Procedure – ..................................................................................................................................................... 60


Postlab Questions ........................................................................................................................................... 67

Experiment: The Mole ............................................................................................................................. 69

Objective ......................................................................................................................................................... 69

Materials .......................................................................................................................................................... 69

Introduction ..................................................................................................................................................... 69

Procedure – ..................................................................................................................................................... 70

Experiment: Using Gas Laws to Identify an Unknown Liquid ................................................................ 79

Objective ......................................................................................................................................................... 79

Materials .......................................................................................................................................................... 79

Introduction ..................................................................................................................................................... 79

Procedure – ..................................................................................................................................................... 81

Experiment 11: Data Sheets .......................................................................................................................... 83


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Experiment: Concentration of a Salt Solution......................................................................................... 89

Objective ......................................................................................................................................................... 89

Materials .......................................................................................................................................................... 89

Introduction ..................................................................................................................................................... 89

Procedure – ..................................................................................................................................................... 90

Experiment 12: Data Sheets .......................................................................................................................... 91


Experiment: Introduction to Equilibrium and LeChatelier’s Principle ..................................................... 95

Objective ......................................................................................................................................................... 95

Materials .......................................................................................................................................................... 95

Introduction ..................................................................................................................................................... 95

Procedure – ..................................................................................................................................................... 98

Experiment 13: Data Sheets ........................................................................................................................ 101

Experiment: Titration of Vinegar ........................................................................................................... 107

Objective ....................................................................................................................................................... 107

Materials ........................................................................................................................................................ 107

Introduction ................................................................................................................................................... 107

Procedure – ................................................................................................................................................... 107

Experiment 14: Data Sheets ........................................................................................................................ 109

Postlab Questions ......................................................................................................................................... 113

A. Appendix ................................................................................................................................................ 115

Common Conversions .................................................................................................................................. 115

Solubility Rules .............................................................................................................................................. 115

Polyatomic Ions ............................................................................................................................................. 116

VSEPR Table ................................................................................................................................................ 120

Stoichiometry Map ........................................................................................................................................ 116

pH, pOH, H3O+, OH- Relationship ................................................................................................................. 120

Periodic Table ............................................................................................................................................... 120

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Check In and Safety Lab Drawer Equipment

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Safety Statement

The safety of yourself and your classmates is of paramount importance while in the laboratory. In the laboratory the chemist works with many potentially dangerous substances and equipment. The most general rules for safe laboratory operations are: be alert, stay alert, and take the trouble to understand what you are doing and the potential hazards associated with the operation you are performing. Some basic rules and precautions are:

1. Always wear safety glasses to protect your eyes from chemicals and broken glassware.

2. Shoes covering the tops of feet must be worn at all times while in the lab.

3. Never work alone in the laboratory - another person should be in the lab room at all times.

4. Use fume-hood when working with poisonous or offensive gases and fumes, or when conducting procedures involving flames.

5. Never heat organic solvents (alcohol, ether, benzene, etc.) in an open vessel over an open flame. Organic solvents are highly flammable so only hot plates or a heating mantle should be used around these flammable liquids.

6. Avoid pointing the mouth of a vessel that is being heated toward any person, including yourself and the instructor.

7. Never heat chemicals of any kind in a fully closed system - the system must be open to air to prevent pressure build-up & explosion.

8. Never add anything (including water) to concentrated acid - instead slowly add acid to other substances to avoid splashing of acid.

9. Lubricate glass tubes & thermometers with glycerol then hold with a towel or thick gloves when pushing through a rubber stopper.

10. Never pipet anything by mouth - especially toxic or corrosive substances.

11. Immediately sweep up spill taking place on the balance. Clean up all spills immediately, even those occurring on your lab bench.

12. Be sure to label all chemical containers correctly.

13. Do not perform any unauthorized experiments.

14. Beware of hot glass tubing - glassware looks cool long before it can be handled safely.

15. Never throw matches, litmus, or any insoluble solids in the sink.

16. Avoid using excessive amounts of reagents - 1 to 3 mL is usually ample for test tube reactions.

17. Do not lay down the bottle stoppers. Impurities may be picked up and contaminate the solution when the stopper is returned.

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18. Do not heat thick glassware such as volumetric flasks, graduated cylinders, or bottles; they break easily with heat. Always check glassware for stress and fatigue such as stars and cracks before using. Do not use these pieces of glassware.

19. Never pour anything back into a reagent stock bottle - take out only the amount that will be used.

20. Tie back long hair and refrain from wearing flowing, fluffy clothing - both are fire hazards in the laboratory.

21. Know the location of exits, fire extinguishers, eye washers, first aid kits, the fire blanket & other safety devices in the laboratory.

22. De-ionized water may be obtained through the curved faucet of each sink. The faucet operated by the foot pedal is the tap water. (Generally, it is good lab technique to do a final rinse with de-ionized water when cleaning glassware.)

23. Never use the thermometer as a stirrer. The thermometer could break and, in some cases, release toxic mercury.

24. No eating, drinking or chewing gum in the lab whatsoever. If you need to eat or take a swig of water,

step outside of the lab and do so. Safety regulations must always be observed as it only takes one accident to cause blindness or serious

permanent injury. Safety glasses must be worn at all times. After the first day of lab a 50% penalty will be assessed to your safety quiz if you come to lab without your safety glasses. If you remove your safety glasses when lab is being conducted your instructor will give you a verbal warning once, the second time 10 points will be reduced from your lab report, and a third offense will result in dismissal from lab.

Initials__________

SAFETY STATEMENT: I am aware that there are hazards associated with being in a chemistry laboratory. I have been made aware of the safety equipment available in the laboratory room and how it is to be used. I have also been made aware of some common hazardous such as: broken glass fire, acids, bases, and the poisonous nature of most chemicals. I will always wear my safety glasses during lab. I understand that special precautions for individual experiments will appear in the lab manual in a section entitled "Safety". (Please sign below).

Signature Date

Print name (CSID) College Student ID Number

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Safety Quiz Name last: _____________________ First _____________ ____ / __ Score Chemistry 100L Date ___________

Indicate in the space provided whether each of the following lab safety statements is True (T) or false (F).

1 Students must purchase and use protection eyewear when conducting experiments. These eyewear must meet the standards of the American National Standards Institute for "Practice for Occupational and Educational Eye and Face Protection" (American National Standard Institute, ANSI Z87.1). Safety eyewear cannot be removed until the whole class is done with the experiment and all chemical and equipment are put away. Only when it is safe will the instructor must give permission for students to remove their safety goggles.

2 Before working on an experiment, students must know the location of the fire extinguisher, first-aid kit, eye wash station, emergency call box, safety shower, and other safety equipment in the laboratory.

3 It is okay to eat or drink in the lab during the prelab discussion but it is never permissible to smoke in class.

4 It is never permissible to taste chemicals in the lab. All chemicals are to be considered hazardous unless instructed otherwise.

5 Wafting is a technique to smell gas, by placing the odorous material directly under your nose.

6 If any chemicals contact your skin or eyes, flush immediately with water, and then notify the instructor.

7 Although most rules indicate you should never place material from the laboratory into your mouth, there are some exceptions that allow some materials to be place in your mouth such as using a pipet.

8 Perform all reactions that involve gases with an unpleasant odor under a fume hood.

9 When heating a test tube, make sure the test tube is completely closed so that nothing spills.

10 Thick, heavy gloves should be used, and glassware should be lubricated with glycerol or water when inserting glass tubing (i.e., thermometer) into a stopper.

11 It is okay to perform unauthorized experiments.

12 Immediately sweep and clean Balances/Scales after use. In other words, the next person to use the balance (or scale) is not responsible for your mess.

13 It is okay to dispose of all solutions including all organic chemicals down the sink.

14 Pour acids into water (not water into acid) because the heat release will be absorbed by the water so that splattering is avoided.

15 It is okay to disregard the special safety precautions mentioned in each experiment.

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16 Wear closed-toe shoes at all times in the laboratory.

17 If glassware breaks while in use, it is okay to leave it alone, not clean it up, and not tell anyone.

18 Read all parts of an experiment (Objectives, Discussion, Procedure, & Pre-lab Assignment), before performing each lab.

19 It is okay to take an excess amount of chemicals for an experimental procedure. If the chemical is unused, it can always be returned back to the original stock container.

20 An important general rule when performing experiments in the laboratory is to be alert, stay alert, and take the trouble to understand the potential hazard associated with each experiment.

21 It is okay to heat organic solvents directly over a Bunsen burner rather than a hot plate.

22 All glassware should be inspected for stars, cracks, or stress before usage and especially before heating.

23 It is permissible to place chemicals directly on the metal platform of the balance pan when using the balance scale.

24 When an experiment calls for water, use distilled or deionized water. Use tap water first to clean the glassware and then do a final rinse with distilled or deionized water.

25 Before leaving the laboratory, all community lab equipment is to be returned to its original place, i.e. ring stand to location under hoods; all excess solid chemical waste should be placed in a designated waste container; each laboratory station should be straightened-out with chairs pushed in; and students should wash their hands.

This quiz is your contract that you will abide by all safety rules of the laboratory. Failure to comply will reduce your grade on this quiz, and may result in being asked to leave the laboratory if your safety or the safety of others is at risk. Initials__________

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Exercise 1A: Math Basics Last Name________________________First_________________________ ___ / __ pts

Lab Partner(s): _________________________________________________ Day _________𝒂𝒎.𝒑𝒎.

Answer all questions with the correct number of significant figures. I. Rounding off Numbers.

a. Round to the tenth place: 21.46499 = ___________

b. Round to the thousandths place: 482.4106 = ___________

c. Round to the tenth place: 4.150279 = ___________

d. Round to the hundredth place: 1.00154 = ___________

e. Round to the ten-thousandths place: 21.873451 = ___________ II. Significant figures. Write the number of significant figures for each of the numbers below.

a. 0.0180 ____3______ b. 970 ___________

c. 0.90 ___________ d. 45.0 ___________

e. 0.0010 ___________ f. 160.900 ___________

g. 904 ___________ h. 6000 ___________

i. 402010 ___________ j. 0.00006140 ___________

III. Scientific Notation. Rewrite the numbers below in scientific notation.

a. 5 ___________ b. 0.000 000 110___________

c. 7,100 ___________ d. 600,600,000 ___________

e. 67,000 ___________ f. 0.010 2 ___________

g. 0.000 100 ___________ h. 0.001 49 ___________

i. 0.000 63 ___________ j. 1,106,140,000___________

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IV. Math operations using significant figures. Answer the following using the correct number of significant figures. For each, write the raw answer first then rewrite the rounded-off answer in the parenthesis. Use scientific notation where appropriate.

a. 2.0 + 3.88= 5.88 . ( 5.9 x 100 ) b. 34.0 − 2.15 = _________ (_________)

c. 6.944 + 3.3= ___________ (__________) d. 114.45 − 10.= _________ (_________)

e. 5.386 − 2.11 = __________ (__________) f. 2.23456.22

= ________ (_________)

g. 2.57422.22289

= ___________ (__________) h. 2.3 × 10;9 × 2.501 = _______ (_______)

i. 7.93×725

57= ___________ (__________) j. 4000 × 2.6789× 106 = ______ (________)

V Math operations using power-of-ten: Calculate each and write your answers in scientific notation .

a. =2.50 ×10−1> ×=0.300 ×10−2> = ________________

b. (100.1 × 109) ÷ (0.200 ×10;9)= ________________

c.3.2B72C

4B72DE = ________________

d. 2.0 × 103 ÷ 4 × 10;7 = ________________

e. (8.00 + 1) × (10;7) = ________________

f. 8.00 ×10;7 + 1 × 10;7 = ________________

g. (10.62 × 102) + (1.63 × 103) = ________________

*HINT: Answers c & d should be different. Answers e & f should be the same.

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Exercise 1B: Dimensional Analysis Last Name________________________First_________________________ ____ / ____ pts

Lab Partner(s): _________________________________________________________ Day _________𝒂𝒎.𝒑𝒎.

Write all answers with the correct number of significant figures and the correct units.

1. Determine the following conversion factors.

Use only the conversions found in the appendix and write your answers to at least 3 significant figures.

a. 1.00 lb = ____________ g b. 1.00 lb = ____________ kg

c. 1.00 oz = ____________ g d. 1.00 qt = ____________ ml

e. 1.00 qt = ____________ L f. 1.00 gal = ____________ L

g. 1.00 in = ____________ cm h. 1.00 m = ____________ in

i. 1.00 cup = ____________ ml j. 1.00 fl. oz = ____________ ml

2. Answer the following using dimensional analysis.

a. How many kilograms in 155.5 lb? b. How long in centimeter is a 30.5-inch waist?

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c1. A Prius automobile requires 11.9 gallons of gasoline (gas) for a full tank. How many mL of gasoline is needed for a full tank in this Prius?

c2 Refer to the above question. If gasoline has a density of 0.75 grams per mL,

what is the mass (in grams) of this volume of gasoline? (Remember that 0.75 g/ml is a measured conversion factor.)

d. How fast is a car moving in cm / sec if its speedometer reads 65.00 mph? e. How many pennies are needed (19.5 mm diameter) to stretch from the earth

to the sun? It takes 8.00 minutes and 15.0 seconds for light to travel from the sun to the earth traveling at 186,282 mi/sec. How much is this many pennies worth in dollars?

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Experiment 2: A Penny for Your Thought; Scientific Method Introduction

Objective The purpose of this experiment is to familiarize students with the Scientific Method, to learn to distinguish between chemical and physical properties, and to learn to write a supported conclusion based on their experimental data & observations. Students will form a hypothesis, test it, revise it and test it again. During this process students gain familiarity with some laboratory equipment and solutions including balances, rulers and working with strong acid in the hood. Critical thinking skills will be developed as the students draw conclusions from their data.

Materials o Pennies from 1960 – present o Metric ruler o Tongs/forceps o Balance o 50 -ml graduated cylinder o Shear metal cutter o Petri dish o 1 M HCl

Introduction Begin by Reading: Ch 1.2 Chemistry in our Lives Timberlake 5th Edition. “Scientific Method: Thinking Like a Scientist”. Familiarize yourself with Questions 1.7 and 1.9 at the end of this section. In addition, read Chapter 2.7 “Density” & Chapter 3.2 “States and Properties of Matter.”

Pennies & Mass For this experiment, we begin with the question, “What happens to the mass of a coin as it ages?” In this experiment a hypothesis is written for what is expected to happen to the mass of a penny as it ages. The mass may increase with dirt and oils adhering to the surface. It may also wear and lose mass. You will be asked to make a prediction. Experiments will follow that may or may not support your hypothesis.

Physical & Chemical Properties In the experiments that follow, you will be asked to record observations about the physical and chemical properties of pennies. Physical properties include color, size, density and hardness. Chemical properties describe reactivity of a substance, including flammability. In a chemical change, a chemical reaction occurs and a substance is converted into one or more new substances.

One physical property that will be determined is density.

𝐷𝑒𝑛𝑠𝑖𝑡𝑦 = 𝑚𝑎𝑠𝑠𝑣𝑜𝑙𝑢𝑚𝑒

Conclusions You will be asked to write a conclusion based on the data and observations from your experiments. You must cite physical properties and chemical changes in your conclusion to support it. If your conclusion does not support your hypothesis, you may need to revise it and design new experiments.

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Procedure – Part I 1 Write a hypothesis using complete sentences on how the mass of a penny changes with age. 2 Sort Pennies.

a. Gather the pennies found at the back station and sort the pennies based on the year minted. b. Note the condition for each side of the coin. Place pennies minted prior to 1980 in one pile and pennies

minted between 1981 - 1983 in another pile and finally, pennies minted after 1983 in the third pile. Keep these piles until the end of the lab.

c. Find two pennies minted in each of the decade 1960's, 70's, 80's... and sort them in increasing order according to the year the coin was minted. You may team up with another group.

3 Determine the mass of various pennies.

a. Weigh each of the 10 pennies on the scale and record the mass in your datasheet from oldest to newest. Write the mass and the unit in your datasheet. Use the forceps (tweezers) when handling the pennies here after.

b. Record the detailed condition for each of the coin, citing the physical properties. c. Prepare a graph with proper labels and headings, including units. d. What are the axes labels in the graph? e. Take the information you collected and place a “o” for the mass of each penny with its year minted. Plot

class data as well if it is aggregated. 4 What conclusion can you draw from your data? Be sure to cite specific evidence from your data to support

your conclusion. Use the class result to provide an overall view of how the mass of a penny changes with age. Does your graph show a gradual increase or decrease of the mass with age? Is something unexpected found from the data?

5 Does your conclusion support your hypothesis? Answer this question in your data sheet.

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Experiment 2: Data Sheets Last Name________________________First_________________________ ____ / _____ pts


Part I

1 Hypothesis: Write a complete hypothesis on what you think happens to the mass of a penny with age? Use complete sentences.

2 Mass of Assorted Pennies Use and write the appropriate units after each data entry. Measure the mass to the precision of the scale. Measure the diameter and thickness to the hundredth of a centimeter.

# Year Mass (g)

Diameter (cm)

Thickness (cm)

Condition of coin

1 Oldest

2

3

4

5

6

7

8

9

10 Newest

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3 Graph of Mass of the penny vs year minted Mass of Penny vs. Year Minted

(Label the graph below, label proper units for mass)

3.8

3.6

3.4

3.2

3.0

2.8

2.6

2.4

2.2

2.0

1950 1960 1970 1980 1990 2000 2010 2020

4 What is your conclusion based on the graph generated from class data?

5 Is your hypothesis consistent with the graph or is there something unexpected that the graph shows?

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Procedure – Part II 1 Write a hypothesis about why the pennies’ mass changes that is consistent with the unexpected result found

in the first part of this lab.

2 Determine the Density of the Pennies Trade pennies with another group until each group has 25 pennies all of which are either older pennies (minted before 1982) or newer pennies (minted after 1983).

a. Take the stack of 25 pre-1982 pennies, place these on the scale and record the total mass of the 25-pennies. Repeat this for the newer pennies and record the mass.

b. Calculate the average mass of the pre and post pennies by dividing the total mass by 25. c. Fill a 100-ml graduated cylinder with water to the 20.0-ml mark. Use a dropper to get the meniscus

exactly to 20.0 mL. d. Carefully place 25 pre-1982 pennies into the cylinder. Avoid trapping air bubbles under the pennies or

splashing water up the sides of the cylinder. e. Accurately measure the total volume of the water and the pennies. Record this measurement in your

datasheet. f. Remove the pennies and thoroughly dry them with paper towel. Return the dry pennies back to their

original container. g. Calculate the average volume of a penny by dividing the volume of the pennies by 25. Repeat this for

the post-1982 pennies. h. Calculate the average density of pre 1982 and post 1982 pennies.

3 Physical Observations of Appearance a. Take 1 penny minted before 1981 and 1 penny minted after 1983 and use a metal cutter to cut these

pennies in half. (Obtain whichever one you don't have from another group). To save on pennies, your

instructor may demonstrate this part for you. b. Examine the interior of the two half-pennies and record your observations in your datasheet.

4 Chemical Properties of the Pennies a. Using a forceps take the half of each type and put them in a Petri dish. Place the Petri dish with the

pieces of pennies under the hood. b. Add a small amount of 1 M hydrochloric acid (CAUTION: This is a corrosive liquid) Observe any

difference in chemical reactivity between each of the halves for the two pennies. Such differences, if they occur, can be taken as evidence of differences in composition. To save on pennies, your

instructor may demonstrate this part for you. 5 Conclusions. What conclusion can you make base on your data? Use the evidence in this experiment to

support your conclusion. Write your answers in complete sentences. 6 Draw a flowchart of the scientific method you employed in this experiment. Be sure you are specific with

details in the flowchart.

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Example of generic scientific flowchart.

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Data Sheets - Part II Last Name________________________First_________________________


1. Hypothesis: Write a complete hypothesis as to why the weight of a penny changed after 1982. State in your hypothesis what you think happened to the mass of the penny after a certain year.

2. Density – Table of Measured Mass & Volume of Pennies (write proper units after each entry) Penny Type

Mass of 25 pennies (to precision of scale)

Vol initial (to tenth ml)

Vol final (to tenth ml)

Vol of 25 pennies

Pre- '82

1981 and earlier

Post - '82

1983 and later

State which group of pennies your group measured.

Show your calculations of the volume of the 25 pennies below.

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Calculation of Density - Avg. mass, volume and density of pre-'82 and post-'82 pennies. Write proper units after each entry and use correct number of significant figures.

Type of penny Average Volume Average Mass Density

pre – 1982

post – 1982

‘

Calculations: Show one sample calculation for the average volume, average mass and density below. Use correct number of significant figures and use correct units.

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3. Describe the inside of the pre-1982 and the post-1982 pennies.

(Use complete sentence in your description.)

4. Describe your observations of what occurs when the pre-1982 and the post-1982 cut pennies are exposed to 1 M HCl. (Use complete sentence in your description)

5. Based on your Observations, Experimental Data and Calculations, write a conclusion for Part II of this experiment. Be sure to cite your data directly to support your conclusion.

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6. Flowchart Create a flow chart that illustrates how the scientific method was used in Part I and Part II. Be sure to follow the scheme shown in the introduction section in this lab. Include details and label each box with its corresponding part of the scientific method. Include both Parts in your flowchart. When submitting reports in this course, only turn in pages that contains information that is to be graded. In other words, do not turn in the intro or procedure pages!

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Experiment 3: Measurements, the Metric System & Density Objective The purpose of experiment is to become familiar with the metric system by taking measurements using metric units. Additionally, the purpose of this experiment is to take measurements using various measuring devices and then report these measurements with the correct precision based on the apparatus used. Each measurement will be converted to other units using dimensional analysis with the answers reported to the correct number of significant figures.

Materials o Alcohol Thermometer o Metric Ruler o Metallic Cube Unknown o Brand new #2 Pencil o Lab Manual o 100 and 250 mL beakers o Empty 2-L soda bottle with cap o 10 mL, 25 mL, 100 mL – prefilled with colored liquid o 500 mL graduated cylinder

Introduction Additional Reading: Ch 2.1 – 2.7, Chemistry in our Lives, 5th Edition Timberlake . “Chemistry and Measurements.” For this experiment, you will be taking measurements with several different laboratory tools. Emphasis in grading will be placed on recording measurements with proper significant figures and with proper units.

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Procedure Part I – Metric Measurements & Estimates Measuring Temperature Temperature is measured in Celsius or Kelvin in a laboratory setting. The lines or graduation marks (Delta, D) on the thermometer represent 1˚C increments in temperature. Your recorded measurement of the temperature should estimate between the lines and carry with it two decimal places (hundredth) or precision to 0.1 ˚C.

Fill a 250 ml beaker 1/2 full (125 mL) with tap water and allow it to reach room temperature, about 30 minutes. Proceed with other parts of the laboratory while you wait for the temperature to equilibrate. Record the temperature of the water with the correct number of significant figures and units. Measuring Mass 1. Measure the mass of the following objects. You must record all significant figures from the balance. This

means measurements to 0.001 g. If the numbers on the scale fluctuate, ask people to move away from the scale and close the doors to the lab if they are open. Your recorded measurement is the midpoint of any fluctuations that continue.

a. Brand new #2 pencil b. 100 ml or 150 ml beaker c. 2-L soda bottle empty with cap

Estimate the mass of your lab manual. Do not measure it on the scale. If you do not have one, see instructor.

Measuring Distance 1. Measure the following lengths, in centimeters, using a meter stick, and record the measurements in your

datasheet with the correct number of significant figures and units. Record the measurement only to the precision that the instrument allows. On the meter stick the lines represent millimeters or 0.1 cm. Therefore, your measurements should be recorded to the 0.01 cm.

a. Your height in centimeters b. The length of a new pencil c. The circumference of a 2-L bottle

Estimate the width of your lab drawer. Do not measure. Measuring Volume For each of the 3 graduated cylinders, do the following procedure.

1. Determine the value of each increment in milliliters. 2. Read and record the volume in milliliters for each of the three

graduated cylinders found at the stations around the laboratory. Record the volume from the bottom of the meniscus.

a. 100 ml graduated cylinder (to nearest 0.1 ml) b. 25 ml graduated cylinder (to nearest 0.05 ml) c. 10 ml graduated cylinder (to nearest 0.01 ml)

Estimate the volume of liquid in the displayed test tube.

A meniscus as seen in a burette of colored water. '20.00 mL' is the correct depth measurement.

By PRHaney (Own work) [CC BY-SA 3.0 (https://creativecommons.org/licenses/by-sa/3.0)], via Wikimedia Commons By PRHaney (Own work)

25

Part II – Determination of Density & Identity of an Unknown object. Obtain an unknown cubic solid from the instructor laboratory table. Record the unknown number.

1. To determine the density by the direct measurement method, measure the mass using the laboratory scale.

Volume by direct measurement Measure the length, width, and height of the cube in centimeters. Calculate the volume by multiplying the length x width x height. Record your measurements in your datasheet

and show your calculations. Use correct significant figures with proper units. Calculate the density of the cube (Equation below)

𝑑𝑒𝑛𝑠𝑖𝑡𝑦 =𝑚𝑎𝑠𝑠𝑣𝑜𝑙𝑢𝑚𝑒

Volume by displacement Take a graduated cylinder and fill it half way to a convenient volume. Read the volume of the water to the precision of the graduated cylinder. Slowly slide the cube down the graduated cylinder being careful not to splash the water to the side. Measure the new displaced volume to the precision of the graduated cylinder. Using the mass from step1, calculate the density using this volume measured by displacement. Record your

work and data in your datasheet. For the Post Lab Questions Identify your unknown based on the experimental density from a list of possible metals.

Table of Possible Unknowns Material Density (g/mL) ALUMINUM 2.70 BRASS 8.70 COPPER 8.90 STEEL 7.60 IRON 7.86

Calculate the percent error from your experimental density from the listed density of the metal you have.

%𝑒𝑟𝑟𝑜𝑟 = V𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 − 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙

𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙V 𝑥100

27



Part I – Metric Measurements & Estimates

Temperature magnitude units

Temperature of water

Mass magnitude units Brand new #2 pencil

100 mL beaker

2-L soda bottle empty w/ cap

Lab Manual ** estimate**

Length magnitude units Your height

Brand new #2 pencil

Circumference of 2-L bottle

Lab Drawer width **estimate**

Volume magnitude units 10 ml graduated cylinder

25 ml graduated cylinder

100 ml graduated cylinder

Displayed test tube **estimate**

Color of Liquid

28

Part II – Determination of Density & Identity of an Unknown object.

Unknown Number of Cube __________

Mass of Cube

Density Calculation Show calculations below. Use correct S.F. and units.

Direct Volume Magnitude Units

Length

Width

Height

Volume

Density

Displacement Magnitude Units Density Calculation Show calculations below. Use correct S.F. and units.

Volume initial

Volume Final

Volume Solid

Density

29

Post Lab Questions. 1. Explain why the two density values from your calculations might be different. Inspect the number of

significant figures for each measurement and speculate which method for the determination of the volume of the cube may be more precise.

2. Based on the measured densities, identify the metal that your cube. Use the density that is the most precise from the metals listed in the introduction to this lab.

________Unknown # of Cube

________Identity of Unknown

Calculate the percent error in your measurement. Show your calculations below.

_________ % Error

Mass Conversions. Do the following conversions for the mass of the pencil. Set up using dimensional analysis.

measurement from part I

Conversion factor: write values in blanks ______mg = ______g

Conversion factor: write values in blanks ______g = _____kg

Conversion factor: write values in blanks ______lb = _____g

milligrams kilograms pounds

Pencil ___________

Beaker ___________

Show your calculations below for the pencil with proper units cancelled. Set up using dimensional analysis and circle your final answer.

30

Length Conversions. Do the following conversions for the length of the pencil.


Conversion factor: write values in blank

______mm = ______cm


______m = _____cm


______in = _____cm

millimeters meters inches

Pencil ___________

Circumference

2L bottle ___________

Show your calculations below for the pencil with proper units cancelled. Set up using dimensional analysis and circle your final answer. Volume Conversions. Do the following conversions for the mas of the pencil.



______mL = ______L

Conversion factor:

write values in blank ______mL = _____gal

Liters gallons 10 mL graduated cylinder ___________

100 mL graduated cylinder

______________

Show your calculations for the 100 mL graduated cylinder with proper units cancelled. Set up using dimensional analysis and circle your final answer.

31

Experiment: Energy of Food Objective The goal of this experiment is to determine the energy content in two types of food, a nut and a snack food. Two trials will be performed on each food, for a better statistical sampling of data. The food will undergo combustion, and the heat will be captured and quantified by monitoring the melting of ice.

Materials o Utility Clamp o Stirring Rod o Small soda can calorimeter o Food holder o Lighter with wand

o 100-mL graduated cylinder o Ring stand o Tongs o Food Samples

Introduction Reading: Ch 3.4 – Ch 3.7, Chemistry in our Lives Timberlake . “Matter & Energy.” Pay special attention to the sections on food and nutrition and using the Heat of Fusion, which you will need to calculate the energy transfer from chemical potential energy to heat energy.

Heat of Fusion In this experiment, you will ignite a nut and another type of snack food, and you will perform 2 trials of each. The heat of the combustion of the food (qfood) will be captured to melt ice (qice) in a soda can calorimeter.

𝑞\]]^ = −𝑞_`a

The phase transition between solid ice to liquid water will be quantified by measuring the mass of the ice that has melted. The heat lost from chemical potential energy, becomes heat during combustion of the food is transferred to the ice in a soda can calorimeter. The heat of fusion for ice is 8.0 x 10 cal/g of H2O. Therefore, the heat or energy content in the food can be calculated in the following way.

𝑞_`a = 80.0𝑐𝑎𝑙𝑔𝑥𝑔𝑟𝑎𝑚𝑠𝑜𝑓𝑖𝑐𝑒𝑚𝑒𝑙𝑡𝑒𝑑

During a phase transition, the temperature of the matter does NOT change. Therefore, we do NOT need to account for heat using the equation q=mCDT because the temperature of the water in the soda can calorimeter remain at a constant 0˚C, as long as there is still ice in the can at the end of the experiment. It is important to make sure that soot has been removed from the bottom of the can because it can insulate the can and prevent heat transfer to the ice. It is important to burn the food for at least a minute so a measurable amount of food is converted to heat. The flame from the food must come in contact with the can calorimeter.

Nutritional Content in Different Foods The nutritional content of food varies depending upon its content of carbohydrate, fat and protein. Fats have a higher calorie content per gram than protein and carbohydrates (Table 1). After your trials with the food, you’ll calculate an average kcal per gram of food and then compare it to the information on the nutritional label on the package.

Table 1 Food Nutrient kcal/g kJ/g Carbohydrate 4 17 Fat 9 38 Protein 4 17 Water 0 -

32

Procedure – Set up of the Calorimeter 1. Obtain a metal can with the top cut off, and the edges folded back so the edges

are safe. This apparatus will serve as the calorimeter. If necessary, wipe any carbon deposit (soot) off the bottom of the can using a moist paper towel. The soot can insulate the calorimeter so that the heat transfer from the food to the ice is not very efficient.

2. Set up the apparatus as shown in the front desk by the instructor, shown in

figure. Use a large adjustable O-ring clamp to hold the metal can. Position the food holder (for holding the nut) at the base of the ring stand, below the O-ring holding the aluminum can calorimeter.

3. Select nuts that are at least 1-gram. Use two pieces if the sample size is too

small. If you do not burn enough material, the heat generated may not be enough to melt the ice.

4. Weigh the mass of the holder and record this in the datasheet (1).

5. Add the nut (or nut pieces) on to the holder and weigh the food holder and nuts. Record this in your

datasheet (2).

6. Before beginning your trial, record the mass of your empty 100-mL graduated cylinder (6).

Combustion of the food sample 7. Add crushed-ice to the can so it is at least three-quarter full. Pack the ice to the bottom of the metal-can-

calorimeter using the bottom of a 150-mL beaker. Be sure not to press so hard as to break the beaker. CAUTION: Use hot mitts or gloves to protect your hands in this procedure so that if the beaker does crack, your hands are not cut.

8. Remove the food and stand from under the can and use the lighter to ignite the food. Once the food is

burning, position it so that the top of the flame just touches the bottom of the can. The food should burn for a minimum of 1 minute. Let the food burn as long as it has flame. If the flame goes out before the nut is completely burned, move the food holder away from the calorimeter to re-ignite the nut, then replace it under the calorimeter (only when the nut is burning again should you place it under the calorimeter). CAUTION: The can and ring clamp will be hot. Remember cool object and hot objects look exactly the same so be cautious when handling the hardware.

9. Record the mass of the food sample and holder after it has finished burning in the datasheet (3). 10. Calculate the mass of the food that has burned and record it in the datasheet (4).

33

11. Being careful not to burn your hands, pour the melted ice out of the soda can into the pre-weighed graduated cylinder. You may use a funnel for this. Do not allow solid ice to fall into the graduated cylinder. Determine the volume of the water and record it in your notebook (5).

12. Measure and record the mass of the water and graduated cylinder directly (7), and subtract your previously

recorded graduated cylinder mass (6) to determine the mass of the melted water (8) (Show your calculation).

13. Clean any soot from the can with a moist paper towel before proceeding with the next trial.

14. Repeat the procedure with the same type of nut for trial 2. After the second trial for the nut, repeat this procedure for 2 trials of one of the snack foods.

Calculation of Heat per gram of food. 15. The total heat produced from burning the food (9) is calculated using the Heat of Fusion of 80.0 calories per

gram of ice melted. Therefore, the calculation is as follows:

𝑞_`a = 80.0𝑐𝑎𝑙𝑔× 𝑔𝑟𝑎𝑚𝑠𝑜𝑓𝑖𝑐𝑒𝑚𝑒𝑙𝑡𝑒𝑑(𝑟𝑜𝑤8)

𝑞_`a = −𝑞\]]^

16. Heat per gram of food burned is then calculated by taking the above total heat (9) produced and dividing by

the mass of the food that burned (4).

𝐻𝑒𝑎𝑡𝑝𝑒𝑟𝑔𝑟𝑎𝑚 =𝑞\]]^

𝑀𝑎𝑠𝑠𝑜𝑓𝑓𝑜𝑜𝑑𝑏𝑢𝑟𝑛𝑒𝑑(𝑟𝑜𝑤4)

17. For row 11, convert the heat per gram from calories per gram to kilocalories per gram.

18. For row 12 in the table, average your two trials of the nut into one number. And average your two trials of the same snack food.

Summary of Results - Comparison to Snack Food Label. 19. Transfer your answer for Row 12 to the Summary of results for the nut and for the snack food.

Remember that 1 Calorie = 1 kilocalorie.

20. Using the nutritional label on the packaging for the nut and the snack food, calculate the Cal/g for each.

21. Calculate the percentage error from the label for each. The theoretical value is the Cal/gram calculated from the nutritional label.

%𝑒𝑟𝑟𝑜𝑟 = V𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙𝑣𝑎𝑙𝑢𝑒 − 𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙𝑣𝑎𝑙𝑢𝑒

𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙𝑣𝑎𝑙𝑢𝑒V × 100

34

Heating Cooling Curve for Water

35



Mass of Burned Food and Mass of Melted Ice :

Sample Nut ________________ Snack ___________

Trial 1 Trial 2 Trial 1 Trial 2

1 Initial mass of food holder

_________g

_________g

_________g

_________g

2 Initial mass of food and holder

_________g

_________g

_________g

_________g

3 Remaining mass of food and holder

_________g

_________g

_________g

_________g

4 Mass of food burned

_________g

_________g

_________g

_________g

5 Volume of water melted

________mL

________mL

________mL

________mL

6 Mass of empty graduated cylinder

_________g

_________g

_________g

_________g

7 Mass of graduated cylinder + water melted

_________g

_________g

_________g

_________g

8 Mass of water melted

_________g

_________g

_________g

_________g

Show calculations for 1 sample with proper labels, significant figures and units. Circle answers.

36

Calculation of Heat per gram of food: 9 Heat produced for total food

ignited:

Write correct units

______

______

______

______ 10 Heat produced by 1 gram of food

(cal/g) ignited:

Write correct units

______

______

______

______

11 Heat produced by 1 gram food (kcal/g) ignited:

Write correct units

______

______

______

______

12 Average heat by 1 gram nut & food sample (kcal/g)

Write correct units.

____________

____________


Summary of Results: Comparison to Food Nutritional Labeling and Percent Error

Nut

q Nut Nut

(Cal/g)

Calorie from package

(Cal / g) % error Snack Food

q food snack

(Cal/g)

Calorie from package

(Cal / g) % error


37

Postlab Questions: 1. The heat of the combusting food is transferred to the contents of the can and melts the ice. If you do not

have enough ice in the can for a trial, it will all melt before the flame extinguishes itself. Once the ice is

melted, what does the excess heat of the burning do? How can you quantify this heat?

2. If all of the ice melts when you were doing this experiment so at the end of the procedure you have a calorimeter full of water, how would this affect your result? Would you calculate a caloric value of your food

to be higher or lower than the true value?

3. What was your percent error for your nut and food sample in terms of caloric value? Explain why your result might be higher or lower than the nutritional label (manufacture label).

39

Exercise 6A: Naming Compounds Last Name________________________First_______________________ _____ / __ pts


1. Naming Type I Ionic Compounds – One metal and a nonmetal

Rb2O Al(HSO4)3

Ca(HCO3)2 Sr3N2

Cd3(PO3)2

2. Naming Type II Ionic Compounds – Containing Variable-Charge Metals

Cu(C2H3O2)2 Hg2S

Co2(CO3)3 CuO

Pb(ClO3)4

3. Naming Type III - Binary Compounds with 2 Nonmetals

N2O2 IF4

SCl2 SF6

SiO2

4. Name the following Mixture of Different Types of Compounds

NaHCO3 NI3

Co2(HPO3)3 NH4ClO4

Pt(CN)2

40

5. Formulas of Type I Ionic Compounds - One metal and a nonmetal

Magnesium chloride

Barium nitrate

Silver sulfite

Aluminum arsenite

Beryllium hydroxide

6. Formulas of Type II Ionic Compounds - Containing Variable Charge Metals

Vanadium(III) nitride

Chromium(III) bisulfite

Molybdenum(VI) oxide

Copper(II) dichromate

Platinum(IV) perbromate

7. Formulas of Type III Compounds - Binary Compounds with 2 Nonmetals

Dinitrogen tetrafluoride

Sulfur hexachloride

Diboron trioxide

Dihydrogen monoxide

Tetraphosphorous heptaoxide _______________________

8. Write the formulas for the following List of Different Types of Compounds.

Calcium carbonate

Ammonium selenide

Cadmium permanganate

Thallium(I) telluride

Tungsten(V) oxalate

41

Exercise 6B: Lewis Structures & VSEPR Last Name________________________First_________________________ ____ / _____ pts

Lab Partner(s): _______________________________________________________Day _________𝒂𝒎.𝒑𝒎.

Valence Electrons & Group Number Element N O P Ca Cl Ga Group Number

# of Valence electrons

Lewis structures: http://faculty.sdmiramar.edu/fgarces/ChemComon/Tutorial/Lewis/LewisTutorial/LewisTutorial.htm

Compound, Chemical formula

Octet/Duet Electrons

Total Valence Electrons

# of Bonds in Structure

# of Nonbonded

Electron pairs on all

central atoms

Lewis Structure

Ammonia, NH3

Fluoromethane, CH3F

Carbon disulfide

42


Octet / Duet Electrons

Total Valence Electrons

# of Bonds in Structure

# of Nonbonded

Electron pairs on all

central atoms

Lewis Structure

Formic acid, HCOOH

Hydroxylamine, NH2OH

Nitrate ion, N𝑂9;7

Sulfur dioxide

Acetate Ion CH3COO-1

⎣⎢⎢⎢⎢⎡ HOH CC O

H ⎦⎥⎥⎥⎥⎤−1

H C

O

O H

H N

H

O H

43

Shapes of Molecules VSEPR Theory. Go to appendix to see VSEPR Table http://faculty.sdmiramar.edu/fgarces/ChemComon/Tutorial/VSEPR/VSEPRTutorial/VSEPRTutorial.htm


# of Valence Electrons Lewis Structure # of Bonded

Atoms Lone Pair on Central

Atom

(1) Electronic Geometry, AE (2) Molecular Geometry, ABE (3) Bond angles (4) Hybridization

OF2

20 4 2

1. Tetrahedral, AE4

2. Bent-Tetrahedral, AE2B2

3. <109.5˚

4. sp3

NH3

1.

2.

3.

4.

CH3F

1.

2.

3.

4.

CS2

1.

2.

3.

4.

44

Shapes of Molecules VSEPR Theory Compound,

Chemical formula

# of Valence

Electrons Lewis Structure # of Bonded

Atoms Lone Pair on Central

Atom

(1) Electronic Geometry, AE (2) Molecular Geometry, ABE (3) Bond angles (4) Hybridization

SiO2

1.

2.

3.

4.

NO9;

1.

2.

3.

4.

NO3;

1.

2.

3.

4.

PBr3

1.

2.

3.

4.

SO2

1.

2.

3.

4.

45

Experiment: Separation of a Ternary Mixture Objective In this experiment, students create a mixture of 3 substances, iodine, sand and cobalt (II) chloride and then separate them using a variety of separation methods. The goal of this experiment is to familiarize students with the methods of separation, identify the isolated substance at each step and also to identify a physical and chemical changes that may occur during the separation of the mixture.

Materials o Evaporating dish o Hot plate o Beakers

o Wash bottle o Deionized water o Iodine crystals

o Sand o Cobalt(II) chloride

hexahydrate (CoCl2•6H2O) Introduction The three substances that will be components in our heterogeneous mixture have some unique chemical and physical properties that will be exploited in the process of isolating each. Here are some relevant definitions for this experiment:

Sublimation - the phase change associated with the transition of a substance from a solid to a gas without passing through a liquid stage. One compound that sublimates at atmospheric pressure is CO2, also known as dry ice. This is a physical change of matter.

Deposition - the phase change associated with the transition of a substance from a gas to a solid without passing through a liquid stage. This is a physical change of matter.

Solubility - the ability of a substance to dissolve in water, creating an aqueous solution. When a substance is dissolved in water, its physical state is termed aqueous and is denoted (aq). It is NOT in a liquid state. This is a physical property of matter.

Density – Density is mass divided by volume (g/ml), and it impacts the ability of a substance to sink or float in a liquid or gas. This is a physical property of matter.

Volatility – The tendency or ability of a substance to enter the vapor phase. This is a physical property of matter.

Miscibility – The ability of two liquids to mix homogeneously in all proportions. Water and ethanol are miscible. This is a physical property of matter.

Isolated – A substance is isolated when it has been removed from all other substances. Sugar dissolved in water has NOT been isolated from the water.

Decantation – The technique of decantation exploits the difference in solubility and density between two or more substances for separation. In a heterogeneous aqueous mixture, where one substance is insoluble and resting on the bottom of the container, the liquid is poured out of the container without transferring the insoluble dense substance, which rests on the bottom. This is a laboratory technique.

Hexahydrate – In this experiment, cobalt (II) chloride is seen in two forms, the anhydrous form (CoCl2) and the hydrated form (CoCl2•6H2O). Both of these forms of cobalt (II)chloride are solids, but they have a very noticeable difference in physical properties. Hexahydrate means that there are six waters in the basic formula of the substance

46

Procedure – Observation of the Starting material. Obtain three substances, iodine (I2), sand (mostly SiO2), and cobalt (II) chloride hexahydrate (CoCl2•6H2O), and place these in three small test tubes. In general, measure out two crystals of iodine about 1- 2 mm in size (tip of spatula). Use an amount of cobalt(II) chloride about 5mm diameter. The amount of sand should be the size of a nickel (1/4 teaspoon). Write a description of the appearance of each substance individually. At each stage of the procedure be sure to identify the isolated species and the physical property that facilitated the isolation. Record by circling and underlining in the Observations Table.

1. Create the mixture. Thoroughly mix together the three substances in a 250 ml beaker. Write a description of the appearance of the mixture in your datasheet in Row 1.

2. Sublimation of Substance A. Place crushed ice in an evaporating dish over the 250ml beaker. Place the beaker with the evaporating dish on a warm hot plate at 3/4-heat setting and observe the mixture while it is heated. Continue to heat the beaker until there is no more signs of gas in the beaker. If you still see gas, in the beaker, you are not done. It must be clear and colorless before you can proceed to the next step. Observe what happens to all three components of the mixture. In order to make careful observation, it is best to follow either iodine or the cobalt substance during the heating process. Inspect the underside of the evaporating dish during this step. Observations. When nothing else appears to be happening, remove the beaker from the hot plate. Be careful that the ice does not melt to the point that some of the ice leaks down into the beaker. Allow the beaker to cool for 2-5 minutes before handling it. Use hot mitts to remove evaporating dish from the 250-mL beaker. Pour the water and ice from the evaporating dish to a 50mL beaker. Place a paper towel on your lab bench top and then invert the evaporating dish and place on the paper towel. What is the identity of the residue that accumulated at the bottom of the evaporating dish? Which substance has been isolated? Record your answer and observation on your datasheet (in row 2 if the data sheet table).

3. Dissolving substance C. Record observations of the remaining mixture. Add approximately 15 mL of water to the 250mL beaker and stir. What is happening? Record the color of the solution and solid inside this beaker. Write your observation on your datasheet (row 3) Note, that if the color of the solution is other than lavender, i.e., brown or black, then you stopped the heating in part 2 prematurely. Your instructor may ask you to repeat the procedure with smaller quantities of iodine and cobalt(II) chloride hexahydrate.

4. Isolated substance B via decantation of C Decant the liquid from the solid in the 250-mL beaker into a 400mL beaker. Add an additional 15mL back to the solid residue that remains in the 250mL beaker in order to wash the solid. Decant the rinse into the 400mL beaker. What substance is separated in this step? What color is the solid component? Write your observations in row 4.

47

5. Evaporating water to recover substance C. Put the content of the 400 mL beaker on a hot plate. Heat until all the liquid has evaporated. Observe and record what happens as the liquid evaporates. Pay close attention to the color change as the last drop of liquid evaporates away. What is the color of the remaining solid? As soon as the last bit of liquid evaporates turn off the hot plate and allow the 400mL beaker to cool for a minute or two. After recording your observation, add a small drop of water to the solid, what do you observe? Record all observations from this step in Row 5.

6. Cleaning Up. Dispose of substances in waste container and wash all your glassware. Wipe down your

station. Wash your hands thoroughly with soap and water before leaving the lab room.

Flowchart of Separation Scheme I

49

Experiment 7: Data Sheets Last Name________________________First_________________________ Day ___________ ____ / _____ pts

Lab Partner(s): ________________________________________________________________ .

Observations Table

Steps

Observations: • Write down what is seen in each procedure. • Be as detailed as possible. • State what component is being isolated.

Basis of separation: Circle the basis of separation.

[Solubility, density, volatility, miscibility, none] Underline the isolated substance (or none).

1

*Circle the correct choice [Solubility-density-volatility-miscibility-none]

-Underline the correct choice a) Isolated I2 b) Isolated CoCl2• 6H2O c) Isolated CoCl2 d) Isolated Sand e) Isolated H2O f) None

2

Circle the correct choice

[Solubility-density-volatility-miscibility-none] -Underline the correct choice a) Isolated I2 b) Isolated CoCl2• 6H2O c) Isolated CoCl2 d) Isolated Sand e) Isolated H2O f) None

3



4



5

Circle the correct choice [Solubility-density-volatility-miscibility-none]


50

Flowchart Questions The flow chart has detailed steps labeled i – x. Answer the following questions concerning procedures labeled i – x and support your explanation using your experimental observations. For example, you may need to describe the color of the gas generated in step 2, question iii. If you are naming a chemical, be sure you spell-out the correct name and do not abbreviate when referring to the chemical.

i) What are the three substances that are mixed? ii) Which chemical is in the majority in the mixture? iii) What is the color of the gas? iv) What color is the bottom of the dish? What is the identity of the chemical isolated? v) What substance is added in order to isolate substance B from C? vi) What is the color of the aqueous solution? vii) What techniques are used to isolate the solid from the liquid in this step? viii) What is the color of the solid residue? What is the chemical identity of this residue? ix) What chemical is removed in order to isolate chemical C? x) What is the color of the solid residue? What is the chemical identity of this residue?

51

Additional Questions 3. Describe one separation technique that you learned about in this experiment and how it was used.

4. In this experiment, physical properties were exploited in the separation of I2, sand and CoCl2. In the separation methods, physical changes were observed with these substances. It is NOT possible to separate cobalt from chloride in CoCl2 using physical changes. What kind of change would be required to separate cobalt from chloride?

5. How does the addition of water change the physical properties of the cobalt(II) chloride substance?

52

Post lab Questions Answer these post lab questions and turn in to your instructor before you leave lab.

1. What is the chemical identity of the purple gas when substance A, B and C are all heated? Justify your answer by citing the evidence as written in your observations in the datasheet.

2. Which one of the three substances do we know is not soluble in water? Justify your answer by citing the evidence as written in your observations in the datasheet.

3. What is the chemical identity of the powder blue solid in the last step? What color does it change to when water is added? Your answer must be unambiguous, so use detailed chemical formulas, and the terms anhydrous and hydrate. Justify your answer by citing the evidence as written in your observations in the datasheet.

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Exercise 8A: Balancing Chemical Equations Last Name________________________First_________________________ ___ / __ pts

Lab Partner(s): ________________________________________________________ Day _________𝒂𝒎.𝒑𝒎.

1. Complete & balance the following double displacement reaction equations:

___NaCl (aq) + ___AgNO3 (aq) à

___NaOH (aq) + ___HCl (aq) à

__BaCl2 (aq) + ___H2SO4 (aq) à

b. Complete, balance & Identify the type of reaction. Product names are named in parentheses (Combination, Decomposition, Single displacement, Double displacement)

___ Zn(s) + ___ Cl2 (g) à

(zinc chloride)

____ H2CO3(aq) à

(carbon dioxide and water)

___ Fe2S3 (aq) + ___ HCN (g) à

(Iron(III)cyanide and dihydrogen monosulfide)

_____ Ni (s) + ____ HCl (aq) à

(nickel(II)chloride and hydrogen)

___ BaO (aq) + __ CO2 (g) à

(barium carbonate)

_____ Co(OH)3 (aq) + ___ Na2CO3 (aq) à

(cobalt(III) carbonate and sodium hydroxide)

___ ZnCl2 (aq) + ____ (NH4)2S (aq) à

(zinc sulfide and ammonium chloride)

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c. Beneath each word equation, write the formula equation and balance the reaction.

Sulfur dioxide(g) + Oxygen(g) Sulfur trioxide(g)

Ammonia(g) + oxygen(g) water(g) + nitrogen monoxide(g)

Ammonium nitrate(s) Nitrogen(g) + Oxygen(g) + Water(g)

Sodium bicarbonate(s) + Acetic acid(aq) à sodium acetate (aq) + carbon dioxide(g) + Water(g)

Sodium bicarbonate(s) sodium carbonate(s) + carbon dioxide(g) + Water(l)

nitric acid(g) + nitrogen monoxide(g) à nitrogen dioxide(g) + Water(l)

€

Δ# → #

€

Δ# → #

€

Δ# → #

€

Δ# → #

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Exercise 8B: Stoichiometry Exercise Last Name________________________First_________________________ ___ / __ pts

Lab Partner(s): ______________________________________________________________ Day _________𝒂𝒎.𝒑𝒎.

1. Molar mass Find or calculate the atomic mass, molar mass, or formula mass for the following: (Include the appropriate unit and round off to the hundredth of a unit, i.e., Atomic Weight of Hydrogen is shown as 1.008 amu and is rounded off as 1.01 amu)

a) Lead

b) Phosphorus

c) Sulfur hexafluoride

d) Sucrose (C12H22O11)

2. Atom, molecules, mole, and Avogadro's Number- a) How many atoms are in 11 molecules of sulfur hexafluoride?

b) What is the mass in grams (g) of a single atom of Phosphorus? Hint: Start with the atomic mass or molar mass of phosphorus and use dimensional analysis to get grams per atom.

c) What is the mass (in grams) of 5.0 moles of sucrose?

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3. Balancing Equations and the Mole Concept a) Write and balance the chemical equation and then determine the moles of potassium chlorate (KClO3) produced, if

17 moles of potassium chloride is combined with excess oxygen?

b) Write and balance the chemical equation and then determine the moles of H2O produced from combustion of 5 moles of sucrose and 39 moles of oxygen? Which chemical is in excess and how many moles remain after the reaction is complete?

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c) How many moles of S8 are needed to produce 4.75 kg of sulfur tetrafluoride according to the reaction:

S8 + F2 à sulfur tetrafluoride. (Hint: First complete the equation and balance if necessary)

d) What is the percent yield for a reaction when 166.3 grams of W2O3 combines with excess carbon monoxide to produce 14.78 g of W ?

___W2O3 + ____ CO à ____ W + _____ CO2

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e) Carbon dioxide, CO2, and ammonia, NH3, combines together to form urea, CH4N2O, plus water. Write a balanced equation and calculate the mass of ammonia (in grams) that would be needed to make 2.0 moles of urea.

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Experiment: Chemical Reactions & Chemical Equations Objective In this experiment, students perform a variety of chemical reactions. For each reaction, student identify the signs that a reaction has occurred, write the balanced chemical equation with appropriate phases and classify the reaction.

Materials o Copper (BB bullets) o Calcium oxide (limewater) o Baking soda (NaHCO3) o Muriatic acid (HCl) o Deionized water o Sugar (sucrose) o Magnesium

o Acetic acid (vinegar) o Small beaker o Spoon o Forceps o Ethanol o Straw o Test tube

Introduction Reading: Chemistry in our Lives Timberlake . Chapter 7.1 Equations for Chemical Reactions, 7.2 Types of Chemical Reactions, and 9.2 Electrolytes and Non-Electrolytes.

In this experiment, you’ll be observing the signs of chemical reactions. These include the following:

1. Color Change 2. Formation of Bubbles, indicates that one of the products is a gas. Ex. H2, O2 or CO2 3. Flame 4. A change in temperature of the reaction mixture 5. Formation of a precipitate.

Write in complete sentence your full observations for each reaction you complete.

Write the balance equation for each reaction and include the phases.

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Procedure – Reaction A: Oxidation of Copper by Air (O2) Add copper shots to a crucible so that it barely covers the bottom. Position the open crucible in a clay triangle supported by an O-ring clamp attached to a ring stand. Heat the bottom of the crucible with a Bunsen burner until the bottom glows red. Turn off the burner and allow the content to cool. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. Empty all chemicals in the proper waste container after you finish.

Reaction B: Combination of CaO with CO2 Pour limewater in to a 250 mL beaker. Using a straw, blow into the solution. Observe what happens to the solution as you continue to blow bubbles. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. Empty all chemicals in the proper waste container after you finish.

Reaction C: Decomposition of H2O Using alligator clips, connect the red wire to the positive terminal of a 9-V battery and the black terminal to the negative terminal. Take care not to touch the bare ends of the terminal. Connect the other ends of the alligator clips to graphite rods. Take a glass of deionized water and add about 1 gram of sodium chloride (about a spatula tip). Place both graphite rods into the water, taking care that they do not touch, and record your observations. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. When done making your observations, return all equipment clean and dry to its original storage location.

Reaction D: Combustion of Sucrose (C12H22O11) Carry out this procedure under the hood. A deflagration spoon is used in this procedure as shown in the figure. Place a teaspoon (~5 grams) of sugar in the deflagration spoon and then heat over a Bunsen burner. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. Empty all chemicals in the proper waste container, clean the deflagration spoon and return to its original location.

Reaction E: Combustion of ethanol (C2H5OH) with oxygen (O2). Carry out this procedure under the hood. A deflagration spoon is used in this procedure as shown in the figure below. Add a small amount of ethanol (~0.5 ml) to a deflagration spoon. Use a Bunsen burner to burn the ethanol. Only a small amount of ethanol is necessary for safety precautions. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. Empty all chemicals in the proper waste container when you finish and return all equipment clean back to its original location.

Reaction F: Magnesium Ribbon with Muriatic Acid, HCl Add 2ml of Muriatic acid (HCl) to an empty test tube. Add a 1-cm strip magnesium ribbon to the test tube using a forceps. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction.

Additional Reaction: Fill a clean test tube with 2mL of deionized water. Place the test tube in a warm bath for 5 min. Add two drops of phenolphthalein indicator to the water. Phenolphthalein is an indicator that changes color (pink) in the presence of hydroxides.* Add a 1-cm strip of magnesium to the test tube and wait 15 minutes for the reaction to take place. After 15 minutes record in your datasheet your observation. Is there the presence of hydroxide formation? Write the chemical reactions that occur in your datasheet.

* In general, alkaloid metals will react with water to form metal hydroxides and hydrogen gas.

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Reaction G: Baking soda with Vinegar (CH3COOH). Halfway fill an evaporating dish with vinegar. Sprinkle baking soda on to the vinegar and record your observations. Write on your datasheet the balanced chemical equation, the type of reaction and the driving force for the reaction. Empty all chemicals in the proper waste container after you finish.

Hint: See Exercise 8A, if you need help writing this equation.

Clean up Wash and dry all your glassware equipment. Dispose of all used chemicals in the proper waste container. Wipe down your station and place your glassware and equipment back in its proper place. Unless your equipment is back in the locker drawer, you are not allowed to remove your safety goggles.

Wash your hands thoroughly with soap and water before leaving your work area.

Example of Complete, Total and Net Ionic Equation:

The equation below is the reaction between Mn(NO3)2 (aq) and K2S (aq). To start both reactant dissolve in water as indicated by the aqueous phase. These combine to form the MnS (s) precipitate. The KNO3 (aq) (other product) remains dissolve in solution as indicated by the aqueous phase. The reaction can be summarized by three equations, the molecular equation, the complete ionic equation and the net ionic equation.

Summary:

a) Molecular equation:

Mn(NO3)2 (aq) + K2S (aq) 2KNO3 (aq) + 2MnS (s)

b) Complete ionic equation:

Mn+2(aq) + 2NO3-(aq) + 2 K+(aq) + S-2 (aq) 2 K+ (aq) + 2NO3- (aq) + MnS (s)

c) Net ionic equation:

Mn+2 (aq) + S-2 (aq) MnS (s)

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Lab Partner(s): ________________________________________________________ Day _________𝒂𝒎.𝒑𝒎.

Reaction A: Oxidation of Copper by Air (O2) to produce copper (II) oxide

Appearance Before

Appearance After

Observation and Evidence of Reaction

Balanced Chemical Equation

Reaction Type (circle one)

Combination - Decomposition - Single Displacement - Combustion - Double Displacement

Reaction B: Combination of CaO with CO2 to produce calcium carbonate

Appearance Before

Appearance After





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Reaction C: Decomposition of H2O into hydrogen and oxygen gas

Appearance Before

Appearance After


Balanced Chemical Equation **



** NaCl is used to help the battery to conduct electrons, e-, and is not part of the chemical equation. Reaction D: Combustion of Sucrose (C12H22O11)

Appearance Before

Appearance After





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Reaction E: Combustion of Ethanol (C2H5OH) with oxygen (O2).

Appearance Before

Appearance After





Reaction F: Magnesium Ribbon with Muriatic Acid, HCl to produce hydrogen and magnesium chloride

Appearance Before

Appearance After


**Balanced Molecular Chemical Equation

Total Ionic Equation

Net Ionic Equation

Reaction Type (circle one) Combination - Decomposition - Single Displacement - Combustion - Double Displacement

** Your instructor will go over how to write a molecular, complete ionic and net ionic equation.

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Reaction G: Baking soda with Vinegar (CH3COOH).

Appearance Before

Appearance After


**Balanced Molecular Chemical Equation

Total Ionic Equation

Net Ionic Equation

Reaction Type (circle one) Combination - Decomposition - Single Displacement - Combustion - Double Displacement

** Your instructor will go over how to write a molecular, complete ionic and net ionic equation.

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Postlab Questions 1. Check all the reactions in this lab, which produced a precipitate.

� Oxidation of Copper � Decomposition of H2O � Reaction Mg ribbon with HCl � Baking soda and Vinegar � Combination of CaO and CO2 � Decomposition of Sugar � Ethanol with Oxygen

2. Check all reactions in this lab that showed evidence of the release or absorption of heat. (Do not confuse the release of heat with providing an external addition of heat, i.e. Heating with a Bunsen burner)


3. Check all the reactions in this lab, which were accompanied by the evolution of gas.


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4. What is the color of the sugar before and after combustion? What is the identity of the gases you observed?

5. What is the gas formed in the reaction between baking soda and vinegar?

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Experiment: The Mole Last Name________________________First_________________________


Objective In this experiment, students use nuts (N) and bolts (B) as model elements to build model compounds using these elements. Calculations are performed to identify unknown compounds and to quantify a known compound. These calculations mirror the process for using molar mass to calculate the number of moles of a substance. This process is then applied to a mixture of lentil beans and to some common substances. The goal is to obtain a deeper functional understanding of the mole concept and its use. Materials

In Bin o Bag of screws o Bag of Lentils & Beans o Bag of bolts o Unknowns in conical vials

In Room o Alloy cube o Sugar packet o Distilled water o Balance

Introduction Required additional reading: Ch 7.4 “the Mole” and 7.5 “Molar Mass” in Chemistry in our Lives Timberlake 5th Edition.

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Procedure – Part I. Establishing the mass of individual elements and molecules. These are found in the clear container. Open the container to remove the 10 HexNuts and 10 Bolts.

A. Element #1 (the HexNut) and Element #2 (the Bolt) Determine and record the following information: i) The average mass of a single HexNut to the precision of the scale. The best way to measure the average mass of a HexNut is to weigh 10 HexNuts and divide the total mass by 10. Show your work here. ii) The average mass of a single Bolt to the precision of the scale. As mentioned above, the best way to measure the average mass of a bolt is to weigh 10 bolts and divide the total mass by 10. Show your work here.

Ai)__________g Aii)__________g

Calculations:

B. Compounds #1 (BN), #2 (BN2), #3 (BN3) and #4(BN4): Bolt – HexNut molecules The combined Bolt (B) and HexNut (N) "elements" will represent our theoretical "compound" and will be represented by BN, BN2, BN3 and BN4. Weigh the container with the compound. Subtract the mass of the container (written on the container), calculate and record the average mass of the following Bolt-HexNut compounds. Show your calculation of the average mass in the right-hand column. NAME Average Mass Calculations

BN

Bi) ________g BN

BN2

Bii)_______g BN2

BN3

Biii)_________BN3

BN4

Biv)_________BN4

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C. The Lentil-Beans LB5 These are found in the plastic baggies labeled LB5. Do NOT open this container. i) Weigh the container containing the 20 Lentil-Beans and record the gross mass in line Ci.

Ci) __________g (gross)

ii) Record the mass of the container in line Cii. This mass is written on the label

Cii) __________g Mass of container

iii) Subtract the mass of the container from the gross mass and write the net mass in line Ciii.

Ciii) __________g 20 LB5 (net)

iv) Calculate the mass of one-Lentil-Bean (LB5) based on the data collected.

Civ) __________g Average Mass of LB5

Calculations:

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Part 2. Counting by Weighing. In this part of this lab, you will determine the number of elements or compounds in a sealed container by weighing the container and its contents and then using the information from part I, you will calculate the number of elements or molecules. D. The number of HexNuts in a container This unknown is found in the green container. Do NOT open this container. i) Record the mass of the “unknown” HexNuts, (Element #1). D i) ______ g Mass of Container+ HexNut (gross) ii) Record the mass of the container as written on the label of the container. D ii)______ g Mass of Container (label on container) iii) Calculate and record the net mass by subtracting the mass of the empty container from the gross mass measured. D iii) ______ g Mass of HexNut in container (net) iv) Using the average mass of a single HexNut from Part I, calculate the number of HexNut elements in your unknown. Show your calculations below. Your answer must be a whole number. D iv) _______ number of HexNut in container

Calculations for D iv.

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E. The number of molecules (BN) in a container This unknown ais found in the orange container labeled “Unknown (BN). Do NOT open this container. Look for the container containing BN items. i) Record the mass of container 2 in line Ei. This is the BN molecules plus the container.

E i) _______g Mass of container + BN (gross)

ii) Record the mass of the container as written on the label of the container in line Eii.

E ii)_______ g Mass of container (label on container)

iii) Calculate the net mass by subtracting the mass of the empty container (as written on the container) from the gross mass measured. Write this mass in line E iii.

E iii) ______ g Mass of BN in container (net)

iv) Using the average mass of a single BN compound, calculate the number of BN compound in your unknown and write this value in line E iv. Show your calculations below. Your answer must be a whole number.

E iv) _____ number of BN compounds in container

Calculations for E iv. F. The number of molecules (Lentil-Bean, LB5) in a container This unknown is found in the purple container labeled “Unknown” (LB5). Do NOT open this container. i) Record the mass of container #3 in line Fi. This is the Lentil-Bean molecules plus the container.

F i) _____ Mass of container + LB5 (gross)

ii) Record the mass of the container as written on the label of the container in line F ii.

F ii) ______ Mass of container (written on label)

iii) Calculate and record net mass of the LB5 in line F iii. F iii) ______ Mass of LB5 (net)

iv) Using the average mass of a single LB5 compound, calculate the number of LB5 compound in your unknown and write this value in line F iv. Show your calculations in the below. Your answer must be a whole number.

F iv) ______ number of Lentil-Bean in container

Calculations for F iv.

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Part 3. Identification of an Unknown based on Average Mass. G & H: Determining the identity of the compound (BN, BN2, BN3 or BN4) These are found in the Blue (G) and Yellow (H) containers labeled “Unknown (BNn). Do NOT open these containers.

Part G Part H

i) Record the number of molecules as written on the containers for this part. Write these in line Gi and Hi. G i) ___ # of molecules H i) ___ # of molecules

ii) Weigh each container and record the mass in line Gii and Hii.

G ii) ______ g Mass of container + Unknown (gross)

H ii) _______ g Mass of container + Unknown (gross)

iii) Record the mass of each container as written on the label of the containers. Write the mass of each container in lines Giii and Hiii.

G iii) _____ g Mass of container

H iii) _____ g Mass container

iv) Calculate the net mass of the content in each container.

G iv) ______g Mass of Unknown1 (net)

H iv) ______ g Mass of Unknown2 (net)

v) Take the net mass and divide by the total number of molecules. This is the average mass of a single molecule in the container. Show your calculations below.

G v) ______g Average mass of Unk1

H v) _______ g Average mass of Unk2

vi) Using the data for the average mass of the BN, BN2, BN3 and BN4 from part 1B, identify the chemical formula of your unknown in the container for this part.

G vi) ________ Formula of molecule in container

H vi) ________ Formula of molecule in container

Calculations for Gv and Hv.

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Part 4. Calculations, Converting from Mass to Moles for some common objects. Using the mass of the iron cube, the 50.0mL of water and the packet of sugar, determine the number of moles of metal, water and sugar. Round of to correct number of significant figures and use scientific notation when appropriate.

I. The number of moles of Fe in a cube

i) Metal cube: Weigh a metal cube and write the mass in line I i.

I i) __________g Fe

ii) Write the average atomic mass of Iron

I ii)______________ g/mole

iii) Calculate the number of moles of the metal in the cube and write your answer in line I ii. (Show calculations below using dimensional analysis)

I iii) ___________moles of iron

iv) How many iron atoms are in the cube?

I iv) ___________ atoms of iron Calculations for Table I.

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J The number of moles of water in 50.0 ml of water i) 50.0 mL of water: Tare a 50mL-graduated cylinder on the scale. Add water and bring to precisely 50.0mL of water using a Berel pipet. Weigh the graduated cylinder with the water and record the mass.

i) ___________g 50.0 mL water

ii) Add up the molar mass of water and write your answer to the right. ii) ____________ g/mole for water

iii) Calculate the number of moles of water in the 50-mL volume of water. Show your calculations below. iii) ___________moles of water

Move to Next Table – answer questions below as a Post Lab

iv) Calculate the number of water molecules in 50.0 mL. Show your calculations below. iv) ___________molecules of water

v) Calculate the number of Oxygen and H atoms in 50 mL of water. Show your calculations below.

v) ___________ O atoms in water ___________H atoms in water

Calculations for Table J

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K. The number of moles of sucrose in a packet of sugar i) Packet of sucrose (sugar), C12H22O11: Weigh a packet of sugar and record the gross mass of the sugar packet.

i) ___________g mass of sugar + bag (gross)

ii) Read and record the mass of the container from the label of the packet.

ii) ___________g mass of bag

iii) Calculate the mass of the sugar in the packet and record in line Kiii.

iii) ___________g mass of sugar (net)

iv) Calculate the molar mass of sucrose (C12H22O11 ). Show your calculations below. iv) ____________ g/mole

v) Calculate the number of moles of sucrose in the packet. Show your calculations below. v) ____________ moles of sugar

Answer questions below as a Post Lab

vi) Calculate the number of molecules of sucrose in the packet of sugar. Show your calculations below. vi) ____________molecules of sugar

vii) Calculate the number of atoms of C, H, and O in the packet of sugar. Show your calculations below.

vii) ___________C atoms ___________H atoms ___________O atoms

Clean up - Replace all items in your bin and bring back to the cart. Complete calculations for table K

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Experiment: Using Gas Laws to Identify an Unknown Liquid Objective In this experiment, you will use the Dumas method to vaporize an unknown volatile liquid into the gas phase. While in the gas phase, the properties of the gas will be measured, applied to a rearrangement of the Ideal Gas Law and used to calculate the molar mass of the unknown. The objective is to identify the unknown liquid based on its molar mass from a list of possible liquids. Materials

o Unknown Liquid o Thermometer o 100-mL graduated cylinder o 125-mL Erlenmeyer Flask o 400 -mL Beaker o Scale o Ring stand o Clamp

o Hot Plate o Paper Clip o Aluminum Foil o Forceps or tongs o Pipet

Safety Do not inhale vapors of the unknown during the heating and evaporation steps of the procedure. Introduction In this experiment, you will exploit the properties of matter in the gas phase to calculate the molar mass (M) of an unknown volatile liquid. The liquid will be heated to evaporation inside a 125-L Erlenmeyer flask so that it is entirely in the gas phase. An algebraic rearrangement of the Ideal Gas Law will be used to calculate the molar mass. Relevant Equations Ideal Gas Law: 𝑷𝑽 = 𝒏𝑹𝑻 P = Pressure in units matching R V = Volume of the gas in liters n = moles of gas T = temperature of the gas in Kelvin R = the ideal gas constant = 0.08206 L*atm/(mol K) = 62.4 L*mmHg/(mol K) The Ideal Gas Law can be rearranged to solve for the number of moles of gas.

𝒏 =𝑷𝑽𝑹𝑻 𝒏 =

𝒎(𝒈𝒓𝒂𝒎𝒔)

𝑴(𝒈𝒓𝒂𝒎𝒔𝒎𝒐𝒍𝒆 )

Therefore, the equation can be rearranged to calculate molar mass (g/mol).

𝑴 =𝒎𝑹𝑻𝑷𝑽

Similarly, this same equation can be rearranged to solve for the density of the gas. Recall the density of any form of matter is the mass divided by the volume (m/V).

𝑫 =𝒎𝑽= 𝑴𝑷𝑹𝑻

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Concepts behind the method All calculations using the Ideal Gas Law equation, or its rearrangements, must use conditions while the unknown is in the vapor phase. There are some assumptions that are made in the calculations of the molar mass (M) of the unknown liquid. Pressure Initially, the unknown is a liquid and exerts only vapor pressure. The flask is filled with air. As the liquid boils, the air will be forced out of the pinhole, and eventually there will no longer be air remaining in the flask, and the only gas will be the unknown. The pressure of the vapor inside the flask is assumed to be the same as atmospheric pressure. To obtain atmospheric pressure, use the barometer in the front of the laboratory room, which reports the pressure in units of mmHg. If there is no barometer in the room, you may use the NOAA.gov website to get the local weather conditions, which will be reported in inches of mercury (inHg) and will need to be converted. Volume The volume of gas is defined by the container in which it is contained. In this experiment, this container is the 125-ml Erlenmeyer flask. The volume of this flask will be accurately determined by measuring the mass of water contained within it, and by transferring the water to a graduated cylinder and measuring. Temperature The temperature of the unknown in the gas phase is assumed to be the same as the water bath, in which it is submerged. For this reason, the temperature of the boiling water is measured and applied to the gas law equations. It is important to have as much as the flask submerged as possible without spilling water into it. Mass of the Vapor It is not possible to put matter on a scale and measure its mass while it is in the vapor phase. However, in this experiment, after the unknown has completely evaporated, the flask is occupied completely by the unknown vapor. Once removed from the heat and allowed to cool, the unknown vapor will recondense and form a liquid at the bottom of the flask. You may or may not be able to see the liquid. However, you can determine the mass by recording the mass of the flask, the foil and the condensed liquid, and subtracting the mass of the flask and the foil that was recorded at the beginning of the experiment. Possible Unknowns This table includes a list of possible unknowns. Verify with your instructor that this list is current as options may change from semester to semester. Calculate the molar mass of each chemical and submit with your results.

Chemicals Formula MM (g/mol)

Acetonitrile C2H3N

Ethanol C2H6O

Ethyl Acetate C4H8O2

2-Propanol C3H8O

Isopropyl Acetate C5H10O2

Hexane C6H14

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Procedure – Evaporation of the Volatile Liquid, Mass of the vapor 1. If your instructor does not provide you the barometric (atmospheric) pressure in the lab on the day of the

experiment, then she/he will show you how to read the barometer. Record in the datasheet the atmospheric pressure.

2. While the water is heating, obtain 3 clean, dry 125-mL Erlenmeyer flasks and weigh it together with a square

of aluminum foil. Record this combined mass in the datasheet (#1).

3. Also, remove the paper backing from 3 squares of parafilm and weigh each flask with one piece of parafilm (without the foil). Record the mass in the datasheet row #8. Set the parafilm aside.

4. Pour ~2-3 mL of your assigned unknown liquid into the flask.

5. Crimp the aluminum foil over the mouth of the flask to form a cap;

try to seal it as completely as possible. Fold the foil up so that it doesn't hang down too far on the neck of the flask (~ 1 cm is far enough). Using a pin, poke a small pinhole at the center top of the foil.

6. Set up the apparatus for the evaporation of your unknown liquid, as shown to the right. Clamp the flask to the ring stand. If you tilt the flask slightly, it may be easier to observe when all the liquid has vaporized. Immerse the flask as much as possible, so that the vapor inside the flask is uniformly heated by the water.

7. Heat the water bath to boiling and measure the temperature of the

boiling water. Be sure you have at least reached the boiling point of water. Record this temperature in your data sheet (#2). Continue heating until all the liquid inside the flask has vaporized, and heat for at least 1 minute after all the liquid has vaporized.

8. Carefully remove the flask from the hot water bath. Allow it to cool to room temperature. Make sure it is

completely dry and carefully reweigh the flask with the foil. Record this mass in the datasheet (#6). Repeat this for the second and third trial. Make sure that you have weighed the other flask before you begin the second and third trial.

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Determination of the Volume of the Flask The volume of the flask is the volume of the vapor. After completing the procedure above, you will also determine the volume of each flask used. This will be accomplished two ways. First, the volume will be calculated from the mass of water held inside the Erlenmeyer flask using density. Volume via mass of water 9. Discard the condensed liquid from your last run and discard the aluminum foil, then rinse the flask 2-3 times

with a small amount of water. Fill the flask completely water to the rim, seal the water in the flask with parafilm wax so the water doesn’t spill and then weight the water in the flask. Weigh the flask, parafilm and water and record in the datasheet in row (#9)

10. Remove the parafilm and discard it, and measure the temperature of the water and record.

11. Using the mass of the flask/water/parafilm (#9) subtract the mass of the flask/parafilm to obtain the mass of the water in the flask (#10).

12. Use the density of water at the temperature of the water to determine the volume of the flask. Volume via volume measurement of water A second method to determine the volume is to measure it directly using a graduated cylinder. 13. Begin with the flask filled completely, all the way to the rim with water. There is more than 100 mL of water in

the Erlenmeyer flask, so you will need to transfer the water to a 100-mL graduated cylinder in 2 stages. Record the volume of water in the flask in row #12 in the data sheets.

Repeat this process two more times Clean up Wash and dry all your glassware equipment. Dispose of all used chemicals in the proper waste container. Wipe down your station and place your glassware and equipment back in its proper place. Unless everyone’s equipment is back in the locker drawer, you are not allowed to remove your safety goggles. Wash your hands thoroughly with soap and water before leaving your work area.

DENSIY OF WATER FROM 0°C TO 40°C °C g / mL °C g / mL °C g / mL 0 0.999868 5 0.999992 10 0.999728 1 0.999926 6 0.999968 11 0.999634 2 0.999968 7 0.999930 12 0.999526 3 0.999992 8 0.99876 13 0.999406 4 1.000000 9 0.999809 14 0.999273

°C 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 15 0.999 129 0.999 113 0.999098 0.999083 0.999067 0.999052 0.999036 0.999020 0.999004 0.998988 16 0.998 972 0.998 956 0.998 939 0.998 923 0.998 906 0.998 890 0.998 873 0.998 856 0.998 839 0.998 821 17 0.998 804 0.998 787 0.998 769 0.998 752 0.998 734 0.998 716 0.998 698 0.998 680 0.998 662 0.998 643 18 0.998 625 0.998 606 0.998 588 0.998 569 0.998 550 0.998 531 0.998 512 0.998 493 0.998 474 0.998 454 19 0.998 435 0.998 415 0.998 395 0.998 375 0.998 356 0.998 336 0.998 315 0.998 295 0.998 275 0.998 254 20 0.998 234 0.998 213 0.998 192 0.998 171 0.998 150 0.998 129 0.998 108 0.998 087 0.998 065 0.998 044 21 0.998 022 0.998 000 0.997 979 0.997 957 0.997 935 0.997 912 0.997 890 0.997 868 0.997 846 0.997 823 22 0.997 800 0.997 778 0.997 755 0.997 732 0.997 709 0.997 686 0.997 662 0.997 639 0.997 616 0.997 592 23 0.997 568 0.997 545 0.997 521 0.997 497 0.997 473 0.997 449 0.997 424 0.997 400 0.997 376 0.997 351 24 0.997 327 0.997 302 0.997 277 0.997 252 0.997 227 0.997 202 0.997 177 0.997 152 0.997 126 0.997 101 25 0.997075 0.997049 0.997024 0.996998 0.996972 0.996946 0.996920 0.996893 0.996867 0.996841

°C g/mL °C g / mL °C g / mL 26 0.996814 31 0.995372 36 0.993716 27 0.996544 32 0.995058 37 0.993360 28 0.996264 33 0.994734 38 0.992997 29 0.995976 34 0.994403 39 0.992626 30 0.995678 35 0.994064 40 0.992247

These densities, given in g / mL, can be converted to g / cm3 by multiplying each value by 0.999972

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Mass of the Vapor of an Unknown Liquid Raw Data: Trial 1 Trial 2 Trial 3

1. Mass of flask & Aluminum foil (g)

2. Temperature of vapor (°C)

3. Temperature of vapor (K)

4. Pressure of vapor (mm Hg)

5. Pressure of vapor (atm)

6. Mass of flask & foil & vapor (g)

7. Mass of vapor (g)

Observations.

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Volume of the Vapor Volume of the flask via mass of the water Trial 1 Trial 2 Trial 3

8. Mass of flask & parafilm (g)

9. Mass of flask & parafilm & water (g)

14. Mass of water (g)

11. Density of water (g/ml) at temperature at 23˚C

0.998 g/ml 0.998 g/ml 0.998 g/ml

12. Use the density and mass of water to calculate the volume of flask (mL)

Volume of the flask via the graduated cylinder Trial 1 Trial 2 Trial 3

13. Volume of flask via graduated cylinder (mL)

Work Space for calculations.

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Calculation of Molar mass and Density: Trial 1 Trial 2 Trial 3

15. The Equation for molar mass of a gas (M)

𝑴 =𝒎𝑹𝑻𝑷𝑽

16. Molar mass of vapor (g /mol) Show sample calculation here.

17. Calculate the mean value of molar mass of vapor for the 3 trials. (g / mol)

18. The Equation for density of a gas 𝑫 =𝒎𝑽 =

𝑴𝑷𝑹𝑻

19. Take the Molar Mass (M) and calculate the Density of the gas (g/L). Show a sample calculation here.

20. Mean value of density of vapor (g / L)

Unknown Number: _________ Chemical name and formula of unknown: Molar mass of unknown:

Use the formula of the molar mass of a gas and the density of a gas to fill the following table. (Don't forget to convert your temperature to the absolute scale: T(K) = T(°C) + 273.15; R = 0.0821 L· atm/K mol)

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Postlab Questions Answer these post lab questions and turn in to your instructor before you leave lab unless told otherwise by instructor. 1. Two methods were used to determine the volume of the flask used for the molar mass determination. Which method will give a more precise determination of the volume? Why? 2. It was important that the flask be completely dry before the unknown liquid was added so that water present would not vaporize when the flask was heated. A typical single drop of liquid water has a volume of approximately 0.050 mL. Assuming the density of liquid water is 0.998 g/mL, how many moles of water are in one drop of liquid, and what volume (mL) would this amount of water occupy when vaporized at 100.°C and 1.0 atm? 3 Calculate the molar mass of each possible unknown below and circle your unknown for this experiment.

Chemicals Formula MM (g/mol)

Acetonitrile C2H3N

Ethanol C2H6O

Ethyl Acetate C4H8O2

2-Propanol C3H8O

Isopropyl Acetate C5H10O2

Hexane C6H14

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Experiment: Concentration of a Salt Solution Objective The purpose of this experiment is to determine the weight & concentration of dissolved solids in salt-tainted water solutions. Concentration by parts (%, ppm & ppb) will be determined. Materials

o Salt-tainted water solutions o Three beakers o Hot plate o Scale o Graduated cylinder o Heat gun

Introduction Additional Reading: Ch 9.4, Chemistry in our Lives, 5th Edition Timberlake. “Solution Concentrations and Reactions.” For this experiment, you will be determining the concentration of a saline solution of unknown concentration. Your results will be reported in a mass percent (m/m %), ppm and ppb. A portion of your grade will reflect the accuracy and precision of your results. Procedure (steps below correspond to data sheet) – 1. Your instructor will assign you an unknown. This unknown is the salt-tainted water sample. Be sure to do

three trials then average the results. Use the three largest beakers in your lab drawer for this procedure. The large beakers provide wider surface area, which allows for faster evaporation of the water during the experiment. Thoroughly clean with soap and water three beaker and be sure it is dry. Inspect the glassware for stress, stars or cracks, if you suspect your beaker will break, have your instructor inspect your beaker and change it out if necessary.

2. Label the beakers with "1", "2" and “3” on the side. Weigh each of the beakers on the balance to the thousandth of a gram (0.001 g). Write the mass of each beaker in your report sheet. This is in row #2 on data sheet.

3. Measure ~100 ml of the salt-tainted water sample that has been assigned to using a 100mL beaker and pour

it into beaker labeled “1”. Do not worry about reading the exact volume since the amount of water will be weighed. Repeat this for trials 2 and 3. Weigh each beaker, which contains the salt-tainted water solution on the balance to the nearest thousandth of a gram (0.001 g). Write this weight as the " Mass Salt-tainted solution + beaker" measurement. This is in row #3 on data sheet. Repeat this for the other two samples. Remember that the salt solution contains water as the solvent and salt as the solute. The salt is the residue left behind after evaporating the solvent.

4. Using the mass of the beaker and the mass of beaker + solution, calculate the mass of the solution. This is

in row #4 on Data sheet. 5. Place each beaker on the hot plate and heat the beaker so that the water boils gently. The rate of boiling can

be controlled by adjusting the heating dial or by moving the beaker away from the center toward the side of the hotplate. When the volume of water sample has finally boiled down to about 5-10 ml, it is advisable to cover the beaker with a watch glass to catch spattering from the concentrated solution. When the beaker is dry, rinse the splattering off the watch glass back into the beaker with a little deionized water using a water bottle. Because the watch glass is hot, it is necessary to hold it with tongs. Evaporate the small amount of water in the beaker to dryness by heating the dish gently without the watch glass.

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When the solution is down to about 0.5 mL, remove reduce the heat by moving the beaker to the side of the hot plate and reduce the temperature setting. When all the water in the beaker has evaporated, carefully use mitts to remove the beaker and place it on the wire screen on the lab bench. If there is still water along the walls of the beaker, use the heat gun to evaporate the residual moisture. * From this point on, use tongs or hot mitts to handle the beaker so that oil from your fingers do not add to the mass of the residue. Use the oven-mitts to handle the hot beaker and do not touch the hot beaker directly. Note and record the color of the residue (the solute salt) and describe this in your data sheet. This is in row #5 on Data sheet.

6. Allow the beaker to cool for an additional 10 - 15 minutes making sure that there are no droplets of water

(solvent) remaining. Weigh each beaker which contains the solute (residue left behind in beaker) and write the result in your datasheet as " Mass beaker + solute ". This is in row #6 on data sheet.

7. Use the "mass of beaker + solute" (row #6) to and the “mass of beaker” (row #2) to calculate the weight of

dissolved solids in the water sample. This is in row #7 on data sheet. The mass is the amount of total dissolved solids (TDS) in your sample. You will express the the TDS in parts per hundred (% mass), parts per million (ppm) and parts per billion (ppb) later in the calculation. Be sure to use the correct number of significant figures in your calculations.

Clean up Wash and dry all your glassware equipment. Dispose of all used chemicals in the proper waste container. Wipe down your station and place your glassware and equipment back in its proper place. Unless your equipment is back in the locker drawer, you are not allowed to remove your safety goggles. Wash your hands thoroughly with soap and water before leaving your work area.

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A. Datasheet

Sample 1 Sample 2 Sample 3 Unknown # ________

Description of sample:

Step1. Approximate volume of salt-tainted solution, measured from graduated cylinder:

Step2. Mass of beaker:

Step3. Mass of solution + beaker:

Step4. Mass of solution:

Step5. Heat Solution to evaporate the water. Residue or solute (salt) is left behind. Write a description of the solute.

Step6. Mass beaker + solute (salt):

Step7. Mass of solute

Show one sample calculation in proper format for Mass of solute (#5) and mass of solution (#6):

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B. % (m/m) of Salt Solution Sample 1 Sample 2 Sample 3

8 Calculation: % concentration (m:m)

9 Calculation: Average % Concentration

Show one sample calculation in proper format for the % concentration of your solution and the calculation of the average. C. Parts per million (ppm) and Parts per billion (ppb)

Sample 1 Sample 2 Sample 3

10a Calculation: ppm

10b Calculation: Average ppm

11a Calculation: ppb

11b Calculation: Average ppb

Show one sample calculation in proper format for the ppm and ppb for one of your sample.

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Postlab Questions 1. If the residue is not heated to dryness but is still somewhat moist, how would this change the final amount of residue calculated in the solution? Would you expect the calculated concentration of the residue to be higher or lower than the true value? (Check mark and provide explanation)

Concentration: Higher ____ Lower ____ No Change _____ Not enough information ____ than true value. 2. If a large piece of dirt, say 10 mg, falls into the beaker just after the sample is evaporated to dryness and is weighed, what effect would this have on the weight of dissolved solids found? Would you expect the calculated concentration of the residue to be higher or lower than the true value? (Check mark and provide explanation)

Concentration: Higher ____ Lower ____ No Change _____ Not enough information ____ than true value. 3. If a large amount of water splatters out of the beaker, say ~10 mL, during the evaporation process. Explain how would this affect the mass of the residue calculated in your result? Would you expect the calculated concentration of the residue to be higher or lower than the true value?

Concentration: Higher ____ Lower ____ No Change _____ Not enough information ____ than true value.

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4. Based on your data, Summarize the result of your experiment by calculating the following for one of your samples. Sample number used for this calculation: Sample No.: _______________ Mass of Solution Mass of solution: _______________ Mass of Solute Mass of solute: _______________ Mass of solvent Mass of solvent: _______________ In your own words and your experience in this experiment, difference between a solution and the solvent? 5. If you calculated a salt-tainted solution of about 3.5% *, what would the molarity of this solution? Assume the density of this solution is 1.03 g/mL. (* BTW, this is the salinity of seawater)

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Experiment: Introduction to Equilibrium and LeChatelier’s Principle Objective The objective of this experiment is to determine the effects of disturbances on chemical systems at equilibrium. The response of the chemical systems will be justified in terms to LeChatelier's principle. Materials

o Bromothymol Blue o 0.1 M Zn(NO3)2 o NH4Cl saturated solution o Phenolphthalein o 15 ml Club Soda o NH4Cl crystals o 0.1 M NaOH

o 6 M NaOH o Distilled water o 0.1M HCl o 6 M HCl o Test tubes o Syringe with cap

Safety Note: 6M HCl is a strong acid and 6M NaOH is a strong base and should be handled with care. 0.1 M Hydrochloric (HCl) acid-corrosive and 0.1 M sodium hydroxide-toxic, corrosive, and irritant. Prevent contact with your eyes, skin, and clothing. Avoid ingesting the substance. If you spill any solution, immediately notify your laboratory instructor. Introduction Required additional reading: For a more complete description of Equilibrium and LeChatelier’s Principle, read Chemistry in our Lives, 5th Edition Timberlake. Ch10.2 “Chemical Equilibrium” and 10.5 “Changing Equilibrium Conditions: LeChatelier’s Principle.” For this experiment, you will be establishing chemical equilibria and observing the effects of perturbations on the equilibria. You will be asked to explain your observations in terms of the chemical equilibrium reaction equation and LeChatelier’s Principle. LeChatelier’s Principle states that when a system at equilibrium is exposed to a stress, the system will shift to re-establish equilibrium. There are many kinds of stresses to systems at equilibrium. Effects of Changing Concentration of Reactants or Products Changing the concentrations of reactants or products of a system at equilibrium will cause the system to respond to reestablish the equilibrium conditions.

aA + bB cC + dD Consider the above reaction with reactants, A and B, and products, C and D, and the related equilibrium expression.

𝑲𝒆𝒒 =[𝑪]𝒄[𝑫]𝒅

[𝑨]𝒂[𝑩]𝒃

Keq is a constant, so a change in the concentrations of one or more of A, B, C or D, must result in a change in the other species involved in the equilibrium to keep the right side of the equation equal to Keq. Excess Stress: Effect of Increasing Concentrations An increase in the concentration of reactant A to the above system at equilibrium will cause an excess stress. The system will reestablish equilibrium by consuming A and shifting the reactant toward products. An increase in product C or D will shift the equilibrium to the left.

Deficit: Effect of Decreasing Concentrations Another means of disturbing a chemical system at equilibrium is to reduce the amount of one of the components, a deficit stress. For instance, we could reduce the amount of product C in the system at equilibrium. The system would respond by “shifting to the right” to replenish C and establish a new equilibrium state. This phenomenon can be used to increase the yield of a product at equilibrium.

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One way to affect a deficit stress is to remove a product or reactant with a secondary or side reaction that consumes the chemical species. In our above system, it would be represented as illustrated below.

X reacts with A to create the product AX. If X is added to this system at equilibrium, it will cause a deficit stress for reactant A. Thus, the equilibrium will shift toward reactants to replenish reactant A. In the reaction for the Haber process shown below, removing the NH3 as it forms during the reaction can increase the yield of NH3, when the equilibrium shifts to replace it.

N2 (g) + 3 H2 (g) 2 NH3(g)

For this experiment, you will need to identify deficit disturbance created this way. It will be useful to know the following reaction.

H+ (aq) + OH- (aq) H2O (l)

The Effect of Pressure LeChatelier’s Principle states that for equilibria containing gases, an increase in pressure, will disturb the system, and shift the reaction to the side with the lesser number of moles of gas. A decrease in the moles of gas will reduce the pressure in the reaction vessel. If the volume of the vessel is increased, the corresponding decrease in pressure will cause a shift in the equilibrium toward the side that has more gas molecules. If a reaction is such that there are equal number of moles of gas in the reactant and the product, then there is no effect on the equilibrium as a result of a volume change in the reaction vessel. The Effect of Temperature A change in temperature will also stress an equilibrium. An exothermic reaction is one that releases heat, that is, the energy of the reactants is greater than that of the products. Since heat is released as the reaction progress in the forward direction then heat can be considered a product. A temperature increase for an exothermic reaction will cause the equilibrium to shift to the to alleviate the excess heat stress caused on the right (product) side of the reaction. For example, the Haber process is exothermic, and can be written:

N2 (g) + 3H2 (g) 2NH3(g) + heat

Heat is released as a product of the forward reaction. A temperature increase causes the equilibrium to shift to the left, resulting in a reduced amount of heat and NH3, and an increase production of N2 and H2. Solubility of gases is another example of exothermic processes. An endothermic reaction is one that absorbs heat. The energy of the products is greater than that of the reactants. Since heat is required for the forward reaction to occur. In this case heat is treated as a reactant in the chemical equation. Therefore, a temperature increase will cause the equilibrium to shift to the right. Table 1 summarizes the stresses and responses to perturbations of a system at equilibrium. In this laboratory experiment, you will need to identify the type of stress that you have affected for several equilibrium reactions and the response. Your answers should cite a stress from this table. Your description must include the specific details, reactants or products, of the reaction you are studying. A clear description of the details as they relate to Table 1, will demonstrate to your instructor that you have a clear understanding of the concepts related to LeChatelier’s Principle.

A + B C + D+X

AX

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Table 13.1 Summary of Equilibrium Stresses and Response A Stress via Concentration Reaction Respond by- The resulting shift A1 Increase [reactants] Consumes the reactants Reaction will shift to the right (product) → A2 Increase [products] Consumes the product Reaction will shift to the left (reactant) ← A3 Decrease [reactants] Replenish the reactants Reaction will shift to the left (reactant) ← A4 Decrease [products] Produce more products Reaction will shift to the right (product) → B Stress via Concentration Reaction Respond by- The resulting shift B1 Increase Pressure Increase side with fewest moles of gas Reaction will shift to side with fewer moles of gas B2 Decrease Pressure Increase side with most moles of gas Reaction will shift to side with more moles of gas

Stress via Temperature / Energy changes* for Endothermic or Exothermic reactions C Stress via Temp Increase Reaction Respond by- The resulting shift C1 - Endothermic Reaction Consume the heat of reactant Reaction will shift to the right (product) → C2 - Exothermic Reaction Consume the heat of product Reaction will shift to the left (reactant) ← C Stress via Temp decrease Reaction Respond by- The resulting shift C3 - Endothermic Reaction Produce energy as the rxn reverses Reaction will shift to the left (reactant) ← C4 - Exothermic Reaction Energy forms as rxn moves forward Reaction will shift to the right (product) →

*Endothermic reactions are reactions in which energy/heat is considered a reactant, energy absorb. DH (+) *Exothermic reactions are reactions in which energy/heat is considered a product, energy release. DH (-) The Observables in this Experiment Many chemical reactions occur without a directly observable change in the reaction vessel. For this experiment, reactions in equilibrium have been selected that have observable physical properties in the reactants, products or both. Table 13.2 Reaction Reactants Products

(Eq. 1) HInd (aq, yellow) H+ (aq) + Ind- (aq, blue) Yellow Blue

(Eq. 2) Zn2+ (aq) + 2OH- (aq) Zn(OH)2 (aq)

Clear, colorless Turbid, White precipitate

(Eq. 3) H2CO3(aq) H2O(l) + CO2 (g)

Clear, Colorless Bubbles

(Eq. 4) NH4Cl(s) NH4+ (aq) + Cl- (aq)

Turbid, White precipitate Clear, Colorless

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Procedure – Prepare a hot water bath before starting these procedures. A. Bromothymol Blue in equilibrium between yellow (Hind) and blue (Ind-) forms:

(Eq. 1) HInd (aq, yellow) H+(aq) + Ind- (aq, blue)

Caution: 0.1 M Hydrochloric (HCl) acid-corrosive and 0.1 M sodium hydroxide-toxic, corrosive, and irritant. Prevent contact with your eyes, skin, and clothing. Avoid ingesting the substance. If you spill

any solution, immediately notify your laboratory instructor. 1. Half-fill a clean, dry test tube with deionized water. Add 2 drops of the indicator, bromothymol blue. Record the

color of the solution on your datasheet (1). 2. Add 0.1 M NaOH, dropwise, while stirring with a clean, dry, glass, stirring rod until a color change occurs. Add 2

more drops. Record your observations on your datasheet (2). 3. Add 0.1 M HCl, dropwise, while stirring with a clean, dry, glass, stirring rod until a color change occurs. Record

your observations on your datasheet (3). 4. Again, add 0.1 M NaOH, dropwise, while stirring with a clean, dry, glass stirring rod until a color change occurs.

Record your observations on your datasheet (4). 5. Experiment by adding the right amount of acid (HCl) dropwise to this test tube to make the solution turn greenish

in color after it is stirred. Record your observations on your datasheet (5). Include how many drops were needed.

6. Pour all waste in the labeled waste container at the appropriate station. Rinse the test tube with 3-5 mL of tap

water and empty the rinse in the waste container. Finally, rinse the test tube with deionized water. B. Solubility of Zinc Hydroxide:

(Eq.2) Zn2+(aq) + 2OH- (aq) Zn(OH)2 (s)

Caution: Concentrated NaOH is toxic, corrosive, an irritant and can cause burns. Prevent contact with your eyes, skin, and clothing. Avoid ingesting the substance. If you spill any solution,

immediately notify your laboratory instructor.

1. Using a Berel pipet, add 20 drops (~1 mL) of 0.1M Zn(NO3)2 to a clean, dry test tube. Add 1 drop of 6M NaOH. Stir with a clean, dry, glass stirring rod. Record your observations on your datasheet (2 a).

2. Add 6M HCl, dropwise, while stirring with a clean, dry, glass, stirring rod until a color change occurs and no more

changes are seen. Record your observations on your datasheet (2 b). 3. Using a Berel pipet, add 20 drops (~1 mL) of 0.1M Zn(NO3)2 to a second clean, dry test tube. Add 1 drop of 6M

NaOH. Stir with a clean, dry, glass stirring rod. 4. Add 6M NaOH, dropwise, while stirring with a clean, dry, glass, stirring rod until a color change occurs and no

more changes are seen. Record your observations on your datasheet (2 d).

5. Pour all waste in the labeled waste container at the appropriate station. Rinse the test tubes with 3-5 mL of tap water and empty the rinse in the waste container. Finally, rinse the test tubes with deionized water.

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C. Carbonic acid and Carbon Dioxide

(Eq. 3) H2CO3 (aq) H2O (l) + CO2 (g)

1. Remove plunger from syringe and place cap on the end. To the syringe, add approximately 15 mL of club soda. Reinsert plunger, invert syringe so that it is pointing up, remove cap and push on plunger so there is no gap between the plunger and the soda. Replace cap. You should see bubbles in the soda solution, if you don’t pull the plunger slightly until some bubbles form.

2. Observe the bubbles within the syringe and record observations on your datasheet (1). 3. Increase the pressure by squeezing down on the plunger. Keep the syringe tip pressed on the bench top so that

the syringe cap does not get pushed off. Observe any changes in the bubbles. Record observations on your datasheet (2).

4. Increase the volume in the syringe and therefore decreasing the pressure in syringe by slowly pulling out the

plunger but do NOT completely remove the plunger from the syringe. Observe any changes in the bubbles. Record observations on your datasheet (3). Note, be sure you start with some bubbles in the syringe otherwise your results may be inconclusive.

5. Carry out a procedure to determine the thermicity (endothermic or exothermic) of equation 3. To do this start

with a syringe that shows bubbles in the soda then place either in hot or cold water. Circle the correct choice for this reaction, endothermic or exothermic. (4)

6. Empty syringe contents into the sink. Thoroughly rinse syringe with tap water, followed by deionized water.

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D. Solubility of Ammonium Chloride

(Eq. 4) NH4Cl(s) NH4+(aq) + Cl-(aq)

Caution: Concentrated HCl is toxic, corrosive, and can cause burns. Prevent contact with your eyes, skin, and clothing. Avoid ingesting the substance. If you spill any solution, immediately notify your

laboratory instructor.

1. Measure 3 mL of saturated NH4Cl solution in a clean test tube. Add 1 drop of concentrated HCl, while stirring the test tube. Continue adding concentrated HCl dropwise, stirring after each addition, until a change occurs. Record your observations on your datasheet (1).

2. Using a test tube holder, place the test tube into the boiling water in the hot-water bath. Stir the solution in the

test tube while heating for 3 min. Record your observations on your datasheet (2). 3. Add ice to a 400mL beaker to the half way mark and then fill with water up to the 300mL level. Using a test tube

holder remove your test tube from the boiling water and place it in the cold-water bath. Leave the test tube in the cold-water bath for 3-5 minutes. Record your observations of your datasheet (3).

4. Using a test tube holder remove the test tube from the cold-water bath and place back in the hot water bath.

Record your observations on your datasheet (4) 5. Turn off the hot plate and allow the hot plate to cool before returning it back to its storage space. 6. Pour all waste in the labeled waste container at the appropriate station. Rinse the graduated cylinder and the

test tube with 3-5 mL of tap water each and then empty the rinse in the waste container. 7. Using a clean spatula, add enough crystals of solid NH4Cl to cover the bottom of a second clean, dry test tube. 8. Add 5 mL of deionized water onto the solid NH4Cl in the test tube. Stir the mixture. Feel the test tube to

determine if the dissolution of NH4Cl produces a temperature change. Record your observations (5). 9. Pour all waste in the labeled waste container at the appropriate station. Rinse the graduated cylinder and the

test tube with 3-5 mL of tap water each and then empty the rinse in the waste container. Wash all your glassware and dispose of all used chemicals in the proper waste containers.

Clean up Wash and dry all your glassware equipment. Dispose of all used chemicals in the proper waste container. Wipe down your station and place your glassware and equipment back in its proper place. Unless your equipment is back in the locker drawer, you are not allowed to remove your safety goggles. Wash your hands thoroughly with soap and water before leaving your work area.

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A. Bromothymol Blue in equilibrium: (Eq. 1) HInd (aq, yellow) H+(aq) + Ind- (aq, blue)

1. Check mark the color of the solution after addition of indicator bromothymol blue. _____ Yellow _____ Blue _____ Colorless

_____ Green _____ Orange _____ White 2. Check mark the change that you observe after addition of NaOH. _____ Turns Yellow _____ Turns Blue _____ Turns Colorless

_____ Turns Green _____ Turns Orange _____ Turns White If none of the choices match, write your own description here. 3. Check mark the change that you observe after addition of HCl. _____ Turns Yellow _____ Turns Blue _____ Turns Colorless

_____ Turns Green _____ Turns Orange _____ Turns White If none of the choices match, write your own description here. 4. Check mark the change that you observe after re-addition of NaOH. _____ Turns Yellow _____ Turns Blue _____ Turns Colorless

_____ Turns Green _____ Turns Orange _____ Turns White If none of the choices match, write your own description here. 5. How many drops were of HCl were needed to obtain a greenish color? 6. Using LeChatelier's Principle, cite letters A1, B1… from Table 13.1 to explain the color changes you observed. a) Briefly explain how the addition of NaOH affects the equilibrium. (Use complete sentences) b) Briefly explain how the addition of HCl affects the equilibrium. (Use complete sentences)

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B. Solubility of Zinc Hydroxide : (Eq.2) Zn2+(aq) + 2OH- (aq) Zn(OH)2 (s)

1. Check mark the color of the solution after addition of 1 drop NaOH. ______ Yellow solution _______ Blue solution ________Colorless solution ______ orange solution _______ Pink solution ________White precipitate 2. Check mark the change that you observe after addition of HCl. ______ solution turns yellow _______ solution turns blue ________solution turns colorless ______ solution turns green _______ solution turns orange ________ white precipitate forms

If none of the choices match, write your own description here. 3. Check mark the change that you observe after addition of Zn(NO3)2. ______ solution turns yellow _______ solution turns blue ________solution turns colorless ______ solution turns green _______ solution turns orange ________ white precipitate forms

If none of the choices match, write your own description here. 4. (2 d) Check mark the change that you observe after dropwise addition of NaOH. ______ solution turns yellow _______ solution turns blue ________solution turns colorless ______ solution turns green _______ solution turns orange ________ white precipitate forms If none of the choices match, write your own description here. 4. Using LeChatelier's Principle cite letters A1, B1… from Table 13.1 to explain the changes you observed. a) Briefly explain how the addition of HCl affects the equilibrium. (Use complete sentences) b) Briefly explain how the addition of a small amount of NaOH affects the equilibrium. (Use complete sentences) c) Briefly explain how the addition of a larger amount of NaOH affects the equilibrium. (Use complete sentences)

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C. Carbonic acid and Carbon Dioxide (Eq. 3) H2CO3 (aq) H2O (l) + CO2 (g)

1. Check mark all that applies for your observations of the fate of the bubbles in this part of the experiment. ______ There are no bubbles _______ More bubbles form ________Less bubbles form ______ Bubbles become smaller _______ Bubbles become bigger ________ Bubbles does not change If none of the choices match, write your own description here. 2. Check mark all that applies for your observations of the fate of the bubbles in this part of the experiment. ______ There are no bubbles _______ More bubbles form ________Less bubbles form ______ Bubbles become smaller _______ Bubbles become bigger ________ Bubbles does not change If none of the choices match, write your own description here. 3. Check mark all that applies for your observations of the fate of the bubbles in this part of the experiment. ______ There are no bubbles _______ More bubbles form ________Less bubbles form ______ Bubbles become smaller _______ Bubbles become bigger ________ Bubbles does not change If none of the choices match, write your own description here. 4. Check mark below the thermicity of this reaction (exothermic or endothermic)? _____ Endothermic ______ Exothermic Support your choice by describing your observations. 5. Using LeChatelier's Principle cite letters A1, B1… from Table 13.1 to explain the changes you observed. a) Briefly explain how the pushing on the plunger affects the equilibrium. (Use complete sentences) b) Briefly explain how pulling back on the plunger affects the equilibrium. (Use complete sentences)

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D. Solubility of Ammonium Chloride (Eq. 4) NH4Cl(s) NH4+(aq) + Cl-(aq)

1. Check mark all the choices below that is consistent of what you observed after addition of HCl. ___ The solution starts clear ___The solution starts white ___The solution starts with different color ___ The solution turns clear ___The solution turns white ___The solution stays the same 2. Check mark all the choices below that are consistent of what you observed after heating the test tube. ___ The solution starts clear ___The solution starts white ___The solution starts with different color ___ The solution turns clear ___The solution turns white ___The solution stays the same If none of the choices match, write your own description here. 3. Check mark all the choices below that is consistent of what you observed after cooling the test tube. ___ The solution starts clear ___The solution starts white ___The solution starts with different color ___ The solution turns clear ___The solution turns white ___The solution stays the same If none of the choices match, write your own description here. 4. Check mark all the choices below that is consistent of what you observed after reheating the test tube. ___ The solution starts clear ___The solution starts white ___The solution starts with different color ___ The solution turns clear ___The solution turns white ___The solution stays the same If none of the choices match, write your own description here. 5. Check mark all the choices below that is consistent of what you observed after dissolving the NH4Cl in deionized water. ___ The solution starts clear ___The solution starts white ___The solution starts with different color ___ The solution turns clear ___The solution turns white ___The solution stays the same If none of the choices match, write your own description here. 6. Was there a temperature change in this procedure?

___Solution warmed up ___ Solution cooled down ___ No temperature change

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7. Is the dissolution (dissolving) of NH4Cl exothermic or endothermic? _____ Endothermic ______ Exothermic 8. Rewrite the net equation for the dissolution of NH4Cl, including heat as either a reactant or product. _______ + _______ D _______ + _______ 9) Using LeChatelier's Principle cite letters A1, B1… from Table 13.1 to explain the changes you observed. a) Briefly explain how the addition of HCl affects the equilibrium. (Use complete sentences) b) briefly explain how the change in temperature affects the equilibrium. (Use complete sentences)

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Post-Laboratory Questions 1. The popular antacid, Milk of Magnesia, is a suspension of magnesium hydroxide, Mg(OH)2. In water, Mg(OH)2 undergoes the reaction shown :

Mg(OH)2 (S, white) Mg2+ (aq) + 2 OH- (aq) a) If an acid (H+) were added to this mixture, which species would it react with? b) Would this cause an excess or deficit stress on that species? c) Predict if the solution will appear cloudier or clearer after the addition of acid. d) Use LeChatelier's Principle to explain your predicted observation in (c). 2. Consider the gas equilibrium that you studied in this experiment. If we are told that the equilibrium favors the “right” side of the equation, is the value of Keq a relatively large number or a relatively small number? Briefly explain using your knowledge of determining Keq. 3. In Step 1 E, after obtaining a greenish color, what does this tell you about the relative concentrations of HInd and Ind- in the solution?

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Experiment: Titration of Vinegar Objective The purpose of this experiment is to determine the molarity and weight/volume percent concentration of acetic acid in household vinegar by titration with sodium hydroxide to a phenolphthalein endpoint. Materials

o Deionized water o vinegar o 0.200 M NaOH o Phenolphthalein o 250-ml beaker

o 50-ml buret o 250-ml Erlenmeyer flask o 400-ml beaker for waste o White paper (2-3 sheets) o 5-ml volumetric pipette

Introduction Required additional reading: For a more complete description of Titrations, read Chemistry in our Lives, 5th Edition Timberlake. Ch11.8 “Acid-Base Titration.” Procedure –

Table A. Setup 1. Record the brand of vinegar and its weight / volume percent of acetic acid in your worksheet. 2. Using a 5 ml volumetric pipette, place 5.00 ml of vinegar in a clean 250 ml Erlenmeyer flask (or a 125mL

Erlenmeyer flask). The flask does not need to be dry but it must be clean and rinsed with deionized water. Record the volume of vinegar in your notebook.

3. To the Erlenmeyer flask containing the vinegar, add approximately 25 ml of

deionized water and swirl gently to mix. Use hydro-ion paper to measure the pH at this point, and record in Table A.

NOTE: Take care to avoid spattering of the vinegar onto the sides of the flask when you are swirling the solution. The addition of water will not change your final results. This is simply to increase the accuracy of the titration.

4. To the same flask, carefully add 3-4 drops of the phenolphthalein indicator using

the dropper provided. NOTE: Do not add more than 4 drops of indicator. This is the critical step in the titration. If you fail to add indicator, then you will never see a color change in your solution.

5. Set up a titration apparatus as shown to the right and place a piece of white paper

underneath the flask. NOTE: This white paper will make the color change from clear to pink easier to see.

6. Using a 150 ml beaker, obtain approximately 75ml of 0.200 M NaOH solution. The exact concentration of the

NaOH will be labeled in the container. Record the molarity of the NaOH solution in Table A. CAUTION: If you spill NaOH solution at any step in this experiment, call your instructor immediately.

7. Making sure that the stopcock of the buret is in the closed position, rinse your buret with 10 ml of the NaOH

solution. NOTE: Your instructor will demonstrate this technique.

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Table B. Titration 8. Using a funnel, slowly fill the buret to slightly above the 0.00 ml mark with the NaOH solution. Carefully open

the stopcock and slowly drain the NaOH solution into a waste beaker until the meniscus of the solution is below the 0.00 ml mark. NOTE: If you have air bubbles in the tip of the buret after this step, call your instructor.

9. Record in your notebook the initial level of the NaOH solution in the buret to the hundredth of a mL.

10. Begin adding the NaOH solution to the vinegar solution prepared in Steps 2-4. To do this, add the NaOH in small amounts (about 1 ml/min) while continuously swirling the flask in a gentle, circular motion. Make sure that the tip of the buret extends slightly into the mouth of the flask, and take care to avoid spattering of the solutions onto the sides of the flask.

Record the final level of NaOH in the buret after the endpoint has been reached. NOTE: While adding the NaOH solution, you will see the formation of temporary pink "clouds" in the vinegar solution. As this pink coloration begins to take longer and longer to disappear, slow the rate of addition of the NaOH solution to about 10 drops/minute. The endpoint of the titration is reached when one drop of the NaOH solution changes the vinegar solution from colorless to a pale, pale pink that permanently remains even after swirling the solution.

11. In your notebook, record the volume of NaOH solution delivered from the buret during the titration. This volume is the difference between the final buret reading (at the endpoint) and the initial buret reading.

12. Use a hydro-ion paper to measure the pH at this point. 13. Repeat steps 1-12 until you have three titrations that agree with 0.50 ml. Clean up Wash and dry all your glassware equipment. Dispose of all used chemicals in the proper waste container. Wipe down your station and place your glassware and equipment back in its proper place. Wash your hands thoroughly with soap and water before leaving your work area. Table C. Calculations of the Molarity of the Acid Show your calculations for each Row of Table C in the space provided in the center column. 14. Calculate average volume of NaOH used in the titration for the 3 trials and record the volume in mL in the top

row of Table C. Show your calculations 15. Convert the average volume from mL to L and record in Table C. 16. Use the concentration of the NaOH to calculate the average number of moles of NaOH used in the titrations and

record in Table C. 17. The number of moles of acid is equal to the number of moles of base at the endpoint of a titration. Your answer

for the number of moles of acid will be the same as the number of moles of base used. 18. To calculate the molarity (moles/liter) of the vinegar, divide the moles of acid by the volume (in liters) of vinegar

that was titrated. Note. Pay attention to significant figures in writing your data and completing your calculations.

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Experiment 14: Data Sheets Last Name________________________First_________________________ Day ___________ ____ / _____ pts

Lab Partner(s): ________________________________________________________________

Note in the calculation section, show setup using dimensional analysis with proper units cancelling out.

Table A: Setup Information

Brand of Vinegar

Weight/Volume % (From Label)

Volume of Vinegar

Molarity of NaOH

pH of Vinegar Round off to hundredth of pH

Table B: Titration Data

Titration #1 Titration #2 Titration #3

9. Initial Buret Reading

10. Endpoint Buret Reading

11. Volume of NaOH

12. pH of Vinegar after neutralization

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Table C: Calculations of the Molarity of Vinegar Calculations Answers

14. Average Volume of NaOH (ml)

15. Average Volume of NaOH (L)

16. Moles of NaOH

17. Moles of Acid

18. Molarity of Acid

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Table D: Calculations for Weight/Volume Percent of Vinegar & Percent Error Calculations Answers

Moles of Acid (copy from Cell C 17)

Molar Mass of Acid

Grams of Acid Titrated

Volume of acid titrated (mL)

Weight/Volume %

Percent Error

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Postlab Questions 1. What other indicator can be used beside phenolphthalein can be used for this experiment? 2. Other than vinegar, name a household product that is acidic and can be used in this experiment. Explain why you have selected this product. 3. What is average [H3O+] of the vinegar solution at the neutralization point in the titration process? 4. Can a solution ever have a pH = 0 ? What does a pH =0 suggest about the concentration of the acid? Can the pH ever be negative? What is the significance of a negative pH for a solution?

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5. Complete the following table. Assume that the solution is 25°C and Kw = 1.00•10-14 Use correct number of significant figures in your answers. See pH, pOH map in the appendix for guidance.

[H3O+] [OH-] pH Acidic Basic or Neutral i.e. 1.00 • 10-10 1.00 • 10-4 10.000 Basic

A

Neutral

B

12.450

C

1.55•10-6

D

1.66 •10-10

Show calculations below.

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A. Appendix Common Conversions

Solubility Rules

Length SI Unit Meter (m) Volume SI Unit Cubic Meter (m3) Mass SI Unit Kilogram (kg)

1 m = 100 cm 1 L = 1000 mL 1 kg = 1000 g

1 m = 1000 mm 1 mL = 1 cm3 1 kg = 2.20 lb

1 cm = 10 mm 1 L = 1.06 qt 1 lb = 454 g

1 km = 0.621 mi 1 qt = 946 mL 1 mole = 6.02 x 1023 particles

1 in. = 2.54 cm (exact) 4 qt = 1 gal Density of water = 1.00 g/mL

16 oz. = 1 lb

Energy SI Unit Joule (J) R = 0.0821 �;��]��

= 62.4 �;��]��

1 calorie = 4.184 J 1 Mole gas = 22.4 L at STP

Specific heat water = 4.184 ��℃

= 1.00 `��℃

Soluble

Exceptions:

All compounds of Li+, Na+, K+, Rb+, Cs+, and NH4+ None

All compounds of NO3− and C2H3O2− None

Compounds of Cl−, Br−, I− Ag+, Hg22+, Pb2+

Compounds of SO42 Sr2+, Pb2+, Ca2+, Ba2+

These compounds generally do not dissolve in water (are insoluble):

Compounds of OH-, S2-, CO32− and PO43−

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Polyatomic Ions

117

VSEPR Table

118

Stoichiometry Map

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Gas law equations:

Ideal Gas Law

STP:

P = 1.0 atm, T = 0°C

1mole = 22.4 L

PV = nRT

Dalton's Law of Partial Pressure

PT = Pa + Pb + Pc + ...

Graham's Law of effusion

pH, pOH, H3O+, OH-, Map

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Periodic Table

chemistry 100 lab - san diego miramar...

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